Revision of Physical Chemistry Rates of reaction, equilibrium, gas laws, energetics.

CHEMISTRY S S 3: Physical Chemistry Revision

Welcome, SS3 Chemistry students! This note will help you revise key concepts in Physical Chemistry: Rates of reaction, chemical equilibrium, gas laws, and energetics. Understanding these topics is crucial for explaining how chemical reactions occur and behave.

Core Concepts

1. Rates of Reaction

The rate of reaction describes how fast reactants are used up or how fast products are formed. It's often measured as the change in concentration of a reactant or product per unit time (e.g., mol/dm³/s).

Factors affecting reaction rate:

• Concentration: Increasing reactant concentration usually increases the rate because there are more particles per unit volume, leading to more frequent collisions.
• Temperature: Raising the temperature increases the kinetic energy of particles, causing them to move faster and collide more frequently and with greater energy. This increases the number of effective collisions.
• Surface Area: For solid reactants, increasing the surface area (e.g., by crushing a lump into powder) exposes more particles to react, thus increasing the rate.
• Catalyst: A catalyst speeds up a reaction without being consumed itself. It does this by providing an alternative reaction pathway with a lower activation energy.
• Pressure (for gases): Increasing the pressure of gaseous reactants brings the particles closer, increasing collision frequency and reaction rate.

Collision Theory explains that reactions occur when reactant particles collide with sufficient energy (activation energy) and in the correct orientation. Example: Sugar dissolves faster in hot water (higher temperature) than in cold water. Powdered sugar also dissolves faster than sugar cubes (larger surface area).

2. Chemical Equilibrium

Chemical equilibrium is a state in a reversible reaction where the rate of the forward reaction equals the rate of the reverse reaction. At equilibrium, the concentrations of reactants and products remain constant, but the reactions are still occurring (dynamic equilibrium).

Le Chatelier's Principle states that if a change of condition (like temperature, pressure, or concentration) is applied to a system at equilibrium, the system will shift in a direction that relieves the stress.

• Concentration: Adding more reactant shifts the equilibrium to favour product formation. Removing product shifts it to favour more product formation.
• Pressure (for gases): Increasing pressure shifts the equilibrium to the side with fewer moles of gas. Decreasing pressure shifts it to the side with more moles of gas.
• Temperature: Increasing temperature shifts equilibrium in the endothermic direction (absorbs heat). Decreasing temperature shifts it in the exothermic direction (releases heat).
• Catalyst: A catalyst speeds up both forward and reverse reactions equally, so it does not shift the equilibrium position but helps the system reach equilibrium faster. Example: The industrial production of ammonia (Haber process) uses high pressure and moderate temperature to shift the equilibrium towards ammonia formation.

3. Gas Laws

Gas laws describe the relationship between pressure (P), volume (V), temperature (T), and the amount of gas (n).

• Boyle's Law: At constant temperature and amount of gas, the volume of a gas is inversely proportional to its pressure (P₁V₁ = P₂V₂).
• Charles's Law: At constant pressure and amount of gas, the volume of a gas is directly proportional to its absolute temperature (V₁/T₁ = V₂/T₂).
• Gay-Lussac's Law: At constant volume and amount of gas, the pressure of a gas is directly proportional to its absolute temperature (P₁/T₁ = P₂/T₂).
• Combined Gas Law: Combines Boyle's, Charles's, and Gay-Lussac's Laws (P₁V₁/T₁ = P₂V₂/T₂).
• Ideal Gas Equation: PV = nRT, where R is the ideal gas constant. This equation relates all four variables for an ideal gas.

Remember to always use absolute temperature (Kelvin) for gas law calculations. K = °C + 273. Example: When you pump air into a bicycle tyre, you increase the amount of air (n) and its pressure (P) inside the fixed volume (V) of the tyre.

4. Energetics (Thermochemistry)

Energetics is the study of heat changes that accompany chemical reactions. Every chemical reaction involves an energy change, usually in the form of heat.

• Enthalpy Change (ΔH): Represents the heat absorbed or released during a chemical reaction at constant pressure.
• Exothermic Reactions: Release heat to the surroundings. The products have lower energy than the reactants, so ΔH is negative (ΔH < 0). Examples include combustion and neutralisation.
• Endothermic Reactions: Absorb heat from the surroundings. The products have higher energy than the reactants, so ΔH is positive (ΔH > 0). Examples include photosynthesis and dissolving ammonium nitrate in water.

Breaking chemical bonds requires energy (endothermic process), while forming new bonds releases energy (exothermic process). The overall energy change of a reaction depends on the balance between energy required for bond breaking and energy released during bond formation. Example: When you light a match, the combustion is an exothermic reaction, releasing heat and light. An instant cold pack, when activated, involves an endothermic reaction that absorbs heat from its surroundings, making it feel cold.

Real-World Example & Practice

The principles of physical chemistry are essential in industrial processes. For instance, in manufacturing chemicals like sulfuric acid (contact process) or ammonia (Haber process), understanding reaction rates helps engineers design efficient reactors, while knowledge of equilibrium and energetics guides them in choosing optimal temperature and pressure conditions to maximise product yield and energy efficiency.

Practice/Reflection Question: Consider a charcoal grill. How do you apply concepts of reaction rate and energetics to get the charcoal burning effectively and cooking your food? (Think about lighting, fanning, and adding more charcoal.)