Chemistry | Reversible reactions and equilibria

Welcome to Reversible Reactions and Equilibria!

Hello future chemist! This chapter is one of the most exciting parts of Physical Chemistry because we stop thinking about reactions that just finish and start thinking about reactions that find a perfect, constant balance.

Understanding reversible reactions and chemical equilibrium is vital. It’s how chemical engineers maximize the production of important materials (like ammonia in the Haber process) and how life processes, like maintaining oxygen levels in your blood, stay perfectly regulated.

Don't worry if the term "equilibrium" sounds complicated—we will break it down using simple analogies that make perfect sense!

1. Understanding Reversible Reactions

So far, most of the reactions you have studied are irreversible. This means they go one way: Reactants turn into products, and the reaction stops when one reactant is used up.

Example of irreversible reaction: Burning wood. You can't turn ash back into wood simply by cooling it down!

What is a Reversible Reaction?

A reversible reaction is a reaction where the products can react together to reform the original reactants. The reaction can proceed in both the forward direction (left to right) and the reverse direction (right to left).

Key Symbol:

We use a special symbol called the equilibrium arrow to show a reversible reaction:

Reactants \(\rightleftharpoons\) Products

Classic IGCSE Example: Hydrated Copper(II) Sulfate

When blue hydrated copper(II) sulfate is heated, it loses water (dehydration) and turns into white anhydrous copper(II) sulfate (the forward reaction).

If you add water to the white anhydrous copper(II) sulfate, the blue colour returns (the reverse reaction).

\( \underset{\text{Blue (Hydrated)}}{CuSO_{4} \cdot 5H_{2}O} \underset{\text{Cooling/Adding water}}{\overset{\text{Heating}}{\rightleftharpoons}} \underset{\text{White (Anhydrous)}}{CuSO_{4}} + 5H_{2}O \)

2. Dynamic Equilibrium: The Perfect Balance

If a reversible reaction is happening in a closed system (where nothing can enter or leave, like a sealed container), it will eventually reach a state called chemical equilibrium.

What does "Dynamic" mean?

The word dynamic means active or constantly changing. In chemical equilibrium, everything looks static (unchanging) but everything is still moving!

Analogy: The Crowded Dance Floor

Imagine a dance party where the main room is the "Reactants" side and the balcony is the "Products" side. People are constantly moving between the two.

• If 10 people move from the main room to the balcony every minute (forward rate)...
• ...and 10 people move from the balcony back to the main room every minute (reverse rate)...

The number of people in the main room and the balcony stays constant. The system looks balanced, but the individual people are constantly moving!

Characteristics of Dynamic Equilibrium

When a system reaches dynamic equilibrium:

1. Rates are Equal: The rate of the forward reaction equals the rate of the reverse reaction.
2. Concentrations are Constant: The concentrations of reactants and products remain constant (but usually not equal!).
3. Closed System: Equilibrium can only be established if the system is closed (no mass or energy is exchanged with the surroundings).

Did you know?

If you see the colour of a reversible reaction stop changing, you know it has reached equilibrium because the concentrations (and therefore the colour intensity) have become constant.

3. Le Chatelier's Principle (LCP)

Dynamic equilibrium is a very stable state, but we can knock it off balance! If we change the conditions (like temperature or pressure), the system reacts immediately to counteract that change.

The Principle of the Chemical Tantrum

Le Chatelier's Principle (often shortened to LCP) states that:

"If a change is made to the conditions of a system at equilibrium, the system responds to counteract the change and restore a new equilibrium."

Think of the reaction mixture as having a chemical "tantrum." Whatever stress you put on it, it tries to do the opposite!

• If you add heat, the system reacts to cool down.
• If you increase pressure, the system reacts to lower the pressure.
• If you add reactants, the system tries to use them up.

This shift in balance is called shifting the position of equilibrium.

4. Applying Le Chatelier's Principle: Three Factors

We will examine how three factors—concentration, temperature, and pressure—cause the equilibrium to shift (either to the right, favouring products, or to the left, favouring reactants).

A. Effect of Changing Concentration

This is the most straightforward factor. The system shifts to consume anything that has been added.

\( A + B \rightleftharpoons C + D \)

1. Adding Reactants (A or B):

   The system has too much A or B. To counteract this, the reaction speeds up the forward reaction to use up the excess A and B.
   Result: Equilibrium shifts to the right (more C and D are made).

2. Adding Products (C or D):

   The system has too much C or D. To counteract this, the reaction speeds up the reverse reaction to use up the excess C and D.
   Result: Equilibrium shifts to the left (more A and B are reformed).

3. Removing Products (C or D):

   If we constantly remove the product (C or D) as soon as it is formed, the system shifts right to replace the lost product. This is a common strategy in industry to maximize yield!
   Result: Equilibrium shifts to the right.

B. Effect of Changing Temperature

Temperature is trickier because you must know whether the forward reaction is exothermic (releases heat, \(\Delta H\) is negative) or endothermic (absorbs heat, \(\Delta H\) is positive).

Remember, if the forward reaction is exothermic, the reverse reaction MUST be endothermic, and vice versa.

Mnemonic: Heat likes Endothermic

When you stress the system with heat, it wants to absorb that heat, so it favors the reaction that requires energy (the endothermic reaction).

Let’s use a reaction where the forward reaction is Exothermic:

Reactants \(\rightleftharpoons\) Products (Energy released)

1. Increasing Temperature (Adding Heat):

   The system attempts to cool down by absorbing the added heat. It favours the endothermic reaction (the reverse reaction).
   Result: Equilibrium shifts to the left (less product).

2. Decreasing Temperature (Removing Heat):

   The system attempts to replace the lost heat. It favours the exothermic reaction (the forward reaction, which releases heat).
   Result: Equilibrium shifts to the right (more product).

C. Effect of Changing Pressure (Gaseous Reactions Only)

Changes in pressure only affect reactions involving gases, and they are caused by changing the volume of the container.

The Rule: Pressure is caused by gas particles hitting the walls. More particles = higher pressure.

Step-by-Step Pressure Analysis:

1. Count the total number of moles of gas particles on the left side (reactants). 2. Count the total number of moles of gas particles on the right side (products).

Example: \( N_{2(g)} + 3H_{2(g)} \rightleftharpoons 2NH_{3(g)} \)

Left side total moles = 1 + 3 = 4 moles of gas.
Right side total moles = 2 moles of gas.

Mnemonic: Pressure Pushes to the Fewest

1. Increasing Pressure (Squeezing the container):

   You stress the system with high pressure. The system fights back by trying to reduce the pressure. It does this by favouring the side with the FEWER moles of gas.
   In the Haber example above (4 moles vs 2 moles):
   Result: Equilibrium shifts to the right (towards the 2 moles of product).

2. Decreasing Pressure (Expanding the container):

   You stress the system with low pressure. The system fights back by trying to increase the pressure. It does this by favouring the side with the MORE moles of gas.
   Result: Equilibrium shifts to the left (towards the 4 moles of reactant).

5. The Role of Catalysts

A catalyst is a substance that speeds up a reaction without being chemically changed itself. So, how does it affect equilibrium?

Catalysts and Equilibrium Position

The crucial thing to remember is that a catalyst speeds up the forward reaction and the reverse reaction by exactly the same amount.

Therefore:

A catalyst DOES NOT affect the position of equilibrium. It does not change the amount of product formed (the yield).

Why use a catalyst then?

Even though it doesn't change the final yield, a catalyst helps the system reach dynamic equilibrium much faster. This is incredibly important in industry, where time is money!

Conclusion: You've mastered the balance!

Don't worry if Le Chatelier's Principle seems like a lot to remember. Focus on the core idea: The system always fights back! If you introduce a stress, the reaction shifts to relieve that stress. This principle is fundamental to managing industrial processes efficiently. Great job mastering this essential concept in physical chemistry!