Acid-Base Equilibrium- Khan Academy Chemistry Guide
What Acid-Base Equilibrium Actually Is
Acid-base equilibrium describes what happens when acids and bases dissolve in water and interact with each other. It's not some abstract concept invented to torture students. It's a fundamental process that explains everything from how your blood maintains its pH to why baking soda makes baked goods rise.
In water, molecules constantly break apart and reform. When you add an acid to this system, it donates hydrogen ions (H⁺). When you add a base, it accepts them. The equilibrium constant tells you how far the reaction goes in either direction.
The Key Players: Ka, Kb, and pH
These three values are the backbone of every acid-base calculation. If you're mixing them up, you're going to fail every problem. Here's the rundown:
- Ka (acid dissociation constant) measures how strongly an acid donates its proton. Bigger Ka = stronger acid.
- Kb (base dissociation constant) measures how strongly a base accepts protons. Bigger Kb = stronger base.
- pH tells you the concentration of H⁺ ions in solution. Lower pH = more acidic.
The relationship between these values matters. For the conjugate acid-base pair of a weak acid and its salt:
Ka × Kb = Kw
Where Kw is the water ion product (1.0 × 10⁻¹⁴ at 25°C). This equation is non-negotiable. Memorize it.
Strong vs. Weak Acids and Bases
Strong acids like HCl, HBr, HI, HNO₃, HClO₄, and H₂SO₄ completely dissociate in water. Their equilibrium lies almost entirely to the right. You can't calculate pH from Ka for these—use the concentration directly.
Weak acids like HF, CH₃COOH (acetic acid), and H₂CO₃ only partially dissociate. Their equilibrium position matters. This is where Ka becomes essential.
Same logic applies to bases. Strong bases like NaOH and KOH fully dissociate. Weak bases like NH₃ and amines don't.
The pKa Trap
Students frequently confuse pH and pKa. Here's the difference:
- pH = -log[H⁺] — depends on actual solution concentration
- pKa = -log[Ka] — an intrinsic property of the acid itself
A lower pKa means a stronger acid. A pKa of 2 is stronger than a pKa of 5. Don't overthink this—just flip the sign and take the log.
The Henderson-Hasselbalch Equation
This equation shows up constantly. It's the shortcut for buffer calculations:
pH = pKa + log([A⁻]/[HA])
Where [A⁻] is the conjugate base concentration and [HA] is the weak acid concentration. When [A⁻] = [HA], pH = pKa. This midpoint is where buffers resist pH changes most effectively.
Buffer capacity depends on the amount of acid and base present, not their ratio. You need sufficient quantities of both to neutralize added acid or base.
Comparing Common Acids
| Acid | Type | Ka | pKa |
|---|---|---|---|
| HCl | Strong | ~10⁷ | -7 |
| HF | Weak | 7.2 × 10⁻⁴ | 3.14 |
| CH₃COOH | Weak | 1.8 × 10⁻⁵ | 4.74 |
| H₂CO₃ | Weak | 4.3 × 10⁻⁷ | 6.37 |
| NH₄⁺ | Weak | 5.6 × 10⁻¹⁰ | 9.25 |
Notice the pattern. Stronger acids have much larger Ka values, which means much smaller pKa values. The math works out cleanly.
How to Solve Acid-Base Equilibrium Problems
Most problems follow the same sequence. Here's how to approach them:
Step 1: Identify the Type
Is this a strong acid/base problem or weak? Strong acids and bases dissociate completely—set up a simple concentration calculation. Weak acids and bases require Ka or Kb.
Step 2: Write the Equilibrium Expression
For acetic acid in water:
CH₃COOH ⇌ H⁺ + CH₃COO⁻
Ka = [H⁺][CH₃COO⁻] / [CH₃COOH]
Step 3: Set Up Your ICE Table
Initial concentrations, Change, Equilibrium concentrations. This organizes your information and shows what you're solving for.
Step 4: Solve for the Unknown
For weak acids where Ka is small, you can often approximate that x (the amount dissociated) is negligible compared to the initial concentration. Check your approximation: x should be less than 5% of the initial value. If it's not, solve the quadratic.
Step 5: Calculate pH
Once you have [H⁺], pH = -log[H⁺]. Done.
Titration Curves: What You Need to Know
Titration problems test whether you understand the equilibrium shifts. The key points on any weak acid-strong base titration curve:
- Initial pH: determined by the weak acid alone
- Before equivalence point: buffer region—use Henderson-Hasselbalch
- At half-equivalence: pH = pKa (the most important landmark)
- At equivalence point: all weak acid converted to conjugate base; pH > 7
- Beyond equivalence point: excess strong base determines pH
The equivalence point for weak acid-strong base titrations isn't neutral. The conjugate base hydrolyzes, producing OH⁻ and making the solution basic.
Common Mistakes That Cost Points
- Using Ka for strong acids (they don't have meaningful Ka values)
- Forgetting that conjugate bases of weak acids are themselves weak bases
- Not checking approximation validity in ICE table problems
- Confusing pKa with pH
- Assuming equivalence point pH = 7 for weak acid-strong base titrations
Where Khan Academy Fits In
Khan Academy's chemistry content provides video walkthroughs of these concepts. The platform works best when you use it actively—pause, attempt problems, then watch the solution. Passive watching doesn't build the problem-solving instincts you need for exams.
The acid-base equilibrium unit covers Ka and Kb calculations, buffer preparation, and titration analysis. Each section includes practice problems with immediate feedback. Use them.
Focus on understanding why the equations work, not just memorizing them. Equilibrium concepts connect to thermodynamics and kinetics later in chemistry curricula. Build the foundation now.