Chemical Equilibria

Introduction to Chemical Equilibria

Welcome to the study of Chemical Equilibria! Up until now, you might have thought of chemical reactions as one-way streets: you mix reactants, they turn into products, and that is the end of the story. However, in the real world, many reactions are reversible. They can go forwards and backwards at the same time!

In this chapter, we will explore what happens when a reaction "settles down" into a state of balance. Understanding this is vital for chemists, especially in industries where making as much product as possible, as quickly as possible, is the goal. Don't worry if it feels a bit "push and pull" at first—it’s actually quite logical once you see the patterns!

1. What is Dynamic Equilibrium?

Before we dive in, let's look at the symbol for a reversible reaction: \(\rightleftharpoons\). This means the reaction can go from left to right (forward) and right to left (backward).

Imagine a busy clothing store. If people are entering the store at the exact same rate that people are leaving, the total number of people inside stays the same. This is exactly what Dynamic Equilibrium is like!

Key Features of Dynamic Equilibrium:

• It only happens in a closed system (where no substances can get in or out).
• The rate of the forward reaction is exactly equal to the rate of the backward reaction.
• The concentrations of reactants and products remain constant (they stop changing).

Common Mistake to Avoid: Many students think that at equilibrium, the concentrations of reactants and products are equal. This is usually not true! They just stay at a steady level. Think of the store analogy: there might be 50 people inside and 1,000 people outside. As long as 5 enter and 5 leave every minute, the numbers are constant, even though they aren't equal.

Quick Review: At equilibrium, the reaction hasn't stopped; it's just moving at the same speed in both directions!

2. Changing the Position of Equilibrium

When we talk about the "position" of equilibrium, we are asking: "Is there more product or more reactant in the mixture?" To predict how the system reacts to changes, we use Le Chatelier’s Principle.

The Principle: If a system at equilibrium is disturbed by a change in conditions, the system shifts to counteract (oppose) that change.

Analogy: Think of a stubborn teenager. If you tell them to turn the music up, they turn it down. If you tell them to move left, they move right. The equilibrium system is just as stubborn!

A. Changing Concentration

If you increase the concentration of a reactant, the system wants to decrease it. It does this by "using up" the extra reactant and moving in the forward direction to make more product.

• Add Reactant: Equilibrium shifts to the right (makes more product).
• Remove Product: Equilibrium shifts to the right (tries to replace what was lost).

B. Changing Pressure

This only affects gases. To figure this out, you must count the number of gas molecules (moles) on each side of the balanced equation.

• Increase Pressure: The system wants to lower the pressure. It shifts to the side with fewer gas molecules.
• Decrease Pressure: The system wants to increase the pressure. It shifts to the side with more gas molecules.

Example: \(N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)\)
Left side = 4 moles of gas. Right side = 2 moles of gas.
If you increase pressure, it shifts to the right (the side with fewer molecules).

C. Changing Temperature

To predict this, you must know if the forward reaction is exothermic (gives out heat, \(-\Delta H\)) or endothermic (takes in heat, \(+\Delta H\)).

• Increase Temperature: The system wants to cool down. It shifts in the endothermic direction to absorb the extra heat.
• Decrease Temperature: The system wants to warm up. It shifts in the exothermic direction to release more heat.

Memory Trick:
Increase Temperature \(\rightarrow\) In the Endothermic direction (I.T.I.E.)

Did you know? Adding a catalyst does NOT change the position of equilibrium. It speeds up both the forward and backward reactions equally, so you just reach the same equilibrium point faster!

Key Takeaway: The system always tries to do the opposite of what you did. If you add heat, it tries to remove it. If you add pressure, it tries to lower it.

3. Industrial Compromises

In a factory, chemists have a difficult choice. They want a high yield (lots of product), but they also need a fast rate (making it quickly) to be profitable. Often, the conditions that give a high yield are very slow, or the conditions that are fast give a low yield.

The Conflict:

Imagine a reaction where the forward step is exothermic.

• For high yield: We should use a low temperature (Le Chatelier’s Principle).
• For high rate: We should use a high temperature (Collision Theory).

The Solution: Scientists use a compromise temperature. This temperature is high enough to make the reaction happen at a decent speed, but low enough to ensure they still get a reasonable amount of product. They also use catalysts to boost the rate without needing extreme temperatures.

Common Exam Question: Why is a pressure of 200 atmospheres used in the Haber Process instead of 1000 atmospheres?
Answer: Even though higher pressure increases yield, it is very expensive to build strong pipes and uses a lot of energy to maintain. 200 atm is a compromise between yield and cost/safety.

4. Practical Examples from the Syllabus

You may be asked to describe or justify observations in these specific systems:

The \(NO_2\) / \(N_2O_4\) System

\(2NO_2(g) \rightleftharpoons N_2O_4(g)\)
(\(NO_2\) is brown; \(N_2O_4\) is colourless. The forward reaction is exothermic.)

• If you heat it: It shifts in the endothermic (backward) direction. The gas becomes more brown.
• If you increase pressure: It shifts to the side with fewer molecules (the right). The gas becomes paler as colourless \(N_2O_4\) forms.

Iodine Chlorides

\(ICl(l) + Cl_2(g) \rightleftharpoons ICl_3(s)\)
(\(ICl\) is a dark brown liquid; \(ICl_3\) is a yellow solid.)

• Add Chlorine gas: The equilibrium shifts to the right to use up the extra \(Cl_2\). You will see more yellow solid form.
• Remove Chlorine gas: The equilibrium shifts to the left. The yellow solid will disappear and more brown liquid will form.

Don't worry if this seems tricky! Just always ask yourself two questions: 1. What change did I make? 2. What is the opposite of that change? That is the direction the equilibrium will move.

Summary: The Equilibrium Cheat Sheet

1. Concentration: Add Reactant \(\rightarrow\) Move Right. Remove Reactant \(\rightarrow\) Move Left.
2. Pressure: Increase Pressure \(\rightarrow\) Move to side with fewer gas moles.
3. Temperature: Increase Temp \(\rightarrow\) Move in Endothermic direction.
4. Catalysts: No change to equilibrium position, just get there faster.
5. Industry: Uses compromise conditions to balance speed (rate), amount (yield), and cost.