Enthalpy change, standard conditions and the sign convention
This unit is Topic 6, Energetics, and its theory is examined on the Unit 2 paper, WCH12 (Energetics, Group Chemistry, Halogenoalkanes and Alcohols).
Across the June 2024, January 2025 and June 2025 papers it earned about a ninth of the WCH12 marks, mostly as calculations.
Its practical side, the calorimetry and Hess's Law experiments of Core Practical 2, is examined again on the Unit 3 practical paper, WCH13.
This first section sets up the two ideas every later section leans on: what an enthalpy change is, and which sign it carries.
The idea: chemical energy you can measure as heat
When methane burns in a gas cooker, the flame heats the pan.
When ammonium nitrate dissolves in water inside a cold pack, the pack goes icy.
In both cases energy moves between the chemicals and everything around them.
Chemists split the world into two parts to keep track of it:
- the system: the reacting chemicals themselves;
- the surroundings: everything else, including the water the chemicals are dissolved in, the container, the thermometer and the air.
Energy is stored in the chemicals as the energy of their bonds and of the attractions between their particles.
Breaking a bond always takes energy in; making a bond always gives energy out.
If the bonds in the products are, overall, stronger than those in the reactants, more energy is released in making them than was taken in breaking the old ones, and the surplus leaves the system as heat.
If the products' bonds are weaker overall, the system must take the difference in from the surroundings.
Enthalpy and enthalpy change
The heat content of a system is called its enthalpy, H. You can never measure H itself, only the change in it when a reaction happens.
Learn this definition of the change precisely:
Learn these word for word: a definition is marked on its key words.
- Enthalpy change, ΔH: the heat energy change measured at constant pressure.
- Standard conditions: 100 kPa and a specified temperature, usually 298 K.
- Exothermic reaction: a reaction in which heat energy is transferred from the system to the surroundings.
- Endothermic reaction: a reaction in which heat energy is transferred from the surroundings to the system.
- Sign convention: by convention, exothermic reactions have a negative enthalpy change and endothermic reactions have a positive enthalpy change.
- Standard state: the physical state of a substance under standard conditions (for example, water is a liquid and oxygen a gas at 100 kPa and 298 K).
Why "at constant pressure"?
Most reactions you meet are done in open containers, so they happen at the pressure of the atmosphere.
If a reaction makes a gas at constant pressure, the gas has to push the atmosphere back to make room, and a little of the energy goes into that push.
Defining ΔH at constant pressure means the number you measure in an open beaker is the number in the data tables.
(Why it works, not examined: the heat change at constant pressure includes this expansion work, so it is the quantity a beaker experiment measures directly.)
Why standard conditions are needed
An enthalpy change depends on the conditions.
Water formed as steam releases less energy than water formed as a liquid, because condensing steam releases extra heat.
Pressure and temperature also make a small difference.
So that values from different laboratories can be compared and combined, they are quoted under standard conditions: 100 kPa and a specified temperature, usually 298 K (25 °C).
Each substance is in its standard state, the state it takes under those conditions: carbon as graphite, bromine as a liquid, sodium chloride as a solid.
An enthalpy change measured under standard conditions is a standard enthalpy change, written with the standard sign: ΔH⦵.
The temperature can be added as a subscript, ΔH⦵₂₉₈, but at AS it is almost always 298 K and left off.
Two details that cost marks:
- Standard pressure is 100 kPa, not 101 kPa. Older books use one atmosphere (101.325 kPa); your course uses 100 kPa.
- 298 K is the usual temperature, but the definition says "a specified temperature". A reaction does not have to be done at 298 K to have a standard value: experimental results are corrected to standard conditions.
For reactions in solution the standard concentration is 1.00 mol dm⁻³.
You need that in Topic 16 (electrode potentials); here it explains why data-book values for neutralisation assume solutions of a stated concentration.
The sign convention
The sign of ΔH is always written from the system's point of view.
- In an exothermic reaction the system loses enthalpy: the products end up with less enthalpy than the reactants. The heat goes into the surroundings, so a thermometer in the solution shows a rise in temperature. ΔH is negative.
- In an endothermic reaction the system gains enthalpy. The heat is taken from the surroundings, so the thermometer shows a fall. ΔH is positive.
This is the source of the most common confusion in the topic: the temperature goes up when ΔH is negative.
The thermometer is part of the surroundings, not the system.
It feels the heat the system has given away.
Process
Temperature of surroundings
Sign of ΔH
Example
combustion of a fuel
rises
negative
propane burning in a camping stove
neutralisation
rises
negative
hydrochloric acid with sodium hydroxide
dissolving ammonium nitrate
falls
positive
instant cold packs
thermal decomposition
heat must be supplied
positive
heating limestone to make lime
breaking any bond
heat must be supplied
positive
Cl₂(g) → 2Cl(g)
making any bond
heat given out
negative
2H(g) → H₂(g)
Reversing a reaction reverses the sign
If a reaction releases 92 kJ when it goes forwards, it must take in exactly 92 kJ when it is reversed, because the same products and reactants are involved.
Otherwise you could run it forwards and backwards and create energy from nothing.
So:
ΔH(reverse) = −ΔH(forward)
The magnitude stays the same; only the sign changes.
This simple rule is the start of Hess's Law (later in this unit), where enthalpy changes are added round cycles.
Connections
The bond-breaking and bond-making picture above is turned into numbers in the bond enthalpy sections at the end of this unit.
In Topic 9 the same energy axis gets a hump, the activation energy, which explains why an exothermic reaction such as burning methane still needs a spark.
At A Level, Topic 12 asks whether a reaction can happen at all, using entropy alongside enthalpy; ΔH alone does not decide that, as the cold pack shows (an endothermic change that happens by itself).
- In the 2024-2025 WCH12 mark schemes each scoring point earned one mark, written "(1)"; labels such as M1 and M2 only numbered the points, with no letter codes.
- Enthalpy answers needed a sign and, where the question asked, a unit (kJ mol⁻¹); a right number with the wrong sign usually kept the other marks of the calculation.
- TE (transferred error): a wrong earlier value used correctly later still earned the later marks, so every step was set out.
- Significant figures were ignored unless the question set them, except that a 1 s.f. answer lost the mark; 2 or 3 s.f. matched to the data was expected.
- Equations: correct formulae and balancing first; state symbols only when asked, and inside an enthalpy cycle the candidate completes.
- One asterisked 6-mark item per paper: six indicative points gave up to 4 marks, plus up to 2 for a linked, logically structured answer.
In four experiments a student mixes chemicals in an insulated cup and records the temperature of the mixture.
Deduce whether each change is exothermic or endothermic, and give the sign of ΔH. (4)
Experiment
Start / °C
End / °C
magnesium ribbon in dilute sulfuric acid
19.5
31.0
potassium nitrate dissolving in water
20.0
13.5
sodium hydrogencarbonate in dilute ethanoic acid
21.0
17.0
potassium hydroxide solution with dilute nitric acid
18.5
25.0
1. Magnesium and acid: temperature rises, heat given to the surroundings: exothermic, ΔH negative (1)
2. Potassium nitrate: temperature falls, heat taken from the surroundings: endothermic, ΔH positive (1)
3. Hydrogencarbonate and acid: temperature falls: endothermic, ΔH positive (1)
4. Neutralisation: temperature rises: exothermic, ΔH negative (1)
Answer: exothermic (ΔH negative): the magnesium and the neutralisation experiments; endothermic (ΔH positive): the potassium nitrate and the hydrogencarbonate experiments.
The rule to apply is: the thermometer is in the surroundings, so a rise means the system lost heat.
State whether each set of conditions is standard, giving a reason where it is not. (3)
(a) 100 kPa and 298 K; (b) 101 kPa and 25 °C; (c) 100 kPa and 20 °C.
1. (a) standard: 100 kPa and 298 K (1)
2. (b) not standard: the pressure is 101 kPa, not 100 kPa (25 °C is 298 K, so the temperature is fine) (1)
3. (c) the pressure is standard, but 20 °C is 293 K, not the usual 298 K, so a value measured here is not the usual standard value; it would have to be quoted as a standard value for 293 K (1)
Answer: only (a) is the usual standard conditions.
Convert °C to K by adding 273 before you compare: 25 °C = 298 K.
In respiration, glucose is oxidised: C₆H₁₂O₆(s) + 6O₂(g) → 6CO₂(g) + 6H₂O(l), ΔH = −2803 kJ mol⁻¹ (NBS tables, 1982).
(a) State the enthalpy change for photosynthesis, the reverse reaction, making one mole of glucose.
(b) Explain why plants need sunlight for it.
(3)
1. (a) reversing a reaction reverses the sign and keeps the size: ΔH = +2803 kJ mol⁻¹ (1)
2. (b) the reaction is endothermic, so the products have more enthalpy than the reactants (1)
3. that energy must be supplied from the surroundings; in a leaf it comes from absorbed light (1)
Answer: (a) +2803 kJ mol⁻¹; (b) photosynthesis is endothermic, so energy (light) must be taken in.
Writing the "+" is part of the answer: an unsigned 2803 does not say which way the heat flows.
When hydrogen burns, 2H₂(g) + O₂(g) → 2H₂O(g), the temperature of the surroundings rises.
Explain, in terms of bonds, why ΔH for this reaction is negative.
(3)
1. energy is taken in to break the H–H and O=O bonds in the reactants (1)
2. energy is given out when the O–H bonds in the products are formed (1)
3. more energy is given out in making bonds than is taken in breaking them, so the system loses enthalpy and ΔH is negative (1)
Answer: bond making releases more energy than bond breaking absorbs.
Note the direction of each energy change: "breaking bonds releases energy" is wrong and loses the point.
Always write the sign of an enthalpy change, including "+" for an endothermic value, and the unit kJ mol⁻¹ when the question asks for units.
When a question describes a temperature change, decide the sign from the system's point of view before you write anything: rise means negative, fall means positive.
- "The temperature went down, so ΔH is negative." It is tempting to match the sign to the thermometer, but the thermometer measures the surroundings. A fall in temperature means the system took in heat: ΔH is positive.
- Quoting standard pressure as 101 kPa or "1 atm". The definition to learn is 100 kPa and a specified temperature, usually 298 K. Learn it in those words.
- "Standard conditions means room temperature." Room temperature is about 20 °C (293 K) and varies; the standard temperature is specified, usually 298 K.
- "Breaking bonds releases energy." Bond breaking always takes energy in. The energy released in a reaction comes from making the new bonds.
Check your understanding
1. Which statement about an exothermic reaction is correct? (1)
- A The temperature of the surroundings falls and ΔH is negative.
- B The temperature of the surroundings rises and ΔH is positive.
- C The products have less enthalpy than the reactants and ΔH is negative.
- D The products have more enthalpy than the reactants and ΔH is negative. Total [1]
2. What are the standard conditions used for enthalpy changes in this course? (1)
- A 101 kPa and 273 K
- B 100 kPa and 273 K
- C 101 kPa and 298 K
- D 100 kPa and a specified temperature, usually 298 K Total [1]
3. Complete the table for three changes carried out in an insulated cup. (3)
Change
Temperature change
Exothermic or endothermic
Sign of ΔH
zinc powder added to copper(II) sulfate solution
rises by 9 °C
ammonium chloride dissolved in water
endothermic
ice melting in a drink
positive
Total [3]
4. A student writes: "When ammonium nitrate dissolves, the beaker gets cold, so the reaction gives out cold and ΔH is negative." Explain what is wrong with this statement. (2) Total [2]
5. A hand warmer contains iron powder, salt solution and air. When it is opened, the iron is oxidised and the pack becomes warm: 4Fe(s) + 3O₂(g) → 2Fe₂O₃(s).
(a) State the sign of ΔH for this reaction. (1)
(b) Explain, in terms of the system and its surroundings, why the pack becomes warm. (2)
(c) The standard enthalpy change of this reaction is −1648 kJ mol⁻¹ (NBS tables, 1982). State the enthalpy change for 2Fe₂O₃(s) → 4Fe(s) + 3O₂(g). (1)
(d) State the meaning of the term enthalpy change. (1)
Total [5]
6. Explain why enthalpy changes are quoted under standard conditions, and give one example of a change of conditions that would alter the value of ΔH for the combustion of hydrogen. (2) Total [2]
Answers
1. C (1). In an exothermic reaction the system loses enthalpy, so the products are lower. A is not correct because the surroundings warm up, not cool down. B is not correct because a rise in temperature goes with a negative ΔH, not a positive one. D is not correct because products with more enthalpy than reactants make ΔH positive.
2. D (1). A is not correct because 101 kPa is the old one-atmosphere value and 273 K is 0 °C. B is not correct because 273 K is 0 °C, not the usual 298 K. C is not correct because the pressure is 100 kPa, not 101 kPa.
3. Zinc and copper(II) sulfate: exothermic, negative (1). Ammonium chloride: temperature falls, positive (1). Ice melting: temperature (of the drink) falls, endothermic (1).
4. Heat cannot be "given out as cold": the solution cools because heat energy is transferred from the surroundings (water and beaker) to the system (1). The system gains enthalpy, so the change is endothermic and ΔH is positive (1).
5. (a) Negative (1). (b) Heat energy is transferred from the reacting iron (the system) to the surroundings, the pack and your hands (1); the products have less enthalpy than the reactants, the difference appearing as heat (1). (c) +1648 kJ mol⁻¹ (1). (d) The heat energy change measured at constant pressure (1).
6. The value of ΔH depends on the conditions, so quoting all values under the same conditions lets them be compared and combined (1). Any one: forming water as steam instead of liquid (a different state), or a different temperature or pressure (1).