Why does the weak acid curve start at a higher pH than the strong acid curve?

A weak acid is only partially dissociated in solution. Its initial pH is higher because the concentration of free H⁺ ions is lower than in a strong acid of the same formal concentration. The pH is calculated using the acid's Ka value and an equilibrium expression, not simply -log[HA]initial.

What exactly is the equivalence point, and how do I find it on the graph?

The equivalence point is the moment in the titration where the moles of added base are stoichiometrically equal to the moles of acid originally present. On the pH curve, it is located at the center of the steepest, nearly vertical rise. It is not necessarily at pH 7; for a weak acid-strong base titration, it occurs in the basic region due to hydrolysis of the conjugate base.

Why is there a flat 'buffer region' in the weak acid titration?

As base is first added to a weak acid, it produces the conjugate base (A⁻), creating a buffer system (HA/A⁻). Buffers resist large pH changes, resulting in a gradual slope. The pH in this region is governed by the Henderson-Hasselbalch equation, and the buffer capacity is greatest when [HA] = [A⁻], which occurs at pH = pKa.

Does this simulator show how to choose an indicator?

Indirectly. A good indicator changes color within the pH range of the vertical portion of the curve (the equivalence point region). By observing where the steep pH jump occurs, you can select an indicator whose color transition interval (e.g., phenolphthalein, ~8.2-10.0) falls entirely within that jump. The model shows why an indicator suitable for a strong acid-strong base titration might fail for a weak acid-strong base one.

Titration Simulator

Add base to acid. pH curve in real time. Find equivalence point.

Flask + indicator

1.00

Acidic

V_a = 25 mL · C_a = 0.10 M · C_b = 0.10 M · V_b = 0.00 mL

Burette

0.0 mL

Live graphs

About this model

A titration simulator visualizes the progressive neutralization of an acid by a base, plotting the resulting pH of the solution against the volume of titrant added. At its core, the model applies principles of acid-base equilibrium and stoichiometry. For a strong acid-strong base titration, the pH is calculated directly from the concentration of excess H⁺ or OH⁻ ions before and after the equivalence point. Near the equivalence point, the autoionization of water (Kw = 1.0 × 10⁻¹⁴ at 25°C) becomes significant. For weak acid-strong base titrations, the calculation involves the acid dissociation constant (Ka) and the Henderson-Hasselbalch equation (pH = pKa + log([A⁻]/[HA])) in the buffer region. The equivalence point pH is determined by the hydrolysis of the conjugate base. Key simplifications include assuming ideal solutions with constant temperature (25°C), neglecting ionic strength effects on activity, and treating all reactions as instantaneous. By interacting with this dynamic model, students learn to predict the shape of titration curves, identify the equivalence point from the steepest pH rise, understand the concept and calculation of buffer regions, and distinguish between titration curves for strong versus weak acid-base pairs.

Who it's for: Advanced high school and undergraduate chemistry students studying acid-base equilibria, buffers, and analytical titration techniques.

Key terms

Titration
Equivalence Point
pH Curve
Acid-Base Neutralization
Buffer Region
Henderson-Hasselbalch Equation
Dissociation Constant (Ka)
Stoichiometry

How it works

Add strong base from a burette into a flask of acid. The pH curve shows how pH changes with titrant volume. For strong–strong titrations the equivalence point is at pH 7. For weak acid / strong base, the equivalence region is basic because the conjugate base hydrolyzes. Use Jump to V_eq to land on the stoichiometric point.

Key equations

Strong–strong: [H⁺] = (C_a V_a − C_b V_b) / (V_a + V_b) before equivalence; after, use [OH⁻] similarly. At C_a V_a = C_b V_b → pH 7.

Weak acid buffer: pH = pK_a + log([A⁻]/[HA]). At equivalence: [OH⁻] ≈ √(K_b · [A⁻]).

Frequently asked questions

Why does the weak acid curve start at a higher pH than the strong acid curve?: A weak acid is only partially dissociated in solution. Its initial pH is higher because the concentration of free H⁺ ions is lower than in a strong acid of the same formal concentration. The pH is calculated using the acid's Ka value and an equilibrium expression, not simply -log[HA]initial.
What exactly is the equivalence point, and how do I find it on the graph?: The equivalence point is the moment in the titration where the moles of added base are stoichiometrically equal to the moles of acid originally present. On the pH curve, it is located at the center of the steepest, nearly vertical rise. It is not necessarily at pH 7; for a weak acid-strong base titration, it occurs in the basic region due to hydrolysis of the conjugate base.
Why is there a flat 'buffer region' in the weak acid titration?: As base is first added to a weak acid, it produces the conjugate base (A⁻), creating a buffer system (HA/A⁻). Buffers resist large pH changes, resulting in a gradual slope. The pH in this region is governed by the Henderson-Hasselbalch equation, and the buffer capacity is greatest when [HA] = [A⁻], which occurs at pH = pKa.
Does this simulator show how to choose an indicator?: Indirectly. A good indicator changes color within the pH range of the vertical portion of the curve (the equivalence point region). By observing where the steep pH jump occurs, you can select an indicator whose color transition interval (e.g., phenolphthalein, ~8.2-10.0) falls entirely within that jump. The model shows why an indicator suitable for a strong acid-strong base titration might fail for a weak acid-strong base one.

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About this model

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Key equations

Frequently asked questions

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Frequently asked questions