Energetics & Enthalpy – Lecture Review

A set of question-and-answer flashcards covering key definitions, calculations, experimental methods, Hess’s Law applications, bond enthalpies, and common sources of error in energetics.

In an exothermic reaction, where does energy flow and what sign does ΔH have?

Energy flows from the system (chemicals) to the surroundings; ΔH is negative.

In an endothermic reaction, where does energy flow and what sign does ΔH have?

Energy is absorbed from the surroundings into the system; ΔH is positive.

Define ‘system’ and ‘surroundings’ in thermochemistry.

The system is the reacting chemicals; the surroundings are everything outside the chemicals.

State the standard conditions used when quoting standard enthalpy changes.

100 kPa pressure, 298 K (25 °C), 1 mol dm⁻³ solutions, and substances in their standard states at 298 K.

What is the definition of the standard enthalpy change of formation (ΔfH°)?

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

What is the value of ΔfH° for any element in its standard state?

0 kJ mol⁻¹.

Give the definition of the standard enthalpy of combustion (ΔcH°).

The enthalpy change when 1 mole of a substance is completely burned in oxygen under standard conditions with all reactants and products in their standard states.

Why is incomplete combustion less exothermic than complete combustion?

Because energy remains stored in products such as CO or C (soot) instead of being fully released as CO₂ and H₂O.

Write the calorimetry equation used to find heat exchanged in a solution.

q (J) = m (g) × cₚ (J g⁻¹ K⁻¹) × ΔT (K).

What specific heat capacity value is normally assumed for aqueous solutions?

4.18 J g⁻¹ K⁻¹ (the value for water).

List two key experimental steps to minimise heat loss in a simple solution calorimetry experiment.

1) Use a polystyrene cup placed in a beaker for insulation; 2) record temperature at regular intervals and extrapolate back to the mixing time.

Name three common sources of error in solution calorimetry.

Heat loss to surroundings, neglecting heat absorbed by calorimeter apparatus, assuming solution has same cₚ and density as water.

Describe the four-step method for calculating ΔH from calorimetry data.

1) Calculate q with m c ΔT; 2) determine moles of limiting reactant; 3) divide q by those moles; 4) convert to kJ mol⁻¹ and add correct sign.

How is the sign of ΔH determined from the temperature change in calorimetry?

Temperature rise → exothermic (negative ΔH); temperature drop → endothermic (positive ΔH).

State Hess’s Law.

The total enthalpy change for a reaction is independent of the route taken

Give the formula for calculating ΔH using standard enthalpies of formation.

ΔH° = Σ ΔfH°(products) − Σ ΔfH°(reactants).

Give the formula for calculating ΔH using standard enthalpies of combustion.

ΔH° = Σ ΔcH°(reactants) − Σ ΔcH°(products).

What is ‘activation energy’ (Eₐ) on an energy profile diagram?

The minimum energy required to start breaking bonds and initiate a reaction, represented by the peak between reactants and products.

Define ‘mean bond enthalpy’.

The average enthalpy required to break a particular type of covalent bond into gaseous atoms, averaged over many compounds.

State the bond-enthalpy method equation for estimating ΔH when all species are gaseous.

ΔH ≈ Σ(bond energies broken) − Σ(bond energies made).

Why are ΔH values from mean bond enthalpies less accurate than those from formation data?

Because mean bond enthalpies are averaged and do not account for the exact molecular environment of each bond.

When moving up a homologous series (e.g., alcohols), how does ΔcH° change and why?

It increases by a roughly constant amount because each extra –CH₂– adds the same set of bonds broken and formed during combustion.

Give two reasons experimental combustion enthalpies measured with a simple burner are lower than literature values.

Heat lost to surroundings and incomplete combustion of fuel.

Why can’t the enthalpy change for hydrating an anhydrous salt be measured directly?

It is difficult to add exactly one mole of water and to measure temperature change in a solid accurately; instead Hess’s Law cycles are used.

In bond-enthalpy calculations, why must all reactants and products be considered in the gaseous state?

Because bond enthalpies are defined for gaseous molecules; phase changes would introduce extra enthalpy terms not accounted for.

What additional enthalpy term must be included when using bond energies for a liquid compound?

The enthalpy of vaporisation/atomisation needed to convert the liquid (or solid) into gaseous atoms.

Explain why the density of aqueous solutions is often assumed to be 1 g cm⁻³ in calorimetry calculations.

It simplifies mass determination (volume in cm³ = mass in g) and introduces minimal error for dilute solutions.

During calorimetry with a solid reactant, how should its mass be determined accurately?

Use a ‘before-and-after’ weighing method: weigh the container with solid before addition and again afterward to find the exact mass transferred.

What graphical technique helps correct for heat loss during a slow reaction in calorimetry?

Taking temperature readings at regular intervals and extrapolating the cooling curve back to the time of mixing.