A candle burns, iron rusts, bread rises, a battery runs down. Each is a chemical reaction, in which starting substances, called reactants, turn into new substances, called products, as old bonds break and new bonds form. The signs are often visible: bubbles of gas, a change of color, heat or light, a solid appearing in a liquid. A chemical equation is a compact record of a reaction. Because atoms are neither created nor destroyed, the same atoms must appear on both sides, only in new arrangements.

Energy in, energy out

Breaking a bond takes energy, and forming a bond releases it. Whether a reaction gives off heat overall depends on the balance. A reaction is exothermic if it releases more energy than it takes in, as burning does, and endothermic if it takes in energy from its surroundings. A campfire is exothermic. Mix baking soda with vinegar and the fizzing mixture turns slightly cold: that reaction draws heat from its surroundings, so it is endothermic. Much of what we call energy, in food, fuel and batteries, is energy stored in the arrangement of chemical bonds.

Getting started

Burning paper releases plenty of energy, yet a book on a shelf does not burst into flame. Activation energy is the push needed to start a reaction, which is why paper does not burn at room temperature even though burning releases energy. Before new bonds can form, some old bonds must be loosened, and that costs energy up front. A match supplies it. Once burning starts, the heat it releases provides the push for the next bit of paper, and the reaction keeps itself going.

How fast

Some reactions finish in a millionth of a second; others, like rusting, take years. For most reactions, the rate rises with higher temperature, higher concentration of the reactants and more surface area. Particles must collide to react, and they must collide hard enough. Heat makes them move faster and hit harder. Crowding them together means more collisions. Grinding a solid into powder exposes far more of it to attack.

The fourth lever is a catalyst. A catalyst speeds up a reaction by offering an easier pathway with a lower activation energy, and it is not used up. The catalytic converter in a car uses platinum, palladium and rhodium to turn poisonous carbon monoxide and other exhaust gases into less harmful ones. Enzymes are the catalysts of living things. Roughly speaking, most of the chemical industry runs on catalysts.

Oxidation and reduction

At first, oxidation meant combining with oxygen, as in rusting or burning. Chemists then realized the key event is the movement of electrons. Oxidation is the loss of electrons and reduction is the gain of electrons, and the two always happen together, since electrons lost by one substance must go somewhere. These reactions, often called redox, power batteries, which let electrons flow from one substance to another through a wire, and they drive the breathing of your cells.

Back and forth

Many reactions can run in both directions. Products can react to rebuild reactants. In a closed container, such a reaction reaches a balance point. At equilibrium the forward and reverse reactions run at equal rates, so the amounts stop changing even though reactions continue. It is like a crowded shop where people enter and leave at the same rate: the number inside stays steady while individuals keep moving.

In 1884 the French chemist Henri Le Chatelier described what happens when such a balance is upset. In modern words, if a system at equilibrium is disturbed, it shifts in the direction that partly counteracts the change. Add more of a reactant and the reaction makes more product to use it up. Raise the pressure on a reaction between gases and it shifts toward the side with fewer gas molecules. Industrial chemists use this principle to push reactions toward the products they want.

Worked example: when flour explodes

On 2 May 1878 the Washburn A Mill in Minneapolis, one of the largest flour mills in the United States, blew apart. The blast killed 18 workers, fires spread to neighboring mills, and the city lost a large share of its milling capacity in a moment. The fuel was flour.

A sack of flour is hard to set alight. Flour is mostly starch, and a heap of it burns slowly because only its surface meets the air. But in the mill, fine flour dust hung in the air. Each tiny particle was surrounded by oxygen, so the total surface open to reaction was enormous. A spark, probably from millstones running dry, supplied the activation energy. The burning of one particle heated its neighbors, which caught in a fraction of a second, and the flame swept through the cloud faster than the heat could escape.

The same chemistry explains explosions in coal mines, grain elevators and sugar refineries, including the 2008 blast at a sugar refinery near Savannah, Georgia, that killed 14 people. After 1878, mills began installing ventilation to remove dust. The safety lesson rests on one idea: a substance that is harmless in a heap can be deadly as a cloud, because surface area controls how fast it reacts.