Every chemical reaction involves a transfer of heat energy between the reacting system and its surroundings. Chemical energetics is the study of these energy changes. By measuring and calculating the enthalpy change, we can determine the energy requirements and stability of chemical substances.
🔑 Key Principle
Energy cannot be created or destroyed. In an exothermic reaction, chemical energy stored in bonds is converted into heat energy and released, raising the temperature of the surroundings. In an endothermic reaction, heat energy is absorbed from the surroundings, lowering their temperature, and stored as chemical energy.
What is Enthalpy Change?
During a chemical reaction at constant pressure, the heat energy change is called the enthalpy change, represented by \( \Delta H \) (read as delta H). It is measured in kilojoules per mole (\( \text{kJ mol}^{-1} \)).
The heat energy change measured under conditions of constant pressure.
Standard Conditions
Because enthalpy changes vary with temperature, pressure, and concentration, chemists measure them under a set of agreed standard conditions. The standard state of a substance is its most stable physical state under these standard conditions:
- A standard pressure of \( 100\text{ kPa} \) (approximately 1 atmosphere).
- A specified temperature, which is almost always \( 298\text{ K} \) (\( 25\text{ }^\circ\text{C} \)).
- Solutions must have a concentration of exactly \( 1.00\text{ mol dm}^{-3} \).
A standard enthalpy change is denoted by a superscript theta symbol: \( \Delta H^\theta \).
Exothermic and Endothermic Reactions
We classify chemical reactions based on the direction of heat transfer:
Exothermic Reactions
Heat energy is released to the surroundings. The reactants have more enthalpy than the products. The value of \( \Delta H \) is negative (\( \Delta H < 0 \)).
Examples: Combustion of fuels, neutralisation of acids and bases.
Endothermic Reactions
Heat energy is absorbed from the surroundings. The products have more enthalpy than the reactants. The value of \( \Delta H \) is positive (\( \Delta H > 0 \)).
Examples: Thermal decomposition of carbonates, photosynthesis.
Enthalpy Profile Diagrams
Enthalpy profiles show the relative enthalpy of reactants and products during a reaction. They also display the activation energy (\( E_a \)), which is the minimum energy required to start a reaction by breaking the initial bonds.
Key Standard Enthalpy Definitions
AQA requires you to learn the following three precise definitions for standard enthalpy changes. You must write them exactly as shown below in exams:
The enthalpy change when one mole of a compound is formed from its constituent elements in their standard states under standard conditions.
The enthalpy change when one mole of a substance is completely burned in oxygen under standard conditions, with all reactants and products in their standard states.
The enthalpy change when an acid and base react to form one mole of water under standard conditions, with all reactants and products in their standard states.
Standard enthalpies of elements in their standard states are defined as zero. For example, \( \Delta_f H^\theta \) of \( \text{O}_2\text{(g)} \) or \( \text{C(s, graphite)} \) is exactly \( 0\text{ kJ mol}^{-1} \). You do not need to form elements from themselves!
Writing Standard Enthalpy Equations
You must practice writing chemical equations representing these standard changes. Let's look at how to do this correctly:
Solution:
1. Start with the product: Under standard conditions, we must produce exactly one mole of the compound. Write \( \text{C}_2\text{H}_5\text{OH(l)} \) on the right-hand side.
\[ \text{...} \rightarrow \text{C}_2\text{H}_5\text{OH(l)} \]
2. Identify constituent elements: The elements that make up ethanol are Carbon, Hydrogen, and Oxygen.
3. Deduce standard states of elements: At \( 298\text{ K} \) and \( 100\text{ kPa} \):
- Carbon is solid graphite, \( \text{C(s)} \)
- Hydrogen is gaseous diatomic molecules, \( \text{H}_2\text{(g)} \)
- Oxygen is gaseous diatomic molecules, \( \text{O}_2\text{(g)} \)
4. Balance the elements to produce exactly one mole of ethanol:
- We need 2 carbon atoms: \( 2\text{C(s)} \)
- We need 6 hydrogen atoms: \( 3\text{H}_2\text{(g)} \)
- We need 1 oxygen atom: This requires a fractional coefficient of \( \frac{1}{2}\text{O}_2\text{(g)} \)
5. Write the final equation:
\[ 2\text{C(s)} + 3\text{H}_2\text{(g)} + \frac{1}{2}\text{O}_2\text{(g)} \rightarrow \text{C}_2\text{H}_5\text{OH(l)} \]
Solution:
1. Start with the reactant: Under standard conditions, we must burn exactly one mole of the fuel. Write \( 1\text{ C}_2\text{H}_6\text{(g)} \) on the left-hand side.
\[ \text{C}_2\text{H}_6\text{(g)} + \text{O}_2\text{(g)} \rightarrow \text{...} \]
2. Identify products of complete combustion: Complete combustion of a hydrocarbon produces carbon dioxide gas, \( \text{CO}_2\text{(g)} \), and liquid water, \( \text{H}_2\text{O(l)} \).
3. Balance the products based on one mole of ethane:
- 2 carbons in ethane produce \( 2\text{CO}_2\text{(g)} \)
- 6 hydrogens in ethane produce \( 3\text{H}_2\text{O(l)} \)
4. Balance oxygen on the reactant side: The products contain \( (2 \times 2) + (3 \times 1) = 7 \) oxygen atoms. This means we require \( \frac{7}{2} \) molecules of oxygen gas on the left.
5. Write the final equation:
\[ \text{C}_2\text{H}_6\text{(g)} + 3\frac{1}{2}\text{O}_2\text{(g)} \rightarrow 2\text{CO}_2\text{(g)} + 3\text{H}_2\text{O(l)} \]
State symbols are mandatory in enthalpy equations. Enthalpy changes are physically different depending on states. For example, water as a liquid has a different enthalpy value than water as a gas because energy is released when steam condenses to liquid water. Always include state symbols: (s), (l), (g), and (aq).
Get flashcards and quizzes in ChemEasy, or plan your revision with ChemPlan IB.