enthalpy change

Question 1

Sure! Here’s a quick revision guide for Enthalpy Change in A-level OCR B Chemistry:

Enthalpy Change (ΔH)
Enthalpy (H): The total energy of a system, including internal energy and the energy required to displace its surroundings to make room for the system.

Enthalpy Change (ΔH): The heat energy change that occurs during a chemical reaction at constant pressure.

If the reaction releases heat to the surroundings, ΔH is negative (exothermic).
If the reaction absorbs heat from the surroundings, ΔH is positive (endothermic).
Types of Enthalpy Changes:
Enthalpy of Combustion (ΔH°_c):

The enthalpy change when 1 mole of a substance is completely burned in oxygen under standard conditions.
Example:
CH
4
(
𝑔
)
+
2
O
2
(
𝑔
)
→
CO
2
(
𝑔
)
+
2
H
2
𝑂
(
𝑔
)
CH
4
​
(g)+2O
2
​
(g)→CO
2
​
(g)+2H
2
​
O(g)
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.
Example:
C(s)
+
O
2
(
𝑔
)
→
CO
2
(
𝑔
)
C(s)+O
2
​
(g)→CO
2
​
(g)
Enthalpy of Neutralisation (ΔH°_neut):

The enthalpy change when 1 mole of water is formed from an acid and an alkali (base) in an aqueous solution.
Example:
H
+
(
𝑎
𝑞
)
+
OH
−
(
𝑎
𝑞
)
→
H
2
𝑂
(
𝑙
)
H
+
(aq)+OH
−
(aq)→H
2
​
O(l)
Typically around -57 kJ/mol for strong acid + strong base reactions.
Ionisation Enthalpy:

The enthalpy change when one mole of electrons is removed from one mole of gaseous atoms to form positive ions.
Example:
Na
(
𝑔
)
→
Na
+
(
𝑔
)
+
𝑒
−
Na(g)→Na
+
(g)+e
−

Calculating Enthalpy Change:
Using Bond Enthalpies (Bond Dissociation Enthalpy):

The enthalpy change when one mole of a particular bond is broken in a molecule in the gas phase.

ΔH = Σ(reactant bond enthalpies) - Σ(product bond enthalpies)

Note: The sign depends on whether bonds are being broken (energy required, endothermic) or formed (energy released, exothermic).

Hess’s Law:

States that the total enthalpy change of a reaction is the sum of the enthalpy changes of the steps into which the reaction can be divided.
ΔH = ΔH1 + ΔH2 + …
Use Hess’s law to find ΔH if direct measurement is not possible. You can add or subtract reactions to get the desired overall reaction.
Standard Conditions:
Pressure: 100 kPa
Temperature: 298 K (25°C)
Concentration: 1 mol/dm³ (for solutions)
Exothermic vs. Endothermic Reactions:
Exothermic (ΔH < 0):
Energy is released, typically as heat.
Common in combustion, neutralisation reactions.
Endothermic (ΔH > 0):
Energy is absorbed from the surroundings.
Common in thermal decomposition reactions.
Important Concepts:
Standard Enthalpy of Reaction (ΔH°_reaction): The enthalpy change for a reaction when all reactants and products are in their standard states.
Heat Capacity: The amount of heat needed to change the temperature of a substance by 1°C.
𝑞
=
𝑚
𝑐
Δ
𝑇
q=mcΔT
q = heat absorbed/released (J)
m = mass (g)
c = specific heat capacity (J/g°C)
ΔT = change in temperature (°C)
Practical Example:
If you’re asked to calculate the enthalpy change for the combustion of methane using bond enthalpy values:
Write the balanced combustion equation:
CH
4
(
𝑔
)
+
2
O
2
(
𝑔
)
→
CO
2
(
𝑔
)
+
2
H
2
𝑂
(
𝑔
)
CH
4
​
(g)+2O
2
​
(g)→CO
2
​
(g)+2H
2
​
O(g)
List the bond enthalpies (for C-H, O=O, C=O, etc.).
Calculate energy required to break bonds (reactants) and energy released when bonds are formed (products).
Apply Hess’s Law if needed.
Key Formulae:
ΔH = q / n

𝑞
q = heat energy (J)
𝑛
n = number of moles of reactant (mol)
q = mcΔT

Where
𝑞
q = heat absorbed,
𝑚
m = mass,
𝑐
c = specific heat capacity, and
Δ
𝑇
ΔT = change in temperature.