🔥 Enthalpy Change, ΔH
1. Understanding Enthalpy Changes
- Enthalpy (H): A measure of the total energy of a system (chemical potential energy + pressure-volume work).
- Enthalpy Change (ΔH): The heat energy change at constant pressure when a chemical reaction occurs.
- Exothermic Reaction: ΔH is negative (releases heat; products have less enthalpy than reactants).
- E.g. Combustion, neutralisation.
- Endothermic Reaction: ΔH is positive (absorbs heat; products have more enthalpy than reactants).
- E.g. Thermal decomposition, photosynthesis.
- Exothermic Reaction: ΔH is negative (releases heat; products have less enthalpy than reactants).
2. Reaction Pathway Diagrams
- Graphical representation showing energy changes during a reaction.
- Y-axis: Enthalpy
- X-axis: Reaction progress
- Shows:
- Reactants’ energy
- Products’ energy
- ΔH (difference between products and reactants)
- Activation energy (Ea): Minimum energy needed to start a reaction (from reactants to peak of energy barrier).
3. Key Definitions
(a) Standard Conditions (⦵):
- Temperature: 298 K (25°C)
- Pressure: 101 kPa (1 atm)
- Concentration (for solutions): 1 mol/dm³
(b) Standard Enthalpy Changes:
| Term | Symbol | Definition |
|---|---|---|
| ΔH⦵r | Standard enthalpy change of reaction | Enthalpy change when the reaction occurs in the molar quantities shown in the equation under standard conditions. |
| ΔH⦵f | Standard enthalpy change of formation | Enthalpy change when 1 mole of a compound is formed from its elements in their standard states. |
| ΔH⦵c | Standard enthalpy change of combustion | Enthalpy change when 1 mole of a substance is completely burned in oxygen under standard conditions. |
| ΔH⦵neut | Standard enthalpy change of neutralisation | Enthalpy change when 1 mole of water is formed by neutralisation of acid and alkali under standard conditions. |
4. Bond Breaking and Making
- Energy is required to break bonds (endothermic).
- Energy is released when new bonds are formed (exothermic).
- Overall enthalpy change: ΔH=∑(Bond energies of bonds broken)−∑(Bond energies of bonds formed)
5. Bond Energies
- Bond Energy (or Bond Enthalpy): Energy required to break 1 mole of a specific bond in a gaseous molecule.
- Always positive (since bond breaking requires energy).
- Enthalpy change can be calculated using bond energies from data tables.
6. Mean Bond Energies
- Some bond energies are averaged from similar compounds.
- E.g. C–H bond energy may vary slightly in methane vs ethane, so an average value is used.
7. Experimental Determination of ΔH
a) Using calorimetry and the formula:
q = mcΔT
- q = heat energy (J)
- m = mass of water (or solution) in g
- c = specific heat capacity (usually 4.18 J/g°C for water)
- ΔT = temperature change (°C or K)
b) Calculating enthalpy change per mole:
ΔH = q /n
- n = number of moles of limiting reagent
- ΔH = kJ/mol (usually divide by 1000 to convert from J)
🔁 Hess’s Law
1. Hess’s Law:
The total enthalpy change of a reaction is the same, no matter what route is taken.
2. Energy Cycles and Calculations
- Useful for finding enthalpy changes that cannot be measured directly.
- Apply Hess’s Law to create energy cycles involving:
- ΔH⦵f values (formation)
- ΔH⦵c values (combustion)
Example Cycle Using ΔH⦵f:
To find ΔH⦵r: ΔHr∘=∑ΔHf (products)∘−∑ΔHf (reactants)
Example Cycle Using ΔH⦵c:
To find ΔH⦵f: ΔHf∘=∑ΔHc (reactants)∘−∑ΔHc (products)
3. Using Bond Energy Data in Hess’s Law
- Alternate approach to Hess’s Law if only bond enthalpy data is available.
- Equation:
ΔH=∑(Bond energies of bonds broken)−∑(Bond energies of bonds formed)
- Apply to reactions involving only gaseous substances for better accuracy.
🧠 Summary Table
| Concept | Formula / Note |
|---|---|
| Heat change (q) | q=mcΔT |
| Enthalpy per mole | ΔH=−q/n |
| Bond energy calc | ΔH=∑Ebonds broken−∑Ebonds formed |
| Hess’s law (formation) | ΔHr=∑ΔHf(products)−∑ΔHf(reactants) |
| Hess’s law (combustion) | ΔHr=∑ΔHc(reactants)−∑ΔHc(products) |
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