Chemistry | Introduction to Kinetics and Equilibria

Welcome to Introduction to Kinetics and Equilibria!

Hello future Chemists! This chapter is incredibly important because it moves us beyond just what products form (stoichiometry) to how fast they form (kinetics) and how far the reaction goes (equilibria). These concepts link directly to how we choose reaction conditions in industry, including how we manufacture key chemicals like alcohols and polymers, and how we control reactions involving halogenoalkanes.

Don't worry if these ideas seem abstract at first. We will break them down using everyday examples. Let's get started!

Part 1: Chemical Kinetics – The Speed of Reaction

Chemical Kinetics is the study of reaction rates and the factors that influence them. In simple terms, it answers the question: How quickly do reactants turn into products?

1.1 Defining the Rate of Reaction

The Rate of Reaction is the change in concentration of a reactant or product over time. We measure this rate by monitoring an observable change, such as:

• The volume of gas produced.
• The change in mass (if gas is escaping).
• The change in colour intensity (spectrometry).
• The change in turbidity (how cloudy the solution is).

\( \text{Rate} = \frac{\text{Change in concentration (or mass/volume)}}{\text{Change in time}} \)

1.2 The Prerequisite: Collision Theory

For a reaction to occur, the particles (atoms, ions, or molecules) must physically collide. However, most collisions are unsuccessful. The Collision Theory states that a reaction will only happen if three conditions are met:

1. Collision: The particles must physically hit each other.
2. Correct Orientation: They must hit each other in the correct alignment so that the parts that need to react actually touch.
   Analogy: Imagine trying to fit a specific key into a lock. You have to put the key in the right way up! A random bump won't work.
3. Sufficient Energy (Activation Energy): They must collide with enough energy to overcome the energy barrier needed to break old bonds and form new ones.

A collision that meets all three criteria is called an Effective Collision.

1.3 Activation Energy (\(E_a\))

The minimum amount of kinetic energy that particles must possess for a collision to result in a reaction is called the Activation Energy (\(E_a\)). Think of it as an energy "hurdle" the particles must jump over.

Did you know? Even highly exothermic reactions (those that release lots of energy overall) still need an initial input of energy to get them started. This is the \(E_a\).

1.4 Temperature and the Boltzmann Distribution

To understand the massive impact temperature has on reaction rate, we use the Boltzmann Distribution curve. This curve shows the distribution of kinetic energies among the molecules in a sample at a given temperature.

Key Features of the Boltzmann Curve:

• Most molecules have energy near the peak of the curve.
• Very few molecules have very low or very high energy.
• We mark the position of the Activation Energy (\(E_a\)) on the curve. Only molecules to the right of this line have enough energy to react.

Effect of Increasing Temperature:

When you increase the temperature (e.g., from \(T_1\) to \(T_2\)):

1. The curve flattens and shifts slightly to the right.
2. The average kinetic energy of the molecules increases.
3. Crucially: A much larger proportion of molecules now have energy greater than or equal to \(E_a\).

Because there are exponentially more high-energy molecules, the number of effective collisions increases dramatically, leading to a much faster reaction rate. A 10 °C rise often doubles the reaction rate!

1.5 Factors Affecting Reaction Rate

Here is how the four main factors increase the rate, explained using Collision Theory:

A. Concentration (or Pressure for Gases)

Effect: Increasing concentration (or pressure) increases the rate.

Why? More particles are packed into the same volume, meaning they are closer together. This leads to a higher frequency of collision. More collisions overall means more effective collisions per second.

B. Temperature

Effect: Increasing temperature increases the rate.

Why? (Two reasons, both major contributors):
1. Increased kinetic energy means higher frequency of collision.
2. A much higher proportion of particles exceed \(E_a\) (as explained by the Boltzmann Distribution).

C. Surface Area (for reactions involving solids)

Effect: Increasing surface area increases the rate (e.g., using powdered zinc instead of a lump).

Why? Only particles on the surface of a solid can collide with the other reactant. Breaking a lump into smaller pieces exposes more internal atoms, providing more sites for collisions, thereby increasing the frequency of collision.

D. Catalysts

Effect: A catalyst increases the rate without being used up itself.

Why? A catalyst provides an alternative reaction pathway that has a lower Activation Energy (\(E_a\)). On the Boltzmann Distribution curve, lowering the \(E_a\) line means a huge increase in the proportion of molecules that can react.

Important note on Catalysts:

• Homogeneous Catalysis: The catalyst is in the same physical state as the reactants (e.g., all liquid).
• Heterogeneous Catalysis: The catalyst is in a different physical state, usually a solid acting on liquid or gas reactants (e.g., using a transition metal surface for gaseous reactions).

Part 2: Chemical Equilibria – Reactions That Settle Down

Not all reactions go to completion. Many are reversible. This leads us to the concept of chemical equilibrium.

2.1 Reversible Reactions and Dynamic Equilibrium

A Reversible Reaction is one where the products can react together to reform the original reactants. We show this using double arrows:

\( \text{Reactants} \rightleftharpoons \text{Products} \)

When a reversible reaction is carried out in a closed system (meaning nothing can enter or leave), it will eventually reach a state called Dynamic Equilibrium.

Characteristics of Dynamic Equilibrium:

1. Dynamic: Reactions are still occurring! The forward reaction (Reactants \(\rightarrow\) Products) and the reverse reaction (Products \(\rightarrow\) Reactants) continue ceaselessly.
2. Equal Rates: The rate of the forward reaction is exactly equal to the rate of the reverse reaction.
3. Constant Concentration: Because the rates are equal, the macroscopic concentrations of the reactants and products remain constant (they don't have to be equal, just unchanging).

Analogy: Imagine a very popular two-story building. If 10 people go up the stairs every minute, and 10 people come down the stairs every minute, the total number of people on the ground floor and the total number on the first floor remain constant, even though there is constant movement. That is dynamic equilibrium.

2.2 Disturbing the Balance: Le Chatelier's Principle (LCP)

If we have a system at dynamic equilibrium, and we change the conditions (temperature, pressure, or concentration), the system will try to counteract that change. This is summarized by Le Chatelier's Principle (LCP):

LCP: When a system at equilibrium is subjected to a change, the system will respond to relieve the stress and restore equilibrium.

The system shifts the equilibrium position either to the right (favouring products) or to the left (favouring reactants).

2.3 Applying LCP: The Three Stresses

A. Change in Concentration

• Stress: Increase the concentration of a reactant (left side).
  Response: The system shifts right to consume the added reactant and make more product.
• Stress: Decrease the concentration of a product (right side).
  Response: The system shifts right to replace the removed product.

Memory Trick: If you add stuff, the reaction runs away from it. If you take stuff away, the reaction runs toward it to replace it.

B. Change in Temperature

This is the trickiest one, as you must know the enthalpy change (\(\Delta H\)) for the reaction:

Scenario 1: Endothermic Reaction (\(\Delta H\) is positive, heat is absorbed)
We can treat heat as a reactant: \(\text{Reactants} + \text{Heat} \rightleftharpoons \text{Products}\)

• Stress: Increase Temperature (Add Heat).
  Response: The system shifts right to use up the added heat. (Favouring products)
• Stress: Decrease Temperature (Remove Heat).
  Response: The system shifts left to generate heat. (Favouring reactants)

Scenario 2: Exothermic Reaction (\(\Delta H\) is negative, heat is released)
We can treat heat as a product: \(\text{Reactants} \rightleftharpoons \text{Products} + \text{Heat}\)

• Stress: Increase Temperature (Add Heat).
  Response: The system shifts left to use up the added heat. (Favouring reactants)
• Stress: Decrease Temperature (Remove Heat).
  Response: The system shifts right to replace the removed heat. (Favouring products)

C. Change in Pressure (Applies only to Gases)

Pressure changes are caused by changing the volume of the container. The key is to count the total number of moles of gas on each side of the equation.

• Stress: Increase Pressure (Squeeze the system).
  Response: The system shifts to the side with the fewer moles of gas to reduce the total number of particles and relieve the pressure.
• Stress: Decrease Pressure (Expand the system).
  Response: The system shifts to the side with the more moles of gas to increase the total number of particles.

Example: \( \text{N}_2(\text{g}) + 3\text{H}_2(\text{g}) \rightleftharpoons 2\text{NH}_3(\text{g}) \)
(Left side has 4 moles of gas; Right side has 2 moles of gas).
If we increase pressure, the system shifts right (favouring ammonia) because 2 moles is less than 4 moles.

Important Exception: If the number of moles of gas is the same on both sides (e.g., \(1+1 \rightleftharpoons 2\)), then changing the pressure has no effect on the equilibrium position.

2.4 The Effect of a Catalyst on Equilibrium

This is a common exam trick! You learned that catalysts speed up reactions, but they affect both the forward and reverse reaction rates equally.

Catalysts DO NOT affect the position of equilibrium.

They simply allow the system to reach the position of dynamic equilibrium much faster. This is vital in industrial chemistry, such as the synthesis of alcohols, where we need high rates but also good yields (i.e., a favourable equilibrium position).

Summary of the Chapter

Kinetics tells us how fast we get there (governed by \(E_a\)). Equilibria tells us where we end up (the ratio of products to reactants), which is governed by LCP. These two concepts are often balanced in industry—for instance, we might choose a compromise temperature that is high enough for a good rate, even if it slightly lowers the yield (the equilibrium position).