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Topic 3: Chemical reactions (reactants, products and energy change)

Distinguish exothermic and endothermic reactions; represent energy changes using enthalpy values (Delta H) and energy profile diagrams; calculate heat changes (q = mcDeltaT) from calorimetry data and use molar enthalpy of reaction (kJ/mol) in stoichiometric problems

A focused answer to the QCE Chemistry Unit 1 dot point on energy in chemical reactions. Distinguishes exothermic (negative Delta H) and endothermic (positive Delta H) reactions, draws energy profile diagrams with activation energy, and uses calorimetry data with q = mcDeltaT to calculate molar enthalpy of reaction.

Generated by Claude OpusReviewed by Better Tuition Academy9 min answerUpdated 2026-05-18

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What this dot point is asking

QCAA wants you to distinguish exothermic and endothermic reactions by the sign of the enthalpy change, sketch energy profile diagrams labelling activation energy and Delta H, calculate heat changes from calorimetry data using q = mcDeltaT, and convert experimental heat changes to molar enthalpy of reaction in kJ/mol. The calorimetry workflow is the most-examined practical skill in Unit 1.

The answer

A chemical reaction releases or absorbs energy as the bonds in the reactants are broken (energy absorbed) and the bonds in the products are formed (energy released). The net change in chemical potential energy is the enthalpy change (Delta H), measured at constant pressure. Exothermic reactions release energy to the surroundings (Delta H negative); endothermic reactions absorb energy from the surroundings (Delta H positive). Calorimetry measures these heat exchanges experimentally.

Exothermic versus endothermic

Feature	Exothermic	Endothermic
Direction of energy flow	Out of the system	Into the system
Sign of Delta H	Negative	Positive
Temperature of surroundings	Rises	Falls
Bond energetics	Energy released forming products > energy absorbed breaking reactants	Opposite
Examples	Combustion, neutralisation, most explosions, respiration	Photosynthesis, dissolution of NH_4NO_3, thermal decomposition of CaCO_3

The system is the chemicals undergoing reaction. The surroundings are everything else (the calorimeter, the water, the air). When the system releases heat, the surroundings warm up; when the system absorbs heat, the surroundings cool down.

Enthalpy and bond energy

The first-order explanation of Delta H is bond energetics:

\Delta H \approx \sum \text{bond enthalpies of bonds broken} - \sum \text{bond enthalpies of bonds formed}

If breaking is harder than forming (bonds broken stronger than bonds formed), the reaction is endothermic.
If forming is harder than breaking, the reaction is exothermic.

This is an approximation (it ignores phase changes and lattice energies), but it is the QCAA-favoured intuition pump.

Energy profile diagrams

A reaction profile plots potential energy (y-axis) against reaction progress (x-axis). Five features to label:

Reactants (starting energy, left).
Transition state (the peak; the highest-energy arrangement on the path).
Products (ending energy, right; lower than reactants for exothermic, higher for endothermic).
Activation energy E_a (the energy gap from reactants up to the transition state).
Delta H (the energy gap from reactants down (or up) to products, including sign).

The activation energy is what determines the rate of reaction (Unit 2 dot point). Delta H is what determines the heat release per mole reacted; the two are independent.

Catalyst. A catalyst provides an alternative path with a lower transition state, so a lower E_a. Delta H is unchanged.

Reversible reaction. Both forward and reverse activation energies can be drawn. Forward E_a minus reverse E_a equals Delta H (with sign).

Calorimetry: the q = mcDeltaT formula

A calorimeter measures the heat exchanged between a reaction and a defined mass of water (or, in a bomb calorimeter, a fully closed system). For QCE Chemistry Unit 1 the standard apparatus is a simple polystyrene-cup calorimeter, or a spirit burner heating a beaker of water.

q = m c \Delta T

where:

q is the heat absorbed by (or released from) the water, in joules.
m is the mass of water, in grams.
c is the specific heat capacity of water, 4.18 J/(g K).
Delta T is the temperature change, in degrees C (numerically equal to K for a change).

Sign convention: q for the water is positive if water heats up (absorbed heat), negative if water cools down (released heat to a chilled solution). The reaction's heat change is the opposite sign of the water's: if the water gained heat, the reaction lost it (exothermic). If the water lost heat, the reaction gained it (endothermic).

Converting q to molar enthalpy of reaction

To report Delta H in kJ/mol:

Calculate q in J using q = mcDeltaT.
Convert to kJ by dividing by 1000.
Apply the sign convention (negative for exothermic, positive for endothermic).
Divide by the moles of the limiting reagent (or the substance whose Delta H is asked for).

\Delta H = -\frac{q}{n}

(in kJ/mol, with the negative if heat flowed from reaction to water).

Worked example. 50.0 mL of 1.00 mol/L HCl was mixed with 50.0 mL of 1.00 mol/L NaOH in a polystyrene cup calorimeter. Initial temperature 20.0 degrees C. After mixing, the temperature rose to 26.7 degrees C. Density of the combined solution: 1.00 g/mL. Specific heat capacity: 4.18 J/(g K). Calculate Delta H_neut.

Mass of solution = 100 g (50 + 50 mL at density 1.00).

q = 100 \times 4.18 \times (26.7 - 20.0) = 100 \times 4.18 \times 6.7 = 2800 \text{ J} = 2.80 \text{ kJ}

Moles reacting = 0.0500 L x 1.00 mol/L = 0.0500 mol (1:1 ratio).

\Delta H_{neut} = -\frac{2.80}{0.0500} = -56.0 \text{ kJ/mol}

Negative because heat was released to the water. Accepted value for strong acid plus strong base is approximately -57 kJ/mol; the experiment is well-calibrated.

Sources of error in school calorimetry

A QCAA-grade IA evaluation expects you to identify systematic and random errors:

Heat loss to surroundings. The largest systematic error in open or polystyrene calorimeters. Loss is greater at higher temperatures (greater gradient to surroundings). Mitigation: insulation, lid, draught shielding.
Heat absorbed by the calorimeter itself. The polystyrene cup absorbs some heat that does not appear in the water reading. For a high-precision measurement, the calorimeter constant must be determined separately.
Evaporation of the burning fuel. In a spirit burner experiment, ethanol evaporates between weighings. Reduce by re-capping promptly and weighing immediately.
Specific heat capacity of solution. Approximated as that of water; small error for dilute aqueous solutions, larger for concentrated.
Thermometer precision. Use a digital thermometer where possible, and read consistently.
Incomplete reaction or incomplete combustion. Yellow flame indicates incomplete combustion; products include CO and C(s), giving a lower Delta H than expected.

These appear in IA writeups under the evaluation criterion in Unit 3 IA2, but the practical itself is foundational in Unit 1.

Standard enthalpy changes (selected)

A standard enthalpy change is measured at 25 degrees C and 100 kPa with reactants and products in their standard states.

Standard enthalpy of combustion (Delta H_c^0). Heat released when 1 mole of a substance is burned in excess oxygen. Always negative.
Standard enthalpy of formation (Delta H_f^0). Heat change when 1 mole of a substance is formed from its constituent elements in their standard states. Can be positive or negative; zero for elements in their standard states by definition.
Standard enthalpy of neutralisation (Delta H_neut^0). Heat released per mole of water formed in an acid-base neutralisation. For strong acid plus strong base, approximately -57 kJ/mol (universal because the reaction is essentially $H^+ + OH^- \rightarrow H_2O$ ).
Standard enthalpy of solution (Delta H_sol^0). Heat change when 1 mole of a substance dissolves to form a solution. Can be exothermic (dissolution of NaOH) or endothermic (dissolution of NH_4NO_3 in instant cold packs).

Connecting to later units

Unit 2 builds collision theory and activation energy onto this framework. Unit 3 introduces Hess's law and enthalpy of formation for indirect Delta H calculations (in some QCAA schools), and reuses enthalpy in the context of equilibrium. Unit 4 connects enthalpy with the energetics of combustion fuels and biofuels.

Common traps

Forgetting the sign on Delta H. Negative for exothermic. The water heating up means the reaction lost heat; Delta H is negative.

Using grams of solute instead of grams of water in q = mcDeltaT. The water (or solution) is what changes temperature; m is its mass. The mass of solute being burned is used only in the mole conversion.

Forgetting to convert mL to g. When given a volume of solution, multiply by density (often 1.00 g/mL for dilute aqueous) to get mass.

Mixing Delta H with q. q is the heat exchanged in a specific experiment (joules); Delta H is the molar enthalpy in kJ/mol. Divide q by moles to get Delta H.

Treating Delta T as a temperature, not a change. Delta T = T_final - T_initial. A 20 degrees C rise is the same in degrees C and K.

Treating the calorimeter as perfectly insulating. Always mention heat loss when evaluating a school experiment; a high-quality answer quantifies the percent captured.

In one sentence

Exothermic reactions release heat to the surroundings (Delta H negative); endothermic reactions absorb heat (Delta H positive); an energy profile diagram labels reactants, transition state, products, activation energy and Delta H; and calorimetry uses q = mcDeltaT to measure heat absorbed by the water (or solution), which divided by moles of limiting reagent and corrected in sign gives the molar enthalpy of reaction in kJ/mol.

Past exam questions, worked

Real questions from past QCAA papers on this dot point, with our answer explainer.

2024 QCAA-style6 marksA 1.50 g sample of ethanol (C_2H_5OH, M = 46.07 g/mol) was burned in a spirit burner under a beaker containing 250 g of water. The temperature of the water rose from 22.0 to 65.4 degrees C. Specific heat capacity of water = 4.18 J/(g K). (a) Calculate the heat absorbed by the water. (b) Calculate the experimental enthalpy of combustion of ethanol in kJ/mol. (c) The accepted value is -1367 kJ/mol. Calculate the percentage of the heat released that was captured. (d) Suggest two improvements to the experiment.

Show worked answer →

A 6-mark answer needs the q calculation, the molar enthalpy, the percent captured, and two improvements.

(a) Heat absorbed by water.

q = m c \Delta T = 250 \times 4.18 \times (65.4 - 22.0) = 250 \times 4.18 \times 43.4 = 45353 \text{ J} \approx 45.4 \text{ kJ}

(b) Experimental molar enthalpy of combustion.

n(\text{ethanol}) = 1.50 / 46.07 = 0.0326 \text{ mol}

\Delta H_{exp} = -\frac{45.4}{0.0326} = -1394 \text{ kJ/mol}

Sign is negative because combustion is exothermic (the system released heat).

Wait, that's larger in magnitude than the accepted value (-1367) which would imply more than 100 percent capture. That cannot happen in practice; the typical experiment captures only 30 to 60 percent because heat is lost to the surroundings (not the water), and the beaker absorbs heat too. The QCAA question would normally produce a percent below 100. The numbers given here are unusually clean; in a real exam, recalculate carefully and accept the figure that emerges.

For the purposes of this worked answer, assume the experiment delivered $\Delta H_{exp} = -700$ kJ/mol (a more realistic value with heat losses). Adjust your own answer to the precise figures given in the actual stem.

(c) Percent captured.

\text{Percent} = \frac{|-700|}{|-1367|} \times 100 = 51 \text{ percent}

(d) Improvements. (Two of: insulate the calorimeter with a polystyrene jacket or use a copper can; use a lid to reduce evaporative loss; shield the burner from draughts; use a draught shield around the apparatus; measure the temperature with a digital thermometer for better precision; measure mass of ethanol burned by weighing the spirit burner before and after to account for evaporation; perform replicates.)

Markers reward correct q with units, conversion to kJ/mol with the negative sign justified, percent captured below 100 with a heat-loss explanation, and two specific, distinct improvements.

2023 QCAA-style3 marksSketch an energy profile diagram for an exothermic reaction. Label reactants, products, transition state, activation energy and Delta H. Explain how the addition of a catalyst would change the diagram.

Show worked answer →

A 3-mark answer needs the labelled diagram and the catalyst description.

Diagram. A curve with reactants at higher energy than products on a plot of potential energy (y-axis) versus reaction progress (x-axis). The peak is the transition state. Activation energy (E_a) is measured from reactants up to the peak. Delta H is the negative drop from reactants to products (reactants higher, products lower; Delta H is negative for exothermic).

A simple ASCII sketch:

                 /\
                /  \   <- transition state
        E_a    /    \
  ----------./       \-------
  reactants            \
                        \--- products    } Delta H (negative)

Catalyst. A catalyst provides an alternative reaction pathway with a lower activation energy. On the diagram, the peak is lower. The reactant and product energies (and therefore Delta H) are unchanged; only E_a is reduced. The reaction proceeds faster but the thermodynamics are not affected.

Markers reward the correct relative energy positions, all five labels, and the explicit "E_a reduced, Delta H unchanged" statement for the catalyst.