Unit Overview

Further chemical equilibria

CIE A-Level Chemistry· 5 min read 📊 n/a

1. Unit at a Glance

This unit follows a logical learning arc that builds from foundational acid-base concepts to applied experimental contexts. We start with core definitions and pH calculations for acids and bases, then move to buffer solutions which are essential for maintaining stable pH in biological and industrial systems. Next, we cover solubility product, a key equilibrium concept for sparingly soluble ionic compounds. Finally, we connect all prior concepts to interpreting titration curves and selecting indicators for quantitative analysis.

This unit is split into 4 connected sub-topics:

Acid-base equilibria

Covers Bronsted-Lowry theory, $K_{a}$ , $p K_{a}$ , $K_{w}$ , and pH calculations for strong/weak acids and bases.

★★★⏱ 12 min

Buffer solutions

Explains buffer action, buffer preparation, and pH calculations for acidic and basic buffers.

★★★⏱ 10 min

Solubility product

Introduces $K_{s p}$ , the common ion effect, and using $K_{s p}$ to predict precipitation reactions.

★★★★⏱ 9 min

pH titration curves and indicators

Covers interpretation of titration curves for different acid-base combinations and indicator selection rules.

★★★⏱ 10 min

2. Common Pitfalls

Wrong move:

Confusing $K_{a}$ with $K_{s p}$ or mixing up solubility and solubility product values.

Why:

These are equilibrium constants for different types of systems, so mixing them leads to incorrect calculations.

Correct move:

Always confirm the type of system you are working with and recall the definition of the equilibrium constant before starting calculations.

Wrong move:

Assuming all buffers are made from a weak acid and its conjugate salt.

Why:

Basic buffers (weak base + conjugate acid) are also commonly tested, requiring slightly different calculation steps.

Correct move:

Identify whether you are working with an acidic or basic buffer first before applying the pH formula.

Wrong move:

Forgetting to adjust concentrations for dilution when mixing solutions for pH calculations.

Why:

Dilution changes the concentration of all species, which alters the final calculated pH.

Correct move:

Always calculate new concentrations after mixing before proceeding with any equilibrium calculation.

3. Quick Reference Cheatsheet

Concept	Key Formula/Relationship
Ionic product of water	$K_{w} = [H^{+}] [O H^{-}] = 1.00 \times 1 0^{- 14}$ at 25°C
pH / pOH relationship	$p H + pO H = 14$ at 25°C
Acid dissociation constant	$K_{a} = \frac{[ H ^{+} ] [ A ^{-} ]}{[ H A ]}$ ; $p K_{a} = - lo g_{10} K_{a}$
Buffer pH (Henderson-Hasselbalch)	$p H = p K_{a} + lo g_{10} \frac{[ A ^{-} ]}{[ H A ]}$
Solubility product for $M_{a} X_{b}$	$K_{s p} = [M^{n +}]^{a} [X^{m -}]^{b}$
Indicator end point range	$p H = p K_{a} (in d i c a t or) \pm 1$
$K_{s p}$ for $A B$ -type salt	$K_{s p} = s^{2}$ , where $s$ = molar solubility

What's Next

Begin this unit by mastering the fundamentals of acid-base equilibria, which forms the foundation for all other topics in this unit. Once you complete all four sub-topics here, you can move on to the next unit on electrochemistry, which builds on many of the equilibrium concepts you will learn in this unit.