Course: General Chemistry (Purdue University) | Section: 16.3 Buffer Solutions
Source: Buffers Lab – Claim, Evidence, Reasoning Activity
Difficulty: Introductory to Intermediate Prerequisites: Basic understanding of acids, bases, pH scale, strong vs. weak acids/bases
A buffer is a solution that resists large pH changes when small amounts of strong acid or strong base are added. Buffers work because they contain both a weak acid and its conjugate base (or a weak base and its conjugate acid), which can neutralise added H⁺ or OH⁻ ions. In this lab, you identify buffers by comparing pH shifts across three solutions: the ones that barely move are the buffers.
Buffer (buffer solution)
A solution that resists significant changes in pH when small amounts of strong acid or strong base are added. It contains a weak acid and its conjugate base, or a weak base and its conjugate acid.
In simple terms, a buffer is a chemical "shock absorber" for pH. You can drip acid or base into it and the pH barely shifts.
Conjugate acid-base pair
A pair of substances related by the gain or loss of a single proton (H⁺). For example, acetic acid (CH₃COOH) and acetate (CH₃COO⁻) form a conjugate pair.
Think of it as two forms of the same molecule: one with a proton attached, one without.
pH
A logarithmic scale (0 to 14) measuring how acidic or basic a solution is. pH = -log[H⁺]. Lower values are acidic, higher values are basic, and 7.00 is neutral.
In simple terms, pH is a number that tells you how acidic or basic something is.
Strong acid
An acid that dissociates completely in water, releasing all of its H⁺ ions. HCl is a common example used in buffer experiments.
Strong base
A base that dissociates completely in water, releasing OH⁻ ions. NaOH is a common example used in buffer experiments.
Buffer capacity
The amount of strong acid or strong base a buffer can absorb before its pH begins to change substantially. Higher concentrations of the buffer components give greater capacity.
A buffer must contain two components working together: a weak acid and its conjugate base, or a weak base and its conjugate acid.
Neither component alone is a buffer. Pure water at pH 7.00 is not a buffer, even though its pH looks "neutral."
The weak acid neutralises any added OH⁻ ions. The conjugate base neutralises any added H⁺ ions. This two-way defence is what keeps the pH stable.
Take a sample of the solution and measure its initial pH.
Add a small amount of strong acid (e.g. 0.001 M HCl) and measure the pH again.
Take a fresh sample, add a small amount of strong base (e.g. 0.001 M NaOH), and measure the pH.
If the pH changes are very small in both directions, the solution is likely a buffer.
If the pH shifts dramatically in either direction, the solution is not a buffer.
Three solutions (A, B, C) were tested. Each started with 5 mL samples.
Solution A (initial pH 3.56, acidic buffer)
After adding HCl: pH dropped to 3.48 (change of 0.08)
After adding NaOH: pH rose to 3.64 (change of 0.08)
Both shifts are tiny. Solution A behaves as a buffer.
Solution B (initial pH 9.98, basic buffer)
After adding HCl: pH dropped to 9.90 (change of 0.08)
After adding NaOH: pH rose to 10.08 (change of 0.10)
Both shifts are small. Solution B behaves as a buffer.
Solution C (initial pH 7.00, not a buffer)
After adding HCl: pH plummeted to 1.63 (change of 5.37)
After adding NaOH: pH jumped to 12.30 (change of 5.30)
Massive shifts in both directions. Solution C is not a buffer. It is likely pure water or a dilute salt solution with no buffering components.
A neutral pH does not make something a buffer. Solution C had no weak acid/conjugate base pair to absorb the added H⁺ or OH⁻. Without those components, even a tiny amount of strong acid or base overwhelms the solution and the pH swings wildly.
Henderson-Hasselbalch equation (for acidic buffers)
pH = pKa + log([A⁻] / [HA])
Where [A⁻] is the concentration of the conjugate base and [HA] is the concentration of the weak acid. This equation lets you calculate the pH of a buffer if you know the pKa and the ratio of base to acid.
Henderson-Hasselbalch equation (for basic buffers)
pOH = pKb + log([BH⁺] / [B])
Where [BH⁺] is the concentration of the conjugate acid and [B] is the concentration of the weak base.
pH change calculation (used in this lab)
ΔpH = pH(after addition) - pH(initial)
A small ΔpH (close to zero) after adding acid or base indicates buffering behaviour. A large ΔpH indicates no buffering.
Buffers are essential in biological systems. Human blood is buffered at around pH 7.4 by the carbonic acid/bicarbonate system (H₂CO₃/HCO₃⁻). Even small deviations from this range can be life-threatening, which is why the body maintains multiple buffering systems.
In industry, buffers are used in pharmaceutical manufacturing to keep drug formulations at a stable pH, in fermentation processes where microorganisms require tightly controlled pH to function, and in water treatment to prevent corrosion in pipes.
"A solution at pH 7 is always a buffer." It is not. pH 7 just means neutral. Pure water sits at pH 7 and has zero buffering capacity. A buffer needs a conjugate acid-base pair, regardless of where the pH lands.
"Buffers prevent all pH change." They do not. Buffers resist change, meaning the shift is small. They still shift, just far less than an unbuffered solution would. Once the buffer capacity is exhausted, the pH will swing freely.
"A buffer only works against acid, not base (or vice versa)." A proper buffer resists pH change in both directions. It must handle added H⁺ and added OH⁻. If a solution only resists one direction, it is not functioning as a true buffer.
"Stronger concentration of acid or base means a better buffer." Buffers are made from weak acids and weak bases, not strong ones. A solution of strong acid is not a buffer, no matter how concentrated.
⚠️ You will almost certainly be asked to define a buffer. The definition must include: (1) resists pH change, (2) when small amounts of strong acid or base are added, (3) contains a weak acid and its conjugate base or a weak base and its conjugate acid.
⚠️ Expect data interpretation questions like the one in this lab. You will be given pH values before and after acid/base addition and asked to identify which solution is a buffer. Calculate ΔpH for each and compare.
⚠️ Know the Henderson-Hasselbalch equation and be able to use it to calculate the pH of a buffer given pKa and concentrations.
⚠️ Be ready to explain why a solution at pH 7 (like pure water) is not a buffer. This is a common exam trap.
⚠️ Understand that buffers work in both directions. If a question asks you to predict what happens when NaOH is added to a buffer, the weak acid component reacts with the OH⁻, converting some HA to A⁻, and the pH rises only slightly.
True or False: A buffer contains a strong acid and a strong base. (False, it contains a weak acid and its conjugate base, or a weak base and its conjugate acid.)
True or False: A solution with a pH of 7.00 that shifts to pH 1.63 when a few drops of HCl are added is a buffer. (False, the large pH shift shows no buffering capacity.)
Fill in the blank: The Henderson-Hasselbalch equation for an acidic buffer is pH = pKa + log( ______ / ______ ). ([A⁻] / [HA])
True or False: A buffer resists pH change only when acid is added, not when base is added. (False, a buffer resists change in both directions.)
Fill in the blank: In this lab, Solutions A and B showed pH changes of roughly ______ units after acid addition, while Solution C showed a change of roughly ______ units. (0.08; 5.37)
Q: Define a buffer solution and state the two essential components it must contain.
A: A buffer solution is one that resists significant changes in pH when small amounts of strong acid or strong base are added. It must contain either a weak acid and its conjugate base, or a weak base and its conjugate acid.
Q: Given the following data, identify which solution is a buffer and explain your reasoning. Solution X has an initial pH of 4.75. After adding HCl, the pH is 4.68. After adding NaOH, the pH is 4.82.
A: Solution X is a buffer. The pH changed by only 0.07 units with acid and 0.07 units with base, indicating the solution resists pH change in both directions. This small, symmetrical shift is characteristic of a buffered solution containing a weak acid/conjugate base pair.
Q: A student claims that pure water (pH 7.00) is a buffer because it is neutral. Explain why this claim is incorrect.
A: Pure water is not a buffer. Although its pH is 7.00, it contains no weak acid/conjugate base pair to neutralise added H⁺ or OH⁻. Adding even a small amount of strong acid or base to water causes a dramatic pH shift, as demonstrated by Solution C in this experiment (pH dropped from 7.00 to 1.63 with HCl).
Q: Using the Henderson-Hasselbalch equation, calculate the pH of a buffer made from a weak acid with pKa = 4.75 when the concentrations of weak acid and conjugate base are equal.
A: When [A⁻] = [HA], the ratio [A⁻]/[HA] = 1, and log(1) = 0. So pH = pKa + 0 = 4.75. When the acid and conjugate base are present in equal concentrations, the buffer pH equals the pKa.
Q: Explain what happens at the molecular level when a small amount of NaOH is added to an acidic buffer (HA/A⁻ system).
A: The OH⁻ ions from NaOH react with the weak acid (HA) in the buffer: HA + OH⁻ → A⁻ + H₂O. This converts some of the weak acid into its conjugate base. Because the OH⁻ is consumed by this reaction rather than accumulating in solution, the pH rises only slightly.
Buffers connect directly to acid-base equilibrium (Ch. 16). Understanding Ka, Kb, and equilibrium expressions is prerequisite to understanding why a buffer holds its pH. If you are shaky on writing equilibrium expressions or calculating Ka from concentration data, revisit Ch. 15 first.
This topic also links to titration curves (Ch. 16.4). The flat region of a titration curve, where pH changes slowly despite adding titrant, is the buffer region. Recognising it on a graph is a common exam question.
In biology and biochemistry, buffer chemistry underpins how enzymes function (most enzymes have a narrow pH optimum) and how the body regulates blood pH through the bicarbonate buffer system.
Buffer solution, buffer system, Henderson-Hasselbalch equation, conjugate acid-base pair, weak acid, conjugate base, weak base, conjugate acid, pH, pKa, pKb, buffer capacity, buffer range, acid-base equilibrium, titration curve, buffer region, carbonic acid bicarbonate buffer, blood pH, resist pH change, neutralisation, HCl, NaOH, strong acid, strong base, pH stability, Le Chatelier's principle and buffers, common ion effect