In 1871 the Russian chemist Dmitri Mendeleev published the properties of a metal no one had ever seen. Four years later the metal was discovered — and it matched his numbers. He read it all off a chart. Reading that chart starts with its rows and columns.
You've already seen the periodic table as the chart that holds every element. Now you read its layout, starting with the rows.
The idea
The periodic table is built from horizontal rows.
Each horizontal row is called a 'period'.
The periods are numbered 1 to 7, from the top row down.
An element's period is simply the row its cell sits in.
Each row is a period, numbered from the top. Sodium sits in Period 3.
Sodium's cell sits in the third row from the top, so sodium is in Period 3.
Carbon's cell sits in the second row, so carbon is in Period 2.
Two extra rows are printed below the main table only to save width — they belong to Periods 6 and 7.
Worked examples
Worked example 1. The periodic table is shown with calcium's cell marked. Which period is calcium in?
Step 1
Find the row that holds the marked cell.
Step 2
Counting from the top, the marked cell sits in row 4.
Step 3
Calcium is in Period 4.
Worked example 2. The periodic table is shown with iodine's cell marked. Which period is iodine in?
Step 1
Find the row that holds the marked cell.
Step 2
Counting from the top, the marked cell sits in row 5.
Step 3
Iodine is in Period 5.
You can now identify the period of an element as the horizontal row it occupies in the periodic table.
Check your understanding
The periodic table outline is shown with phosphorus's cell marked. Which period is phosphorus in?
Answer: 3
Find the row that holds the marked cell. Counting from the top, phosphorus's cell sits in row 3. Phosphorus is in Period 3.
Check your understanding
The periodic table outline is shown with lithium's cell marked. Which period is lithium in?
Answer: 2
Find the row that holds the marked cell. Counting from the top, lithium's cell sits in row 2. Lithium is in Period 2.
Check your understanding
The periodic table outline is shown with barium's cell marked. Which period is barium in?
Answer: 6
Find the row that holds the marked cell. Counting from the top, barium's cell sits in row 6. Barium is in Period 6.
Lesson 2 of 50 · PTB-002
Groups
Learning Journey
Read · Check · Unlock
Did You Know?
You've already read the table's rows — the numbered periods. You've also already used the table's columns once, when you read valence-electron counts from them in the atoms unit. The columns have their own name and numbering.
The idea
The periodic table is also built from vertical columns.
Each vertical column is called a 'group'.
The groups are numbered 1 to 18, from the left edge to the right edge.
Count every column, including the ten columns of the table's middle block.
Each column is a group, numbered 1 to 18 from the left. Chlorine sits in Group 17.
An element's group is the numbered column its cell sits in.
Chlorine's cell sits in the seventeenth column, so chlorine is in Group 17.
Potassium's cell sits in the first column, so potassium is in Group 1.
Worked examples
Worked example 1. The periodic table is shown with nitrogen's cell marked. Which group is nitrogen in?
Step 1
Find the column that holds the marked cell.
Step 2
Counting from the left and including the middle block, the marked cell sits in column 15.
Step 3
Nitrogen is in Group 15.
Worked example 2. The periodic table is shown with iron's cell marked. Which group is iron in?
Step 1
Find the column that holds the marked cell.
Step 2
The middle-block columns have numbers of their own — iron's column is the eighth from the left.
Step 3
Iron is in Group 8.
You can now identify the group of an element as the numbered vertical column it occupies in the periodic table, using the 1 to 18 numbering.
Check your understanding
The periodic table outline is shown with neon's cell marked. Which group is neon in?
Answer: 18
Find the column that holds the marked cell. Counting from the left and including the middle block, neon's cell sits in the last column — column 18. Neon is in Group 18.
Check your understanding
The periodic table outline is shown with manganese's cell marked. Which group is manganese in?
Answer: 7
Find the column that holds the marked cell. Manganese sits in the middle block, and those columns have numbers of their own — its column is the seventh. Manganese is in Group 7.
Check your understanding
The periodic table outline is shown with bromine's cell marked. Which group is bromine in?
Answer: 17
Find the column that holds the marked cell. Counting from the left and including the middle block, bromine's cell sits in column 17. Bromine is in Group 17.
Lesson 3 of 50 · PTB-003
Mendeleev's table
Learning Journey
Read · Check · Unlock
Did You Know?
Wonder this:
By 1869, chemists had discovered 63 elements — a drawer full of facts with no order to them. Dmitri Mendeleev wrote each element on its own card and dealt them out on a table, looking for a pattern.
Each card carried the element's atomic mass and its known properties.
The idea
Mendeleev laid the cards out in order of increasing atomic mass.
Whenever properties began to repeat, he started a new row.
Starting new rows made elements with similar chemical and physical properties line up in the same column.
Lithium, sodium, and potassium landed in one column, because all three react in similar ways.
Mass order, wrapped into rows
Li 6.9
Be 9.0
B 10.8
C 12.0
N 14.0
O 16.0
F 19.0
Na 23.0
Mg 24.3
Al 27.0
Si 28.1
P 31.0
S 32.1
Cl 35.5
One stretch of the mass-ordered list, wrapped into rows. Lithium stacks above sodium — similar elements line up in columns.
The result was the first periodic table that chemists widely accepted.
Worked examples
Worked example 1. Mendeleev placed chlorine, bromine, and iodine in a single column. What put them there?
Step 1
His rows were cut so that similar elements stacked up in columns.
Step 2
Chlorine, bromine, and iodine behave in similar ways, so the mass-ordered rows stacked them in one column.
You can now state that Mendeleev built the first widely accepted periodic table by listing the elements found by then in order of increasing atomic mass and starting a new row so that elements with similar chemical and physical properties lined up in the same column.
Check your understanding
In what order did Mendeleev list the elements when he built his table?
AIn order of increasing atomic mass.correct
BIn order of increasing atomic number.
This option is wrong — you used the modern ordering — atomic numbers were not measured until decades after Mendeleev's table.
CIn alphabetical order of their names.
This option is wrong — you ordered by name — names carry no chemical information, and his columns came from a mass-ordered list.
DIn the order they were discovered.
This option is wrong — you ordered by discovery date — the pattern he found lives in atomic mass, not in history.
Mendeleev laid the element cards out in order of increasing atomic mass. Starting new rows in that mass-ordered list is what made similar elements line up in columns.
Check your understanding
Why did Mendeleev start new rows instead of leaving the elements in one long line?
ASo that elements with similar properties lined up in the same column.correct
BSo that the chart would fit on a single page.
This option is wrong — you treated the rows as a layout convenience — the row breaks were placed where properties began to repeat.
CSo that elements with similar masses stayed close together.
This option is wrong — you kept mass as the reason for the rows — mass set the order along the line, but properties set where each row ended.
DSo that each row held exactly the same number of elements.
This option is wrong — you made row length the goal — his rows varied in length because properties, not counting, decided the breaks.
The order along the line came from atomic mass. The row breaks came from properties: whenever behavior began to repeat, a new row started. That is what stacked similar elements into shared columns.
Check your understanding
Beryllium and magnesium react in similar ways. Where did Mendeleev's method put them?
AIn the same column.correct
BIn the same row.
This option is wrong — you swapped rows and columns — rows follow increasing mass, and it is columns that collect similar elements.
CFar apart, because their atomic masses are different.
This option is wrong — you let the mass gap separate them — the new-row move is exactly what brings similar elements back together, one above the other.
DSide by side in one row.
This option is wrong — you placed similar elements as row neighbors — row neighbors have consecutive masses, not similar behavior.
Mendeleev's columns collect elements that behave alike. Beryllium and magnesium react in similar ways, so the mass-ordered rows stack them in one column.
Lesson 4 of 50 · PTB-004
Gaps that made predictions
Learning Journey
Read · Check · Unlock
Did You Know?
Wonder this:
A chart with holes in it looks unfinished. Mendeleev said the holes were the best part.
You've already seen how Mendeleev's mass-ordered rows stacked similar elements into columns. Sometimes the next card in the mass order did not fit the column it landed in.
The idea
When the next known element did not match a column's properties, Mendeleev left the cell empty instead of forcing the element in.
He read each empty cell as an element that existed but had not been discovered yet.
For each gap he predicted the missing element's properties from the elements around the gap.
Below aluminum he predicted a soft metal with an atomic mass near 68 and a density near 5.9 g/cm³.
The gap below aluminum. Mendeleev filled it with a prediction, not an element.
In 1875 that metal was discovered — gallium, with an atomic mass of 69.7 and a density of 5.91 g/cm³.
Discoveries that matched his predictions became strong evidence that the table's pattern was real.
Worked examples
Worked example 1. Mendeleev predicted an undiscovered element below silicon, with an atomic mass near 72. In 1886 germanium was discovered, with an atomic mass of 72.6. Where had the predicted numbers come from?
Step 1
No sample of the element existed, so nothing could be measured.
Step 2
The gap's neighbors — silicon above it and the elements beside it — set the pattern the missing element had to continue.
Step 3
The predictions came from the properties of the elements around the gap, and germanium matched them.
You can now explain that Mendeleev left gaps in his table for undiscovered elements and predicted the properties of the missing elements from the elements around each gap, and that later discoveries matching those predictions became strong evidence for his table.
Check your understanding
Why did Mendeleev leave gaps in his table?
AHe kept cells empty as deliberate placeholders for elements not yet discovered.correct
BThe gaps marked places where his mass measurements had failed.
This option is wrong — you read the gaps as errors — they were deliberate placeholders for missing elements.
CHe left gaps so every row would come out the same length.
This option is wrong — you made the gaps a layout device — each gap sat exactly where the next known element failed to match the column.
DHe removed elements that did not fit his pattern.
This option is wrong — you had him deleting data — no element was removed; the empty cells stood for elements no one had found yet.
Forcing the next element into a column it did not match would have broken the pattern. So Mendeleev left the cell empty and read it as an undiscovered element. The known elements stayed in columns that matched their properties.
Check your understanding
Where did the predicted properties for a gap element come from?
AFrom the properties of the elements around the gap.correct
BFrom measurements on a small sample of the new element.
This option is wrong — you assumed a sample existed — the whole point of a gap is that no one had the element to measure.
CFrom the average of all known elements' properties.
This option is wrong — you averaged the whole table — only the gap's neighbors set the pattern the missing element had to continue.
DFrom random guesses that were corrected after the discovery.
This option is wrong — you made the predictions guesswork — they were read off the surrounding pattern before any discovery, which is why the matches counted as evidence.
No sample existed, so nothing could be measured. The elements around the gap set the pattern, and the missing element had to continue it. That is why a matching discovery counted as evidence for the table.
Check your understanding
In 1879 scandium was discovered, and its properties matched what Mendeleev had predicted for one of his gaps. Why did this discovery strengthen chemists' trust in his table?
AIt showed the table's pattern was real, not an accident of the elements already known.correct
BIt proved that every atomic mass in the table had been measured correctly.
This option is wrong — you turned the evidence into a mass check — the prediction was about properties read from the pattern, not about re-measuring masses.
CIt completed the table, leaving nothing else to discover.
This option is wrong — you treated one filled gap as completion — other gaps remained, and each matching discovery added evidence of its own.
DIt showed that new elements can only be found by using tables.
This option is wrong — you overreached — the discovery supported the table's pattern; it did not make tables the only route to new elements.
The prediction was written before anyone had seen the element. A real element then matched it. A pattern that predicts the unknown is strong evidence of being real.
Lesson 5 of 50 · PTB-005
Moseley's atomic-number ordering
Learning Journey
Read · Check · Unlock
Did You Know?
Mendeleev's mass order worked almost everywhere — almost. A few element pairs refused to cooperate.
The idea
Tellurium's atomic mass is 127.6 and iodine's is 126.9, so mass order puts iodine before tellurium.
But iodine behaves like the elements in the column after tellurium's, so mass order drops both into columns that do not match their properties.
One stubborn pair
Element
Atomic mass (amu)
Atomic number
tellurium
127.6
52
iodine
126.9
53
Mass order puts iodine first. Atomic-number order puts tellurium first — and matches both elements' behavior.
Mendeleev could only swap such pairs by hand and hope a reason would turn up.
In 1913 Henry Moseley measured each element's atomic number — its proton count — using X-rays.
Tellurium's atomic number is 52 and iodine's is 53, so atomic-number order puts tellurium first.
Ordered by atomic number, every element lands in the column that matches its properties, with no hand swaps.
The modern periodic table keeps Moseley's order: elements run by increasing atomic number.
Worked examples
Worked example 1. Argon's atomic mass is 39.95 and potassium's is 39.10. By mass, potassium would come first — landing it among gases that almost never react, with argon among metals that react violently with water. How does atomic-number ordering fix the pair?
Step 1
Argon's atomic number is 18 and potassium's is 19.
Step 2
Atomic-number order puts argon first and potassium second.
Step 3
Each element then sits in the column that matches its behavior.
Step 4
Ordering by atomic number places argon before potassium, and both land in matching columns.
You can now explain that Moseley reordered the periodic table by atomic number instead of atomic mass, which fixed the element pairs that mass ordering had placed out of line with their columns' properties.
Check your understanding
Cobalt's atomic mass is 58.93 and nickel's is 58.69, yet cobalt comes before nickel in the modern periodic table. Why?
ACobalt has the smaller atomic number, and the modern table orders elements by atomic number.correct
BCobalt was discovered before nickel, and earlier discoveries come first.
This option is wrong — you ordered by history — discovery dates play no part in the table's order.
CThe pair was swapped by hand to keep similar elements lined up.
This option is wrong — you reached for Mendeleev's old patch — hand swaps were the problem; atomic-number ordering removed the need for them.
DAtomic masses that close count as a tie, so either order is allowed.
This option is wrong — you invented a tie rule — the table's order is fixed by atomic number, however close the masses are.
The modern table orders elements by atomic number. Cobalt's atomic number is 27 and nickel's is 28. So cobalt comes first, even though its atomic mass is larger.
Check your understanding
What did Moseley measure, and what did the table do with it?
AHe measured each element's atomic number, and the table was reordered by it.correct
BHe measured each element's atomic mass more precisely, and the mass order was corrected.
This option is wrong — you kept mass as the ordering — better masses would not have fixed the stubborn pairs; a different quantity did.
CHe measured each element's density, and the table was reordered by it.
This option is wrong — you swapped in a bulk property — density never ordered the table.
DHe measured each element's number of neutrons, and the table was reordered by it.
This option is wrong — you counted the wrong nuclear particle — the atomic number counts protons.
Moseley's X-ray work measured each element's atomic number — its proton count. Reordering the table by atomic number put every element in a column matching its properties.
Check your understanding
After the table was reordered by atomic number, what happened to the element pairs that mass order had placed badly?
AEach landed in the column matching its properties, with no hand swaps needed.correct
BThey still had to be swapped by hand, but the swaps now had a name.
This option is wrong — you kept the hand swaps — atomic-number order needs none; that is exactly why it won.
CThey were removed from the table until their masses could be re-measured.
This option is wrong — you made it a mass problem — the masses were fine; the ordering quantity was wrong.
DThey were placed in a separate block below the table.
This option is wrong — you confused the misfit pairs with the detached bottom rows — those rows are a printing convenience, not a fix for bad placements.
Mass order dropped a few pairs into columns that did not match their behavior. Ordered by atomic number, every element lands in a matching column. The hand swaps disappeared — which is why chemists accepted the new order.
Lesson 6 of 50 · PTB-006
The periodic law
Learning Journey
Read · Check · Unlock
Did You Know?
Forget the table's shape for a moment and stretch the elements out into one long line, in order of increasing atomic number.
The idea
Walk along the line and similar elements keep coming back.
Lithium is element 3, sodium is element 11, and potassium is element 19 — and all three react in similar ways.
Elements with similar properties appear at regular intervals along the line.
One long line, and the same behavior keeps coming back: lithium at 3, sodium at 11, potassium at 19.
In the early stretch of the line, the repeat comes every 8 elements; farther along, the intervals grow longer, but the repeating never stops.
This repeating pattern is called the 'periodic law': listed in order of increasing atomic number, elements with similar properties appear at regular intervals.
The table's rows are just this one line, cut wherever the pattern starts over.
Worked examples
Worked example 1. Helium (element 2), neon (element 10), and argon (element 18) are all gases that almost never react. How does the periodic law describe their positions?
Step 1
Similar properties appearing at regular intervals along the atomic-number line.
Step 2
The three gases that almost never react sit at regular intervals — 2, 10, 18 — exactly the periodic law's pattern.
You can now state that when the elements are listed in order of increasing atomic number, elements with similar properties appear at regular intervals, a repeating pattern called the periodic law.
Check your understanding
When the elements are listed in order of increasing atomic number, what does the periodic law say happens?
AElements with similar properties appear at regular intervals.correct
BElements with similar properties always sit next to each other in the list.
This option is wrong — you put the similar elements side by side — the pattern is a repeat at intervals, not a clustering of neighbors.
CProperties change steadily from the first element to the last, with no repeats.
This option is wrong — you expected one long trend — the striking fact is that behavior keeps coming back.
DEvery element behaves differently from every other, in no pattern at all.
This option is wrong — you denied the pattern — similar behavior returns at regular intervals, which is the law's whole content.
List the elements by increasing atomic number. Similar properties keep reappearing at regular intervals along the list. That repeating pattern is the periodic law.
Check your understanding
Beryllium is element 4, magnesium is element 12, and calcium is element 20 — and the three behave alike. Which pattern is this an example of?
ASimilar properties appearing at regular intervals in the atomic-number list — the periodic law.correct
BA coincidence among three unrelated elements.
This option is wrong — you dismissed the spacing — 4, 12, 20 is a regular interval, exactly the law's pattern.
CSimilar properties belonging to elements with nearly equal atomic masses.
This option is wrong — you moved the similarity onto mass — the three have very different masses; it is their positions in the list that repeat.
DProperties improving as atomic number grows.
This option is wrong — you read the list as one long trend — the pattern is a repeat, not a steady climb.
The three sit at 4, 12, and 20 — regular intervals of 8. Similar behavior returning at regular intervals is the periodic law.
Check your understanding
Which observation is an example of the periodic law?
AFluorine's kind of behavior reappearing at chlorine, a regular interval later in the atomic-number list.correct
BFluorine's kind of behavior fading gradually across the elements after it in the list.
This option is wrong — you described a fade-out — the law is about behavior returning, not weakening.
CEvery ninth element in the whole list being a gas.
This option is wrong — you invented a fixed universal count — the intervals are regular but grow longer farther along the list.
DNeighboring elements in the list always behaving alike.
This option is wrong — you moved the similarity to next-door neighbors — neighbors usually differ; it is elements an interval apart that match.
The periodic law says similar properties reappear at regular intervals along the atomic-number list. Fluorine and chlorine are such a pair — the same kind of behavior, one interval apart.
Lesson 7 of 50 · PTB-007
How the table developed
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen each chapter of the table's story on its own. This lesson walks the whole story once, naming the evidence that drove each change.
The idea
In 1869 Mendeleev listed the elements by increasing atomic mass and started new rows so similar elements shared columns.
He left gaps for undiscovered elements and predicted their properties from the elements around each gap.
The table's story: built on mass and properties, tested by predictions, fixed by atomic number.
Gallium's discovery in 1875, matching his predictions, became strong evidence that the table's pattern was real.
But mass order held a flaw: pairs like tellurium and iodine landed in columns that did not match their behavior, and only hand swaps could fix them.
In 1913 Moseley measured atomic numbers, and ordering by atomic number put every element in a matching column with no swaps.
Each change in the table was driven by evidence: confirmed predictions supported the pattern, and the stubborn pairs forced the change of ordering.
Worked examples
Worked example 1. Germanium was discovered in 1886, matching Mendeleev's prediction for the gap below silicon. Which part of the table's design did this discovery support?
Step 1
The prediction had been read from the elements around the gap.
Step 2
A real element matching it showed the property-grouped columns carried real information.
Step 3
It supported the property-grouped columns and the gap method built on them.
Worked example 2. Argon's atomic mass (39.95) is larger than potassium's (39.10), yet argon must come first to match both elements' behavior. Which change in the table did cases like this force?
Step 1
Mass order puts potassium first — into the wrong column.
Step 2
Atomic-number order (argon 18, potassium 19) puts both elements in matching columns.
Step 3
They forced the switch from mass ordering to atomic-number ordering.
You can now explain how the periodic table changed over time, from Mendeleev's mass-ordered table with property-grouped columns and predictive gaps to Moseley's atomic-number ordering, naming the evidence that drove each change.
Check your understanding
Which change did Moseley's measurements make to the periodic table?
AThe elements were reordered by atomic number instead of atomic mass.correct
BThe elements were reordered by the number of neutrons.
This option is wrong — you counted the wrong nuclear particle — the atomic number counts protons.
CThe gaps were removed because the predictions had failed.
This option is wrong — you crossed two chapters of the story — the predictions succeeded, and Moseley's fix was about ordering, not gaps.
DThe rows were removed so all elements sat in one line.
This option is wrong — you undid the table's shape — the rows stayed; only the ordering quantity changed.
Moseley measured each element's atomic number. The table was reordered by atomic number, and every element landed in a column matching its behavior.
Check your understanding
Scandium was discovered in 1879, matching a Mendeleev gap prediction. What role did discoveries like this play in the table's story?
AThey were evidence that the table's pattern was real.correct
BThey forced the switch to atomic-number ordering.
This option is wrong — you matched the evidence to the wrong change — the stubborn mass-order pairs forced the reordering; confirmed predictions supported the pattern.
CThey showed the table needed no further changes.
This option is wrong — you ended the story early — the ordering flaw remained until 1913.
DThey filled the last empty cells of the table.
This option is wrong — you closed the table — other gaps remained open for decades.
Gap predictions were written before the elements were found. Each matching discovery — scandium among them — was evidence the pattern was real. The reordering, a separate change, was forced by the misfit pairs.
Check your understanding
Which sequence puts the periodic table's development in order?
AMass-ordered table with gaps — predicted elements discovered — reordering by atomic number.correct
BReordering by atomic number — mass-ordered table with gaps — predicted elements discovered.
This option is wrong — you started with the ending — atomic numbers could not be measured until 1913.
CMass-ordered table with gaps — reordering by atomic number — predicted elements discovered.
This option is wrong — you moved the reordering before the confirmations — gallium arrived in 1875, decades before Moseley's work.
DPredicted elements discovered — mass-ordered table with gaps — reordering by atomic number.
This option is wrong — you put the confirmations before the table — the predictions only existed because the gapped table came first.
First the 1869 table: mass order, property columns, gaps. Then the confirmations, starting in 1875. Last, in 1913, the reordering by atomic number.
Hit a copper wire with a hammer and it flattens into a ribbon. Hit a lump of coal the same way and it shatters. Two solids, two completely different answers to the same blow.
Copper's kind of answer marks a whole class of elements — the metals. Their properties come as a package.
The idea
Metals are shiny.
Metals conduct heat well and conduct electricity well.
The metal package
Property
Seen in copper
shiny
polished copper gleams
conducts heat and electricity
copper wiring and copper pans
malleable — hammers into sheets
copper sheeting
ductile — draws into wires
copper wire
solid at room temperature
all metals except liquid mercury
Five properties, one package. Copper shows them all.
Metals can be hammered into sheets — this property is called 'malleable'.
Metals can be drawn into wires — this property is called 'ductile'.
Metals are solids at room temperature, except mercury, which is a liquid.
Copper shows the whole package: it is shiny, it conducts, and it bends into sheets and wires without shattering.
Worked examples
Worked example 1. Gold can be hammered into sheets thin enough to let light through. Which metal property is that?
Step 1
Hammered into sheets is the malleable behavior.
Step 2
Malleability — metals can be hammered into sheets.
Worked example 2. Overhead power lines are made of aluminum. Which metal property is the reason?
Step 1
Power lines exist to carry current.
Step 2
Metals conduct electricity well.
You can now state the characteristic properties of metals: metals are shiny, conduct heat and electricity well, can be hammered into sheets (malleable) and drawn into wires (ductile), and are solids at room temperature except mercury.
Check your understanding
Which behavior would you expect from a typical metal?
AIt conducts electricity well.correct
BIt shatters when hammered.
This option is wrong — you picked the opposite of the metal behavior — metals flatten under a hammer rather than shattering.
CIt blocks heat from passing through.
This option is wrong — you reversed the heat behavior — metals conduct heat well, which is why metal pans heat food.
DIt is a gas at room temperature.
This option is wrong — you missed the state rule — metals are solids at room temperature, with liquid mercury the one exception.
The metal package: shiny, conducts heat and electricity, malleable, ductile, solid at room temperature except mercury. Conducting electricity well is part of the package.
Check your understanding
A jeweler draws platinum into a fine wire without it breaking. Which property name fits that behavior?
ADuctile.correct
BMalleable.
This option is wrong — you mixed the two shaping words — malleable is hammered into sheets; drawn into wires is ductile.
CConducts electricity well.
This option is wrong — you named a different package property — conducting electricity is about current passing through, not about changing shape.
DShiny.
This option is wrong — you named a surface property — shine says nothing about how a solid changes shape.
Drawn into wires is the ductile behavior. Hammered into sheets is the malleable behavior — the two words split the two shaping jobs.
Check your understanding
A metalworker hammers an iron bar flat into a blade. Which property name fits that behavior?
AMalleable.correct
BDuctile.
This option is wrong — you mixed the two shaping words — ductile is drawn into wires; hammered flat is malleable.
CConducts heat well.
This option is wrong — you named a different package property — conduction is about heat passing through, not about changing shape.
DSolid at room temperature.
This option is wrong — you named the state rule — true of iron, but the question asks which property name fits how the metal changes shape.
Hammered into sheets or flattened under blows is the malleable behavior. Iron flattening into a blade instead of shattering is malleability at work.
Lesson 9 of 50 · PTB-009
What nonmetals are like
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen the metal package: shiny, conducting, bendable, solid. A second class of elements answers every one of those points the opposite way.
The idea
These elements are the nonmetals.
Nonmetals are dull rather than shiny.
Nonmetals conduct heat poorly and conduct electricity poorly.
Two packages, point for point
Property
Metals
Nonmetals
surface
shiny
dull
conducts heat and electricity
well
poorly
when hammered (solid)
flattens — malleable
shatters — brittle
state at room temperature
solid (except mercury)
many are gases
The nonmetal package answers the metal package point for point.
Solid nonmetals shatter rather than bend — this property is called 'brittle'.
Many nonmetals are not solid at all: at room temperature, many are gases — most of the air around you is two of them.
Sulfur shows the solid version of the package: a dull yellow solid that shatters when struck and does not conduct electricity.
Worked examples
Worked example 1. Solid iodine crumbles into flakes when pressed with a spoon. Which nonmetal property is that?
Step 1
Shattering or crumbling instead of bending is the brittle behavior.
Step 2
Brittleness — solid nonmetals are brittle.
Worked example 2. Chlorine is a nonmetal you will never pick up as a lump at room temperature. Why not?
Step 1
Many nonmetals are gases at room temperature, and chlorine is one of them.
Step 2
Chlorine is a gas at room temperature.
You can now state the characteristic properties of nonmetals: nonmetals are dull rather than shiny, poor conductors of heat and electricity, brittle (they shatter rather than bend) when solid, and many are gases at room temperature.
Check your understanding
Which behavior would you expect from a solid nonmetal?
AIt shatters when struck.correct
BIt flattens into a sheet when struck.
This option is wrong — you gave the metal answer — solid nonmetals are brittle and shatter.
CIt conducts electricity well.
This option is wrong — you gave the metal answer — nonmetals conduct poorly.
DIt can be drawn into wire.
This option is wrong — you gave the metal answer — a brittle solid snaps instead of drawing out.
The nonmetal package: dull, poor conductor, brittle when solid, and many are gases at room temperature. A solid nonmetal struck with a hammer shatters.
Check your understanding
Red phosphorus is a dull solid that does not conduct electricity. When a lump of it is crushed, it breaks into powder. Which property name fits that breaking behavior?
ABrittle.correct
BMalleable.
This option is wrong — you named the metal shaping behavior — a malleable solid flattens; this one broke to powder.
CDuctile.
This option is wrong — you named the wire-drawing behavior — a solid that crumbles cannot be drawn out.
DDull.
This option is wrong — you named the surface property from the setup — the question asks about how the solid breaks.
Shattering or crumbling instead of bending is the brittle behavior. Solid nonmetals such as phosphorus are brittle.
Check your understanding
Nitrogen makes up most of the air you breathe. Which part of the nonmetal package does nitrogen show at room temperature?
AMany nonmetals are gases at room temperature.correct
BNonmetals are brittle.
This option is wrong — you applied the solid rule to a gas — brittleness only describes solid nonmetals.
CNonmetals are dull.
This option is wrong — you reached for the surface rule — an invisible gas shows the state rule, not the surface rule.
DNonmetals are solids except one.
This option is wrong — you borrowed the metal state rule — it is metals that are all solid but one; many nonmetals are gases.
Metals are solids at room temperature, except mercury. Nonmetals break that pattern: many, including nitrogen, are gases at room temperature.
Lesson 10 of 50 · PTB-010
What metalloids are like
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen two opposite packages — metal and nonmetal. A handful of elements refuse to pick a side.
The idea
These elements are the metalloids.
A metalloid shows a mix: some of its properties are metal-like, others are nonmetal-like.
Silicon looks shiny like a metal, but it is brittle like a solid nonmetal.
Three packages side by side
Property
Metals
Metalloids
Nonmetals
surface
shiny
often shiny
dull
when hammered (solid)
flattens
shatters
shatters
conducts electricity
well
in between — semiconductor
poorly
Metalloids take their surface from one side and their breaking behavior from the other, with conduction in between.
Silicon also conducts electricity in between: better than a nonmetal, but far worse than a metal.
An in-between conductor like this is called a 'semiconductor', and several metalloids are semiconductors.
That in-between conduction is exactly what computer chips need, which is why chips are built on silicon.
Worked examples
Worked example 1. Germanium conducts electricity far worse than copper but far better than sulfur. What does that in-between conduction make it?
Step 1
Conduction between a metal's and a nonmetal's is the semiconductor behavior.
Step 2
A semiconductor — the mark of several metalloids.
You can now state that metalloids have a mix of metal and nonmetal properties, and that several are semiconductors, conducting electricity better than nonmetals but not as well as metals.
Check your understanding
What mix of properties marks a metalloid?
ASome properties like a metal's and others like a nonmetal's.correct
BAll the metal properties, just in weaker versions.
This option is wrong — you made metalloids weak metals — a metalloid's brittleness is a full nonmetal property, not a weak metal one.
CAll the nonmetal properties, just in stronger versions.
This option is wrong — you made metalloids strong nonmetals — the shine of silicon is a genuine metal-like property.
DNo properties in common with either class.
This option is wrong — you pushed metalloids outside both packages — they borrow from both, which is the whole idea of the mix.
A metalloid takes some properties from the metal package and others from the nonmetal package. Silicon is the model case: shiny like a metal, brittle like a nonmetal, conduction in between.
Check your understanding
Antimony is a metalloid with a silvery shine. When a sample is struck sharply, what is the likely result?
AIt shatters.correct
BIt flattens into a sheet.
This option is wrong — you let the shine promise the whole metal package — a metalloid's breaking behavior is the nonmetal kind.
CIt stretches into a wire.
This option is wrong — you gave the ductile metal behavior — a brittle solid snaps instead of stretching.
DNothing happens — metalloids cannot break.
This option is wrong — you made the mix mean indestructible — the mix means metal-like in some properties and nonmetal-like in others.
Shine is antimony's metal-like property. Its breaking behavior comes from the other package: metalloids are brittle and shatter.
Check your understanding
Boron conducts electricity poorly compared with silver, but clearly better than phosphorus does. What is this in-between conduction called?
ABeing a semiconductor.correct
BBeing a full conductor.
This option is wrong — you rounded the conduction up — a semiconductor falls well short of a metal's conduction.
CBeing a non-conductor.
This option is wrong — you rounded the conduction down — boron clearly beats a nonmetal's near-zero conduction.
DBeing an alloy.
This option is wrong — you named a metal mixture — alloy describes what something is made of, not how it conducts.
Conduction between a metal's and a nonmetal's is the semiconductor behavior. Several metalloids, boron among them, are semiconductors.
Lesson 11 of 50 · PTB-011
Classify from properties
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen the three property packages. Now you work backward: given the measured properties of an unknown element, name its class.
The idea
Sample A is shiny, flattens under a hammer, and conducts electricity well — A is a metal.
Sample B is also shiny, but it shatters under the hammer and conducts only weakly — B is a metalloid.
Sample C is dull, shatters, and does not conduct at all — C is a nonmetal.
Notice what changed between A and B: the shine stayed, but the breaking and the conduction switched sides.
So judge by conduction and breaking first: conducts well and bends means metal; conducts weakly and shatters means metalloid; barely conducts, dull or gas, means nonmetal.
Shine can mislead — a metalloid can shine like a metal — and one more shortcut always holds: an element that is a gas at room temperature is a nonmetal.
Judge by conduction and breaking
Conducts electricity
When hammered
Class
well
flattens
metal
weakly
shatters
metalloid
barely or not at all
shatters — or the element is a gas
nonmetal
Conduction and breaking decide. Shine can mislead.
Worked examples
Worked example 1. An element is a gas at room temperature. Classify it.
Step 1
Only nonmetals are gases at room temperature — metals are solid (except liquid mercury), and metalloids are solids too.
Step 2
The element is a nonmetal.
Worked example 2. An element is shiny and conducts electricity — but far worse than copper wiring — and it shatters when struck. Classify it.
Step 1
Judge by conduction and breaking first: weak conduction and shattering point the same way.
Step 2
The shine does not overrule them — metalloids can shine.
Step 3
The element is a metalloid.
You can now classify an element as a metal, nonmetal, or metalloid from supplied property data.
Check your understanding
An element is shiny, can be hammered flat without cracking, and conducts electricity well. Classify it.
AA metal.correct
BA metalloid.
This option is wrong — you held back because shine can mislead — but here conduction and breaking both give the metal answer, and they decide.
CA nonmetal.
This option is wrong — you went against every listed property — dullness, brittleness, and poor conduction are the nonmetal marks, and none is present.
DNone of the above.
This option is wrong — you rejected all three classes — every element lands in one of them, and this data set is the full metal package.
Judge by conduction and breaking first. Conducts well and flattens — that is the metal pattern. The shine agrees, so the element is a metal.
Check your understanding
An element is a dull yellow solid that shatters when struck and does not conduct electricity. Classify it.
AA nonmetal.correct
BA metal.
This option is wrong — you went against every listed property — metals are shiny, flatten, and conduct well.
CA metalloid.
This option is wrong — you gave the in-between class to a clear case — a metalloid conducts weakly but measurably, and this element does not conduct at all.
DNone of the above.
This option is wrong — you rejected all three classes — dull, brittle, and non-conducting is the nonmetal package exactly.
Judge by conduction and breaking first. No conduction and shattering — that is the nonmetal pattern. The dull surface agrees, so the element is a nonmetal.
Check your understanding
A silvery-gray element is brittle and conducts electricity, though far worse than silver wiring does. Classify it.
AA metalloid.correct
BA metal.
This option is wrong — you let the silvery shine decide — conduction and breaking outrank shine, and both point away from metal.
CA nonmetal.
This option is wrong — you rounded the weak conduction down to none — a nonmetal barely conducts at all, and this element measurably does.
DNone of the above.
This option is wrong — you rejected all three classes — shiny but brittle with weak conduction is the metalloid mix itself.
Judge by conduction and breaking first. Weak conduction with shattering is the metalloid pattern. The shine does not overrule them — metalloids can shine.
Lesson 12 of 50 · PTB-012
Where each class lives on the table
Learning Journey
Read · Check · Unlock
Did You Know?
You've already classified elements from their measured properties. The periodic table saves you the measuring: each class lives in its own territory.
The idea
Metals fill the left side and the center of the periodic table.
Nonmetals sit in the upper right corner.
Three territories: metals left and center, nonmetals upper right, metalloids along the stair-step line. Hydrogen is the stray.
One nonmetal is out of place: hydrogen sits at the top left, in metal territory, but hydrogen is a nonmetal.
Between the two territories runs a zigzag boundary called the 'stair-step line'.
Metalloids sit along the stair-step line.
Silicon, for example, lies right on the line.
Worked examples
Worked example 1. Which class of elements takes up the largest share of the periodic table?
Step 1
The metals' territory is the whole left side plus the center.
Step 2
Metals — most of the table is metal territory.
Worked example 2. Where on the table do you look to find the metalloids?
Step 1
The metalloids' home is not a broad region but the boundary itself.
Step 2
Along the stair-step line, between the metals and the nonmetals.
You can now identify the regions of the periodic table where metals (left and center), nonmetals (upper right, plus hydrogen), and metalloids (along the stair-step line between them) are found.
Check your understanding
The periodic table is shown with one region shaded: the left side and center. Which class of elements lives there?
AThe metals.correct
BThe nonmetals.
This option is wrong — you swapped the territories — the nonmetals hold the upper right corner, not the left and center.
CThe metalloids.
This option is wrong — you gave the boundary dwellers a whole region — metalloids sit only along the stair-step line.
DNone of the above.
This option is wrong — you rejected all three classes — the left and center is the metals' territory.
Metals fill the left side and the center of the table. The shaded region is metal territory — the table's largest share.
Check your understanding
The periodic table is shown with one region shaded: the upper right corner. Which class of elements lives there?
AThe nonmetals.correct
BThe metals.
This option is wrong — you swapped the territories — metals hold the left and center, and the upper right is the far side of the stair-step line.
CThe metalloids.
This option is wrong — you widened the boundary into a corner — metalloids sit only along the stair-step line itself.
DNone of the above.
This option is wrong — you rejected all three classes — the upper right corner is nonmetal territory.
Nonmetals sit in the upper right corner of the table. The shaded corner lies past the stair-step line, on the nonmetal side.
Check your understanding
The periodic table is shown with the cells along the zigzag boundary shaded. Which class of elements lives in those cells?
AThe metalloids.correct
BThe metals.
This option is wrong — you extended metal territory onto the boundary — the metals stop where the stair-step line begins.
CThe nonmetals.
This option is wrong — you extended nonmetal territory onto the boundary — the boundary cells belong to the in-between class.
DNone of the above.
This option is wrong — you rejected all three classes — the boundary cells are exactly where the metalloids live.
The zigzag boundary is the stair-step line. Metalloids — the in-between class — sit along it.
Lesson 13 of 50 · PTB-013
Classify from position
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen the map: metals left and center, nonmetals upper right, metalloids along the stair-step line. Now the map does the classifying for you.
The idea
To classify an element, find its cell and ask which territory the cell sits in.
Iron's cell sits in the center block — metal territory — so iron is a metal.
Find the cell, name the territory: iron in metal country, bromine past the line, antimony on it.
Bromine's cell sits near the upper right corner, past the stair-step line, so bromine is a nonmetal.
Antimony's cell sits right on the stair-step line, so antimony is a metalloid.
The position alone answers the question — no hammer, no wires, no measuring.
Worked examples
Worked example 1. The table outline is shown with barium's cell marked, in the second column from the left. Classify barium.
Step 1
The cell sits on the left side — metal territory.
Step 2
Barium is a metal.
Worked example 2. The outline is shown with phosphorus's cell marked, in the upper right, past the stair-step line. Classify phosphorus.
Step 1
The cell sits past the stair-step line, in the upper right corner — nonmetal territory.
Step 2
Phosphorus is a nonmetal.
Worked example 3. The outline is shown with tellurium's cell marked, sitting on the stair-step line. Classify tellurium.
Step 1
A cell on the stair-step line is a metalloid seat.
Step 2
Tellurium is a metalloid.
You can now classify an element as a metal, nonmetal, or metalloid from its position in the periodic table.
Check your understanding
The table outline is shown with cesium's cell marked, in the first column, near the bottom. Classify cesium from its position.
AA metal.correct
BA nonmetal.
This option is wrong — you put nonmetal territory on the left — the left side, top to bottom, is metal country.
CA metalloid.
This option is wrong — you placed the cell on the boundary — the stair-step line runs far to the right of column 1.
DNone of the above.
This option is wrong — you rejected all three classes — every position on the map belongs to one of them, and this one is metal territory.
Find the cell: first column, near the bottom — the left side of the table. The left side and center are metal territory. Cesium is a metal.
Check your understanding
The table outline is shown with selenium's cell marked, in the upper right, past the stair-step line. Classify selenium from its position.
AA nonmetal.correct
BA metal.
This option is wrong — you kept metal territory going past the boundary — the stair-step line ends it, and this cell sits beyond.
CA metalloid.
This option is wrong — you widened the boundary — near the line is not on it; metalloid seats are the line's own cells.
DNone of the above.
This option is wrong — you rejected all three classes — a cell past the line in the upper right is nonmetal territory.
Find the cell: upper right, past the stair-step line. Past the line is nonmetal territory. Selenium is a nonmetal.
Check your understanding
The table outline is shown with boron's cell marked, sitting at the very top of the stair-step line. Classify boron from its position.
AA metalloid.correct
BA nonmetal.
This option is wrong — you read high-and-right as nonmetal automatically — this cell sits ON the line, and the line's cells are metalloid seats.
CA metal.
This option is wrong — you kept metal territory running to the line's top — cells on the line itself belong to the metalloids.
DNone of the above.
This option is wrong — you rejected all three classes — on the stair-step line means metalloid.
Find the cell: it sits on the stair-step line itself. Cells on the line are metalloid seats, top to bottom of the zigzag. Boron is a metalloid.
Lesson 14 of 50 · PTB-014
Main-group elements and transition metals
Learning Journey
Read · Check · Unlock
Did You Know?
You've already numbered the groups 1 to 18. Chemists split those eighteen columns into two named blocks.
The idea
Groups 1, 2, and 13 through 18 are the 'main-group elements'.
Groups 3 through 12 — the table's middle block — are the 'transition metals'.
The group number alone decides which block an element belongs to.
Two blocks: groups 1, 2, and 13-18 are the main-group elements; groups 3-12 are the transition metals.
Calcium sits in Group 2, so calcium is a main-group element.
Iron sits in Group 8, so iron is a transition metal.
Worked examples
Worked example 1. Silver sits in Group 11. Is silver a main-group element or a transition metal?
Step 1
Group 11 falls inside the 3-through-12 middle block.
Step 2
Silver is a transition metal.
Worked example 2. Tin sits in Group 14. Is tin a main-group element or a transition metal?
Step 1
Group 14 falls in the 13-through-18 run, outside the middle block.
Step 2
Tin is a main-group element.
You can now identify an element as a main-group element (Groups 1, 2, and 13 through 18) or a transition metal (Groups 3 through 12) from its position in the periodic table.
Check your understanding
Four cells are marked on the table outline: chromium (Group 6), rubidium (Group 1), iodine (Group 17), and lead (Group 14). Which marked element is a transition metal?
AChromium.correct
BRubidium.
This option is wrong — you put Group 1 in the middle block — Groups 1 and 2 are main-group columns.
CIodine.
This option is wrong — you put Group 17 in the middle block — Groups 13 through 18 are main-group columns.
DLead.
This option is wrong — you put Group 14 in the middle block — the block ends at Group 12.
The transition metals are Groups 3 through 12. Chromium's Group 6 falls inside that block; Groups 1, 14, and 17 fall outside it.
Check your understanding
Which groups hold the transition metals?
AGroups 3 through 12.correct
BGroups 1 and 2.
This option is wrong — you picked the left-edge columns — those are main-group columns.
CGroups 13 through 18.
This option is wrong — you picked the right-hand run — those are main-group columns too.
DGroups 1 through 18.
This option is wrong — you took the whole table — the transition metals are only the middle block.
The middle block — Groups 3 through 12 — holds the transition metals. Groups 1, 2, and 13 through 18 are the main-group elements.
Check your understanding
Four cells are marked on the table outline: barium (Group 2), manganese (Group 7), nickel (Group 10), and copper (Group 11). Which marked element is a main-group element?
ABarium.correct
BManganese.
This option is wrong — you called a Group 7 element main-group — Group 7 sits inside the 3-through-12 block.
CNickel.
This option is wrong — you called a Group 10 element main-group — Group 10 sits inside the 3-through-12 block.
DCopper.
This option is wrong — you called a Group 11 element main-group — Group 11 sits inside the 3-through-12 block.
Main-group columns are 1, 2, and 13 through 18. Barium's Group 2 is main-group; Groups 7, 10, and 11 are all middle-block columns.
Lesson 15 of 50 · PTB-015
Shells from the period number
Learning Journey
Read · Check · Unlock
Did You Know?
You can already write an atom's electron configuration, and each block's leading number names its shell — the main energy level it belongs to. For a main-group element, the table hands you the shell count with no configuration-writing at all.
The idea
For a main-group element, the period number equals the number of occupied electron shells.
Sulfur sits in Period 3, so a sulfur atom has three occupied shells.
For a main-group element, the period number is the shell count. Sulfur, Period 3: three occupied shells.
Check it against a configuration you can already write: sulfur is 1s²2s²2p⁶3s²3p⁴. Group the blocks by their leading number — level 1 holds 2 electrons, level 2 holds 8, level 3 holds 6 — and exactly three shells are occupied.
The rule works all the way down: cesium sits in Period 6, so a cesium atom has six occupied shells — no assigning needed.
Read the period, and you have the shell count.
Worked examples
Worked example 1. Bromine sits in Period 4. How many occupied electron shells does a bromine atom have?
Step 1
Bromine is a main-group element in Period 4.
Step 2
The period number equals the number of occupied shells.
Step 3
A bromine atom has four occupied shells.
Worked example 2. Strontium sits in Period 5. How many occupied electron shells does a strontium atom have?
Step 1
Strontium is a main-group element in Period 5.
Step 2
The period number equals the number of occupied shells — the row is deep, but the rule is the same.
Step 3
A strontium atom has five occupied shells.
You can now identify the number of occupied electron shells in an atom of a main-group element from its period number.
Check your understanding
Nitrogen sits in Period 2 of the periodic table. How many occupied electron shells does a nitrogen atom have?
Answer: 2
Nitrogen is a main-group element in Period 2. The period number equals the number of occupied shells. A nitrogen atom has two occupied shells.
Check your understanding
Potassium sits in Period 4 of the periodic table. How many occupied electron shells does a potassium atom have?
Answer: 4
Potassium is a main-group element in Period 4. The period number equals the number of occupied shells. A potassium atom has four occupied shells.
Check your understanding
Iodine sits in Period 5 of the periodic table. How many occupied electron shells does an iodine atom have?
Answer: 5
Iodine is a main-group element in Period 5. The period number equals the number of occupied shells. An iodine atom has five occupied shells.
Lesson 16 of 50 · PTB-016
Valence electrons from the group number
Learning Journey
Read · Check · Unlock
Did You Know?
You've already read valence-electron counts from the table's columns, back in the atoms unit. Those columns now have proper names — groups — and the reading deserves its full statement.
The idea
For a main-group element, the group number tells you the valence-electron count.
Groups 1 and 2 give the count directly: 1 and 2 valence electrons.
Groups 1 and 2 give the count directly; Groups 13-18 give it after subtracting 10. Chlorine, Group 17: seven valence electrons.
Groups 13 through 18 give the count after subtracting 10: 3 through 8 valence electrons.
Chlorine sits in Group 17, and 17 − 10 = 7, so a chlorine atom has seven valence electrons.
Magnesium sits in Group 2, so a magnesium atom has two valence electrons.
Helium is the one exception in Group 18: it has only two electrons in total, so it has two valence electrons.
The rule holds in every period, because the group fixes the valence count wherever the row sits.
Worked examples
Worked example 1. Aluminum sits in Group 13. How many valence electrons does an aluminum atom have?
Step 1
Aluminum is a main-group element in Group 13.
Step 2
Groups 13 through 18 give the group number minus 10: 13 − 10 = 3.
Step 3
An aluminum atom has three valence electrons.
Worked example 2. Krypton sits in Group 18, in Period 4. How many valence electrons does a krypton atom have?
Step 1
Krypton is a main-group element in Group 18.
Step 2
Groups 13 through 18 give the group number minus 10: 18 − 10 = 8.
Step 3
The period does not matter — Group 18 gives 8 in every row (helium excepted).
Step 4
A krypton atom has eight valence electrons.
You can now identify the number of valence electrons in an atom of a main-group element from its group number.
Check your understanding
Selenium sits in Group 16 of the periodic table. How many valence electrons does a selenium atom have?
Answer: 6
Selenium is a main-group element in Group 16. Groups 13 through 18 give the group number minus 10. 16 − 10 = 6 valence electrons.
Check your understanding
Rubidium sits in Group 1 of the periodic table, in Period 5. How many valence electrons does a rubidium atom have?
Answer: 1
Rubidium is a main-group element in Group 1. Groups 1 and 2 give the count directly. A rubidium atom has one valence electron, whatever its period.
Check your understanding
Gallium sits in Group 13 of the periodic table. How many valence electrons does a gallium atom have?
Answer: 3
Gallium is a main-group element in Group 13. Groups 13 through 18 give the group number minus 10. 13 − 10 = 3 valence electrons.
Lesson 17 of 50 · PTB-017
Why a group behaves alike
Learning Journey
Read · Check · Unlock
Have You Ever Wondered?
Wonder this:
Sodium and potassium are different elements — different sizes, different masses. Yet drop a piece of either into water and you get the same show: fizzing, and the same gas released. Why should two different elements put on the same show?
The answer sits in the valence electrons the table shows you.
The idea
An atom meets other atoms with its outermost electrons — the valence electrons are the ones that take part when atoms react.
So an atom's chemical behavior is set mainly by how many valence electrons it has.
Elements in the same group have the same number of valence electrons.
So elements in the same group behave in similar ways.
1 valence electron each
Different sizes, different shell counts — but the same single valence electron. That match is what water reacts with.
Sodium and potassium are both in Group 1, each with one valence electron — and that shared count is why water gets the same show from both.
Worked examples
Worked example 1. Magnesium and calcium react with the same partners in similar ways. Why?
Step 1
Both sit in Group 2, so each atom has two valence electrons.
Step 2
The valence electrons are the ones that take part when atoms react.
Step 3
Magnesium and calcium behave alike because their atoms have the same number of valence electrons.
You can now explain that elements in the same group behave similarly because their atoms have the same number of valence electrons.
Check your understanding
Fluorine is a pale gas and iodine is a dark solid, yet the two react with the same metals in similar ways. Why?
ABoth are in Group 17, so their atoms have the same number of valence electrons.correct
BTheir atoms have the same number of occupied electron shells.
This option is wrong — you counted shells — fluorine has two and iodine has five; it is the valence count they share.
CTheir atomic masses are nearly equal.
This option is wrong — you reached for mass — their masses differ hugely, and mass does not set chemical behavior.
DBoth sit in the same period of the table.
This option is wrong — you swapped group for period — they sit in different rows, and it is the shared column that matters.
The valence electrons are the ones that take part when atoms react. Fluorine and iodine are both in Group 17, so each atom has seven valence electrons. They behave alike because their atoms have the same number of valence electrons.
Check your understanding
Beryllium and barium behave in similar ways. What do their atoms share?
AThe same number of valence electrons — two each, as Group 2 elements.correct
BThe same number of occupied electron shells in their atoms.
This option is wrong — you counted shells — beryllium has two and barium has six; the shared count is the valence electrons.
CThe same atomic size.
This option is wrong — you reached for size — a barium atom is far larger, and size is not what sets the similar behavior.
DThe same atomic number.
This option is wrong — you gave them the same identity — different elements always have different atomic numbers; what they share is the valence count.
Both elements sit in Group 2, so each atom has two valence electrons. The valence electrons are the ones that take part when atoms react. They behave alike because their atoms have the same number of valence electrons.
Check your understanding
Which two elements should behave most alike?
ACalcium and strontium — the same group, so the same number of valence electrons.correct
BCalcium and potassium — next-door neighbors in the same period.
This option is wrong — you trusted closeness — period neighbors differ by one valence electron, and that one electron changes the behavior.
CCalcium and bromine — the same period.
This option is wrong — you used the row — same period means the valence count is different, so the behavior is too.
DPotassium and bromine — the two ends of one period.
This option is wrong — you paired across a whole row — their valence counts differ by six, about as unalike as main-group atoms get.
Similar behavior follows the shared valence count. Calcium and strontium are both in Group 2, with two valence electrons each. Same group, same valence count, similar behavior.
Lesson 18 of 50 · PTB-018
Why properties repeat
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen the periodic law: list the elements by atomic number, and elements with similar properties appear at regular intervals. You've also seen that elements in the same group behave alike because their atoms have the same number of valence electrons. Put those two ideas together, and the repetition itself has an explanation.
The idea
Walk across Period 2 from lithium to neon, and the valence-electron count climbs one step at a time: 1, 2, 3, 4, 5, 6, 7, 8.
Walk across Period 3 from sodium to argon, and the count runs through exactly the same climb: 1 up to 8.
Each new period starts the valence-electron count over at 1 and climbs again.
An atom's chemical behavior follows its valence electrons.
So when the same valence-electron count comes back around, the same kind of behavior comes back with it.
Properties repeat at regular intervals because the number of valence electrons runs through the same sequence in every period.
Worked examples
Worked example 1. Fluorine in Period 2 and chlorine in Period 3 behave in very similar ways. Use the valence-electron sequence to explain why.
Step 1
Fluorine sits at the 7-valence-electron step of Period 2's climb.
Step 2
Chlorine sits at the 7-valence-electron step of Period 3's climb.
Step 3
The same count returns, so the same behavior returns.
Step 4
Fluorine and chlorine behave alike because each period runs the same valence-electron sequence, and both elements sit at the seven-electron step of their period.
Worked example 2. A student lists the elements in atomic-number order and marks every element that behaves like beryllium. The marks land at regular intervals. Why?
Step 1
Beryllium's atoms have two valence electrons.
Step 2
Each period restarts the count, so the two-electron step returns once in every period.
Step 3
Elements matching beryllium's behavior appear once per period — at regular intervals — because every period passes through the two-valence-electron step exactly once.
You can now explain that properties repeat at regular intervals across the periodic table because the number of valence electrons runs through the same sequence in every period.
Check your understanding
When the elements are listed in order of increasing atomic number, elements with similar properties appear at regular intervals. What causes the repetition?
AEach period runs the valence-electron count through the same sequence, and behavior follows the valence electrons.correct
BAtomic mass grows by the same amount from each element to the next, and behavior follows the mass.
This option is wrong — you tied the repeats to atomic mass — mass just keeps growing and never repeats; it is the valence-electron count that starts over each period.
CAll the elements in a period have the same number of valence electrons, so each period behaves as one block.
This option is wrong — you swapped rows and columns — a period climbs through the counts 1 to 8; it is a group whose members share one count.
DThe number of occupied shells starts over at 1 in every period.
This option is wrong — you restarted the shell count — each new period adds one more shell and keeps it; it is the valence-electron count that starts over.
Each period runs the valence-electron count through the same climb, starting over at 1. An atom's behavior follows its valence electrons. So the same behavior returns each time the same count returns — at regular intervals.
Check your understanding
Nitrogen and phosphorus are both in Group 15 — nitrogen in Period 2, phosphorus in Period 3 — and they behave in similar ways. Which statement explains the similarity?
AEach period runs the same valence-electron sequence, and both elements sit at the five-electron step of their period.correct
BBoth elements hold the same number of occupied shells, so their electrons are arranged identically.
This option is wrong — you matched their shells — nitrogen has two occupied shells and phosphorus has three; it is their valence-electron count of five that matches.
CBoth elements have the same atomic number.
This option is wrong — you matched their atomic numbers — every element has its own atomic number; the shared feature is the valence-electron count.
This option is wrong — you nested one element inside the other — each element's atoms are its own; the two elements simply share a valence-electron count of five.
Period 2 and Period 3 each climb through the valence-electron counts 1 to 8. Nitrogen and phosphorus both sit at the five-electron step of their period. The same count means the same kind of behavior.
Check your understanding
Aluminum behaves much like boron, and boron's atoms have three valence electrons. Using the reason properties repeat, how many valence electrons does an aluminum atom have?
AThree — matching behavior means the same valence-electron count has come back around.correct
BFour — aluminum sits one period lower, so it carries one more valence electron.
This option is wrong — you let the count grow with the period — moving down a group adds a shell, not a valence electron; the count that returns is the same three.
CMore than boron, because aluminum's atoms hold more electrons in total.
This option is wrong — you counted all the electrons — most of aluminum's extra electrons sit in inner shells; behavior follows the three in the outer shell.
DIt cannot be worked out from behavior alone.
This option is wrong — you separated behavior from the valence count — matching behavior is exactly the sign of a matching valence-electron count; that link is the periodic law's explanation.
Behavior follows the valence electrons. Boron's behavior comes with three valence electrons, so an element that behaves the same way carries the same count. Aluminum's atoms have three valence electrons.
Lesson 19 of 50 · PTB-019
The alkali metals
Learning Journey
Read · Check · Unlock
Did You Know?
Wonder this:
In a chemistry stockroom, sodium is kept in a jar of oil. Not to keep it clean — to keep the air away from it.
Sodium sits in Group 1, at the far left of the periodic table. What it does in open air is the family trademark.
The idea
Sodium is a soft, shiny metal — soft enough to cut with a table knife.
Left in open air, sodium's fresh surface dulls within minutes, and in water it reacts violently.
An element that reacts quickly and easily like this is called 'reactive'.
The other Group 1 metals — lithium, potassium, rubidium, and cesium — are also soft, shiny, and highly reactive.
Together, the Group 1 elements except hydrogen are the 'alkali metals'.
Hydrogen sits at the top of Group 1, but it is a nonmetal gas, not an alkali metal.
Alkali metals are so reactive that none of them is ever found uncombined in nature — every sample dug from the ground is already combined with other elements.
Worked examples
Worked example 1. Which elements are the alkali metals?
Step 1
Answer: the Group 1 elements except hydrogen.
Worked example 2. Why is potassium never found as a pure metal in nature?
Step 1
Answer: potassium is an alkali metal — so reactive that in nature it is always found combined with other elements.
You can now state that the Group 1 elements except hydrogen are the alkali metals: soft, shiny metals that react quickly with water and air, so reactive that they are never found uncombined in nature.
Check your understanding
Which elements make up the alkali metals?
AThe Group 1 elements except hydrogen.correct
BAll of the Group 1 elements, including hydrogen.
This option is wrong — you included hydrogen — it sits at the top of Group 1 but is a nonmetal gas, not an alkali metal.
CThe Group 2 elements.
This option is wrong — you moved one column over — Group 2 is a different family; the alkali metals fill Group 1 below hydrogen.
DThe elements of Groups 3 through 12.
This option is wrong — you took the transition-metal block — the alkali metals are the single leftmost column, Group 1, minus hydrogen.
The alkali metals are the Group 1 elements except hydrogen. Hydrogen shares the column but is a nonmetal gas. Lithium, sodium, potassium, rubidium, and cesium are the family.
Check your understanding
A freshly cut piece of lithium is shiny, but its surface turns dull within minutes in open air. Which alkali-metal property does this show?
AAlkali metals react quickly with air.correct
BAlkali metals are soft enough to cut.
This option is wrong — you answered from the cutting, not the dulling — softness let the student cut it, but the question asks about the surface change in air.
CAlkali metals melt at room temperature.
This option is wrong — you turned dulling into melting — the lithium stays solid; its surface dulls because the metal reacts with the air.
DAlkali metals are shiny only in the dark.
This option is wrong — you made shine depend on light — the fresh surface is shiny anywhere; it dulls because the metal reacts with air, not because of lighting.
A fresh alkali-metal surface is shiny. In open air the metal reacts quickly, and the reacted surface is dull. Fast dulling in air is the family's reactivity on display.
Check your understanding
Rubidium is shipped in sealed containers with all air and moisture removed. Why?
ARubidium reacts quickly with air and with water, so both must be kept away from it.correct
BRubidium is a gas and would escape an unsealed container.
This option is wrong — you made rubidium a gas — it is a soft solid metal; the sealing keeps air and moisture away from it.
CRubidium loses its shine in light, and the container blocks the light.
This option is wrong — you blamed the light — the danger is air and moisture, which the metal reacts with quickly.
DRubidium is too soft to hold its shape without a container.
This option is wrong — you answered from softness — soft metals still hold their shape; the seal is there because rubidium reacts quickly with air and water.
Rubidium is an alkali metal. Alkali metals react quickly with air and with water. Sealing out both is the only way to keep the metal unreacted.
Lesson 20 of 50 · PTB-020
Why alkali metals are so reactive
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen what alkali metals do: fresh surfaces dull within minutes in open air, and water sets off violent reactions. You've also seen that a Group 1 atom has one valence electron, and that an atom which loses an electron becomes an ion. Here is why the family is so reactive.
The idea
A metal atom reacts by giving away its outer electrons, and it reacts more easily when those electrons are easier to remove.
An alkali metal atom has a single valence electron.
That single outer electron is easy to remove.
With its one valence electron given away so easily, an alkali metal atom reacts quickly — that is what makes the whole family so reactive.
Worked examples
Worked example 1. Why does potassium react rapidly with water?
Step 1
A metal atom reacts by giving away its outer electrons, and it reacts more easily when those electrons are easier to remove.
Step 2
A potassium atom has a single valence electron, and that electron is easy to remove.
Step 3
Potassium reacts rapidly because its atoms give away their single, easily removed valence electron.
Worked example 2. A strip of lithium dulls quickly in open air. Explain the dulling at the level of lithium's atoms.
Step 1
Each lithium atom has one valence electron that is easy to remove.
Step 2
At the surface, lithium atoms give that electron away to the air's particles and end up combined with them.
Step 3
The surface dulls because lithium atoms react quickly, giving away their single easily removed valence electron to combine with the air.
You can now explain that alkali metals are very reactive because a metal atom reacts by giving away its outer electrons, and it reacts more easily when those electrons are easier to remove — an alkali metal atom has a single valence electron that is easy to remove.
Check your understanding
Why are the alkali metals so reactive?
AA metal atom reacts by giving away its outer electrons, and an alkali metal atom's single valence electron is easy to remove.correct
BAn alkali metal atom pulls in extra electrons very strongly, filling its outer shell from its surroundings.
This option is wrong — you gave the nonmetal account — a metal atom reacts by giving electrons away, not by pulling more in.
CTheir atoms have many valence electrons to give away.
This option is wrong — you multiplied the valence electrons — an alkali metal atom has exactly one, and its easy removal is the point.
DThey are soft, and soft metals react faster.
This option is wrong — you used softness as the cause — softness and reactivity are two family traits, but the reacting is done by the easily removed valence electron.
A metal atom reacts by giving away its outer electrons, and it reacts more easily when those electrons are easier to remove. An alkali metal atom has a single valence electron that is easy to remove. That easy give-away is the family's reactivity.
Check your understanding
Cesium reacts explosively with water. What do its atoms do during the reaction?
AEach atom gives away its single valence electron.correct
BEach atom pulls an extra electron in from the water.
This option is wrong — you ran the reaction backwards — cesium is a metal, and a metal atom reacts by giving its outer electron away.
CEach atom gives away all of its electrons.
This option is wrong — you emptied the whole atom — only the single valence electron in the outer shell is given away; the inner shells keep their electrons.
DEach atom splits into smaller atoms.
This option is wrong — you broke the atom apart — the atom stays whole; it becomes an ion by losing its one valence electron.
A metal atom reacts by giving away its outer electrons. A cesium atom has a single valence electron, and it is easy to remove. Giving that one electron away is the whole event at the atomic level.
Check your understanding
An alkali metal atom's valence electron is easy to remove. What does that mean for how the metal behaves?
AThe metal reacts quickly, because a metal atom reacts more easily when its outer electrons are easier to remove.correct
BThe metal conducts electricity poorly, because it keeps losing its electrons.
This option is wrong — you turned easy removal into poor conduction — conduction is not this lesson's claim; easy removal makes the atom react more easily.
CThe metal is harder, because atoms missing an electron pack more tightly.
This option is wrong — you connected electron loss to hardness — alkali metals are in fact soft; easy electron removal shows up as fast reaction.
DThe metal reacts slowly, because each atom must find its lost electron again before it can react again.
This option is wrong — you made the lost electron a delay — giving the electron away IS the reaction, so easy removal means fast reaction.
A metal atom reacts by giving away its outer electrons, and it reacts more easily when those electrons are easier to remove. Easy removal therefore means quick reaction. That is exactly the alkali metals' behavior in air and water.
Lesson 21 of 50 · PTB-021
The alkaline earth metals
Learning Journey
Read · Check · Unlock
Did You Know?
One column to the right of the alkali metals sits Group 2. Its members are metals with a family personality of their own — recognizably reactive, but tamer.
The idea
The Group 2 elements are the 'alkaline earth metals': beryllium, magnesium, calcium, strontium, barium, and radium.
Alkaline earth metals are harder than the alkali metals.
They also react with water and with air, but less vigorously than the alkali metals do.
Magnesium, for example, reacts with water far more slowly than sodium does — a magnesium ribbon can sit in water for days while sodium reacts within seconds.
Worked examples
Worked example 1. Which elements are the alkaline earth metals?
Step 1
Answer: the Group 2 elements.
Worked example 2. Calcium dropped into water fizzes steadily rather than exploding. Is that in keeping with its family?
Step 1
Answer: yes — calcium is an alkaline earth metal, and the family reacts with water less vigorously than the alkali metals do.
You can now state that the Group 2 elements are the alkaline earth metals: metals that are harder than the alkali metals and also react with water and air, but less vigorously.
Check your understanding
Which elements are the alkaline earth metals?
AThe Group 2 elements.correct
BThe Group 1 elements except hydrogen.
This option is wrong — you named the alkali metals — the alkaline earth metals are the next column over, Group 2.
CThe elements of Groups 3 through 12.
This option is wrong — you took the transition-metal block — the alkaline earth metals are the single column labeled Group 2.
DThe Group 17 elements.
This option is wrong — you picked a nonmetal column — Group 17 holds nonmetals; the alkaline earth metals are the Group 2 metals.
The alkaline earth metals are the Group 2 elements. Beryllium, magnesium, calcium, strontium, barium, and radium make up the family.
Check your understanding
How does an alkaline earth metal's reaction with water compare with an alkali metal's?
AIt reacts, but less vigorously than an alkali metal does.correct
BIt reacts more violently than an alkali metal does.
This option is wrong — you flipped the comparison — Group 2 metals are the tamer family; the alkali metals react more violently.
CIt does not react with water at all.
This option is wrong — you switched the family off — alkaline earth metals do react with water, just less vigorously than the alkali metals.
DIt reacts identically — the two families cannot be told apart in water.
This option is wrong — you erased the difference — the pace of the water reaction is exactly how the two families differ.
Alkaline earth metals react with water and with air. The reactions run less vigorously than the alkali metals' do. Reactive, but tamer, is the Group 2 profile.
Check your understanding
A strip of strontium reacts steadily with water, and the metal is too hard to cut with a table knife. Which family fits?
AThe alkaline earth metalscorrect
BThe alkali metals
This option is wrong — you ignored the hardness — an alkali metal is knife-soft and reacts violently; harder metal plus a steadier reaction is Group 2's profile.
CThe transition metals
This option is wrong — you reached for Groups 3 through 12 — transition metals do not react steadily with water; a water-reactive but hard metal fits Group 2.
DThe Group 17 nonmetals
This option is wrong — you chose a nonmetal column — Group 17 sits past the stair-step line; the sample described is a metal.
Reacts with water, but not violently: that rules out the alkali metals. Harder than knife-soft also points away from Group 1. Water-reactive, harder metal is the alkaline earth profile — Group 2.
Lesson 22 of 50 · PTB-022
Why Group 2 reacts less vigorously than Group 1
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen the two families' water reactions side by side: a Group 1 metal erupts while its Group 2 neighbor fizzes steadily. Both are metals giving away electrons — so why is one hand-over so much faster?
The idea
A metal atom reacts by giving away its outer electrons, and it reacts more easily when those electrons are easier to remove.
A Group 1 atom gives away one valence electron; a Group 2 atom must give away two.
Removing a Group 2 atom's two valence electrons takes more energy than removing a Group 1 atom's one.
The harder hand-over makes the reaction slower — so alkaline earth metals react less vigorously than alkali metals.
Worked examples
Worked example 1. Why does calcium react with water more slowly than potassium does?
Step 1
A metal atom reacts by giving away its outer electrons, and it reacts more easily when those electrons are easier to remove.
Step 2
A potassium atom gives away one valence electron; a calcium atom must give away two.
Step 3
Removing calcium's two valence electrons takes more energy than removing potassium's one.
Step 4
Calcium reacts more slowly because giving away its two valence electrons takes more energy than potassium's single easy hand-over.
You can now explain that alkaline earth metals react less vigorously than alkali metals because a metal atom reacts by giving away its outer electrons, and it reacts more easily when those electrons are easier to remove — removing a Group 2 atom's two valence electrons takes more energy than removing a Group 1 atom's one.
Check your understanding
Why do alkaline earth metals react less vigorously than alkali metals?
ARemoving a Group 2 atom's two valence electrons takes more energy than removing a Group 1 atom's one.correct
BA Group 2 atom has fewer valence electrons to give away.
This option is wrong — you reversed the counts — a Group 2 atom has two valence electrons to the Group 1 atom's one, and removing two takes more energy.
CGroup 2 metals are harder, and harder metals always react more slowly.
This option is wrong — you used hardness as the cause — hardness is another family trait; the pace is set by how much energy the electron hand-over takes.
DGroup 2 atoms react by pulling extra electrons in from outside instead of giving their own away.
This option is wrong — you gave Group 2 the nonmetal account — both families are metals, and both react by giving electrons away; Group 2's two-electron hand-over just costs more energy.
A metal atom reacts by giving away its outer electrons, and it reacts more easily when those electrons are easier to remove. A Group 2 atom must give away two valence electrons, and removing two takes more energy than removing one. The costlier hand-over is the slower reaction.
Check your understanding
A Group 1 metal and its Group 2 neighbor in the same period both meet water. Which atom's valence electrons take more energy to remove, and which metal reacts faster?
AThe Group 2 atom's — so the Group 1 metal reacts faster.correct
BThe Group 1 atom's — so the Group 2 metal reacts faster.
This option is wrong — you flipped both halves — the Group 1 atom gives away one easily removed electron and reacts faster; the Group 2 atom's two-electron removal takes more energy.
CThe Group 2 atom's — and its harder removal makes the Group 2 metal react faster.
This option is wrong — you attached harder removal to faster reaction — a metal reacts more easily when its electrons are EASIER to remove, so the harder removal is the slower reaction.
DThey take the same energy — the two metals react with water at exactly the same pace.
This option is wrong — you leveled the two atoms — removing two valence electrons takes more energy than removing one, and the water reactions differ accordingly.
The Group 1 atom's single valence electron is easier to remove than the Group 2 atom's two. A metal atom reacts more easily when its outer electrons are easier to remove. So the Group 1 metal reacts faster in water.
Check your understanding
Rubidium and strontium are Period 5 neighbors, in Group 1 and Group 2. Predict the difference in their reactions with water, with the reason.
ARubidium reacts more vigorously, because giving away its one valence electron takes less energy than strontium giving away two.correct
BStrontium reacts more vigorously, because giving away two electrons instead of one means twice as much reaction.
This option is wrong — you counted electrons as reaction strength — what sets the pace is the energy the removal takes, and removing two costs more, slowing strontium down.
CThe two react identically, because both are metals that give away electrons.
This option is wrong — you stopped at the shared mechanism — both give electrons away, but rubidium's single electron is easier to remove, so its reaction is more vigorous.
DNeither reacts with water, because Period 5 metals are unreactive.
This option is wrong — you switched both families off by period — family behavior follows the group; both react, rubidium the more vigorously.
A metal atom reacts by giving away its outer electrons, and it reacts more easily when those electrons are easier to remove. Rubidium hands over one valence electron; strontium must hand over two, which takes more energy. Rubidium therefore reacts more vigorously.
Lesson 23 of 50 · PTB-023
The halogens
Learning Journey
Read · Check · Unlock
Did You Know?
Wonder this:
The white crystals on your table — table salt — were made by letting a soft silvery metal meet a pale green toxic gas.
The gas is chlorine, from Group 17 on the right-hand side of the periodic table. Its column is a family of nonmetals with a sharp group personality.
The idea
The Group 17 elements are the 'halogens': fluorine, chlorine, bromine, and iodine.
Halogens are very reactive nonmetals.
Halogens travel as two-atom molecules — each moving particle is a pair of identical atoms.
Halogens react with metals, and the compound a halogen forms with a metal is called a 'salt'.
Chlorine reacting with sodium gives the salt sodium chloride — table salt itself.
Worked examples
Worked example 1. Bromine reacts with potassium. What kind of compound forms, and what is it called?
Step 1
Answer: a salt — potassium bromide.
Worked example 2. A sample of pure fluorine contains ten fluorine atoms. How many molecules is that?
Step 1
Answer: five — halogens travel as two-atom molecules, so ten atoms move as five pairs.
You can now state that the Group 17 elements are the halogens: very reactive nonmetals that travel as two-atom molecules and react with metals to form salts.
Check your understanding
Which elements are the halogens?
AThe Group 17 elements.correct
BThe Group 18 elements.
This option is wrong — you moved one column right — Group 18 is a different family; the halogens are the Group 17 column.
CThe Group 1 elements except hydrogen.
This option is wrong — you named the alkali metals — those are reactive METALS on the left; the halogens are reactive nonmetals in Group 17.
DAll the nonmetals in the periodic table.
This option is wrong — you widened the family to every nonmetal — the halogens are only the Group 17 column: fluorine, chlorine, bromine, and iodine.
The halogens are the Group 17 elements. Fluorine, chlorine, bromine, and iodine make up the family.
Check your understanding
Iodine reacts with lithium. What forms?
AA salt — lithium iodide.correct
BA mixture of lithium and iodine that could be separated back out.
This option is wrong — you mingled the elements without reacting them — a halogen reacts with a metal, and the product is a new compound, a salt.
CA new metal.
This option is wrong — you kept the product metallic — the reaction of a metal with a halogen makes a salt, a compound unlike either element.
DNothing — iodine does not react with metals.
This option is wrong — you switched off a halogen's trademark reaction — reacting with metals to form salts is exactly what the family does.
Halogens react with metals. The compound a halogen forms with a metal is called a salt. Lithium and iodine give the salt lithium iodide.
Check your understanding
In a container of pure iodine vapor, how do the iodine atoms travel?
AIn pairs — each molecule is two iodine atoms.correct
BAlone — each particle of the vapor is a single iodine atom.
This option is wrong — you split the pairs — halogens travel as two-atom molecules, so the moving particles are atom pairs.
CIn clusters of eight atoms each.
This option is wrong — you oversized the molecule — a halogen molecule is exactly two atoms.
DJoined to potassium atoms.
This option is wrong — you put a metal inside pure iodine — a metal only enters when the halogen reacts; pure iodine is pairs of iodine atoms.
Halogens travel as two-atom molecules. Pure iodine is pairs of iodine atoms moving together.
Lesson 24 of 50 · PTB-024
Why halogens are so reactive
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen that halogens are very reactive nonmetals, and that a Group 17 atom has seven valence electrons. The alkali metals reacted by giving an electron away — a halogen atom does the opposite.
The idea
A nonmetal atom reacts by pulling in extra electrons, and it reacts more easily when that pull is stronger.
A halogen atom has seven valence electrons — it needs just one more electron to fill its outer shell.
Being one electron short, a halogen atom pulls in that missing electron very readily.
That ready pull is why the halogens are so reactive.
Worked examples
Worked example 1. Why does fluorine react so readily with metals?
Step 1
A nonmetal atom reacts by pulling in extra electrons, and it reacts more easily when that pull is stronger.
Step 2
A fluorine atom has seven valence electrons and needs just one more to fill its outer shell.
Step 3
Fluorine reacts readily because each atom pulls in the one electron it needs to fill its outer shell.
Worked example 2. When bromine meets a metal, which way do electrons move, and why?
Step 1
The metal atom reacts by giving away its outer electrons.
Step 2
The bromine atom, one electron short of a full outer shell, pulls an electron in.
Step 3
Electrons move from the metal atoms to the bromine atoms — the metal gives, and the one-electron-short bromine pulls in.
You can now explain that halogens are very reactive because a nonmetal atom reacts by pulling in extra electrons, and it reacts more easily when that pull is stronger — a halogen atom needs just one more electron to fill its outer shell.
Check your understanding
Why are the halogens so reactive?
AA nonmetal atom reacts by pulling in extra electrons, and a halogen atom needs just one more electron to fill its outer shell.correct
BA halogen atom gives away its seven valence electrons easily, emptying its outer shell when it reacts.
This option is wrong — you gave the metal account — a halogen atom keeps its seven and pulls in the one electron it is short.
CTheir atoms already have full outer shells.
This option is wrong — you filled the shell too early — a full outer shell is the profile of the UNreactive family next door; a halogen atom is one electron short.
DThey travel as two-atom molecules, and paired atoms react faster.
This option is wrong — you used the pairing as the cause — the molecules are a family trait, but the reactivity comes from each atom's strong pull for its missing electron.
A nonmetal atom reacts by pulling in extra electrons, and it reacts more easily when that pull is stronger. A halogen atom has seven valence electrons — one short of a full outer shell. Pulling in that one missing electron is the family's reactivity.
Check your understanding
When iodine reacts with a metal, what does each iodine atom do?
AIt pulls in one electron from a metal atom.correct
BIt gives one of its valence electrons to the metal atom.
This option is wrong — you reversed the electron's direction — the metal gives; the iodine atom, one short of a full shell, pulls the electron in.
CIt pulls in seven electrons to double its outer shell.
This option is wrong — you filled the wrong gap — the atom already holds seven valence electrons and needs just one more.
DIt trades its nucleus with the metal atom's nucleus.
This option is wrong — you moved the nuclei — nuclei stay put in chemistry; the change is one electron pulled in from the metal.
A halogen atom has seven valence electrons and needs one more. A nonmetal atom reacts by pulling in extra electrons. Each iodine atom pulls one electron in from a metal atom.
Check your understanding
Potassium meets iodine, and the two react. Which account names what BOTH kinds of atoms are doing?
APotassium atoms give away their single valence electron, and iodine atoms pull in the one electron they are short.correct
BBoth kinds of atoms give electrons away.
This option is wrong — you made iodine a giver — a nonmetal atom reacts by pulling electrons in; only the metal gives.
CBoth kinds of atoms pull electrons in.
This option is wrong — you made potassium a puller — a metal atom reacts by giving its outer electron away; only the iodine pulls.
DPotassium atoms pull in seven electrons from the iodine, and each iodine atom gives away one of its own.
This option is wrong — you swapped the two roles — the metal gives away one electron; the halogen pulls in one.
A metal atom reacts by giving away its outer electrons. A nonmetal atom reacts by pulling in extra electrons. Potassium gives one; iodine, one short of a full shell, pulls one in.
Lesson 25 of 50 · PTB-025
The noble gases
Learning Journey
Read · Check · Unlock
Did You Know?
Wonder this:
A party balloon filled with helium can sit against a birthday candle's smoke, night after night, and nothing about the gas ever changes.
Helium heads the far-right column of the periodic table, Group 18. The whole column shares its indifference.
The idea
The Group 18 elements are the 'noble gases': helium, neon, argon, krypton, xenon, and radon.
Noble gases are colorless gases at room temperature.
Noble gases almost never react with anything — not with metals, not with nonmetals, not with each other.
Helium is used in balloons partly because it does not react or burn.
Worked examples
Worked example 1. Which elements are the noble gases?
Step 1
Answer: the Group 18 elements.
Worked example 2. Light bulbs are filled with argon rather than air. What noble-gas property makes argon the safer filling?
Step 1
Answer: argon almost never reacts, so it leaves the hot metal parts of the bulb unchanged where air would react with them.
You can now state that the Group 18 elements are the noble gases: colorless gases that almost never react with anything.
Check your understanding
Which elements are the noble gases?
AThe Group 18 elements.correct
BThe Group 17 elements.
This option is wrong — you moved one column left — Group 17 holds the very reactive halogens; the noble gases fill Group 18.
CAll the elements that are gases at room temperature.
This option is wrong — you widened the family to every gas — chlorine and others are gases too; the noble gases are only the Group 18 column.
DThe Group 1 elements except hydrogen.
This option is wrong — you named the alkali metals — those are the most reactive metals; the noble gases are the least reactive elements of all.
The noble gases are the Group 18 elements. Helium, neon, argon, krypton, xenon, and radon make up the family.
Check your understanding
What happens when krypton is mixed with a hot, freshly cut metal surface?
ANothing — krypton almost never reacts with anything.correct
BA salt forms on the metal's surface.
This option is wrong — you gave krypton the halogen's trademark — salts come from Group 17 reacting with metals; a noble gas leaves the metal unchanged.
CThe krypton burns with a colored flame.
This option is wrong — you burned an unreactive gas — noble gases do not burn; burning is itself a reaction.
DThe metal dissolves into the gas.
This option is wrong — you dissolved a solid metal into a gas — the two simply share the space; the noble gas reacts with nothing.
Krypton is a noble gas. Noble gases almost never react — not with metals, not with nonmetals. The hot metal stays unchanged.
Check your understanding
The gas sealed inside a glass display tube has stayed chemically unchanged for forty years. The gas is colorless when the power is off. Which family does the filling most likely come from?
AThe noble gasescorrect
BThe halogens
This option is wrong — you filled the tube with the most reactive nonmetals — a halogen would have attacked the tube's metal parts decades ago.
CThe alkali metals
This option is wrong — you filled a tube with metals — alkali metals are solids, and famously reactive ones.
DThe alkaline earth metals
This option is wrong — you chose Group 2 — solid reactive metals, not colorless unreactive gases.
A colorless gas that stays unchanged for decades matches one family. Noble gases are colorless and almost never react. Group 18 is the filling.
Lesson 26 of 50 · PTB-026
Why noble gases do not react
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen the pattern: alkali metals give an electron away, halogens pull one in, and both react eagerly. Noble gases do neither. The reason sits in their outer shells.
The idea
A noble-gas atom has a full outer shell.
For helium that full outer shell is the first shell, complete with just two electrons; for the other noble gases the full outer shell holds eight.
A full outer shell is a stable arrangement, so the atom has little tendency to gain, lose, or share electrons.
An atom that neither gains, loses, nor shares electrons takes part in almost no reactions — that is the noble gases' unreactivity.
Worked examples
Worked example 1. Why does neon form no everyday compounds?
Step 1
A neon atom's outer shell is full, with eight electrons.
Step 2
A full outer shell is a stable arrangement, so the atom has little tendency to gain, lose, or share electrons.
Step 3
Neon forms no everyday compounds because its atoms, with full outer shells, have little tendency to gain, lose, or share electrons — and compounds only form when atoms do one of those.
Worked example 2. Argon sits in the same air as fiercely reactive oxygen and never combines with it. Explain argon's indifference.
Step 1
An argon atom's outer shell is full.
Step 2
With a full outer shell, the atom has little tendency to gain, lose, or share electrons.
Step 3
Argon stays uncombined because its full outer shell is a stable arrangement — the atom has almost no tendency to trade electrons with anything, oxygen included.
You can now explain that noble gases are unreactive because a full outer shell is a stable arrangement, so the atom has little tendency to gain, lose, or share electrons.
Check your understanding
Why are the noble gases unreactive?
AA full outer shell is a stable arrangement, so the atom has little tendency to gain, lose, or share electrons.correct
BTheir atoms have no electrons at all to trade.
This option is wrong — you emptied the atoms — noble-gas atoms have plenty of electrons; it is the FULL outer shell that removes the tendency to trade them.
CTheir atoms need one more electron and cannot find it.
This option is wrong — you gave them the halogen's shell — a halogen atom is one electron short; a noble-gas atom's outer shell is already full.
DTheir particles spread too far apart as gases to ever meet another particle and react with it.
This option is wrong — you used gas spacing as the cause — gas particles collide constantly; the atoms stay unreacted because their full outer shells are stable.
A noble-gas atom's outer shell is full. A full outer shell is a stable arrangement, so the atom has little tendency to gain, lose, or share electrons. No gaining, losing, or sharing means almost no reactions.
Check your understanding
A xenon atom meets a sodium atom, which readily gives away its valence electron. Why does the xenon atom still not react?
AIts outer shell is already full, so it has little tendency to gain the offered electron.correct
BIts outer shell is one electron short, but sodium's electron is the wrong kind.
This option is wrong — you made xenon a halogen with picky taste — electrons have no kinds, and xenon's outer shell is full, not one short.
CXenon atoms are too heavy to take part in reactions.
This option is wrong — you used weight as the cause — heavy atoms react fine elsewhere; xenon's full outer shell is what removes the tendency.
DThe xenon atom gives its own electrons to the sodium instead.
This option is wrong — you turned xenon into the giver — sodium is the metal doing the giving, and xenon, with a full stable shell, has little tendency to lose electrons either.
A xenon atom's outer shell is full. A full outer shell is a stable arrangement — little tendency to gain, lose, or share. Even a freely offered electron finds no vacancy worth taking.
Check your understanding
Helium's outer shell holds only two electrons, yet helium is just as unreactive as the rest of its family. Why?
AHelium's outer shell is the first shell, which is full with two — and a full outer shell is a stable arrangement.correct
BHelium is unreactive despite its outer shell sitting six electrons short of the eight a full shell needs.
This option is wrong — you measured helium's shell against eight — its outer shell is the FIRST shell, and the first shell is complete at two.
CHelium constantly borrows six electrons from nearby atoms to stay full.
This option is wrong — you had helium borrowing — no borrowing happens; two electrons already fill its first shell completely.
DHelium is unreactive only because it is the lightest noble gas.
This option is wrong — you used lightness as the cause — the stability comes from the full outer shell, which for helium is the two-electron first shell.
The first shell is full with two electrons. Helium's outer shell IS the first shell, so it is full. A full outer shell is a stable arrangement — helium has little tendency to gain, lose, or share electrons.
Lesson 27 of 50 · PTB-027
The transition metals
Learning Journey
Read · Check · Unlock
Did You Know?
Between Group 2 and Group 13 sits the wide middle block of the periodic table — Groups 3 through 12. You've already seen the block's name: the transition metals. Here is the family profile.
The idea
The transition metals are typical metals: hard, dense, and good conductors of heat and electricity.
Iron, copper, silver, and gold are all transition metals.
Many transition metals form colorful compounds.
Copper's compounds, for example, are often blue or green — the green skin on old copper roofs is a copper compound.
Worked examples
Worked example 1. Chromium is in Group 6. What properties would you expect of the metal itself?
Step 1
Answer: a typical metal — hard, dense, and a good conductor; chromium's compounds are also often colorful.
Worked example 2. A jeweler's silver wire conducts electricity extremely well and holds its shape under hard wear. Which two transition-metal traits is the wire showing?
Step 1
Answer: good conduction and hardness — both typical of the transition metals.
You can now state that the transition metals, in Groups 3 through 12, are typical metals — hard, dense, and good conductors — and that many of them form colorful compounds.
Check your understanding
Which properties are typical of the transition metals?
AHard, dense, and good conductors of heat and electricity.correct
BSoft, knife-cuttable, and violently reactive with water.
This option is wrong — you gave the alkali-metal profile — that is Group 1; transition metals are hard and are not violently water-reactive.
CDull, brittle, and poor conductors.
This option is wrong — you gave the nonmetal profile — transition metals are typical METALS: hard, dense, conducting.
DColorless gases that almost never react.
This option is wrong — you gave the noble-gas profile — transition metals are solid, dense metals.
Transition metals are typical metals. Hard, dense, and good conductors of heat and electricity. Many of them also form colorful compounds.
Check your understanding
A pinch of a nickel compound stirred into molten glass turns the whole batch green. Which transition-metal trait is on display?
AMany transition metals form colorful compounds.correct
BTransition metals conduct electricity well.
This option is wrong — you answered with conduction — the visible clue is the colored compound in the glass, the family's colorful-compound trait.
CTransition metals are dense.
This option is wrong — you answered with density — the visible clue is the colored compound, the family's colorful-compound trait.
DTransition metals and their compounds are colorless in every form.
This option is wrong — you removed the family's color trait — the metals look metallic, but their COMPOUNDS are often strongly colored, as the green glass shows.
The green tint comes from the nickel compound. Many transition metals form colorful compounds. Colorful compounds are a transition-metal trademark.
Check your understanding
A metal is hard and dense, conducts well, shows no reaction when dropped into water, and its compounds are a vivid green. Which family fits best?
AThe transition metalscorrect
BThe alkali metals
This option is wrong — you overlooked the water test and hardness — an alkali metal is soft and reacts vigorously with water.
CThe alkaline earth metals
This option is wrong — you picked Group 2 — its metals do react with water, if less vigorously; a hard, water-indifferent metal with colorful compounds is a transition metal.
DThe noble gases
This option is wrong — you chose a family of gases — the sample is a hard, dense metal.
Hard, dense, good conductor: typical metal. No water reaction rules out Groups 1 and 2. Vivid green compounds seal it: a transition metal.
Lesson 28 of 50 · PTB-028
Predict a reaction by family pattern
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen why a group behaves alike: same valence electrons, same behavior. That sameness is a prediction machine — know ONE reaction, and you can write the family's whole page.
The idea
Sodium reacts with chlorine to form the salt sodium chloride.
Swap the metal for its groupmate potassium, and the pattern holds: potassium reacts with chlorine to form potassium chloride.
Swap the halogen instead: sodium reacts with bromine to form sodium bromide.
Elements in the same group have the same number of valence electrons, so they react in the same pattern.
To predict a reaction, take a known reaction and swap an element for one of its groupmates — the product follows the same pattern, with the new element's name in place.
Worked examples
Worked example 1. Magnesium burns in oxygen to form magnesium oxide. Calcium is in magnesium's group. Predict the product when calcium burns in oxygen.
Step 1
Calcium is magnesium's groupmate, so it reacts in the same pattern.
Step 2
Swap magnesium's name for calcium's in the product.
Step 3
Calcium burns in oxygen to form calcium oxide.
Worked example 2. Magnesium reacts with chlorine to form magnesium chloride. Fluorine is chlorine's groupmate. Predict the product of magnesium reacting with fluorine.
Step 1
Fluorine sits in chlorine's group, so it reacts with the metal in the same pattern.
Step 2
Swap chlorine's part of the product name for fluorine's.
Step 3
Magnesium reacts with fluorine to form magnesium fluoride.
You can now predict the product pattern of a reaction between elements by analogy with a familiar reaction of elements from the same groups.
Check your understanding
Lithium reacts with iodine to form lithium iodide. Cesium is in lithium's group. What does cesium form with iodine?
ACesium iodidecorrect
BCesium bromide
This option is wrong — you swapped the nonmetal partner too — only the metal changed, so the product keeps iodine's part: cesium iodide.
CLithium cesium
This option is wrong — you combined the two metals — the analogy replaces lithium WITH cesium in the reaction, pairing cesium with the iodine.
DNo product — cesium does not react with iodine.
This option is wrong — you dropped the family pattern — cesium is lithium's groupmate, so it reacts with iodine in the same salt-forming pattern.
Cesium sits in lithium's group, so it reacts in the same pattern. Swap lithium's name for cesium's in the product. Cesium reacts with iodine to form cesium iodide.
Check your understanding
Strontium reacts with iodine to form strontium iodide. Barium is in strontium's group. What does barium form with iodine?
ABarium iodidecorrect
BBarium chloride
This option is wrong — you changed the halogen as well — iodine stayed in the reaction, so the product is barium iodide.
CStrontium barium
This option is wrong — you joined the two metals — barium REPLACES strontium in the reaction; the product pairs barium with iodine.
DThe pattern cannot say — barium's reactions are unrelated to strontium's.
This option is wrong — you cut the group tie — groupmates share a valence-electron count, so barium follows strontium's pattern with iodine.
Barium is strontium's groupmate. Same group, same valence electrons, same reaction pattern. Barium reacts with iodine to form barium iodide.
Check your understanding
Lithium reacts with sulfur to form lithium sulfide. Rubidium is in lithium's group. What forms when rubidium reacts with sulfur?
ARubidium sulfidecorrect
BLithium sulfide
This option is wrong — you kept the old metal's name in the product — rubidium replaced lithium, so the product carries rubidium's name.
CRubidium lithium
This option is wrong — you paired the two groupmates together — the analogy swaps rubidium INTO lithium's place, keeping sulfur as the partner.
DNo product — rubidium is too far down its group to follow lithium's pattern.
This option is wrong — you let distance down the column break the pattern — groupmates share the pattern top to bottom; only the product's name changes.
Rubidium sits in lithium's group, so it reacts with sulfur in the same pattern. Swap lithium's part of the product name for rubidium's. Rubidium sulfide forms.
Lesson 29 of 50 · PTB-029
Predict an element's behavior from its family
Learning Journey
Read · Check · Unlock
Did You Know?
Rubidium is an element most people never meet. Yet a chemist who has never touched it can tell you how it behaves — because the periodic table says which family it belongs to.
The idea
To predict an unfamiliar element's behavior, first find its group on the periodic table.
The group names the chemical family, and the family's shared valence-electron count sets the behavior.
Then apply the family profile to the new element.
Rubidium sits in Group 1, so it is an alkali metal with one valence electron — predict a soft metal that reacts vigorously with water and air and is never found uncombined in nature.
Radon sits in Group 18, so it is a noble gas with a full outer shell — predict a colorless gas that almost never reacts.
Group to family, family to behavior: that chain is the prediction.
Worked examples
Worked example 1. Strontium sits in Group 2. Predict its behavior.
Step 1
Group 2 makes strontium an alkaline earth metal, its atoms holding two valence electrons.
Step 2
Apply the family profile.
Step 3
Strontium is predicted to be a metal harder than the Group 1 metals that reacts with water and air, but less vigorously than an alkali metal.
Worked example 2. Astatine, at the bottom of Group 17, is so rare that its chemistry has barely been observed. Predict its reaction with metals.
Step 1
Group 17 makes astatine a halogen, its atoms one electron short of a full outer shell.
Step 2
Apply the family profile.
Step 3
Astatine is predicted to react with metals to form salts, as its halogen family does.
You can now predict the properties of an unfamiliar element from its chemical family and the family's valence-electron pattern.
Check your understanding
Barium sits in Group 2. Predict its behavior.
AA metal that reacts with water and air, but less vigorously than a Group 1 metal.correct
BA soft metal that reacts violently with water and is never found uncombined.
This option is wrong — you applied the Group 1 profile — Group 2's family is harder and reacts less vigorously.
CA colorless gas that almost never reacts.
This option is wrong — you applied the Group 18 profile — barium's group holds reactive metals, not unreactive gases.
DA very reactive nonmetal that forms salts with metals.
This option is wrong — you applied the Group 17 profile — barium IS a metal; halogens are the salt-forming nonmetals.
Group 2 names the family: the alkaline earth metals. The family profile: metals, harder than Group 1, reacting with water and air less vigorously. That profile is the prediction for barium.
Check your understanding
Xenon sits in Group 18. Predict its behavior.
AA colorless gas that takes part in almost no reactions.correct
BA very reactive nonmetal that travels as two-atom molecules.
This option is wrong — you applied the halogen profile — that is Group 17; Group 18's atoms have full outer shells and almost never react.
CA soft, shiny metal that must be stored under oil.
This option is wrong — you applied the alkali-metal profile — Group 18 holds unreactive gases, not reactive metals.
DA hard, dense metal whose compounds are colorful.
This option is wrong — you applied the transition-metal profile — xenon's group is the far-right gas family, which forms no everyday compounds at all.
Group 18 names the family: the noble gases. Full outer shells give the family almost no tendency to react. Predict a colorless, unreactive gas.
Check your understanding
Francium sits at the bottom of Group 1, and no classroom has ever kept a usable sample. Which prediction about francium is justified?
AA soft metal that would react vigorously with water and would never be found uncombined in nature.correct
BA hard metal that would resist reacting with water.
This option is wrong — you hardened and calmed a Group 1 element — the alkali-metal profile predicts a soft, vigorously water-reactive metal.
CNo prediction is justified for an element that no chemist has ever kept a sample of.
This option is wrong — you demanded a sample — the group-to-family chain is exactly how chemists predict unhandled elements.
DA reactive nonmetal that would form salts with metals.
This option is wrong — you applied the halogen profile — francium's column is Group 1, the alkali metals.
Group 1, below hydrogen, means alkali metal. The family profile: soft, vigorously reactive with water and air, never found uncombined. The chain from group to family to behavior needs no sample.
Summary video — Element classes, electron patterns, and chemical families
A chlorine atom has 17 protons and 17 electrons; a sodium atom has only 11 of each. More particles, bigger atom — surely? Measure them, and the sodium atom is nearly twice as wide.
An atom's size is measured by its 'atomic radius' — the distance from the center of the nucleus out to the atom's outer electrons. Both atoms sit in Period 3, and their positions in the row decide the sizes.
The idea
Move left to right across a period, and each element's atoms hold one more proton than the last — the nuclear charge increases.
Across that same period, the outer electrons stay in the same shell; no new shell is added until the next period.
Atomic radius decreases from left to right across a period because across a period the nuclear charge increases while the outer electrons stay in the same shell, so the outer electrons are pulled in more strongly.
Electrons pulled in more strongly sit closer to the nucleus, and the atom is smaller for it.
That is why sodium, near the left of Period 3, has larger atoms than chlorine, near the right — despite chlorine's greater particle count.
Worked examples
Worked example 1. Magnesium and sulfur are both in Period 3, with sulfur farther right. Which element has the smaller atoms, and why?
Step 1
Across a period the nuclear charge increases while the outer electrons stay in the same shell, so the outer electrons are pulled in more strongly.
Step 2
Sulfur sits farther right, so its atoms' outer electrons are pulled in more strongly than magnesium's.
Step 3
Sulfur has the smaller atoms — its greater nuclear charge pulls the same-shell outer electrons in more strongly.
Worked example 2. Lithium and oxygen sit in Period 2, with oxygen farther right. Compare their atomic radii.
Step 1
Both elements' outer electrons occupy the same shell, the second.
Step 2
Oxygen's atoms carry the greater nuclear charge, so their outer electrons are pulled in more strongly.
Step 3
Oxygen's atomic radius is smaller than lithium's — atomic radius decreases from left to right across the period.
You can now explain that atomic radius decreases from left to right across a period because across a period the nuclear charge increases while the outer electrons stay in the same shell, so the outer electrons are pulled in more strongly.
Check your understanding
Why does atomic radius decrease from left to right across a period?
AAcross a period the nuclear charge increases while the outer electrons stay in the same shell, so the outer electrons are pulled in more strongly.correct
BEach element across a period adds a whole new electron shell, and the added shells squeeze the inner shells into a smaller space.
This option is wrong — you added shells across the row — a new shell comes only with a new period; across a period the outer electrons stay in the same shell.
CAtoms lose electrons one by one as you move right, so less remains.
This option is wrong — you removed electrons — each element to the right has MORE electrons, but they join the same shell while the growing nuclear charge pulls them in.
DAtoms to the right are heavier, and heavier atoms pack themselves smaller.
This option is wrong — you made mass do the shrinking — weight does not pull electrons in; the increasing nuclear charge acting on the same shell does.
Across a period, each element adds a proton: the nuclear charge increases. The outer electrons stay in the same shell the whole way across. A stronger pull on the same shell draws the outer electrons closer — the atoms shrink.
Check your understanding
Potassium and bromine are both in Period 4, with bromine far to the right. Which atom has the smaller atomic radius?
AThe bromine atom.correct
BThe potassium atom.
This option is wrong — you shrank the leftward atom — atomic radius decreases toward the RIGHT, where the greater nuclear charge pulls the same-shell electrons in harder.
CThey are the same size, because both atoms' outer electrons share one shell.
This option is wrong — you turned same-shell into same-size — the shell is shared, but bromine's greater nuclear charge pulls that shell in more strongly.
DIt cannot be decided from their positions.
This option is wrong — you gave up on the trend — same-period positions decide it: farther right means greater nuclear charge and a smaller atom.
Both atoms' outer electrons sit in the fourth shell. Bromine, far to the right, has the greater nuclear charge, so its outer electrons are pulled in more strongly. The bromine atom is the smaller one.
Check your understanding
Beryllium's atoms are larger than fluorine's, though the two elements sit in the same period. What accounts for the difference?
AFluorine's greater nuclear charge pulls its same-shell outer electrons in more strongly.correct
BFluorine's outer electrons occupy a shell closer to the nucleus than beryllium's.
This option is wrong — you moved fluorine's electrons to an inner shell — both elements' outer electrons occupy the second shell; the stronger pull, not a different shell, makes fluorine smaller.
CBeryllium's atoms hold more electrons, and more electrons take more room.
This option is wrong — you gave beryllium the bigger electron count — fluorine holds more electrons, yet its atoms are smaller; the nuclear charge's pull decides, not the count.
DBeryllium's atoms are larger because metals are always larger than nonmetals.
This option is wrong — you used class as the cause — the size follows position in the period: increasing nuclear charge on the same shell, pulled in more strongly toward the right.
Across a period the nuclear charge increases while the outer electrons stay in the same shell. The outer electrons are pulled in more strongly toward the right of the row. Fluorine, far to the right, has the smaller atoms.
Lesson 31 of 50 · PTB-031
Atomic radius down a group
Learning Journey
Read · Check · Unlock
Did You Know?
Across a row, you've already seen atoms shrink. Now turn the corner and move down a column — sodium down to potassium — and the size question has a different answer, for a different reason.
The idea
Move down a group, and each new row starts a new electron shell — a potassium atom has four occupied shells to sodium's three.
Atomic radius increases from top to bottom down a group because down a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly.
Outer electrons sitting farther out make a larger atom.
That is why a potassium atom is larger than a sodium atom.
Worked examples
Worked example 1. Chlorine and iodine are in the same group, iodine two rows lower. Which has the larger atoms, and why?
Step 1
Down a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly.
Step 2
Iodine sits lower, so its atoms' outer electrons occupy a farther shell.
Step 3
Iodine has the larger atoms — its outer electrons sit in a farther shell, added row by row down the group.
Worked example 2. Magnesium and barium share a group, barium far below. Compare their atomic radii.
Step 1
Each row down the group adds an electron shell.
Step 2
Barium's outer electrons therefore sit several shells farther from the nucleus than magnesium's.
Step 3
Barium's atomic radius is larger — atomic radius increases from top to bottom down the group.
You can now explain that atomic radius increases from top to bottom down a group because down a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly.
Check your understanding
Why does atomic radius increase from top to bottom down a group?
ADown a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly.correct
BAtoms lower in a group hold more valence electrons than atoms above them, and the bigger crowd spreads outward.
This option is wrong — you grew the valence count down the column — groupmates share the same valence count; what grows is the number of shells.
CHeavier atoms swell because their nuclei push electrons away.
This option is wrong — you made the nucleus push — a nucleus attracts electrons; the atoms grow because each row's outer electrons occupy a farther, newly added shell.
DThe nuclear charge shrinks down a group, loosening its grip.
This option is wrong — you shrank the nuclear charge — the proton count grows downward too; the size comes from the added shells placing outer electrons farther out.
Each row down a group adds an occupied electron shell. The outer electrons therefore sit farther from the nucleus and are held more weakly. Farther-out outer electrons make a larger atom.
Check your understanding
Oxygen and sulfur are in the same group, with sulfur one row lower. Which atom has the larger atomic radius?
AThe sulfur atom.correct
BThe oxygen atom.
This option is wrong — you grew the atom upward — atomic radius increases DOWNWARD, where each row adds a shell and the outer electrons sit farther out.
CThey are the same size, because both atoms have the same number of valence electrons.
This option is wrong — you turned same-valence into same-size — the valence count matches, but sulfur's outer electrons occupy a farther shell.
DIt cannot be decided from their positions.
This option is wrong — you gave up on the trend — same-group positions decide it: the lower element's outer electrons sit in a farther, added shell.
Sulfur sits one row below oxygen, so its atoms hold one more occupied shell. Outer electrons in a farther shell sit farther from the nucleus. The sulfur atom is the larger one.
Check your understanding
A carbon atom is smaller than a silicon atom, though both elements share a group. What accounts for the difference?
ASilicon's outer electrons occupy an added shell, farther from the nucleus and held more weakly.correct
BSilicon's atoms hold more valence electrons than carbon's.
This option is wrong — you grew the valence count — carbon and silicon share four valence electrons each; silicon's four simply sit in a farther, added shell.
CCarbon's nucleus pulls electrons harder than silicon's does, because carbon is a nonmetal.
This option is wrong — you used element class as the pull — the size difference comes from silicon's extra occupied shell, not from metal-versus-nonmetal character.
DSilicon's atoms are larger only when silicon is heated.
This option is wrong — you made the size conditional on temperature — the added shell makes silicon's atoms larger at any temperature.
Down a group each row adds an electron shell. Silicon sits one row below carbon, so its outer electrons occupy a farther shell, held more weakly. That farther shell is the larger radius.
Lesson 32 of 50 · PTB-032
Rank elements by atomic radius
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen both size trends one pair at a time. Ranking three or more elements needs nothing new — just the two directions, applied to positions read off the table.
The idea
First read the elements' positions: are they in the same period, or the same group?
In the same period, atomic radius decreases from left to right — the leftmost element has the largest atoms.
In the same group, atomic radius increases from top to bottom — the lowest element has the largest atoms.
Lithium, nitrogen, and fluorine sit left to right in Period 2, so the order from smallest to largest atomic radius is fluorine, nitrogen, lithium.
Beryllium, calcium, and barium sit top to bottom in Group 2, so the order from smallest to largest atomic radius is beryllium, calcium, barium.
Worked examples
Worked example 1. Sodium, magnesium, and chlorine sit in Period 3, in that left-to-right order. Rank them from smallest to largest atomic radius.
Step 1
Same period: atomic radius decreases from left to right.
Step 2
Chlorine sits farthest right, so its atoms are smallest; sodium sits farthest left, so its atoms are largest.
Step 3
Smallest to largest: chlorine, magnesium, sodium.
Worked example 2. Chlorine, bromine, and iodine sit in Group 17, in that top-to-bottom order. Rank them from smallest to largest atomic radius.
Step 1
Same group: atomic radius increases from top to bottom.
Step 2
Chlorine sits highest, so its atoms are smallest; iodine sits lowest, so its atoms are largest.
Step 3
Smallest to largest: chlorine, bromine, iodine.
You can now rank elements from the same period or the same group by atomic radius using their positions in the periodic table.
Check your understanding
The periodic-table snippet highlights oxygen, sulfur, and selenium. Rank them from smallest to largest atomic radius.
ASmallest to largest: oxygen, sulfur, selenium.correct
BSmallest to largest: selenium, sulfur, oxygen.
This option is wrong — you ran the group trend upward — each row DOWN the group adds a shell, so the lowest element, selenium, has the largest atoms.
CAll three have the same atomic radius, because they share a group.
This option is wrong — you turned shared valence electrons into shared size — groupmates behave alike, but each row's added shell grows the radius downward.
DThey cannot be ranked from their positions alone.
This option is wrong — you asked for measured values — same-group positions rank the radii by themselves: size grows down the column.
The snippet shows the three elements in one column — the same group. In the same group, atomic radius increases from top to bottom. Oxygen at the top is smallest; selenium at the bottom is largest.
Check your understanding
The periodic-table snippet highlights boron, carbon, and oxygen. Rank them from smallest to largest atomic radius.
ASmallest to largest: oxygen, carbon, boron.correct
BSmallest to largest: boron, carbon, oxygen.
This option is wrong — you let the radius grow to the right — across a period the growing nuclear charge pulls the same shell in, so the rightmost element, oxygen, is smallest.
CAll three have the same atomic radius, because their outer electrons share one shell.
This option is wrong — you turned same-shell into same-size — the shell is shared, but the nuclear charge grows to the right and pulls it in more strongly.
DThey cannot be ranked from their positions alone.
This option is wrong — you asked for measured values — same-period positions rank the radii by themselves: size shrinks toward the right.
The snippet shows the three elements in one row — the same period. In the same period, atomic radius decreases from left to right. Oxygen, farthest right, is smallest; boron, farthest left, is largest.
Check your understanding
The periodic-table snippet highlights aluminum, phosphorus, and sulfur. Which element has the largest atomic radius?
AAluminumcorrect
BSulfur
This option is wrong — you grew the atoms toward the right — across a period the radius shrinks that way; the leftmost element, aluminum, has the largest atoms.
CPhosphorus
This option is wrong — you picked the middle of the row — in a same-period set the LEFTMOST element has the largest atoms, and aluminum sits farthest left.
DAll three are the same size.
This option is wrong — you leveled the row — the nuclear charge grows left to right and pulls the shared shell in, so sizes differ across the period.
The snippet shows the three elements in one row — the same period. Atomic radius decreases from left to right across a period. Aluminum, farthest left, has the largest atoms.
Lesson 33 of 50 · PTB-033
Ionization energy
Learning Journey
Read · Check · Unlock
Did You Know?
Wonder this:
An electron does not simply fall off an atom. The nucleus's positive charge holds it, and prying one loose means fighting that attraction all the way out.
You've already seen that an atom which loses an electron becomes an ion. The prying-loose itself has a price, and the price has a name.
The idea
Removing one electron from an atom takes energy, because the nucleus's attraction must be overcome.
removing one electron takes energy
The energy needed to remove one electron from an atom is called the atom's 'ionization energy'.
The name fits: removing the electron turns the atom into an ion.
Every element's atoms have their own ionization energy.
Helium's is the highest of all — removing one electron from a helium atom takes more energy than removing one from any other kind of atom.
Worked examples
Worked example 1. What is ionization energy?
Step 1
Answer: the energy needed to remove one electron from an atom.
Worked example 2. Removing one electron from a magnesium atom takes a certain amount of energy. What is that quantity called?
Step 1
Answer: magnesium's ionization energy.
You can now state that ionization energy is the energy needed to remove one electron from an atom.
Check your understanding
What is an atom's ionization energy?
AThe energy needed to remove one electron from the atom.correct
BThe energy needed to add one electron to the atom.
This option is wrong — you ran the change in reverse — ionization energy is the price of REMOVING an electron, not of adding one.
CThe energy holding the protons together inside the nucleus.
This option is wrong — you moved the energy into the nucleus — ionization energy concerns an outer electron's removal, and the nucleus stays untouched.
DThe energy released when the atom becomes an ion.
This option is wrong — you paid the energy in the wrong direction — removing the electron takes energy IN; nothing is released by the removal.
The nucleus attracts every electron, so removal has a price. The energy needed to remove one electron from an atom is the ionization energy. Removing the electron turns the atom into an ion — the name's origin.
Check your understanding
Removing one electron from an aluminum atom requires a specific amount of energy. What is that amount called?
AAluminum's ionization energycorrect
BAluminum's atomic radius
This option is wrong — you named the size quantity — atomic radius is the atom's size; the removal price is the ionization energy.
CAluminum's nuclear charge
This option is wrong — you named the cause instead of the quantity — the nuclear charge is what makes removal costly; the energy itself is the ionization energy.
DAluminum's atomic mass
This option is wrong — you reached for the mass — atomic mass weighs the atom; the energy to remove one electron is the ionization energy.
The energy needed to remove one electron from an atom is its ionization energy. For an aluminum atom, that amount is aluminum's ionization energy.
Check your understanding
Removing one electron takes more energy for which kind of atom than for any other?
AA helium atomcorrect
BA hydrogen atom
This option is wrong — you picked the smallest element — small, but not the record holder; helium's atoms hold the highest ionization energy of all.
CA cesium atom
This option is wrong — you picked a large metal atom — a cesium atom's outer electron is famously EASY to remove, the opposite of the record.
DAn iron atom
This option is wrong — you picked a hard, tough metal — an object's hardness is not its atoms' grip on electrons; helium holds the record.
Ionization energy is the energy needed to remove one electron. Helium's is the highest of any element. No atom holds its electron more expensively than helium's.
Lesson 34 of 50 · PTB-034
Ionization energy across a period
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen why atoms shrink across a period: the nuclear charge grows while the outer electrons stay in the same shell. The same cause sets the price of removing an electron.
The idea
Move left to right across a period, and the nuclear charge increases while the outer electrons stay in the same shell.
Ionization energy generally increases from left to right across a period because across a period the nuclear charge increases while the outer electrons stay in the same shell, so the outer electrons are pulled in more strongly and take more energy to remove.
Sodium, at the left of Period 3, has a low ionization energy; chlorine, near the right, has a much higher one.
Worked examples
Worked example 1. Lithium and fluorine sit in Period 2, fluorine far to the right. Which element's ionization energy is higher, and why?
Step 1
Across a period the nuclear charge increases while the outer electrons stay in the same shell, so the outer electrons are pulled in more strongly.
Step 2
Fluorine sits far right, so its outer electrons are pulled in more strongly and take more energy to remove.
Step 3
Fluorine's ionization energy is higher — its outer electrons, pulled in more strongly by the greater nuclear charge, cost more energy to remove.
Worked example 2. Silicon and chlorine sit in Period 3, chlorine farther right. Compare their ionization energies.
Step 1
Both elements' outer electrons occupy the third shell.
Step 2
Chlorine's greater nuclear charge pulls those electrons in more strongly, so they take more energy to remove.
Step 3
Chlorine's ionization energy is generally higher than silicon's — the period trend.
You can now explain that ionization energy generally increases from left to right across a period because across a period the nuclear charge increases while the outer electrons stay in the same shell, so the outer electrons are pulled in more strongly and take more energy to remove.
Check your understanding
Why does ionization energy generally increase from left to right across a period?
AThe nuclear charge increases while the outer electrons stay in the same shell, so they are pulled in more strongly and take more energy to remove.correct
BAtoms to the right of a period hold their outer electrons in farther shells, and electrons in farther shells are harder to remove.
This option is wrong — you moved the electrons outward — across a period the outer electrons stay in the SAME shell, and a farther shell would make removal easier, not harder.
CAtoms to the right have fewer electrons, so each one is guarded more closely.
This option is wrong — you shrank the electron count — elements to the right have MORE electrons; the stronger pull comes from the growing nuclear charge on the same shell.
DAtoms to the right are larger, and large atoms grip their electrons more tightly.
This option is wrong — you grew the atoms rightward — they shrink that way, and the shrinking and the higher removal price share one cause: the stronger pull of the greater nuclear charge.
Across a period the nuclear charge increases while the outer electrons stay in the same shell. The outer electrons are pulled in more strongly. A more strongly held electron takes more energy to remove — so ionization energy generally rises across the row.
Check your understanding
Magnesium and sulfur sit in Period 3, sulfur farther right. Which element's ionization energy is generally higher?
ASulfur'scorrect
BMagnesium's
This option is wrong — you ran the trend leftward — the greater nuclear charge sits to the right, holding sulfur's same-shell electrons more strongly.
CThey are equal, because both elements' outer electrons occupy the same shell.
This option is wrong — you turned same-shell into same-price — the shell matches, but sulfur's greater nuclear charge pulls it in more strongly.
DIt cannot be judged from their positions.
This option is wrong — you gave up on the trend — same-period positions decide it: farther right generally means a higher ionization energy.
Both elements keep their outer electrons in the third shell. Sulfur, farther right, has the greater nuclear charge, so its outer electrons are pulled in more strongly. More strongly held electrons take more energy to remove — sulfur's ionization energy is generally higher.
Check your understanding
Removing an electron from an oxygen atom takes more energy than removing one from a carbon atom, though both elements sit in Period 2. What accounts for the difference?
AOxygen's greater nuclear charge pulls its same-shell outer electrons in more strongly, so removal costs more.correct
BOxygen's outer electrons sit closer to the nucleus because oxygen has fewer shells than carbon.
This option is wrong — you dropped a shell from oxygen — both elements' outer electrons occupy the second shell; the stronger pull of oxygen's greater nuclear charge is the cause.
COxygen's atoms hold more electrons, and a bigger crowd is harder to pull from.
This option is wrong — you priced the removal by crowd size — the count does not set the grip; the nuclear charge pulling on the same shell does.
DCarbon gives electrons away because it is a metal.
This option is wrong — you reclassified carbon — it is a nonmetal, and class is beside the point; position in the period sets the pull on the outer electrons.
Across a period the nuclear charge increases while the outer electrons stay in the same shell. Oxygen, farther right than carbon, pulls its outer electrons in more strongly. The stronger hold makes the removal cost more energy.
Lesson 35 of 50 · PTB-035
Ionization energy down a group
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen that ionization energy is the energy needed to remove one electron from an atom, and why it generally climbs from left to right across a period. Now compare two atoms in the same column: sodium and potassium both sit in Group 1, but potassium sits one row lower.
The idea
Each new row down a group adds one more electron shell, so potassium's outer electron sits in a shell farther from the nucleus than sodium's.
An outer electron that sits farther from the nucleus is held more weakly.
A weakly held electron takes less energy to remove, so potassium's ionization energy is lower than sodium's.
Ionization energy decreases from top to bottom down a group because down a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly and take less energy to remove.
down Group 1: ionization energy decreases
Worked examples
Worked example 1. Chlorine and iodine both sit in Group 17. Why does iodine have a lower ionization energy than chlorine?
Step 1
Iodine sits three rows below chlorine, so iodine's outer electrons sit in a shell much farther from the nucleus.
Step 2
Outer electrons that sit farther from the nucleus are held more weakly.
Step 3
So removing one electron from an iodine atom takes less energy — iodine's ionization energy is lower, because down a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly.
You can now explain that ionization energy decreases from top to bottom down a group because down a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly and take less energy to remove.
Check your understanding
Lithium and rubidium both sit in Group 1. From which atom does it take less energy to remove one electron?
ARubidiumcorrect
BLithium
This option is wrong — you inverted the group direction — down a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly, and rubidium sits lower.
CBoth take the same energy, because each atom has one outer electron
This option is wrong — you counted outer electrons instead of shells — the outer electron's distance from the nucleus, not its count, sets how tightly it is held.
DIt cannot be predicted from their positions
This option is wrong — you gave up on the positions — same-group positions decide this: the lower element's outer electron sits farther out and is held more weakly.
Rubidium sits three rows below lithium in Group 1. Down a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly. A weakly held electron takes less energy to remove, so rubidium's ionization energy is lower.
Check your understanding
Magnesium and barium both sit in Group 2, and barium's ionization energy is lower. Why?
ABecause down a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly and take less energy to remove.correct
BBecause barium has more protons, which push the outer electrons away from the atom.
This option is wrong — you made protons repel electrons — protons attract electrons; what lowers barium's ionization energy is its outer electrons sitting in a shell farther from the nucleus.
CBecause across a period the nuclear charge increases while the outer electrons stay in the same shell, so barium's outer electrons are pulled in more strongly than magnesium's.
This option is wrong — you applied the across-a-period mechanism — magnesium and barium are in the same group, and down a group each row adds a shell.
DBecause barium has fewer outer electrons to remove than magnesium.
This option is wrong — you miscounted — both Group 2 elements have two outer electrons; the difference is how far from the nucleus those electrons sit.
Barium sits four rows below magnesium in Group 2. Down a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly. So barium's outer electrons take less energy to remove.
Check your understanding
Sulfur and tellurium both sit in Group 16. From which atom is the outermost electron harder to remove?
ASulfurcorrect
BTellurium
This option is wrong — you inverted the group direction — tellurium sits lower, so its outer electrons sit farther from the nucleus and are held more weakly, not more strongly.
CBoth are equally hard, because both atoms have six outer electrons
This option is wrong — you counted outer electrons instead of shells — the added shells down the group, not the electron count, decide how tightly the outer electrons are held.
DNeither — electrons cannot be removed from nonmetal atoms
This option is wrong — you treated ionization as metals-only — removing an electron takes energy for any atom, and that energy is its ionization energy.
Sulfur sits above tellurium in Group 16. Down a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly. Sulfur's outer electrons sit closer in and are held more tightly, so they are harder to remove.
Lesson 36 of 50 · PTB-036
Rank elements by ionization energy
Learning Journey
Read · Check · Unlock
Did You Know?
You've seen the two directions: ionization energy generally increases from left to right across a period, and decreases from top to bottom down a group. Ranking several elements is just those two directions read off a periodic table.
The idea
To rank elements by ionization energy, first find their positions: are they in the same period or the same group?
down Group 1: ionization energy decreases
In the same period, ionization energy generally increases from left to right — the farther right, the higher.
In the same group, ionization energy decreases from top to bottom — the farther down, the lower.
Sodium, aluminum, and chlorine share Period 3, and chlorine sits farthest right — so from lowest to highest the order is sodium, aluminum, chlorine.
Lithium, potassium, and cesium share Group 1, and cesium sits lowest — so from lowest to highest the order is cesium, potassium, lithium.
Worked examples
Worked example 1. Lithium, nitrogen, and fluorine all sit in Period 2. Rank them from lowest to highest ionization energy.
Step 1
All three share Period 2, so use the period direction: ionization energy generally increases from left to right.
Step 2
Lithium sits farthest left, then nitrogen, then fluorine farthest right.
Step 3
From lowest to highest: lithium, nitrogen, fluorine.
Worked example 2. Chlorine, bromine, and iodine all sit in Group 17. Rank them from lowest to highest ionization energy.
Step 1
All three share Group 17, so use the group direction: ionization energy decreases from top to bottom.
Step 2
Iodine sits lowest, then bromine, then chlorine at the top of the three.
Step 3
From lowest to highest: iodine, bromine, chlorine.
You can now rank elements from the same period or the same group by ionization energy using their positions in the periodic table.
Check your understanding
The figure highlights magnesium, sulfur, and argon in Period 3. Rank them from lowest to highest ionization energy.
AMagnesium, sulfur, argoncorrect
BArgon, sulfur, magnesium
This option is wrong — you ran the period direction backwards — ionization energy generally increases from left to right, so the leftmost element is lowest.
CSulfur, magnesium, argon
This option is wrong — you misread the positions — magnesium sits farthest left of the three, so it is lowest.
DThey cannot be ranked from their positions alone
This option is wrong — you gave up on the table — same-period positions rank these three: ionization energy generally increases from left to right.
All three sit in Period 3, so use the period direction. Ionization energy generally increases from left to right across a period. Magnesium is farthest left, then sulfur, then argon — lowest to highest: magnesium, sulfur, argon.
Check your understanding
The figure highlights beryllium, calcium, and barium in Group 2. Rank them from lowest to highest ionization energy.
ABarium, calcium, berylliumcorrect
BBeryllium, calcium, barium
This option is wrong — you ran the group direction backwards — ionization energy decreases from top to bottom, so the lowest element in the column has the lowest value.
CCalcium, barium, beryllium
This option is wrong — you misread the positions — barium sits lowest in the column, so it comes first in a lowest-to-highest ranking.
DThey are equal, because all three have two outer electrons
This option is wrong — you counted outer electrons instead of shells — down the group the outer electrons sit farther from the nucleus and take less energy to remove.
All three sit in Group 2, so use the group direction. Ionization energy decreases from top to bottom down a group. Barium sits lowest, then calcium, then beryllium — lowest to highest: barium, calcium, beryllium.
Check your understanding
The figure highlights oxygen, selenium, and tellurium in Group 16. Which of the three has the highest ionization energy?
AOxygencorrect
BTellurium
This option is wrong — you inverted the group direction — ionization energy decreases from top to bottom, so the top element of the column has the highest value.
CSelenium
This option is wrong — you picked the middle element — in a same-group set the extremes sit at the top and bottom, and the top has the highest ionization energy.
DAll three are the same
This option is wrong — you treated same-group elements as identical — their outer electrons sit in different shells, so the energies differ down the column.
All three sit in Group 16, so use the group direction. Ionization energy decreases from top to bottom down a group. Oxygen sits at the top of the three, so oxygen's ionization energy is the highest.
Lesson 37 of 50 · PTB-037
Electronegativity
Learning Journey
Read · Check · Unlock
Have You Ever Wondered?
Wonder this:
In a water molecule, an oxygen atom and two hydrogen atoms share electrons. Sharing sounds fair — but is it? Does each atom pull equally hard on the electrons they share?
The sharing is not fair. Some kinds of atoms pull harder on shared electrons than others, and chemists needed a way to say how hard each kind pulls.
The idea
Each kind of atom pulls on shared electrons with its own characteristic strength.
That pulling strength is called 'electronegativity' — a measure of how strongly an atom pulls on electrons it shares with another atom.
Every element is assigned an electronegativity number; the whole scale only runs from about 0.7 to 4, and a bigger number means a stronger pull.
Fluorine has the highest electronegativity of any element, 3.98 — fluorine pulls shared electrons more strongly than any other atom.
In the water molecule, oxygen's 3.44 beats hydrogen's 2.20, so the shared electrons are pulled closer to the oxygen atom.
Worked examples
Worked example 1. Nitrogen's electronegativity is 3.04 and carbon's is 2.55. When a carbon atom and a nitrogen atom share electrons, which atom pulls the shared electrons more strongly?
Step 1
The bigger electronegativity number means the stronger pull on the shared electrons.
Step 2
Nitrogen — 3.04 beats 2.55, so the nitrogen atom pulls the shared electrons more strongly.
You can now state that electronegativity is a measure of how strongly an atom pulls on electrons it shares with another atom.
Check your understanding
What does an element's electronegativity measure?
AHow strongly its atoms pull on electrons shared with another atomcorrect
BHow much energy it takes to remove an electron from its atoms
This option is wrong — you described ionization energy — electronegativity is about pulling on shared electrons, not about removing an atom's own electron.
CHow many electrons its atoms hold in the outer shell
This option is wrong — you described the valence-electron count — electronegativity is a pulling strength, not an electron count.
DHow strongly its atoms attract other whole atoms
This option is wrong — you moved the pull onto whole atoms — the pull that electronegativity measures acts on the shared electrons between atoms.
Atoms in a molecule share electrons, and some atoms pull harder on those shared electrons than others. Electronegativity is the measure of how strongly an atom pulls on electrons it shares with another atom.
Check your understanding
Phosphorus has an electronegativity of 2.19 and sulfur has 2.58. When a phosphorus atom and a sulfur atom share electrons, which statement is correct?
AThe sulfur atom pulls the shared electrons more strongly.correct
BThe phosphorus atom pulls the shared electrons more strongly.
This option is wrong — you picked the smaller number — a bigger electronegativity number means a stronger pull, and 2.58 is bigger than 2.19.
CThey pull equally, because both elements are nonmetals.
This option is wrong — you used the element class instead of the numbers — nonmetals differ in electronegativity, and the numbers say sulfur pulls harder.
DNeither atom pulls on the shared electrons.
This option is wrong — you dropped the pull entirely — every atom pulls on shared electrons; electronegativity says how strongly.
A bigger electronegativity number means a stronger pull on shared electrons. Sulfur's 2.58 is bigger than phosphorus's 2.19. So the sulfur atom pulls the shared electrons more strongly.
Check your understanding
One element pulls shared electrons more strongly than any other element. Which is it?
AFluorinecorrect
BOxygen
This option is wrong — you stopped one place short — oxygen's 3.44 is high, but fluorine's 3.98 is the highest of all.
CCesium
This option is wrong — you picked one of the weakest pullers — cesium's atoms barely pull on shared electrons at all.
DNeon
This option is wrong — you went to the far right of the table — neon atoms barely share electrons with other atoms, so no strong pull on shared electrons is in play.
Fluorine has the highest electronegativity of any element, 3.98. So fluorine pulls shared electrons more strongly than any other atom.
Lesson 38 of 50 · PTB-038
Electronegativity across a period
Learning Journey
Read · Check · Unlock
Did You Know?
Lithium sits at the left end of Period 2 and fluorine near the right end. Lithium's electronegativity is 0.98; fluorine's is 3.98. That gap is not an accident of the two elements — it is a direction that runs across the whole row.
The idea
Moving from lithium to fluorine across Period 2, each element has more protons in its nucleus, while the outer electrons stay in the same shell.
A bigger nuclear charge pulling on outer electrons in the same shell pulls them in more strongly.
A stronger pull on the outer electrons also acts on electrons shared with another atom, so the pull on shared electrons climbs from lithium's 0.98 to fluorine's 3.98.
Electronegativity increases from left to right across a period, leaving out the noble gases, because across a period the nuclear charge increases while the outer electrons stay in the same shell, so the outer electrons are pulled in more strongly — and a stronger pull also acts on shared electrons.
The noble gases are left out because their atoms barely share electrons with other atoms, so the trend skips them.
down Group 1: ionization energy decreases
Worked examples
Worked example 1. Sodium's electronegativity is 0.93 and chlorine's is 3.16. Both sit in Period 3. Why is chlorine's so much higher?
Step 1
Chlorine sits far to the right of sodium in Period 3, so a chlorine atom has many more protons pulling on outer electrons in the same shell.
Step 2
Across a period the nuclear charge increases while the outer electrons stay in the same shell, so the outer electrons are pulled in more strongly — and a stronger pull also acts on shared electrons.
Step 3
So chlorine pulls shared electrons much more strongly, and its electronegativity is far higher.
You can now explain that electronegativity increases from left to right across a period, leaving out the noble gases, because across a period the nuclear charge increases while the outer electrons stay in the same shell, so the outer electrons are pulled in more strongly — and a stronger pull also acts on shared electrons.
Check your understanding
Magnesium and sulfur both sit in Period 3. Which element's atoms pull shared electrons more strongly?
ASulfurcorrect
BMagnesium
This option is wrong — you inverted the period direction — electronegativity increases from left to right, and sulfur sits farther right.
CThey pull equally, because both sit in the same period
This option is wrong — you treated a shared period as a shared value — the pull climbs across the row as the nuclear charge grows.
DIt cannot be judged from their positions
This option is wrong — you gave up on the positions — same-period positions decide this: the farther right, the stronger the pull on shared electrons.
Sulfur sits well to the right of magnesium in Period 3. Across a period the nuclear charge increases while the outer electrons stay in the same shell, so the outer electrons are pulled in more strongly — and a stronger pull also acts on shared electrons. So sulfur pulls shared electrons more strongly.
Check your understanding
Oxygen's electronegativity (3.44) is far higher than beryllium's (1.57); both sit in Period 2. Why?
ABecause across a period the nuclear charge increases while the outer electrons stay in the same shell, so the outer electrons are pulled in more strongly — and a stronger pull also acts on shared electrons.correct
BBecause down a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly — and a weaker hold means a weaker pull on shared electrons.
This option is wrong — you applied the group mechanism — beryllium and oxygen sit in the same period, so the shell count does not change between them.
CBecause oxygen atoms are bigger than beryllium atoms, and bigger atoms pull shared electrons harder.
This option is wrong — you tied the pull to size and got the size backwards — oxygen's atoms are smaller, and the pull comes from the bigger nuclear charge acting on the same shell.
DBecause oxygen has more outer electrons, and each outer electron adds its own pull on the shared pair.
This option is wrong — you made the outer electrons do the pulling — the nucleus does the pulling, and it is the growing nuclear charge that strengthens the pull.
Oxygen sits well to the right of beryllium in Period 2, so its nucleus holds more protons. Across a period the nuclear charge increases while the outer electrons stay in the same shell, so the outer electrons are pulled in more strongly — and a stronger pull also acts on shared electrons. That stronger pull is oxygen's higher electronegativity.
Check your understanding
Which Period 3 element has the highest electronegativity?
AChlorinecorrect
BArgon
This option is wrong — you went all the way to the right — the noble gases are left out of the electronegativity trend because their atoms barely share electrons.
CSodium
This option is wrong — you inverted the period direction — electronegativity increases from left to right, so the far-left element has the weakest pull.
DSilicon
This option is wrong — you stopped in the middle of the row — the pull keeps climbing to the right, up to chlorine.
Electronegativity increases from left to right across a period, leaving out the noble gases. In Period 3 the last element before the noble gas is chlorine. So chlorine has the highest electronegativity in Period 3.
Lesson 39 of 50 · PTB-039
Electronegativity down a group
Learning Journey
Read · Check · Unlock
Did You Know?
Fluorine sits at the top of Group 17 and iodine three rows below it. Fluorine's electronegativity is 3.98; iodine's is only 2.66. The drop follows the column, the same way the size and ionization-energy trends did.
The idea
Each row down from fluorine to iodine adds an electron shell, so iodine's outer electrons sit much farther from the nucleus.
Outer electrons that sit farther from the nucleus are held more weakly.
A weaker hold on its own outer electrons means a weaker pull on shared electrons too, so iodine's 2.66 sits well below fluorine's 3.98.
Electronegativity decreases from top to bottom down a group because down a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly — and a weaker hold means a weaker pull on shared electrons.
down Group 1: ionization energy decreases
Worked examples
Worked example 1. Oxygen's electronegativity is 3.44 and sulfur's is 2.58. Both sit in Group 16. Why is sulfur's lower?
Step 1
Sulfur sits one row below oxygen, so sulfur's outer electrons sit in a shell farther from the nucleus.
Step 2
Down a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly — and a weaker hold means a weaker pull on shared electrons.
Step 3
So a sulfur atom pulls shared electrons more weakly than an oxygen atom, and its electronegativity is lower.
You can now explain that electronegativity decreases from top to bottom down a group because down a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly — and a weaker hold means a weaker pull on shared electrons.
Check your understanding
Nitrogen and phosphorus both sit in Group 15. Which element's atoms pull shared electrons more strongly?
ANitrogencorrect
BPhosphorus
This option is wrong — you inverted the group direction — electronegativity decreases from top to bottom, and phosphorus sits lower.
CThey pull equally, because both have five outer electrons
This option is wrong — you counted outer electrons instead of shells — phosphorus's outer electrons sit a shell farther out, where they are held more weakly.
DIt cannot be judged from their positions
This option is wrong — you gave up on the positions — same-group positions decide this: the higher element pulls shared electrons more strongly.
Nitrogen sits above phosphorus in Group 15. Down a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly — and a weaker hold means a weaker pull on shared electrons. So nitrogen, the higher element, pulls shared electrons more strongly.
Check your understanding
Lithium's electronegativity (0.98) is higher than cesium's (0.79); both sit in Group 1. Why?
ABecause down a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly — and a weaker hold means a weaker pull on shared electrons.correct
BBecause across a period the nuclear charge increases while the outer electrons stay in the same shell, so the outer electrons are pulled in more strongly — and a stronger pull also acts on shared electrons.
This option is wrong — you applied the period mechanism — lithium and cesium sit in the same group, so it is the added shells, not a same-shell charge change, that separates them.
CBecause cesium has more protons, and more protons must mean a stronger pull — the two values must be listed in the wrong order.
This option is wrong — you counted protons only — cesium's outer electron sits many shells farther out, and that distance outweighs the extra protons.
DBecause metals do not pull on shared electrons at all, so their values mean nothing.
This option is wrong — you wrote metals out of the scale — metal atoms pull weakly, but they do pull, and the numbers rank those pulls.
Cesium sits four rows below lithium in Group 1. Down a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly — and a weaker hold means a weaker pull on shared electrons. So lithium, the higher element, keeps the higher electronegativity.
Check your understanding
Carbon sits above silicon in Group 14. Which atom pulls shared electrons more strongly?
ACarboncorrect
BSilicon
This option is wrong — you inverted the group direction — silicon's outer electrons sit a shell farther from the nucleus, so its pull on shared electrons is weaker.
CBoth pull equally, because both have four outer electrons
This option is wrong — you counted outer electrons instead of shells — the distance of the outer shell from the nucleus, not the electron count, sets the pull.
DNeither, because Group 14 elements do not share electrons
This option is wrong — you removed sharing from Group 14 — carbon and silicon atoms share electrons constantly; the question is how hard each pulls.
Carbon sits above silicon in Group 14. Down a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly — and a weaker hold means a weaker pull on shared electrons. So carbon pulls shared electrons more strongly.
Lesson 40 of 50 · PTB-040
Rank elements by electronegativity
Learning Journey
Read · Check · Unlock
Did You Know?
You've seen both electronegativity directions: up across a period, down a group. Ranking several elements by electronegativity is the same position-reading skill you used for ionization energy — with one extra rule about the noble gases.
The idea
To rank elements by electronegativity, first find their positions: same period or same group?
down Group 1: ionization energy decreases
In the same period, electronegativity increases from left to right — leaving out the noble gases.
In the same group, electronegativity decreases from top to bottom.
Bromine, chlorine, and fluorine share Group 17, and fluorine sits highest — so from lowest to highest the order is bromine, chlorine, fluorine.
Sodium, silicon, and sulfur share Period 3, and sulfur sits farthest right — so from lowest to highest the order is sodium, silicon, sulfur.
Worked examples
Worked example 1. Lithium, carbon, and oxygen all sit in Period 2. Rank them from lowest to highest electronegativity.
Step 1
All three share Period 2, so use the period direction: electronegativity increases from left to right, leaving out the noble gases.
Step 2
Lithium sits farthest left, then carbon, then oxygen.
Step 3
From lowest to highest: lithium, carbon, oxygen.
Worked example 2. Cesium, potassium, and sodium all sit in Group 1. Rank them from lowest to highest electronegativity.
Step 1
All three share Group 1, so use the group direction: electronegativity decreases from top to bottom.
Step 2
Cesium sits lowest, then potassium, then sodium at the top of the three.
Step 3
From lowest to highest: cesium, potassium, sodium.
You can now rank elements from the same period or the same group by electronegativity using their positions in the periodic table, leaving out the noble gases.
Check your understanding
The figure highlights magnesium, phosphorus, and chlorine in Period 3. Rank them from lowest to highest electronegativity.
AMagnesium, phosphorus, chlorinecorrect
BChlorine, phosphorus, magnesium
This option is wrong — you ran the period direction backwards — electronegativity increases from left to right, so the leftmost element is lowest.
CPhosphorus, magnesium, chlorine
This option is wrong — you misread the positions — magnesium sits farthest left of the three, so it is lowest.
DThey cannot be ranked from their positions alone
This option is wrong — you gave up on the table — same-period positions rank these three: electronegativity increases from left to right.
All three sit in Period 3, so use the period direction. Electronegativity increases from left to right across a period, leaving out the noble gases. Magnesium is farthest left, then phosphorus, then chlorine — lowest to highest: magnesium, phosphorus, chlorine.
Check your understanding
The figure highlights oxygen, selenium, and tellurium in Group 16. Rank them from lowest to highest electronegativity.
ATellurium, selenium, oxygencorrect
BOxygen, selenium, tellurium
This option is wrong — you ran the group direction backwards — electronegativity decreases from top to bottom, so the lowest element in the column has the lowest value.
CSelenium, tellurium, oxygen
This option is wrong — you misread the positions — tellurium sits lowest in the column, so it comes first in a lowest-to-highest ranking.
DThey are equal, because all three have six outer electrons
This option is wrong — you counted outer electrons instead of shells — down the group the outer electrons sit farther out and pull shared electrons more weakly.
All three sit in Group 16, so use the group direction. Electronegativity decreases from top to bottom down a group. Tellurium sits lowest, then selenium, then oxygen — lowest to highest: tellurium, selenium, oxygen.
Check your understanding
The figure highlights potassium, calcium, and bromine in Period 4. Which of the three pulls shared electrons the most strongly?
ABrominecorrect
BPotassium
This option is wrong — you inverted the period direction — potassium sits farthest left, so its pull on shared electrons is the weakest of the three.
CCalcium
This option is wrong — you stopped near the left of the row — the pull keeps climbing to the right, and bromine sits far to the right.
DThey pull equally, because all three sit in Period 4
This option is wrong — you treated a shared period as a shared value — electronegativity climbs across the row.
All three sit in Period 4, so use the period direction. Electronegativity increases from left to right across a period, leaving out the noble gases. Bromine sits farthest right of the three, so bromine pulls shared electrons the most strongly.
Lesson 41 of 50 · PTB-041
Metallic character
Learning Journey
Read · Check · Unlock
Have You Ever Wondered?
Wonder this:
Hold a copper wire next to a lump of sulfur. Copper is obviously a metal; sulfur obviously is not. But the periodic table is full of in-between cases — so is there a way to say HOW strongly an element behaves like a metal?
There is. Chemists treat metal behavior as something an element can show strongly, weakly, or barely at all — and they tie it to one thing the atoms do.
The idea
Some elements show metal behavior strongly, some weakly, and some barely at all.
How strongly an element shows metal behavior is called its 'metallic character'.
What makes behavior metallic is what the atoms do with their outer electrons: an element is more metallic the more easily its atoms give away their outer electrons.
Cesium's single outer electron is very easy to remove, so cesium is one of the most metallic elements of all.
Worked examples
Worked example 1. Rubidium atoms give up their single outer electron very easily. What does that say about rubidium's metallic character?
Step 1
Atoms that give away their outer electrons easily are the mark of a highly metallic element.
Step 2
Rubidium's metallic character is strong.
You can now state that metallic character means how strongly an element shows metal behavior, and that an element is more metallic the more easily its atoms give away their outer electrons.
Check your understanding
What does an element's metallic character describe?
AHow strongly the element shows metal behaviorcorrect
BHow shiny the element's surface is
This option is wrong — you reduced metal behavior to shine — luster is one metal property, but metallic character is the whole pattern, rooted in how easily the atoms give away outer electrons.
CHow dense the element is
This option is wrong — you tied metal behavior to density — some metals are very light, and density is not what metallic character measures.
DWhether the element is attracted to a magnet
This option is wrong — you tied metal behavior to magnetism — most metals are not magnetic at all.
Metallic character means how strongly an element shows metal behavior. An element is more metallic the more easily its atoms give away their outer electrons.
Check your understanding
Potassium atoms give away their outer electron very easily; chlorine atoms almost never give electrons away — they pull electrons in instead. Which element has the stronger metallic character?
APotassiumcorrect
BChlorine
This option is wrong — you rewarded the electron-pulling atom — metallic character grows with giving electrons away, not pulling them in.
CBoth are equally metallic, because both react with other elements
This option is wrong — you equated reacting with being metallic — what matters is HOW the atoms react: giving electrons away is the metallic way.
DNeither has any metallic character
This option is wrong — you treated metallic character as all-or-nothing — it is a strength, and potassium's easy electron loss makes its metallic character strong.
An element is more metallic the more easily its atoms give away their outer electrons. Potassium's atoms give their outer electron away easily; chlorine's atoms pull electrons in. So potassium has the stronger metallic character.
Check your understanding
An element's atoms hold their outer electrons tightly and tend to pull in extra electrons instead of giving any away. What does that say about the element?
AIts metallic character is weak — it behaves as a nonmetal.correct
BIts metallic character is strong, because pulling electrons is a form of reacting.
This option is wrong — you inverted the definition — metallic character grows with giving outer electrons away, and this element does the opposite.
CIt must be a noble gas.
This option is wrong — you overreached — noble-gas atoms tend to leave electrons alone entirely; an atom that pulls electrons in is an ordinary reactive nonmetal.
DIt must be a liquid at room temperature.
This option is wrong — you reached for the element's state — metallic character is about electron behavior, not about being solid, liquid, or gas.
An element is more metallic the more easily its atoms give away their outer electrons. These atoms hold their electrons tightly and pull more in — the opposite behavior. So the element's metallic character is weak: nonmetal behavior.
Lesson 42 of 50 · PTB-042
Metallic character across a period
Learning Journey
Read · Check · Unlock
Did You Know?
Sodium and aluminum both sit in Period 3, and both are metals. Yet sodium gives its outer electron away far more readily than aluminum parts with its three — sodium is the more metallic of the two.
The idea
Moving from sodium to aluminum across Period 3, the nuclear charge increases while the outer electrons stay in the same shell.
Outer electrons that are pulled in more strongly are harder to give away.
Atoms that find it harder to give away outer electrons show weaker metal behavior, so metallic character fades from left to right along the row.
Metallic character decreases from left to right across a period because across a period the nuclear charge increases while the outer electrons stay in the same shell, so the outer electrons are pulled in more strongly and are harder to give away.
down Group 1: ionization energy decreases
Worked examples
Worked example 1. Magnesium and chlorine both sit in Period 3. Why is magnesium far more metallic than chlorine?
Step 1
Chlorine sits far to the right of magnesium, so its nucleus pulls on outer electrons in the same shell with a much bigger charge.
Step 2
Across a period the nuclear charge increases while the outer electrons stay in the same shell, so the outer electrons are pulled in more strongly and are harder to give away.
Step 3
Magnesium's outer electrons, feeling the smaller pull, are given away far more easily — so magnesium shows the stronger metal behavior.
You can now explain that metallic character decreases from left to right across a period because across a period the nuclear charge increases while the outer electrons stay in the same shell, so the outer electrons are pulled in more strongly and are harder to give away.
Check your understanding
Lithium and carbon both sit in Period 2. Which element shows the stronger metal behavior?
ALithiumcorrect
BCarbon
This option is wrong — you inverted the period direction — metallic character decreases from left to right, and carbon sits farther right.
CBoth equally, because both sit in Period 2
This option is wrong — you treated a shared period as shared behavior — the ease of giving away outer electrons fades across the row.
DIt cannot be judged from their positions
This option is wrong — you gave up on the positions — same-period positions decide this: the farther left, the more metallic.
Lithium sits at the left end of Period 2; carbon sits near the middle. Across a period the nuclear charge increases while the outer electrons stay in the same shell, so the outer electrons are pulled in more strongly and are harder to give away. Lithium's outer electron goes more easily, so lithium is the more metallic element.
Check your understanding
Potassium is far more metallic than bromine; both sit in Period 4. Why?
ABecause across a period the nuclear charge increases while the outer electrons stay in the same shell, so the outer electrons are pulled in more strongly and are harder to give away.correct
BBecause down a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly and are easier to give away.
This option is wrong — you applied the group mechanism — potassium and bromine sit in the same period, so the shell count does not separate them.
CBecause potassium atoms are heavier than bromine atoms, and heavier atoms shed electrons more easily.
This option is wrong — you tied electron loss to mass and reversed the masses — bromine's atoms are heavier, and mass is not the mechanism anyway.
DBecause bromine's outer electrons sit farther from its nucleus than potassium's do.
This option is wrong — you moved bromine's electrons outward — both elements keep outer electrons in the same shell; bromine's are pulled in more strongly by its bigger nuclear charge.
Bromine sits far to the right of potassium in Period 4. Across a period the nuclear charge increases while the outer electrons stay in the same shell, so the outer electrons are pulled in more strongly and are harder to give away. Potassium's outer electron, feeling the smaller pull, goes easily — that is its stronger metallic character.
Check your understanding
Beryllium and oxygen both sit in Period 2. Which element's atoms give away outer electrons more easily?
ABerylliumcorrect
BOxygen
This option is wrong — you inverted the period direction — oxygen's bigger nuclear charge pulls its outer electrons in more strongly, so they are harder to give away.
CBoth equally, because both are Period 2 elements
This option is wrong — you treated a shared period as shared behavior — the pull on outer electrons grows across the row.
DNeither — only Group 1 elements can give electrons away
This option is wrong — you fenced electron loss into Group 1 — many atoms can give away outer electrons; how easily is what varies with position.
Beryllium sits well to the left of oxygen in Period 2. Across a period the nuclear charge increases while the outer electrons stay in the same shell, so the outer electrons are pulled in more strongly and are harder to give away. Beryllium's outer electrons feel the smaller pull, so they are given away more easily.
Lesson 43 of 50 · PTB-043
Metallic character down a group
Learning Journey
Read · Check · Unlock
Did You Know?
Sodium and potassium are both Group 1 metals, but they are not equally metallic: potassium, one row lower, gives its outer electron away even more readily than sodium does.
The idea
Potassium sits one row below sodium, so its outer electron sits in a shell farther from the nucleus.
An outer electron that is held more weakly is easier to give away.
Atoms that give their outer electrons away more easily show stronger metal behavior, so potassium is the more metallic of the two.
Metallic character increases from top to bottom down a group because down a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly and are easier to give away.
down Group 1: ionization energy decreases
Worked examples
Worked example 1. Carbon sits at the top of Group 14 and lead near the bottom. Carbon is a nonmetal, while lead is a metal. Why does metal behavior grow down this group?
Step 1
Each row down from carbon to lead adds an electron shell, so lead's outer electrons sit far from the nucleus.
Step 2
Down a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly and are easier to give away.
Step 3
Lead's weakly held outer electrons go easily, so lead behaves as a metal while carbon, holding its electrons tightly, does not.
You can now explain that metallic character increases from top to bottom down a group because down a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly and are easier to give away.
Check your understanding
Lithium and cesium both sit in Group 1. Which element's atoms part with their outer electron more easily?
ACesiumcorrect
BLithium
This option is wrong — you inverted the group direction — cesium sits lower, so its outer electron sits farther from the nucleus and is held more weakly.
CBoth equally, because each atom has one outer electron
This option is wrong — you counted outer electrons instead of shells — the outer electron's distance from the nucleus, not its count, sets how easily it goes.
DIt cannot be judged from their positions
This option is wrong — you gave up on the positions — same-group positions decide this: the lower element's outer electron is held more weakly.
Cesium sits four rows below lithium in Group 1. Down a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly and are easier to give away. So cesium's outer electron goes more easily — cesium is the more metallic element.
Check your understanding
Barium is more metallic than beryllium; both sit in Group 2. Why?
ABecause down a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly and are easier to give away.correct
BBecause across a period the nuclear charge increases while the outer electrons stay in the same shell, so the outer electrons are pulled in more strongly.
This option is wrong — you applied the period mechanism — barium and beryllium sit in the same group, so the added shells are what separate them.
CBecause barium has more protons, and more protons make an atom more metallic.
This option is wrong — you counted protons only — more protons pull electrons in harder; what makes barium more metallic is its outer electrons sitting farther out.
DBecause barium atoms have more outer electrons to give away than beryllium atoms.
This option is wrong — you miscounted — both Group 2 elements have two outer electrons; the difference is how weakly barium holds them.
Barium sits four rows below beryllium in Group 2. Down a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly and are easier to give away. Electrons that go easily are the mark of strong metallic character.
Check your understanding
Silicon and tin both sit in Group 14. Which element shows the stronger metal behavior?
ATincorrect
BSilicon
This option is wrong — you inverted the group direction — tin sits lower, so its outer electrons are held more weakly and go more easily.
CBoth equally, because both sit in Group 14
This option is wrong — you treated a shared group as shared behavior — metal behavior strengthens with every added row down the column.
DNeither — Group 14 elements are all nonmetals
This option is wrong — you flattened the group — Group 14 runs from nonmetal carbon at the top to metals at the bottom, exactly because metallic character grows downward.
Tin sits two rows below silicon in Group 14. Down a group each row adds an electron shell, so the outer electrons sit farther from the nucleus and are held more weakly and are easier to give away. So tin's atoms part with outer electrons more easily — stronger metal behavior.
Lesson 44 of 50 · PTB-044
Nonmetal reactivity down a group
Learning Journey
Read · Check · Unlock
Did You Know?
Fluorine, at the top of Group 17, reacts with almost everything it touches, sometimes explosively. Iodine, three rows below it, is tame enough to swab on a cut. Same family — very different vigor.
The idea
A nonmetal atom reacts by pulling in extra electrons, and it reacts more easily when that pull is stronger.
Down the group the outer shell sits farther from the nucleus, so the pull on an incoming electron is weaker.
Fluorine's pull on an incoming electron is the group's strongest, and iodine's is far weaker — so fluorine reacts fiercely while iodine reacts gently.
Reactivity decreases from top to bottom down a group of nonmetals because a nonmetal atom reacts by pulling in extra electrons, and it reacts more easily when that pull is stronger — down the group the outer shell sits farther from the nucleus, so the pull on an incoming electron is weaker.
down Group 1: ionization energy decreases
Worked examples
Worked example 1. Chlorine reacts vigorously with hot iron wool; bromine reacts noticeably more slowly with it. Both are Group 17 nonmetals. Why?
Step 1
Bromine sits one row below chlorine, so its outer shell sits farther from the nucleus.
Step 2
A nonmetal atom reacts by pulling in extra electrons, and it reacts more easily when that pull is stronger — bromine's pull on an incoming electron is the weaker of the two.
Step 3
So chlorine grabs electrons from the iron more readily, and its reaction runs more vigorously.
You can now explain that reactivity decreases from top to bottom down a group of nonmetals because a nonmetal atom reacts by pulling in extra electrons, and it reacts more easily when that pull is stronger — down the group the outer shell sits farther from the nucleus, so the pull on an incoming electron is weaker.
Check your understanding
Oxygen and sulfur both sit in Group 16. Which reacts more readily with a metal like iron?
AOxygencorrect
BSulfur
This option is wrong — you inverted the nonmetal direction — sulfur's outer shell sits farther from the nucleus, so its pull on an incoming electron is weaker.
CBoth equally, because both are Group 16 nonmetals
This option is wrong — you treated a shared group as shared vigor — the pull on incoming electrons weakens with every added row.
DNeither reacts with metals
This option is wrong — you cut nonmetals off from metals — grabbing electrons from metal atoms is exactly how these nonmetals react, rust being the everyday example.
Oxygen sits above sulfur in Group 16. A nonmetal atom reacts by pulling in extra electrons, and it reacts more easily when that pull is stronger — oxygen's outer shell sits closer to the nucleus, so its pull is the stronger. So oxygen reacts with iron more readily.
Check your understanding
Selenium sits above tellurium in Group 16, and selenium reacts more easily. Why?
ABecause a nonmetal atom reacts by pulling in extra electrons, and it reacts more easily when that pull is stronger — down the group the outer shell sits farther from the nucleus, so the pull on an incoming electron is weaker.correct
BBecause an atom reacts by giving away its outer electrons, and it reacts more easily when those electrons are easier to remove — selenium, higher in the group, holds its outer electrons more loosely than tellurium does.
This option is wrong — you used the metal mechanism — nonmetals react by pulling electrons IN, and the higher element pulls harder.
CBecause selenium's atoms have more protons than tellurium's, and an atom with more protons always reacts more readily.
This option is wrong — you counted protons and reversed the counts — tellurium has more protons; what matters is that selenium's outer shell sits closer to the nucleus.
DBecause tellurium's outer shell is already full, so its atoms have no room to take in any more electrons.
This option is wrong — you gave tellurium a full shell — both elements have six outer electrons and room for more; tellurium's pull on them is just weaker.
Selenium sits one row above tellurium in Group 16. A nonmetal atom reacts by pulling in extra electrons, and it reacts more easily when that pull is stronger. Down the group the outer shell sits farther from the nucleus, so tellurium's pull on an incoming electron is weaker.
Check your understanding
Astatine sits at the very bottom of Group 17. Predict how its reactivity compares with the other elements in the group.
AThe least reactive of the groupcorrect
BThe most reactive of the group
This option is wrong — you inverted the nonmetal direction — the pull on an incoming electron is weakest at the bottom of the group, so the bottom element reacts least readily.
CAbout the same as the elements above it
This option is wrong — you flattened the group — the pull on incoming electrons weakens with every added row, so vigor fades down the column.
DImpossible to say without an experiment
This option is wrong — you gave up on the trend — the group direction predicts it: bottom of a nonmetal group, weakest pull, least reactive.
Astatine sits at the bottom of Group 17, a nonmetal group. A nonmetal atom reacts by pulling in extra electrons, and it reacts more easily when that pull is stronger — down the group the outer shell sits farther from the nucleus, so the pull on an incoming electron is weaker. The weakest pull of the group means the least reactive member.
Lesson 45 of 50 · PTB-045
Rank elements by reactivity
Learning Journey
Read · Check · Unlock
Did You Know?
You've now seen two reactivity directions that point opposite ways: metals grow more reactive down a group, nonmetals less. Ranking elements by reactivity therefore starts with a question none of the other rankings needed.
The idea
Before ranking elements by reactivity, ask one question first: are they metals or nonmetals?
down Group 1: ionization energy decreases
For metals, reactivity increases down the group — a metal reacts by giving away outer electrons, and lower down those electrons are held more weakly.
For nonmetals, reactivity decreases down the group — a nonmetal reacts by pulling in extra electrons, and lower down that pull is weaker.
Lithium, sodium, and potassium are Group 1 metals, so the most reactive is the lowest: from least to most reactive the order is lithium, sodium, potassium.
Chlorine, bromine, and iodine are Group 17 nonmetals, so the most reactive is the highest: from least to most reactive the order is iodine, bromine, chlorine.
Worked examples
Worked example 1. Rank magnesium, calcium, and barium (all Group 2 metals) from least to most reactive.
Step 1
All three are metals, so reactivity increases down the group.
Step 2
Magnesium sits highest, then calcium, then barium at the bottom.
Step 3
From least to most reactive: magnesium, calcium, barium.
Worked example 2. Rank oxygen, sulfur, and selenium (all Group 16 nonmetals) from least to most reactive.
Step 1
All three are nonmetals, so reactivity decreases down the group.
Step 2
Selenium sits lowest, then sulfur, then oxygen at the top.
Step 3
From least to most reactive: selenium, sulfur, oxygen.
You can now rank elements from the same group by reactivity using their positions in the periodic table, applying the direction that fits the element class — metals grow more reactive down a group, nonmetals less.
Check your understanding
Rubidium and cesium both sit in Group 1. Which reacts more vigorously with water?
ACesiumcorrect
BRubidium
This option is wrong — you inverted the metal direction — metals grow MORE reactive down a group, and cesium sits lower.
CBoth react identically, because both are alkali metals
This option is wrong — you flattened the family — the family shares a reaction pattern, but the vigor grows down the column.
DNeither reacts with water
This option is wrong — you forgot the alkali metals' signature behavior — Group 1 metals react with water, increasingly violently down the group.
Both are Group 1 metals, so use the metal direction: more reactive down the group. Cesium sits below rubidium. So cesium reacts more vigorously.
Check your understanding
Fluorine and bromine both sit in Group 17. Which reacts more vigorously?
AFluorinecorrect
BBromine
This option is wrong — you applied the metal direction to a nonmetal group — nonmetals grow LESS reactive down a group, so the top element is the fiercest.
CBoth react identically, because both are halogens
This option is wrong — you flattened the family — halogens share a reaction pattern, but the vigor fades down the column.
DIt cannot be judged without knowing what they react with
This option is wrong — you demanded the partner — for same-group comparisons the position already ranks the vigor, whatever the partner.
Both are Group 17 nonmetals, so use the nonmetal direction: less reactive down the group. Fluorine sits at the top of the group. So fluorine reacts more vigorously.
Check your understanding
Potassium sits below sodium in Group 1, and bromine sits below chlorine in Group 17. Which pair of statements is correct?
APotassium is more reactive than sodium, and chlorine is more reactive than bromine.correct
BPotassium is more reactive than sodium, and bromine is more reactive than chlorine.
This option is wrong — you applied the metal direction to the halogens — nonmetals grow less reactive down a group, so chlorine beats bromine.
CSodium is more reactive than potassium, and chlorine is more reactive than bromine.
This option is wrong — you applied the nonmetal direction to the metals — metals grow more reactive down a group, so potassium beats sodium.
DSodium is more reactive than potassium, and bromine is more reactive than chlorine.
This option is wrong — you inverted both directions — metals strengthen downward and nonmetals weaken downward, the opposite of both claims.
Ask the class question first: Group 1 elements are metals; Group 17 elements are nonmetals. Metals grow more reactive down a group, so potassium is more reactive than sodium. Nonmetals grow less reactive down a group, so chlorine is more reactive than bromine.
Lesson 46 of 50 · PTB-046
Identify the trend in a data table
Learning Journey
Read · Check · Unlock
Did You Know?
So far the trends have come as statements. Real chemistry hands you the opposite: a table of measured values, and the job of saying what trend the numbers show. The table lists atomic radius for the Period 3 elements, in order from sodium to chlorine.
The idea
First check how the elements are ordered: sodium through chlorine runs left to right across Period 3.
Atomic radius across Period 3
Element
Atomic radius (pm)
Na
186
Mg
160
Al
143
Si
118
P
110
S
104
Cl
99
Then read the values in that order: 186, 160, 143, 118, 110, 104, 99 — every step is a decrease.
State what the values show: atomic radius decreases from left to right across Period 3.
A complete trend statement always names three things: the property, the direction of change, and the path through the table.
Worked examples
Worked example 1. A table lists ionization energy (kJ/mol) for the Group 1 elements in order down the group — lithium 520, sodium 496, potassium 419, rubidium 403, cesium 376. What trend does the table show?
Step 1
The elements run top to bottom down Group 1.
Step 2
Reading the values in that order, every step is a decrease: 520, 496, 419, 403, 376.
Step 3
So the table shows that ionization energy decreases from top to bottom down Group 1.
You can now identify the trend shown by a table of elemental data (atomic radius, atomic mass, electronegativity, ionization energy, or reactivity) for elements listed in period or group order.
Check your understanding
The table lists electronegativity values for the Period 2 elements in order from lithium to fluorine. What trend does the table show?
Electronegativity across Period 2
Element
Electronegativity
Li
0.98
Be
1.57
B
2.04
C
2.55
N
3.04
O
3.44
F
3.98
AElectronegativity increases from left to right across the period.correct
BElectronegativity decreases from left to right across the period.
This option is wrong — you read the direction backwards — each value in the table is larger than the one before it.
CElectronegativity rises and then falls across the period.
This option is wrong — you saw a peak that is not there — check every step: no value in the list is smaller than the one before it.
DThe values show no clear trend.
This option is wrong — you stopped before comparing the steps — every step from 0.98 up to 3.98 is an increase, which is as clear as a trend gets.
The elements run left to right across Period 2. Reading the values in that order, every step is an increase: 0.98 up to 3.98. So the table shows electronegativity increasing from left to right across the period.
Check your understanding
The table lists atomic radius for four Group 17 elements in order from fluorine down to iodine. What trend does the table show?
Atomic radius down Group 17
Element
Atomic radius (pm)
F
72
Cl
99
Br
114
I
133
AAtomic radius increases from top to bottom down the group.correct
BAtomic radius decreases from top to bottom down the group.
This option is wrong — you read the direction backwards — each value in the table is larger than the one above it.
CAtomic radius stays roughly constant down the group.
This option is wrong — you flattened real growth — the values nearly double from 72 pm to 133 pm.
DAtomic radius increases and then decreases down the group.
This option is wrong — you saw a turn that is not there — every step in the list is an increase.
The elements run top to bottom down Group 17. Reading the values in that order, every step is an increase: 72, 99, 114, 133. So the table shows atomic radius increasing from top to bottom down the group.
Check your understanding
The table lists ionization energy for the Group 2 elements in order from beryllium down to barium. What trend does the table show?
Ionization energy down Group 2
Element
Ionization energy (kJ/mol)
Be
899
Mg
738
Ca
590
Sr
549
Ba
503
AIonization energy decreases from top to bottom down the group.correct
BIonization energy increases from top to bottom down the group.
This option is wrong — you read the direction backwards — each value in the table is smaller than the one above it.
CIonization energy decreases and then increases down the group.
This option is wrong — you saw a turn that is not there — check every step: each of the five values is smaller than the last.
DThe values show no clear trend.
This option is wrong — you stopped before comparing the steps — a steady fall from 899 to 503 is a clear trend.
The elements run top to bottom down Group 2. Reading the values in that order, every step is a decrease: 899 down to 503. So the table shows ionization energy decreasing from top to bottom down the group.
Lesson 47 of 50 · PTB-047
The repeating pattern in a property graph
Learning Journey
Read · Check · Unlock
Did You Know?
Wonder this:
As atomic number climbs from 1 to 20, does ionization energy simply climb with it? Plot one point per element and see.
The graph shows ionization energy against atomic number for the first 20 elements — one point per element, joined in order.
The idea
The graph does not climb steadily — it rises to a sharp peak, crashes, and rises to a peak again.
Every peak sits on a noble gas: helium, neon, argon.
Every crash lands on a Group 1 metal: lithium, sodium, potassium.
Each crash marks the start of a new period — the pattern repeats, period after period.
A property that repeats like this as atomic number grows is showing the periodic pattern — the same repetition the periodic law describes.
Worked examples
Worked example 1. The second panel plots atomic radius against atomic number (noble gases left out). Where are the peaks, and what do they mark?
Step 1
The tallest points sit at lithium, sodium, and potassium — the Group 1 metals.
Step 2
After each peak the values fall across the period, then jump back up.
Step 3
Each jump up to a new peak marks the start of a new period — the same repeating pattern, with the peaks on different elements than the ionization-energy graph.
You can now identify the repeating periodic pattern in a graph of a property against atomic number, including where each new period begins.
Check your understanding
The graph plots ionization energy against atomic number for elements 1 through 36 (transition metals omitted). After argon, which element gives the next peak?
AKryptoncorrect
BBromine
This option is wrong — you stopped one element short — bromine's point is high but the graph keeps rising one more step, to krypton.
CRubidium
This option is wrong — you overshot into the next crash — the element after a peak starts the new period at a low point.
DCalcium
This option is wrong — you picked a point on the early rise — calcium's value sits far below the peak at the period's end.
Peaks in this graph sit at the end of each period, on the noble gases. After argon, the plotted values climb across Period 4 and top out at krypton. So krypton gives the next peak.
Check your understanding
On the same graph of elements 1 through 36 (transition metals omitted), which element marks the start of Period 4 — the sharp drop right after argon's peak?
APotassiumcorrect
BArgon
This option is wrong — you picked the peak itself — argon ENDS Period 3; the crash to the next point is where Period 4 begins.
CSodium
This option is wrong — you picked the previous period's start — sodium's crash follows neon's peak, one period earlier.
DKrypton
This option is wrong — you jumped to the next peak — krypton ends Period 4 rather than starting it.
Each crash right after a peak marks the start of a new period. Argon's peak ends Period 3, and the value crashes at the next plotted element. That element is potassium — the start of Period 4.
Check your understanding
This graph plots atomic radius against atomic number for elements 3 through 38, with the noble gases left out. The tallest points recur at four elements. Which family are they?
AThe alkali metals (Group 1)correct
BThe noble gases (Group 18)
This option is wrong — you carried the ionization-energy graph's peaks over — this radius graph leaves the noble gases out, and its peaks sit at the period STARTS.
CThe halogens (Group 17)
This option is wrong — you picked the period's right-hand end — radius shrinks across a period, so the halogens sit near each period's low point.
DThe alkaline earth metals (Group 2)
This option is wrong — you landed one element after each peak — the very largest atom in each period is the Group 1 metal that starts it.
The labeled tallest points are lithium, sodium, potassium, and rubidium. All four sit in Group 1 — each one starts a new period as the period's largest atom. So the peaks of the radius graph are the alkali metals.
Lesson 48 of 50 · PTB-048
Estimate a missing value from group neighbors
Learning Journey
Read · Check · Unlock
Have You Ever Wondered?
Wonder this:
Suppose chlorine's ionization energy went missing from every data table. Chlorine sits between fluorine and bromine in Group 17 — could you still say something useful about the missing number?
You could, because of what you've seen about group data: when a group's values follow a steady trend, position pins a missing value down. The table shows Group 17's ionization energies with chlorine's entry blank.
The idea
Group 17's ionization energies fall steadily down the column: 1681 for fluorine, the blank, 1140 for bromine, 1008 for iodine.
Ionization energies down Group 17
Element
Ionization energy (kJ/mol)
F
1681
Cl
? (highlighted blank)
Br
1140
I
1008
When a group's data follow a steady trend, an element's value lies between the values of its neighbors above and below.
So chlorine's ionization energy must lie between bromine's 1140 and fluorine's 1681 kJ/mol.
The prediction is a range, not an exact number — the data alone do not say where inside the range the value falls.
The measured value, 1251 kJ/mol, does fall inside that range.
Worked examples
Worked example 1. A Group 1 table lists atomic radii of sodium (186 pm) and rubidium (248 pm), but potassium's entry — between them — is blank. Between which two values must potassium's radius lie?
Step 1
The group's radii grow steadily down the column, and potassium sits between sodium and rubidium.
Step 2
When a group's data follow a steady trend, an element's value lies between the values of its neighbors above and below.
Step 3
So potassium's radius must lie between those neighbors' values — the data give the range; they do not pick a number inside it.
Step 4
Potassium's radius lies between 186 pm and 248 pm — and the measured value, 227 pm, falls inside that range.
You can now predict that an element's property value lies between the values of its neighbors above and below in the same group, when the group data follow a steady trend.
Check your understanding
The table shows ionization energies for three Group 2 elements, with calcium's entry blank between magnesium's and strontium's. What can you predict about calcium's ionization energy?
Ionization energies down Group 2
Element
Ionization energy (kJ/mol)
Mg
738
Ca
? (highlighted blank)
Sr
549
AIt lies between 549 and 738 kJ/mol.correct
BIt is above 738 kJ/mol.
This option is wrong — you put the middle element outside its neighbors — in a steady group trend the between element takes a between value.
CIt is below 549 kJ/mol.
This option is wrong — you carried the fall past the blank — the value below calcium's is strontium's; calcium itself sits between its neighbors.
DIt is exactly halfway between them, 643.5 kJ/mol.
This option is wrong — you computed a midpoint — the data give a range, not a formula; nothing says the value sits exactly halfway.
The group's values fall steadily, and calcium sits between magnesium and strontium. When a group's data follow a steady trend, an element's value lies between the values of its neighbors above and below. So calcium's value lies between 549 and 738 kJ/mol — the measured value is 590.
Check your understanding
The table shows atomic radii for three Group 16 elements, with sulfur's entry blank between oxygen's and selenium's. What can you predict about sulfur's atomic radius?
Atomic radii down Group 16
Element
Atomic radius (pm)
O
73
S
? (highlighted blank)
Se
117
AIt lies between 73 and 117 pm.correct
BIt is below 73 pm.
This option is wrong — you ran the growth downward — radii grow down this group, and sulfur sits below oxygen, so its value sits above 73, not below.
CIt is above 117 pm.
This option is wrong — you carried the growth past the blank — selenium's 117 belongs to the element BELOW sulfur, so sulfur's value stays under it.
DIt is exactly 95 pm, the halfway point.
This option is wrong — you computed a midpoint — the data give a range, not a formula; the measured value, 104 pm, is not the halfway point.
The group's radii grow steadily down the column, and sulfur sits between oxygen and selenium. When a group's data follow a steady trend, an element's value lies between the values of its neighbors above and below. So sulfur's radius lies between 73 and 117 pm — the measured value is 104.
Check your understanding
The table shows melting points for three Group 1 metals, with sodium's entry blank between lithium's and potassium's. What can you predict about sodium's melting point?
Melting points down Group 1
Element
Melting point (°C)
Li
181
Na
? (highlighted blank)
K
63
AIt lies between 63 and 181 °C.correct
BIt is above 181 °C.
This option is wrong — you put the middle element outside its neighbors — the values fall steadily, so sodium's sits below lithium's 181.
CIt is below 63 °C.
This option is wrong — you carried the fall past the blank — 63 belongs to potassium, the element BELOW sodium, so sodium's value stays above it.
DIt is exactly 122 °C, the halfway point.
This option is wrong — you computed a midpoint — the data give a range, not a formula; the measured value, 98 °C, is nowhere near halfway.
The group's melting points fall steadily, and sodium sits between lithium and potassium. When a group's data follow a steady trend, an element's value lies between the values of its neighbors above and below. So sodium's melting point lies between 63 and 181 °C — the measured value is 98.
Lesson 49 of 50 · PTB-049
Judge data against a claimed trend
Learning Journey
Read · Check · Unlock
Did You Know?
A claim about a trend is cheap to make. The table gives you the means to check one: it lists the Group 1 melting points, and the claim to test is 'melting point decreases down Group 1'.
The idea
Take the claim apart into its three parts: the property (melting point), the direction (decreases), and the path (down Group 1).
Claim: melting point decreases down Group 1
Element
Melting point (°C)
Li
181
Na
98
K
63
Rb
39
Cs
28
Check every step of the data against the claim: 181 to 98 to 63 to 39 to 28 — each step is a decrease.
Every value agrees with the claim, so the data support it.
One disagreeing value is enough to break a claim, so a fair judgment checks every step and cites the values that agree or disagree.
Worked examples
Worked example 1. Claim: 'atomic radius decreases down Group 17.' The data: fluorine 72 pm, chlorine 99 pm, bromine 114 pm, iodine 133 pm. Do the data support the claim?
Step 1
The claim says the values should fall down the group.
Step 2
Check every step: 72 to 99 to 114 to 133 — every step is an INCREASE.
Step 3
The data contradict the claim at every step — atomic radius grows down this group, with 72 pm against 133 pm at the extremes.
You can now evaluate whether a supplied set of elemental data supports a stated trend, citing the values that agree or disagree.
Check your understanding
Claim: 'density increases down Group 1.' The table lists the five metals' densities in order down the group. Do the data support the claim?
Claim: density increases down Group 1
Element
Density (g/cm³)
Li
0.53
Na
0.97
K
0.86
Rb
1.53
Cs
1.87
ANot fully — potassium's 0.86 g/cm³ is lower than sodium's 0.97, breaking the rise.correct
BYes — every value increases down the group.
This option is wrong — you skipped a step — check sodium to potassium: 0.97 down to 0.86 is a decrease, and one disagreeing step breaks a claim.
CNo — every value decreases down the group.
This option is wrong — you flipped the whole table — four of the five steps rise; the single potassium dip is what spoils the claim, not a general fall.
DThe data cannot test the claim, because density is not a group property.
This option is wrong — you dismissed the data — any measured property listed in group order can be checked against a claim about it.
Check every step: 0.53 to 0.97 rises, but 0.97 to 0.86 FALLS, then 0.86 to 1.53 to 1.87 rise again. One disagreeing value is enough to break a claim. So the data do not fully support the claim — cite potassium's 0.86 against sodium's 0.97.
Check your understanding
Claim: 'boiling point decreases down Group 17.' The table lists four halogens' boiling points in order down the group. Do the data support the claim?
Claim: boiling point decreases down Group 17
Element
Boiling point (°C)
F
-188
Cl
-34
Br
59
I
184
ANo — every step is an increase, from -188 °C up to 184 °C, contradicting the claim.correct
BYes — read from iodine at the bottom of the table up to fluorine, the values go from 184 down to -188.
This option is wrong — you read the table from the bottom up — in group order, fluorine to iodine, every step rises.
CYes — negative values mean the boiling points are decreasing.
This option is wrong — you read the minus signs as a direction — -34 is HIGHER than -188; compare the values, not their signs.
DThe data are mixed, with some steps rising and some falling.
This option is wrong — you found disagreement that is not there — all three steps rise: -188 to -34 to 59 to 184.
Check every step in group order: -188 to -34 rises, -34 to 59 rises, 59 to 184 rises. The claim said decrease; every step increases instead. So the data contradict the claim at every step.
Check your understanding
Claim: 'electronegativity increases across Period 2 from lithium to fluorine.' The table lists the seven values in order. Do the data support the claim?
Claim: electronegativity increases across Period 2
Element
Electronegativity
Li
0.98
Be
1.57
B
2.04
C
2.55
N
3.04
O
3.44
F
3.98
AYes — every step rises, from 0.98 up to 3.98, so every value supports the claim.correct
BNo — the values fall from fluorine back to lithium.
This option is wrong — you read the table against the claim's path — the claim runs lithium to fluorine, and in that order every step rises.
CNot fully — one of the middle values breaks the rise.
This option is wrong — you assumed a break without finding one — check each of the six steps; none falls.
DThe data cannot test the claim, because electronegativity has no units.
This option is wrong — you demanded units — the claim is about direction, and unitless values can rise or fall like any others.
The claim's path is left to right, lithium to fluorine. Check every step: 0.98, 1.57, 2.04, 2.55, 3.04, 3.44, 3.98 — all six steps rise. Every value agrees, so the data support the claim.
Lesson 50 of 50 · PTB-050
Name the property behind an unlabeled data set
Learning Journey
Read · Check · Unlock
Have You Ever Wondered?
Wonder this:
A data table arrives with its label torn off: seven Period 3 elements, listed sodium to chlorine, next to seven rising numbers. Nobody wrote down which property was measured. Can the numbers themselves tell you?
They can, because you now know how each property behaves across the table. The torn-label table reads: 0.93, 1.31, 1.61, 1.90, 2.19, 2.58, 3.16.
The idea
Read the elements and their order first: sodium through chlorine runs left to right across Period 3.
Property: ______ (label torn off)
Element
Value
Na
0.93
Mg
1.31
Al
1.61
Si
1.90
P
2.19
S
2.58
Cl
3.16
Read the direction next: every value rises.
Match the direction against the trends you've seen: across a period, atomic radius falls, while electronegativity rises.
Check the size of the numbers too: electronegativity's whole scale runs from about 0.7 to 4, while ionization energies are hundreds of kJ/mol and radii are around 100 pm.
Rising across the period, on a scale that tops out near 4 — the torn label said electronegativity, not atomic radius.
Worked examples
Worked example 1. An unlabeled table lists the Group 1 elements from lithium down to cesium beside the values 520, 496, 419, 403, 376. Which property was most likely measured?
Step 1
The elements run down Group 1, and every value falls.
Step 2
Down a group, atomic radius, metallic character, and reactivity all grow — a falling data set rules them out.
Step 3
Ionization energy and electronegativity both fall down a group, but electronegativity never leaves the 0.7-to-4 scale — values in the hundreds fit ionization energy in kJ/mol.
Step 4
So the data set most likely shows ionization energy.
You can now identify which property an unlabeled elemental data set most likely shows by matching its pattern against the established periodic trends.
Check your understanding
An unlabeled table lists the Period 2 elements from lithium to fluorine beside the values shown in the figure. Which property was most likely measured?
Property: ______ (label torn off)
Element
Value
Li
152
Be
112
B
85
C
77
N
75
O
73
F
72
AAtomic radiuscorrect
BElectronegativity
This option is wrong — you matched the wrong direction — electronegativity RISES across a period, and these values fall; the 0.7-to-4 scale does not fit values above 70 either.
CIonization energy
This option is wrong — you matched the wrong direction — ionization energy generally rises across a period, and these values fall.
DAtomic mass
This option is wrong — you ignored both cues — atomic mass rises across a period, and lithium's mass is about 7, nothing like 152.
The elements run left to right across Period 2, and every value falls. Across a period only atomic radius, of these candidates, falls — electronegativity, ionization energy, and mass all rise. Values near 100 also fit radii in picometers, so the property is atomic radius.
Check your understanding
An unlabeled table lists four Group 17 elements from fluorine down to iodine beside the values shown in the figure. Which property was most likely measured?
Property: ______ (label torn off)
Element
Value
F
3.98
Cl
3.16
Br
2.96
I
2.66
AElectronegativitycorrect
BAtomic radius
This option is wrong — you matched the wrong direction — radius GROWS down a group, and these values fall; radii also run near 100 pm, not below 4.
CIonization energy
This option is wrong — you matched the direction but not the scale — ionization energies are hundreds of kJ/mol, while values between 2 and 4 sit on the electronegativity scale.
DAtomic mass
This option is wrong — you ignored both cues — mass grows down a group, and fluorine's mass is about 19, not 3.98.
The elements run down Group 17, and every value falls. Falling down a group fits ionization energy or electronegativity — radius and mass grow downward. The values sit between 2 and 4, on electronegativity's 0.7-to-4 scale, so the property is electronegativity.
Check your understanding
An unlabeled table lists four Group 1 metals in order down the group beside the recorded behaviors shown in the figure. Which property was most likely recorded?
Property: ______ (label torn off)
Element
Recorded behavior
Li
fizzes steadily in water
Na
fizzes vigorously, melts into a ball
K
bursts into a lilac flame
Cs
explodes on contact with water
AReactivitycorrect
BMelting point
This option is wrong — you seized on one melting detail — sodium melting is a side effect of its vigorous reaction; every row records how violently the metal reacts.
CElectronegativity
This option is wrong — you named a numeric property for observational data — electronegativity is a number on a 0.7-to-4 scale, not a description of fizzing and flames.
DAtomic radius
This option is wrong — you named a size for behavioral records — nothing in the rows measures how big the atoms are.
Each row records how vigorously the metal reacts with water. The vigor grows down the group — steady fizz to explosion — matching the metal reactivity trend. So the recorded property is reactivity.
Summary video — Periodic trends and reading elemental data