Why does a firework explode in a flash while iron rusts over years, and why does a hand warmer heat up while a cold pack cools? Two ideas, collision theory and bond energy, explain both.
Some chemical reactions happen in a flash, such as a firework exploding. Others take years, such as iron slowly turning to rust. Some reactions make the surroundings hot, such as a burning candle, and others make them cold, such as the chemicals in an instant cold pack. Two ideas explain all of this: the rate of a reaction, which tells us how fast it goes, and the energy change, which tells us how much energy it takes in or gives out. This guide explains both, shows how to measure and calculate them, and links them to everyday life, with each scientific word explained.
Reactions happen when particles collide with enough energy. The rate rises with higher concentration, higher temperature, larger surface area and the use of a catalyst, because these make successful collisions more frequent or easier. Every reaction either releases energy (exothermic) or absorbs it (endothermic): breaking bonds takes in energy and making bonds gives it out. Energy changes can be measured by temperature change and calculated from bond energies.
The rate of reaction is a measure of how quickly the reactants are used up or how quickly the products are made. A fast reaction has a high rate, and a slow reaction has a low rate. We can write it as:
Rate = change in amount ÷ time taken
The unit of rate depends on what you measure, for example g/s (grams per second) or cm³/s (cubic centimetres per second). If you need a reminder about reactions and equations, read our article Chemical Reactions and Equations.
To follow a reaction, we choose something that changes steadily and can be measured. Four common methods are shown below.
| Method | What you measure | Suits reactions that... |
|---|---|---|
| Collect the gas | Volume of gas produced over time, using a gas syringe or an upturned measuring cylinder in water | Produce a gas, such as acid with a metal or a carbonate |
| Measure mass loss | The mass on a balance as the gas escapes | Release a gas that is not very soluble |
| Time a cloudy change | How long a solid (precipitate) takes to hide a cross under the flask | Produce a solid that makes the liquid cloudy |
| Observe a colour change | Time for a colour to appear or vanish, or readings from a colour meter | Involve a change of colour |
Chemistry experiments can involve glassware, hot water, acids and gases, so always follow your teacher's instructions and wear eye protection. Our article on lab safety explains the basics.
A student reacts magnesium with acid and collects the hydrogen gas. The total volume is recorded every 10 seconds.
| Time (s) | Total gas (cm³) | Gas made in this 10 s (cm³) | Rate (cm³/s) |
|---|---|---|---|
| 0 | 0 | – | – |
| 10 | 22 | 22 | 2.2 |
| 20 | 38 | 16 | 1.6 |
| 30 | 48 | 10 | 1.0 |
| 40 | 54 | 6 | 0.6 |
| 50 | 58 | 4 | 0.4 |
| 60 | 60 | 2 | 0.2 |
Average rate over the whole experiment: 60 cm³ ÷ 60 s = 1.0 cm³/s.
The rate is highest at the start (2.2 cm³/s) and falls steadily, because the reactants are being used up and there are fewer particles to collide. On a graph of total gas against time, the line starts steep and then flattens. It becomes flat when one reactant has run out, and the reaction has stopped.
On such a graph, a steeper line means a faster rate. If you compare two experiments on the same graph, the steeper line belongs to the faster reaction, and both lines level off at the same height if they started with the same amount of the limiting reactant. For practice with reading graphs, see Graphs and Pie Charts.
The collision theory explains reaction rates. It says that a reaction occurs only when reacting particles collide, and only some collisions lead to a reaction. A successful (effective) collision needs two things:
Most collisions do not have enough energy and simply bounce apart. So the rate of reaction depends on how often particles collide and how many of those collisions are successful. Every factor that speeds up a reaction does so by increasing one of these two things.
| Factor | Effect on particles | Effect on rate |
|---|---|---|
| Higher concentration (or higher pressure for gases) | More particles in the same space, so collisions are more frequent | Faster |
| Higher temperature | Particles move faster and have more energy: collisions are more frequent and more of them are successful | Faster |
| Larger surface area (smaller pieces) | More particles of a solid are exposed to collisions | Faster |
| A catalyst | Gives the reaction an easier path with a lower activation energy | Faster |
| Light (for some reactions) | Provides the energy for the reaction | Faster |
Concentration is the amount of a substance dissolved in a given volume. In a more concentrated solution there are more particles in every cubic centimetre, so they meet more often. For gases, increasing the pressure squeezes the particles closer together, which has the same effect.
Raising the temperature has two effects. The particles move faster, so they collide more often, and, more importantly, a larger fraction of them have at least the activation energy. For many reactions at around room temperature, the rate roughly doubles for each 10°C rise, but this is only a rule of thumb, not an exact law. Cooling slows reactions, which is why we keep food in a refrigerator: the chemical reactions that spoil food, and the growth of microorganisms, are slowed.
Only the particles at the surface of a solid can collide with a liquid or gas. Breaking a solid into smaller pieces exposes more particles. Consider a cube with sides of 1 cm.
One large cube (1 cm × 1 cm × 1 cm): it has 6 faces, each 1 cm², so the surface area is 6 × 1 = 6 cm².
Cut it into 8 small cubes (each 0.5 cm on a side): each face is 0.5 × 0.5 = 0.25 cm², so each small cube has 6 × 0.25 = 1.5 cm². Together, 8 × 1.5 = 12 cm².
The total volume is unchanged, but the surface area has doubled, so the reaction can happen about twice as fast, if other things stay the same. Powdered marble reacts with acid much faster than a large lump of marble.
This idea explains why a cloud of fine flour or coal dust in the air can explode: the huge surface area exposed to oxygen allows combustion to spread extremely rapidly. Factories that handle powders, such as flour mills and grain stores, must control dust to prevent this danger.
A catalyst is a substance that speeds up a reaction without being used up. It provides an alternative route for the reaction with a lower activation energy, so a larger fraction of collisions are successful. At the end of the reaction, the catalyst is chemically the same as it was at the start and can be used again. A catalyst does not change the amount of product that can be made, only how fast it forms.
| Catalyst | Used for |
|---|---|
| Manganese(IV) oxide | Speeding up the breakdown of hydrogen peroxide into water and oxygen |
| Iron | Making ammonia in industry (the Haber process) |
| Platinum, palladium and rhodium | Catalytic converters in car exhausts, which change harmful gases into less harmful ones |
| Enzymes | Speeding up the reactions in living things, such as digestion |
Enzymes are biological catalysts, made of protein. Each works best at a certain temperature and pH. To explore this, see Enzyme Activity and Biological Molecules, Enzymes and Food Tests.
Every chemical reaction involves energy. The energy is stored in the chemical bonds between atoms (read more in Chemical Bonding). During a reaction, the bonds in the reactants are broken, and new bonds in the products are made. The key rule is:
Breaking bonds takes in energy. Making bonds gives out energy.
Energy is never created or destroyed in a reaction. It is only transferred between the chemicals and the surroundings, a principle called the conservation of energy. Our article on energy types and transformations explains this idea more widely.
| Exothermic | Endothermic |
|---|---|
| Energy goes out to the surroundings | Energy comes in from the surroundings |
| Surroundings get warmer | Surroundings get colder |
| Making bonds releases more energy than breaking bonds takes in | Breaking bonds takes in more energy than making bonds releases |
| Examples: burning fuels, respiration, neutralisation, slow rusting | Examples: photosynthesis, heating calcium carbonate, dissolving ammonium nitrate |
| Everyday use: hand warmers, heaters | Everyday use: instant cold packs |
A helpful memory aid: exo means "out" and endo means "in", like exit and entrance. Notice that "endothermic" means that the reaction takes in energy, so the surroundings get colder. The reaction itself does not become cold.
If a reaction is exothermic in one direction, it is endothermic by the same amount in the reverse direction. A classic example is blue hydrated copper(II) sulfate. When it is heated, water is driven off and the powder turns white (an endothermic change). When water is added to the white powder, it turns blue again and gets warm (an exothermic change).
An energy profile diagram shows how the energy of the chemicals changes during a reaction. The vertical axis is energy, and the horizontal axis is the progress of the reaction. The "hump" in the middle is the activation energy, the energy barrier that must be overcome. The difference in height between reactants and products is the energy change, written ΔH (pronounced "delta H"). The symbol Δ means "change in".
Two important points follow from these diagrams. First, even an exothermic reaction needs a push to start: a match must be struck and a fuel must be lit, because the activation energy must be supplied first. Second, a catalyst lowers the height of the hump but does not change the heights of the reactants and products, so ΔH stays the same.
A bond energy is the average energy needed to break one mole of a particular bond (see The Mole and Chemical Calculations for the meaning of a mole). It is measured in kilojoules per mole (kJ/mol). The method has three steps:
H2 + Cl2 → 2HCl
Bond energies: H–H = 436, Cl–Cl = 243, H–Cl = 432 (all in kJ/mol).
Bonds broken: one H–H and one Cl–Cl: 436 + 243 = 679 kJ.
Bonds made: two H–Cl: 2 × 432 = 864 kJ.
ΔH = 679 − 864 = −185 kJ/mol. The negative sign shows the reaction is exothermic, because more energy is released than is taken in.
CH4 + 2O2 → CO2 + 2H2O
Bond energies: C–H = 413, O=O = 498, C=O = 805, O–H = 464 (kJ/mol).
Bonds broken: 4 C–H and 2 O=O: (4 × 413) + (2 × 498) = 1,652 + 996 = 2,648 kJ.
Bonds made: 2 C=O and 4 O–H: (2 × 805) + (4 × 464) = 1,610 + 1,856 = 3,466 kJ.
ΔH = 2,648 − 3,466 = −818 kJ/mol, so the reaction is exothermic.
Bond energies are average values, so a calculated answer is an estimate and may differ somewhat from a value measured in an experiment.
The energy released or absorbed by a reaction in water can be found by measuring the temperature change. This method is called calorimetry. The formula is:
Q = m × c × ΔT
A student adds an acid to an alkali in 50 cm³ of water (assume 50 g). The temperature rises from 20.0°C to 26.0°C.
ΔT = 26.0 − 20.0 = 6.0°C.
Q = 50 × 4.18 × 6.0 = 1,254 J, which is 1.254 kJ.
If 0.025 mol of reactant took part, the energy per mole is 1.254 ÷ 0.025 = 50.16 kJ/mol, or about 50.2 kJ/mol. The reaction is exothermic, so ΔH is about −50.2 kJ/mol.
This calculation assumes that 1 cm³ of water has a mass of 1 g, and that no heat is lost to the air. In practice some heat is always lost, so the result is a little too small. Using a lid and an insulated cup reduces the error. If you want to learn about the units used here, see Scientific Measurement and SI Units.
For more, see Chemistry in Everyday Life.
A good experiment changes only one thing at a time. In an investigation of how temperature affects the rate of the reaction between sodium thiosulfate solution and dilute hydrochloric acid, you would measure the time for a cross under the flask to disappear as the mixture turns cloudy.
| Type of variable | In this experiment |
|---|---|
| Independent variable (what you change) | The temperature of the mixture |
| Dependent variable (what you measure) | The time for the cross to disappear |
| Control variables (keep the same) | Concentration and volume of both solutions, the flask, the person judging the cross |
Repeat each temperature at least twice and calculate a mean to improve reliability. Because the rate is the opposite of the time (a shorter time means a faster rate), you can plot 1 ÷ time to show rate on a graph. This reaction produces a small amount of sulfur dioxide gas, so it must be done only in a well-ventilated room under teacher supervision, with eye protection and with care when handling hot water.
| Misconception | Correct idea |
|---|---|
| A catalyst is used up | A catalyst is not used up, and can be reused |
| A catalyst gives more product | It only makes the same amount of product form faster |
| Breaking bonds releases energy | Breaking bonds takes in energy; making bonds releases it |
| Exothermic reactions need no energy to start | They still need activation energy, such as a flame or spark |
| Endothermic reactions are themselves cold | They take energy from the surroundings, which then cool |
| A faster reaction makes more product | Rate is how fast; the amount of product depends on the amount of reactants |
| The rate always exactly doubles for each 10°C | This is only a rough rule for many reactions |
| Heating only makes particles collide more often | It also gives more particles enough energy to react |
| Term | Meaning |
|---|---|
| Rate of reaction | How quickly reactants are used up or products are made |
| Collision theory | Reactions happen when particles collide with enough energy and the right orientation |
| Activation energy | The minimum energy needed for a collision to be successful |
| Catalyst | A substance that speeds up a reaction without being used up |
| Exothermic | A reaction that releases energy to the surroundings |
| Endothermic | A reaction that absorbs energy from the surroundings |
| ΔH | The overall energy change of a reaction |
| Bond energy | The energy needed to break one mole of a particular bond |
| Calorimetry | Measuring energy changes using temperature change |
| Surface area | The total area of the outside of a solid that can meet other substances |
Why does a lump of coal burn slowly while coal dust can explode? Why does a cold pack get cold and a hand warmer get warm? The answers come from just two ideas: particles must collide with enough energy to react, and bonds take in energy when they break and give it out when they form. Once you understand these, a very large number of everyday observations make sense, from cooking and refrigeration to engines and the chemistry of your own body. Try drawing both energy profile diagrams from memory today, and explain each one aloud in your own words.
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