Bonding, Formulas & Molecular Structure — Print · Physical Science: Chemistry, Grade 11

Ionic Bonding

Learning goals

• I can explain how atoms transfer electrons to form cations and anions, and predict a monatomic ion's charge from its group.
• I can write formulas and names for simple ionic compounds by balancing positive and negative charges.

Readings for this unit

• Why three solids need three different explanations

A chemical bond describes interactions that hold atoms or ions in a stable arrangement under suitable conditions. Bond formation can lower potential energy relative to separated particles, but whether an overall process occurs also depends on entropy, temperature, and the complete chemical system. The octet rule is a useful main-group pattern with exceptions, not a goal that atoms try to achieve. In ionic compounds, ionic bonding is electrostatic attraction throughout a lattice; electron transfer helps account for ion formation but is not itself the lattice bond.

Many common metals form cations and many common nonmetals form anions. Removing an electron from an isolated atom requires energy; it is not automatically cheap or favorable merely because the atom is a metal. Electron attachment, lattice formation, and other changes contribute to the full energy account. Losing electrons leaves a positive ion because protons outnumber electrons; gaining electrons gives a negative ion. Charge balance and measured properties help identify the resulting compound.

In a simple electron-count account for sodium chloride, a neutral sodium atom becomes Na⁺ by losing one electron and a neutral chlorine atom becomes Cl⁻ by gaining one. Na⁺ has 11 protons and 10 electrons; Cl⁻ has 17 protons and 18 electrons. Their noble-gas-like configurations are useful main-group patterns. In solid NaCl, opposite charges attract in an extended lattice. This ionic model describes charge and structure without requiring isolated atom pairs or claiming that real electron density has perfectly sharp boundaries.

It would be a mistake to picture salt as tidy little Na⁺–Cl⁻ pairs floating around. Because the attraction of a charge reaches out in every direction, each Na⁺ pulls on every nearby Cl⁻ and vice versa, and the ions stack into a vast, orderly, three-dimensional crystal lattice — in salt, each sodium is surrounded by six chlorines and each chlorine by six sodiums. A single grain of salt holds billions upon billions of ions locked in this repeating grid. That's why chemists don't call 'NaCl' a molecule; it's a formula unit, the simplest whole-number ratio of ions in the lattice.

Many ionic solids have relatively high melting temperatures because separating ions from an ordered lattice requires energy. They are often brittle: shifting layers can bring like charges closer and promote fracture. Their ions are not free to carry sustained current through an ordinary solid lattice, while a molten ionic compound or a solution containing mobile ions can conduct. Not every ionic compound dissolves readily in water, and conductivity depends on how many mobile charged particles are present.

An ionic formula balances total cation and anion charge in the simplest whole-number ratio. Mg²⁺ requires two Cl⁻ ions, giving MgCl₂. Al³⁺ and O²⁻ require a 2:3 ratio, giving Al₂O₃. A criss-cross shortcut must be followed by reduction to the simplest ratio; Mg²⁺ and O²⁻ give MgO, not Mg₂O₂. Preserve a polyatomic ion as a group and use parentheses when needed, as in Ca(NO₃)₂. Variable-charge metals require an appropriate Roman numeral in the name.

Model check: Ionic solids are extended lattices, not isolated little NaCl molecules. Ca²⁺ needs two Cl⁻ ions for neutrality; a 1:1 formula would carry +1 net charge.

Charge Balance in an Ionic Formula

(number of cations) × (cation charge) = (number of anions) × (anion charge)

an ionic compound must be electrically neutral, so the total positive charge equals the total negative charge. Example: for Al³⁺ and O²⁻, 2 × 3 = 3 × 2, giving Al₂O₃. Always reduce the subscripts to the simplest whole-number ratio.

Worked examples

Avoid these mistakes

• Ionic solids are extended lattices, not isolated little NaCl molecules.

Words to know

ionic bond

Electrostatic attraction among oppositely charged ions in the specified arrangement; ion formation by electron transfer is a distinct part of the account.

ion

An atom (or group of atoms) that carries an electric charge because it has lost or gained electrons.

cation

A positively charged ion, formed when an atom loses one or more electrons (typically a metal).

anion

A negatively charged ion, formed when an atom gains one or more electrons (typically a nonmetal).

octet rule

The tendency of atoms to gain, lose, or share electrons until their outermost shell holds a stable set of eight valence electrons.

crystal lattice

The orderly, repeating three-dimensional arrangement of ions in an ionic solid.

formula unit

The simplest whole-number ratio of ions in an ionic compound (for example, NaCl).

polyatomic ion

A charged group of covalently bonded atoms that acts as a single unit, such as nitrate (NO₃⁻) or sulfate (SO₄²⁻).

How electricity revealed the ions hiding inside salt

For centuries no one knew that tame table salt hid a violent metal and a poison gas. In 1807 the chemist Humphry Davy passed electricity through molten compounds and tore them apart, isolating pure sodium and potassium for the first time — soft, silvery metals so reactive they can burst into flame on water. The experiment worked precisely because melting frees the ions to move toward the electrodes, the very same reason molten salt conducts electricity while the solid does not.

Why a 'molecule of salt' doesn't exist

It's tempting to talk about 'a molecule of NaCl,' but there's no such thing. In a salt crystal every Na⁺ is surrounded by six Cl⁻ and every Cl⁻ by six Na⁺, an endless three-dimensional grid with no way to point to one sodium and say that chlorine belongs to it. That's why chemists use the term formula unit instead of molecule: NaCl only tells you the 1-to-1 ratio of ions, not a count of atoms locked into a separate particle.

What makes one salt melt hotter than another

Table salt melts at about 801°C, but magnesium oxide (MgO) doesn't melt until a staggering 2852°C — hot enough to line furnaces. The difference is charge. In NaCl the ions carry ±1, but in MgO they carry ±2, and the force between charges climbs steeply as the charges grow and the ions pack closer together. A stronger electrostatic attraction means a stronger lattice, which means a far higher melting point — the melting point is a direct readout of how tightly the ions are gripped.

How water pulls a crystal apart

Drop salt into water and the rigid lattice dissolves away at room temperature — how? Water molecules are polar (a fact you'll meet in the geometry lesson), with a slightly negative oxygen end and slightly positive hydrogen ends. The negative ends swarm each Na⁺ and the positive ends surround each Cl⁻, tugging the ions out of the lattice and wrapping them in a hydration shell. Once free and drifting, those ions can carry a current — which is exactly why saltwater conducts electricity and pure water barely does.

The criss-cross trap: don't forget to reduce

The criss-cross shortcut — drop each ion's charge as the other ion's subscript — is handy but has one famous trap. Cross Mg²⁺ with O²⁻ and you'd write 'Mg₂O₂,' yet the real formula is MgO. Because a formula unit is the simplest whole-number ratio, you must always reduce the subscripts to their lowest terms. Criss-cross gives you a correct ratio, but only after you divide out any common factor does it give the correct formula.

In the real world

Ionic compounds are everywhere: table salt (NaCl) seasons and preserves food, and the calcium phosphate that stiffens your bones is an ionic mineral. The dissolved ions called electrolytes — sodium, potassium, and calcium — carry the electrical signals that fire your nerves and beat your heart.

Exit check

Why does Ca form CaCl₂ rather than CaCl?

Covalent Bonding

Learning goals

• I can explain how nonmetal atoms share electron pairs to form single, double, and triple covalent bonds and complete their octets.
• I can draw simple Lewis structures and identify bonding pairs and lone pairs.

Readings for this unit

• Why three solids need three different explanations

A covalent bond involves shared electron density in an interacting atomic arrangement. A Lewis line represents a shared pair for bookkeeping, not a rigid connection or atoms consciously reaching a compromise. Covalent bonding can produce discrete molecules, such as H₂O, or extended network structures, such as diamond. The octet rule is a useful pattern for many main-group examples, with known exceptions; it does not itself supply a physical cause or a complete energy account.

Picture two hydrogen atoms drifting together. Each brings a single electron, and rather than one stealing from the other, they pool their two electrons into one shared bonding pair that both nuclei attract. That shared pair — one pair, two electrons — is a single covalent bond, often drawn as a simple line: H–H. Not every electron gets shared, though. In water, the oxygen atom shares one pair with each hydrogen but keeps two pairs entirely to itself; electrons that aren't shared are called lone pairs, and, as you'll see in the next lesson, they quietly shape the whole molecule.

A double bond is represented by two shared pairs in a simple Lewis model, and a triple bond by three. Comparable atom pairs often have shorter, stronger bonds at higher order, but this is not a universal comparison across unrelated species or a claim that bond energy is an exact multiple of pair count. Nitrogen’s strong N≡N bond helps explain some kinetic barriers, while actual reactivity still depends on pathway and conditions.

Chemists keep track of all these shared and unshared electrons with Lewis structures — simple maps that show every valence electron as a dot and every bonding pair as a line. To draw one, you add up the valence electrons from all the atoms, connect the atoms with single bonds, and then hand out the leftover electrons as lone pairs until each atom (except hydrogen, which is content with just two) is surrounded by eight. If there aren't enough electrons to give everyone an octet with single bonds, you pull lone pairs in to make double or triple bonds. The Lewis structure of water, for instance, shows two O–H bonding pairs and two lone pairs resting on the oxygen.

In many molecular substances, ordinary melting or boiling changes arrangements and intermolecular interactions while the molecules’ internal covalent connections remain intact. Extended covalent networks require another structural account. Melting, boiling and conductivity depend on size, shape, phase and electronic/ionic carriers, so low boiling or poor conduction is not a definition of every covalent substance.

So covalently bonded atoms build an astonishing range of substances through varied bonding arrangements — the O₂ you breathe, the CO₂ you exhale, the water that fills the oceans, and the enormous molecules of DNA and protein that run every living cell. Understanding *how many* pairs atoms share, and *where* the lone pairs sit, is the launch pad for the final piece of the bonding story: the three-dimensional shapes these molecules fold into, and why those shapes matter so much.

Model check: Sharing need not be equal; polar covalent bonds share electron density unevenly. Three shared pairs give six bonding electrons, not three.

Total Valence Electrons (for a Lewis structure)

total valence e⁻ = sum of every atom's valence electrons (add 1 for each unit of negative charge, subtract 1 for each unit of positive charge)

the first step in drawing a Lewis structure. Example: H₂O = 6 (from O) + 1 + 1 (from the two H) = 8 electrons = 4 pairs.

Electrons Shared in a Bond

shared electrons = 2 × (number of shared pairs)

a single bond shares 1 pair = 2 electrons, a double bond shares 2 pairs = 4 electrons, and a triple bond shares 3 pairs = 6 electrons.

Worked examples

Avoid these mistakes

• Sharing need not be equal; polar covalent bonds share electron density unevenly.

Words to know

covalent bond

Bonding involving shared electron density between atoms; Lewis lines represent shared pairs in a simplified model.

molecule

A neutral group of atoms held together by covalent bonds.

bonding pair

A pair of electrons shared between two atoms, forming a covalent bond.

lone pair

A pair of valence electrons on an atom that is not shared in a bond.

Lewis structure

A diagram that shows an atom's or molecule's valence electrons as dots and each bonding pair as a line.

double bond

A covalent bond in which two pairs of electrons (four electrons) are shared.

triple bond

A covalent bond in which three pairs of electrons (six electrons) are shared.

diatomic molecule

A molecule made of exactly two atoms, such as O₂ or N₂.

The chemist who invented the dot diagram

The shared-pair picture you use today came from Gilbert N. Lewis, who in 1916 proposed that atoms bond by sharing pairs of electrons — years before physicists fully understood what an electron even was. He first sketched electrons as dots on the corners of a cube, and that humble idea grew into the Lewis structures chemists still draw a century later. Lewis was nominated for the Nobel Prize dozens of times but never won it, one of the most famous oversights in the history of chemistry.

When the octet rule bends

The octet rule is a superb guide, not an unbreakable law. Hydrogen is satisfied with just two electrons (a 'duet'), not eight. Boron happily settles for only six electrons around it in a molecule like BF₃, staying electron-hungry. And larger atoms from the third row down, like the sulfur in SF₆, can hold more than eight in an 'expanded octet.' Knowing when the rule bends is part of thinking like a chemist rather than memorizing a slogan.

The covalent solids that refuse to melt

You just learned that molecular substances melt at low temperatures — so how can diamond, pure covalently bonded carbon, be the hardest natural material, melting only above 3500°C? The trick is that diamond isn't built from separate little molecules. Every carbon is bonded to four others in one continuous covalent network that runs through the whole crystal, so melting it means shattering countless strong bonds at once. Quartz (SiO₂) and the silicon carbide in sandpaper are network solids too — a giant exception well worth remembering.

The strongest bond in the air you breathe

About four-fifths of every breath is nitrogen gas, N₂, locked by a triple bond so strong (roughly 945 kilojoules per mole) that the gas is almost completely unreactive. That's mostly a blessing — but it also means plants can't use the nitrogen floating right above them. Only lightning and certain bacteria can crack the triple bond naturally, and in the early 1900 s the Haber–Bosch process learned to do it industrially under crushing heat and pressure. The fertilizer it makes now helps feed roughly half the people on Earth.

Why carbon builds a whole branch of chemistry

One covalent atom is so versatile that an entire science is named after it. Carbon has four valence electrons, so it forms four strong covalent bonds and, uniquely, bonds happily to itself — chaining into rings, sheets, and branches of almost unlimited length. That single talent, called catenation, is why there are millions of carbon compounds, from the methane in a stove to the DNA in your cells, and why organic chemistry exists as its own enormous field.

In the real world

The oxygen you breathe (O₂, held by a double bond) and the nitrogen that makes up most of the air (N₂, held by a very strong triple bond) are both covalent molecules. That tough nitrogen triple bond is why the gas is so unreactive — and why turning it into fertilizer to feed the world takes enormous energy.

Exit check

How many electrons are in a triple bond?

Metallic Bonding

Learning goals

• I can describe metallic bonding using the electron-sea model of positive ions in a shared pool of delocalized electrons.
• I can use the electron-sea model to explain metals' conductivity, malleability, ductility, and luster.

Readings for this unit

• Why three solids need three different explanations

You've now seen two ways atoms bond — ionic (electrostatic attraction among ions) and covalent (sharing them between two atoms). But there's a third kind, and it hides in plain sight in the spoon in your drawer and the wire in your wall: metallic bonding, the bond that holds a chunk of pure metal together. Metals are almost all cations-in-waiting — they cling to their valence electrons so loosely that, when millions of metal atoms pack together, something remarkable happens to those outer electrons.

Instead of clinging to individual atoms, valence electron density is delocalized in the simplified model and pool together into a shared, mobile crowd that belongs to the whole piece of metal at once. Chemists picture this as a 'sea of electrons': a lattice of positively charged metal ions sitting in a fluid ocean of delocalized electrons — 'delocalized' meaning the electrons aren't tied to any one atom but are free to roam everywhere. The metallic bond is the strong, reaches-in-every-direction attraction between those fixed positive ions and the negative electron sea washing around them, gluing the metal together.

This one picture explains nearly every famous property of metals. Start with electrical conductivity. Because the electrons in the sea are free to move, the tiniest push from a voltage sends them drifting through the metal — and moving electrons are exactly what an electric current is. The very same mobile electrons also ferry heat quickly from a hot spot to a cold one, which is why a metal spoon left in hot soup soon warms all the way up the handle while a plastic one stays cool.

Next come malleability and ductility — the ability to be hammered into thin sheets or drawn into long wires without shattering. Here the electron sea works like a flexible cushion. When you strike a metal, layers of positive ions slide to new positions, but the electron sea simply flows along with them and keeps holding everything together; no bond is permanently broken. Contrast that with a brittle ionic crystal, where sliding a layer lines up like charges that repel and split the crystal apart. The metal bends; the salt crystal cracks.

The mirror-bright luster of a polished metal comes from those same free electrons: they readily absorb light and re-emit it, bouncing it back to your eye as a shiny surface. And because the metallic bond reaches in every direction and grips the whole lattice at once, most metals are also strong, dense, and high-melting — you have to pour in a lot of energy to tear the ions loose from their electron sea. As a rule, the more valence electrons an atom pours into the sea (and the higher its ion's charge), the stronger the metallic bond: magnesium, giving two electrons per atom, is harder and higher-melting than sodium, which gives only one.

Pure metals are sometimes too soft or too plain for the job, so we mix them into alloys — a metal blended with one or more other elements. Adding carbon to iron makes steel, far stronger than iron alone; copper plus zinc makes brass; copper plus tin makes bronze, the alloy so important that it named an age of human history. The differently sized atoms in an alloy jam up the sliding layers and make the metal harder. From the power grid to skyscrapers to the coins in your pocket, the loose, shared electron sea of metallic bonding is quietly doing the work.

Model check: Not all metals have the same melting point or strength; electron-sea pictures are simplified. Copper has delocalized mobile electrons; salt’s ions are fixed in a lattice under ordinary conditions.

Worked examples

Avoid these mistakes

• Not all metals have the same melting point or strength; electron-sea pictures are simplified.

Words to know

metallic bond

The attraction between positively charged metal ions and the surrounding sea of delocalized electrons.

delocalized electrons

Valence electrons that are not tied to a single atom but move freely throughout a piece of metal.

electron sea

The model that pictures a metal's valence electrons as a mobile 'sea' shared by all of the metal ions at once.

malleable

Able to be hammered or rolled into thin sheets without breaking.

ductile

Able to be drawn out into a wire.

luster

The shiny appearance of a polished metal surface.

alloy

A mixture of a metal with one or more other elements, such as steel or brass.

A 'sea of electrons' proposed before the electron was even understood

The image of metal ions floating in a shared electron ocean dates to Paul Drude in 1900, just three years after the electron was discovered. His free-electron model treated a metal's outer electrons as a kind of gas free to roam the whole solid, and it neatly explained why metals conduct heat and electricity so well. Modern physics has since refined it into 'band theory,' but the simple electron-sea picture is still the first and most useful way to understand a metal.

Metals really are bonded — and strongly

A common misconception is that a lump of metal is just atoms stacked like marbles, held together by nothing but their weight. In fact the attraction between the positive ions and the surrounding electron sea is a genuine, powerful bond — which is exactly why most metals are solid, dense, and hard to melt. Tungsten doesn't melt until 3422°C, the highest melting point of any metal, which is why it is spun into the glowing filaments of old light bulbs, where anything weaker would simply vaporize.

The metal that's liquid in your hand

Not every metallic bond is equally strong, and the range is astonishing. Mercury bonds so weakly that it stays liquid at room temperature, pooling in old thermometers; gallium melts at just under 30°C, so a solid chunk turns to liquid in your palm. At the far extreme sits tungsten with its stubborn 3422°C. As a rule, the more valence electrons an atom pours into the sea — and the higher its ion's charge — the stronger the metallic bond and the higher the melting point.

Two different ways to toughen a metal

Alloys are stronger than pure metals because their mismatched atoms jam the layers so they cannot slide as freely, and there are two neat ways to arrange this. In steel, small carbon atoms squeeze into the gaps between the iron ions — an interstitial alloy. In brass, some copper ions are swapped out for similar-sized zinc — a substitutional alloy. Either way the smooth sliding that makes a pure metal soft is disrupted, and the metal grows harder and stronger.

Why wires warm up — and a few lose all resistance

If the electron sea flows so freely, why does a wire have any resistance at all? Because the positive ions are constantly vibrating, and those jiggling ions scatter the drifting electrons and slow them down. Heat the wire and the ions vibrate harder, so its resistance actually rises with temperature. Cool certain metals close to absolute zero, though, and the vibrations nearly stop — a few become superconductors with exactly zero resistance, a startling effect Heike Kamerlingh Onnes discovered in mercury in 1911.

In the real world

Metallic bonding is why copper wires carry electricity through your walls, why gold can be beaten into leaf far thinner than paper, and why a blacksmith can hammer glowing iron into a new shape without it crumbling. Alloys like steel and bronze let us tune those metallic properties for bridges, tools, and instruments.

Exit check

Why can solid copper conduct while solid salt generally cannot?

Molecular Geometry & Polarity

Learning goals

• I can use VSEPR theory to predict the shape of a simple molecule from its bonding groups and lone pairs.
• I can decide whether a bond and a whole molecule are polar or nonpolar, and name the main intermolecular forces.

Readings for this unit

• Why three solids need three different explanations

A Lewis structure drawn flat on paper hides one of the most important facts about a molecule: it is a real, three-dimensional object with a definite shape. And shape is not a minor detail — it decides how a molecule packs, dissolves, smells, and reacts, and whether a life-saving drug fits its target or misses entirely. The tool chemists use to predict that shape from a simple Lewis structure is VSEPR theory, which stands for Valence Shell Electron Pair Repulsion — a long name for one commonsense idea.

The idea is this: the groups of electrons around a central atom all carry negative charge, so they repel one another and spread out as *far apart as possible*, and that spreading fixes the molecule's shape. Count the 'electron groups' on the central atom — each single, double, or triple bond counts as one group, and so does each lone pair. Two groups push to opposite sides and give a linear shape (180°), as in carbon dioxide, O=C=O. Three groups spread into a flat triangle, or trigonal planar (120°). Four groups point to the corners of a tetrahedron (about 109.5°), the shape of methane, CH₄.

Lone pairs are the plot twist. They take up space and repel just like bonding groups — actually a little *more* strongly — but you don't 'see' them when you name the shape, because the name describes only where the *atoms* are. The oxygen in water has four electron groups (two bonds plus two lone pairs), so they arrange tetrahedrally, but with only two hydrogens attached the visible molecule is bent, its H–O–H angle squeezed to about 104.5°. The nitrogen in ammonia has three bonds and one lone pair, giving a trigonal pyramidal shape. The invisible lone pairs quietly push the atoms into these bent and pyramidal forms.

Shape becomes truly powerful when you combine it with polarity. Recall that in a covalent bond atoms share electrons — but they don't always share them *equally*. Electronegativity measures how strongly an atom pulls on that shared pair, and when two bonded atoms differ in electronegativity, the greedier one hogs the electrons, taking on a slight negative charge while its partner is left slightly positive. That lopsided bond is a polar bond, and the little separation of charge is a dipole. When the two atoms are identical, as in O₂, the sharing is perfectly even and the bond is nonpolar.

Here's the subtle part that trips up many students: a molecule can be full of polar bonds and yet be nonpolar *overall* — it all depends on the shape. Think of each polar bond as an arrow (a dipole) pointing toward the greedier atom. If the molecule is symmetric, those arrows point in balanced directions and cancel out, leaving no net pull. Carbon dioxide is the perfect example: its two C=O bonds are polar, but the molecule is linear, so the two arrows point in exactly opposite directions and cancel — CO₂ is nonpolar. Water has equally polar O–H bonds, but because it is *bent*, the arrows don't cancel; they add up to a net dipole, so water is polar. Same kind of bonds, opposite result — all because of shape.

Molecular polarity, in turn, controls the intermolecular forces — the attractions *between* whole molecules that decide whether a substance is a gas, a liquid, or a solid. The weakest, present in everything, are fleeting London dispersion forces; polar molecules feel stronger dipole–dipole attractions; and the strongest of all is the hydrogen bond, a powerful attraction between a hydrogen bonded to nitrogen, oxygen, or fluorine and a lone pair on a neighboring molecule. Water's abundant hydrogen bonds are why it boils at a startlingly high temperature for such a small molecule, why ice floats, and why it dissolves so much. A simple rule falls out of all this — 'like dissolves like': polar dissolves polar and nonpolar dissolves nonpolar, which is exactly why oil and water refuse to mix.

Model check: Polar bonds do not automatically make a polar molecule. CO₂ is linear with equal opposite bond dipoles that cancel; bent water has a nonzero molecular dipole.

Electron Groups on the Central Atom (VSEPR)

electron groups = (atoms bonded to the central atom) + (lone pairs on the central atom)

each single, double, or triple bond counts as ONE group. The number of groups sets the shape: 2 = linear, 3 = trigonal planar, 4 = tetrahedral. Example: the O in H₂O has 2 bonded atoms + 2 lone pairs = 4 groups.

Worked examples

Avoid these mistakes

• Polar bonds do not automatically make a polar molecule.

Words to know

VSEPR theory

Valence Shell Electron Pair Repulsion — the idea that electron groups around a central atom spread out as far apart as possible, which sets the molecule's shape.

molecular geometry

The three-dimensional shape formed by the atoms in a molecule.

electronegativity

A measure of how strongly an atom pulls on the shared electrons in a bond.

polar bond

A covalent bond in which electrons are shared unequally, creating a slightly positive end and a slightly negative end.

nonpolar bond

A covalent bond in which electrons are shared equally, with no separation of charge.

dipole

A separation of partial positive and partial negative charge across a bond or a whole molecule.

intermolecular force

An attraction between separate molecules; these are much weaker than the bonds within a molecule.

hydrogen bond

A strong intermolecular attraction between a hydrogen bonded to N, O, or F and a lone pair on the N, O, or F of a neighboring molecule.

The two Ronalds who named the shapes

The tidy rule that electron groups spread out to set a molecule's shape has a name and an origin: VSEPR theory, developed in 1957 by Ronald Gillespie and Ronald Nyholm. Their insight was that you don't need heavy quantum calculations to predict a molecule's geometry — you just count the electron groups around the central atom and let their mutual repulsion do the rest. It remains one of chemistry's best examples of a simple idea with enormous predictive power.

How lone pairs squeeze the angles

Lone pairs don't just bend a molecule — they push harder than bonding pairs, because a lone pair is held by only one nucleus and so spreads out more. That extra shove shows up in the numbers: methane (no lone pairs) sits at a roomy 109.5°, ammonia (one lone pair) is squeezed to about 107°, and water (two lone pairs) is pinched all the way down to 104.5°. The order of pushiness runs lone-pair–lone-pair > lone-pair–bond > bond–bond.

The bent molecule that lets fish survive winter

Water's humble bent shape has outsized consequences. Because the molecule is bent and polar, its many hydrogen bonds lock solid water into an open, roomy lattice — so ice is less dense than liquid water and floats. A freezing pond therefore forms an insulating lid of ice on top while liquid water survives beneath it, letting fish live through the winter. Almost every other substance sinks when it freezes; water's shape makes it a life-preserving exception.

Why soap can wash away grease

'Like dissolves like' says polar mixes with polar and nonpolar with nonpolar — so how does water ever wash off greasy oil? The secret is soap, whose long molecules have a polar head that loves water and a nonpolar tail that clings to grease. The tails bury themselves in the oil while the heads face the water, breaking the grease into tiny droplets that rinse away. A single cleverly shaped molecule bridges two worlds that normally refuse to mix.

Shape, polarity, and the warming planet

A molecule can trap heat only if its shape lets its charges shift as it vibrates. Symmetric O₂ and N₂ — the bulk of the air — stay perfectly balanced when they wiggle, so they let Earth's infrared heat pass straight by. But bent water, and even linear carbon dioxide when it bends and flexes, develop a changing dipole that absorbs that infrared and radiates it back down. That difference in molecular shape is precisely why CO₂ and water vapor are greenhouse gases while the main gases of the air are not.

In the real world

A molecule's shape and polarity decide how it behaves. Water's bent, polar shape lets it dissolve salts and sugars and gives it the surface tension that lets insects walk on a pond, while nonpolar oils won't mix with it — 'like dissolves like.' Drug designers obsess over molecular shape because a medicine works only if it fits the exact target in your body.

Exit check

Compare CO₂ and water polarity.

Chemical Names, Formulas & Polyatomic Ions

Learning goals

• I can write neutral ionic formulas with polyatomic ions.
• I can distinguish molecular prefixes, subscripts, and metal-charge numerals.

Readings for this unit

• Why three solids need three different explanations

A chemical formula states relative particle counts, not a recipe for mixing arbitrary amounts. An ionic compound must be electrically neutral overall. Combine charge magnitudes to find the smallest whole-number ratio, then check the total charge. Ca²⁺ and Cl⁻ give CaCl₂ because +2 + 2(−1) = 0.

A polyatomic ion is a bonded group with net charge that acts as a unit in many reactions. Common examples are ammonium NH₄⁺, hydroxide OH⁻, nitrate NO₃⁻, carbonate CO₃²⁻, sulfate SO₄²⁻, and phosphate PO₄³⁻. Parentheses preserve a group when more than one is needed: Ca(NO₃)₂ contains two nitrate ions.

Name a simple ionic compound with the cation first and anion second. A monatomic anion ends in -ide; oxygen gives oxide and chlorine chloride. Keep the established polyatomic-ion name. Metals with multiple common charges use Roman numerals to specify the cation charge: FeCl₃ is iron(III) chloride, not three iron atoms.

Binary molecular compounds usually use numerical prefixes: mono 1, di 2, tri 3, tetra 4, penta 5, hexa 6. CO is carbon monoxide and CO₂ carbon dioxide. Prefixes specify atom numbers, while Roman numerals specify ionic charge. Do not apply molecular prefixes indiscriminately to salts.

Formula subscripts count atoms inside one formula unit or molecule; a coefficient multiplies the entire formula. In 3Ca(NO₃)₂ there are 3 calcium atoms, 6 nitrogen atoms, and 18 oxygen atoms in the represented particle count. Changing a subscript changes substance identity; changing a coefficient changes amount.

Build a small ion-reference notebook and check neutrality instead of memorizing a criss-cross trick alone. Charge ratios must be reduced: Mg²⁺ and O²⁻ yield MgO, not Mg₂O₂. Naming is a way to communicate structure and composition; household product names do not necessarily identify pure substances.

Worked examples

Avoid these mistakes

• Do not break a polyatomic group apart when balancing charge.

Words to know

polyatomic ion

A bonded group of atoms with an overall charge.

Roman numeral

A notation specifying charge of a variable-charge metal cation.

subscript

A number giving atom counts within a formula.

In the real world

Use the chapter investigation to connect this unit to measured evidence and a practical decision.

Exit check

Compare coefficients, subscripts, and Roman numerals.

Lewis Structures, Resonance & Bond Models

Learning goals

• I can count valence electrons and check a Lewis structure’s total charge.
• I can distinguish resonance, electron-domain geometry, and molecular shape.

Readings for this unit

• Why three solids need three different explanations

A Lewis structure is a model of valence electrons: dots represent lone electrons and lines represent shared pairs. Count all valence electrons first; add electrons for negative charge and subtract for positive charge. Identify a central atom, usually a less electronegative non-hydrogen atom, and connect terminal atoms with single bonds.

Distribute remaining electrons to complete terminal octets, with hydrogen needing only two. Put leftovers on the central atom; if its octet is incomplete, consider a double or triple bond. Recount electrons. This procedure explains CO₂ with two double bonds and H₂O with two O–H bonds and two lone pairs on oxygen.

The octet rule is a useful pattern, not a universal law. Hydrogen follows a duet; boron compounds can have incomplete octets; odd-electron molecules exist; some third-period and heavier atom structures are conventionally drawn with expanded valence shells. A model should be applied within its scope.

Formal charge is bookkeeping: valence electrons minus nonbonding electrons minus half of bonding electrons. It is not the same as measured partial charge. For NH₄⁺, nitrogen has 5 − 0 − 4 = +1 and each hydrogen has 0 formal charge; the total is +1.

Resonance structures use the same atom arrangement but different electron placement. The actual bonding is described by a delocalized structure, not a molecule rapidly switching between drawings. In nitrate, three equivalent N–O arrangements contribute; the real bonds are equivalent in the idealized ion.

Use Lewis structure to count electron domains before predicting shape with VSEPR. A double bond counts as one direction of electron density. Separate electron-domain geometry from molecular shape, because lone pairs also occupy domains. Molecular polarity depends on the vector sum of bond dipoles and geometry.

Worked examples

Avoid these mistakes

• Formal charge is bookkeeping; resonance drawings are not separate alternating molecules.

Words to know

lone pair

A valence electron pair not represented as a bond.

formal charge

An electron-counting assignment, distinct from partial charge.

resonance

Multiple electron drawings representing delocalized bonding in one structure.

In the real world

Use the chapter investigation to connect this unit to measured evidence and a practical decision.

Exit check

How do you check a proposed Lewis structure?