Nota Chemistry Tingkatan 5 Bab 3

3.1 Exothermic and Endothermic Reactions

Thermochemistry is the study of heat changes during chemical reactions.

Exothermic Reaction: A chemical reaction that releases heat to the surroundings. The surrounding temperature rises. ($\Delta H$ is negative, e.g., combustion, neutralization, reaction between acid and metal).
Endothermic Reaction: A chemical reaction that absorbs heat from the surroundings. The surrounding temperature drops. ($\Delta H$ is positive, e.g., thermal decomposition, photosynthesis, dissolving ammonium salts in water).

Exothermic Reaction: Total energy of reactants is higher than total energy of products ($E_{\text{reactants}} > E_{\text{products}}$). Heat of reaction, $\Delta H = E_{\text{products}} - E_{\text{reactants}} < 0$.
Endothermic Reaction: Total energy of reactants is lower than total energy of products ($E_{\text{reactants}} < E_{\text{products}}$). Heat of reaction, $\Delta H = E_{\text{products}} - E_{\text{reactants}} > 0$.
Activation Energy ($E_a$): The minimum energy that reactant particles must possess in order to react upon collision.

The heat of reaction is the heat change when 1 mole of reactant reacts or 1 mole of product is formed according to the chemical equation.

Heat released or absorbed by solution ($Q$):

$$Q = mc\Delta T$$

$m$ = Mass of solution in grams ($\text{g}$), assuming density of aqueous solution $= 1.0\text{ g cm}^{-3}$ ($1\text{ cm}^3 \approx 1\text{ g}$).
$c$ = Specific heat capacity of solution (typically $4.2\text{ J g}^{-1}\text{ }^\circ\text{C}^{-1}$).
$\Delta T$ = Temperature change in $^\circ\text{C}$ or $\text{K}$ ($\Delta T = T_{\text{final}} - T_{\text{initial}}$).

Heat of reaction ($\Delta H$):

$$\Delta H = \pm \frac{Q}{n}$$

where $n$ is the number of moles of reactant or product formed, measured in $\text{kJ mol}^{-1}$.

The heat change when 1 mole of precipitate is formed from its ions in an aqueous solution.

Example: $\text{Ag}^+\text{(aq)} + \text{Cl}^-\text{(aq)} \rightarrow \text{AgCl(s)}$ ($\Delta H$ is exothermic).

The heat change when 1 mole of a metal is displaced from its salt solution by a more electropositive metal.

Example: $\text{Zn(s)} + \text{Cu}^{2+}\text{(aq)} \rightarrow \text{Zn}^{2+}\text{(aq)} + \text{Cu(s)}$
A metal higher in the electrochemical series causes a larger magnitude of heat release during displacement.

The heat change when 1 mole of water ($\text{H}_2\text{O}$) is formed from the reaction between an acid and an alkali.

$\text{H}^+\text{(aq)} + \text{OH}^-\text{(aq)} \rightarrow \text{H}_2\text{O(l)}$
Strong Acid + Strong Alkali: $\Delta H \approx -57.3\text{ kJ mol}^{-1}$ (constant, as complete ionisation occurs).
Weak Acid or Weak Alkali: Magnitude of $\Delta H < 57.3\text{ kJ mol}^{-1}$ because some heat released is absorbed to completely ionise the weak acid/alkali in water.

The heat released when 1 mole of a substance is completely burnt in excess oxygen ($\text{O}_2$).

Example: $\text{CH}_4\text{(g)} + 2\text{O}_2\text{(g)} \rightarrow \text{CO}_2\text{(g)} + 2\text{H}_2\text{O(l)}$
As the number of carbon atoms per molecule of alcohol increases, the heat of combustion increases because more carbon dioxide and water molecules are produced, forming more bonds and releasing more heat.
Fuel Value ($\text{kJ g}^{-1}$): The amount of heat energy released when $1\text{ g}$ of fuel is completely burnt in excess oxygen.

Bond Breaking: Endothermic process (energy is absorbed to break chemical bonds).
Bond Formation: Exothermic process (energy is released when new bonds are formed).
In an exothermic reaction, total energy released during bond formation is greater than total energy absorbed for bond breaking.