Strong Acid Strong Base Titration Curve Explained

What Is a Strong Acid Strong Base Titration Curve?

A strong acid strong base titration curve is a graph that shows how the pH of a solution changes as you slowly add a strong base to a strong acid (or vice versa). The curve has a distinctive S-shape with a steep jump at the equivalence point.

You plot volume of titrant on the x-axis and pH on the y-axis. The result tells you exactly where the neutralization reaction completes.

The Chemistry Behind It

Strong acids (like HCl, HBr, HI, HNO₃, HClO₄) and strong bases (like NaOH, KOH, LiOH, Ca(OH)₂) dissociate completely in water. This means every molecule donates or accepts a proton. There's no equilibrium to worry about.

The net reaction is:

H⁺ + OH⁻ → H₂O

That's it. Water and salt. No weak acid or weak base complications. This simplicity is why the titration curve looks so clean and predictable.

Key Regions of the Curve

1. Initial Point (Before Any Titrant Added)

The pH starts low for a strong acid. For 0.1 M HCl, pH = 1. The solution is just acid in water. Adding a strong base will start neutralizing those H⁺ ions immediately.

2. Before the Equivalence Point

You've added some base, but not enough to neutralize all the acid. The solution still contains excess H⁺. The pH rises slowly at first, then faster as you approach the equivalence point.

Calculate pH here using the excess moles of H⁺ or OH⁻ remaining after neutralization.

3. The Equivalence Point

This is where the moles of acid equal the moles of base added. For a strong acid-strong base titration, the pH is exactly 7.00 at 25°C. The solution is neutral.

The equivalence point is marked by the steepest part of the curve—a near-vertical line on the graph. One or two drops of titrant can change the pH by several units.

4. After the Equivalence Point

You've added more base than acid. Excess OH⁻ ions float around. The pH is now above 7 and climbs as you add more base.

Calculate pH here using the excess moles of OH⁻.

Calculating pH at Different Points

Here's the straightforward method:

Before Equivalence Point:

• Find moles of acid originally present
• Find moles of base added
• Subtract: moles H⁺ remaining = moles acid - moles base added
• Divide by total volume to get [H⁺]
• pH = -log[H⁺]

After Equivalence Point:

• Find moles of base added
• Find moles of acid originally present
• Subtract: moles OH⁻ excess = moles base added - moles acid
• Divide by total volume to get [OH⁻]
• pOH = -log[OH⁻]
• pH = 14 - pOH

At Equivalence Point:

pH = 7.00 (for strong acid + strong base at 25°C)

Practical Example: Titration of HCl with NaOH

Problem: 50.0 mL of 0.100 M HCl is titrated with 0.100 M NaOH. Calculate the pH after adding 0, 25, 50, 75, and 100 mL of NaOH.

Step 1: Initial (0 mL NaOH added)

[H⁺] = 0.100 M
pH = -log(0.100) = 1.00

Step 2: 25 mL NaOH added

Moles HCl = 0.0500 L × 0.100 M = 0.00500 mol
Moles NaOH = 0.0250 L × 0.100 M = 0.00250 mol
Excess H⁺ = 0.00500 - 0.00250 = 0.00250 mol
Total volume = 0.0750 L
[H⁺] = 0.00250 / 0.0750 = 0.0333 M
pH = -log(0.0333) = 1.48

Step 3: 50 mL NaOH added (Equivalence Point)

Moles NaOH = 0.0500 L × 0.100 M = 0.00500 mol
This equals moles HCl (0.00500 mol).
pH = 7.00

Step 4: 75 mL NaOH added

Moles NaOH = 0.0750 L × 0.100 M = 0.00750 mol
Excess OH⁻ = 0.00750 - 0.00500 = 0.00250 mol
Total volume = 0.125 L
[OH⁻] = 0.00250 / 0.125 = 0.0200 M
pOH = -log(0.0200) = 1.70
pH = 14 - 1.70 = 12.30

Step 5: 100 mL NaOH added

Moles NaOH = 0.100 L × 0.100 M = 0.0100 mol
Excess OH⁻ = 0.0100 - 0.00500 = 0.00500 mol
Total volume = 0.150 L
[OH⁻] = 0.00500 / 0.150 = 0.0333 M
pOH = -log(0.0333) = 1.48
pH = 14 - 1.48 = 12.52

How to Read a Strong Acid Strong Base Titration Curve

Here's what to look for when you see one:

• Flat starting region: Buffer region doesn't exist here. The pH changes immediately because strong acids and bases don't form buffer systems.
• Steep vertical section: This is your equivalence point zone. The pH jumps from acidic to basic over a tiny volume change.
• Flat ending region: Adding more base barely changes pH. The solution is already strongly basic.
• Midpoint: The pH at half the equivalence volume equals the pKa or pKb of the conjugate species. For strong acid-strong base, this happens at pH 7.

Strong Acid Strong Base vs. Weak Acid Strong Base

Common Mistakes Students Make

• Assuming pH = 7 at equivalence for weak acid titrations. This is only true for strong acid-strong base.
• Forgetting to account for total volume when calculating concentrations after mixing.
• Using the wrong formula before vs. after equivalence. Excess H⁺ before, excess OH⁻ after.
• Skipping the Henderson-Hasselbalch equation for weak acid titrations while applying it to strong acid titrations where it doesn't work.

The Indicator Cheat Sheet

You need an indicator that changes color within the steep portion of the curve. For strong acid-strong base titrations, these work best:

Phenolphthalein is the most common choice. It goes from colorless to pink right at the equivalence point.

Quick Reference: pH at Key Points

Bottom Line

The strong acid strong base titration curve is simple because the chemistry is simple. Complete dissociation means no equilibrium complications. Calculate excess H⁺ or OH⁻, divide by total volume, take the -log.

The equivalence point is always pH 7 for this combination. The curve jumps steeply through that point. Pick your indicator based on which direction you're titrating.