Buffer Solutions

Submitted by ChemPRIME Staff on Thu, 12/16/2010 - 14:54


So far in discussing pHA logarithmic measure of the concentration of hydrogen (hydronium) ion; pH = -log10([H+]) or pH = -log10([H3O+]). we have dealt only with solutions obtained by adding a single acidIn Arrhenius theory, a substance that produces hydrogen ions (hydronium ions) in aqueous solution. In Bronsted-Lowry theory, a hydrogen-ion (proton) donor. In Lewis theory, a species that accepts a pair of electrons to form a covalent bond., such as acetic acid, or a single baseIn Arrhenius theory, a substance that increases the concentration of hydroxide ions in an aqueous solution. In Bronsted-Lowry theory, a hydrogen-ion (proton) acceptor. In Lewis theory, a species that donates a pair of electrons to form a covalent bond., such as the acetate ion, to water. We must now turn to a consideration of solutions to which both an acid and a base have been added. The simplest case of such a solution occurs when the acid and base are conjugate to each other and also present in comparable amounts. Solutions of this special kind are called buffer solutions because, as we shall shortly see, it is difficult to change their pH even when an appreciable amount of strong acidAn acid that ionizes completely in a particular solvent. or strong baseA base that dissociates completely or ionizes completely in a particular solvent. is added.

As a typical example of a bufferA solution to which strong acid or strong base can be added without significant change in pH; a solution containing a weak acid and its conjugate base or a weak base and its conjugate acid. solution, let us consider the solution obtained when 3.00 mol acetic acid (HC2H3O2) and 2.00 mol sodium acetate (Na C2H3O2) are added to sufficient water to produce a solution of total volume 1 dm³. The stoichiometric concentrationA measure of the ratio of the quantity of a substance to the quantity of solvent, solution, or ore. Also, the process of making something more concentrated. of acetic acid, namely, ca, is then 3.00 mol dm–3, while the stoichiometric concentration of sodium acetate, cb, is 2.00 mol dm–3. As a result of mixing the two components, some of the acetic acid, say x mol dm–3, is converted to acetate ion and hydronium ion. We can now draw up a table in order to find the equilibriumA state in which no net change is occurring, that is, in which the concentrations of reactants and products remain constant; chemical equilibrium is characterized by forward and reverse reactions occurring at the same rate. concentrations in the usual way.


Species Initial Concentration

mol dm-3

Change in Concentration

mol dm-3

Equilibrium Concentration

mol dm-3

  H3O+       10-7 (negligible) x x
  CH3COO-       2.00 x 2.00 + x
  CH2COOH       3.00 (-x) 3.00 - x



We can now substitute concentrations in the equilibrium expression


K_{a}=\frac{[\text{CH}_{3}\text{COO}^{-}][\text{H}_{3}\text{O}^{+}]}{[\text{CH}_{3}\text{COOH}]}


from which we obtain


\text{1.8}\times \text{10}^{-5}\text{ mol dm}^{-3}=\frac{\text{(2}\text{.00 + }x\text{)}x}{\text{3.00}-x}\text{ mol dm}^{-3}...(1)


In order to solve this equation, we make the approximation that x is negligibly small compared with both 2.00 and 3.00, that is, that only a minute fraction of acetic acid has converted to acetate ion. We then have


         \frac{\text{2.00}x}{\text{3.00}}=\text{1.8}\times \text{10}^{-5}


   or            \begin{align}x&=\frac{\text{3.00}}{\text{2.00}}\times \text{ 1.8 }\times \text{ 10}^{-5}\
\text{ }&=\text{2.7}\times \text{10}^{-5}\end{align}


Obviously, our approximation is a very good one. Since x is only 0.001 percent of 2.00 or 3.00, there really is no point in obtaining a second approximation by feeding x back into Eq. (1). We can thus safely conclude that


[\text{H}_{3}\text{O}^{+}] = \text{2.7} \times \text{10}^{-5}\text{mol dm}^{-3}       and      \text{pH} = \text{4.57}\,


The example we have just considered demonstrates two obvious features:


1 When the acid and its conjugate baseThe base formed when an acid releases a hydrogen ion (proton). are mixed, very little of the acid is converted to base, or vice versa. (x was small compared with 2.00 and 3.00.)

2 In a buffer mixtureA combination of two or more substances in which the substances retain their chemical identity., the hydronium-ion concentration and the hydroxide-ion concentration are small compared with the concentrations of acid and conjugate base. ([H3O+] = 2.7 × 10–5 mol dm–3; [HO] = 3.7 × 10–10 mol dm–3 as compared with [CH3COO] = 2.00 mol dm–3 and [CH3COOH] = 3.00 mol dm–3)


By assuming that these features are common to all buffer solutions, we make it very easy to handle them from a mathematical standpoint.

Let us now consider the general problem of finding the pH of a buffer solution which is a mixture of a weak acidAn acid that ionizes only partially in a given solvent. HA, of stoichiometric concentration ca, and its conjugate base A, of stoichiometric concentration cb. We can rearrange the expression for Ka of the weak acid (Equation 2 on the pH of solutions of weak acids) as follows:


             [\text{H}_{3}\text{O}^{+}]=K_{a}\times \frac{[\text{HA}]}{[\text{A}^{-}]}             (2)


Taking negative logarithms of both sides, we obtain


-\text{log }[\text{H}_{3}\text{O}^{+}]=-\text{log }K_{a}-\text{log}\frac{[\text{HA}]}{[\text{A}^{-}]}


                    \text{pH}=\text{p}K_{a}\text{+ log}\frac{[\text{A}^{-}]}{[\text{HA}]}       (3)


Equation (3) is called the Henderson-Hasselbalch equation and is often used by chemists and biologists to calculate the pH of a buffer.

As we saw in the case of the acetic acid―sodium acetate buffer described earlier, the equilibrium concentrations of HA and A are usually almost identical to the stoichiometric concentrations. That is,


[\text{HA}] \approx \text{c}_{a}      and       [\text{A}^{-}]\approx\text{c}_{b}


We can substitute these values into Eqs. (2) and (3) to obtain the very useful approximations


           [\text{H}_{3}\text{O}^{+}]\approx K_{a}\times \frac{c_{a}}{c_{b}}      (4)


   and           \text{pH}\approx \text{ p}K_{a}\text{ + log}\frac{c_{b}}{c_{a}}      (5)


EXAMPLE Find the pH of the solution obtained when 1.00 mol NH3 and 0.40 mol NH4Cl are mixed to give 1 dm3 of solution. Kb(NH3) = 1.8 × 10–5 mol dm–3.


Solution In order to use Eq. (4),we need first to have the value of


\begin{align}K_{a}\left(\text{NH}_{4}^{+}\right)&=\frac{K_{w}}{K_{b}\left(\text{NH}_{3}\right)}\
\text{ }&=\frac{\text{1.00}\times \text{ 10}^{-14}\text{ mol}^{2}\text{ dm}^{6}}{\text{1.8 }\times \text{ 10}^{-5}\text{ mol dm}^{-3}}\
\text{ }&=\text{5.56}\times \text{ 10}^{-10}\text{ mol dm}^{-3}\end{align}


We also have ca = 0.40 mol dm–3 and cb = 1.00 mol dm–3. Thus


\begin{align}\left[\text{H}_{3}\text{O}^{+}\right]&=K_{a}\times \frac{c_{a}}{c_{b}}\
\text{ }&=\text{5.56}\times \text{ 10}^{-10}\text{ mol dm}^{3}\times \frac{\text{  0.4 mol dm}^{-3}}{\text{1.0 mol dm}^{-3}}\
\text{ }&=\text{2.22 }\times \text{ 10}^{-10}\text{ mol dm}^{-3}\end{align}

from which


 and   \text{pH} = \text{9.65}\,



To see why a mixture of an acid and its conjugate base is resistant to a change in pH, let us go back to our first example: a mixture of acetic acid (3 mol dm–3)and sodium acetate (2 mol dm–3). What would happen if we now added 0.50 mol sodium hydroxide to 1 dm3 of this mixture? The added hydroxide ion will attack both the acids present, namely, the hydronium ion and acetic acid. Since the hydronium-ion concentration is so small, very little hydroxide ion will be consumed by reaction with the hydronium ion. Most will be consumed by reaction with acetic acid. Further, since the hydroxide ion is such a strong base, the reaction


\text{CH}_{3}\text{COOH}+ \text{OH}^{-}  \rightarrow  \text{CH}_{3}\text{COO}^{-} + \text{H}_{2}\text{O}


will go virtually to completion, and 0.50 mol acetic acid will be consumed. The same amount of acetate ion will be produced. In tabular form:


Species Initial Concentration

mol dm-3

Change in Concentration

mol dm-3

Equilibrium Concentration

mol dm-3

  H3O+       2.7 x 10-5 Small approx. 2.7 x 10-5
  CH3COO-       2.00 0.50 2.50 + 2.7 x 10-5 = 2.50
  CH2COOH       3.00 (-0.50) 2.50 - 2.7 x 10-5 = 2.50


Substituting the equilibrium concentrations of base (acetate ion) and conjugate acidThe acid formed when a base accepts a hydrogen ion (proton). (acetic acid) into the Henderson-Hasselbalch equation, Eq. (3), we have


\begin{align}\text{pH}&=\text{p}K_{a}\text{ + log}\frac{[\text{A}^{-}]}{[\text{HA}]}\
\text{ }&=-\text{log(1.8} \times \text{10}^{-5}\text{) + log}\frac{\text{(2.50 mol dm}^{-3}\text{)}}{\text{(2.50 mol dm}^{-3}\text{)}}\
\text{ }&=-\left(\text{0.25}-\text{5} \right)+ \text{log}\left(\text{1}\right)\
\text{ }&=\text{4.74 + 0}=\text{4.74}\end{align}


The addition of 0.5 mol sodium hydroxide to buffer mixture has thus succeeded in raising its pH from 4.57 to only 4.74. If the same 0.5 mol had been added to a cubic decimeter of pure water, the pH would have jumped all the way from 7.00 up to 13.7! The buffer is extremely effective at resisting a change in pH because the added hydroxide ion attacks the weak acid (in very high concentration) rather than the hydronium ion (in very low concentration). The major effect of the addition of the hydroxide ion is thus to change the ratio of acid to conjugate base, i.e., to change the value of


\frac{[\text{CH}_{3}\text{COOH}]}{[\text{CH}_{3}\text{COO}^{-}]}


As long as the amount of weak acid is much larger than the amount of base added, this ratio is not altered by very much. Since the hydronium-ion concentration is governed by


[\text{H}_{3}\text{O}^{+}]=K_{a}\frac{[\text{CH}_{3}\text{COOH}]}{[\text{CH}_{3}\text{COO}^{-}]}


the hydronium-ion concentration and pH are also altered to only a small extent.

The ability of a buffer solution to resist large changes in pH has a great many chemical applications, but perhaps the most obvious examples of buffer action are to be found in living matterAnything that occupies space and has mass; contrasted with energy.. If the pH of human blood, for instance, gets outside the range 7.2 to 7.6, the results are usually fatal. The pH of blood is controlled by the buffering action of several conjugate acid-base pairs. The most important of these is undoubtedly the H2CO3/HCO3 pair, but side chains of the amino acidA carboxylic acid containing an amino group (-NH2). In an alpha amino acid, the amino group is attached to the carbon atom adjacent to the carboxyl group. histidine in the hemoglobin moleculeA set of atoms joined by covalent bonds and having no net charge. also play a part. (Hemoglobin, a proteinA biological polymer of amino acids joined by peptide bonds., is the red substanceA material that is either an element or that has a fixed ratio of elements in its chemical formula. in the blood. It is responsible for carrying oxygen away from the lungs.) Most enzymes (biological catalysts) can only function inside a rather limited pH range and must therefore operate in a buffered environment. The enzymes which start the process of digestion in the mouth at a pH of around 7 become inoperative in the stomach at a pH of 1.4. The stomach enzymes in turn cannot function in the slightly basic environment of the intestines.