Welcome to the World of Equilibrium!
In most of your chemistry journey so far, you’ve probably looked at reactions as a one-way street: Reactants go in, and Products come out. But in the real world, many reactions are actually a two-way street. Think of it like a busy shop—people are walking in through the front door while others are walking out. If the number of people inside stays the same because people enter and leave at the same speed, you’ve reached equilibrium!
In this chapter, we will learn how reactions find this balance, how we can "nudge" them to get more of what we want, and how this applies to the acids and bases you use every day. Don't worry if this seems tricky at first—once you see the patterns, it all clicks into place!
7.1 Chemical Equilibria: The Balancing Act
1. What is a Reversible Reaction?
A reversible reaction is one where the products can react together to reform the original reactants. We show this using a special double arrow: \( \rightleftharpoons \).
Example: \( N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g) \)
2. Dynamic Equilibrium
When a reversible reaction happens in a closed system (where nothing can get in or out), it eventually reaches dynamic equilibrium.
Two things happen at dynamic equilibrium:
1. The rate of the forward reaction is exactly equal to the rate of the reverse reaction.
2. The concentrations of reactants and products stay constant (they don't change anymore).
Analogy: Imagine walking up an "up" escalator while it’s moving down. If you walk up at the exact same speed the escalator moves down, you stay in the same spot. You are moving (dynamic), but your position doesn't change (equilibrium)!
3. Le Chatelier’s Principle: The "Stubborn" Rule
Definition: If a change is made to a system at dynamic equilibrium, the position of equilibrium moves to minimise this change.
Think of the system like a stubborn teenager: whatever you try to do to it, it will try to do the opposite to cancel you out!
- Change in Concentration: If you add more reactant, the system tries to remove it by moving to the right (making more product).
- Change in Pressure: (Only affects gases!) If you increase pressure, the system moves to the side with fewer gas molecules to lower the pressure back down.
- Change in Temperature:
- If you increase heat, the system tries to cool down by moving in the endothermic direction.
- If you decrease heat, the system tries to warm up by moving in the exothermic direction. - Catalysts: A catalyst increases the rate of both the forward and reverse reactions equally. Therefore, a catalyst does not change the position of equilibrium; it just helps you get there faster!
Quick Review:
- Closed system: Required for equilibrium.
- Equal rates: Required for equilibrium.
- Catalyst: No effect on equilibrium position.
4. Equilibrium Constants: \( K_c \) and \( K_p \)
We use constants to describe exactly where the balance lies.
For \( K_c \) (using Concentrations):
For the reaction: \( aA + bB \rightleftharpoons cC + dD \)
\( K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b} \)
(Square brackets \([ ]\) mean concentration in \( mol\ dm^{-3} \)).
For \( K_p \) (using Partial Pressures):
First, you need to know Mole Fraction and Partial Pressure:
- Mole Fraction: \( \frac{\text{moles of a gas}}{\text{total moles of all gases}} \)
- Partial Pressure: \( \text{mole fraction} \times \text{total pressure} \)
\( K_p \) is calculated just like \( K_c \), but using partial pressures (\( p \)) instead of concentrations:
\( K_p = \frac{(pC)^c(pD)^d}{(pA)^a(pB)^b} \)
Important! Only temperature changes the actual value of \( K_c \) or \( K_p \). Concentration, pressure, and catalysts have no effect on the value of the constant.
5. Industrial Applications: Haber and Contact Processes
In industry, we want the most product for the least money. We use Le Chatelier’s principle to find compromise conditions.
The Haber Process (Making Ammonia)
\( N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g) \) (\( \Delta H \) is negative/exothermic)
- Temperature: Low temp would give more product but would be too slow. We use a compromise temp (400–450°C).
- Pressure: High pressure (20,000 kPa) shifts equilibrium to the right (fewer molecules).
- Catalyst: Iron.
The Contact Process (Making Sulfur Trioxide)
\( 2SO_2(g) + O_2(g) \rightleftharpoons 2SO_3(g) \) (\( \Delta H \) is negative/exothermic)
- Temperature: Compromise (400–450°C).
- Pressure: Just above atmospheric (100–200 kPa) because the yield is already very high.
- Catalyst: Vanadium(V) oxide (\( V_2O_5 \)).
Key Takeaway: Equilibrium is a balance. If you change the conditions, the system shifts to restore the balance. \( K_c \) and \( K_p \) tell us mathematically how much product we have compared to reactants.
7.2 Brønsted–Lowry Theory: Acids and Bases
1. Defining Acids and Bases
Forget what you learned in middle school! In AS Chemistry, we define acids and bases by what they do with protons (an \( H^+ \) ion is just a proton).
- Brønsted–Lowry Acid: A proton (\( H^+ \)) donor.
- Brønsted–Lowry Base: A proton (\( H^+ \)) acceptor.
2. Common Acids and Alkalis You Must Know
Acids:
- Hydrochloric Acid: \( HCl \)
- Sulfuric Acid: \( H_2SO_4 \)
- Nitric Acid: \( HNO_3 \)
- Ethanoic Acid: \( CH_3COOH \)
Alkalis (Bases that dissolve in water):
- Sodium Hydroxide: \( NaOH \)
- Potassium Hydroxide: \( KOH \)
- Ammonia: \( NH_3 \)
3. Strong vs. Weak
This is where many students get confused. "Strong" does not mean "concentrated." It refers to dissociation (splitting up into ions).
- Strong Acids/Bases: Fully dissociate into ions in water. (e.g., \( HCl \rightarrow H^+ + Cl^- \))
- Weak Acids/Bases: Only partially dissociate. This creates an equilibrium! (e.g., \( CH_3COOH \rightleftharpoons CH_3COO^- + H^+ \))
Did you know? Because weak acids have fewer ions in solution, they are poorer conductors of electricity compared to strong acids of the same concentration!
4. The pH Scale
The pH scale measures how acidic or alkaline a solution is.
- pH < 7: Acidic
- pH = 7: Neutral (Pure water)
- pH > 7: Alkaline
5. Neutralisation and Titrations
Neutralisation happens when an acid and a base react to form a salt and water. The ionic equation is almost always:
\( H^+(aq) + OH^-(aq) \rightarrow H_2O(l) \)
When we perform a titration, we can plot the pH on a graph. The shape of the titration curve depends on the strength of the acid and base:
- Strong Acid + Strong Base: Large vertical section, equivalence point at pH 7.
- Weak Acid + Strong Base: Equivalence point at pH > 7.
- Strong Acid + Weak Base: Equivalence point at pH < 7.
Memory Trick for Indicators:
- Methyl Orange: Use for Strong Acid titrations (Red in acid, Yellow in alkali).
- Phenolphthalein: Use for Strong Base titrations (Colourless in acid, Pink in alkali).
Common Mistake: Thinking \( NH_3 \) is an acid because it has H atoms. It is actually a base because its lone pair of electrons can accept a proton to become \( NH_4^+ \)!
Key Takeaway: Acids donate protons; bases accept them. Strong species split up completely, while weak species only split up a little bit, creating an equilibrium.
Congratulations! You've just covered the essentials of Equilibrium and Acids/Bases. Keep practicing those \( K_c \) expressions, and you'll be an expert in no time!