A buffered vs. a non-buffered solution

Students often struggle with the difference between a solution that is buffered and one that is not. This demonstration clearly distinguishes between the two, both qualitatively and quantitatively. It is simple to prepare and perform, and uses common materials that can be stored for future use.

A mixture of water, methyl orange and bromothymol blue — a non-buffered solution — is divided into two parts and placed on an overhead projector or a document camera. Dilute HCl (1.0 M) is added to one of the mixtures, a drop at a time, until a color change occurs — usually 1 to 2 drops — signifying a significant drop in pH. NaOH solution (1.0 M) is added dropwise to the other mixture until a color change occurs — again usually 1 or 2 drops — signifying a significant rise in pH.

A second mixture of water, methyl orange and bromothymol blue is prepared. This time, 50 mL of acetic acid (6.0 M) and 10.0 g of solid sodium acetate trihydrate (CH3COONa.3H2O) is added to the mixture forming the buffered solution. Again, the mixture is divided into two portions. The HCl and NaOH are added to separate portions until a color change occurs. It will require large quantities of acid or base to obtain a color change.

In the non-buffered solution, the added hydronium or hydroxide ions have nothing to react with so the concentrations increase rapidly, changing the pH significantly.

However, in the buffered solution, the acetate ions react with the hydronium ions from the HCl to neutralize the added acid. Similarly, hydroxide ions from the strong base, NaOH, are neutralized by the acetic acid, forming water and the acetate ion. The equations are:

 H3O+(aq) + CH3COO-(aq)    →    CH3COOH(aq) + H2O(l)

(Strong acid is completely converted into weak acid, K = 1/Ka ~ 105)

OH- (aq) + CH3COOH(aq)    →    H2O(l) + CH3COO-(aq)

(Strong base is completely converted into weak base, K = 1/Kb ~ 109)

In both cases, the equilibrium lies far to the right, thus the concentration of the hydronium or hydroxide ions is reduced, preventing the pH from changing significantly. It is important to note that the pH does change but not significantly. This can be shown via calculations — shown later in this article — or by the use of a pH meter.

four images showing buffered and non buffered solutions

Chemical concepts

  • Buffered vs. non-buffered solutions
  • Strong vs. weak acids and bases
  • pH, Ka and Kb, pKa and pKb
  • Calculation of pH for buffered solutions
  • Calculation of pH using the Henderson-Hasselbalch equation
  • Calculation of buffer range and buffer capacity


  • Deionized (DI) water ~ 750 mL
  • Bromothymol blue indicator solution
  • Methyl orange indicator solution
  • Two 600 mL beakers
  • 10.0 g sodium acetate trihydrate
  • Glacial acetic acid to prepare 6.0 M acetic acid
  • Two disposable pipets for indicator solutions
  • Six glass stirring rods
  • Four glass containers for mixtures; we use Pyrex® No. 3140 100 x 50 crystallization dishes
  • Dropper bottle filled with 1.0 M HCl
  • Dropper bottle filled with 1.0 M NaOH
  • Overhead projector or document camera — cover the latter with clear wrap to protect the surface


Wear protective safety glasses/goggles and gloves.

Hydrochloric acid and sodium hydroxide solutions are corrosive; avoid contact or ingestion.

Bromothymol blue and methyl orange solutions are mildly toxic and will stain; avoid contact or ingestion.

Advance preparation

6.0 M acetic acid may be prepared from glacial acetic acid in a fume hood. Add 37.4 mL of glacial acetic acid to every 62.6 mL of DI water and mix thoroughly. The resulting mixture can be divided into 50 mL portions and placed in designated bottles.

Measure 10.0 g portions of sodium acetate trihydrate and place in designated bottles.

Fill designated dropper bottles with 1 M HCl and NaOH.

Fill a 1 L designated bottle with DI water.


Part 1: Non-buffered solution

1.  Place ~400 mL of DI water into one of the clean beakers.

2.  Add methyl orange and bromothymol blue solutions to the water — should require ~5 mL of each — and stir.

3.  Divide the mixture into two crystallization dishes.

Set dishes on an overhead projector. Turn on the projector.

5.  Add 1.0 M HCl to one of the two dishes dropwise until the mixture changes color — should only require 1 or 2 drops. Use one of the stirring rods to mix after each drop.

6.  To the other dish add 1.0 M NaOH solution dropwise until the mixture changes color — again, should only require 1 or 2 drops. Use a clean stirring rod to mix after each drop.

Part 2: Buffered solution

1.  Place ~350 mL of DI water in the other clean beaker.

2.  Add methyl orange and bromothymol blue solutions to the water – about the same amount as for the non-buffered solution.

3.  Add 50 mL of 6.0 M acetic acid to the mixture and stir — should turn orange indicating the solution is acidic.

4.  Add the 10.0 g of solid sodium acetate trihydrate. Stir until dissolved.

5.  Repeat steps 3 through 6 described for the non-buffered solution. It should require a great deal more of the 1.0 M HCl and 1.0 M NaOH to produce a color change than needed for the non-buffered mixture.


Add sodium bicarbonate to the solutions to neutralize the mixtures. The mixtures can then be disposed of or poured into the sink. Check with your local community regulations about disposing of the dyes before doing so.


Calculation of initial pH of the buffered solution:

initial pH of a buffered solution calculation set up

“x” can be considered to be very small due to the presence of the acetate ions that suppress the ionization of acetic acid, which is a weak acid — little of it is ionized.

x = 7.3 x 1O-5 M = [H3O+]

pH = 4.13

Alternate calculations

Henderson-Hasselbalch Equation

pH of buffer = pKa + log(#moles of A-/#moles of HA)

pKa acetic acid = -log(Ka) acetic acid = -log(1.8 x 10-5) = 4.74

pH of buffer = 4.74 + log(0.0735/0.30) = 4.13

Buffer range

The buffer range is defined as the range in pH that the buffer will be effective as a buffer; +/- 1 pH unit of the pKa. This is when the ratio of moles of A-/HA is between 1/10 and 10/1. The pKa of acetic acid is 4.74 so the buffer range would be between 3.74 and 5.74. For the prepared buffer in this instance the ratio is 0.0735/0.30 or 2.45/10, within the desired range.

The capacity depends on the overall number of moles of acid (HA) and base (A-) present in the mixture capable of converting strong acid or base into weak acid or base without going outside the buffer range.

In this buffer solution there were 0.30 moles of acetic acid and 0.0735 moles of acetate ions. Clearly, this buffer solution is more resistant to additions of strong base than of strong acids.


Challenge your students to prepare other buffered solutions — different initial pH values, different buffer ranges or capacities, different materials other than sodium acetate/acetic acid, etc. Divide your class into teams and give each team a different challenge. Once completed, have each team present their challenge, the procedures attempted — successes and failures, and their final results to the other teams as if in a research group meeting.

Dr. Kenneth Lyle is a Lecturing Fellow at Duke University. The Powell Family Trust funds the Duke Chemistry Program.