1.1 Measuring enthalpy changes

A chemical reaction transfers energy between the system and surroundings. Total energy is conserved. Heat is energy transferred because of a temperature difference; temperature describes the thermal state.

State which part is the system, which part is the surroundings, and the direction of energy transfer before interpreting a temperature change.

For an exothermic hand-warmer reaction, define the reacting chemicals as the system: energy leaves that system and enters your hand and the air as surroundings. A temperature rise is evidence about the surroundings, while heat names the energy crossing the boundary; neither quantity is 'stored temperature'.

Representative question

[Maximum number: 1]

Consider a reaction mixture that is in thermal equilibrium with the surroundings. When a reaction takes place, the temperature of the mixture decreases.

Which row correctly shows the changes in energy when the new thermal equilibrium is established?

Energy of system

Energy of surroundings

decreases

increases

decreases

increases

decreases

Type	Energy direction	Surroundings temperature
Endothermic	Absorbed by the system	Decreases
Exothermic	Released by the system	Increases

Classify from the direction of energy transfer, then check that the observed temperature change of the surroundings agrees.

Connect three representations: exothermic means energy flows out of the system, the surroundings warm, and ΔH for the system is negative; endothermic gives the opposite pattern and positive ΔH. State the observed part before inferring the reaction type.

Representative question

[Maximum number: 1]

Which statement about an exothermic reaction is correct?

Temperature decreases and the products have higher enthalpy than the reactants.

Temperature decreases and the products have lower enthalpy than the reactants.

Temperature increases and the products have higher enthalpy than the reactants.

Temperature increases and the products have lower enthalpy than the reactants.

Lower relative energy corresponds to greater relative stability. In an exothermic reaction the products are lower in energy than the reactants; in an endothermic reaction the products are higher.

Read reactant and product levels, identify the sign and direction of ΔH, and label the energy profile clearly.

On a profile, ΔH is the vertical difference between product and reactant levels, while activation energy rises from reactants to the peak. A catalyst lowers the peak by changing the pathway but leaves the two energy levels, ΔH and the relative stability of reactants and products unchanged.

An energy profile identifies an activation barrier but does not by itself determine an observed rate: temperature, particle concentrations/collision frequency and the available pathway also matter. Use the diagram to compare energetic barriers only when the profiles and conditions make that comparison valid.

Representative question

[Maximum number: 3]

The forward reaction is endothermic, uses iron(III) oxide as a catalyst, and takes place at 900 K .

Sketch the energy profile for the reaction, both with and without the catalyst, labelling $\Delta H$ and the activation energies.

$Q = mcΔT and ΔH = −Q/n$

At constant pressure, calculate heat transferred from mass, specific heat capacity, and temperature change, then divide by reacting moles and apply the sign convention. Check standard conditions and units.

If 100.0 g of solution warms by 6.0 K and c = 4.18 J g⁻¹ K⁻¹, q(solution) = +2.51 kJ, so q(reaction) = −2.51 kJ. Divide by moles of the limiting reacting amount for kJ mol⁻¹, and identify heat loss or an ignored calorimeter heat capacity as reasons an experimental magnitude may be too small.

A standard molar enthalpy change belongs to the balanced reaction as written, with substances in their stated standard states under standard conditions. A classroom calorimetry value is an experimental estimate: report its conditions and uncertainty separately from a data-book or theoretical standard value. Heat loss to the surroundings or ignored calorimeter heat capacity commonly makes the measured magnitude too small.

Representative question

[Maximum number: 1]

What is the standard enthalpy change, $\Delta H_{\text {combustion }}^{\ominus}$ , according to the data?

Amount of fuel burned $=0.110 \mathrm{~mol}$
Mass of water $=200 \mathrm{~g}$
Initial temperature of water $=21.0^{\circ} \mathrm{C}$
Final temperature of water $=25.0^{\circ} \mathrm{C}$
Specific heat capacity of water, $c_{\mathrm{w}}=4.18 \mathrm{~J} \mathrm{~g}^{-1} \mathrm{~K}^{-1}$
$Q=m c \Delta T$

$-2110 \mathrm{~kJ} \mathrm{~mol}^{-1}$

$-30.4 \mathrm{~kJ} \mathrm{~mol}^{-1}$

$+30.4 \mathrm{~kJ} \mathrm{~mol}^{-1}$

$+2110 \mathrm{~kJ} \mathrm{~mol}^{-1}$

Retrieve the route: define system/surroundings transfer, classify endothermic or exothermic, read stability from energy profiles, then calculate Q and ΔH with the correct sign.

Check the energy direction, surroundings temperature, reactant/product levels, moles, units, and ΔH sign.

ConceptIB Chemistry SL

1.1 Measuring enthalpy changes

Energy Transfer in Reactions

Describing Energy Transfer

Assessment in practice

Endothermic and Exothermic

Classifying Energy Changes

Assessment in practice

Stability and Energy Profiles

Interpreting Energy Profiles

Assessment in practice

Standard Enthalpy from Calorimetry

Calculating Enthalpy Change

Assessment in practice

Measuring Enthalpy Summary

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