Every reaction either releases or absorbs energy — enthalpy is how chemists measure that energy change, and energy profile diagrams make it visible. Hold a hand warmer and you're feeling a negative ΔH. Crack an instant cold pack and you're feeling a positive one.
Use the PDF for classwork, homework or revision. It includes key ideas, activities, questions, an extend task and success-criteria proof.
You've probably held a hand warmer on a cold day — it heats up without any battery or flame. And you've seen an instant cold pack snap cold the moment you crack it. Both are just chemicals reacting. Why does one reaction heat up and the other cool down?
Before we name anything — write down what you think the reaction must be doing with energy in each case. What is the system doing with energy when it gets hot? When it gets cold?
Type your initial response below — you will revisit this at the end of the lesson.
Write your initial response in your book. You will revisit it at the end of the lesson.
📚 Core Content
Enthalpy (H) is a measure of the total energy stored in the bonds of a chemical system at constant pressure — and ΔH tells us how much of that energy is transferred when a reaction occurs.
We can't measure the absolute enthalpy of any substance — but we can measure the change in enthalpy (ΔH) when a reaction occurs. ΔH is defined as the enthalpy of products minus the enthalpy of reactants:
ΔH = H(products) − H(reactants)
If products have less stored energy than reactants, the difference is released to the surroundings as heat — the reaction feels warm. If products have more stored energy, the reaction must absorb heat from the surroundings — it feels cold. Enthalpy is a state function, meaning ΔH depends only on the initial and final states, not the pathway taken. Standard enthalpy changes are measured at 25°C and 100 kPa (standard conditions), denoted ΔH°.
Whether a reaction releases or absorbs heat depends entirely on whether the products sit at lower or higher enthalpy than the reactants.
In an exothermic reaction, energy is released to the surroundings — ΔH is negative because H(products) < H(reactants). The surroundings heat up; a thermometer in the solution rises. Examples include combustion of fuels, neutralisation of strong acids and bases, and cellular respiration.
In an endothermic reaction, energy is absorbed from the surroundings — ΔH is positive because H(products) > H(reactants). The surroundings cool down; a thermometer drops. Examples include dissolving ammonium nitrate, photosynthesis, and thermal decomposition.
| Feature | Exothermic | Endothermic |
|---|---|---|
| ΔH sign | Negative (< 0) | Positive (> 0) |
| Energy flow | System → surroundings | Surroundings → system |
| Surroundings temperature | Increases (warms up) | Decreases (cools down) |
| Product energy vs reactants | Lower | Higher |
| Common examples | Combustion, neutralisation, respiration | Dissolving NH₄NO₃, photosynthesis, thermal decomposition |
An energy profile diagram is a graph of enthalpy versus reaction progress — it makes the invisible energy landscape of a reaction visible at a glance.
The x-axis is "reaction coordinate" (progress of reaction, not time). The y-axis is enthalpy (kJ mol⁻¹). Reactants start at one enthalpy level; products end at another. Between them is a peak — the transition state — representing the activation energy (Ea), the minimum energy required to break bonds and start the reaction.
Five features to label on every
A thermochemical equation is a balanced chemical equation that also states the enthalpy change for the reaction as written — change either the equation or the coefficients, and ΔH changes too.
Thermochemical equations must include:
Example:
CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l) ΔH = −890 kJ mol⁻¹
This means 890 kJ is released per mole of methane burned. Two rules apply universally:
🔧 Worked Examples
The thermochemical equation for the combustion of propane is:
C₃H₈(g) + 5O₂(g) → 3CO₂(g) + 4H₂O(l) ΔH = −2220 kJ mol⁻¹
(a) Is this reaction exothermic or endothermic? (b) How much energy is released when 2.00 mol of propane is burned? (c) Write the thermochemical equation for the reverse reaction.
A reaction has ΔH = +85 kJ mol⁻¹ and a forward activation energy of 120 kJ mol⁻¹.
(a) Is the reaction exo- or endothermic? (b) What is the activation energy for the reverse reaction? (c) Describe the key features of the energy profile diagram for this reaction.
❌ "The reaction gets hot/cold" → The surroundings change temperature. The system (the chemicals reacting) loses or gains enthalpy. Precision in this language is an HSC marker requirement.
❌ "Activation energy is the height of the peak from the x-axis" → Ea is the energy difference from the reactant level to the transition state only. The position of the reactant level on the y-axis is arbitrary.
❌ "Reversing the equation just removes the negative sign" → Reversing flips the sign — so a positive becomes negative and vice versa. ΔH = −890 becomes +890 when the equation is reversed.
🧪 Activities
| Reaction | Exo or Endo? | ΔH sign | Energy flow direction |
|---|---|---|---|
| Combustion of ethanol: C₂H₅OH(l) + 3O₂(g) → 2CO₂(g) + 3H₂O(l) | Your answer | Your answer | Your answer |
| Photosynthesis: 6CO₂(g) + 6H₂O(l) → C₆H₁₂O₆(s) + 6O₂(g) | Your answer | Your answer | Your answer |
| Dissolving NaOH(s) in water — the solution warms up | Your answer | Your answer | Your answer |
| Dissolving NH₄Cl(s) in water — the solution cools down | Your answer | Your answer | Your answer |
Type your table answers and extension response below:
Complete this table in your workbook.
1 Student A says: "For a reaction with ΔH = −60 kJ mol⁻¹ and Ea = 100 kJ mol⁻¹, the activation energy for the reverse reaction is 100 − (−60) = 160 kJ mol⁻¹." Identify the error.
2 Student B says: "I drew my energy profile diagram with the transition state peak touching the top of the y-axis to show it has maximum energy. The Ea arrow goes from the x-axis to the top of the peak." Identify the error.
3 Student C says: "I doubled the equation 2CH₄(g) + 4O₂(g) → 2CO₂(g) + 4H₂O(l) but kept ΔH = −890 kJ mol⁻¹ because the substances are the same." Identify the error.
Type your error identification below before revealing answers:
Write your corrections in your workbook.
Look back at what you wrote in the Think First section. What has changed? What did you get right? What surprised you?
Wrong: An exothermic reaction has ΔH > 0 because it releases heat.
Right: Exothermic reactions have ΔH < 0 (negative) because the system loses energy to the surroundings. Endothermic reactions have ΔH > 0 (positive) because the system gains energy. The sign convention refers to the system, not the surroundings.
5 random questions from a replayable lesson bank — feedback shown immediately
✍️ Short Answer
6. Distinguish between exothermic and endothermic reactions in terms of (a) the sign of ΔH, (b) the direction of energy flow, and (c) how each would appear on an energy profile diagram. Use a specific example of each. 4 MARKS
Type your answer below:
Answer in your workbook.
7. The thermochemical equation for the combustion of ethane is:
2C₂H₆(g) + 7O₂(g) → 4CO₂(g) + 6H₂O(l) ΔH = −3120 kJ mol⁻¹
(a) Calculate the enthalpy change when 0.500 mol of ethane is burned. (1 mark)
(b) Write the thermochemical equation for the reverse of this reaction. (1 mark)
(c) Explain why water being produced as liquid (l) rather than gas (g) affects the value of ΔH. (2 marks)
4 MARKS
Type your answer below:
Answer in your workbook.
8. Real-World Application: Instant cold packs used by sports trainers contain solid ammonium nitrate (NH₄NO₃) separated from water by a thin inner bag. When the bag is cracked, NH₄NO₃ dissolves in the water and the pack becomes very cold.
(a) Is the dissolution of ammonium nitrate exothermic or endothermic? Justify using the observation described. (2 marks)
(b) Draw a labelled energy profile diagram for the dissolution of NH₄NO₃. Include: reactants, products, transition state, Ea, and ΔH arrows. Show whether products sit above or below reactants. (3 marks)
5 MARKS
Type (a) below and sketch diagram (b) in your workbook or on graph paper:
Answer both parts in your workbook — draw the diagram carefully.
Go back to your Think First response at the top of this lesson. Now that you've studied enthalpy and energy flow:
Type your reflection below:
Write your reflection in your book.
Row 1 (ethanol combustion): Exothermic | ΔH < 0 (negative) | System → surroundings (heat released)
Row 2 (photosynthesis): Endothermic | ΔH > 0 (positive) | Surroundings (light) → system (energy stored in glucose)
Row 3 (NaOH dissolving): Exothermic | ΔH < 0 | System → surroundings (solution warms up)
Row 4 (NH₄Cl dissolving): Endothermic | ΔH > 0 | Surroundings → system (solution cools down)
Extension: Photosynthesis and respiration have equal and opposite ΔH values — they are the reverse of each other. Combustion of glucose (respiration) ΔH = −2803 kJ mol⁻¹; photosynthesis ΔH = +2803 kJ mol⁻¹. This is a direct consequence of Hess's Law (path independence of enthalpy), which you'll formalise in Lesson 8.
1. Student A: The arithmetic is actually correct but the conceptual reasoning is confused. Ea(rev) = Ea(fwd) − ΔH = 95 − (−40) = 135 kJ mol⁻¹. For an exothermic reaction (ΔH = −40), products sit below reactants, so the reverse reaction must climb a larger hill: 95 + 40 = 135 kJ mol⁻¹.
2. Student B: (1) Peak height from x-axis is meaningless — only differences between levels matter. (2) Ea arrow must start from the reactant enthalpy level, not from zero or the x-axis.
3. Student C: Doubling the equation doubles ΔH. Correct thermochemical equation: 2C₂H₆(g) → ... wait, this applies to CH₄ example. For Student C: 2 × (−890) = −1780 kJ mol⁻¹.
1. B — Endothermic: ΔH > 0 (positive), energy flows from surroundings into the system, surroundings cool down.
2. D — Ea(rev) = 95 + 40 = 135 kJ mol⁻¹. For the reverse reaction, you start from the products (40 kJ mol⁻¹ lower than reactants since ΔH = −40) and climb to the same peak — a larger gap.
3. C — Doubling the equation doubles ΔH: 2 × (−572) = −1144 kJ mol⁻¹.
4. A — Exothermic: products are below reactants. ΔH < 0 because H(products) < H(reactants).
5. B — The reverse reaction absorbs 890 kJ mol⁻¹. Reversing an equation flips the sign of ΔH: the decomposition of CO₂ and H₂O back to CH₄ and O₂ is endothermic (ΔH = +890 kJ mol⁻¹). Energy is not "released the same way" — the direction of flow reverses.
Q6 (4 marks): Exothermic: ΔH < 0; energy flows from system to surroundings; on energy profile diagram, products sit below reactants, ΔH arrow points downward; example — combustion of methane, ΔH = −890 kJ mol⁻¹ [2 marks]. Endothermic: ΔH > 0; energy flows from surroundings to system; products sit above reactants, ΔH arrow points upward; example — dissolving ammonium nitrate, ΔH > 0 [2 marks].
Q7 (4 marks): (a) Equation shows 2 mol ethane releases 3120 kJ. For 0.500 mol: ΔH = (0.500/2) × 3120 = 780 kJ released [1]. (b) 4CO₂(g) + 6H₂O(l) → 2C₂H₆(g) + 7O₂(g), ΔH = +3120 kJ mol⁻¹ (sign flips) [1]. (c) Converting water from liquid to gas requires energy (latent heat of vaporisation = 44 kJ mol⁻¹ per mol H₂O) [1]. If water were produced as gas, less energy would be released to the surroundings — ΔH would be less negative. The liquid (l) state releases the full amount including the condensation energy [1].
Q8 (5 marks): (a) Endothermic [1]. Justification: the pack becomes cold, meaning heat is being transferred from the surroundings (including the injured area) into the system — the reaction absorbs heat rather than releasing it. This means ΔH > 0 [1]. (b) Diagram requirements: reactants (solid NH₄NO₃ + liquid H₂O) at a lower enthalpy level [1]; products (NH₄⁺(aq) + NO₃⁻(aq)) at a higher enthalpy level — ΔH arrow points upward [1]; transition state at the peak above products; Ea arrow from reactant level to peak [1]. Accept: x-axis labelled "Reaction coordinate", y-axis "Enthalpy (kJ mol⁻¹)".
Climb platforms, hit checkpoints, and answer questions on Enthalpy & Energy Profile Diagrams. Quick recall from lessons 1–1.
Tick when you've finished all activities and checked your answers.