A particulate diagram draws each atom as a circle and marks its outer electrons as dots.

In some diagrams the two circles overlap, and electrons sit in the overlap region.

Electrons in the overlap region belong to both atoms at once, so the diagram shows 'electron sharing'.

In other diagrams the two circles stay separate, and an electron has moved from one circle to the other.

A moved electron belongs only to the atom that received it, so the diagram shows 'electron transfer'.

To classify any diagram, find the electrons: sitting in the overlap region means sharing, moved across to one atom means transfer.

In Na–Cl, one electron moves from the sodium atom to the chlorine atom — the atoms do not share.

In Cl–Cl, the two chlorine atoms hold one pair of electrons in the overlap region — the atoms share.

In F–Cl, two different elements still hold a shared pair between them — the atoms share.

In O–O, the two oxygen atoms hold two shared pairs between them — the atoms share.

In Mg–O, two electrons move from the magnesium atom to the oxygen atom — the atoms do not share.

A bond in which two atoms share one or more pairs of electrons is called a 'covalent bond'.

Cl–Cl, F–Cl, and O–O are covalent bonds; Na–Cl and Mg–O are not, because their electrons moved instead of being shared.

Every atomic nucleus is positively charged, and electrons are negatively charged.

Opposite charges attract, so each nucleus attracts the shared electrons.

The shared pair sits between the two atoms, inside both circles at once.

Two covalently bonded atoms stay together because both nuclei attract the same shared electrons, which pulls the two atoms together.

It works like two people holding the same shopping bag — each grips the bag, so the two stay side by side.

Pulling the two atoms apart means pulling both nuclei away from electrons they attract, and that takes an energy supply from outside.

A group of atoms held together by covalent bonds is called a 'molecule'.

A water molecule is three atoms — one oxygen and two hydrogens — held together by covalent bonds.

A substance made of molecules is called a 'molecular substance'.

Water is a molecular substance: a sample of water is a crowd of separate water molecules.

A molecule can be as small as two atoms or hold thousands.

A molecular substance's formula tells you what one molecule contains.

A formula read this way is called a 'molecular formula'.

Each element's symbol appears once, and the subscript after a symbol counts that element's atoms in one molecule.

A symbol with no subscript means one atom.

CO₂ reads: one carbon atom and two oxygen atoms in one CO₂ molecule.

The counts belong to a single molecule — every molecule of the substance contains exactly these atoms.

Each counting prefix stands for one number.

mono- means 1, di- means 2, tri- means 3, tetra- means 4, and penta- means 5.

hexa- means 6, hepta- means 7, octa- means 8, nona- means 9, and deca- means 10.

penta- always signals five, whatever name it appears in.

Two pairs are easy to mix up: hexa- is 6 but hepta- is 7, and nona- is 9 but deca- is 10.

A molecular compound of exactly two elements is called a 'binary molecular compound'.

To name one from its formula, name the first element with a prefix for its count — but drop mono- on the first element.

Then name the second element with a prefix for its count, and change its ending to -ide; mono- is kept on the second element.

When a prefix ending in a or o runs into oxide, drop the prefix's last letter: mono- + oxide becomes monoxide, and tetra- + oxide becomes tetroxide.

CO₂ is carbon dioxide: mono- is dropped on the one carbon, and di- counts the two oxygens.

N₂O₄ is dinitrogen tetroxide: di- counts the two nitrogens, and tetra- + oxide gives tetroxide.

Each prefix in the name becomes a subscript in the formula.

Write the element symbols in the order the name gives them.

A prefix's number becomes the subscript after that element's own symbol.

A count of one — mono-, or a first element with no prefix — gets no subscript.

Sulfur hexafluoride: S for the one sulfur, then F with subscript 6 — SF₆.

Dinitrogen trioxide: N with subscript 2, then O with subscript 3 — N₂O₃.

Check the first element in the formula.

A metal first means an ionic compound — use ionic naming, with no prefixes.

Ionic names need no counting, because the ion charges already fix the ratio.

Nonmetals only means a molecular compound — use prefix naming.

Hydrogen counts as a nonmetal here.

CaCl₂ starts with the metal calcium, so it is calcium chloride — no prefixes.

SCl₂ starts with the nonmetal sulfur, so it is sulfur dichloride — prefixes count the atoms.

A 'Lewis diagram' draws a molecule with element symbols, lines, and dots.

Each element symbol stands for one atom of that element.

A line drawn between two symbols stands for one shared pair of electrons — two electrons, held by both atoms.

Each dot stands for one valence electron that is not shared, and unshared dots are drawn in pairs.

Only valence electrons appear — inner electrons are never drawn.

In the Lewis diagram of HCl, the line stands for the shared pair, and the three dot pairs on the Cl are chlorine's six unshared valence electrons.

An electron pair drawn as a line — shared between two atoms — is called a 'bonding pair'.

An electron pair drawn as two dots on one atom — not shared — is called a 'lone pair'.

Every electron pair in a Lewis diagram is one or the other.

In the Lewis diagram of water, the two lines are two bonding pairs, and the two dot pairs on the oxygen are two lone pairs.

To classify a pair, check where it is drawn: a line between two symbols is a bonding pair, and dots sitting on one atom are a lone pair.

Two atoms can share one, two, or three pairs of electrons.

One shared pair is a 'single bond', drawn as one line.

Two shared pairs are a 'double bond', drawn as two lines.

Three shared pairs are a 'triple bond', drawn as three lines.

The two lines in O=O show a double bond — the two oxygen atoms share two pairs.

The count of lines always equals the count of shared pairs.

Classify each bond in a diagram separately.

Count the lines between one pair of atoms: one line is a single bond, two lines a double bond, three lines a triple bond.

Dots never count toward a bond — only lines between two symbols do.

In the Lewis diagram of CO₂, two lines join the carbon to each oxygen — both carbon–oxygen bonds are double bonds.

In the Lewis diagram of N₂, three lines join the two nitrogens — one triple bond.

In most stable molecules, each atom is surrounded by eight electrons.

The twist: a shared pair counts for BOTH atoms it joins.

Each chlorine atom in Cl₂ is surrounded by eight electrons — six in its own lone pairs, plus the two shared electrons, counted for both atoms.

Hydrogen is the exception: a hydrogen atom in a molecule is surrounded by two electrons.

Two is hydrogen's full count, matching helium, just as eight matches the other noble gases.

A few molecules break the octet pattern — this course never tests them; the sentence is here only so an exception never surprises you.

To count the electrons around an atom, add its lone-pair electrons to every electron in its shared pairs.

Each lone pair on the atom contributes two electrons.

Each shared pair also contributes two electrons, and both bonded atoms count the same shared pair in full.

A double bond is two shared pairs, so it contributes four electrons, and a triple bond contributes six.

Count only the pairs touching the atom in question — lone pairs sitting on other atoms belong to those atoms.

Panel 1 of the figure shows the water Lewis diagram: the oxygen holds two lone pairs and two single bonds.

The lone pairs give 2 + 2 = 4 electrons, the two single bonds give 2 + 2 = 4 more, so the oxygen is surrounded by 8 electrons.

Each hydrogen in the same diagram shares one pair and holds no lone pairs, so each hydrogen is surrounded by 2 electrons.

In a covalent bond, the partner atom supplies one of the two shared electrons.

So every bond an atom forms raises its electron count by one.

An atom therefore typically forms one bond for every electron it needs to reach eight.

typical bonds = 8 − valence electrons

Nitrogen has five valence electrons, needs three more, and typically forms three bonds.

Carbon has four valence electrons, needs four more, and typically forms four bonds.

Hydrogen's full count is two, not eight, so hydrogen needs one electron and forms one bond.

The table gives the typical bond count for hydrogen, carbon, nitrogen, oxygen, and the halogens.

Every atom in a molecule brings its own valence electrons.

total valence electrons = (atoms of element 1 × its valence electrons) + (atoms of element 2 × its valence electrons) + …

Read each element's atom count from the formula's subscripts.

Multiply each element's atom count by the valence electrons of one atom of that element.

Add the results for every element in the formula.

Methane is CH₄: one carbon with four valence electrons, and four hydrogens with one each.

total valence electrons = (1 × 4) + (4 × 1) = 8

Count the molecule's total valence electrons first.

Draw the two element symbols side by side, joined by one line for the shared pair.

The line uses two of the total electrons.

Place every remaining electron as lone pairs around the two atoms.

Draw each lone pair as two dots side by side, pointing straight out from its atom.

Give each atom lone pairs until its count reaches eight, remembering hydrogen is already full at two.

Hydrogen chloride, HCl, has 1 + 7 = 8 valence electrons.

One line joins H and Cl, using two electrons, and the remaining six form three lone pairs on the chlorine.

Finish by checking: every electron used, hydrogen at two, chlorine at eight.

On paper, draw the Lewis diagram of iodine, I₂. Iodine is in Group 17. Then select Continue to compare your drawing with the model answer.

A molecule with three or more atoms has one atom in the middle, bonded to the others.

The middle atom is called the 'central atom'.

The central atom is the atom that can form the most bonds.

Typical bond counts decide it: carbon's four beats nitrogen's three, which beats oxygen's two, which beats a halogen's one.

Hydrogen forms only one bond, so hydrogen is never the central atom.

A shortcut: the central atom is usually the atom that appears only once in the formula.

In CF₄, carbon forms four bonds and fluorine forms one, so carbon is the central atom.

Count the molecule's total valence electrons.

Place the central atom in the middle, with the other atoms spaced around it.

Join each outer atom to the central atom with one line.

Each line uses two electrons from the total.

Place the remaining electrons as lone pairs, giving every atom eight, with hydrogen full at two.

Draw each lone pair as two dots side by side, pointing straight out from its atom.

Ammonia, NH₃, has 5 + 3 = 8 valence electrons.

Three N–H lines use six electrons, and the remaining two form one lone pair on the nitrogen.

Finish by checking: all eight electrons drawn, each hydrogen at two, the nitrogen at eight.

On paper, draw the Lewis diagram of phosphine, PH₃. Phosphorus is in Group 15 and hydrogen has one valence electron. Then select Continue to compare your drawing with the model answer.

Sometimes every electron is placed and an atom still sits below eight.

The fix: move a lone pair from a neighboring atom into the space between the two atoms, making it a second shared pair.

A second shared pair turns the single bond into a double bond, drawn as two lines.

A third shared pair makes a triple bond, drawn as three lines.

Each moved pair raises the short atom's count by two without changing the diagram's total.

Keep converting until every atom reaches eight.

Carbon dioxide, CO₂, has 16 valence electrons: two C–O lines and three lone pairs on each oxygen use all 16, but the carbon sits at only four.

Moving one lone pair in from each oxygen makes two double bonds, and every atom reaches eight.

On paper, draw the Lewis diagram of nitrogen, N₂. Nitrogen is in Group 15. Then select Continue to compare your drawing with the model answer.

Check 1: count every electron in the diagram — each line is two — and compare with the molecule's total valence electrons.

Check 2: count the electrons around each atom — eight for most atoms, two for hydrogen.

A diagram that fails either check is wrong, and the failed check names the error.

The figure shows a student's water diagram with one lone pair on the oxygen.

Check 1: two lines (4) plus one lone pair (2) is 6 electrons, but H₂O has 8 — two electrons are missing.

Check 2 agrees: the oxygen sits at six.

The diagram is wrong; the oxygen needs a second lone pair.

Step 1: count the molecule's total valence electrons.

Step 2: place the central atom, with the other atoms around it.

Step 3: join each outer atom to the central atom with a single line.

Step 4: place the remaining electrons as lone pairs, outer atoms first, until every atom has eight, with hydrogen full at two.

Step 5: if an atom is still short, convert a lone pair from a neighboring atom into an extra shared pair.

Finish with the two checks: the total matches, and every atom is full.

Methanal, H₂CO, runs the routine like this: total 1 + 1 + 4 + 6 = 12; carbon central; three single lines use 6; the remaining 6 make three lone pairs on the oxygen; the carbon sits at six, so one oxygen pair converts to a C=O double bond.

Both checks now pass: 12 electrons drawn, hydrogens at two, carbon and oxygen at eight.

On paper, use the five-step routine to draw the Lewis diagram of carbonyl sulfide, OCS — a real gas with a carbon central atom. Carbon is in Group 14; oxygen and sulfur are in Group 16. Then select Continue to compare your drawing with the model answer.

A real molecule is a three-dimensional object, and its atoms sit at definite positions.

The bonding pairs and lone pairs around a central atom are called its 'electron groups'.

Electron groups are all made of negative electrons, and like charges repel.

A molecule takes its shape because electron groups repel each other, so they take up positions as far apart as possible around the central atom.

That is why the four bonding pairs of CH₄ spread out evenly in three dimensions instead of crowding onto one side.

The balloons behave the same way: squeezed together at one point, they push one another apart until they are as spread out as they can get.

Look only at the central atom.

Count each bond as one electron group, whether it is single, double, or triple.

A double or triple bond is one group because its shared pairs sit together in the same space between the same two atoms.

Count each lone pair on the central atom as one electron group.

Lone pairs on the outer atoms do not count — they crowd their own atom, not the central one.

In CO₂, the carbon has two double bonds and no lone pairs, so the carbon has two electron groups.

Two electron groups sit on exactly opposite sides of the central atom, because electron groups repel each other, so they take up positions as far apart as possible around the central atom.

With the two groups opposite each other, the three atoms lie in a straight line.

A molecule whose central atom has two electron groups and no lone pairs is 'linear'.

Both conditions matter: count the groups, then check that none of them is a lone pair — in the molecules you will meet in this course, linear means exactly this pair of conditions.

In CO₂, the carbon has two electron groups — the two double bonds — and no lone pairs, so CO₂ is linear.

The figure shows the shape: O, C, and O in one straight line.

Three electron groups spread toward the corners of a flat triangle, because electron groups repel each other, so they take up positions as far apart as possible around the central atom.

A molecule whose central atom has three electron groups and no lone pairs is 'trigonal planar'.

'Trigonal' means three-cornered, and 'planar' means flat — the central atom and all three attached atoms lie in one flat plane.

Both conditions matter: count three groups, then check that none of them is a lone pair.

In methanal, H₂CO, the carbon has three electron groups — two single bonds and one double bond — and no lone pairs, so H₂CO is trigonal planar.

The figure shows the shape: the carbon in the middle, the two hydrogens and the oxygen at the triangle's corners.

Four electron groups spread out farthest by pointing at the four corners of a 'tetrahedron' — a solid with four triangular faces, like a pyramid on a triangular base.

They do this because electron groups repel each other, so they take up positions as far apart as possible around the central atom.

In three dimensions the four corners of a tetrahedron are farther apart than any flat arrangement allows.

A molecule whose central atom has four electron groups and no lone pairs is 'tetrahedral'.

Both conditions matter: count four groups, then check that none of them is a lone pair.

In methane, CH₄, the carbon has four electron groups — the four single bonds — and no lone pairs, so CH₄ is tetrahedral.

The figure shows the shape: the carbon in the middle, the four hydrogens at the tetrahedron's corners.

In the Lewis diagram of NH₃, the nitrogen holds three bonding pairs and one lone pair — four electron groups.

All four groups spread out, because electron groups repel each other, so they take up positions as far apart as possible around the central atom.

Only three of the four positions hold atoms — the fourth holds the lone pair.

The three hydrogen atoms sit as a low tripod with the nitrogen above them.

A molecule shaped this way is called 'trigonal pyramidal' — a pyramid with a triangle for its base and the central atom at its peak.

Classify a molecule as trigonal pyramidal when its central atom has four electron groups and exactly one of them is a lone pair.

In the Lewis diagram of H₂O, the oxygen holds two bonding pairs and two lone pairs — four electron groups.

All four groups spread out, because electron groups repel each other, so they take up positions as far apart as possible around the central atom.

Only two of the four positions hold atoms — the other two hold lone pairs.

The three atoms make a wide open V.

A molecule shaped this way is called 'bent'.

Classify a molecule as bent when its central atom has four electron groups and exactly two of them are lone pairs.

CH₄ and NH₃ both hold four electron groups around the central atom, spread as far apart as possible.

In CH₄ every group holds an atom; in NH₃ one group is a lone pair.

A shape's name describes only the positions of the atoms.

The lone pair still occupies its own position — it is an electron group like any other — but it does not appear in the name.

So NH₃ is named for what its atoms show: three hydrogens in a tripod under the nitrogen, which is trigonal pyramidal rather than tetrahedral.

The lone pair does one more thing: it pushes the three bonding pairs slightly closer together, squeezing the pyramid narrower than CH₄'s spread.

Count the electron groups around the central atom, counting each bond — single, double, or triple — as one group and each lone pair as one group.

Then count how many of those groups are lone pairs.

Two groups with no lone pair is linear.

Three groups with no lone pair is trigonal planar.

Four groups with no lone pair is tetrahedral.

Four groups with one lone pair is trigonal pyramidal.

Four groups with two lone pairs is bent.

Take H₂S: the sulfur holds two bonding pairs and two lone pairs — four groups, two of them lone pairs — so H₂S is bent.

A 'shape diagram' shows where a molecule's atoms sit: draw each atom as its element symbol, join the symbols with plain lines for the bonds, and write the shape's name under the drawing.

Leave lone pairs out of a shape diagram, because a shape's name describes only the positions of the atoms.

Linear: draw the atoms in one straight line.

Bent: draw the two bonds as a wide V with the central atom at the point.

Trigonal planar: draw the three bonds out to the corners of a flat triangle around the central atom.

Tetrahedral: draw three bonds down as a low tripod and one bond straight up.

Trigonal pyramidal: draw the central atom at the top with its three bonds going down to a low tripod — a pyramid with the central atom at the peak.

Take NH₃: its Lewis diagram shows four groups with one lone pair, so it is trigonal pyramidal — draw the N at the top, three H at the tripod's corners, and label it 'trigonal pyramidal'.

The figure shows the Lewis diagram of SiH₄: the silicon makes single bonds to four hydrogens and holds no lone pair. On paper, draw the shape diagram of SiH₄, label every atom, and write the shape's name under your drawing. Then select Continue to compare your drawing with the model answer.

Start from the formula and run the full Lewis routine: total the valence electrons, place the central atom, connect the atoms with single bonds, complete octets with lone pairs, and convert lone pairs into extra shared pairs if any octet falls short.

Then count the electron groups and the lone pairs on the central atom.

Read the shape from the counts, and draw or name it.

Take CF₄: the total is 4 + 4 × 7 = 32 valence electrons.

Carbon is the central atom, four C–F single bonds use 8 electrons, and the remaining 24 complete the fluorines' octets as lone pairs.

The carbon holds four groups and no lone pair, so CF₄ is tetrahedral.

On paper, draw the Lewis diagram of H₂Se from its formula, then classify its shape. Selenium is in Group 16 and hydrogen has one valence electron. Then select Continue to compare your work with the model answer.

Every covalent bond is a tug of war over the shared electrons.

When the two atoms have the same electronegativity, they pull equally, and the sharing is equal.

Two atoms of the same element always have the same electronegativity, so the sharing in Cl–Cl is equal.

When the two atoms have different electronegativities, the stronger puller drags the shared electrons its way, and the sharing is unequal.

Chlorine is more electronegative than hydrogen, so the sharing in H–Cl is unequal.

Classify the sharing in a covalent bond as equal when the two atoms have the same electronegativity, and as unequal when their electronegativities differ.

The size of the electronegativity difference sets the kind of covalent bond.