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Molecules and chemistry

October 16, 2026

Chemistry is the study of how atoms join into molecules and how those molecules swap partners. It covers nearly every material and process around you: cooking, rusting, burning fuel, charging a battery, digesting food, making plastic and making medicine. The reason it matters is that a substance's properties come from which atoms are bonded to which, and how. Carbon atoms bonded one way make diamond, the hardest natural material. Bonded another way they make graphite, soft enough to write with. Ethanol and dimethyl ether have exactly the same atoms, two carbons, six hydrogens and an oxygen, but one is a drink that boils at 78 °C and the other is a gas that boils at −24 °C.

All of chemistry runs on the outer electrons of atoms, the ones described in What matter is made of. Atoms bond by sharing or trading those electrons, and reactions break bonds and make new ones, taking in or giving out energy as they go.

Why atoms bond

Atoms bond because bonded atoms have less energy than separate ones, and anything left to itself settles into its lowest-energy state, the way a ball rolls downhill. When two hydrogen atoms come close, each electron is attracted by both nuclei instead of one, and the electrons gather in the space between the nuclei, where they pull on both at once. That shared pair is the bond, and it holds the atoms together because pulling them apart means raising the energy again.

The atoms don't merge completely, because their nuclei are both positive and repel strongly at close range. The bond settles where attraction and repulsion balance: for hydrogen, 0.074 nm apart and 4.5 eV lower than two free atoms. Chemists usually quote bond energies per mole of bonds (a mole is a fixed count of 6.022 × 10²³ things, explained below), and 4.5 eV per bond is 436 kJ per mole. Breaking a mole of hydrogen molecules into atoms takes 436 kJ, and forming them gives the same 436 kJ back as heat.

Which atoms bond, and how many bonds each makes, comes down to the outer shell. An atom is most stable with a full outer shell of 8 electrons (2 for hydrogen and helium). Noble gases like neon already have one, which is why they barely bond at all. Every other atom bonds in a way that gets it a full shell, by sharing electrons, giving them away, or taking them.

Three kinds of bond

Atoms can fill their shells by sharing electrons, by handing them over, or by pooling them, and each gives a different kind of bond.

Electronegativity decides the bond

Electronegativity is how strongly an atom pulls on the electrons in a bond. It rises toward the top right of the periodic table, because atoms there have more protons pulling on outer electrons that sit close to the nucleus, and fluorine is the strongest of all at 3.98. Metals on the left are weak, around 1.

The gap between two atoms' electronegativities predicts the kind of bond they form:

The cut-offs are rough, because the types blend into each other, but the gap tells you where a bond sits on the scale.

Counting bonds

Each atom makes as many covalent bonds as it needs to complete its outer shell. Hydrogen has 1 electron and needs 1 more, so it makes 1 bond. Oxygen has 6 outer electrons and needs 2, so it makes 2. Nitrogen needs 3, and carbon, with 4 outer electrons, needs 4. Those four numbers (H 1, O 2, N 3, C 4) are enough to draw most of the molecules in your body.

These drawings are Lewis structures. Each line is a shared pair of electrons, and the dots are lone pairs, outer electrons that belong to one atom and aren't shared. When an atom runs short of partners it shares more than one pair with the same atom. Carbon dioxide's carbon shares two pairs with each oxygen (a double bond), and the two atoms in nitrogen gas share three pairs (a triple bond). That triple bond takes 945 kJ/mol to break, which is why the nitrogen that makes up 78% of the air is so unreactive.

The shapes of molecules

Molecules have definite 3D shapes, set by one rule: the electron pairs around an atom repel each other, so they spread out as far apart as they can. This rule has a name, VSEPR (valence shell electron pair repulsion), and it predicts the shape of most small molecules from the number of pairs around the central atom.

Methane's carbon has four bonding pairs, and four pairs spread evenly around a point end up at the corners of a tetrahedron, 109.5° apart. Water's oxygen also has four pairs, two bonds and two lone pairs, so it starts from the same tetrahedron. Lone pairs sit closer to their atom and take up more room, so they squeeze the two O–H bonds together to 104.5°, and the molecule is bent rather than straight.

Shape matters because it decides whether a molecule's charges cancel. Carbon dioxide has two polar C=O bonds, but they point in opposite directions along a straight line, so their pulls cancel and the molecule as a whole is non-polar. Water's two polar bonds point at an angle, so their pulls add up, and the oxygen side of the molecule is negative while the hydrogen side is positive. Almost every unusual property of water comes from that bend.

Water and hydrogen bonds

A polar molecule's positive end attracts the negative end of its neighbours. In water the attraction is especially strong, because hydrogen is tiny and its positive charge is barely shielded. The hydrogen of one water molecule attracts a lone pair on the oxygen of the next, and that attraction is a hydrogen bond. At about 20 kJ/mol it is about a twentieth as strong as the O–H bond inside the molecule, but every water molecule can make up to four of them.

Hydrogen bonds explain why water is a liquid at room temperature at all. Hydrogen sulfide, H₂S, is the same shape and nearly twice as heavy, but sulfur is barely more electronegative than hydrogen, so H₂S makes almost no hydrogen bonds and boils at −60 °C. Water, held together by hydrogen bonds, boils at 100 °C. Hydrogen bonds also give water its high surface tension, let ice float (covered in States of matter), and hold the two strands of DNA together.

Polarity also decides what dissolves in what. Water's charged ends surround ions and other polar molecules and pull them apart from each other, so salt and sugar dissolve in water. Oil is non-polar, so water molecules gain nothing by surrounding it and stay bonded to each other instead, and the oil is pushed out into a separate layer. Chemists sum this up as "like dissolves like."

Non-polar molecules still attract each other weakly. Their electron clouds shift at random, and for an instant one side of a molecule becomes slightly negative, which pushes the electrons in its neighbour the other way, so the two attract. These van der Waals forces are weak, but they grow with the size of the molecule, which is why methane (one carbon) is a gas, petrol (about eight) is a liquid, and candle wax (about twenty-five) is a solid.

Counting atoms: the mole

Atoms are far too small to count one by one, so chemists count them in batches called moles. One mole is 6.022 × 10²³ of anything, a number chosen so that one mole of a substance weighs its molecular mass in grams. A carbon atom has a mass of 12 atomic mass units, so a mole of carbon weighs 12 g. A water molecule weighs 18 units (16 for oxygen plus 1 for each hydrogen), so a mole of water is 18 g, about a tablespoon.

The mole turns a balanced equation into quantities you can weigh out. A litre of water weighs 1,000 g, so it holds 1,000 / 18 = 55.5 moles of water molecules, which is 3.3 × 10²⁵ molecules.

Reactions rearrange atoms

A chemical reaction breaks some bonds and makes new ones, so the same atoms end up in different molecules. No atom is created or destroyed in the process, so every atom on the left of a reaction equation has to appear on the right. Writing the equation so the counts match is called balancing it.

When natural gas burns, one methane molecule reacts with two oxygen molecules:

The energy released is the difference between the bonds broken and the bonds made. Breaking a bond always costs energy, and making one always releases it, so a reaction releases heat when the new bonds are stronger than the old ones in total:

For methane, breaking four C–H bonds (413 kJ/mol each) and two O=O bonds (498 each) costs 2,648 kJ. Making two C=O bonds (799 each) and four O–H bonds (463 each) releases 3,450 kJ. The new bonds win by 802 kJ per mole of methane, which is released as heat, and since a mole of methane weighs 16 g, that is 50 MJ per kilogram. A negative ΔH like this one means the reaction gives out heat, and is called exothermic. A reaction that takes in heat, like the one that makes a cold pack go cold, is endothermic.

Reaction rates

Even a reaction that releases a lot of energy can be very slow. Natural gas and air can sit mixed for years without burning, because to react, the molecules first have to start breaking their old bonds, and that costs energy, called the activation energy, before the new bonds pay it back. A spark supplies it locally, the heat from the reaction then supplies it to the next molecules, and the flame keeps itself going.

The energy to get over the hump comes from collisions. Molecules move at hundreds of metres per second and collide billions of times a second, but only collisions energetic enough to clear the barrier lead to a reaction. Heating speeds up every molecule, and it increases the number of collisions that clear the barrier much more steeply, because those come from the high-speed tail of the speed distribution. The rate follows the Arrhenius equation, where k is the rate constant, Ea the activation energy, R the gas constant and T the absolute temperature:

For a typical activation energy of 50 kJ/mol, warming from 25 °C to 35 °C multiplies the rate by 1.9. That is the origin of the rule of thumb that 10 °C doubles a reaction rate, and the reason food keeps longer in a fridge.

A catalyst speeds up a reaction without being used up, by giving it a route with a lower barrier. It doesn't change the start or end energies, so the reaction releases the same heat, only faster. The enzymes in your body are catalysts, and so are the platinum and rhodium in a car's catalytic converter, which turn carbon monoxide and unburned fuel into carbon dioxide and water.

Many reactions also run backwards. When the forward and backward rates become equal, the amounts stop changing, and the reaction has reached equilibrium. Pushing on a reaction at equilibrium makes it shift to push back, which is Le Chatelier's principle: adding more of a reactant drives it forwards, and removing the product does the same. The Haber process uses this to make ammonia from nitrogen and hydrogen, at high pressure with an iron catalyst. Fertiliser made from that ammonia grows the food for about half the people alive today.

Acids and bases

An acid is a substance that gives hydrogen ions (H⁺, bare protons) to water, and a base is one that takes them, usually by releasing hydroxide ions (OH⁻). Hydrochloric acid, HCl, splits apart in water into H⁺ and Cl⁻. Sodium hydroxide, NaOH, splits into Na⁺ and OH⁻. Mix the two and the H⁺ and OH⁻ combine into water, leaving salt water behind. That reaction, called neutralisation, is why an antacid tablet, which is a mild base, settles an upset stomach.

The concentration of hydrogen ions varies over a huge range, so it is measured on a logarithmic scale, the pH. The pH is minus the power of ten of the hydrogen ion concentration in moles per litre:

Pure water splits very slightly into H⁺ and OH⁻, giving 10⁻⁷ mol/L of each, so its pH is 7, which is called neutral. Lower numbers are acidic and higher ones basic, and each step is a factor of ten, so stomach acid at pH 1.5 holds about 300,000 times as many hydrogen ions as pure water.

Redox: moving electrons

Many reactions work by moving electrons from one substance to another. Losing electrons is called oxidation, gaining them is called reduction, and the two always happen together, because electrons don't vanish: whatever one atom loses, another gains. Reactions like this are called redox reactions. Burning, rusting, breathing and every battery are redox reactions. When iron rusts, iron atoms lose electrons to oxygen and become iron ions, which bind with the oxygen and water into rust.

Normally the electrons pass directly between atoms that touch. A battery separates the two halves of the reaction so that the electrons have to travel through a wire to get from one to the other, and on the way they can power whatever the wire runs through.

In this zinc–copper cell, zinc gives up electrons more readily than copper does, and the difference in that readiness sets the voltage, 1.10 V. A lithium-ion battery works the same way with different materials: lithium ions shuttle through the electrolyte between a graphite electrode and a metal-oxide electrode, while the electrons go round the outside circuit. Each cell gives about 3.7 V, and the best store about 250 Wh per kilogram. Petrol holds about 12,800 Wh per kilogram, fifty times as much. An electric motor turns about 90% of its energy into motion and a petrol engine about 30%, which narrows the gap. How the electrons in the wire actually carry that energy is the subject of Electricity.

Carbon chemistry

Carbon makes four bonds, and it bonds strongly to itself, so carbon atoms can link into chains, branches and rings of almost any length. That makes millions of distinct carbon compounds possible, more than every other element combined, and it is why life is built on carbon. The chemistry of carbon compounds is called organic chemistry.

The simplest are hydrocarbons, chains of carbon with hydrogen filling the remaining bonds: methane, propane, petrol and diesel are all hydrocarbons of increasing length. Polymers are very long chains made by linking small molecules end to end. Polyethylene, the plastic of shopping bags and milk bottles, is a chain of thousands of CH₂ units. Living things use the same idea: proteins are chains of amino acids, starch and cellulose are chains of sugars, and DNA is two chains of nucleotides held together by hydrogen bonds.