Welcome to the World of Balance: Equilibria
In your chemistry journey so far, you might have thought that reactions only go one way—from reactants to products. But in the real world, many reactions are like a two-way street! In this chapter, we will explore Chemical Equilibria. You will learn how reactions find a "balance," how we can nudge them to produce more of what we want, and how this applies to the acids and bases you use every day.
Don’t worry if this seems a bit abstract at first; we’ll use plenty of analogies to keep things grounded!
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 reactants. We use a special double arrow \(\rightleftharpoons\) to show this.
Example: \(N_{2}(g) + 3H_{2}(g) \rightleftharpoons 2NH_{3}(g)\)
2. Dynamic Equilibrium
Imagine you are running up an escalator that is moving down. If you run at the exact same speed the escalator moves down, you stay in the same place. This is Dynamic Equilibrium!
In a chemical system at equilibrium:
• The rate of the forward reaction is equal to the rate of the reverse reaction.
• The concentrations of reactants and products remain constant (they aren't changing anymore, but the reaction hasn't stopped!).
• This can only happen in a closed system (where no substances can get in or out).
3. Le Chatelier’s Principle
This is a fancy way of saying: "If you change the conditions, the reaction will try its best to undo what you did."
The Definition: If a change is made to a system at dynamic equilibrium, the position of equilibrium moves to minimise this change.
How can we "annoy" an equilibrium?
• Concentration: If you add more reactant, the system tries to remove it by moving to the right (making more product).
• Pressure (Gases only): If you increase the pressure, the system moves to the side with fewer moles of gas to reduce the pressure.
• Temperature:
- If you increase temperature, the system tries to cool down by moving in the endothermic direction (\(\Delta H\) is positive).
- If you decrease temperature, it moves in the exothermic direction (\(\Delta H\) is negative).
• Catalysts: A catalyst does not change the position of equilibrium. It just helps the system reach equilibrium faster by speeding up both the forward and backward reactions equally.
Quick Review: Equilibrium is a "dynamic" balance in a closed box. Le Chatelier’s Principle helps us predict which way the "seesaw" will tilt if we push it.
7.2 Equilibrium Constants: \(K_c\) and \(K_p\)
To be precise chemists, we need numbers to describe the equilibrium. We use Equilibrium Constants.
1. \(K_c\) (Concentration Constant)
For a general reaction: \(aA + bB \rightleftharpoons cC + dD\)
The expression is: \(K_{c} = \frac{[C]^{c}[D]^{d}}{[A]^{a}[B]^{b}}\)
Note: Square brackets \([ ]\) mean concentration in \(mol\ dm^{-3}\). Always put Products on top!
2. \(K_p\) (Gas Pressure Constant)
When dealing with gases, it’s easier to use partial pressures (\(p\)).
The expression looks the same, but uses pressure: \(K_{p} = \frac{p(C)^{c} \times p(D)^{d}}{p(A)^{a} \times p(B)^{b}}\)
Key Terms for \(K_p\):
• Mole Fraction: The proportion of a specific gas in a mixture.
\(\text{Mole fraction of gas A} = \frac{\text{moles of gas A}}{\text{total moles of all gases}}\)
• Partial Pressure: The pressure exerted by one specific gas in a mixture.
\(p(A) = \text{mole fraction of A} \times \text{total pressure}\)
Common Mistake to Avoid: Only Temperature changes the numerical value of \(K_c\) or \(K_p\). Changing concentration or pressure might shift the "position" of equilibrium, but the constant \(K\) stays the same unless the temperature changes!
7.3 Equilibrium in Industry
In industry, time is money. Scientists use Le Chatelier’s Principle to get the best "yield" (amount of product).
1. The Haber Process (Making Ammonia)
\(N_{2}(g) + 3H_{2}(g) \rightleftharpoons 2NH_{3}(g)\) (\(\Delta H = -92\ kJ\ mol^{-1}\))
• Pressure: High pressure (approx 200 atm) shifts equilibrium to the right (fewer gas moles).
• Temperature: The forward reaction is exothermic. Low temperature would give a high yield but would be too slow. A compromise temperature (approx 450°C) is used.
• Catalyst: Iron is used to speed up the process.
2. The Contact Process (Making Sulfur Trioxide)
\(2SO_{2}(g) + O_{2}(g) \rightleftharpoons 2SO_{3}(g)\) (\(\Delta H = -197\ kJ\ mol^{-1}\))
• Similar logic to Haber: A compromise temperature (450°C) and a catalyst (Vanadium(V) oxide, \(V_{2}O_{5}\)) are used.
7.4 Brønsted–Lowry Theory: Acids and Bases
Forget what you learned in middle school! At AS Level, we define acids and bases by what they do with protons (\(H^{+}\) ions).
1. Definitions
• Brønsted–Lowry Acid: A proton donor (it gives away \(H^{+}\)).
• Brønsted–Lowry Base: A proton acceptor (it takes in \(H^{+}\)).
2. Common Chemicals to Know
Acids: Hydrochloric acid (\(HCl\)), Sulfuric acid (\(H_{2}SO_{4}\)), Nitric acid (\(HNO_{3}\)), and Ethanoic acid (\(CH_{3}COOH\)).
Alkalis/Bases: Sodium hydroxide (\(NaOH\)), Potassium hydroxide (\(KOH\)), and Ammonia (\(NH_{3}\)).
3. Strong vs. Weak
This is about dissociation (splitting up in water):
• Strong Acids/Bases: They split up 100% into ions.
Example: \(HCl \rightarrow H^{+} + Cl^{-}\)
• Weak Acids/Bases: They only split up a tiny bit (the reaction is reversible!).
Example: \(CH_{3}COOH \rightleftharpoons CH_{3}COO^{-} + H^{+}\)
How to tell them apart in a lab?
• pH: Strong acids have lower pH (0-2) than weak acids (3-5) of the same concentration.
• Conductivity: Strong acids conduct electricity better because they have more free-moving ions.
• Reactivity: Strong acids react much faster with metals or carbonates.
Did you know? Water can act as both an acid and a base! It has a pH of 7 because it contains a tiny, equal amount of \(H^{+}\) and \(OH^{-}\) ions.
7.5 Titrations and Indicators
When you add an acid to a base, they neutralise each other: \(H^{+}(aq) + OH^{-}(aq) \rightarrow H_{2}O(l)\). This produces a salt and water.
1. Titration Curves
A titration curve is a graph of pH vs. Volume of acid/base added. They usually look like an "S" shape.
• The Vertical Section: This is where the pH changes very rapidly. The middle of this section is the equivalence point (where acid and base have reacted in the exact stoichiometric ratio).
• Strong Acid + Strong Base: Vertical section is long (pH 3 to 11).
• Weak Acid + Strong Base: Vertical section is shorter (pH 7 to 11).
2. Choosing an Indicator
An indicator is just a weak acid that changes colour at a specific pH. To pick the right one, its working range must fall within the vertical section of your titration curve.
• If your vertical section is between pH 8 and 10, use Phenolphthalein.
• If your vertical section is between pH 3 and 5, use Methyl Orange.
Key Takeaway: Acids donate protons; bases accept them. Strength is about how much they split up. In titrations, the indicator must "flip" its colour during the steep pH jump!