Energetics

What is enthalpy (H)?

A measure of the total energy of a thermodynamic system.

What is enthalpy change (ΔH)?

The heat energy change (of a reaction) measured at constant pressure.

What equation relates ΔH to the enthalpies of reactants and products?

ΔH = H(products) − H(reactants)

What does a negative ΔH mean?

The reaction is exothermic and releases heat to the surroundings.

What does a positive ΔH mean?

The reaction is endothermic and absorbs heat from the surroundings.

In an exothermic reaction, which has greater enthalpy: reactants or products?

Reactants have greater enthalpy than products.

In an endothermic reaction, which has greater enthalpy: reactants or products?

Products have greater enthalpy than reactants.

What happens to the surroundings during an exothermic reaction?

They warm up because energy is transferred from the reaction to the surroundings.

What happens to the surroundings during an endothermic reaction?

They cool down because energy is transferred from the surroundings to the reaction.

What are standard conditions for standard enthalpy changes?

A pressure of 100 kPa and a stated temperature, commonly 298 K (25°C).

What is the standard concentration of a solution?

1 mol dm⁻³.

Why are standard conditions used?

To ensure enthalpy changes can be consistently measured and compared.

What is the standard enthalpy of formation, ΔH°f?

The enthalpy change when 1 mole of a compound is formed from its elements in their standard states under standard conditions.

What is the standard enthalpy of combustion, ΔH°c?

The enthalpy change when 1 mole of a substance is completely burned in oxygen under standard conditions.

What is the standard enthalpy of formation of an element in its standard state?

0 kJ mol⁻¹.

Give an example of a standard enthalpy of formation equation for water.

H₂(g) + ½O₂(g) → H₂O(l)

What is ΔH°f for H₂O(l)?

−286 kJ mol⁻¹.

Give an example of a standard enthalpy of combustion equation for methane.

CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l)

What is ΔH°c for methane?

Approximately −890 kJ mol⁻¹.

Give an example of an endothermic reaction, with CaCO₃(s) as the reactant

CaCO₃(s) → CaO(s) + CO₂(g)

What is the approximate enthalpy change for the decomposition of CaCO₃?

+178 kJ mol⁻¹.

What equation is used to calculate the heat energy transferred in calorimetry?

q = mcΔT

What does q represent?

Heat energy transferred, in J.

What does m represent?

Mass of the substance whose temperature changes, in g.

What does c represent?

Specific heat capacity, in J g⁻¹ K⁻¹.

What does ΔT represent?

Temperature change, in K or °C.

What is the temperature change equation?

ΔT = final temperature − initial temperature

What happens to the numerical value of a temperature change when °C is converted to K?

stays the same

How do you calculate molar enthalpy change from q?

ΔH = −q/n, when q is in kJ and n is the number of moles reacting.

Why is there a negative sign in ΔH = −q/n?

Because q usually represents the heat gained by the surroundings. For an exothermic reaction, the reaction's ΔH must therefore be negative.

What must you do if q is calculated in joules but ΔH is required in kJ mol⁻¹?

Divide q by 1000 to convert J → kJ.

What units are normally used for molar enthalpy change?

kJ mol⁻¹.

What must you identify before calculating ΔH = −q/n?

The number of moles of the limiting reactant.

What apparatus is commonly used for solution calorimetry?

• Polystyrene cup

• Lid

• Thermometer

• Measuring cylinder/pipette

• Balance

• Stirring rod

Why is a polystyrene cup used?

It provides insulation, reducing heat transfer between the reaction mixture and surroundings.

Why is a lid used?

To reduce heat loss to the surroundings and reduce evaporation.

Why should the reaction mixture be stirred?

To ensure the temperature is uniform throughout the solution.

Why should the initial temperature be measured before the reactants are mixed?

To determine the temperature change caused by the reaction.

Why should temperature be recorded at regular time intervals?

To allow a temperature–time graph to be plotted and the temperature change to be determined more accurately.

What is plotted on a temperature–time graph?

Temperature on the y-axis and time on the x-axis.

Why might the maximum/minimum recorded temperature not be the true temperature change?

Heat may be lost to or gained from the surroundings while the experiment is taking place.

How can a temperature–time graph improve the accuracy of ΔT?

Draw best-fit lines before and after the reaction and extrapolate them back to the time the reactants were mixed.

What does the difference between the extrapolated temperatures give?

more accurate value of ΔT.

Why does extrapolation improve the accuracy of ΔT?

It estimates the temperature change at the moment of mixing, before heat has been significantly transferred to or from the surroundings.

How do you calculate the heat energy transferred?

q = mcΔT

How do you calculate molar enthalpy change from q?

ΔH = −q/n

What must you do to q before using ΔH = −q/n if q is in joules?

Convert J → kJ by dividing by 1000.

Why is there a negative sign in ΔH = −q/n?

Because q represents heat transferred to the surroundings. Therefore, heat released by the reaction gives a negative ΔH.

What units should the final molar enthalpy change have?

kJ mol⁻¹.

Which number of moles should be used when calculating ΔH?

The number of moles of the limiting reactant.

What are common sources of error in calorimetry experiments?

Heat loss to the surroundings
Incomplete combustion
Evaporation
Heat absorbed by the apparatus

Why does heat loss make an experimental exothermic ΔH less negative?

Less heat is measured as being transferred to the water, so the calculated magnitude of ΔH is too small.

How can heat loss be reduced?

Use insulation
Use a lid
Use a polystyrene cup
Minimise exposure to the surroundings

How can incomplete combustion be reduced in combustion calorimetry?

Ensure there is an adequate supply of oxygen and that the fuel burns completely.

How can evaporation of fuel be reduced?

Keep the fuel covered/contained when not being used and minimise the time it is exposed to the air.

What is AQA Required Practical 2?

Measurement of an enthalpy change.

What types of reactions can be investigated using calorimetry in AQA?

Dissolution of KCl
Dissolution of Na₂CO₃
Neutralisation of NaOH + HCl
Displacement: CuSO₄ + Zn
Combustion of alcohols

What practical skills should you be able to apply to calorimetry questions?

Selecting appropriate apparatus
Measuring temperature accurately
Recording measurements
Plotting temperature–time graphs
Determining ΔT
Calculating q using q = mcΔT
Calculating molar ΔH
Identifying sources of error
Suggesting improvements

What is the general sequence for a calorimetry calculation?

1. Find ΔT
↓
2. Calculate q = mcΔT
↓
3. Convert q from J → kJ
↓
4. Find moles of the limiting reactant
↓
5. Calculate ΔH = −q/n
↓
6. Add the correct units: kJ mol⁻¹
↓
7. Check the sign makes sense

What sign should ΔH have for an exothermic calorimetry experiment?

Negative.

What sign should ΔH have for an endothermic calorimetry experiment?

Positive.

What is the most important thing to remember about calorimetry practicals?

The measured temperature change is used to calculate q, which is then converted into a molar enthalpy change.

What is Hess's law?

The total enthalpy change for a reaction is independent of the route taken, provided the initial and final conditions are the same.

Why does Hess's law work?

Enthalpy is a state function, so the overall enthalpy change depends only on the initial and final states, not the pathway.

When is Hess's law particularly useful?

When an enthalpy change cannot be measured directly or is difficult/impractical to measure.

What happens to ΔH when a chemical equation is reversed?

The sign of ΔH changes.

What happens to ΔH when all coefficients in an equation are multiplied by a factor?

ΔH is multiplied by the same factor.

What is the key rule when constructing a Hess cycle?

The enthalpy change along one route must equal the enthalpy change along the other route.

What should you do if an equation in a Hess cycle is reversed?

Reverse the sign of its ΔH.

What should you do if an equation is multiplied by 2?

Multiply its ΔH by 2.

What should you do if an equation is divided by 2?

Divide its ΔH by 2.

What equation calculates ΔH°r using enthalpies of formation?

ΔH°r = ΣΔH°f(products) − ΣΔH°f(reactants)

What does Σ mean in the formation enthalpy equation?

Sum of all the relevant enthalpy changes, including stoichiometric coefficients.

What value is used for elements in their standard states when calculating using formation enthalpies?

0 kJ mol⁻¹.

What is the most common mistake when using the formation enthalpy equation?

Forgetting to multiply each ΔH°f by its stoichiometric coefficient.

What equation calculates ΔH°r using enthalpies of combustion?

ΔH°r = ΣΔH°c(reactants) − ΣΔH°c(products)

Why are the reactants and products reversed in the combustion equation compared with the formation equation?

Combustion enthalpies represent substances being combusted to the same final products, so the cycle requires the difference between combustion of the reactants and combustion of the products.

What should you always check before doing a Hess calculation?

Balanced equations
Correct states
Correct coefficients
Correct direction of each reaction
Correct sign of each ΔH

Use Hess's law to calculate the enthalpy change for C(s) + O₂(g) → CO₂(g), given: C(s) + ½O₂(g) → CO(g) ΔH = −110.5 kJ mol⁻¹ CO(g) + ½O₂(g) → CO₂(g) ΔH = −283.0 kJ mol⁻¹

ΔH = −110.5 + (−283.0) = −393.5 kJ mol⁻¹

What equation would you use to calculate the enthalpy change for a reaction from formation enthalpy data?

ΔH°r = ΣΔH°f(products) − ΣΔH°f(reactants)
For:
SO₂(g) + 2H₂S(g) → 3S(s) + 2H₂O(l)

Given: ΔH°f[SO₂(g)] = −296.8 kJ mol⁻¹ ΔH°f[H₂S(g)] = −20.6 kJ mol⁻¹ ΔH°f[H₂O(l)] = −i285.8 kJ mol⁻¹ ΔH°f[S(s)] = 0 Calculate ΔH°r.

Products: (3 × 0) + (2 × −285.8) = −571.6 Reactants: −296.8 + (2 × −20.6) = −338.0 ΔH°r = −571.6 − (−338.0) = −233.6 kJ mol⁻¹

What types of Hess calculations are specifically required by AQA?

Calculating enthalpy changes from enthalpies of formation
Calculating enthalpy changes from enthalpies of combustion
Constructing/using Hess cycles
AQA gives examples including thermal decomposition of NaHCO₃ and hydration of MgSO₄ and CuSO₄.

What is bond enthalpy?

The enthalpy change required to break one mole of a particular covalent bond in gaseous molecules.

What is mean bond enthalpy?

The average enthalpy change required to break one mole of a particular type of covalent bond in gaseous molecules, averaged over a range of compounds.

Why are bond enthalpies called "mean" values?

The same type of bond has slightly different bond enthalpies in different molecules.

Why can the same bond have different enthalpies in different molecules?

The strength of a bond is affected by its molecular environment and surrounding atoms.

Is bond breaking endothermic or exothermic?

Endothermic — energy is required to overcome the attraction between bonded atoms.

Is bond formation endothermic or exothermic?

Exothermic — energy is released when new attractive forces form between atoms.

What happens to energy when bonds are broken?

Energy is absorbed.

What happens to energy when bonds are formed?

Energy is released.

What equation is used to calculate ΔH from mean bond enthalpies?

ΔH = Σ(bonds broken) − Σ(bonds formed)

Why are bond enthalpies for bonds broken added together?

Energy must be supplied to break all the bonds.

Why are bond enthalpies for bonds formed subtracted?

Energy is released when the new bonds form.

When can mean bond enthalpies be used to calculate ΔH according to AQA?

For reactions in the gaseous phase.

Worked example — ammonia- For: N₂(g) + 3H₂(g) → 2NH₃(g)

1. Which bonds must be broken?

2. Which bonds are formed?

3. Using N≡N = 941 kJ mol⁻¹ and H–H = 436 kJ mol⁻¹, how much energy is required to break the reactant bonds?

4. Using N–H = 391 kJ mol⁻¹, how much energy is released when the product bonds form?

5. Calculate ΔH for the formation of ammonia.

1. 1 × N≡N 3 × H–H

2. 6 × N–H

3. (1 × 941) + (3 × 436) = 2249 kJ mol⁻¹

4. 6 × 391 = 2346 kJ mol⁻¹

5. ΔH = 2249 − 2346 = −97 kJ mol⁻¹ Therefore the reaction is exothermic.