Oceans (O): Equilibria (Acid–Base)

Welcome to one of the most vital chapters in your A Level Chemistry journey! In the Oceans (O) storyline, we explore how the vast waters of our planet stay balanced. This chapter is all about Acid–Base Equilibria. We will look at how scientists define acids and bases, how we calculate the "strength" of a solution (pH), and how the ocean acts as a giant "buffer" to protect marine life from changes in acidity. Don't worry if the math looks a bit scary at first—we'll break it down step-by-step!

1. The Brønsted–Lowry Theory

Before we can calculate anything, we need to know what we are looking at. In the Brønsted–Lowry theory, the whole game is about a single particle: the proton (a hydrogen ion, \(H^+\)).

Proton Donors and Acceptors

In any acid–base reaction, think of the proton like a tennis ball being hit across a net:

  • Acids are proton donors. They "hit" the proton away.
  • Bases are proton acceptors. They "catch" the proton.

Conjugate Acid–Base Pairs

When an acid gives away its proton, it doesn't just disappear—it becomes a species that is now capable of taking a proton back. This is called a conjugate pair. They are like a "before and after" photo of the same molecule.

Example: When \(HCl\) reacts with water:
\(HCl (aq) + H_2O (l) \rightleftharpoons H_3O^+ (aq) + Cl^- (aq)\)

  • \(HCl\) is the acid (it donated \(H^+\)).
  • \(Cl^-\) is its conjugate base (it’s what’s left over).
  • \(H_2O\) is the base (it accepted \(H^+\)).
  • \(H_3O^+\) (the hydronium ion) is its conjugate acid.

Memory Aid: B.A.D. — Bases Accept, Donors are Acids.

Quick Review:

An acid–base reaction is just a proton transfer. To find a conjugate base, just remove one \(H^+\). To find a conjugate acid, just add one \(H^+\).

2. Strong Acids and Bases

In Chemistry, "Strong" doesn't mean "highly concentrated." It refers to how much the substance ionises (splits apart) in water.

Strong Acids

A strong acid fully ionises in water. It is "all-in." Every single molecule of \(HCl\) you put into water will turn into \(H^+\) and \(Cl^-\).

Equation: \(HCl (aq) \rightarrow H^+ (aq) + Cl^- (aq)\)

Strong Bases

Similarly, a strong base (like \(NaOH\)) fully ionises into metal ions and hydroxide ions (\(OH^-\)).

Equation: \(NaOH (aq) \rightarrow Na^+ (aq) + OH^- (aq)\)

Key Takeaway: Because they ionise 100%, the concentration of \(H^+\) in a strong monoprotic acid is exactly the same as the concentration of the acid itself. If you have \(0.1\ mol\ dm^{-3}\) of \(HCl\), you have \(0.1\ mol\ dm^{-3}\) of \(H^+\).

3. The pH Scale and Calculations

The pH scale is a way of measuring the concentration of \(H^+\) ions. Because these concentrations are often tiny (like \(0.0000001\)), we use a "log" scale to make the numbers easier to handle.

The Formulas You Need:

1. \(pH = -\log_{10}[H^+]\)
2. \([H^+] = 10^{-pH}\)

Calculating pH for Strong Acids:

Since a strong acid like \(HCl\) fully splits, \([H^+] = [Acid]\).

Step-by-step example: Find the pH of \(0.05\ mol\ dm^{-3}\) \(HNO_3\).
1. Identify it’s a strong acid, so \([H^+] = 0.05\).
2. \(pH = -\log_{10}(0.05) = 1.30\).

Calculating pH for Strong Bases (using \(K_w\)):

Water always has a tiny bit of equilibrium happening: \(H_2O \rightleftharpoons H^+ + OH^-\).
The ionic product of water, \(K_w\), is a constant: \(K_w = [H^+][OH^-]\).
At \(298K\), \(K_w\) is always \(1.0 \times 10^{-14}\ mol^2\ dm^{-6}\).

Step-by-step example: Find the pH of \(0.1\ mol\ dm^{-3}\) \(NaOH\).
1. Since \(NaOH\) is a strong base, \([OH^-] = 0.1\).
2. Use \(K_w\) to find \([H^+]\): \([H^+] = \frac{K_w}{[OH^-]} = \frac{1.0 \times 10^{-14}}{0.1} = 1.0 \times 10^{-13}\).
3. \(pH = -\log_{10}(1.0 \times 10^{-13}) = 13\).

Common Mistake: Don't forget that pH + pOH = 14 at room temperature. If you calculate a pH of 1 for a base, something has gone wrong!

4. Weak Acids and the Acidity Constant (\(K_a\))

Weak acids (like ethanoic acid or the carbonic acid in the oceans) only partially ionise. Most of the molecules stay stuck together. This is a dynamic equilibrium.

The \(K_a\) Expression

For a weak acid \(HA \rightleftharpoons H^+ + A^-\):
\(K_a = \frac{[H^+][A^-]}{[HA]}\)

The larger the \(K_a\), the stronger the acid (the more it ionises). Because \(K_a\) values are often tiny, we use \(pK_a\):

\(pK_a = -\log_{10}K_a\)

Trick: A lower \(pK_a\) means a stronger weak acid.

Calculating pH of a Weak Acid

We make two "simplifying assumptions" here:
1. \([H^+] = [A^-]\) (we assume all \(H^+\) comes from the acid).
2. \([HA]_{equilibrium} \approx [HA]_{initial}\) (we assume so little ionises that the concentration doesn't really change).

This gives us: \(K_a = \frac{[H^+]^2}{[Acid]}\), which rearranges to \([H^+] = \sqrt{K_a \times [Acid]}\).

Did you know? The pH of the ocean is currently around 8.1. As we add more \(CO_2\) from the atmosphere, it reacts with water to form a weak acid (carbonic acid), which lowers the pH. This is called ocean acidification.

5. Buffer Solutions

A buffer is a chemical "hero"—it minimises pH changes when small amounts of acid or base are added. The ocean is a massive buffer system that keeps the water slightly alkaline so shells can grow.

How they are made:

A buffer is usually a mixture of a weak acid (e.g., \(CH_3COOH\)) and its conjugate base (e.g., \(CH_3COONa\)).

How they work (The Le Chatelier way):

1. Add \(H^+\) (acid): The extra \(H^+\) reacts with the conjugate base in the buffer to form more weak acid molecules. The \(H^+\) is "mopped up."
2. Add \(OH^-\) (alkali): The \(OH^-\) reacts with the \(H^+\) in the buffer to make water. The weak acid then ionises slightly more to replace the lost \(H^+\).

Buffer Calculations

We use the \(K_a\) expression, but we cannot assume \([H^+] = [A^-]\) because we added extra base! Instead, use:
\([H^+] = K_a \times \frac{[Acid]}{[Salt]}\)

Buffer Quick Summary:

A buffer needs a high concentration of both the weak acid and its conjugate base to be able to deal with additions of either acid or alkali.

Key Takeaway for Oceans: The carbonate buffer system (\(CO_2 / HCO_3^- / CO_3^{2-}\)) is essential for marine life. If the pH drops too much, the buffer capacity is overwhelmed, and organisms like coral cannot form their calcium carbonate shells.

Summary Checklist

  • Can you identify Brønsted–Lowry acids and bases in an equation?
  • Do you know the difference between a strong acid (full ionisation) and a weak acid (partial ionisation)?
  • Can you calculate the pH of a strong acid and a strong base (using \(K_w\))?
  • Can you calculate the pH of a weak acid (using the square root formula)?
  • Can you explain how a buffer works in the context of the ocean?

Keep practicing those log calculations on your calculator—you've got this!