Chem 221 Experiment 01 pH and Buffers Version 2

Background:  Water, the most common solvent for biochemical experiments, has an equal concentration of H+ (more properly expressed as H3O+) ions and OH ions, arising from the autoionization of water:

H2O(l)  +  H2O(l)  ↔  H3O+(aq)  + OH (aq);  Kw = 1.0 x 1014 at 25°C.

Addition of H+ from an external source suppresses the ionization of water, reducing the OH concentration and creating a solution where [H+] > [OH]; an acidic solution.  Similarly, addition of OH creates a solution where [OH] > [H+]; a basic solution.  The pH scale represents [H+] by a logarithmic quantity: pH = – log [H+].

Weak acids dissociate slightly in aqueous solution, creating a situation with relatively high concentration of undissociated acid mixed with small concentrations of H+ and the conjugate base of the acid.  The final pH of a solution of weak monoprotic acid (only) depends on the acid dissociation constant of the acid (Ka or its related pKa), the concentration of the acid and the temperature of the solution.  (Remember these equilibrium problems from Chem 102?)  The use of a polyprotic acid (an acid with more than one ionizable hydrogen) creates another complication.  Each of the ionizations has its own dissociation constant.  However, in general, one need only worry about a single dissociation at a time within a given pH range near one species’ pKa.  One can safely assume that the ionization occurs stepwise with each step going essentially to completion in sequence.  If you initially dissolve NaH2PO4 in water, you will deal with the equilibria:

H2PO4‾(aq)  ↔  H+(aq)  +  HPO42‾ (aq)

or alternatively:

H2PO4‾(aq)  +  H2O ↔ H3PO4(aq) + OH‾ (aq)

The pH of this solution will be controlled solely by the two acidic species’ pKa values and is independent of concentration.  pH = (pKa1 + pKa2) / 2.

A solution simultaneously containing appreciable quantities of both a weak acid (HA) and its conjugate base (A‾ ) is a buffer.  One of the buffer components, the weak acid HA, can dissociate H+ if mixed with base and the other buffer component, A‾, can accept H+ if mixed with acid.  Either event alters the ratio of the buffer components but only slightly changes the solution pH, which is dependent on the log of the ratio of the buffer components.   The Henderson-Hasselbalch equation can be used to calculate pH when you have a buffer:

pH = pKa + log ([A‾ ]/[HA])

A buffer effectively controls pH within a range of ±1 pH unit of its pK­a (where the mixture ranges from 10HA per A‾ up to 10 A‾ per HA) and a buffer has a buffering capacity determined by the available concentration of the buffer components.

The original buffers used in the lab were made from familiar weak acids mixed with their salts, like an acetic acid/sodium acetate mixture or a monobasic phosphate/dibasic phosphate mixture.  However, these had some problems.  They changed pH too much when diluted or if the temperature changed.  They often permeated the cell, changing its internal concentrations, or participated in biochemical reactions.  Contemporary synthetic buffers (Good’s buffers) are zwitterions, highly soluble in aqueous systems, are chemically stable, do not permeate the cell, are resistant to concentration and temperature changes, have insignificant light absorption in the ultraviolet and visible regions (so they don’t interfere with spectroscopic analysis of biomolecules) and have a variety of pKa values within the range (pH 6 – 8) normally employed in biochemistry.  Most can be purchased in either the acid form or the basic form to facilitate achieving the target solution pH.  “Tris” [tris(hydroxymethyl)aminomethane] is probably the most commonly used.  A table of structures and properties of several Good’s buffers is given.

Tris is useful in the range of pHs from 7.5 to 9.0.  It is available in basic form as highly purified crystals with molar mass 121.1 g/mol.  To prepare 1.0 L of 0.10 M buffer, place 0.10 mole of Tris base in 950-975 mL of water in a volumetric flask, add acid to achieve the desired pH then add water to a final volume of 1.0 L and recheck the pH (and readjust slightly if needed).  Tris has several disadvantages, including (1) pH dependence on concentration; pH decreases 0.1 pH unit for each 10-fold dilution, (2) interference with some pH electrodes, and (3) a large ΔpKa/˚C compared to most other buffers.  These drawbacks can be minimized by (1) adjusting the pH after dilution to the appropriate concentration and (2) preparing the buffer at the temperature at which it will be used.


After this experiment, you will be able to

  • Calculate the pH of solutions of strong acids or bases, weak acids or bases, buffers, and combinations of these.
  • Explain how a buffer resists pH change.
  • Prepare an appropriate buffer for a given pH.
  • Predict the effect of changing temperature and concentration on buffer pH.
  • Standardize and operate a pH meter.


  • Glacial acetic acid, HC2H3O2 at 17.416 M and/or other concentrated acetic acid
  • Solid NaC2H3O2, anhydrous
  • Solid NaH2PO4 (monobasic phosphate), trihydrate
  • Solid Na2HPO4 (dibasic phosphate), anhydrous & heptahydrate
  • Solid Tris (base form) or solid Tris HCl (acid form)
  • 0.1 & 0.5 M NaOH(aq)
  • 0.1 & 0.5 M HCl(aq)
  • 3 vials of standard buffer solutions for pH meter calibration, pH 4.00, 7.00, and 10.00. Keep these in your drawer for the rest of the semester.
  • Accumet pH meter (Accumet-AB15-PH-Meter User Manual)
  • or Venier Labquest + pH probe (cannot be used for Tris buffer) (Calibration Instructions)


You will each prepare 3 buffer solutions at a given pH target. For example, acetate buffer at pH 5.20, phosphate buffer at pH 7.00, and Tris buffer at pH 8.50.  You will make 100.0 mL of each buffer, with a total combine buffer component concentration of 0.100 M solution. Your instructor will give you a card with this information.

You must provide your instructor with the calculations and instructions on how to make your buffers before proceeding. 

Preparation of the Buffers

To prepare the acetate and phosphate buffers, calculate the proportions of buffer components needed to achieve your target pH, the actual masses and/or volume needed, and then prepare the mixture.

Set up and calibrate the pH meter.  Measure the pH reading of your final buffer.  If it does not have a value close to your target, what might have gone wrong and how could you correct the pH value without substantially changing the concentration?

Make 100.00 mL of a 0.100 M buffer starting with Tris alone.  Calculate the mass of Tris needed for the solution, make the solution initially in 50 mL of water or less in a beaker over a stirplate. Insert the pH meter and add 0.10 M HCl or 0.10 M NaOH until you achieve the target pH. Transfer the solution to a volumetric flask and complete the dilution.  Recheck the pH to make sure it has not changed. If it has, correct it. Warning: you may want to use a lesser concentration of HCl or NaOH for this. Record the final temperature of the buffer.

I. Effect of concentration on pH (prepare a data table for this part).

  1. Take 10.0 mL of one of the three buffers and dilute 10-fold with deionized water. What is the new concentration of the buffer? Save it in any of the containers in your drawer.
  2. Take 10.0 mL of the diluted buffer from step 1 and dilute 10-fold once again. What is the new concentration of the buffer? Save this solution.
  3. Measure the pHs of the diluted solutions and compare them to the value for the undiluted solution.
  4. Repeat steps 1-3 for the remaining two buffers. Once you have measured the pHs of the diluted buffer solutions, you can discard them. Save the original buffers.

II. Demonstration of buffering action

  1. Put 50.0 mL of one of your original buffers in a beaker.  If your buffer has a pH higher than its pKa, add 0.50 mL of 0.10 M HCl to it.  Record the new pH.  If your buffer has a pH lower than its pKa, add 0.50 mL of 0.10 M NaOH to it.  Record the new pH.  Add additional portions of titrant until you have exceeded the buffering capacity of your buffer.  Record the volume of the titrant that you have added and note the pH.  What could be done to increase the buffering capacity without changing initial pH?
  2. Repeat step 1 for the other 2 buffers.
  3. Repeat step 1 but use 50 mL of water instead of your buffer and add either acid or base, as you did in step 1.  Record the pH values.

Pre-Lab Questions

For the acetate buffer, calculate the mass of sodium acetate and volume of

Calculations and Questions

  1. For each buffer, what is the ratio of Aˉ to HA at the target pH?  How many moles of Aˉ and HA are present in the final solution?
  2. If you add 3.0 mL of 0.10 M NaOH to each of your buffers, what is the new pH of each solution. Will they still be effective buffers?  Support your answer with a calculation.
  3. What is the effect of dilution on each buffer?  What is the origin of this effect?

Follow up Questions

  1. Calculate the theoretical pH of your Tris buffer at 0˚C.
  2. What is the most efficient way to make a HEPES buffer at pH 8.5?  You have HEPES in both zwitterionic form and in anionic form plus all common laboratory reagents.  The pKa of HEPES is 7.55 at 20˚C.
  3. If you mix 50. mL of 0.10 M solution of Tris acid with 60. mL of 0.20 M solution of Tris base, what will be the resulting pH?
  4. An enzyme-catalyzed reaction is carried out in a 50. mL solution containing 0.10 M Tris buffer.  The initial buffer pH was 8.0.  As a result of the enzymatic reaction, 0.00020 mole of H+ was produced in the solution.  What is the ratio of Tris base to Tris acid at the start of the experiment?  What is the ratio at the end?  What is the final pH of the solution?