Exothermic and endothermic reactions

System and surroundings, enthalpy and the sign of ΔH, exothermic and endothermic reactions and physical changes, thermochemical equations, and the practical investigation of temperature changes in combustion and in dissolving ionic substances.

Snap the metal disc inside a reusable hand warmer and it turns cloudy and hot in seconds. Squeeze an instant cold pack and it goes icy. Neither has a battery. All the energy comes from chemical changes inside the pack: one gives energy out, the other takes it in. By the end of this page you'll be able to say which is which, and why, from a thermometer reading alone.

System and surroundings

To keep track of energy, chemists split the world in two:

The system is the chemicals that are reacting: the reactants turning into products.
The surroundings are everything else: the water they are dissolved in, the beaker, the thermometer, the air and you.

Energy can move between the two as heat, but the total is conserved (the law of conservation of energy). A thermometer sits in the surroundings, so it tells you what happened to the system indirectly: if the surroundings warm up, the system must have given out heat.

In an exothermic change heat flows out of the system and the surroundings warm up. In an endothermic change heat flows in and the surroundings cool down.

Enthalpy and ΔH

Every substance stores chemical energy in its bonds and in the attractions between its particles. The total heat content of a system at constant pressure is called its enthalpy, $H$. You can't measure $H$ itself, only the change when reactants become products:

$$\Delta H = H_\text{products} - H_\text{reactants}$$

For a reaction in an open beaker (constant pressure), $\Delta H$ equals the heat absorbed by the system. Its sign is the whole story:

	Exothermic	Endothermic
Heat flow	System → surroundings	Surroundings → system
Temperature of surroundings	Rises	Falls
Enthalpy of products	Lower than reactants	Higher than reactants
Sign of ΔH	Negative	Positive
Examples	Combustion, neutralisation, respiration, metals with acid, NaOH or CaCl₂ dissolving, condensing, freezing	Photosynthesis, thermal decomposition of CaCO₃, NH₄NO₃ or KNO₃ dissolving, melting, evaporating

Enthalpy level diagrams. The arrow always points from reactants to products, so it points down when ΔH is negative and up when ΔH is positive.

The sign tells you the direction of heat flow for the system. It is easy to get backwards: an endothermic reaction feels cold to touch precisely because it is pulling heat out of your hand.

Thermochemical equations

A thermochemical equation is a balanced equation with states and the enthalpy change for the amounts written:

$$\text{CH}_4(g) + 2\text{O}_2(g) \rightarrow \text{CO}_2(g) + 2\text{H}_2\text{O}(l) \qquad \Delta H = -890 \text{ kJ mol}^{-1}$$

"Per mole" here means per mole of the reaction as written: 890 kJ is released when 1 mol of CH₄ reacts with 2 mol of O₂. Three rules follow, and you'll use them all through this module:

Double the equation, double ΔH. 2CH₄ + 4O₂ → 2CO₂ + 4H₂O has ΔH = −1780 kJ.
Reverse the equation, reverse the sign. Turning CO₂ and water back into methane would need +890 kJ.
States matter. If the water forms as steam, ΔH is only −802 kJ, because 88 kJ that would be released by condensing the steam stays in the gas.

Worked example: heat from a mass of fuel

How much heat is released when 4.00 g of methane burns completely? ΔH = −890 kJ mol⁻¹.

$n(\text{CH}_4) = \dfrac{4.00}{16.04} = 0.249$ mol

moles

Heat released $= 0.249 \times 890 = 222$ kJ

scale ΔH

222 kJ of heat is released (the system's enthalpy falls by 222 kJ).

answer

Worked example: which way is it?

When solid potassium nitrate is stirred into water, the temperature falls from 21.4 °C to 16.9 °C. Classify the change and give the sign of ΔH.

The surroundings (the water) got colder, so heat flowed from the water into the system.

observe

Heat absorbed by the system: endothermic, ΔH positive.

classify

The energy needed to separate K⁺ and NO₃⁻ ions from the lattice is more than the energy released when the ions are hydrated by water molecules.

explain

Dissolving ionic substances

Dissolving is a physical change, but it still has an enthalpy change, the heat of solution, $\Delta_\text{sol}H$. Two energy terms compete:

breaking up the ionic lattice absorbs energy (endothermic);
water molecules clustering around the separated ions (hydration) releases energy (exothermic).

Whichever is bigger decides the sign. For NaOH, CaCl₂ and anhydrous MgSO₄ hydration wins and the solution warms up. For NH₄NO₃, KNO₃ and NH₄Cl the lattice wins and the solution cools, which is how instant cold packs work. (Why they dissolve at all when it costs energy is the story of 4H and 4I.)

Practical: measuring temperature changes

Investigation: temperature changes in physical and chemical processes

Aim: to classify a set of changes as exothermic or endothermic from temperature measurements.

Method: For each solid change, put 50 mL of water in a polystyrene cup, record its temperature, add about 2 g of the solid, stir and record the highest or lowest temperature. Use NaOH, CaCl₂, NH₄NO₃ and KNO₃. For reactions, mix 25 mL of 1 mol L⁻¹ HCl with a 3 cm strip of magnesium ribbon, and 25 mL of 1 mol L⁻¹ citric acid with 2 g of sodium hydrogen carbonate. For combustion, heat 100 mL of water in a beaker over a Bunsen burner for one minute and record the rise.

Variables: independent: the substance; dependent: temperature change; controlled: volume of water, mass of solid, starting temperature, type of cup.

Typical results: NaOH +10 °C, CaCl₂ +8 °C, Mg + HCl +9 °C (exothermic); NH₄NO₃ −5 °C, KNO₃ −6 °C, citric acid + NaHCO₃ −6 °C (endothermic).

Risks: solid NaOH is corrosive and gets hot as it dissolves: wear gloves and goggles and add it to water in small amounts. Magnesium with acid gives flammable hydrogen: no flames nearby. NH₄NO₃ is an oxidiser: keep it away from fuels and heat.

Validity: this is a fair comparison only if the same amount of solid and water is used each time. To compare enthalpies properly you need moles and $q = mc\Delta T$ (4B).

InteractiveReaction heat lab

Try this

Choose Dissolve NaOH and press Add reactant. Which way do the heat arrows point, and what is the sign of ΔH?
Choose Dissolve NH₄NO₃. The temperature falls by about 5.5 °C. Is the system gaining or losing enthalpy?
Run Mg + HCl with the Insulated cup switched off. How does the curve differ, and why would the highest reading underestimate ΔT?
Try Ba(OH)₂·8H₂O + NH₄SCN. The temperature falls by more than 20 °C, to around 0 °C: this is two solids reacting with no water added at all.

Test yourself

When ammonium chloride dissolves in water, the temperature of the solution falls. Which statement is correct?
2H₂(g) + O₂(g) → 2H₂O(l), ΔH = −572 kJ mol⁻¹. What is ΔH for H₂O(l) → H₂(g) + ½O₂(g)?
Which of these changes is endothermic?
C₃H₈(g) + 5O₂(g) → 3CO₂(g) + 4H₂O(l), ΔH = −2220 kJ mol⁻¹. How much heat is released by burning 11.0 g of propane?

Label the system, the surroundings and the direction of heat flow.

Exothermic or endothermic?

The chemicals that are reacting are called the , and everything else is the . In an exothermic reaction heat is and ΔH is . The heat content of a substance at constant pressure is its .

4 marks

A student adds 5 g of ammonium nitrate to water in a polystyrene cup and the temperature falls by 4.8 °C. Explain this observation in terms of the energy changes involved in dissolving, and state the sign of ΔH.

Your answer