Bake a pizza at 300 °C and the cheese melts, the dough browns, and the sugar in the sauce caramelises. The salt sprinkled on top comes out exactly as it went in. Salt does not melt until 801 °C. What holds its particles together so stubbornly?
You've already seen that atoms can lose or gain electrons to become ions. This lesson answers a more basic question: why do particles stick to each other at all?
The idea
Particles stick together when being together gives them lower energy than being apart.
Chemists call this sticking a 'chemical bond'.
Atoms bond because the atoms end up with lower energy bonded together than they have when separate.
A ball behaves the same way: it rolls downhill and rests in the valley, because the valley is the lowest-energy place available.
Bonded particles sit at lower energy than the same particles kept apart.
Pulling bonded particles apart means pushing them back uphill in energy, and that always takes an energy supply from outside.
A sodium ion and a chloride ion bonded inside solid sodium chloride have lower energy than the same two particles kept apart.
That is why the salt survived the oven: at 300 °C, the heat supplied was not enough energy to pull its bonded particles apart.
Worked examples
Worked example 1. Solid magnesium oxide holds a magnesium ion Mg²⁺ bonded next to an oxide ion O²⁻. Compare the energy of the two bonded ions with the energy of the same two ions kept far apart.
Step 1
The two ions are bonded, and atoms bond because the atoms end up with lower energy bonded together than they have when separate.
Step 2
So the bonded Mg²⁺ and O²⁻ have lower energy than the same two ions kept far apart.
Step 3
The bonded pair sits at lower energy; the separated pair sits at higher energy.
Worked example 2. To pull apart the bonded ions in solid potassium fluoride, you must keep supplying energy. Why does separating bonded particles need an energy supply?
Step 1
Bonded particles sit at lower energy than separated particles.
Step 2
Moving from lower energy to higher energy is an uphill change.
Step 3
An uphill change only happens when energy is supplied from outside.
Step 4
Separating bonded particles moves them uphill in energy, so it needs an outside energy supply.
You can now explain that atoms bond because the atoms end up with lower energy bonded together than they have when separate.
Check your understanding
Why do atoms form chemical bonds?
ABecause the atoms end up with lower energy bonded together than they have when separate.correct
BBecause the atoms end up with higher energy bonded together, which makes them harder to pull apart.
This option is wrong — you flipped the energy direction — bonded particles sit at LOWER energy, and that lower energy is exactly what makes them hard to pull apart.
CBecause energy must be added to the atoms to make them stick together.
This option is wrong — you reversed the energy flow — bonding releases energy as the atoms drop to a lower level; it is pulling them apart that needs added energy.
DBecause the particles in any solid must touch each other.
This option is wrong — you treated touching as the cause — particles can touch without bonding; bonding happens when being together lowers their energy.
Bonding is a downhill change in energy. Atoms bond because the atoms end up with lower energy bonded together than they have when separate. Pulling bonded atoms apart is the uphill change, and that is the one that needs added energy.
Check your understanding
A lithium ion and a chloride ion sit bonded inside solid lithium chloride. Which statement about their energy is correct?
AThe bonded pair has lower energy than the same two ions kept apart.correct
BThe bonded pair has higher energy than the same two ions kept apart.
This option is wrong — you flipped the energy comparison — bonding is a drop to lower energy, like a ball rolling into a valley.
CThe bonded pair has the same energy as the two ions kept apart.
This option is wrong — you treated bonding as an energy-neutral change — if the energies were equal, nothing would hold the ions together.
DThe two ions had to gain energy before they could become bonded.
This option is wrong — you reversed the energy flow — the ions release energy when they bond; energy is needed to separate them, not to join them.
Atoms bond because the atoms end up with lower energy bonded together than they have when separate. The bonded Li⁺ and Cl⁻ therefore sit at lower energy than the separated pair. Bonding is the downhill direction; separating is the uphill direction.
Check your understanding
What must happen for two bonded ions to be pulled apart from each other?
AEnergy must be supplied to the ions from outside.correct
BEnergy must be released by the ions as they separate.
This option is wrong — you ran the energy flow backwards — energy is released when the ions bond; separating them absorbs energy.
CThe ions must lose energy until the bond breaks.
This option is wrong — you treated separation as a downhill change — separated ions sit at HIGHER energy, so the ions must be pushed uphill, not allowed to fall.
DNothing needs to happen, because bonded ions drift apart on their own over time.
This option is wrong — you ignored the energy valley — the bonded pair rests at the lowest energy available, so it stays put until outside energy pushes it uphill.
Bonded ions sit at lower energy than separated ions. Moving from lower energy to higher energy needs an energy supply from outside. That is why heat, which supplies energy, is what eventually pulls bonded particles apart.
Lesson 2 of 55 · IMB-002
The octet pattern
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen why atoms bond: the atoms end up with lower energy bonded together than they have when separate. Looking at what atoms end up WITH after bonding reveals a pattern.
The idea
Check the electrons of main-group atoms — the elements of Groups 1–2 and 13–18 — after they have bonded.
A main-group atom usually finishes bonding with the same electron arrangement as a noble gas.
Most often that means eight outer electrons, so the pattern is called the 'octet pattern'.
A sodium atom has 11 electrons; the sodium ion Na⁺ has 10 — the same electron arrangement as neon.
The octet pattern describes the OUTCOME of bonding, not the reason bonding happens.
The octet pattern
Ion after bonding
Electrons
Noble gas with the same arrangement
Na⁺
10
neon (10 electrons)
S²⁻
18
argon (18 electrons)
After bonding, main-group atoms usually match a noble gas's electron arrangement. Noble-gas electron counts: helium 2, neon 10, argon 18, krypton 36.
The reason is still energy: atoms bond because the atoms end up with lower energy bonded together than they have when separate.
The octet pattern is a description of where atoms tend to land, the way 'balls end up in valleys' describes where balls tend to land.
Worked examples
Worked example 1. A potassium atom has 19 electrons. After bonding, potassium exists as the ion K⁺. Which noble gas has the same electron arrangement as K⁺?
Step 1
K⁺ has lost one electron: 19 − 1 = 18 electrons.
Step 2
The noble gas with 18 electrons is argon.
Step 3
K⁺ has the same electron arrangement as argon — the octet pattern holds.
Worked example 2. An oxygen atom has 8 electrons. After bonding, oxygen exists as the ion O²⁻. Which noble gas has the same electron arrangement as O²⁻?
Step 1
O²⁻ has gained two electrons: 8 + 2 = 10 electrons.
Step 2
The noble gas with 10 electrons is neon.
Step 3
O²⁻ has the same electron arrangement as neon — gaining electrons reaches the pattern too.
You can now state that atoms of the main groups (Groups 1–2 and 13–18) usually finish bonding with the same electron arrangement as a noble gas — most often eight outer electrons, the octet pattern — and that this pattern describes the outcome of bonding, not the reason bonding happens.
Check your understanding
A calcium atom has 20 electrons. After bonding, calcium exists as the ion Ca²⁺. Which noble gas has the same electron arrangement as Ca²⁺?
AArgon, which has 18 electrons.correct
BNeon, which has 10 electrons.
This option is wrong — you matched the wrong noble gas — Ca²⁺ has 20 − 2 = 18 electrons, and 18 electrons is argon's arrangement, not neon's.
CKrypton, which has 36 electrons.
This option is wrong — you moved up to the noble gas after calcium — a 2+ ion has LOST two electrons, leaving 18 (argon's arrangement), not gained sixteen to reach krypton's 36.
DCalcium keeps its own 20-electron arrangement, which matches no noble gas.
This option is wrong — you ignored the charge — the 2+ superscript means two electrons are gone, and 18 electrons is exactly argon's arrangement.
A 2+ charge means the atom has lost two electrons. Ca²⁺ has 20 − 2 = 18 electrons. The noble gas with 18 electrons is argon, so Ca²⁺ fits the octet pattern.
Check your understanding
What does the octet pattern describe?
AThe electron arrangement main-group atoms usually END UP with after bonding.correct
BThe reason that bonding happens in the first place.
This option is wrong — you promoted the pattern to a cause — bonding happens because bonded atoms end up with lower energy, and the octet pattern only describes the outcome.
CThe number of bonds every atom in the periodic table must form.
This option is wrong — you read 'octet' as a bond count — the pattern is about eight OUTER ELECTRONS after bonding, and it covers main-group atoms, not every atom.
DThe energy released when two atoms bond together.
This option is wrong — you swapped the pattern for the cause's currency — energy explains WHY atoms bond; the octet pattern describes the electron arrangement they land in.
The octet pattern is a description of outcomes: main-group atoms usually finish bonding with a noble-gas electron arrangement, most often eight outer electrons. It is not the reason bonding happens. The reason is energy — atoms bond because the atoms end up with lower energy bonded together than they have when separate.
Check your understanding
A fluorine atom has 9 electrons. After bonding, fluorine exists as the ion F⁻. How does F⁻ fit the octet pattern?
AF⁻ has 10 electrons — the same electron arrangement as neon.correct
BF⁻ has 8 electrons — the same electron arrangement as oxygen.
This option is wrong — you subtracted an electron instead of adding one — a 1− charge means one electron GAINED, and oxygen is not a noble gas.
CF⁻ has 9 electrons — fluorine's own arrangement, unchanged by bonding.
This option is wrong — you ignored the charge — the 1− superscript means one extra electron, taking fluorine from 9 electrons to 10.
DF⁻ has 18 electrons — the same electron arrangement as argon.
This option is wrong — you jumped to the next noble gas down — fluorine needs only one extra electron to reach neon's 10, not nine extras to reach argon's 18.
A 1− charge means one electron gained. F⁻ has 9 + 1 = 10 electrons. Ten electrons is neon's arrangement, so F⁻ fits the octet pattern.
Lesson 3 of 55 · IMB-003
Metal ion charges from group number
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen that a metal atom loses electrons to reach a noble-gas arrangement. For the metals of Groups 1, 2, and 13, the periodic table tells you the resulting charge directly.
The idea
Find the metal's group in the periodic table — the group tells you the ion's charge: Group 1 → 1+, Group 2 → 2+, Group 13 → 3+.
A Group 1 metal forms a 1+ ion.
Metal ion charges read straight off the group: Group 1 → 1+, Group 2 → 2+, Group 13 → 3+.
A Group 2 metal forms a 2+ ion.
A Group 13 metal forms a 3+ ion.
Sodium is in Group 1, so sodium forms Na⁺.
Magnesium is in Group 2, so magnesium forms Mg²⁺.
Aluminum is in Group 13, so aluminum forms Al³⁺.
Worked examples
Worked example 1. Potassium is in Group 1. What ion does potassium form?
Step 1
Group 1 metals form 1+ ions.
Step 2
Potassium forms K⁺.
Worked example 2. Gallium is in Group 13. What ion does gallium form?
Step 1
Group 13 metals form 3+ ions — the charge is 3+, not 13+.
Step 2
Gallium forms Ga³⁺.
You can now predict the charge of the ion a Group 1, Group 2, or Group 13 metal forms — 1+, 2+, and 3+ in turn — from its group number.
Check your understanding
Barium is a metal in Group 2. What ion does barium form?
ABa²⁺correct
BBa⁺
This option is wrong — you gave the Group 1 charge — a Group 2 metal forms a 2+ ion.
CBa³⁺
This option is wrong — you gave the Group 13 charge — a Group 2 metal forms a 2+ ion.
DBa²⁻
This option is wrong — you flipped the sign — metals lose electrons, so their ions are positive.
Find the group: barium is in Group 2. A Group 2 metal forms a 2+ ion. So barium forms Ba²⁺.
Check your understanding
Cesium is a metal in Group 1. What ion does cesium form?
ACs⁺correct
BCs²⁺
This option is wrong — you gave the Group 2 charge — a Group 1 metal forms a 1+ ion.
CCs⁻
This option is wrong — you flipped the sign — metals lose electrons, so their ions are positive.
DCs³⁺
This option is wrong — you gave the Group 13 charge — a Group 1 metal forms a 1+ ion.
Find the group: cesium is in Group 1. A Group 1 metal forms a 1+ ion. So cesium forms Cs⁺.
Check your understanding
Indium is a metal in Group 13. What ion does indium form?
AIn³⁺correct
BIn¹³⁺
This option is wrong — you used the group number 13 itself as the charge — a Group 13 metal loses three electrons, forming a 3+ ion.
CIn³⁻
This option is wrong — you flipped the sign — metals lose electrons, so their ions are positive.
DIn⁺
This option is wrong — you gave the Group 1 charge — a Group 13 metal forms a 3+ ion.
Find the group: indium is in Group 13. A Group 13 metal forms a 3+ ion — the 13 in the group name never becomes the charge. So indium forms In³⁺.
Lesson 4 of 55 · IMB-004
Nonmetal ion charges from group number
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen that a nonmetal atom gains electrons to reach a noble-gas arrangement, and that Group 1, 2, and 13 metals take their charge from their group. The nonmetal groups have their own read-off pattern.
The idea
Find the nonmetal's group in the periodic table, and count how far it sits from Group 18.
A Group 15 nonmetal forms a 3− ion.
Nonmetal ion charges read straight off the group: Group 15 → 3−, Group 16 → 2−, Group 17 → 1−.
A Group 16 nonmetal forms a 2− ion.
A Group 17 nonmetal forms a 1− ion.
Nitrogen is in Group 15, so nitrogen forms N³⁻.
Oxygen is in Group 16, so oxygen forms O²⁻.
Chlorine is in Group 17, so chlorine forms Cl⁻.
Worked examples
Worked example 1. Phosphorus is in Group 15. What ion does phosphorus form?
Step 1
Group 15 nonmetals form 3− ions.
Step 2
Phosphorus forms P³⁻.
Worked example 2. Bromine is in Group 17. What ion does bromine form?
Step 1
Group 17 nonmetals form 1− ions — the charge is 1−, not 17−.
Step 2
Bromine forms Br⁻.
You can now predict the charge of the ion a Group 15, Group 16, or Group 17 nonmetal forms — 3−, 2−, and 1− in turn — from its group number.
Check your understanding
Sulfur is a nonmetal in Group 16. What ion does sulfur form?
AS²⁻correct
BS²⁺
This option is wrong — you flipped the sign — nonmetals gain electrons, so their ions are negative.
CS⁻
This option is wrong — you gave the Group 17 charge — a Group 16 nonmetal forms a 2− ion.
DS⁶⁻
This option is wrong — you counted sulfur's six valence electrons as the charge — the charge comes from the group's read-off pattern, and Group 16 gives 2−.
Find the group: sulfur is in Group 16. A Group 16 nonmetal forms a 2− ion. So sulfur forms S²⁻.
Check your understanding
Iodine is a nonmetal in Group 17. What ion does iodine form?
AI⁻correct
BI²⁻
This option is wrong — you gave the Group 16 charge — a Group 17 nonmetal forms a 1− ion.
CI⁺
This option is wrong — you flipped the sign — nonmetals gain electrons, so their ions are negative.
DI¹⁷⁻
This option is wrong — you used the group number 17 itself as the charge — a Group 17 nonmetal gains one electron, forming a 1− ion.
Find the group: iodine is in Group 17. A Group 17 nonmetal forms a 1− ion. So iodine forms I⁻.
Check your understanding
Arsenic is a nonmetal in Group 15. What ion does arsenic form?
AAs³⁻correct
BAs³⁺
This option is wrong — you flipped the sign — nonmetals gain electrons, so their ions are negative.
CAs⁵⁻
This option is wrong — you counted arsenic's five valence electrons as the charge — the read-off pattern gives Group 15 a 3− charge.
DAs⁻
This option is wrong — you gave the Group 17 charge — a Group 15 nonmetal forms a 3− ion.
Find the group: arsenic is in Group 15. A Group 15 nonmetal forms a 3− ion. So arsenic forms As³⁻.
Lesson 5 of 55 · IMB-005
Transition metals form more than one ion
Learning Journey
Read · Check · Unlock
Did You Know?
The group read-off rules you've just used cover Groups 1, 2, and 13–17. The transition metals — the wide middle block of the periodic table — play by a different rule.
The idea
Many transition metals form more than one stable ion.
Iron forms both Fe²⁺ and Fe³⁺.
Copper forms both Cu⁺ and Cu²⁺.
One metal, more than one ion
Transition metal
Stable ions
iron
Fe²⁺ and Fe³⁺
copper
Cu⁺ and Cu²⁺
Many transition metals form more than one stable ion.
Because one metal can form two different ions, a transition metal's charge cannot be read off its group number.
When you need a transition metal's charge, the compound it sits in has to tell you — later lessons show you how.
Worked examples
Worked example 1. Iron forms Fe²⁺ and one other common stable ion. What is the other ion?
Step 1
Answer: Fe³⁺.
Worked example 2. Name the two common stable ions of copper.
Step 1
Answer: Cu⁺ and Cu²⁺.
You can now state that many transition metals form more than one stable ion.
Check your understanding
Which pair shows the two common stable ions of iron?
AFe²⁺ and Fe³⁺correct
BFe⁺ and Fe²⁺
This option is wrong — you recalled copper's charge pattern for iron — iron's pair is Fe²⁺ and Fe³⁺.
CFe²⁻ and Fe³⁻
This option is wrong — you flipped the signs — iron is a metal, and metal ions are positive.
DFe³⁺ and Fe⁴⁺
This option is wrong — you shifted both charges up by one — iron's pair is Fe²⁺ and Fe³⁺.
Iron is one of the transition metals that forms more than one stable ion. Iron forms Fe²⁺ and Fe³⁺.
Check your understanding
Which pair shows the two common stable ions of copper?
ACu⁺ and Cu²⁺correct
BCu²⁺ and Cu³⁺
This option is wrong — you recalled iron's charge pattern for copper — copper's pair is Cu⁺ and Cu²⁺.
CCu⁻ and Cu²⁻
This option is wrong — you flipped the signs — copper is a metal, and metal ions are positive.
DCu⁺ and Cu³⁺
This option is wrong — you skipped over copper's second ion — the pair is Cu⁺ and Cu²⁺, one charge apart.
Copper is one of the transition metals that forms more than one stable ion. Copper forms Cu⁺ and Cu²⁺.
Check your understanding
Can a transition metal's ion charge be read off its group number, the way a Group 2 metal's can?
ANo — many transition metals form more than one stable ion, so no single charge follows from the group.correct
BYes — every metal's ion charge equals its group number, wherever the metal sits in the table.
This option is wrong — you stretched the read-off rule over the whole table — it covers Groups 1, 2, and 13, and transition metals form more than one stable ion.
CNo — transition metals never form ions at all, so the question of a charge does not arise.
This option is wrong — you turned 'unpredictable charge' into 'no ions' — transition metals form ions readily; the issue is that many form more than one.
DYes — every transition metal forms exactly a 2+ ion, so one charge covers the whole block.
This option is wrong — you promoted one common charge to a rule — iron also forms Fe³⁺ and copper also forms Cu⁺, so no single charge covers them.
The group read-off rule covers Groups 1, 2, and 13. Many transition metals form more than one stable ion — iron forms Fe²⁺ and Fe³⁺, copper forms Cu⁺ and Cu²⁺. So a transition metal's charge cannot be read off its group number.
Lesson 6 of 55 · IMB-006
The ionic bond
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen that atoms bond because bonded atoms end up with lower energy, and that metals form positive ions while nonmetals form negative ones. Put a positive ion next to a negative ion, and the bond appears.
The idea
A positive ion and a negative ion pull toward each other, because opposite charges attract, and the attraction is stronger when the charges are larger or closer together.
That pull between oppositely charged ions is a chemical bond.
The ionic bond: oppositely charged ions pull toward each other.
A bond of this kind is called an 'ionic bond' — the attraction between oppositely charged ions.
In sodium chloride, the ionic bond is the attraction between the Na⁺ ions and the Cl⁻ ions.
Pulled together this way, the ions sit at lower energy than they have apart — which is why the bond forms and holds.
Worked examples
Worked example 1. Solid potassium bromide contains K⁺ ions and Br⁻ ions. What is the ionic bond in potassium bromide?
Step 1
An ionic bond is the attraction between oppositely charged ions.
Step 2
The oppositely charged ions here are K⁺ and Br⁻.
Step 3
The ionic bond in potassium bromide is the attraction between the K⁺ ions and the Br⁻ ions.
Worked example 2. A Mg²⁺ ion sits the same distance from an O²⁻ ion as a Na⁺ ion sits from a Cl⁻ ion. Which pair pulls together more strongly?
Step 1
The attraction is stronger when the charges are larger or closer together.
Step 2
The distances are equal, so compare the charges: 2+ and 2− are larger than 1+ and 1−.
Step 3
The Mg²⁺ and O²⁻ pair pulls together more strongly, because its charges are larger.
You can now state that an ionic bond is the attraction between oppositely charged ions.
Check your understanding
Solid lithium fluoride contains Li⁺ ions and F⁻ ions. Which pair of particles pulls toward each other?
AA Li⁺ ion and an F⁻ ion.correct
BA Li⁺ ion and another Li⁺ ion.
This option is wrong — you paired like charges — two positive ions repel each other; attraction needs opposite charges.
CAn F⁻ ion and another F⁻ ion.
This option is wrong — you paired like charges — two negative ions repel each other; attraction needs opposite charges.
DNo pair — ions in a solid neither attract nor repel.
This option is wrong — you switched the charges off — the ions keep their charges in the solid, and opposite charges attract.
Opposite charges attract, and the attraction is stronger when the charges are larger or closer together. Li⁺ is positive and F⁻ is negative, so this pair pulls together. That pull is the ionic bond in lithium fluoride.
Check your understanding
Solid calcium sulfide contains Ca²⁺ ions and S²⁻ ions. What is the ionic bond in calcium sulfide?
AThe attraction between the Ca²⁺ ions and the S²⁻ ions.correct
BThe repulsion between the Ca²⁺ ions and the S²⁻ ions.
This option is wrong — you swapped attraction for repulsion — opposite charges attract; repulsion happens between like charges.
CThe attraction between one Ca²⁺ ion and another Ca²⁺ ion.
This option is wrong — you paired like charges — two positive ions repel; the bond is the pull between the oppositely charged Ca²⁺ and S²⁻.
DThe noble-gas electron arrangement that each ion has reached.
This option is wrong — you named the octet pattern as the bond — the arrangement describes the ions' electrons; the bond is the attraction between the oppositely charged ions.
An ionic bond is the attraction between oppositely charged ions. In calcium sulfide the oppositely charged ions are Ca²⁺ and S²⁻. Their mutual pull is the bond.
Check your understanding
Two ions are attracting each other. Which single change makes their attraction stronger?
AMoving the two ions closer together.correct
BMoving the two ions farther apart.
This option is wrong — you reversed the distance effect — the attraction is stronger when the charges are closer together.
CMaking both ions' charges smaller.
This option is wrong — you reversed the charge effect — the attraction is stronger when the charges are larger.
DSwapping both ions' signs at once.
This option is wrong — you changed nothing that matters — after swapping both signs the pair is still one positive and one negative charge at the same distance, so the attraction is unchanged.
Opposite charges attract, and the attraction is stronger when the charges are larger or closer together. Closing the distance strengthens the pull. Growing the charges would strengthen it too; shrinking them or separating the ions weakens it.
Lesson 7 of 55 · IMB-007
Transfer makes the ions; attraction makes the bond
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen what the ionic bond is: the attraction between oppositely charged ions. But before that attraction can exist, the ions themselves have to come from somewhere.
The idea
The ions come from a transfer: a metal atom transfers one or more electrons to a nonmetal atom.
Sodium transfers one electron to chlorine, forming Na⁺ and Cl⁻.
Transfer makes the ions; attraction makes the bond.
The transfer explains how the ions FORM — and that is all it explains.
The ionic bond itself is the attraction between the oppositely charged ions, not the transfer.
Keep the two jobs separate: transfer makes the ions; attraction makes the bond.
Worked examples
Worked example 1. A magnesium atom meets an oxygen atom. Describe how the ions form, and then say what the ionic bond between them is.
Step 1
The metal atom transfers electrons to the nonmetal atom: magnesium transfers two electrons to oxygen.
Step 2
The transfer forms the ions Mg²⁺ and O²⁻.
Step 3
The ionic bond is the attraction between the oppositely charged Mg²⁺ and O²⁻ ions — not the transfer.
Step 4
Transfer of two electrons makes the ions; the attraction between Mg²⁺ and O²⁻ is the bond.
Worked example 2. In lithium bromide, lithium transferred one electron to bromine. Which part of the story is the ionic bond?
Step 1
The transfer formed the ions Li⁺ and Br⁻ — that part is over once the ions exist.
Step 2
What remains, and what holds the compound together, is the attraction between the oppositely charged ions.
Step 3
The ionic bond is the attraction between the Li⁺ and Br⁻ ions.
You can now explain that electron transfer from a metal atom to a nonmetal atom explains how the ions form, and that the ionic bond itself is the attraction between the oppositely charged ions, not the transfer.
Check your understanding
Which statement correctly separates the two jobs when a metal and a nonmetal form an ionic bond?
AElectron transfer makes the ions; attraction between the ions is the bond.correct
BAttraction between the atoms makes the ions; electron transfer is the bond.
This option is wrong — you swapped the two jobs — the transfer creates the charged ions, and the attraction between those ions is the bond.
CElectron transfer both makes the ions and is itself the bond.
This option is wrong — you let the transfer play both parts — once the ions exist the transfer is over, and what holds them together is the attraction.
DAttraction between the ions both makes the ions and is the bond.
This option is wrong — you let attraction play both parts — attraction cannot create charges; the ions come from electron transfer.
Two different jobs, two different actors. The transfer of electrons from the metal atom to the nonmetal atom makes the ions. The attraction between the oppositely charged ions is the bond.
Check your understanding
In calcium sulfide, calcium transferred two electrons to sulfur, forming Ca²⁺ and S²⁻. What is the ionic bond in calcium sulfide?
AThe attraction between the Ca²⁺ and S²⁻ ions.correct
BThe movement of the two electrons from calcium to sulfur.
This option is wrong — you called the transfer the bond — the transfer only explains how the ions formed; the bond is the attraction between them.
CThe two electrons themselves, sitting between the ions.
This option is wrong — you placed the transferred electrons between the ions — they moved fully onto the sulfur, and the bond is the attraction between the resulting charges.
DThe noble-gas arrangement each ion reached after the transfer.
This option is wrong — you named the octet pattern as the bond — the arrangement describes the ions' electrons; the bond is the attraction between the oppositely charged ions.
The transfer formed the ions Ca²⁺ and S²⁻ — that job is done. The ionic bond is the attraction between the oppositely charged ions. Transfer makes the ions; attraction makes the bond.
Check your understanding
In potassium fluoride, which event does the electron transfer explain?
AHow the K⁺ and F⁻ ions formed.correct
BWhy the K⁺ and F⁻ ions attract each other.
This option is wrong — you stretched the transfer over the attraction — the ions attract because their charges are opposite, whatever made the charges.
CWhat holds the K⁺ and F⁻ ions together.
This option is wrong — you made the transfer the glue — the attraction between the oppositely charged ions is what holds them together.
DWhy potassium fluoride is electrically neutral overall.
This option is wrong — you reached past the transfer's job — the transfer explains where the K⁺ and F⁻ ions came from; the compound's overall neutrality is a charge-balance story, taken up in a later lesson.
The transfer's one job is to explain where the ions came from: potassium transferred one electron to fluorine, forming K⁺ and F⁻. The attraction between the oppositely charged ions is the bond that holds them together. Transfer makes the ions; attraction makes the bond.
Lesson 8 of 55 · IMB-008
The crystal lattice
Learning Journey
Read · Check · Unlock
Have You Ever Wondered?
Wonder this:
A single grain of table salt holds about a billion billion ions. How do that many charged particles arrange themselves inside one tiny cube?
You've already seen that oppositely charged ions attract. Inside a solid, that attraction produces a strikingly tidy result.
The idea
A compound made of ions is called an 'ionic compound'.
In a solid ionic compound, the ions do not sit in loose pairs or a random jumble.
They pack into a 'crystal lattice' — a repeating three-dimensional arrangement in which each ion is surrounded by ions of the opposite charge.
In solid sodium chloride, every Na⁺ ion sits among Cl⁻ neighbours, and every Cl⁻ ion sits among Na⁺ neighbours.
A flat fragment of the sodium chloride lattice. Each ion sits among neighbours of the opposite charge, and the pattern repeats in every direction.
The same alternating pattern repeats over and over, in every direction, out to the edges of the crystal.
The figure shows a flat fragment of the pattern — the real lattice extends up, down, and outward in three dimensions, far beyond any diagram.
Worked examples
Worked example 1. Solid magnesium oxide packs Mg²⁺ and O²⁻ ions into a crystal lattice. Which ions surround each O²⁻ ion?
Step 1
In a crystal lattice, each ion is surrounded by ions of the opposite charge.
Step 2
The opposite charge to O²⁻ here is Mg²⁺.
Step 3
Each O²⁻ ion is surrounded by Mg²⁺ ions.
Worked example 2. A diagram of solid lithium fluoride shows a 4 × 4 fragment of its lattice. Does the crystal end where the diagram ends?
Step 1
A lattice diagram shows a fragment of a repeating arrangement.
Step 2
The real pattern continues in every direction, in three dimensions, to the edges of the crystal.
Step 3
No — the drawn fragment is a small sample of a pattern that keeps repeating far beyond the diagram.
You can now state that the ions in a solid ionic compound (a compound made of ions) pack into a crystal lattice — a repeating three-dimensional arrangement in which each ion is surrounded by ions of the opposite charge.
Check your understanding
Solid potassium fluoride contains K⁺ and F⁻ ions packed in a crystal lattice. Which particles sit directly around each K⁺ ion?
AF⁻ ions.correct
BOther K⁺ ions.
This option is wrong — you surrounded a positive ion with positives — in a lattice each ion is surrounded by ions of the OPPOSITE charge.
CA random mix of K⁺ and F⁻ ions with no pattern.
This option is wrong — you made the lattice a jumble — the arrangement repeats in a strict alternating pattern, opposite charges adjacent.
DNeutral potassium and fluorine atoms.
This option is wrong — you filled the solid with neutral atoms — a solid ionic compound is made of ions, and the neighbours of K⁺ are the F⁻ ions.
A crystal lattice is a repeating arrangement in which each ion is surrounded by ions of the opposite charge. K⁺ is positive, so its immediate neighbours are the negative F⁻ ions. The same holds in reverse: each F⁻ sits among K⁺ neighbours.
Check your understanding
Which description fits the arrangement of ions in any solid ionic compound?
AA repeating three-dimensional pattern with each ion surrounded by oppositely charged ions.correct
BSeparate pairs, each one positive ion stuck to one negative ion.
This option is wrong — you broke the solid into private pairs — no ion has a single partner; each one is surrounded by several oppositely charged neighbours.
CA random pile of positive and negative ions with no repeating pattern.
This option is wrong — you removed the order — the ions pack into a strict repeating pattern, which is why the solid forms crystals.
DA flat, single-layer sheet of alternating ions.
This option is wrong — you stopped at the diagram — the drawn fragment is flat, but the real lattice repeats in three dimensions.
The ions pack into a crystal lattice. A crystal lattice is a repeating three-dimensional arrangement in which each ion is surrounded by ions of the opposite charge. Diagrams show a flat fragment of it; the real pattern fills the crystal.
Check your understanding
The figure shows a fragment of the lattice of solid calcium oxide, with one empty site marked X between four labeled ions. Which particle belongs at X?
AAn O²⁻ ion.correct
BA Ca²⁺ ion.
This option is wrong — you placed a positive ion among positive neighbours — the four ions around X are Ca²⁺, and each site must be surrounded by the opposite charge.
CA neutral oxygen atom.
This option is wrong — you inserted a neutral atom — the lattice is built of ions, and the site among Ca²⁺ neighbours takes the negative ion O²⁻.
DEither ion — the lattice accepts any particle at any site.
This option is wrong — you let the pattern go — the alternating arrangement is strict: every site's neighbours carry the opposite charge.
Read the neighbours of X in the figure: all four are Ca²⁺. Each ion in a lattice is surrounded by ions of the opposite charge. So the site at X takes the negative ion, O²⁻.
Lesson 9 of 55 · IMB-009
Why the lattice holds together
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen the lattice: every ion surrounded by neighbours of the opposite charge. That arrangement is also the answer to a natural question — why do the ions stay put instead of drifting apart?
The idea
Pick any single ion in the lattice and count what pulls on it.
Every one of its oppositely charged neighbours attracts it — from above, below, left, right, in front, and behind.
The ion is not held by one partner; it is held by attractions acting in every direction at once.
One ion, many holders: oppositely charged neighbours attract it from every direction.
A solid ionic compound holds together because the ions are held in a repeating lattice by attraction between opposite charges acting in every direction.
No single attraction does the job alone — the whole lattice shares the work of holding every ion in place.
Worked examples
Worked example 1. Why do the ions in solid lithium fluoride stay in place instead of drifting apart?
Step 1
Each Li⁺ ion is surrounded by F⁻ neighbours, and each F⁻ ion by Li⁺ neighbours.
Step 2
Every neighbour attracts the ion between them, so the pulls come from every direction at once.
Step 3
Held by all those attractions together, no ion can drift away.
Step 4
The ions stay in place because they are held in a repeating lattice by attraction between opposite charges acting in every direction.
Worked example 2. In solid calcium oxide, one single Ca²⁺–O²⁻ attraction is not what holds the crystal together. What does?
Step 1
Each ion sits among several oppositely charged neighbours, and every one of them pulls on it.
Step 2
The attractions act in every direction at once, throughout the whole lattice.
Step 3
The crystal is held together by the many attractions between opposite charges acting in every direction — the lattice as a team, not any single pair.
You can now explain that a solid ionic compound holds together because the ions are held in a repeating lattice by attraction between opposite charges acting in every direction.
Check your understanding
Why does solid potassium bromide hold together?
ABecause the ions are held in a repeating lattice by attraction between opposite charges acting in every direction.correct
BBecause each K⁺ ion is bonded to exactly one Br⁻ ion, forming a pair that cannot separate.
This option is wrong — you gave each ion a single partner — every ion is held by several oppositely charged neighbours at once, not by one.
CBecause the ions have reached noble-gas electron arrangements, which lock them in place.
This option is wrong — you used the octet pattern as glue — the arrangement describes the ions' electrons; what holds the solid together is the attraction between opposite charges.
DBecause the ions are packed so tightly that there is no room for any of them to move.
This option is wrong — you made crowding the cause — the ions stay put because attractions pull them from every direction, not because space runs out.
Pick any ion in the lattice and count what pulls on it. Every oppositely charged neighbour attracts it, from every direction at once. The solid holds together because the ions are held in a repeating lattice by attraction between opposite charges acting in every direction.
Check your understanding
In solid magnesium sulfide, how many directions do the attractions holding one Mg²⁺ ion come from?
AEvery direction — each of the ion's oppositely charged neighbours pulls on it.correct
BOne direction — from its single partner ion.
This option is wrong — you gave the ion one partner — in a lattice each ion sits among several oppositely charged neighbours, and all of them pull.
CTwo directions — from the left and the right.
This option is wrong — you read the flat diagram as the whole story — the lattice is three-dimensional, so neighbours also pull from above, below, in front, and behind.
DNo direction — attractions switch off once the solid has formed.
This option is wrong — you switched the charges off — the ions keep their charges in the solid, and the attractions keep acting.
Each ion in the lattice is surrounded by oppositely charged neighbours. Every neighbour attracts it, so the pulls come from every direction at once. Those ever-present attractions are what hold each ion in place.
Check your understanding
Pulling one Rb⁺ ion away from a single Cl⁻ partner would take far less effort than pulling it out of a rubidium chloride crystal. Why?
AEach ion is held by many attractions acting in every direction, so the lattice shares the work of holding every ion.correct
BThe attraction between each ion pair inside the crystal is far stronger than the attraction in a lone pair.
This option is wrong — you strengthened each attraction — a pair's attraction is only ONE attraction; in the lattice the ion is held by attractions from every direction at once.
CThe ions merge into neutral particles inside the crystal, and neutral particles resist separation.
This option is wrong — you neutralised the ions — they keep their charges in the solid, and those charges are exactly what the holding attractions act on.
DThe crystal's outer surface forms a sealed shell that protects the ions inside.
This option is wrong — you invented a shell — there is no special surface layer; every ion, inside and out, is held by attractions from its oppositely charged neighbours.
A lone partner holds an ion with just one attraction; the crystal holds it with many. In the lattice, every ion is pulled by all of its oppositely charged neighbours at once. The ions are held in a repeating lattice by attraction between opposite charges acting in every direction — the whole lattice shares the work.
Lesson 10 of 55 · IMB-010
Draw an ionic lattice fragment
Learning Journey
Read · Check · Unlock
Did You Know?
You've already read lattice diagrams and seen why the arrangement holds together. Now you draw a lattice fragment yourself.
The idea
Before you draw, choose two circle styles — one per ion — that differ in size AND shading.
Label every circle with its ion's symbol and charge, such as Na⁺ or Cl⁻.
Place the ions in a grid so that every neighbour — above, below, left, and right — carries the opposite charge.
Never let two like charges sit next to each other.
The drawing target: two distinct circle styles, every ion labeled, opposite charges adjacent everywhere.
The circles may touch: in a lattice drawing, touching shows how the ions pack.
Fill the whole fragment with the same alternating pattern, so the drawing reads as a piece of something that repeats.
Worked examples
Worked example 1. Draw a 4 × 4 fragment of the potassium bromide lattice, using K⁺ and Br⁻ ions.
Step 1
Choose the circle styles: K⁺ as small light circles, Br⁻ as large dark circles.
Step 2
Start the first row alternating: K⁺, Br⁻, K⁺, Br⁻.
Step 3
Start the second row on the OTHER ion — Br⁻, K⁺, Br⁻, K⁺ — so every up-down neighbour pair is also opposite.
Step 4
Repeat rows three and four the same way, and label every circle with its symbol and charge.
Step 5
A 4 × 4 alternating grid: no two like charges adjacent, two visibly different circle styles, every ion labeled.
Worked example 2. Draw a 4 × 4 fragment of the magnesium oxide lattice, using Mg²⁺ and O²⁻ ions.
Step 1
The charges are 2+ and 2−, but the drawing rule is unchanged: opposite charges adjacent everywhere.
Step 2
Alternate Mg²⁺ and O²⁻ along each row, and start each new row on the other ion.
Step 3
Label every circle — the labels carry the 2+ and 2− charges.
Step 4
A 4 × 4 alternating grid of labeled Mg²⁺ and O²⁻ circles, no like charges adjacent.
You can now draw a fragment of an ionic lattice as a repeating arrangement that places each ion next to oppositely charged neighbours, never like charge beside like charge.
Check your understanding
Panels A to D each show a student's drawing of an ionic lattice fragment. Which panel shows a correct fragment?
APanel Acorrect
BPanel B
This option is wrong — you grouped like charges into blocks — every neighbour in a lattice drawing must carry the opposite charge.
CPanel C
This option is wrong — you drew separate ion pairs — a lattice is one continuous alternating grid, not floating pairs.
DPanel D
This option is wrong — you drew the alternation but made the ions indistinguishable — the two ions must differ in size and shading and carry their charge labels.
Check three things: alternation, distinct circle styles, and labels. Panel A alternates opposite charges everywhere, uses two visibly different circle styles, and labels every ion. Blocks of like charges, floating pairs, and unlabeled identical circles each break one of the drawing rules.
Your turn
On paper, draw a 4 × 4 fragment of the lithium fluoride lattice, using Li⁺ and F⁻ ions. Use two circle styles that differ in size and shading, and label every ion. Then select Continue to compare your drawing with the model answer.
Model answer. A 4 × 4 grid of alternating circles: small light circles labeled Li⁺ and large dark circles labeled F⁻. Each row alternates the two ions, and each new row starts on the other ion, so every up-down and left-right neighbour pair is oppositely charged. The circles touch.
Two circle styles appear, differing in size and shading — one for Li⁺, one for F⁻.
Every circle is labeled with its symbol and charge.
Every neighbour pair — up-down and left-right — is one positive and one negative ion.
The alternating pattern fills the whole 4 × 4 fragment, with no like charges adjacent anywhere.
Your turn
On paper, draw a 4 × 4 fragment of the calcium oxide lattice, using Ca²⁺ and O²⁻ ions. Use two circle styles that differ in size and shading, and label every ion with its charge. Then select Continue to compare your drawing with the model answer.
Model answer. A 4 × 4 grid of alternating circles: small light circles labeled Ca²⁺ and large dark circles labeled O²⁻. Rows alternate the two ions and consecutive rows start on opposite ions, so no like charges are adjacent. The circles touch.
Two circle styles appear, differing in size and shading — one for Ca²⁺, one for O²⁻.
Every circle is labeled, and the labels carry the full charges 2+ and 2−.
Every neighbour pair — up-down and left-right — is one positive and one negative ion.
The alternating pattern fills the whole 4 × 4 fragment.
Formulas like H₂O appear on labels and throughout science. For an ionic compound, a formula is a precise count — and it reads by one simple rule.
The idea
A 'chemical formula' lists the ions in a compound using their element symbols and small lowered numbers.
The small lowered number after a symbol is a 'subscript'.
A subscript counts the ion directly to its LEFT.
No subscript means a count of one.
A subscript counts the ion directly to its left; no subscript means one.
CaCl₂ reads as one Ca²⁺ ion for every two Cl⁻ ions.
The 2 belongs to Cl, the symbol on its left — not to Ca.
Worked examples
Worked example 1. What ratio of ions does the formula Na₂S represent?
Step 1
The subscript 2 sits after Na, so it counts the Na⁺ ions.
Step 2
S carries no subscript, so its count is one.
Step 3
Na₂S represents two Na⁺ ions for every one S²⁻ ion.
Worked example 2. What ratio of ions does the formula AlBr₃ represent?
Step 1
Al carries no subscript, so its count is one.
Step 2
The subscript 3 sits after Br, so it counts the Br⁻ ions.
Step 3
AlBr₃ represents one Al³⁺ ion for every three Br⁻ ions.
You can now identify the ratio of ions represented by a chemical formula, reading each subscript as the count of the ion to its left and no subscript as a count of one.
Check your understanding
What ratio of ions does the formula K₂O represent?
ATwo K⁺ ions for every one O²⁻ ion.correct
BOne K⁺ ion for every two O²⁻ ions.
This option is wrong — you attached the subscript to the symbol on its right — a subscript counts the ion directly to its LEFT, and the 2 sits after K.
CTwo K⁺ ions for every two O²⁻ ions.
This option is wrong — you spread the subscript over both symbols — the 2 counts only the K on its left, and O with no subscript counts as one.
DOne K⁺ ion for every one O²⁻ ion.
This option is wrong — you ignored the subscript — the 2 after K means two K⁺ ions.
The subscript 2 sits after K, so it counts the K⁺ ions. O carries no subscript, so its count is one. K₂O: two K⁺ ions for every one O²⁻ ion.
Check your understanding
What ratio of ions does the formula MgF₂ represent?
AOne Mg²⁺ ion for every two F⁻ ions.correct
BTwo Mg²⁺ ions for every one F⁻ ion.
This option is wrong — you attached the subscript to the symbol on its right — the 2 sits after F, so it counts the F⁻ ions.
COne Mg²⁺ ion for every one F⁻ ion.
This option is wrong — you ignored the subscript — the 2 after F means two F⁻ ions.
DTwo Mg²⁺ ions for every two F⁻ ions.
This option is wrong — you spread the subscript over both symbols — the 2 counts only the F on its left, and Mg with no subscript counts as one.
Mg carries no subscript, so its count is one. The subscript 2 sits after F, so it counts the F⁻ ions. MgF₂: one Mg²⁺ ion for every two F⁻ ions.
Check your understanding
What ratio of ions does the formula Li₃N represent?
AThree Li⁺ ions for every one N³⁻ ion.correct
BOne Li⁺ ion for every three N³⁻ ions.
This option is wrong — you attached the subscript to the symbol on its right — the 3 sits after Li, so it counts the Li⁺ ions.
CThree Li⁺ ions for every three N³⁻ ions.
This option is wrong — you spread the subscript over both symbols — the 3 counts only the Li on its left, and N with no subscript counts as one.
DOne Li⁺ ion for every one N³⁻ ion.
This option is wrong — you ignored the subscript — the 3 after Li means three Li⁺ ions.
The subscript 3 sits after Li, so it counts the Li⁺ ions. N carries no subscript, so its count is one. Li₃N: three Li⁺ ions for every one N³⁻ ion.
Lesson 12 of 55 · IMB-012
Formula unit, not molecule
Learning Journey
Read · Check · Unlock
Have You Ever Wondered?
Wonder this:
The formula H₂O names a real particle: any single water molecule is exactly two hydrogen atoms and one oxygen atom. So does NaCl name a real two-atom particle of salt?
You've already seen the lattice: every ion surrounded by oppositely charged neighbours, with no ion belonging to any single partner.
The idea
Search the sodium chloride lattice for a private Na–Cl pair and you will not find one.
The lattice contains no separate molecules — every Na⁺ touches several Cl⁻ neighbours, none of them its own.
So the formula of an ionic compound cannot describe a molecule.
Instead, the formula gives a 'formula unit' — the simplest whole-number ratio of ions in the lattice.
NaCl means a 1:1 ratio of Na⁺ to Cl⁻ ions throughout the crystal, not a two-atom particle.
The dashed oval marks one formula unit of NaCl: a 1:1 counting ratio. The two ions inside it are not a separate particle — each also touches other neighbours.
A formula unit is a counting statement about the whole lattice — you cannot pick one up, because it is a ratio, not a thing.
Worked examples
Worked example 1. What does the formula MgO of solid magnesium oxide tell you, and what does it NOT tell you?
Step 1
MgO gives the formula unit: a 1:1 ratio of Mg²⁺ to O²⁻ ions throughout the crystal.
Step 2
It does not describe a two-atom MgO particle, because the lattice contains no separate molecules.
Step 3
MgO states the 1:1 ion ratio of the lattice — it names no discrete particle.
Worked example 2. Solid potassium sulfide has the formula K₂S. What does its formula unit say about the crystal?
Step 1
The formula unit is the simplest whole-number ratio of ions in the lattice.
Step 2
K₂S says the lattice holds two K⁺ ions for every one S²⁻ ion, throughout the crystal.
Step 3
Everywhere in the crystal, K⁺ outnumbers S²⁻ two to one — and no three-ion K₂S particle exists.
You can now explain that the formula of an ionic compound gives a formula unit — the simplest whole-number ratio of ions in the lattice — and does not describe a molecule, because the lattice contains no separate molecules.
Check your understanding
Solid lithium bromide has the formula LiBr. Which statement about the crystal is correct?
ALiBr gives the 1:1 ratio of Li⁺ to Br⁻ ions in the lattice — the crystal contains no separate LiBr particles.correct
BThe crystal is a stack of two-atom LiBr molecules.
This option is wrong — you read the formula as a particle — the lattice contains no separate molecules; LiBr states a ratio of ions.
CEach Li⁺ ion in the crystal is permanently attached to one particular Br⁻ ion.
This option is wrong — you gave each ion a private partner — every ion touches several oppositely charged neighbours, none of them its own.
DLiBr describes one region of the crystal where a lithium ion and a bromide ion happen to sit together.
This option is wrong — you shrank the formula to one spot — the 1:1 ratio holds throughout the whole crystal, not in one place.
The formula of an ionic compound gives a formula unit — the simplest whole-number ratio of ions in the lattice. LiBr means one Li⁺ for every Br⁻, everywhere in the crystal. No separate LiBr particle exists anywhere in the lattice.
Check your understanding
What is a formula unit?
AThe simplest whole-number ratio of ions in an ionic compound's lattice.correct
BThe smallest particle of an ionic compound that can exist on its own.
This option is wrong — you turned the ratio into a particle — a formula unit is a counting statement about the lattice, not a thing you could pick up.
CA molecule made of one positive ion and one negative ion.
This option is wrong — you called it a molecule — the lattice of an ionic compound contains no molecules at all.
DThe total number of ions in one crystal.
This option is wrong — you counted the whole crystal — the formula unit is the simplest RATIO of ions, not a total count.
A formula unit is the simplest whole-number ratio of ions in the lattice. It is a count, not a particle. NaCl, for instance, means a 1:1 ratio of Na⁺ to Cl⁻ throughout the crystal.
Check your understanding
Solid calcium fluoride has the formula CaF₂. What does the formula unit CaF₂ state about the crystal?
AThe lattice holds one Ca²⁺ ion for every two F⁻ ions, throughout the crystal.correct
BThe crystal is built from three-atom CaF₂ particles.
This option is wrong — you read the formula as a particle — no three-atom CaF₂ unit exists; the formula states the 1:2 ion ratio of the lattice.
CEach Ca²⁺ ion in the lattice touches exactly two F⁻ ions and no others.
This option is wrong — you turned the ratio into a touching rule — the formula counts ions across the lattice; it does not say how many neighbours touch each ion.
DOne region of the crystal holds calcium ions and a different region holds fluoride ions, in a 1:2 split.
This option is wrong — you separated the ions into regions — the ions alternate through one shared lattice, and the 1:2 ratio holds everywhere in it.
The formula unit is the simplest whole-number ratio of ions in the lattice. CaF₂ means one Ca²⁺ for every two F⁻, everywhere in the crystal. It is a counting statement — not a particle, and not a map of who touches whom.
Lesson 13 of 55 · IMB-013
Ionic compounds are neutral
Learning Journey
Read · Check · Unlock
Did You Know?
A crystal of an ionic compound is packed with charged particles. Yet a salt shaker neither attracts your hand nor gives you a shock — the crystal as a whole carries no charge.
The idea
An ionic compound is electrically neutral overall.
Its total positive charge equals its total negative charge, so the two cancel.
The formula shows the balance: count each ion's charge as many times as the ion appears.
In MgO, one Mg²⁺ balances one O²⁻: (2+) + (2−) = 0.
In Li₂O, two Li⁺ balance one O²⁻: 2 × (1+) + (2−) = 0.
Whatever the compound, the totals always cancel to zero.
Worked examples
Worked example 1. Calcium bromide is CaBr₂, made of Ca²⁺ and Br⁻ ions. Show that it is electrically neutral.
Step 1
Total positive charge: 1 × (2+) = 2+.
Step 2
Total negative charge: 2 × (1−) = 2−.
Step 3
The totals cancel: (2+) + (2−) = 0.
Step 4
CaBr₂ is neutral — its total positive and negative charges cancel to zero.
Worked example 2. Potassium sulfide is K₂S, made of K⁺ and S²⁻ ions. Show that it is electrically neutral.
Step 1
Total positive charge: 2 × (1+) = 2+.
Step 2
Total negative charge: 1 × (2−) = 2−.
Step 3
The totals cancel: (2+) + (2−) = 0.
Step 4
K₂S is neutral — two 1+ charges balance one 2− charge.
You can now state that an ionic compound is electrically neutral overall because its total positive charge equals its total negative charge.
Check your understanding
What is the overall electric charge of any ionic compound?
AZero — the total positive charge equals the total negative charge.correct
BPositive — compounds contain more metal ions than nonmetal ions.
This option is wrong — you let the positives win — the compound contains exactly enough of each ion for the charge totals to cancel.
CNegative — the gained electrons give the compound extra negative charge.
This option is wrong — you counted the transferred electrons twice — every electron a nonmetal gained, a metal lost, so the totals still cancel.
DIt depends on the compound — some ionic compounds are charged overall.
This option is wrong — you allowed exceptions — neutrality is universal: every ionic compound's positive and negative totals cancel to zero.
An ionic compound is electrically neutral overall. Its total positive charge equals its total negative charge. The two totals cancel to zero in every ionic compound.
Check your understanding
Sodium sulfide is Na₂S, made of Na⁺ and S²⁻ ions. Which line shows why it is electrically neutral?
A2 × (1+) + (2−) = 0correct
B(1+) + (2−) = 0
This option is wrong — you counted only one Na⁺ — the subscript 2 means two sodium ions, so the positive total is 2 × (1+).
C2 × (1+) + 2 × (2−) = 0
This option is wrong — you doubled the sulfide too — S carries no subscript, so there is one S²⁻, and 2+ with 4− would not cancel.
D2 × (2+) + (2−) = 0
This option is wrong — you gave sodium a 2+ charge — sodium is a Group 1 metal, so each ion is 1+, and the subscript supplies the count of two.
Count each ion's charge as many times as the ion appears. Two Na⁺ give 2 × (1+) = 2+; one S²⁻ gives 2−. The totals cancel: 2 × (1+) + (2−) = 0.
Check your understanding
Aluminum chloride is AlCl₃, made of Al³⁺ and Cl⁻ ions. Why does the compound carry no overall charge?
AOne 3+ charge is cancelled by three 1− charges: (3+) + 3 × (1−) = 0.correct
BThe aluminum ion's 3+ charge fades away once the ions pack into the solid.
This option is wrong — you let the charges fade — every ion keeps its full charge in the solid; the compound is neutral because the totals cancel.
CEach Cl⁻ cancels the Al³⁺ completely, so the compound is neutral three times over.
This option is wrong — you let one 1− cancel a 3+ — cancelling a 3+ charge takes all three of the 1− charges together.
DThe compound is neutral because it contains equal NUMBERS of positive and negative ions.
This option is wrong — you balanced counts instead of charges — AlCl₃ has one positive ion and three negative ions; what balance are the charge TOTALS: (3+) + 3 × (1−) = 0.
Neutrality is about charge totals, not ion counts. Total positive: 1 × (3+) = 3+. Total negative: 3 × (1−) = 3−. (3+) + 3 × (1−) = 0, so AlCl₃ carries no overall charge.
Lesson 14 of 55 · IMB-014
Write a formula by balancing charges
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen that an ionic compound contains exactly enough of each ion for the charges to cancel. That fact turns formula writing into a short routine.
The idea
Start from the two ions of a two-element — 'binary' — ionic compound.
Write the positive ion's symbol first, then the negative ion's symbol.
Take each ion's charge NUMBER, ignoring the sign.
Each ion's charge number becomes the OTHER ion's subscript — the numbers criss-cross.
A subscript of 1 is never written.
This 'criss-cross method' works because the crossed counts make the positive and negative charge totals cancel.
For Al³⁺ and O²⁻: the 3 crosses to O, the 2 crosses to Al, giving Al₂O₃.
Check the totals: 2 × (3+) + 3 × (2−) = 0.
Worked examples
Worked example 1. Write the formula of the ionic compound of Mg²⁺ and Cl⁻.
Step 1
Positive ion first: Mg, then Cl.
Step 2
Charge numbers, signs ignored: 2 and 1.
Step 3
Criss-cross: the 2 crosses to Cl, the 1 crosses to Mg.
Step 4
A subscript of 1 is never written, so Mg keeps no subscript.
Step 5
MgCl₂
Worked example 2. Write the formula of the ionic compound of Li⁺ and N³⁻.
Step 1
Positive ion first: Li, then N.
Step 2
Charge numbers, signs ignored: 1 and 3.
Step 3
Criss-cross: the 1 crosses to N, the 3 crosses to Li.
Step 4
A subscript of 1 is never written, so N keeps no subscript.
Step 5
Li₃N
You can now write the formula of a two-element (binary) ionic compound from its two ions, choosing subscripts so the positive and negative charges cancel — each ion's charge number becomes the other ion's subscript (the criss-cross method).
Check your understanding
Write the formula of the compound formed from K⁺ ions and S²⁻ ions.
Accepted answer: K₂S
Charge numbers, signs ignored: 1 for K⁺ and 2 for S²⁻. Criss-cross: the 2 crosses to K, the 1 crosses to S and goes unwritten. K₂S — check: 2 × (1+) + (2−) = 0.
Check your understanding
Which formula is correct for the ionic compound of Ba²⁺ and I⁻?
ABaI₂correct
BBa₂I
This option is wrong — you criss-crossed in reverse — barium's 2 becomes IODINE's subscript, not barium's own.
CBaI
This option is wrong — you ignored the charges — one 1− iodide cannot balance a 2+ barium; the 2 must cross to iodine.
DBa₂I₂
This option is wrong — you gave both ions a subscript 2 — only barium's charge number crosses to iodine; iodide's 1 goes unwritten on barium.
Charge numbers, signs ignored: 2 for Ba²⁺ and 1 for I⁻. Criss-cross: the 2 crosses to I, the 1 crosses to Ba and goes unwritten. BaI₂ — check: (2+) + 2 × (1−) = 0.
Your turn
On paper, write the formula of the ionic compound of Al³⁺ and F⁻, showing the criss-cross step. Then select Continue to compare your work with the model answer.
Model answer. Positive ion first: Al, then F. Charge numbers, signs ignored: 3 and 1. Criss-cross: the 3 crosses to F, the 1 crosses to Al and goes unwritten. Formula: AlF₃. Check: (3+) + 3 × (1−) = 0.
The positive ion's symbol is written first.
Each charge number crossed to the OTHER ion's subscript position.
No subscript of 1 appears anywhere.
The finished formula is AlF₃, and the charge totals cancel: (3+) + 3 × (1−) = 0.
Lesson 15 of 55 · IMB-015
Reduce subscripts to the simplest ratio
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen the criss-cross method. For some ion pairs it hands back subscripts that are technically balanced but not finished.
The idea
Run the criss-cross for Ca²⁺ and O²⁻ and you get Ca₂O₂.
The charges do cancel — but the formula is still wrong.
A formula gives the SIMPLEST whole-number ratio of ions, because it is a ratio, not a particle count.
So after criss-crossing, check the subscripts: if both can be divided by the same whole number, divide them.
In Ca₂O₂ both subscripts divide by 2, giving CaO.
Write Ca²⁺ and O²⁻ as CaO, never Ca₂O₂.
Worked examples
Worked example 1. Write the formula of the ionic compound of Mg²⁺ and S²⁻.
Step 1
Criss-cross: the 2 crosses each way, giving Mg₂S₂.
Step 2
Check the subscripts: both divide by 2.
Step 3
Divide: 2 ÷ 2 = 1 for each ion, and a subscript of 1 is never written.
Step 4
MgS
Worked example 2. Write the formula of the ionic compound of Al³⁺ and N³⁻.
Step 1
Criss-cross: the 3 crosses each way, giving Al₃N₃.
Step 2
Check the subscripts: both divide by 3.
Step 3
Divide: 3 ÷ 3 = 1 for each ion, and a subscript of 1 is never written.
Step 4
AlN
You can now write the formula of a binary ionic compound whose criss-crossed subscripts must be reduced to the simplest whole-number ratio.
Check your understanding
Write the formula of the compound formed from Ba²⁺ ions and S²⁻ ions.
Accepted answer: BaS
Criss-cross: Ba₂S₂. Both subscripts divide by 2 — divide them. BaS: one 2+ balances one 2−.
Check your understanding
A student criss-crosses Sr²⁺ and O²⁻ and writes Sr₂O₂. What is the correct final formula, and why?
ASrO — a formula gives the simplest whole-number ratio, and both subscripts divide by 2.correct
BSr₂O₂ — whatever the criss-cross method hands back is always the final formula.
This option is wrong — you skipped the reduction check — the charges cancel in Sr₂O₂, but the formula must show the SIMPLEST ratio, which is 1:1.
CSrO₂ — only the metal's subscript reduces, because the metal is written first.
This option is wrong — you reduced one subscript alone — dividing must apply to BOTH subscripts, or the ratio changes and the charges no longer cancel.
DSr₂O — only the nonmetal's subscript reduces, because it sits at the end.
This option is wrong — you reduced one subscript alone — dividing must apply to BOTH subscripts, or the ratio changes and the charges no longer cancel.
Criss-crossing 2 and 2 gives Sr₂O₂ — balanced, but not finished. A formula gives the simplest whole-number ratio of ions. Divide both subscripts by 2: SrO.
Check your understanding
Which formula is correct for the ionic compound of Ca²⁺ and S²⁻?
ACaScorrect
BCa₂S₂
This option is wrong — you stopped after criss-crossing — both subscripts divide by 2, so the formula reduces to CaS.
CCaS₂
This option is wrong — you crossed calcium's 2 to sulfur but dropped sulfide's 2 — the two charge numbers cross both ways and then reduce to 1:1.
DCa₂S₄
This option is wrong — you doubled instead of reducing — the criss-cross gives 2 and 2, and the only allowed change afterwards is dividing both by their shared divisor.
Criss-cross the two 2s: Ca₂S₂. Both subscripts divide by 2. CaS — check: (2+) + (2−) = 0.
Lesson 16 of 55 · IMB-016
Check a formula by charge balance
Learning Journey
Read · Check · Unlock
Did You Know?
Writing formulas is one direction. Just as useful is the reverse skill: looking at a formula someone hands you and judging whether it can be right.
The idea
The test is the neutrality principle you've already seen: the charge totals must cancel.
Multiply each ion's charge by its count in the formula.
Add the positive total and the negative total.
If the sum is zero, the formula is consistent with its ions.
If the sum is anything other than zero, the formula is wrong.
For MgCl₂ from Mg²⁺ and Cl⁻: (2+) + 2 × (1−) = 0 — the formula checks out.
Worked examples
Worked example 1. A worksheet gives K₂S for the ions K⁺ and S²⁻. Check the formula by totalling the charges.
Step 1
Total positive charge: 2 × (1+) = 2+.
Step 2
Total negative charge: 1 × (2−) = 2−.
Step 3
Sum: 2 × (1+) + (2−) = 0.
Step 4
The sum is zero, so K₂S is correct for K⁺ and S²⁻.
Worked example 2. A worksheet gives CaCl₃ for the ions Ca²⁺ and Cl⁻. Check the formula by totalling the charges.
Step 1
Total positive charge: 1 × (2+) = 2+.
Step 2
Total negative charge: 3 × (1−) = 3−.
Step 3
Sum: (2+) + 3 × (1−) = 1−.
Step 4
The sum is 1−, not zero, so CaCl₃ is wrong — the correct formula is CaCl₂.
You can now judge whether a chemical formula is correct for a given pair of ions by totalling the positive and negative charges and checking that they cancel.
Check your understanding
A worksheet gives AlCl₂ for the ions Al³⁺ and Cl⁻. Total the charges. Is the formula correct?
ANo — the sum is (3+) + 2 × (1−) = 1+, not zero.correct
BYes — the sum is (3+) + 2 × (1−) = 0.
This option is wrong — you called 3+ and 2− a cancellation — the totals differ by one whole charge, so the sum is 1+, and the formula fails the check.
CYes — a formula is correct whenever it contains both ions.
This option is wrong — you skipped the check entirely — containing the right ions is not enough; their charge totals must cancel to zero.
DNo — the sum is (3+) + 2 × (1−) = 1−.
This option is wrong — you got the right verdict with the wrong arithmetic — 3+ combined with 2− leaves 1+ (a positive excess), not 1−.
Total positive charge: 1 × (3+) = 3+. Total negative charge: 2 × (1−) = 2−. Sum: (3+) + 2 × (1−) = 1+ — not zero, so AlCl₂ is wrong; the correct formula is AlCl₃.
Check your understanding
A worksheet gives BaBr₂ for the ions Ba²⁺ and Br⁻. Total the charges. Is the formula correct?
AYes — the sum is (2+) + 2 × (1−) = 0.correct
BNo — the sum is (2+) + 2 × (1−) = 2−.
This option is wrong — you left the barium out of the total — one Ba²⁺ contributes 2+, which exactly cancels the two bromides' 2−.
CNo — the formula has more bromide ions than barium ions, so it cannot be neutral.
This option is wrong — you judged by ion counts — neutrality is about charge totals, and one 2+ ion balances two 1− ions exactly.
DYes — but only because the charges of the ions vanish inside a solid.
This option is wrong — you credited the right verdict to vanished charges — the ions keep their charges; the formula is right because the totals cancel: (2+) + 2 × (1−) = 0.
Total positive charge: 1 × (2+) = 2+. Total negative charge: 2 × (1−) = 2−. Sum: (2+) + 2 × (1−) = 0 — the formula checks out.
Check your understanding
A worksheet gives NaO for the ions Na⁺ and O²⁻. Total the charges. Is the formula correct?
ANo — the sum is (1+) + (2−) = 1−, so the correct formula needs two Na⁺: Na₂O.correct
BYes — the sum is (1+) + (2−) = 0.
This option is wrong — you cancelled unequal totals — 1+ against 2− leaves 1− uncancelled, so NaO fails the check.
CNo — the sum is (1+) + (2−) = 1+, so the formula needs another oxide ion.
This option is wrong — you flipped the leftover's sign — the negative total is the larger one, so the excess is 1− and the fix is MORE sodium, not more oxide.
DYes — one symbol of each element always makes a correct formula.
This option is wrong — you defaulted to 1:1 — the ion charges set the ratio, and a 1+ ion cannot balance a 2− ion one to one.
Total positive charge: 1 × (1+) = 1+. Total negative charge: 1 × (2−) = 2−. Sum: (1+) + (2−) = 1− — not zero, so NaO is wrong; two Na⁺ are needed, giving Na₂O.
Lesson 17 of 55 · IMB-017
Polyatomic ions
Learning Journey
Read · Check · Unlock
Did You Know?
Every ion you've met so far has been a single atom carrying a charge — Na⁺, Cl⁻, O²⁻. Ions come in a second variety.
The idea
A 'polyatomic ion' is a group of atoms bonded together that carries one overall charge.
The whole group moves as a single unit — it forms, travels, and settles into lattices in one piece.
The charge belongs to the whole group, not to any single atom inside it.
The nitrate ion NO₃⁻ is one 1− unit made of one nitrogen atom and three oxygen atoms.
Read the formula like any other: subscripts count the atoms, and the raised charge at the end belongs to the whole unit.
The nitrate ion: four bonded atoms, one overall 1− charge belonging to the whole unit.
Worked examples
Worked example 1. The carbonate ion has the formula CO₃²⁻. How many atoms does one carbonate ion contain, and what is its charge?
Step 1
Count the atoms: one carbon and three oxygens — four atoms in the group.
Step 2
The raised 2− at the end is the charge of the whole unit.
Step 3
One carbonate ion is a four-atom unit carrying one overall 2− charge.
Worked example 2. The sulfate ion has the formula SO₄²⁻. Which atom in the unit carries the 2− charge?
Step 1
The charge of a polyatomic ion belongs to the whole group, not to any single atom.
Step 2
The five bonded atoms — one sulfur and four oxygens — share one overall 2− charge.
Step 3
No single atom carries it — the 2− belongs to the whole sulfate unit.
You can now state that a polyatomic ion is a group of atoms bonded together that carries one overall charge and moves as a single unit.
Check your understanding
The phosphate ion has the formula PO₄³⁻. How many atoms make up one phosphate ion, and what is its charge?
AFive atoms, carrying one overall 3− charge.correct
BFour atoms, carrying one overall 3− charge.
This option is wrong — you forgot the atom with no subscript — the P counts as one, and with four O atoms the unit holds five.
CFive atoms, each carrying its own 3− charge.
This option is wrong — you copied the charge onto every atom — the 3− belongs to the whole unit once, not to each atom.
DFive atoms with no charge, because groups of atoms are neutral molecules.
This option is wrong — you filed the group as a molecule — a molecule is neutral, but this group carries an overall 3− charge, which makes it a polyatomic ion.
Count the atoms: one P and four O — five atoms. The raised 3− at the end belongs to the whole unit. A group of bonded atoms with one overall charge is a polyatomic ion.
Check your understanding
Which of these particles is a polyatomic ion?
AClO₃⁻ — a group of bonded atoms carrying one overall charge.correct
BBr⁻ — a single particle carrying a clear negative charge.
This option is wrong — you stopped at 'charged' — Br⁻ is a single atom with a charge; a polyatomic ion needs a GROUP of bonded atoms.
CO₂ — a group of atoms bonded firmly to one another.
This option is wrong — you stopped at 'group of atoms' — O₂ carries no overall charge, which makes it a neutral molecule, not an ion.
DNaCl — a compound whose formula joins two element symbols.
This option is wrong — you picked a neutral compound — NaCl is a compound of two ions with charges that cancel, not a single charged unit.
A polyatomic ion needs both features at once: a group of bonded atoms AND one overall charge. ClO₃⁻ has four bonded atoms and an overall 1− charge. A lone charged atom, a neutral molecule, and a neutral compound each lack one of the two features.
Check your understanding
The sulfite ion has the formula SO₃²⁻. In this ion, where does the 2− charge sit?
AOn the whole four-atom unit, shared as one overall charge.correct
BOn the sulfur atom, because it is written first.
This option is wrong — you pinned the charge to the first atom — the charge of a polyatomic ion belongs to the whole group, not to any single atom.
COn each oxygen atom, giving three separate 2− charges.
This option is wrong — you copied the charge onto every oxygen — the raised 2− is written once and belongs to the whole unit once.
DNowhere — the atoms in a bonded group lose their charge.
This option is wrong — you erased the charge — the group genuinely carries 2−; that overall charge is exactly what makes it an ion.
The raised charge at the end of a polyatomic formula belongs to the whole unit. SO₃²⁻ is one 2− unit made of one sulfur atom and three oxygen atoms. No single atom inside owns the charge.
Lesson 18 of 55 · IMB-018
Common polyatomic ions: ammonium, hydroxide, nitrate
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen what a polyatomic ion is. A handful of them appear constantly in this course, and their names, formulas, and charges are worth committing to memory now. Here are the first three.
The idea
The 'ammonium' ion is NH₄⁺ — one nitrogen, four hydrogens, charge 1+.
Ammonium is the odd one out among common polyatomic ions: it is positive.
Three polyatomic ions
Name
Formula
Charge
ammonium
NH₄⁺
1+
hydroxide
OH⁻
1−
nitrate
NO₃⁻
1−
Ammonium, hydroxide, nitrate — the first three of the course's common polyatomic ions.
The 'hydroxide' ion is OH⁻ — one oxygen, one hydrogen, charge 1−.
Hydroxide is the smallest of the common polyatomic ions: just two atoms.
The 'nitrate' ion is NO₃⁻ — one nitrogen, three oxygens, charge 1−.
Ammonium and nitrate both contain nitrogen; the hydrogens mark ammonium, the oxygens mark nitrate.
Worked examples
Worked example 1. What is the formula and charge of the hydroxide ion?
Step 1
Answer: OH⁻, charge 1−.
Worked example 2. Which polyatomic ion has the formula NH₄⁺?
Step 1
Answer: ammonium, charge 1+.
You can now state the name, formula, and charge of the ammonium ion NH₄⁺, the hydroxide ion OH⁻, and the nitrate ion NO₃⁻.
Check your understanding
What is the formula and charge of the ammonium ion?
ANH₄⁺, charge 1+correct
BNH₄⁻, charge 1−
This option is wrong — you flipped the sign — ammonium is the positive one of the common polyatomic ions.
CNH₃⁺, charge 1+
This option is wrong — you dropped a hydrogen — ammonium carries FOUR hydrogens: NH₄⁺.
DNO₃⁻, charge 1−
This option is wrong — you recalled nitrate — the other nitrogen-containing ion; the hydrogens mark ammonium, NH₄⁺.
Ammonium is NH₄⁺ — one nitrogen, four hydrogens. Its charge is 1+, the odd one out among the common polyatomic ions.
Check your understanding
Which polyatomic ion has the formula OH⁻?
AHydroxidecorrect
BAmmonium
This option is wrong — you matched the wrong name — ammonium is NH₄⁺; the two-atom OH⁻ is hydroxide.
CNitrate
This option is wrong — you matched the wrong name — nitrate is NO₃⁻; the two-atom OH⁻ is hydroxide.
DNone of the above — OH⁻ is not a polyatomic ion.
This option is wrong — you required more atoms — two bonded atoms with one overall charge already make a polyatomic ion, and OH⁻ is hydroxide.
OH⁻ is the hydroxide ion — one oxygen, one hydrogen, charge 1−. It is the smallest of the common polyatomic ions.
Check your understanding
What is the formula and charge of the nitrate ion?
ANO₃⁻, charge 1−correct
BNO₃²⁻, charge 2−
This option is wrong — you doubled the charge — nitrate carries 1−.
CNH₄⁺, charge 1+
This option is wrong — you recalled ammonium — the other nitrogen-containing ion; the oxygens mark nitrate, NO₃⁻.
DNO₂⁻, charge 1−
This option is wrong — you dropped an oxygen — nitrate carries THREE oxygens: NO₃⁻.
Nitrate is NO₃⁻ — one nitrogen, three oxygens. Its charge is 1−.
Lesson 19 of 55 · IMB-019
Common polyatomic ions: carbonate, sulfate, phosphate
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen three polyatomic ions — ammonium NH₄⁺, hydroxide OH⁻, and nitrate NO₃⁻. Here are the next three, and these carry charges bigger than 1−.
The idea
The 'carbonate' ion is CO₃²⁻ — one carbon atom and three oxygen atoms sharing a 2− charge.
Carbonate is the 2− ion in chalk, limestone, and seashells.
The 'sulfate' ion is SO₄²⁻ — one sulfur atom and four oxygen atoms sharing a 2− charge.
Sulfate is the 2− ion in plaster and in Epsom salt.
The 'phosphate' ion is PO₄³⁻ — one phosphorus atom and four oxygen atoms sharing a 3− charge.
Phosphate is the 3− ion in fertilizer and in your bones.
Notice the pattern in this trio: carbonate and sulfate each carry 2−, and phosphate carries 3−.
Each ion moves as one unit and keeps its charge — exactly as you've already seen with nitrate and hydroxide.
Worked examples
Worked example 1. What is the formula of the sulfate ion, including its charge?
Step 1
SO₄²⁻
Worked example 2. Which polyatomic ion has the formula PO₄³⁻?
Step 1
The phosphate ion.
You can now state the name, formula, and charge of the carbonate ion CO₃²⁻, the sulfate ion SO₄²⁻, and the phosphate ion PO₄³⁻.
Check your understanding
Write the formula of the carbonate ion, including its charge.
Accepted answer: CO₃²⁻
The carbonate ion is one carbon atom and three oxygen atoms. Its charge is 2−. Write it: CO₃²⁻.
Check your understanding
A fertilizer label lists the ion PO₄³⁻. What is this ion's name?
APhosphate.correct
BSulfate.
This option is wrong — you matched the four oxygen atoms, but the central atom here is phosphorus — sulfate is SO₄²⁻, with sulfur.
CCarbonate.
This option is wrong — you picked the ion with three oxygen atoms and a 2− charge — carbonate is CO₃²⁻, with carbon.
DNitrate.
This option is wrong — you picked a 1− ion with three oxygen atoms — nitrate is NO₃⁻, with nitrogen.
PO₄³⁻ has phosphorus at its center and four oxygen atoms. The ion with phosphorus and a 3− charge is phosphate.
Check your understanding
What is the charge on the sulfate ion? Write the number with its sign.
Accepted answer: 2−
The sulfate ion is SO₄²⁻. The superscript 2− is its charge: 2−.
Three more polyatomic ions complete the roster you'll use for formulas and names. All three carry a 1− charge.
The idea
The 'hydrogen carbonate' ion is HCO₃⁻ — a carbonate ion carrying one extra hydrogen atom, with a 1− charge.
Hydrogen carbonate is the ion in baking soda.
The 'acetate' ion is C₂H₃O₂⁻ — two carbon, three hydrogen, and two oxygen atoms sharing a 1− charge.
Acetate is the ion in the salt inside reusable heat packs.
The 'chlorate' ion is ClO₃⁻ — one chlorine atom and three oxygen atoms sharing a 1− charge.
Chlorate is the oxygen-rich ion in match heads.
All three carry 1−, like nitrate and hydroxide.
Watch the lookalikes: HCO₃⁻ is carbonate plus one hydrogen, and its charge drops from carbonate's 2− to 1−.
ClO₃⁻ has the same shape as nitrate's formula — chlorine sits where nitrate has nitrogen.
Worked examples
Worked example 1. What is the formula of the acetate ion, including its charge?
Step 1
C₂H₃O₂⁻
Worked example 2. Which polyatomic ion has the formula ClO₃⁻?
Step 1
The chlorate ion.
You can now state the name, formula, and charge of the hydrogen carbonate ion HCO₃⁻, the acetate ion C₂H₃O₂⁻, and the chlorate ion ClO₃⁻.
Check your understanding
Write the formula of the hydrogen carbonate ion, including its charge.
Accepted answer: HCO₃⁻
Hydrogen carbonate is a carbonate ion carrying one extra hydrogen atom. Its charge is 1−. Write it: HCO₃⁻.
Check your understanding
A weed-killer label lists the ion ClO₃⁻. What is this ion's name?
AChlorate.correct
BNitrate.
This option is wrong — you matched the three oxygen atoms and the 1− charge, but the central atom here is chlorine — nitrate is NO₃⁻.
CCarbonate.
This option is wrong — carbonate is CO₃²⁻ — carbon at the center and a 2− charge.
DAcetate.
This option is wrong — acetate is C₂H₃O₂⁻ — it contains carbon and hydrogen, and ClO₃⁻ contains neither.
ClO₃⁻ is one chlorine atom and three oxygen atoms with a 1− charge. That ion is chlorate.
Check your understanding
What is the charge on the acetate ion? Write the number with its sign.
Accepted answer: 1−
The acetate ion is C₂H₃O₂⁻. The superscript minus with no number means a charge of 1−.
Lesson 21 of 55 · IMB-021
Spot the polyatomic ion in a formula
Learning Journey
Read · Check · Unlock
Did You Know?
You've now seen nine polyatomic ions. Real formulas mix them with single-atom ions, so the first skill is spotting the group inside a formula.
The idea
A polyatomic ion appears inside a formula as its whole block of atoms, written together.
The polyatomic ion sits in the formula as one whole block.
In KNO₃, the block NO₃ is the nitrate ion — the K in front is a potassium ion.
In Na₂CO₃, the block CO₃ is the carbonate ion.
When a formula needs more than one of a polyatomic ion, the block sits inside parentheses — in Ca(OH)₂, the OH block is the hydroxide ion.
A polyatomic ion can also lead the formula: in NH₄Cl, the NH₄ block is the ammonium ion.
To spot the ion, scan the formula for a block you've already seen — the leftover symbols belong to the partner ion.
Charges are never written inside a compound's formula — you recall the charge from the ion itself.
Worked examples
Worked example 1. Which polyatomic ion is present in Na₃PO₄?
Step 1
Scan for a familiar block: PO₄ is the phosphate block.
Step 2
The leftover symbols, Na₃, are three sodium ions.
Step 3
The phosphate ion, PO₄³⁻.
Worked example 2. Which polyatomic ion is present in Sr(OH)₂?
Step 1
The parentheses wrap the block OH — the hydroxide ion.
Step 2
The Sr in front is a strontium ion.
Step 3
The hydroxide ion, OH⁻.
Worked example 3. Which polyatomic ion is present in Li₂SO₄?
Step 1
The block SO₄ is the sulfate ion.
Step 2
The leftover Li₂ is two lithium ions.
Step 3
The sulfate ion, SO₄²⁻.
You can now identify the polyatomic ion present in a chemical formula.
Check your understanding
Which polyatomic ion is present in MgSO₄?
ASulfate.correct
BCarbonate.
This option is wrong — you matched an oxygen-carrying block, but carbonate is CO₃ — this block's central atom is sulfur.
CPhosphate.
This option is wrong — phosphate is PO₄ — the block here starts with S, not P.
DHydroxide.
This option is wrong — hydroxide is OH — this formula has no OH block.
Scan MgSO₄ for a familiar block: SO₄ is the sulfate block. The leftover Mg is a magnesium ion.
Check your understanding
Which polyatomic ion is present in NaHCO₃? Write the ion's name.
Accepted answer: hydrogen carbonate
The block after Na is HCO₃ — hydrogen carbonate. The leftover Na is a sodium ion.
Check your understanding
Which polyatomic ion is present in Al(NO₃)₃?
ANitrate.correct
BCarbonate.
This option is wrong — carbonate is CO₃, with carbon — the block in the parentheses is NO₃, with nitrogen.
CChlorate.
This option is wrong — chlorate is ClO₃, with chlorine — the block here starts with N.
DAmmonium.
This option is wrong — ammonium is NH₄ and leads a formula — the block here is NO₃, nitrogen with oxygen.
The parentheses wrap the block NO₃ — the nitrate ion. The Al in front is an aluminum ion.
Lesson 22 of 55 · IMB-022
Formulas with parentheses
Learning Journey
Read · Check · Unlock
Did You Know?
You've already written formulas for two-element compounds by balancing charges. The same balancing runs the show when a polyatomic ion is involved — with one new writing rule.
The idea
Magnesium forms Mg²⁺ and nitrate is NO₃⁻, so one 2+ ion needs two 1− ions before the charges cancel.
Two whole nitrate ions balance one Mg²⁺ — parentheses keep each block intact.
The formula must show two whole nitrate ions — but a bare 2 written after the O would read as two oxygen atoms only.
So the polyatomic ion goes inside parentheses, and the count goes outside: Mg(NO₃)₂.
Parentheses mean the whole block repeats — Mg(NO₃)₂ contains one magnesium ion and two complete nitrate ions.
Use parentheses only when the formula contains more than one of that polyatomic ion.
Sodium nitrate needs just one nitrate ion for one Na⁺, so it is written NaNO₃ — no parentheses.
The check step is unchanged: for Mg(NO₃)₂, (2+) + 2 × (1−) = 0.
Worked examples
Worked example 1. Write the formula of the compound formed from Ca²⁺ ions and OH⁻ ions.
Step 1
Write the ions: Ca²⁺ and OH⁻.
Step 2
Balance the charges: one 2+ ion needs two 1− ions.
Step 3
More than one hydroxide ion, so parentheses: Ca(OH)₂.
Step 4
Ca(OH)₂
Worked example 2. Write the formula of the compound formed from NH₄⁺ ions and S²⁻ ions.
Step 1
Write the ions: NH₄⁺ and S²⁻.
Step 2
Balance the charges: two 1+ ions cancel one 2− ion.
Step 3
The repeated ion is the polyatomic one, so the parentheses wrap NH₄: (NH₄)₂S.
Step 4
(NH₄)₂S
Worked example 3. Write the formula of the compound formed from Sr²⁺ ions and NO₃⁻ ions.
Step 1
Write the ions: Sr²⁺ and NO₃⁻.
Step 2
Balance the charges: one 2+ ion needs two 1− ions.
Step 3
Two nitrate ions, so parentheses: Sr(NO₃)₂.
Step 4
Sr(NO₃)₂
You can now write the formula of an ionic compound that needs parentheses because it contains more than one of a polyatomic ion.
Check your understanding
Write the formula of the compound formed from Ba²⁺ ions and NO₃⁻ ions.
Accepted answer: Ba(NO₃)₂
Write the ions: Ba²⁺ and NO₃⁻. Balance the charges: one 2+ ion needs two 1− ions. Two nitrate ions, so parentheses: Ba(NO₃)₂. Check: (2+) + 2 × (1−) = 0.
Check your understanding
Write the formula of the compound formed from Mg²⁺ ions and OH⁻ ions.
Accepted answer: Mg(OH)₂
Write the ions: Mg²⁺ and OH⁻. Balance the charges: one 2+ ion needs two 1− ions. Two hydroxide ions, so parentheses: Mg(OH)₂. Check: (2+) + 2 × (1−) = 0.
Check your understanding
Write the formula of the compound formed from NH₄⁺ ions and SO₄²⁻ ions.
Accepted answer: (NH₄)₂SO₄
Write the ions: NH₄⁺ and SO₄²⁻. Balance the charges: two 1+ ions cancel one 2− ion. The repeated ion is polyatomic, so parentheses wrap NH₄: (NH₄)₂SO₄. Check: 2 × (1+) + (2−) = 0.
Lesson 23 of 55 · IMB-023
Count atoms through parentheses
Learning Journey
Read · Check · Unlock
Did You Know?
A formula with parentheses packs a lot of atoms into a small space. Counting them correctly is a two-step read.
The idea
The subscript outside parentheses multiplies every atom inside them.
The 2 outside the parentheses multiplies every atom inside.
In Ca(NO₃)₂, each nitrate block holds one nitrogen atom and three oxygen atoms, and the 2 outside doubles the whole block.
So Ca(NO₃)₂ contains one calcium atom, 2 × 1 = 2 nitrogen atoms, and 2 × 3 = 6 oxygen atoms.
Atoms outside the parentheses are not multiplied — the Ca counts once.
A formula with no parentheses reads as you've already seen: Na₂SO₄ has two sodium atoms, one sulfur atom, and four oxygen atoms.
Worked examples
Worked example 1. How many atoms of each element are in Ba(OH)₂?
Step 1
Inside the parentheses: one oxygen atom and one hydrogen atom.
Step 2
The 2 outside doubles both: 2 oxygen atoms and 2 hydrogen atoms.
Step 3
The Ba outside the parentheses counts once.
Step 4
1 barium, 2 oxygen, and 2 hydrogen atoms
Worked example 2. How many atoms of each element are in Al₂(SO₄)₃?
Step 1
Inside the parentheses: one sulfur atom and four oxygen atoms.
Step 2
The 3 outside triples both: 3 sulfur atoms and 3 × 4 = 12 oxygen atoms.
Step 3
The Al₂ outside the parentheses counts as written: 2 aluminum atoms.
Step 4
2 aluminum, 3 sulfur, and 12 oxygen atoms
Worked example 3. How many atoms of each element are in Mg₃(PO₄)₂?
Step 1
Inside the parentheses: one phosphorus atom and four oxygen atoms.
Step 2
The 2 outside doubles both: 2 phosphorus atoms and 2 × 4 = 8 oxygen atoms.
Step 3
The Mg₃ outside counts as written: 3 magnesium atoms.
Step 4
3 magnesium, 2 phosphorus, and 8 oxygen atoms
You can now calculate the number of atoms of each element in a formula containing parentheses by multiplying each inner subscript by the outer subscript.
Check your understanding
How many oxygen atoms are in one formula unit of Sr(NO₃)₂?
Answer: 6atoms
Inside the parentheses each nitrate block holds 3 oxygen atoms. The 2 outside doubles the block: 2 × 3 = 6. Sr(NO₃)₂ contains 6 oxygen atoms.
Check your understanding
How many hydrogen atoms are in one formula unit of Al(OH)₃?
Answer: 3atoms
Each hydroxide block holds 1 hydrogen atom. The 3 outside triples the block: 3 × 1 = 3. Al(OH)₃ contains 3 hydrogen atoms.
Check your understanding
How many oxygen atoms are in one formula unit of Ca₃(PO₄)₂?
Answer: 8atoms
Each phosphate block holds 4 oxygen atoms. The 2 outside doubles the block: 2 × 4 = 8. Ca₃(PO₄)₂ contains 8 oxygen atoms.
Lesson 24 of 55 · IMB-024
Write any ionic formula from its ions
Learning Journey
Read · Check · Unlock
Did You Know?
You've already balanced charges, reduced subscripts, and used parentheses — each in its own lesson. This lesson chains them into one routine that writes any ionic formula from its ions.
The idea
Step 1 — write both ions with their charges.
Step 2 — find the smallest counts of each ion that make the charges cancel.
The criss-cross shortcut gets you there fast: each ion's charge number becomes the other ion's subscript.
Step 3 — if the two subscripts share a common factor, reduce them to the simplest ratio.
Step 4 — if the formula contains more than one of a polyatomic ion, wrap that ion in parentheses.
Step 5 — check: the total positive charge and the total negative charge must cancel.
Watch the routine run on Ca²⁺ and PO₄³⁻.
Criss-cross: the 2 from Ca²⁺ becomes phosphate's count, and the 3 from PO₄³⁻ becomes calcium's — three calcium ions and two phosphate ions.
3 and 2 share no common factor, and phosphate repeats, so it takes parentheses: Ca₃(PO₄)₂.
Check: 3 × (2+) + 2 × (3−) = (6+) + (6−) = 0.
Worked examples
Worked example 1. Write the formula of the compound formed from Ba²⁺ ions and O²⁻ ions.
Step 1
Write the ions: Ba²⁺ and O²⁻.
Step 2
Criss-cross gives Ba₂O₂ — but 2 and 2 reduce to 1 and 1.
Step 3
One 2+ ion cancels one 2− ion, so the simplest ratio is 1:1.
Step 4
BaO
Worked example 2. Write the formula of the compound formed from Al³⁺ ions and SO₄²⁻ ions.
Step 1
Write the ions: Al³⁺ and SO₄²⁻.
Step 2
Balance: two 3+ ions give 6+, and three 2− ions give 6−.
Step 3
2 and 3 share no common factor, and sulfate repeats, so parentheses: Al₂(SO₄)₃.
Step 4
Al₂(SO₄)₃
Worked example 3. Write the formula of the compound formed from Mg²⁺ ions and P³⁻ ions.
Step 1
Write the ions: Mg²⁺ and P³⁻.
Step 2
Balance: three 2+ ions give 6+, and two 3− ions give 6−.
Step 3
Phosphorus here is a single-atom ion, so no parentheses: Mg₃P₂.
Step 4
Mg₃P₂
You can now write the formula of an ionic compound from its ions for any combination of single-atom and polyatomic ions, reducing subscripts and adding parentheses where needed.
Check your understanding
Write the formula of the compound formed from K⁺ ions and SO₄²⁻ ions.
Accepted answer: K₂SO₄
Write the ions: K⁺ and SO₄²⁻. Balance the charges: two 1+ ions cancel one 2− ion. Sulfate appears once, so no parentheses: K₂SO₄. Check: 2 × (1+) + (2−) = 0.
Check your understanding
Write the formula of the compound formed from Ca²⁺ ions and O²⁻ ions.
Accepted answer: CaO
Write the ions: Ca²⁺ and O²⁻. Criss-cross gives Ca₂O₂ — reduce 2:2 to 1:1. Write the formula: CaO. Check: (2+) + (2−) = 0.
Check your understanding
Write the formula of the compound formed from Al³⁺ ions and NO₃⁻ ions.
Accepted answer: Al(NO₃)₃
Write the ions: Al³⁺ and NO₃⁻. Balance the charges: one 3+ ion needs three 1− ions. Nitrate repeats, so parentheses: Al(NO₃)₃. Check: (3+) + 3 × (1−) = 0.
Summary video — Ionic formulas and polyatomic ions
Every ionic compound has a name as well as a formula. Two-element compounds get named first — the pattern is short.
The idea
Name the metal first, unchanged: the metal in K₂S is just potassium.
Then name the nonmetal with its ending changed to '-ide': sulfur becomes sulfide.
The -ide names
Nonmetal
Name in the compound
fluorine
fluoride
chlorine
chloride
bromine
bromide
iodine
iodide
oxygen
oxide
sulfur
sulfide
nitrogen
nitride
phosphorus
phosphide
The nonmetal's ending changes to -ide; everything else about its name stays.
K₂S is potassium sulfide.
The name carries no counting words — the subscripts are recovered from the charges whenever the formula is needed.
MgBr₂ follows the same pattern: magnesium bromide.
The common -ide names: fluoride, chloride, bromide, iodide, oxide, sulfide, nitride, phosphide.
Worked examples
Worked example 1. Name the compound Li₂O.
Step 1
Metal first, unchanged: lithium.
Step 2
Nonmetal with -ide: oxygen becomes oxide.
Step 3
Lithium oxide.
Worked example 2. Name the compound Ca₃N₂.
Step 1
Metal first, unchanged: calcium.
Step 2
Nonmetal with -ide: nitrogen becomes nitride.
Step 3
The subscripts 3 and 2 never enter the name.
Step 4
Calcium nitride.
Worked example 3. Name the compound AlCl₃.
Step 1
Metal first, unchanged: aluminum.
Step 2
Nonmetal with -ide: chlorine becomes chloride.
Step 3
Aluminum chloride.
You can now name a binary ionic compound from its formula by giving the metal's name followed by the nonmetal's name with its ending changed to -ide.
Check your understanding
Name the compound CaF₂.
Accepted answer: calcium fluoride
Metal first, unchanged: calcium. Nonmetal with -ide: fluorine becomes fluoride. CaF₂ is calcium fluoride.
Check your understanding
Name the compound Na₂O.
Accepted answer: sodium oxide
Metal first, unchanged: sodium. Nonmetal with -ide: oxygen becomes oxide. Na₂O is sodium oxide.
Check your understanding
Name the compound BaI₂.
Accepted answer: barium iodide
Metal first, unchanged: barium. Nonmetal with -ide: iodine becomes iodide. BaI₂ is barium iodide.
Lesson 26 of 55 · IMB-026
Formula from a binary name
Learning Journey
Read · Check · Unlock
Did You Know?
Naming ran formula → name. Now run the other direction: the name hands you the ions, and charge balancing does the rest.
The idea
The metal's name gives the positive ion — sodium means Na⁺, from Group 1.
The -ide name gives the nonmetal's ion — sulfide means S²⁻, from sulfur in Group 16.
Recall each charge from the element's group position, as you've already seen.
Balance the charges: two 1+ ions cancel one 2− ion, so sodium sulfide is Na₂S.
Check: 2 × (1+) + (2−) = 0.
The name never shows the subscripts — the balancing recovers them.
Worked examples
Worked example 1. Write the formula of calcium iodide.
Step 1
Turn the names into ions: calcium is Ca²⁺ (Group 2), and iodide is I⁻ (Group 17).
Step 2
Balance the charges: one 2+ ion needs two 1− ions.
Step 3
CaI₂
Worked example 2. Write the formula of strontium oxide.
Step 1
Turn the names into ions: strontium is Sr²⁺ (Group 2), and oxide is O²⁻ (Group 16).
Step 2
Criss-cross gives Sr₂O₂ — reduce 2:2 to 1:1.
Step 3
SrO
Worked example 3. Write the formula of aluminum bromide.
Step 1
Turn the names into ions: aluminum is Al³⁺ (Group 13), and bromide is Br⁻ (Group 17).
Step 2
Balance the charges: one 3+ ion needs three 1− ions.
Step 3
AlBr₃
You can now write the formula of a binary ionic compound from its name by recalling each ion's charge and balancing the charges.
Check your understanding
Write the formula of potassium fluoride.
Accepted answer: KF
Turn the names into ions: potassium is K⁺ (Group 1), and fluoride is F⁻ (Group 17). One 1+ ion cancels one 1− ion. Write the formula: KF. Check: (1+) + (1−) = 0.
Check your understanding
Write the formula of calcium sulfide.
Accepted answer: CaS
Turn the names into ions: calcium is Ca²⁺ (Group 2), and sulfide is S²⁻ (Group 16). One 2+ ion cancels one 2− ion — the ratio reduces to 1:1. Write the formula: CaS. Check: (2+) + (2−) = 0.
Check your understanding
Write the formula of lithium nitride.
Accepted answer: Li₃N
Turn the names into ions: lithium is Li⁺ (Group 1), and nitride is N³⁻ (Group 15). Balance the charges: three 1+ ions cancel one 3− ion. Write the formula: Li₃N. Check: 3 × (1+) + (3−) = 0.
Lesson 27 of 55 · IMB-027
Find the metal's charge from the formula
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen that many transition metals form more than one ion — iron forms Fe²⁺ and Fe³⁺. A formula like CoCl₂ doesn't print the metal's charge, but the formula itself lets you work it out.
The equation
Every ionic compound is neutral: total positive charge = total negative charge.
The nonmetal ion's charge is always recallable from its group — chloride is Cl⁻.
In CoCl₂, the two chloride ions contribute 2 × (1−) = 2−.
The single cobalt ion must contribute 2+ to cancel it, so the metal in CoCl₂ is Co²⁺.
When the formula holds more than one metal ion, the balancing positive charge is shared equally among them.
Work the negative side first, then divide the balancing positive charge by the number of metal ions.
Worked examples
Worked example 1. What is the charge of the copper ion in Cu₂S?
Step 1
The sulfide ion, from Group 16, is S²⁻.
Step 2
Total negative charge: 1 × (2−) = 2−.
Step 3
Total positive charge must be 2+.
Step 4
Two copper ions share it: 2+ ÷ 2 = 1+ each.
Step 5
Each copper ion is Cu⁺ — a 1+ charge.
Worked example 2. What is the charge of the chromium ion in CrCl₃?
Step 1
The chloride ion, from Group 17, is Cl⁻.
Step 2
Total negative charge: 3 × (1−) = 3−.
Step 3
One chromium ion must contribute 3+.
Step 4
The chromium ion is Cr³⁺ — a 3+ charge.
Worked example 3. What is the charge of the manganese ion in MnO₂?
Step 1
The oxide ion, from Group 16, is O²⁻.
Step 2
Total negative charge: 2 × (2−) = 4−.
Step 3
One manganese ion must contribute 4+.
Step 4
The manganese ion is Mn⁴⁺ — a 4+ charge.
You can now calculate the charge of the metal ion in an ionic compound's formula from the given charge of the other ion and the requirement that the charges cancel.
Check your understanding
What is the charge of the nickel ion in NiBr₂? Write the number with its sign.
Accepted answer: 2+
Bromide, from Group 17, is Br⁻. Total negative charge: 2 × (1−) = 2−. One nickel ion must contribute 2+. Check: (2+) + 2 × (1−) = 0.
Check your understanding
What is the charge of the copper ion in CuO? Write the number with its sign.
Accepted answer: 2+
Oxide, from Group 16, is O²⁻. Total negative charge: 1 × (2−) = 2−. One copper ion must contribute 2+. Check: (2+) + (2−) = 0.
Check your understanding
What is the charge of each copper ion in Cu₂O? Write the number with its sign.
Accepted answer: 1+
Oxide, from Group 16, is O²⁻. Total negative charge: 1 × (2−) = 2−, so the total positive charge is 2+. Two copper ions share it: 2+ ÷ 2 = 1+ each. Check: 2 × (1+) + (2−) = 0.
Lesson 28 of 55 · IMB-028
Name a transition-metal compound
Learning Journey
Read · Check · Unlock
Did You Know?
Wonder this:
A bottle labeled just 'iron chloride' is ambiguous — FeCl₂ and FeCl₃ are different compounds with different colors and different behavior.
Chemistry's fix is to write the metal ion's charge into the name.
The idea
For a metal that forms more than one ion, the name shows the ion's charge as a Roman numeral in parentheses right after the metal's name.
Roman numerals in names
Metal ion charge
Numeral in the name
1+
(I)
2+
(II)
3+
(III)
4+
(IV)
The numeral is the charge on one metal ion, not a count of atoms.
Find the charge from the formula first — the routine you've already seen.
In CuCl₂, the two Cl⁻ ions total 2−, so the copper ion is Cu²⁺.
Cu²⁺ is written copper(II), so CuCl₂ is copper(II) chloride.
The numeral is the charge on one metal ion — never the count of metal atoms.
The numerals used here: I = 1, II = 2, III = 3, IV = 4.
Worked examples
Worked example 1. Name the compound FeO.
Step 1
Oxide is O²⁻, so the single iron ion must be 2+.
Step 2
Fe²⁺ is written iron(II); oxide stays oxide.
Step 3
Iron(II) oxide.
Worked example 2. Name the compound CrBr₃.
Step 1
Three Br⁻ ions total 3−, so the chromium ion is 3+.
Step 2
Cr³⁺ is written chromium(III).
Step 3
Chromium(III) bromide.
Worked example 3. Name the compound Cu₂S.
Step 1
Sulfide is S²⁻, and the 2+ that balances it is shared by two copper ions: 1+ each.
Step 2
The numeral shows that 1+ charge — not the two copper atoms.
Step 3
Copper(I) sulfide.
You can now name an ionic compound of a transition metal from its formula, showing the metal ion's charge as a Roman numeral in parentheses after the metal's name.
Check your understanding
Name the compound CuO.
Accepted answer: copper(II) oxide
Oxide is O²⁻, so the single copper ion must be 2+. Cu²⁺ is written copper(II). CuO is copper(II) oxide.
Check your understanding
Name the compound FeBr₃.
Accepted answer: iron(III) bromide
Bromide is Br⁻; three of them total 3−. The single iron ion must be 3+ — iron(III). FeBr₃ is iron(III) bromide.
Check your understanding
Name the compound CoO.
Accepted answer: cobalt(II) oxide
Oxide is O²⁻, so the single cobalt ion must be 2+. Co²⁺ is written cobalt(II). CoO is cobalt(II) oxide.
Lesson 29 of 55 · IMB-029
Formula from a Roman-numeral name
Learning Journey
Read · Check · Unlock
Did You Know?
With the Roman numeral in the name, name → formula gets easier, not harder: the numeral hands you the metal ion's charge directly.
The idea
The numeral is the charge on the metal ion: cobalt(II) means Co²⁺.
Turn the -ide name into its ion as before: fluoride is F⁻.
Balance the charges exactly as you've already seen: one 2+ ion needs two 1− ions, so cobalt(II) fluoride is CoF₂.
Check: (2+) + 2 × (1−) = 0.
The numeral never becomes a subscript by itself — the subscripts still come from the balancing.
Worked examples
Worked example 1. Write the formula of copper(II) bromide.
Step 1
The numeral gives the metal ion: copper(II) is Cu²⁺; bromide is Br⁻.
Step 2
Balance the charges: one 2+ ion needs two 1− ions.
Step 3
CuBr₂
Worked example 2. Write the formula of iron(III) chloride.
Step 1
The numeral gives the metal ion: iron(III) is Fe³⁺; chloride is Cl⁻.
Step 2
Balance the charges: one 3+ ion needs three 1− ions.
Step 3
FeCl₃
Worked example 3. Write the formula of chromium(II) chloride.
Step 1
The numeral gives the metal ion: chromium(II) is Cr²⁺; chloride is Cl⁻.
Step 2
Balance the charges: one 2+ ion needs two 1− ions.
Step 3
CrCl₂
You can now write the formula of a transition-metal ionic compound from its name, using the Roman numeral as the metal ion's charge.
Check your understanding
Write the formula of iron(II) sulfide.
Accepted answer: FeS
Iron(II) is Fe²⁺; sulfide, from Group 16, is S²⁻. One 2+ ion cancels one 2− ion — a 1:1 ratio. Write the formula: FeS. Check: (2+) + (2−) = 0.
Check your understanding
Write the formula of copper(II) fluoride.
Accepted answer: CuF₂
Copper(II) is Cu²⁺; fluoride, from Group 17, is F⁻. Balance the charges: one 2+ ion needs two 1− ions. Write the formula: CuF₂. Check: (2+) + 2 × (1−) = 0.
Check your understanding
Write the formula of nickel(II) oxide.
Accepted answer: NiO
Nickel(II) is Ni²⁺; oxide, from Group 16, is O²⁻. One 2+ ion cancels one 2− ion — a 1:1 ratio. Write the formula: NiO. Check: (2+) + (2−) = 0.
Lesson 30 of 55 · IMB-030
Name a compound containing a polyatomic ion
Learning Journey
Read · Check · Unlock
Did You Know?
Compounds built around polyatomic ions are named with pieces you've already seen — and the polyatomic ion's name never changes.
The idea
Spot the polyatomic block first: in K₂CO₃ the block CO₃ is the carbonate ion.
Name the metal, then give the polyatomic ion its own name, unchanged: K₂CO₃ is potassium carbonate.
The '-ide' ending is only for single-atom negative ions — a polyatomic ion's name is never rebuilt.
Hydroxide's name happens to end in -ide already; use it as it is.
Ammonium can be the positive partner: NH₄Cl is ammonium chloride.
Parentheses and subscripts never enter the name.
Worked examples
Worked example 1. Name the compound NaC₂H₃O₂.
Step 1
The block C₂H₃O₂ is the acetate ion.
Step 2
Metal first, then the ion's own name.
Step 3
Sodium acetate.
Worked example 2. Name the compound Ba(OH)₂.
Step 1
The parentheses wrap OH — the hydroxide ion.
Step 2
The subscript 2 stays out of the name.
Step 3
Barium hydroxide.
Worked example 3. Name the compound (NH₄)₂SO₄.
Step 1
Two polyatomic ions: ammonium leads, sulfate closes.
Step 2
Both keep their own names.
Step 3
Ammonium sulfate.
You can now name an ionic compound containing a polyatomic ion from its formula, keeping the polyatomic ion's name unchanged.
Check your understanding
Name the compound KClO₃.
Accepted answer: potassium chlorate
The block ClO₃ is the chlorate ion. Metal first, then the ion's own name. KClO₃ is potassium chlorate.
Check your understanding
Name the compound Mg₃(PO₄)₂.
Accepted answer: magnesium phosphate
The block in the parentheses, PO₄, is the phosphate ion. Metal first, then the ion's own name — no counting words. Mg₃(PO₄)₂ is magnesium phosphate.
Check your understanding
Name the compound Ca(NO₃)₂.
Accepted answer: calcium nitrate
The block in the parentheses, NO₃, is the nitrate ion. Metal first, then the ion's own name. Ca(NO₃)₂ is calcium nitrate.
Lesson 31 of 55 · IMB-031
Formula from a name containing a polyatomic ion
Learning Journey
Read · Check · Unlock
Did You Know?
Running name → formula with a polyatomic ion adds one recall step: the ion's formula and charge come from the roster you've already seen.
The idea
Turn each name into an ion: sodium is Na⁺, and phosphate is PO₄³⁻.
Balance the charges: three 1+ ions cancel one 3− ion.
Write it: Na₃PO₄ — one phosphate, so no parentheses.
Check: 3 × (1+) + (3−) = 0.
When the polyatomic ion repeats, wrap it in parentheses — exactly as you've already seen.
Worked examples
Worked example 1. Write the formula of calcium acetate.
Step 1
Turn the names into ions: calcium is Ca²⁺, and acetate is C₂H₃O₂⁻.
Step 2
Balance the charges: one 2+ ion needs two 1− ions.
Step 3
Acetate repeats, so parentheses: Ca(C₂H₃O₂)₂.
Step 4
Ca(C₂H₃O₂)₂
Worked example 2. Write the formula of lithium carbonate.
Step 1
Turn the names into ions: lithium is Li⁺, and carbonate is CO₃²⁻.
Step 2
Balance the charges: two 1+ ions cancel one 2− ion.
Step 3
One carbonate, so no parentheses: Li₂CO₃.
Step 4
Li₂CO₃
Worked example 3. Write the formula of barium chlorate.
Step 1
Turn the names into ions: barium is Ba²⁺, and chlorate is ClO₃⁻.
Step 2
Balance the charges: one 2+ ion needs two 1− ions.
Step 3
Chlorate repeats, so parentheses: Ba(ClO₃)₂.
Step 4
Ba(ClO₃)₂
You can now write the formula of an ionic compound from a name that includes a polyatomic ion.
Check your understanding
Write the formula of potassium hydroxide.
Accepted answer: KOH
Turn the names into ions: potassium is K⁺, and hydroxide is OH⁻. One 1+ ion cancels one 1− ion. One hydroxide, so no parentheses: KOH. Check: (1+) + (1−) = 0.
Check your understanding
Write the formula of magnesium carbonate.
Accepted answer: MgCO₃
Turn the names into ions: magnesium is Mg²⁺, and carbonate is CO₃²⁻. One 2+ ion cancels one 2− ion — a 1:1 ratio. One carbonate, so no parentheses: MgCO₃. Check: (2+) + (2−) = 0.
Check your understanding
Write the formula of aluminum sulfate.
Accepted answer: Al₂(SO₄)₃
Turn the names into ions: aluminum is Al³⁺, and sulfate is SO₄²⁻. Balance the charges: two 3+ ions give 6+, and three 2− ions give 6−. Sulfate repeats, so parentheses: Al₂(SO₄)₃. Check: 2 × (3+) + 3 × (2−) = 0.
Lesson 32 of 55 · IMB-032
Hydrates
Learning Journey
Read · Check · Unlock
Did You Know?
Wonder this:
Blue crystals of copper(II) sulfate look bone-dry — yet locked inside every crystal, five water molecules travel with each CuSO₄.
Formulas have a way of showing that built-in water.
The idea
A 'hydrate' is an ionic compound whose crystal contains a fixed number of water molecules for every formula unit.
A hydrate's formula: the compound, the raised dot, and the fixed water count.
The formula shows the water after a raised dot: CuSO₄·5H₂O.
Read the dot as 'together with' — every formula unit of CuSO₄ comes with five water molecules.
The number after the dot counts whole water molecules; the 5 belongs to the H₂O, not to the sulfate.
The water is part of the crystal itself — the compound is not wet, and it is not dissolved.
Because the count is fixed, a hydrate is a definite compound: CuSO₄·5H₂O always carries exactly five waters per formula unit.
Worked examples
Worked example 1. How many water molecules travel with each formula unit in MgSO₄·7H₂O?
Step 1
The number after the raised dot counts water molecules.
Step 2
Seven water molecules per formula unit.
Worked example 2. In CaCl₂·2H₂O, which part is the ionic compound?
Step 1
Everything before the raised dot is the ionic compound; everything after it is the built-in water.
Step 2
CaCl₂ — with two water molecules riding along per formula unit.
You can now state that a hydrate is an ionic compound whose crystal contains a fixed number of water molecules per formula unit, written after a raised dot in the formula.
Check your understanding
What does the ·2H₂O in the formula BaCl₂·2H₂O tell you?
ATwo water molecules are built into the crystal for every formula unit of BaCl₂.correct
BThe BaCl₂ has been dissolved in water, two parts of water for every part of solid.
This option is wrong — you read the dot as a solution — the water is locked in the solid crystal, not surrounding it.
CThe crystal's outer surface is wet, coated with a two-molecule layer of water.
This option is wrong — you put the water on the outside — the water molecules sit inside the crystal in a fixed ratio.
DTwo extra hydrogen atoms are bonded directly onto the barium ion in the crystal.
This option is wrong — you split up the H₂O — the dot attaches whole water molecules, not separate atoms.
The number after the raised dot counts whole water molecules. BaCl₂·2H₂O has two water molecules for every formula unit of BaCl₂, inside the crystal.
Check your understanding
How many water molecules travel with each formula unit in CoCl₂·6H₂O?
Answer: 6water molecules
The number after the raised dot counts water molecules. CoCl₂·6H₂O carries 6 water molecules per formula unit.
Check your understanding
Which part of the formula Na₂CO₃·10H₂O is the ionic compound?
ANa₂CO₃correct
B10H₂O
This option is wrong — you picked the built-in water — the ionic compound sits before the raised dot.
CCO₃²⁻
This option is wrong — you picked just the polyatomic ion — the whole Na₂CO₃ before the dot is the compound.
DNa₂CO₃·10H₂O
This option is wrong — the full formula is the hydrate; the ionic compound alone is the part before the dot.
Everything before the raised dot is the ionic compound: Na₂CO₃. The 10H₂O after the dot is the built-in water.
Lesson 33 of 55 · IMB-033
Name a hydrate
Learning Journey
Read · Check · Unlock
Did You Know?
A hydrate's name has two halves: the ionic compound's name, which you can already build, plus a counting word for the water.
The idea
Name the ionic part with the rules you've already seen — the CoCl₂ in CoCl₂·6H₂O is cobalt(II) chloride.
Then add a number prefix plus 'hydrate' for the water count.
Six waters means hexahydrate, so CoCl₂·6H₂O is cobalt(II) chloride hexahydrate.
The prefix counts water molecules only — everything else about the formula is already inside the compound's name.
Worked examples
Worked example 1. Name the hydrate CaSO₄·2H₂O. (Prefixes: 1 mono-, 2 di-, 3 tri-, 4 tetra-, 5 penta-, 6 hexa-, 7 hepta-, 8 octa-, 9 nona-, 10 deca-)
Step 1
Name the ionic part: CaSO₄ is calcium sulfate.
Step 2
Two waters: dihydrate.
Step 3
Calcium sulfate dihydrate.
Worked example 2. Name the hydrate Na₂CO₃·10H₂O. (Prefixes: 1 mono-, 2 di-, 3 tri-, 4 tetra-, 5 penta-, 6 hexa-, 7 hepta-, 8 octa-, 9 nona-, 10 deca-)
Step 1
Name the ionic part: Na₂CO₃ is sodium carbonate.
Step 2
Ten waters: decahydrate.
Step 3
Sodium carbonate decahydrate.
Worked example 3. Name the hydrate Ba(OH)₂·8H₂O. (Prefixes: 1 mono-, 2 di-, 3 tri-, 4 tetra-, 5 penta-, 6 hexa-, 7 hepta-, 8 octa-, 9 nona-, 10 deca-)
Step 1
Name the ionic part: Ba(OH)₂ is barium hydroxide.
Step 2
Eight waters: octahydrate.
Step 3
Barium hydroxide octahydrate.
You can now name a hydrate from its formula by naming the ionic compound and adding a number prefix plus 'hydrate' for the water count, using a supplied prefix table.
Name the ionic part: LiCl is lithium chloride. One water: monohydrate. LiCl·H₂O is lithium chloride monohydrate.
Lesson 34 of 55 · IMB-034
Formula from a hydrate name
Learning Journey
Read · Check · Unlock
Did You Know?
Reverse the hydrate name: the prefix tells you exactly how many waters to write after the raised dot.
The idea
Write the ionic compound's formula from its name, as you've already seen — copper(II) sulfate is CuSO₄.
Turn the prefix into the water count: penta- means 5.
Attach the water after a raised dot: CuSO₄·5H₂O.
A count of one is written with no number — monohydrate means ·H₂O.
The prefix touches only the water — it never changes the compound's own subscripts.
Worked examples
Worked example 1. Write the formula of cobalt(II) nitrate hexahydrate. (Prefixes: 1 mono-, 2 di-, 3 tri-, 4 tetra-, 5 penta-, 6 hexa-, 7 hepta-, 8 octa-, 9 nona-, 10 deca-)
Step 1
Ionic part: cobalt(II) is Co²⁺ and nitrate is NO₃⁻, so Co(NO₃)₂.
Step 2
Hexa- means 6 waters.
Step 3
Co(NO₃)₂·6H₂O
Worked example 2. Write the formula of barium chlorate monohydrate. (Prefixes: 1 mono-, 2 di-, 3 tri-, 4 tetra-, 5 penta-, 6 hexa-, 7 hepta-, 8 octa-, 9 nona-, 10 deca-)
Step 1
Ionic part: barium is Ba²⁺ and chlorate is ClO₃⁻, so Ba(ClO₃)₂.
Step 2
Mono- means 1 water, written with no number.
Step 3
Ba(ClO₃)₂·H₂O
Worked example 3. Write the formula of sodium acetate trihydrate. (Prefixes: 1 mono-, 2 di-, 3 tri-, 4 tetra-, 5 penta-, 6 hexa-, 7 hepta-, 8 octa-, 9 nona-, 10 deca-)
Step 1
Ionic part: sodium is Na⁺ and acetate is C₂H₃O₂⁻, so NaC₂H₃O₂.
Step 2
Tri- means 3 waters.
Step 3
NaC₂H₃O₂·3H₂O
You can now write the formula of a hydrate from its name, converting the number prefix into the count of water molecules after the raised dot.
Check your understanding
Write the formula of strontium chloride hexahydrate. (Prefixes: 1 mono-, 2 di-, 3 tri-, 4 tetra-, 5 penta-, 6 hexa-, 7 hepta-, 8 octa-, 9 nona-, 10 deca-)
Accepted answer: SrCl₂·6H₂O
Ionic part: strontium is Sr²⁺ and chloride is Cl⁻, so SrCl₂. Hexa- means 6 waters. Write it: SrCl₂·6H₂O.
Check your understanding
Write the formula of lithium sulfate monohydrate. (Prefixes: 1 mono-, 2 di-, 3 tri-, 4 tetra-, 5 penta-, 6 hexa-, 7 hepta-, 8 octa-, 9 nona-, 10 deca-)
Accepted answer: Li₂SO₄·H₂O
Ionic part: lithium is Li⁺ and sulfate is SO₄²⁻, so Li₂SO₄. Mono- means 1 water, written with no number. Write it: Li₂SO₄·H₂O.
Check your understanding
Write the formula of calcium chloride dihydrate. (Prefixes: 1 mono-, 2 di-, 3 tri-, 4 tetra-, 5 penta-, 6 hexa-, 7 hepta-, 8 octa-, 9 nona-, 10 deca-)
Accepted answer: CaCl₂·2H₂O
Ionic part: calcium is Ca²⁺ and chloride is Cl⁻, so CaCl₂. Di- means 2 waters. Write it: CaCl₂·2H₂O.
Lesson 35 of 55 · IMB-035
Name any ionic compound
Learning Journey
Read · Check · Unlock
Did You Know?
Every naming skill so far handles one compound type — but a real label never announces its type. Three quick questions sort any ionic formula, and then the right rule takes over.
The idea
Question 1 — is there a raised dot? If yes, name the part before the dot first, then add the water prefix plus 'hydrate'.
Question 2 — is the metal fixed-charge (Groups 1, 2, 13) or ammonium? Name it straight. A transition metal instead? Work out its charge and add the Roman numeral.
Question 3 — is the negative ion a single nonmetal atom or a polyatomic ion? A single atom takes -ide; a polyatomic ion keeps its own name.
Watch the routine on CuSO₄: no dot; copper is a transition metal, and one SO₄²⁻ makes it Cu²⁺ — copper(II); the negative ion is the polyatomic sulfate.
CuSO₄ is copper(II) sulfate.
Worked examples
Worked example 1. Name the compound SrBr₂.
Step 1
No dot; strontium is a fixed-charge Group 2 metal; bromine is a single-atom ion, so -ide.
Step 2
Strontium bromide.
Worked example 2. Name the compound Mn(C₂H₃O₂)₂.
Step 1
No dot; manganese is a transition metal — two acetate ions total 2−, so the manganese is 2+.
Step 2
The negative ion is the polyatomic acetate, keeping its own name.
Step 3
Manganese(II) acetate.
Worked example 3. Name the compound Ni(NO₃)₂·6H₂O. (Prefixes: 1 mono-, 2 di-, 3 tri-, 4 tetra-, 5 penta-, 6 hexa-, 7 hepta-, 8 octa-, 9 nona-, 10 deca-)
Step 1
A raised dot — name the part before it first.
Step 2
Nickel is a transition metal: two NO₃⁻ ions total 2−, so nickel(II) nitrate.
Step 3
Six waters: hexahydrate.
Step 4
Nickel(II) nitrate hexahydrate.
You can now name any ionic compound from its formula — binary, transition-metal, polyatomic, or hydrate.
Check your understanding
Name the compound CaBr₂.
Accepted answer: calcium bromide
No dot; calcium is a fixed-charge Group 2 metal. Bromine is a single-atom ion: bromide. CaBr₂ is calcium bromide.
Check your understanding
Name the compound CoBr₂.
Accepted answer: cobalt(II) bromide
No dot; cobalt is a transition metal. Two Br⁻ ions total 2−, so the cobalt ion is 2+ — cobalt(II). CoBr₂ is cobalt(II) bromide.
Check your understanding
Name the compound KC₂H₃O₂.
Accepted answer: potassium acetate
No dot; potassium is a fixed-charge Group 1 metal. The block C₂H₃O₂ is the polyatomic acetate ion — its name is unchanged. KC₂H₃O₂ is potassium acetate.
Lesson 36 of 55 · IMB-036
Write any ionic formula from its name
Learning Journey
Read · Check · Unlock
Did You Know?
The final skill runs the whole unit in reverse: from any ionic compound's name to its formula.
The idea
Step 1 — turn each name piece into an ion: a fixed-charge metal from its group, a Roman numeral directly, a polyatomic ion from the roster, an -ide name from the nonmetal's group.
Step 2 — balance the charges to the smallest whole-number ratio.
Step 3 — wrap a polyatomic ion in parentheses if it repeats.
Step 4 — turn a hydrate name's prefix into the water count after a raised dot.
Step 5 — check that the charges cancel.
Watch the routine on copper(II) phosphate: the numeral gives Cu²⁺, and the roster gives PO₄³⁻.
Criss-cross: three copper ions and two phosphate ions — Cu₃(PO₄)₂.
Check: 3 × (2+) + 2 × (3−) = (6+) + (6−) = 0.
Worked examples
Worked example 1. Write the formula of barium nitride.
Step 1
Turn the names into ions: barium is Ba²⁺ (Group 2), and nitride is N³⁻ (Group 15).
Step 2
Balance: three 2+ ions give 6+, and two 3− ions give 6−.
Step 3
Ba₃N₂
Worked example 2. Write the formula of ammonium acetate.
Step 1
Turn the names into ions: ammonium is NH₄⁺, and acetate is C₂H₃O₂⁻.
Step 2
One 1+ cancels one 1− — and with nothing repeating, no parentheses are needed.
Step 3
NH₄C₂H₃O₂
Worked example 3. Write the formula of manganese(II) sulfate monohydrate. (Prefixes: 1 mono-, 2 di-, 3 tri-, 4 tetra-, 5 penta-, 6 hexa-, 7 hepta-, 8 octa-, 9 nona-, 10 deca-)
Step 1
The numeral gives Mn²⁺; sulfate is SO₄²⁻ — one 2+ cancels one 2−, so MnSO₄.
Step 2
Mono- means one water after the raised dot.
Step 3
MnSO₄·H₂O
You can now write the formula of any ionic compound from its name — binary, transition-metal, polyatomic, or hydrate.
Check your understanding
Write the formula of calcium phosphide.
Accepted answer: Ca₃P₂
Turn the names into ions: calcium is Ca²⁺, and phosphide is P³⁻. Balance: three 2+ ions give 6+, and two 3− ions give 6−. Write the formula: Ca₃P₂. Check: 3 × (2+) + 2 × (3−) = 0.
Check your understanding
Write the formula of chromium(III) fluoride.
Accepted answer: CrF₃
The numeral gives the metal ion: chromium(III) is Cr³⁺; fluoride is F⁻. Balance: one 3+ ion needs three 1− ions. Write the formula: CrF₃. Check: (3+) + 3 × (1−) = 0.
Check your understanding
Write the formula of sodium chlorate.
Accepted answer: NaClO₃
Turn the names into ions: sodium is Na⁺, and chlorate is ClO₃⁻. One 1+ cancels one 1−, and nothing repeats, so no parentheses. Write the formula: NaClO₃. Check: (1+) + (1−) = 0.
Summary video — Naming ionic compounds and hydrates
Butter melts in a warm pan within seconds, and sugar melts and browns soon after — but the salt sprinkled in stays gritty and solid no matter how hot the stove gets.
Salt is an ionic compound, and its refusal to melt is your first look at how this whole class of solids behaves.
The idea
Ionic compounds have high melting points — typically many hundreds of degrees Celsius.
Melting points of ionic compounds
Ionic compound
Melting point
sodium chloride
801 °C
potassium bromide
734 °C
calcium fluoride
1418 °C
magnesium oxide
2852 °C
Typical ionic compounds melt at many hundreds of degrees Celsius.
Sodium chloride melts at 801 °C, roughly four times hotter than a kitchen oven runs.
The pattern holds across the class: potassium bromide melts at 734 °C, calcium fluoride at 1418 °C, and magnesium oxide at 2852 °C.
So at everyday temperatures an ionic compound is a solid, and it stays solid far beyond cooking and boiling-water temperatures.
When you meet an ionic compound, expect its melting point in the hundreds-to-thousands of degrees Celsius range.
Worked examples
Worked example 1. Calcium chloride is an ionic compound. Would you expect it to melt in boiling water at 100 °C?
Step 1
No — ionic compounds melt at many hundreds of degrees Celsius. Calcium chloride melts at 772 °C.
Worked example 2. Lithium fluoride is an ionic compound. Give a reasonable estimate of its melting point.
Step 1
Many hundreds of degrees Celsius — lithium fluoride melts at 845 °C.
You can now state that ionic compounds have high melting points, typically many hundreds of degrees Celsius.
Check your understanding
Potassium chloride is an ionic compound. Which is the best estimate of its melting point?
AAbout 770 °C.correct
BAbout 100 °C, like boiling water.
This option is wrong — you anchored on boiling water — ionic compounds stay solid far past 100 °C.
CAbout 37 °C, near body temperature.
This option is wrong — you expected it to melt like butter or wax — ionic solids need many hundreds of degrees.
DAbout 0 °C, like melting ice.
This option is wrong — you gave it ice's melting point — an ionic compound is nowhere near melting at 0 °C.
Ionic compounds melt at many hundreds of degrees Celsius. Of the choices, only 770 °C sits in that range — and it is potassium chloride's real melting point.
Check your understanding
Which melting behavior is typical of ionic compounds as a class?
AThey melt at many hundreds of degrees Celsius.correct
BThey melt a little above room temperature.
This option is wrong — you expected everyday-solid behavior — ionic compounds stay solid hundreds of degrees past room temperature.
CThey melt near the boiling point of water.
This option is wrong — you capped the class at 100 °C — sodium chloride alone needs 801 °C.
DThey cannot be melted at any temperature.
This option is wrong — you stretched high into impossible — every ionic compound melts once its high melting point is reached.
The class pattern: ionic compounds melt at many hundreds of degrees Celsius. They do melt — sodium chloride at 801 °C, magnesium oxide at 2852 °C.
Check your understanding
A lab hotplate reaches 300 °C. Potassium iodide, an ionic compound, is placed on it in a metal dish. What happens to the crystals?
AThey stay solid.correct
BThey melt into a liquid.
This option is wrong — you melted an ionic compound at hotplate temperature — potassium iodide needs 681 °C.
CThey soften and turn syrupy.
This option is wrong — you gave crystals sugar-like softening — an ionic solid stays fully solid until its melting point.
DThey turn straight into a gas.
This option is wrong — you skipped past melting entirely — at 300 °C the crystals do not even melt.
Ionic compounds melt at many hundreds of degrees Celsius — potassium iodide at 681 °C. A 300 °C hotplate is far below that, so the crystals stay solid.
Lesson 38 of 55 · IMB-038
Why ionic compounds melt high
Learning Journey
Read · Check · Unlock
Have You Ever Wondered?
Wonder this:
A pizza oven runs hot enough to char the crust and boil off every drop of water — yet the grains of salt on that crust come out exactly as they went in, still solid. What is holding them together so firmly?
You've already seen that ionic compounds melt at very high temperatures — sodium chloride at 801 °C. Now comes the reason.
The idea
Melting sodium chloride means pulling its ions out of their fixed places so they can move past one another.
Ionic compounds have high melting points because the ions are held in a repeating lattice by attraction between opposite charges acting in every direction.
Every ion is pulled toward all of its oppositely charged neighbours at once, not toward just one partner.
Each ion is attracted to all of its oppositely charged neighbours at once.
Pulling an ion away from the lattice means working against all of those attractions at the same time, and that takes a large amount of energy.
Heat supplies that energy, so the temperature must climb many hundreds of degrees before the ions come loose.
Worked examples
Worked example 1. Magnesium oxide is used to line the inside walls of furnaces that run at 2000 °C. Explain why the lining stays solid at that temperature.
Step 1
The Mg²⁺ and O²⁻ ions are held in a repeating lattice by attraction between opposite charges acting in every direction.
Step 2
Melting means pulling every ion away from all of its oppositely charged neighbours at once.
Step 3
A 2000 °C furnace does not supply enough energy to do that — magnesium oxide melts at 2852 °C.
Step 4
The attractions acting in every direction make separating the ions cost a large amount of energy, so magnesium oxide stays solid at 2000 °C.
Worked example 2. Calcium fluoride melts at 1418 °C. Explain why melting it takes so much heat.
Step 1
The Ca²⁺ and F⁻ ions are held in a repeating lattice by attraction between opposite charges acting in every direction.
Step 2
Each ion must be pulled away from all of its oppositely charged neighbours at once.
Step 3
Only a very large amount of energy does that, and the heat has to supply it.
Step 4
The temperature must reach 1418 °C before the added energy is enough to pull the ions out of the lattice.
You can now explain that ionic compounds have high melting points because the ions are held in a repeating lattice by attraction between opposite charges acting in every direction, so separating the ions takes a large amount of energy.
Check your understanding
Potassium bromide melts at 734 °C. Why does melting it take so much heat?
AIts ions are held in a lattice by attractions acting in every direction, and pulling them apart takes a large amount of energy.correct
BMelting has to break electron pairs shared between the potassium and bromine atoms.
This option is wrong — you pictured shared electron pairs — potassium bromide contains K⁺ and Br⁻ ions held by attraction between opposite charges, not shared pairs.
CEach K⁺ ion is attracted to one Br⁻ partner, and that single attraction is very hard to break.
This option is wrong — you pictured the lattice as separate ion pairs — every ion is attracted to all of its oppositely charged neighbours at once, and that is what makes separation so costly.
DThe K⁺ and Br⁻ ions are heavy particles, and heavier particles always take more heat before they can melt.
This option is wrong — you reasoned from particle mass — melting point depends on the strength of the attractions holding the particles, not on how heavy the particles are.
Melting means pulling the ions out of their fixed places. The ions are held in a repeating lattice by attraction between opposite charges acting in every direction. Pulling each ion away from all of its oppositely charged neighbours at once takes a large amount of energy, so the melting point is high.
Check your understanding
A pan on a kitchen stove reaches about 250 °C. Calcium chloride, an ionic compound, melts at 772 °C. Why does calcium chloride stay solid in the pan?
AIts ions are held in a lattice by attractions acting in every direction, and 250 °C does not supply enough energy to pull them apart.correct
BCalcium chloride's particles are far too large and heavy to start moving at 250 °C, so the crystals simply stay frozen in place.
This option is wrong — you reasoned from particle size — what keeps the ions in place is the attraction between opposite charges acting in every direction, not how big the ions are.
CThe pan heats only the outside of each crystal, so the inside never receives any energy.
This option is wrong — you explained the observation with uneven heating — heat spreads through the crystal, but even a fully heated crystal at 250 °C lacks the energy to separate the ions.
DSolid calcium chloride contains no charges, so heat has nothing to act on.
This option is wrong — you removed the ions from the solid — the solid is made entirely of Ca²⁺ and Cl⁻ ions; their opposite-charge attractions are exactly why so much energy is needed.
The ions are held in a repeating lattice by attraction between opposite charges acting in every direction. Separating them takes a large amount of energy — enough only when the temperature reaches 772 °C. A 250 °C pan falls far short, so the crystals stay solid.
Check your understanding
Lithium fluoride melts at 845 °C. At the melting point, what has the added heat finally supplied enough energy to do?
APull each ion away from all of its oppositely charged neighbours so the ions can move past one another.correct
BSplit each lithium fluoride unit into neutral lithium and fluorine atoms.
This option is wrong — you turned melting into a chemical change — the Li⁺ and F⁻ ions keep their charges in the liquid; they are freed from their places, not converted into atoms.
CBreak the shared electron pairs that join each lithium atom to its own fluorine atom.
This option is wrong — you pictured shared electron pairs — lithium fluoride is held by attraction between oppositely charged ions, not by shared pairs.
DBoil the lithium fluoride into a gas.
This option is wrong — you jumped a state too far — at the melting point the solid becomes a liquid; the ions come loose from the lattice but stay close together.
Melting frees the ions from their fixed places in the lattice. The ions are held in a repeating lattice by attraction between opposite charges acting in every direction. At 845 °C the added energy is finally enough to pull each ion away from its oppositely charged neighbours, so the solid melts.
Lesson 39 of 55 · IMB-039
Ionic solids are brittle
Learning Journey
Read · Check · Unlock
Did You Know?
Wonder this:
Rock salt is hard — squeeze a crystal and it does not squish. So it seems like it should stand up to a hammer. It doesn't.
You've already seen that ionic compounds form crystals with high melting points. Here is how those crystals respond to a sharp blow.
The idea
Ionic solids are 'brittle' — struck hard, they crack or shatter rather than bending.
Hardness and brittleness are different behaviours: an ionic crystal resists scratching and squeezing, yet one sharp blow breaks it.
The broken pieces have flat faces — the crystal splits along straight planes instead of denting or crumpling.
So the prediction for any ionic solid under a hammer blow is the same: it shatters; it never flattens.
Worked examples
Worked example 1. A crystal of calcium fluoride is struck sharply with a hammer. Predict what happens to the crystal.
Step 1
Calcium fluoride is an ionic solid, and ionic solids are brittle.
Step 2
A brittle solid cracks or shatters under a sharp blow rather than bending.
Step 3
The crystal shatters into smaller pieces with flat faces.
Worked example 2. A hard white pellet of magnesium oxide is squeezed slowly in a clamp, and then a second identical pellet is hit sharply. Predict each result.
Step 1
Magnesium oxide is an ionic solid: hard, so it resists the slow squeeze without squashing.
Step 2
Ionic solids are brittle, so the sharp blow cracks the second pellet apart.
Step 3
The squeezed pellet holds its shape; the struck pellet cracks into pieces.
You can now state that ionic solids are brittle — when struck hard they crack or shatter rather than bending.
Check your understanding
A crystal of potassium bromide, an ionic compound, is struck sharply with a hammer. What happens to the crystal?
AIt shatters into smaller pieces.correct
BIt flattens into a thin sheet.
This option is wrong — you predicted the behaviour of a metal — an ionic solid never flattens under a blow; it cracks or shatters.
CIt bends, then springs back to its original shape.
This option is wrong — you predicted rubber-like behaviour — an ionic crystal does not bend under a sharp blow; it breaks.
DIt squashes into a rounded lump.
This option is wrong — you treated the crystal as soft — ionic solids are hard, and under a sharp blow they crack or shatter rather than deforming.
Potassium bromide is an ionic solid. Ionic solids are brittle — struck hard, they crack or shatter rather than bending. So the crystal shatters into smaller pieces with flat faces.
Check your understanding
A crystal of lithium fluoride, an ionic compound, is dropped onto a concrete floor and lands hard. What is the most likely result?
AIt cracks or shatters into pieces.correct
BIt dents at the corner that hit the floor.
This option is wrong — you predicted the behaviour of a metal — denting means deforming without breaking, and ionic solids break instead.
CIt bounces without any damage, because ionic crystals are hard.
This option is wrong — you used hardness to rule out breaking — hardness and brittleness are different behaviours, and a hard ionic crystal still cracks under a sharp impact.
DIt flattens slightly where it landed.
This option is wrong — you predicted the behaviour of a metal — an ionic solid does not flatten under impact; it cracks or shatters.
A hard landing is a sharp blow. Ionic solids are brittle — struck hard, they crack or shatter rather than bending. So the lithium fluoride crystal cracks or shatters.
Check your understanding
Two identical crystals of sodium carbonate, an ionic compound, are tested. Crystal 1 is squeezed slowly and steadily in a clamp. Crystal 2 is hit once, sharply. Which pair of results is correct?
ACrystal 1 keeps its shape; crystal 2 cracks apart.correct
BCrystal 1 squashes into a lump; crystal 2 cracks apart.
This option is wrong — you treated the crystal as soft under the slow squeeze — ionic solids are hard and resist steady pressure without squashing.
CCrystal 1 keeps its shape; crystal 2 flattens into a sheet.
This option is wrong — you predicted metal behaviour for the blow — an ionic solid shatters under a sharp blow; it never flattens.
DBoth crystals are unaffected, because ionic solids are hard.
This option is wrong — you used hardness to rule out breaking — a hard ionic crystal still cracks or shatters under a sharp blow; hardness and brittleness are different behaviours.
Hardness handles the slow squeeze: crystal 1 resists and keeps its shape. Ionic solids are brittle — struck hard, they crack or shatter rather than bending. So crystal 2 cracks apart, while crystal 1 survives.
Lesson 40 of 55 · IMB-040
Why ionic solids are brittle
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen that a struck ionic crystal shatters along flat planes. The lattice itself explains why — and the figure shows the whole story in two frames.
The idea
In an untouched sodium chloride crystal, every ion sits beside oppositely charged neighbours, so every neighbouring pair attracts.
Ionic solids are brittle because a sharp blow slides one layer of ions, bringing like charges next to each other, and their repulsion splits the crystal apart.
After the slide, Na⁺ sits beside Na⁺ and Cl⁻ beside Cl⁻ all along the shifted plane.
Like charges repel, so the two halves of the crystal push each other away along that plane.
That is why the crystal splits along a flat face instead of denting.
Worked examples
Worked example 1. A crystal of potassium chloride shatters when struck. Use the lattice to explain why the blow breaks the crystal.
Step 1
The blow slides one layer of ions past its neighbour.
Step 2
The slide brings K⁺ next to K⁺ and Cl⁻ next to Cl⁻ along the shifted plane.
Step 3
Like charges repel, and their repulsion splits the crystal apart along that plane.
Step 4
One shifted layer turns attraction into repulsion all along the plane, so the crystal splits.
Worked example 2. A struck magnesium oxide crystal splits along a flat face rather than crumbling into powder. Explain the flat face.
Step 1
The blow slides a whole layer of ions at once, so the like-charge lineup — Mg²⁺ beside Mg²⁺, O²⁻ beside O²⁻ — forms along one entire plane.
Step 2
The repulsion acts along that whole plane at once.
Step 3
The crystal therefore separates along the plane, leaving a flat face on each half.
Step 4
The split follows the shifted layer, so each new surface is a flat plane of the lattice.
You can now explain that ionic solids are brittle because a sharp blow slides one layer of ions, bringing like charges next to each other, and their repulsion splits the crystal apart.
Check your understanding
A crystal of lithium fluoride shatters when hit with a hammer. Why does the blow split the crystal?
AThe blow slides one layer of ions, bringing like charges next to each other, and their repulsion splits the crystal apart.correct
BThe blow knocks electrons off the ions, so the attractions holding the crystal vanish.
This option is wrong — you invented an electron transfer — a hammer blow moves whole layers of ions; the ions keep their charges throughout.
CThe blow melts a thin layer inside the crystal, and that liquid layer lets the two halves slide apart from each other.
This option is wrong — you brought in melting — no heat change is needed; the shifted layer's like-charge repulsion does the splitting.
DThe blow squeezes the oppositely charged ions so close that they begin to repel each other.
This option is wrong — you made opposite charges repel — opposite charges attract at any distance; the repulsion comes from LIKE charges lining up after the layer slides.
A sharp blow slides one layer of ions past its neighbour. The slide brings like charges next to each other — Li⁺ beside Li⁺, F⁻ beside F⁻. Like charges repel, and their repulsion splits the crystal apart.
Check your understanding
The figure shows a calcium oxide lattice just after a hammer blow has slid the top layer one position sideways. What happens next, and why?
AThe layers push apart, because the slide has put like charges next to each other and like charges repel.correct
BThe layers snap back to their original positions, because opposite charges pull them back into place.
This option is wrong — you let the old attractions win — after the slide, like charges face each other along the whole plane, and their repulsion pushes the halves apart before anything can snap back.
CThe layers keep sliding smoothly, because the ions can roll past one another.
This option is wrong — you described metal-like sliding — in an ionic lattice the slide creates a like-charge lineup that repels and splits the crystal instead of letting it keep deforming.
DNothing changes, because the total number of positive and negative ions is still equal.
This option is wrong — you reasoned from overall neutrality — the crystal is still neutral overall, but the ARRANGEMENT now puts like charges together, and it is the local repulsion that splits it.
Read the shifted frame: Ca²⁺ now sits beside Ca²⁺ and O²⁻ beside O²⁻ along the slide plane. Like charges repel. Their repulsion splits the crystal apart along that plane.
Check your understanding
Before a potassium bromide crystal is struck, its lattice holds together firmly. What changes at the moment a blow slides one layer of ions a single position sideways?
ALike charges come to sit next to each other along the slide plane, so attraction there is replaced by repulsion.correct
BThe ions along the slide plane lose their charges, so nothing holds the halves together anymore.
This option is wrong — you discharged the ions — the ions keep their charges; the problem is that like charges now face each other and repel.
CThe crystal becomes soft along the slide plane, so later blows can bend it.
This option is wrong — you turned shattering into bending — the like-charge repulsion splits the crystal at once; there is no soft, bendable stage.
DThe layer that moved gains extra energy and begins to vibrate until the crystal melts.
This option is wrong — you rerouted the story through melting — no melting occurs; the repulsion between the aligned like charges does the splitting directly.
One slid layer is all it takes. K⁺ now sits beside K⁺ and Br⁻ beside Br⁻ along the plane. Like charges repel, and their repulsion splits the crystal apart.
Lesson 41 of 55 · IMB-041
When ionic compounds conduct
Learning Journey
Read · Check · Unlock
Did You Know?
Wonder this:
Touch a conductivity tester to a pile of dry table salt: the bulb stays dark. Stir that same salt into the tester's beaker of water: the bulb lights. Same compound — opposite results.
A conductivity tester is two metal contacts wired to a bulb; the bulb lights only when electric current flows through whatever sits between the contacts.
The idea
A solid ionic compound does not conduct electricity — solid sodium chloride leaves the bulb dark.
The same compound melted — 'molten' — does conduct.
The same compound dissolved in water also conducts.
So one rule covers every ionic compound: no conduction as a solid; conduction when molten or dissolved.
Worked examples
Worked example 1. A conductivity tester is dipped into three samples of potassium chloride: dry crystals, potassium chloride melted in a crucible, and potassium chloride dissolved in water. Predict the bulb's behaviour for each sample.
Step 1
Dry crystals — a solid ionic compound does not conduct: bulb dark.
Dark, lit, lit — no conduction as a solid; conduction when molten or dissolved.
Worked example 2. Lithium chloride is heated past its melting point. A tester's contacts are lowered into the liquid. Predict the bulb's behaviour.
Step 1
The liquid is molten lithium chloride — an ionic compound past its melting point.
Step 2
An ionic compound conducts when molten.
Step 3
The bulb lights.
You can now state that ionic compounds do not conduct electricity as solids but do conduct when molten (melted) or when dissolved in water.
Check your understanding
A conductivity tester's contacts touch a dry crystal of calcium chloride, an ionic compound. What does the bulb do?
AIt stays dark.correct
BIt lights at full brightness.
This option is wrong — you predicted conduction for a solid ionic compound — an ionic compound conducts only when molten or dissolved, never as a solid.
CIt glows dimly.
This option is wrong — you gave the solid partial conduction — a solid ionic compound does not conduct at all, so the bulb stays fully dark.
DIt flickers on and off.
This option is wrong — you gave the solid intermittent conduction — a solid ionic compound carries no current, so the bulb never lights.
The sample is a solid ionic compound. The rule: no conduction as a solid; conduction when molten or dissolved. So the bulb stays dark.
Check your understanding
Magnesium chloride, an ionic compound, is stirred into a beaker of water until it dissolves. A conductivity tester's contacts are lowered into the beaker. What does the bulb do?
AIt lights.correct
BIt stays dark.
This option is wrong — you carried the solid's behaviour into the solution — a dissolved ionic compound conducts, so the bulb lights.
CThe bulb flickers — solid crystals conduct, but only weakly.
This option is wrong — you gave the solid partial conductivity — a solid ionic compound carries no current at all.
DIt lights only if the water is first heated to boiling.
This option is wrong — you added a heat requirement — a dissolved ionic compound conducts at room temperature; no heating is needed.
The magnesium chloride is dissolved in water. The rule: no conduction as a solid; conduction when molten or dissolved. Dissolved is one of the conducting states, so the bulb lights.
Check your understanding
Which sample of potassium iodide, an ionic compound, leaves a conductivity tester's bulb dark?
ADry potassium iodide crystals.correct
BPotassium iodide melted in a crucible.
This option is wrong — you kept the solid's behaviour after melting — a molten ionic compound conducts, so this sample lights the bulb.
CPotassium iodide dissolved in a beaker of water.
This option is wrong — you kept the solid's behaviour after dissolving — a dissolved ionic compound conducts, so this sample lights the bulb.
DNone of the above — every sample of an ionic compound conducts.
This option is wrong — you made conduction universal — the solid state is the exception, and dry crystals leave the bulb dark.
Run the rule over the three samples: no conduction as a solid; conduction when molten or dissolved. Molten conducts. Dissolved conducts. Only the dry crystals — the solid — leave the bulb dark.
Lesson 42 of 55 · IMB-042
Why conductivity depends on state
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen the pattern: an ionic compound leaves the bulb dark as a solid but lights it when molten or dissolved. The same particles are present in all three samples — so what changes?
The idea
Electric current is a flow of charged particles.
For a sample to conduct, it must contain charged particles that are free to move between the tester's contacts.
Ionic compounds conduct only when molten or dissolved because electric current is carried by charged particles that are free to move — in the solid the ions are locked in the lattice, but when melted or dissolved the ions can move.
In solid sodium chloride, every Na⁺ and Cl⁻ ion is held in its place by its oppositely charged neighbours, so no charge travels anywhere.
Melting or dissolving frees the ions to drift, and the drifting ions carry the current.
Worked examples
Worked example 1. Molten calcium chloride lights a conductivity tester's bulb. Explain why the molten compound conducts.
Step 1
Electric current is carried by charged particles that are free to move.
Step 2
In molten calcium chloride, the Ca²⁺ and Cl⁻ ions are no longer held in a lattice — they can move.
Step 3
The moving ions carry charge between the contacts.
Step 4
Free-moving ions carry the current, so the bulb lights.
Worked example 2. A single dry crystal of potassium bromide leaves the bulb dark. Explain why the solid does not conduct, even though it is packed with charged ions.
Step 1
Charged particles alone are not enough — the charges must be FREE TO MOVE.
Step 2
In the solid, every K⁺ and Br⁻ ion is locked in the lattice by its oppositely charged neighbours.
Step 3
Locked ions cannot travel between the contacts, so no current flows.
Step 4
The ions are present but immobile, so the crystal cannot carry a current.
You can now explain that ionic compounds conduct only when molten or dissolved because electric current is carried by charged particles that are free to move — in the solid the ions are locked in the lattice, but when melted or dissolved the ions can move.
Check your understanding
Molten lithium fluoride lights a conductivity tester's bulb. Why does the molten compound conduct?
ACurrent is carried by charged particles that are free to move, and in the molten compound the ions can move.correct
BMelting releases free electrons from the compound, and the electrons carry the current.
This option is wrong — you handed the current to electrons — in a molten ionic compound the moving charged particles are the ions themselves.
CThe heat of the molten liquid pushes the current through on its own.
This option is wrong — you made heat the charge carrier — heat only melts the compound; the current is carried by the freed ions.
DMelting gives the ions their charges, and charged particles conduct.
This option is wrong — you delayed the charges until melting — the ions are charged in the solid too; melting changes their freedom to move, not their charges.
Electric current is carried by charged particles that are free to move. In molten lithium fluoride the Li⁺ and F⁻ ions are no longer locked in a lattice. The moving ions carry the current, so the bulb lights.
Check your understanding
A dry crystal of sodium bromide is packed with charged Na⁺ and Br⁻ ions from end to end, yet it leaves a conductivity tester's bulb dark. Why doesn't the solid conduct?
AThe ions are locked in the lattice, and charges that cannot move cannot carry a current.correct
BThe positive and negative charges cancel out, so the crystal contains no usable charge.
This option is wrong — you cancelled the charges — the ions keep their individual charges; the problem is that the lattice locks them in place.
CSolid sodium bromide contains no charged particles until it is melted.
This option is wrong — you emptied the solid of ions — the solid is built entirely of Na⁺ and Br⁻ ions; they are simply not free to move.
DThe crystal's surface blocks the current from entering the sample.
This option is wrong — you blamed a surface barrier — the contacts touch the crystal fine; no current flows because the ions inside cannot travel.
Current needs charged particles that are FREE TO MOVE. In the solid, every ion is locked in the lattice by its oppositely charged neighbours. Charges are present but immobile, so the bulb stays dark.
Check your understanding
Potassium sulfate, an ionic compound, is stirred into water and dissolves. The solution lights a conductivity tester's bulb. Which particles carry the current through the solution?
AThe K⁺ and SO₄²⁻ ions, drifting freely through the water.correct
BElectrons flowing through the water from one contact to the other.
This option is wrong — you handed the current to electrons — in a solution the moving charged particles are the dissolved ions, not free electrons.
CThe water itself, since water is a natural conductor.
This option is wrong — you credited the solvent — the water without the dissolved compound would leave the bulb dark; the added ions are the charge carriers.
DWhole potassium sulfate units, moving as neutral pieces.
This option is wrong — you kept the compound in neutral units — a neutral particle carries no charge; the current is carried by the separated K⁺ and SO₄²⁻ ions.
Electric current is carried by charged particles that are free to move. Dissolving frees the K⁺ and SO₄²⁻ ions to drift through the water. Those drifting ions carry the current between the contacts.
Lesson 43 of 55 · IMB-043
The metallic bond and the electron sea
Learning Journey
Read · Check · Unlock
Have You Ever Wondered?
Wonder this:
Sodium chloride holds together because opposite charges attract. But a copper bar contains only copper — one element, no second partner to swap electrons with. What holds a metal together?
The answer is a new arrangement of the same familiar pieces: positive ions and electrons.
The idea
In a metal, each atom releases its outer electrons into a shared pool.
The atoms left behind are positive metal ions, and they sit in fixed positions in a regular arrangement.
The released electrons move freely among the positive ions — chemists call this pool the 'electron sea'.
Positive copper ions sit in fixed positions; the released outer electrons move freely among them.
The 'metallic bond' is the attraction between the positive ions and the electron sea around them.
So solid copper is positive copper ions held in place by a sea of shared, free-moving electrons.
Worked examples
Worked example 1. Describe solid zinc using the electron-sea model.
Step 1
Each zinc atom releases its outer electrons into the shared pool.
Step 2
The positive zinc ions sit in fixed positions in a regular arrangement.
Step 3
The released electrons move freely among the positive ions.
Step 4
Solid zinc is positive zinc ions held together by their attraction to the shared electron sea.
Worked example 2. The figure shows an electron-sea diagram of solid magnesium. What do the large circles and the small dots represent?
Step 1
The large circles, arranged in rows, are the positive magnesium ions.
Step 2
The small dots scattered among them are the free-moving outer electrons — the electron sea.
Step 3
Large circles: fixed positive ions. Small dots: the shared electrons moving freely among them.
You can now state that in a metal the atoms release their outer electrons into a shared 'sea' of electrons that moves freely among the resulting positive metal ions, and that the metallic bond is the attraction between the positive ions and this electron sea.
Check your understanding
In the electron-sea model of solid silver, what is the metallic bond?
AThe attraction between the positive silver ions and the sea of free-moving electrons.correct
BThe attraction between positive silver ions and negative silver ions.
This option is wrong — you added negative ions to the metal — a metal contains no negative ions; the negative charge is the shared electron sea.
CThe shared electron pairs joining each silver atom to its neighbour.
This option is wrong — you pinned the electrons into fixed pairs between particular atoms — in a metal the released electrons belong to the whole sample and move freely.
DThe transfer of electrons from one silver atom to another, forming ion pairs.
This option is wrong — you used ionic transfer — every silver atom releases electrons into the same shared pool; no atom hands electrons to another particular atom.
Each silver atom releases its outer electrons into a shared pool — the electron sea. The atoms left behind are positive ions in fixed positions. The metallic bond is the attraction between those positive ions and the electron sea.
Check your understanding
The figure shows a particle diagram of solid iron. What do the small dots scattered between the large circles represent?
AOuter electrons that move freely among the ions.correct
BNegative iron ions filling the gaps between the positive ones.
This option is wrong — you added negative ions to the metal — a metal contains only positive ions plus the shared electrons; the dots are those electrons.
COuter electrons fixed in place, each belonging to one iron ion.
This option is wrong — you pinned the electrons down — the released electrons belong to no single ion and move freely through the whole metal.
DTiny iron atoms that lost their place in the rows.
This option is wrong — you read the dots as displaced atoms — the ordered large circles are the iron ions, and the dots are the electrons they released.
The large circles in rows are the fixed positive iron ions. The small dots are the outer electrons the atoms released. Those electrons move freely among the ions — the electron sea.
Check your understanding
Which statement correctly describes the particles in a bar of solid aluminum?
APositive aluminum ions sit in fixed positions while their released outer electrons move freely among them.correct
BNeutral aluminum atoms sit in fixed positions with all of their electrons attached.
This option is wrong — you kept the atoms neutral — in a metal each atom has released its outer electrons into the shared sea, leaving positive ions.
CPositive aluminum ions drift freely through a fixed framework of electrons.
This option is wrong — you swapped the moving and fixed parts — the ions are fixed and the electrons move.
DPositive and negative aluminum ions alternate in a repeating lattice.
This option is wrong — you built an ionic lattice out of one element — a metal has no negative ions; the negative charge is the free-moving electron sea.
Each aluminum atom releases its outer electrons into the shared pool. The positive ions stay in fixed positions in a regular arrangement. The released electrons move freely among them — that pairing is the electron-sea picture.
Lesson 44 of 55 · IMB-044
Metals conduct electricity
Learning Journey
Read · Check · Unlock
Did You Know?
You've already seen that an ionic compound conducts only when molten or dissolved. Metals answer the same conductivity test differently.
The idea
A solid metal conducts electricity — a copper wire carries current exactly as it is, no melting or dissolving needed.
A molten metal conducts too.
So a metal conducts in both states: solid and molten.
That is the opposite of an ionic solid, which never conducts until it is melted or dissolved.
Worked examples
Worked example 1. An overhead power line is solid aluminum. Does the line conduct, and does it need any preparation first?
Step 1
Aluminum is a metal, and a metal conducts as a solid.
Step 2
No melting or dissolving is needed.
Step 3
The solid aluminum line conducts just as it is.
Worked example 2. In a foundry, a vat of molten iron sits at 1600 °C. A conductivity probe is lowered into the liquid. Predict the result.
Step 1
Iron is a metal, and a metal conducts when molten as well as when solid.
Step 2
The state change does not switch conduction off.
Step 3
The molten iron conducts, and the probe registers a current.
You can now state that metals conduct electricity in both the solid and molten states.
Check your understanding
A conductivity tester's contacts touch a solid bar of zinc. What does the bulb do?
AIt lights.correct
BIt stays dark.
This option is wrong — you applied the ionic-solid rule to a metal — a metal conducts as a solid, no melting or dissolving needed.
CIt stays dark unless the bar is first heated.
This option is wrong — you added a heat requirement — a solid metal conducts at room temperature.
DIt lights only if the bar is first dissolved in water.
This option is wrong — you carried the dissolved-ionic condition over to a metal — a metal conducts as the solid bar it already is.
Zinc is a metal. A metal conducts in both states — solid and molten. So the solid bar conducts and the bulb lights.
Check your understanding
Tin melts at 232 °C. A crucible of tin is heated to 300 °C, and a conductivity probe is lowered into the liquid. What does the probe register?
AA current — the molten tin conducts.correct
BNo current — melting a metal stops it from conducting.
This option is wrong — you switched conduction off at the melting point — a metal conducts when molten as well as when solid.
CNo current — liquids cannot carry electricity.
This option is wrong — you ruled out all liquids — molten metals and molten ionic compounds both conduct; the state alone does not decide.
DA current, but only because the probe itself is metal.
This option is wrong — you credited the probe — the probe registers current only when the sample between its contacts conducts, and molten tin does.
Tin is a metal. A metal conducts in both states — solid and molten. At 300 °C the tin is molten, and it conducts.
Check your understanding
Two white solids sit on a bench: a bar of magnesium and a crystal of magnesium chloride. A conductivity tester touches each in turn. Which pair of results is correct?
AThe magnesium bar lights the bulb; the magnesium chloride crystal leaves it dark.correct
BBoth light the bulb.
This option is wrong — you let the ionic solid conduct — an ionic compound conducts only when molten or dissolved, so the crystal leaves the bulb dark.
CNeither lights the bulb.
This option is wrong — you took conduction away from the solid metal — a metal conducts as a solid, so the magnesium bar lights the bulb.
DThe magnesium chloride crystal lights the bulb; the magnesium bar leaves it dark.
This option is wrong — you swapped the two rules — the metal conducts as a solid and the ionic solid does not.
Magnesium is a metal: it conducts as a solid — bulb lit. Magnesium chloride is an ionic compound: no conduction as a solid — bulb dark. Same test, opposite results, because the two solids follow different rules.
Lesson 45 of 55 · IMB-045
Why metals conduct
Learning Journey
Read · Check · Unlock
Did You Know?
A copper wire conducts while its ions stay exactly where they are. Something else must be doing the moving — and the electron-sea model says what.
The idea
Electric current is a flow of charged particles, and a sample conducts only if it holds charged particles that are free to move.
Metals conduct electricity because the metal's outer electrons are free to move among the fixed positive ions.
Connect a wire into a circuit and the sea of electrons drifts through it — that drift is the current.
The electron sea drifts through the wire; the positive ions stay in place.
The positive ions never leave their positions; the wire stays a solid bar while its electrons flow.
The electron sea survives melting, which is why a molten metal keeps conducting.
Worked examples
Worked example 1. A solid aluminum power line carries current across a valley. Explain how the current travels through the solid metal.
Step 1
Aluminum's outer electrons are released into a shared sea and are free to move among the fixed positive ions.
Step 2
In the circuit, that sea of electrons drifts along the line.
Step 3
The drifting electrons are the current — the aluminum ions stay in place.
Step 4
Free-moving electrons carry the current through the solid line.
Worked example 2. Molten silver in a crucible conducts electricity. Explain why melting did not stop the conduction.
Step 1
Conduction needs charged particles that are free to move.
Step 2
In silver those carriers are the outer electrons, free to move among the positive ions in the solid and in the melt alike.
Step 3
Melting loosens the ions, but the electron sea is there either way.
Step 4
The free-moving electrons carry current in both states, so molten silver still conducts.
You can now explain that metals conduct electricity because the metal's outer electrons are free to move among the fixed positive ions.
Check your understanding
A solid gold contact in a circuit carries current. Why does the solid metal conduct?
AThe metal's outer electrons are free to move among the fixed positive ions.correct
BThe positive gold ions flow through the solid from one end to the other.
This option is wrong — you set the ions moving — in a solid metal the ions stay in fixed positions; the free-moving electrons carry the current.
CWhole gold atoms drift along the contact, carrying their charges with them.
This option is wrong — you moved neutral atoms — a neutral atom carries no charge, and the metal's atoms have already released their outer electrons; those electrons are the carriers.
DThe solid contains no moving particles, but current can pass through empty spaces between the ions.
This option is wrong — you let current flow without a carrier — current IS moving charge, and in a metal the moving charges are the free electrons.
Electric current is a flow of charged particles. In gold, the outer electrons are free to move among the fixed positive ions. That electron drift is the current through the contact.
Check your understanding
A zinc bar conducts electricity while remaining a rigid solid — nothing about its shape changes. Which statement resolves the apparent contradiction?
AThe positive ions stay fixed, keeping the bar rigid, while the free electrons move among them, carrying the current.correct
BAll of the bar's particles drift slightly along with the current, so the bar is never truly rigid while it conducts.
This option is wrong — you loosened the ions to allow conduction — the ions never need to move; the electron sea does all the moving.
CCurrent only flows along the bar's outer surface, leaving the rigid inside untouched.
This option is wrong — you confined the current to the surface — free electrons move throughout the metal, inside and out.
DThe bar conducts only for short bursts, between which its particles lock back in place.
This option is wrong — you made conduction intermittent — the electrons are permanently free to move; rigidity and conduction coexist continuously.
The bar has two kinds of particles with two different jobs. The fixed positive ions give the bar its rigid shape. The metal's outer electrons are free to move among the fixed positive ions — they carry the current.
Check your understanding
Which particles carry the current through a solid magnesium ribbon in a circuit?
AThe free-moving outer electrons.correct
BThe positive magnesium ions.
This option is wrong — you picked the fixed framework — the ions hold their positions; the released outer electrons do the moving.
CNegative magnesium ions drifting between the positive ones.
This option is wrong — you invented negative ions — a metal contains no negative ions; its negative charge is the free electron sea.
DNeutral magnesium atoms passing charge along a chain.
This option is wrong — you moved neutral atoms — neutral particles carry no charge; the current is the drift of the free electrons.
In the ribbon, each magnesium atom has released its outer electrons into the shared sea. Those electrons are free to move among the fixed positive ions. Their drift through the ribbon is the current.
Lesson 46 of 55 · IMB-046
Metals bend, not break
Learning Journey
Read · Check · Unlock
Did You Know?
Wonder this:
A hammer blow shatters a salt crystal. The same blow lands on a piece of copper — and the copper just flattens.
Metals respond to force in their own way, and the behaviour has two everyday faces.
The idea
Hammer a metal and it flattens instead of breaking — metals can be beaten into thin sheets, a property called being 'malleable'.
Pull a metal through a narrow opening and it stretches instead of snapping — metals can be drawn into wires, a property called being 'ductile'.
Both faces are one underlying behaviour: under force, a metal changes shape rather than shattering.
That is the opposite of an ionic solid, which cracks or shatters under the same blow.
Worked examples
Worked example 1. A goldsmith hammers a small lump of gold thinner and thinner until it becomes gold leaf, a sheet far thinner than paper. Which property of metals is the goldsmith using?
Step 1
The gold is being beaten into a thin sheet.
Step 2
Beaten into sheets means malleable.
Step 3
The goldsmith is using gold's malleability.
Worked example 2. In a wire mill, a thick rod of aluminum is pulled through a series of smaller and smaller openings until it becomes fine wire. Which property of metals is the mill using?
Step 1
The aluminum is being drawn out into wire.
Step 2
Drawn into wires means ductile.
Step 3
The mill is using aluminum's ductility.
You can now state that metals are malleable (they can be hammered into sheets) and ductile (they can be drawn into wires).
Check your understanding
A silversmith hammers a bar of silver into a flat, curved plate without a single crack appearing. Which property of silver does this show?
AMalleability.correct
BDuctility.
This option is wrong — you picked the wire-drawing property — hammering into sheets is malleability; ductility is being drawn into wires.
CBrittleness.
This option is wrong — you picked the shattering behaviour — a brittle solid would have cracked under the hammer, and the silver flattened instead.
DConductivity.
This option is wrong — you picked an electrical property — carrying current has nothing to do with how the metal responds to a hammer.
The silver was beaten flat with a hammer and did not crack. Beaten into sheets means malleable. So the plate shows silver's malleability.
Check your understanding
A factory pulls thick zinc rod through smaller and smaller openings until it becomes thin zinc wire. Which property of zinc does this show?
ADuctility.correct
BMalleability.
This option is wrong — you picked the hammering property — being drawn into wires is ductility; malleability is being beaten into sheets.
CBrittleness.
This option is wrong — you picked the shattering behaviour — a brittle solid would snap in the opening instead of stretching into wire.
DLuster (a shiny surface).
This option is wrong — you picked the shiny-surface property — shine has nothing to do with how the metal responds to being pulled.
The zinc was pulled through openings and stretched instead of snapping. Drawn into wires means ductile. So the wire shows zinc's ductility.
Check your understanding
A hammer strikes a small bar of lead and, separately, a crystal of potassium bromide. What happens to each?
AThe lead flattens; the potassium bromide crystal shatters.correct
BBoth flatten.
This option is wrong — you gave metal behaviour to the ionic solid — an ionic crystal cracks or shatters under a blow; it never flattens.
CBoth shatter.
This option is wrong — you gave ionic behaviour to the metal — a metal changes shape under force instead of shattering.
DThe lead shatters; the potassium bromide crystal flattens.
This option is wrong — you swapped the two behaviours — the metal flattens and the ionic crystal shatters.
Lead is a metal: under force it changes shape — it flattens. Potassium bromide is an ionic solid: struck hard, it cracks or shatters. Same blow, opposite outcomes.
Lesson 47 of 55 · IMB-047
Why metals bend
Learning Journey
Read · Check · Unlock
Did You Know?
A blow slides a layer of particles in a metal, just as it does in an ionic crystal. The ionic crystal split apart. The metal doesn't — and the electron sea is the reason.
The idea
A blow or a pull slides one layer of metal ions past its neighbours.
Metals are malleable and ductile because layers of metal ions can slide past each other while the electron sea keeps holding the ions together.
After the slide, the arrangement looks just as it did before: positive ions surrounded by the same free-moving electrons, so nothing repels.
In an ionic crystal, the same slide lines up like charges, which repel and split the crystal; in a metal, no like-against-like lineup ever forms.
So the metal changes shape and holds together, blow after blow.
Worked examples
Worked example 1. A goldsmith's hammer flattens a lump of gold into leaf without cracking it. Explain, at the particle level, why the gold flattens instead of shattering.
Step 1
Each hammer blow slides layers of gold ions past each other.
Step 2
The electron sea surrounds the ions in their new positions just as it did in the old ones, so the attraction is unbroken.
Step 3
With no like-charge lineup to repel, nothing splits.
Step 4
The layers slide while the electron sea keeps holding the ions together, so the gold thins into leaf instead of cracking.
Worked example 2. A thick rod of aluminum is drawn through a narrow opening and stretches into wire. Explain why the rod stretches rather than snapping.
Step 1
The pull slides layers of aluminum ions past each other, lengthwise.
Step 2
The electron sea keeps holding the ions together in every new arrangement.
Step 3
Layer after layer slides, and the rod lengthens without any plane of repulsion forming.
Step 4
Sliding layers plus an unbroken electron-sea attraction let the rod stretch into wire.
You can now explain that metals are malleable and ductile because layers of metal ions can slide past each other while the electron sea keeps holding the ions together.
Check your understanding
A hammer blow flattens a bar of silver instead of cracking it. Why does the silver hold together while its shape changes?
ALayers of silver ions slide past each other while the electron sea keeps holding the ions together.correct
BThe blow pushes the silver ions closer together, and closer ions always bond more strongly.
This option is wrong — you explained the survival with compression — the layers SLIDE sideways, and it is the ever-present electron sea that keeps the ions held.
CThe silver ions are locked so rigidly that a hammer cannot move them at all.
This option is wrong — you froze the layers — the blow does slide the layers; the flattening IS the layers moving, held together by the electron sea.
DEach silver ion is bonded to one particular neighbour and drags it along wherever it moves.
This option is wrong — you built fixed pairs — no ion is tied to a particular partner; every ion is held by the shared electron sea, whichever neighbours it has.
The blow slides layers of silver ions past each other. The electron sea surrounds the ions in every new position, so the attraction never breaks. The shape changes; the metal holds together.
Check your understanding
A blow slides one layer of particles in a zinc bar, and the same blow slides one layer in a potassium chloride crystal. The zinc dents; the crystal splits. What makes the outcomes differ?
AIn zinc every slid ion is still held by the electron sea, but in the crystal the slide lines up like charges, which repel and split it.correct
BZinc's particles are held together more strongly than the crystal's are, and stronger attractions always survive a hammer blow.
This option is wrong — you turned the difference into strength — the difference is what the slide CREATES: an unchanged environment in the metal, a like-charge lineup in the ionic crystal.
CThe blow slides layers in the metal only — layers in an ionic crystal cannot move.
This option is wrong — you froze the ionic layers — the blow slides them too; the split happens BECAUSE they slide into a like-charge lineup.
DZinc has no charged particles, so there is nothing inside it to repel.
This option is wrong — you emptied the metal of charges — zinc is positive ions in an electron sea; nothing repels because every slid ion is still surrounded by that sea, not because charges are absent.
Same slide, different aftermath. Zinc: layers of metal ions slide past each other while the electron sea keeps holding the ions together — no repulsion forms. Potassium chloride: the slide brings like charges next to each other, and their repulsion splits the crystal apart.
Check your understanding
The figure shows a metal lattice just after a blow has slid the top layer one position. Why does no crack form along the slide plane?
AEach ion in its new position is still a positive ion surrounded by the electron sea, so the attraction continues and nothing repels.correct
BThe ions along the plane swapped their charges to keep the rows attracting.
This option is wrong — you re-charged the ions — no charges change; the ions were positive before and after, and the electron sea holds them either way.
CThe electron sea rushed into the slide plane and glued the crack shut after it formed.
This option is wrong — you let a crack form first — no crack ever opens, because the electron sea already surrounds every ion throughout the slide.
DThe slide was simply too small to matter — a full-position slide would split a metal apart just like an ionic crystal.
This option is wrong — you kept the ionic outcome for metals — the figure shows a full one-position slide, and the metal still holds, because like charges never line up against each other in an electron sea.
Compare the two frames: before and after, every ion has the same surroundings. Positive ions in an electron sea — no like-charge lineup can form. Layers of metal ions can slide past each other while the electron sea keeps holding the ions together.
Lesson 48 of 55 · IMB-048
Metals are shiny
Learning Journey
Read · Check · Unlock
Did You Know?
One more everyday metal property, and then the why. Look at polished silverware: it shows your reflection.
The idea
A clean metal surface is shiny — polished silver reflects a clear image.
This shine is called 'luster'.
Luster belongs to the clean metal surface: tarnish or dirt on top hides it, and polishing or cutting brings it back.
So a dull metal object is dull because of its surface coating, not because the metal underneath lost its shine.
Worked examples
Worked example 1. A copper pot has turned dull brown all over. After a minute with polishing paste, the surface is bright and reflective again. What does the polishing show about the copper's luster?
Step 1
The dull brown layer was a surface coating sitting on top of the metal.
Step 2
Polishing removed the coating and exposed clean copper.
Step 3
The exposed clean surface is shiny.
Step 4
The copper kept its luster all along — the coating was hiding it.
Worked example 2. A stick of sodium metal looks dull grey. A knife slices it, and the fresh cut surface is bright and mirror-like for a few moments. Explain the bright surface.
Step 1
The dull grey outside is a coating that formed on the stored metal.
Step 2
The cut exposes clean sodium metal underneath.
Step 3
A clean metal surface is shiny.
Step 4
The fresh cut shows sodium's luster — the clean metal surface reflects light.
You can now state that clean metal surfaces are shiny, a property called luster.
Check your understanding
A freshly polished gold ring reflects a small, clear image of a window. Which property of gold is this?
ALuster.correct
BMalleability.
This option is wrong — you picked the shape-change property — being beaten into sheets is malleability; reflecting light from a clean surface is luster.
CDuctility.
This option is wrong — you picked the wire-drawing property — being drawn into wires is ductility; the shiny reflection is luster.
DConductivity.
This option is wrong — you picked the electrical property — carrying current is conductivity; the shiny reflection is luster.
A clean metal surface is shiny. That shine — reflecting a clear image — is called luster. The polished ring is showing gold's luster.
Check your understanding
An old aluminum pan looks grey and dull. A student rubs one patch with fine sandpaper, and the patch turns bright and reflective. Why did the patch become shiny?
ASanding removed the dull surface coating and exposed clean aluminum, and a clean metal surface is shiny.correct
BThe rubbing heated the aluminum, and warm metal glows brightly.
This option is wrong — you turned reflection into glowing — the patch is reflecting room light from a clean surface, not giving off its own light.
CThe sandpaper deposited a shiny layer of its own material onto the pan.
This option is wrong — you added a coating instead of removing one — sanding strips the dull layer away; the shine is the aluminum itself.
DThe rubbing flattened the aluminum surface, and flattened metal always looks shinier.
This option is wrong — you reached for malleability — the change is at the surface: the dull coating came off, and the clean metal beneath has luster.
Luster belongs to the clean metal surface. The grey layer was a coating hiding it. Sanding exposed clean aluminum, and the clean surface is shiny.
Check your understanding
Which statement about a dull, tarnished strip of nickel is correct?
AThe nickel under the tarnish still has its luster — polishing the strip will expose a shiny surface.correct
BThe nickel has permanently lost its luster, and no treatment can bring the shine back.
This option is wrong — you assigned the dullness to the metal itself — the dullness is the surface coating, and removing it restores the shine.
CThe strip was never really nickel, because a true metal cannot turn dull.
This option is wrong — you made luster tarnish-proof — real metals grow dull coatings; the metal beneath stays shiny.
DThe tarnish itself will turn shiny if the strip is left alone long enough.
This option is wrong — you gave the coating luster — the coating is not clean metal; only removing it exposes the shine.
Luster belongs to the clean metal surface. Tarnish is a coating sitting on top, hiding the shine. Polish it off and the clean nickel beneath reflects again.
Lesson 49 of 55 · IMB-049
Why metals are shiny
Learning Journey
Read · Check · Unlock
Did You Know?
Polished aluminum reflects your face; polished rock salt never will. The electron sea explains the difference.
The idea
Light lands on the metal's surface.
Metals are shiny because the free-moving electrons at the metal's surface reflect light.
The same electron sea that carries current through the metal is what turns light back at the surface.
A coating of tarnish or dirt is not metal — it has no electron sea, so it reflects poorly until it is removed.
Worked examples
Worked example 1. The reflective backing of a bathroom mirror is a thin layer of silver. Explain why silver makes the mirror reflect.
Step 1
Light passes through the glass and lands on the silver surface.
Step 2
The free-moving electrons at the silver's surface reflect the light back.
Step 3
The silver layer's surface electrons send the light back, so the mirror shows an image.
Worked example 2. A freshly polished gold bracelet gleams. Explain the gleam at the particle level.
Step 1
Polishing exposes clean gold — a surface with its electron sea at the top.
Step 2
The free-moving electrons at that surface reflect the light that lands on it.
Step 3
The free-moving surface electrons reflect the light, and the bracelet gleams.
You can now explain that metals are shiny because the free-moving electrons at the metal's surface reflect light.
Check your understanding
Why does a polished copper tray reflect light?
AThe free-moving electrons at the copper's surface reflect light.correct
BThe tray's surface is perfectly smooth, and any perfectly smooth material reflects like a metal.
This option is wrong — you made smoothness the whole cause — polished rock salt is smooth too and never gleams like copper; the metal's free surface electrons do the reflecting.
CCopper absorbs the light and then glows with its own new light.
This option is wrong — you turned reflection into glowing — the tray sends back the light that lands on it; it does not produce light.
DThe copper ions at the surface bounce the light off their positive charges.
This option is wrong — you gave the job to the ions — it is the free-moving electrons at the surface, not the fixed ions, that reflect the light.
Light lands on the copper's surface. The free-moving electrons at the metal's surface reflect light. That reflected light is the shine you see.
Check your understanding
Chromium is used for shiny bumpers and taps. Which feature of chromium makes its clean surface reflective?
AFree-moving electrons at its surface, which reflect light.correct
BIts hardness — harder materials always reflect more light.
This option is wrong — you tied shine to hardness — hard nonmetal solids can still be dull; the reflection comes from the metal's free surface electrons.
CIts high melting point, which keeps the surface solid and mirror-like.
This option is wrong — you tied shine to melting point — staying solid does not create reflection; the free-moving surface electrons do.
DA thin layer of glass that forms naturally on chromium.
This option is wrong — you added a glassy coating — no such layer exists, and coatings HIDE luster; the clean metal surface itself reflects because of its free electrons.
Chromium is a metal: positive ions in a sea of free-moving electrons. The free-moving electrons at the metal's surface reflect light. That is what makes the clean bumper mirror-bright.
Check your understanding
A tarnished brass doorknob reflects poorly, but polishing makes it gleam. At the particle level, what did polishing change?
AIt removed the coating, exposing a metal surface whose free-moving electrons reflect light.correct
BIt pushed extra electrons into the metal, giving the surface more charge to reflect with.
This option is wrong — you added electrons — polishing adds nothing; it strips the coating so the metal's own surface electrons can face the light.
CIt smoothed the tarnish itself until the coating became reflective.
This option is wrong — you kept the coating — tarnish is not metal and has no free surface electrons; the gleam comes from the exposed metal beneath.
DIt heated the surface, and warm surfaces reflect better.
This option is wrong — you brought in heat — temperature is beside the point; the change is bare metal, with free-moving electrons, now facing the light.
The coating is not metal — it has no electron sea, so it reflects poorly. Polishing strips the coating away. The free-moving electrons at the exposed metal surface reflect light, and the knob gleams.
Lesson 50 of 55 · IMB-050
Covalent bonding: sharing electrons
Learning Journey
Read · Check · Unlock
Have You Ever Wondered?
Wonder this:
Chlorine gas is made of Cl₂ — two chlorine atoms bound tightly together. But both atoms are nonmetals: each one gains electrons in every compound you have met. Neither will hand an electron over. So what holds Cl₂ together?
Transfer is off the table when no atom will give. The way out is the third option.
The idea
Two nonmetal atoms both tend to gain electrons, so neither atom gives its electrons away.
Instead, the two atoms share a pair of electrons — the shared pair sits between the two atoms and belongs to both at once.
Bonding by sharing electron pairs is called 'covalent' bonding.
In sodium chloride the metal atom transfers its electron to the nonmetal; in Cl₂ the two atoms share a pair — transfer makes ions, sharing does not.
Different nonmetals share too: the pair does not need two identical atoms, only two atoms that both hold on to their electrons.
Worked examples
Worked example 1. Fluorine gas is made of F₂ particles. What holds the two fluorine atoms together, and what is this bonding called?
Step 1
Both atoms are nonmetals, so neither gives an electron away.
Step 2
The two atoms share a pair of electrons, and the shared pair belongs to both.
Step 3
A shared electron pair holds F₂ together — covalent bonding.
Worked example 2. Hydrogen chloride, HCl, is made of one hydrogen atom joined to one chlorine atom — two different nonmetals. Is the bonding covalent, and why?
Step 1
Hydrogen and chlorine are both nonmetals, so neither hands an electron to the other.
Step 2
The two atoms share a pair of electrons instead.
Step 3
Yes — the shared pair makes the H–Cl bond covalent, even though the atoms differ.
You can now state that nonmetal atoms bond with each other by sharing pairs of electrons, and that bonding by sharing is called covalent bonding.
Check your understanding
In a molecule of bromine, Br₂, the two bromine atoms are held together by a pair of electrons that sits between them and belongs to both. What is this bonding called?
ACovalent bonding.correct
BIonic bonding.
This option is wrong — you named the transfer category — ionic bonding is the attraction between oppositely charged ions, and no electron was transferred here.
CMetallic bonding.
This option is wrong — you named the electron-sea category — metallic bonding is positive ions in a shared sea of free electrons, not a single pair shared between two atoms.
DNone of the above.
This option is wrong — you rejected all three categories — sharing an electron pair between two atoms is exactly the covalent case.
Both bromine atoms are nonmetals — neither gives an electron away. They share a pair of electrons, and the pair belongs to both. Bonding by sharing electron pairs is covalent bonding.
Check your understanding
In a water molecule, each hydrogen atom is joined to the oxygen atom by one shared pair of electrons. Which statement about the oxygen–hydrogen bonding is correct?
AIt is covalent — the atoms share electron pairs, and no ions form.correct
BIt is ionic — the oxygen atom takes both electrons of each pair completely.
This option is wrong — you turned sharing into transfer — the pair sits between the two atoms and belongs to both; no electron leaves either atom.
CIt is covalent, but only because the two atoms are the same element.
This option is wrong — you required identical atoms — different nonmetals also share; hydrogen and oxygen sharing a pair is still covalent.
DIt is metallic — the electrons in water move freely among the atoms.
This option is wrong — you gave water an electron sea — the shared pairs stay between particular atoms; nothing moves freely as in a metal.
Hydrogen and oxygen are both nonmetals — neither gives its electrons away. Each O–H pair is shared and belongs to both atoms. Sharing electron pairs is covalent bonding, whatever the two nonmetals are.
Check your understanding
The figure shows two particle diagrams. Diagram 1: two identical circles overlap, with two electron dots in the overlap. Diagram 2: a small circle marked + sits apart from a large circle marked −, with an arrow showing one electron having moved from the small circle to the large one. Which diagram shows covalent bonding?
ADiagram 1 only.correct
BDiagram 2 only.
This option is wrong — you picked the transfer picture — a moved electron and separated ions is the ionic case; covalent bonding is the shared pair in the overlap.
CBoth diagrams.
This option is wrong — you merged the two categories — only the overlap with a shared pair is covalent; the transfer picture makes ions.
DNeither diagram.
This option is wrong — you rejected the shared pair — two dots in the overlap belonging to both atoms is exactly what covalent bonding looks like.
Covalent bonding is bonding by sharing electron pairs. Diagram 1 puts the pair in the overlap — shared, belonging to both atoms. Diagram 2 moves an electron across and leaves charged ions — that is transfer, not sharing.
Lesson 51 of 55 · IMB-051
Predict bond type from the elements' positions
Learning Journey
Read · Check · Unlock
Have You Ever Wondered?
Wonder this:
Magnesium and oxygen react in a flash of white light. Sulfur and oxygen react too. Can you tell, before either reaction happens, which product will be ionic and which covalent — just from a periodic table?
You can, because the periodic table already tells you which atoms give electrons away and which hold on to them.
The idea
Find each element in the periodic table and note whether it is a metal or a nonmetal.
A metal with a nonmetal generally bonds ionically — the metal transfers electrons and oppositely charged ions attract.
Metal with nonmetal: ionic. Nonmetal with nonmetal: covalent.
A nonmetal with a nonmetal bonds covalently — both atoms hold their electrons, so they share pairs.
Magnesium (metal) with oxygen (nonmetal): ionic. Sulfur (nonmetal) with oxygen (nonmetal): covalent. Magnesium (metal) with sulfur (nonmetal): ionic.
'Generally' matters — position gives a first-pass prediction, and measured properties are better evidence than position alone.
Worked examples
Worked example 1. Predict whether potassium and bromine bond ionically or covalently.
Step 1
Potassium is on the far left of the periodic table — a metal.
Step 2
Bromine is on the upper right — a nonmetal.
Step 3
Metal with nonmetal: ionic.
Step 4
Potassium and bromine bond ionically.
Worked example 2. Predict whether carbon and hydrogen bond ionically or covalently.
Step 1
Carbon is a nonmetal, and hydrogen is a nonmetal.
Step 2
Nonmetal with nonmetal: covalent.
Step 3
Carbon and hydrogen bond covalently.
You can now predict whether two elements will bond ionically or covalently from their positions in the periodic table — a metal with a nonmetal generally bonds ionically, and a nonmetal with a nonmetal bonds covalently.
Check your understanding
Calcium is a metal and iodine is a nonmetal. Predict the bonding between calcium and iodine. Ionic or covalent?
AIonic.correct
BCovalent.
This option is wrong — you applied the nonmetal-with-nonmetal rule — calcium is a metal, and a metal with a nonmetal generally bonds ionically.
CMetallic.
This option is wrong — you applied the metal-with-metal case — only one of the two elements is a metal, so the pair bonds ionically, not metallically.
DNone of the above — a metal and a nonmetal cannot bond.
This option is wrong — you ruled the pair out — metal-with-nonmetal is the classic ionic pairing.
Calcium: metal. Iodine: nonmetal. A metal with a nonmetal generally bonds ionically. So the prediction is ionic.
Check your understanding
Nitrogen and chlorine are both nonmetals. Predict the bonding between nitrogen and chlorine. Ionic or covalent?
ACovalent.correct
BIonic.
This option is wrong — you applied the metal-with-nonmetal rule — both elements are nonmetals, and nonmetal with nonmetal bonds covalently.
CMetallic.
This option is wrong — you applied the metal-with-metal case — neither element is a metal, so no electron sea forms; the atoms share pairs.
DNone of the above — two different nonmetals cannot bond to each other.
This option is wrong — you required identical atoms — different nonmetals share electron pairs just as identical ones do.
Nitrogen: nonmetal. Chlorine: nonmetal. A nonmetal with a nonmetal bonds covalently. So the prediction is covalent.
Check your understanding
Two compounds form: one from barium and oxygen, the other from phosphorus and oxygen. Barium is a metal; phosphorus and oxygen are nonmetals. Which pair of predictions is correct?
This option is wrong — you swapped the two rules — the metal-with-nonmetal pair is the ionic one, and the two-nonmetal pair is covalent.
CBoth ionic.
This option is wrong — you let oxygen force every compound ionic — the partner matters, and phosphorus with oxygen is nonmetal with nonmetal: covalent.
DBoth covalent.
This option is wrong — you treated barium as a nonmetal — barium sits on the far left of the table, a metal, so barium with oxygen is ionic.
Same oxygen, different partners. Barium (metal) with oxygen (nonmetal): ionic. Phosphorus (nonmetal) with oxygen (nonmetal): covalent.
Lesson 52 of 55 · IMB-052
Predict bond type from electronegativity
Learning Journey
Read · Check · Unlock
Did You Know?
Wonder this:
The position rule says 'generally'. When you want firmer evidence than a glance at the table, numbers are available.
You've already seen electronegativity: how strongly an atom pulls shared electrons toward itself. Each element has a value, and the values decide bond type more directly than position.
The idea
Find the two elements' electronegativity values and take the difference.
A large difference means the stronger puller strips the electrons away completely — electron transfer, so the bonding is ionic.
A small difference means neither atom can strip electrons from the other — the atoms share, so the bonding is covalent.
Sodium (0.9) and chlorine (3.0): difference 2.1 — large, so chlorine takes the electron and the bonding is ionic.
Chlorine (3.0) and chlorine (3.0): difference 0.0 — small, so the atoms share and the bonding is covalent.
You will only be asked about clearly large and clearly small differences — no single cutoff number decides the borderline cases.
Worked examples
Worked example 1. Calcium's electronegativity is 1.0 and fluorine's is 4.0. Predict the bonding between calcium and fluorine.
Step 1
Difference: 4.0 − 1.0 = 3.0.
Step 2
A difference of 3.0 is large — the fluorine strips the electrons away: transfer.
Step 3
Transfer makes ions, and the bonding is ionic.
Step 4
Large difference (3.0) → electron transfer → ionic bonding.
Worked example 2. Carbon's electronegativity is 2.5 and hydrogen's is 2.1. Predict the bonding between carbon and hydrogen.
Step 1
Difference: 2.5 − 2.1 = 0.4.
Step 2
A difference of 0.4 is small — neither atom strips electrons from the other: sharing.
Step 3
Sharing pairs is covalent bonding.
Step 4
Small difference (0.4) → electron sharing → covalent bonding.
You can now predict that a large electronegativity difference between two bonded atoms leads to electron transfer (ionic bonding) and a small difference leads to electron sharing (covalent bonding), without using an exact numerical cutoff.
Check your understanding
Lithium's electronegativity is 1.0 and fluorine's is 4.0. Predict the bonding between lithium and fluorine. Ionic or covalent?
AIonic.correct
BCovalent.
This option is wrong — you read a large difference as sharing — a difference of 3.0 is large, and a large difference means the stronger puller strips the electrons away: transfer, so ionic.
CMetallic.
This option is wrong — you reached for the electron-sea category — electronegativity differences sort ionic from covalent; a metal-nonmetal transfer pair is ionic.
DNone of the above — electronegativity values cannot predict bond type.
This option is wrong — you dismissed the method — the difference is exactly what predicts transfer versus sharing.
Difference: 4.0 − 1.0 = 3.0. A difference of 3.0 is clearly large — fluorine strips the electron away: transfer. Transfer makes ions, so the bonding is ionic.
Check your understanding
Nitrogen's electronegativity is 3.0 and oxygen's is 3.5. Predict the bonding between nitrogen and oxygen. Ionic or covalent?
ACovalent.correct
BIonic.
This option is wrong — you read a small difference as transfer — a difference of 0.5 is small, so neither atom strips electrons from the other and the atoms share: covalent.
CMetallic.
This option is wrong — you reached for the electron-sea category — two nonmetals with a small difference share electron pairs; no electron sea forms.
DNone of the above — atoms with different electronegativities cannot bond.
This option is wrong — you required equal values — a small difference still allows sharing; only a large difference tips the pair into transfer.
Difference: 3.5 − 3.0 = 0.5. A difference of 0.5 is clearly small — neither atom strips electrons from the other. The atoms share a pair, so the bonding is covalent.
Check your understanding
Four element pairs are listed with their electronegativity values. Which pair is predicted to bond ionically?
ABarium (0.9) and chlorine (3.0).correct
BSulfur (2.5) and chlorine (3.0).
This option is wrong — you called a small difference ionic — 3.0 − 2.5 = 0.5 is clearly small, so this pair shares: covalent.
CCarbon (2.5) and sulfur (2.5).
This option is wrong — you called a zero difference ionic — equal values mean neither atom can strip electrons from the other: covalent.
DHydrogen (2.1) and bromine (2.8).
This option is wrong — you called a small difference ionic — 2.8 − 2.1 = 0.7 is clearly small, so this pair shares: covalent.
Take each difference: Ba–Cl 2.1, S–Cl 0.5, C–S 0.0, H–Br 0.7. Only 2.1 is clearly large. A large difference means transfer, so barium and chlorine bond ionically.
Lesson 53 of 55 · IMB-053
Molecular substances: the third profile
Learning Journey
Read · Check · Unlock
Did You Know?
Wonder this:
A tealight candle turns to liquid wax from the heat of its own small flame. Salt would need a furnace four times hotter than a kitchen oven. Same job — melting — wildly different price.
Wax belongs to a third family of substances, different from ionic compounds and metals.
The idea
Some substances are made of separate molecules — small groups of atoms; these are 'molecular substances'.
Molecular substances generally melt at low temperatures — candle wax melts at about 55 °C.
Molecular substances do not conduct electricity in any state: not as solids, not melted, and not dissolved.
Compare the profiles: ionic compounds melt high and conduct when molten or dissolved; metals conduct as solids; molecular substances melt low and never conduct.
Why molecular substances behave this way is a later story — for now, the profile itself is the tool.
Worked examples
Worked example 1. Naphthalene, the strong-smelling molecular substance in old-fashioned mothballs, is tested. Predict its melting behaviour and its conductivity.
Step 1
Naphthalene is a molecular substance.
Step 2
Molecular substances generally melt low — naphthalene melts at 80 °C.
Step 3
Molecular substances never conduct: solid, melted, or dissolved.
Step 4
Low melting point (80 °C) and no conduction in any state.
Worked example 2. Table sugar is a molecular substance that dissolves well in water. A conductivity tester is placed in the sugar solution. Predict the bulb's behaviour, and contrast it with salt water.
Step 1
Sugar is molecular, and molecular substances do not conduct even when dissolved.
Step 2
The bulb stays dark in sugar water.
Step 3
Salt water lights the bulb, because salt is ionic and dissolving frees its ions.
Step 4
Sugar water leaves the bulb dark; dissolving changes nothing for a molecular substance.
You can now state that substances made of separate molecules (molecular substances) generally have low melting points and do not conduct electricity in any state.
Check your understanding
Glucose is a molecular substance. A conductivity tester is placed in a beaker of molten glucose. What does the bulb do?
AIt stays dark.correct
BIt lights, because melting frees charged particles to move.
This option is wrong — you applied the ionic rule — melting frees ions in an IONIC compound, but a molecular substance has no ions to free, so it never conducts.
CIt lights, because every liquid conducts electricity.
This option is wrong — you made all liquids conductors — molten molecular substances carry no current; molten ionic compounds and metals do.
DIt flickers as the glucose cools and hardens.
This option is wrong — you gave the substance partial conduction — a molecular substance does not conduct in any state, warm or cool.
Glucose is molecular. Molecular substances do not conduct in any state — solid, melted, or dissolved. So the bulb stays dark even in the molten liquid.
Check your understanding
Ice is solid water, a molecular substance, and it melts at 0 °C. Which melting-point pattern does water fit?
AThe molecular pattern — molecular substances generally melt at low temperatures.correct
BThe ionic pattern — compounds of two elements melt at many hundreds of degrees.
This option is wrong — you matched water to the ionic profile — water is made of molecules, not ions, and its 0 °C melting point is exactly the low molecular pattern.
CNo pattern — a melting point of 0 °C means water never truly freezes.
This option is wrong — you misread the value — 0 °C is a genuine, low melting point; below it, water is solid ice.
DThe metallic pattern — water melts low the way metals do.
This option is wrong — you built a metallic pattern from one case — metals span low to very high melting points; the reliably LOW melters are the molecular substances.
Water is made of separate molecules — a molecular substance. Molecular substances generally melt at low temperatures. An everyday freezer reaches water's melting point; no furnace needed.
Check your understanding
Ethanol is a molecular substance. Ethanol is mixed thoroughly into a beaker of water, and a conductivity tester is placed in the mixture. What does the bulb do, and what would salt do in its place?
AThe bulb stays dark with ethanol; the same test with dissolved salt would light it.correct
BThe bulb lights with ethanol, just as it would with salt.
This option is wrong — you let dissolving create conduction for any substance — dissolving only frees charge carriers that exist, and a molecular substance has none.
CThe bulb stays dark in both cases, because water shuts down all conduction.
This option is wrong — you turned water into an insulator for everything — dissolved salt conducts; it is the molecular ethanol that adds no charge carriers.
DThe bulb lights with ethanol but would stay dark with salt.
This option is wrong — you swapped the two families — the ionic compound conducts when dissolved and the molecular substance never does.
Ethanol is molecular: no conduction in any state, dissolved included. Salt is ionic: dissolved, it conducts. Same water, same tester — the family of the dissolved substance decides.
Lesson 54 of 55 · IMB-054
Argue the bond type
Learning Journey
Read · Check · Unlock
Did You Know?
Bond type can be predicted two ways: from the elements' positions, and from their electronegativity values. A full argument uses both and says so.
The idea
Make the call twice — once from positions, once from electronegativity values.
Position evidence: a metal with a nonmetal points ionic; a nonmetal with a nonmetal points covalent.
Electronegativity evidence: a large difference points to transfer, so ionic; a small difference points to sharing, so covalent.
A complete argument states the call and both pieces of evidence.
Worked once in full — magnesium chloride: magnesium is a metal and chlorine is a nonmetal, so position points ionic; the electronegativity difference is 3.0 − 1.2 = 1.8, a large difference, so transfer points ionic; both lines agree — the bonding is ionic.
Worked examples
Worked example 1. Judge whether the bonding between phosphorus and chlorine is ionic or covalent, and defend the call with both lines of evidence. (Electronegativities: P 2.1, Cl 3.0.)
Step 1
Position: phosphorus is a nonmetal and chlorine is a nonmetal — nonmetal with nonmetal points covalent.
Step 2
Electronegativity: 3.0 − 2.1 = 0.9, a small difference — neither atom strips electrons from the other, so sharing.
Step 3
Both lines point the same way.
Step 4
The bonding is covalent: two nonmetals, and a small electronegativity difference of 0.9.
Worked example 2. Judge whether the bonding between barium and fluorine is ionic or covalent, and defend the call with both lines of evidence. (Electronegativities: Ba 0.9, F 4.0.)
Step 1
Position: barium is a metal and fluorine is a nonmetal — metal with nonmetal points ionic.
Step 2
Electronegativity: 4.0 − 0.9 = 3.1, a large difference — fluorine strips the electrons away, so transfer.
Step 3
Both lines point the same way.
Step 4
The bonding is ionic: metal with nonmetal, and a large electronegativity difference of 3.1.
Worked example 3. A student writes, 'Nitrogen and hydrogen bond covalently because both are gases.' Improve the argument. (Electronegativities: N 3.0, H 2.1.)
Step 1
The call is right but the evidence is not — being a gas is not bonding evidence.
Step 2
Position: nitrogen and hydrogen are both nonmetals — points covalent.
Step 3
Electronegativity: 3.0 − 2.1 = 0.9, a small difference — points to sharing, so covalent.
Step 4
The bonding is covalent: two nonmetals, and a small electronegativity difference of 0.9 — that is the defensible argument.
You can now judge whether the bonding between two given elements will be ionic or covalent, and support the judgment with the elements' positions and supplied electronegativity values.
Check your understanding
Which argument correctly classifies the bonding between calcium and bromine? (Electronegativities: Ca 1.0, Br 2.8.)
AIonic — calcium is a metal with a nonmetal partner, and the electronegativity difference of 1.8 is large, pointing to transfer.correct
BCovalent — calcium and bromine sit far apart in the periodic table, and far-apart elements share electrons.
This option is wrong — you invented a distance rule and reached the wrong call — far-apart here means metal with nonmetal and a large difference, both of which point to transfer: ionic.
CIonic — calcium and bromine are both metals, and metals always bond ionically with each other.
This option is wrong — you got the call right with broken evidence — bromine is a nonmetal, and two metals would bond metallically, not ionically; a defensible argument needs correct evidence.
DCovalent — the difference of 1.8 is small, so the atoms share.
This option is wrong — you misjudged the difference — 1.8 is a clearly large difference, which points to transfer and an ionic bond.
Position: calcium (metal) with bromine (nonmetal) points ionic. Electronegativity: 2.8 − 1.0 = 1.8, a large difference, points to transfer — ionic. Both lines agree: the bonding is ionic, and the argument must state both correctly.
Check your understanding
Which argument correctly classifies the bonding between carbon and sulfur? (Electronegativities: C 2.5, S 2.5.)
ACovalent — both elements are nonmetals, and the electronegativity difference of 0.0 means neither atom can strip electrons from the other, so they share.correct
BIonic — carbon and sulfur have identical electronegativities, and equal pulls mean electrons transfer freely back and forth between the atoms.
This option is wrong — you attached transfer to equal pulls — a zero difference is the clearest sharing case there is; transfer needs one atom to out-pull the other by a large margin.
CCovalent — carbon and sulfur are both solids at room temperature, and solids bond covalently.
This option is wrong — you argued from room-temperature state — the call is right, but state is not bonding evidence; positions (two nonmetals) and the 0.0 difference are.
DIonic — sulfur sits below oxygen, and oxygen forms ionic compounds.
This option is wrong — you borrowed a neighbour's behaviour — oxygen bonds ionically only with metals, and carbon is a nonmetal; the evidence for THIS pair points covalent both ways.
Position: carbon (nonmetal) with sulfur (nonmetal) points covalent. Electronegativity: 2.5 − 2.5 = 0.0, a small difference, points to sharing — covalent. Both lines agree, and both belong in the argument.
Check your understanding
A student classifies the bonding between lithium and oxygen: 'Ionic, because lithium is a metal and oxygen is a nonmetal.' The electronegativities are Li 1.0 and O 3.5. What is the strongest fair criticism of the answer?
AThe call is right, but only position evidence was given — the large difference of 2.5, pointing to transfer, belongs in the argument too.correct
BThe call is wrong — lithium and oxygen bond covalently.
This option is wrong — you overturned a correct call — metal with nonmetal and a difference of 2.5 both point ionic; the answer's weakness is incomplete evidence, not the conclusion.
CNothing is missing — one piece of evidence is always enough for a full argument.
This option is wrong — you accepted a one-legged argument — the routine makes the call twice, from positions AND from electronegativity values, and states both.
DThe argument fails because electronegativity, not position, is the only evidence that ever counts in a classification.
This option is wrong — you discarded position evidence entirely — position is genuine evidence; the complete argument uses both lines, not one.
The call — ionic — is correct. A complete argument states the call and BOTH pieces of evidence. Position: metal with nonmetal, ionic. Electronegativity: 3.5 − 1.0 = 2.5, large, transfer — ionic. Both belong in the answer.
Lesson 55 of 55 · IMB-055
Classify bonding from property data
Learning Journey
Read · Check · Unlock
Did You Know?
Wonder this:
A jar in a storeroom has lost its label. White crystals — but so are salt, sugar, and a dozen other things. You cannot see bonding. You can measure it.
You've already seen each family's property profile one property at a time. A property card puts the measurements side by side, and the card is enough to classify the substance.
The idea
Read the conductivity rows first — they separate all three families.
Conducts as a solid: metallic.
No conduction as a solid, but conduction when molten or dissolved: ionic.
No conduction in any state: molecular.
Check the call against the other rows — a high melting point fits ionic or metallic, a low one fits molecular, shattering under a hammer fits ionic, and flattening fits metallic.
When rows seem to pull in different directions, the conductivity pattern outranks the melting point — some metals melt low.
The example card: potassium chloride melts at 770 °C, does not conduct as a solid, conducts when molten and when dissolved, and shatters when struck — the ionic profile, top to bottom.
Property card: potassium chloride
Property
Measurement
Melting point
770 °C
Conducts as a solid?
no
Conducts when molten?
yes
Conducts when dissolved in water?
yes
When struck with a hammer
shatters
The ionic profile, top to bottom.
Worked examples
Worked example 1. Substance A is a white solid: melting point 993 °C; no conduction as a solid; conducts when molten; conducts when dissolved in water; shatters when struck. Classify the bonding in substance A.
Step 1
Conductivity pattern: none as a solid, but conducts molten and dissolved — the ionic pattern.
Step 2
Check the other rows: melting point 993 °C is high (fits ionic), and shattering fits ionic.
Step 3
Every row agrees.
Step 4
Substance A is ionic.
Worked example 2. Substance B is a silvery solid: melting point 232 °C; conducts as a solid; conducts when molten; does not dissolve in water; flattens when struck. Classify the bonding in substance B.
Step 1
Conductivity pattern: conducts as a solid — the metallic pattern, and the deciding row.
Step 2
The melting point, 232 °C, is low — but the conductivity pattern outranks the melting point, because some metals melt low.
Step 3
Check the last row: flattening under the hammer fits metallic.
Step 4
Substance B is metallic, despite its low melting point.
Worked example 3. Substance C is a yellow solid: melting point 115 °C; no conduction as a solid, molten, or dissolved; shatters when struck. Classify the bonding in substance C.
Step 1
Conductivity pattern: no conduction in any state — the molecular pattern.
Step 2
Check: the melting point, 115 °C, is low, which fits molecular.
Step 3
The shattering row does not overturn the call — the conductivity pattern is the separator, and brittle molecular solids exist.
Step 4
Substance C is molecular.
You can now classify a substance as ionic, metallic, or molecular from measured property data such as melting point, conductivity as a solid, molten, and dissolved, and response to hammering.
Check your understanding
The property card shows the measurements for substance X, a white solid. Classify the bonding in substance X. Ionic, metallic, or molecular?
Property card: substance X
Property
Measurement
Melting point
772 °C
Conducts as a solid?
no
Conducts when molten?
yes
Conducts when dissolved in water?
yes
When struck with a hammer
shatters
AIonic.correct
BMetallic.
This option is wrong — you overlooked the solid-state row — a metallic solid conducts as a solid, and substance X does not.
CMolecular.
This option is wrong — you stopped at the solid-state row — a molecular substance conducts in NO state, but substance X conducts when molten and dissolved, which is the ionic pattern.
DNone of the above.
This option is wrong — you rejected all three categories — the card matches the ionic profile row for row: high melting point, conduction only when molten or dissolved, shattering.
Conductivity rows: no as a solid, yes molten, yes dissolved — the ionic pattern. Check: 772 °C is a high melting point, and shattering fits ionic. Every row agrees: substance X is ionic.
Check your understanding
The property card shows the measurements for substance Y, a silvery solid. Classify the bonding in substance Y. Ionic, metallic, or molecular?
Property card: substance Y
Property
Measurement
Melting point
660 °C
Conducts as a solid?
yes
Conducts when molten?
yes
Conducts when dissolved in water?
does not dissolve
When struck with a hammer
flattens
AMetallic.correct
BIonic.
This option is wrong — you skipped the deciding row — an ionic solid does not conduct as a solid, and substance Y does.
CMolecular.
This option is wrong — you classified against every conductivity row — a molecular substance never conducts, and substance Y conducts in both tested states.
DNone of the above.
This option is wrong — you rejected all three categories — conducting as a solid is the metallic signature, and the flattening row agrees.
Conducts as a solid: yes — the metallic signature, and the deciding row. Check: 660 °C melting fits a metal, and flattening under the hammer fits metallic. Substance Y is metallic.
Check your understanding
The property card shows the measurements for substance Z, a white solid with a strong smell. Classify the bonding in substance Z. Ionic, metallic, or molecular?
Property card: substance Z
Property
Measurement
Melting point
80 °C
Conducts as a solid?
no
Conducts when molten?
no
Conducts when dissolved in water?
no
When struck with a hammer
shatters
AMolecular.correct
BIonic.
This option is wrong — you predicted molten and dissolved conduction that the card rules out — an ionic solid conducts once molten or dissolved, and substance Z conducts in no state.
CMetallic.
This option is wrong — you passed over the solid-state row — a metallic solid conducts as a solid, and substance Z does not conduct at all.
DNone of the above.
This option is wrong — you rejected all three categories — no conduction in any state plus a low melting point is the molecular profile exactly.
Conductivity rows: no, no, and no — the molecular pattern. Check: 80 °C is a low melting point, which fits molecular. Substance Z is molecular.
Summary video — Properties of ionic and metallic substances, and classifying bonding