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

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

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:

• 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:

1. Reactants (starting energy, left).
2. Transition state (the peak; the highest-energy arrangement on the path).
3. Products (ending energy, right; lower than reactants for exothermic, higher for endothermic).
4. Activation energy E_a (the energy gap from reactants up to the transition state).
5. $\Delta H$ (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.

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:

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

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).

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

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.

Examples in context

Example 1. Ethanol combustion bomb-calorimetry at a Logan biofuels pilot. A Logan-based pilot plant tests cane-molasses ethanol. A bomb calorimeter burns $0.500 \, \text{g}$ of ethanol ($M = 46.07 \, \text{g mol}^{-1}$); the water bath temperature ($1.20 \, \text{kg}$ water) rises by $2.74^{\circ}\text{C}$. Using $q = mc \Delta T$ with $c = 4.18 \, \text{J g}^{-1} \, ^{\circ}\text{C}^{-1}$ gives $q = 1200 \times 4.18 \times 2.74 = 13{,}740 \, \text{J}$. Moles of ethanol $= 0.500 / 46.07 = 0.01086 \, \text{mol}$, so $\Delta H_c = -13{,}740 / 0.01086 = -1{,}266 \, \text{kJ mol}^{-1}$ (the literature value is $-1{,}367 \, \text{kJ mol}^{-1}$; the $7\%$ shortfall is heat loss, a common IA2 discussion point).

Example 2. Aluminium production energy at Gladstone's Boyne smelter. Boyne uses the Hall-Heroult process to reduce $\text{Al}_2 \text{O}_3$ to molten aluminium at $950^{\circ}\text{C}$. Hess's law lets students combine the formation enthalpies $\Delta H_f(\text{Al}_2 \text{O}_3) = -1{,}676 \, \text{kJ mol}^{-1}$ and $\Delta H_f(\text{CO}_2) = -394 \, \text{kJ mol}^{-1}$ to obtain the overall electrolysis enthalpy. The reaction $2 \text{Al}_2 \text{O}_3 + 3 \text{C} \rightarrow 4 \text{Al} + 3 \text{CO}_2$ has $\Delta H = +2{,}170 \, \text{kJ}$ per mole reaction, explaining why Boyne consumes about $14 \, \text{MWh}$ per tonne of aluminium and is built next to dedicated Callide-derived power.

Try this

Q1. Define $\Delta H_c$ and state the sign convention for exothermic reactions. [2 marks]

• Cue. Enthalpy change per mole of fuel completely combusted at standard conditions. Exothermic $\Delta H < 0$.

Q2. $25.0 \, \text{mL}$ of $1.00 \, \text{mol L}^{-1}$ $\text{HCl}$ is mixed with $25.0 \, \text{mL}$ of $1.00 \, \text{mol L}^{-1}$ $\text{NaOH}$; temperature rises by $6.50^{\circ}\text{C}$. Calculate $\Delta H$ per mole of water formed. Assume $c = 4.18 \, \text{J g}^{-1} \, ^{\circ}\text{C}^{-1}$ and density $= 1.00 \, \text{g mL}^{-1}$. [4 marks]

• Cue. $q = 50.0 \times 4.18 \times 6.50 = 1{,}359 \, \text{J}$; $n = 0.0250 \, \text{mol}$; $\Delta H = -54.4 \, \text{kJ mol}^{-1}$.

Q3. A student measures $\Delta H_c$ of methanol. Their value is $25\%$ less negative than the data-book figure. (a) State two heat-loss sources. (b) Suggest a calorimeter improvement. (c) Identify one assumption challenged by IA process limitations. [2+2+2 marks]

• Cue. (a) Heat to surroundings; incomplete combustion. (b) Insulated bomb. (c) Constant specific heat. Communication and judgement criteria.