1. What happens to ΔH when a reaction is reversed?
- Its sign reverses
- It becomes zero
- It doubles always
Check the answer and reasoning
Its sign reverses
The state change is reversed, reversing the energy transfer.
CHEM / AS level
Enthalpy and Hess cycles
Enthalpy changes compare states under specified conditions. Formation, combustion, atomisation and bond enthalpies use different definitions; their signs cannot be assigned from names alone.
Learning objectives
- Explain enthalpy and hess cycles using a concrete example.
- Apply the worked method and justify each step.
- Recognise the limitation described in the final check.
The big idea
Enthalpy changes compare states under specified conditions. Formation, combustion, atomisation and bond enthalpies use different definitions; their signs cannot be assigned from names alone.
Hess’s law uses enthalpy as a state function to combine reactions. Reverse an equation and reverse its enthalpy sign; multiply an equation and scale its enthalpy change.
Equations and relationships
ΔHreaction ≈ ΣE(bonds broken) − ΣE(bonds formed)
Mean bond enthalpies give an estimate, usually for gaseous species.
Worked example
Breaking bonds costs 500 kJ/mol; formation releases 650 kJ/mol. Estimate ΔH.
Follow the solution
500−650=−150 kJ/mol, so the model reaction is exothermic.
Investigate and explain
Predict the worked result before calculating. Change one input, record the new result, and explain which assumptions still hold.
Use the evidence activity below; no exact PhET model is claimed.
Check your understanding
2. Breaking bonds costs 500 kJ/mol; formation releases 650 kJ/mol. Estimate ΔH.
- +150 kJ/mol
- −150 kJ/mol
- −1150 kJ/mol
Check the answer and reasoning
−150 kJ/mol
500−650=−150 kJ/mol, so the model reaction is exothermic.
3. Are mean bond enthalpies exact for every compound?
- Yes
- Only for liquids
- No
Check the answer and reasoning
No
They average gaseous bond environments, so they estimate rather than exactly reproduce many reaction enthalpies.