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15_applications_of_aqueous_equilibria_print

# 15_applications_of_aqueous_equilibria_print - The Common...

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10/11/2009 1 Chapter 15 Applications of Aqueous Equilibria The Common Ion Effect This occurs when a salt, NaA is added to the solution of a weak acid, HA, sharing the common ion A - Addition of the common ion is causes the dissociation equilibrium of the acid to shift to the left (according to Le Châtelier’s Principle), making the solution less acidic than if the ion were not present Example: If the salt NaF is added to the solution of the weak acid, the presence of the common ion F - causes the equilibrium to shift to the left and the [H + ] concentration to decrease and hence the pH to increase: The procedure for finding the pH of a weak acid solution containing a common ion are similar to the calculations for weak acids covered in Chapter 14. The only difference is that the initial concentration of the anion A - , is not zero due to the addition of the salt NaA Calculate [H + ] and the percent dissociation of HF in a solution containing 1.0 M HF ( K a = 7.2 x 10 -4 ) and 1.0 M NaF. How does this compare to the 2.7% dissociation of a 1.0 M HF solution alone? [H + ] = 7.2 x 10 -4 M (pH 3.14), % dissociation = 0.072% Buffer Solutions Normally the pH changes by a large amount even when a small amount of acid or base is added to pure water: A buffer solution is a solution which resists a change in pH when small amounts of acid or base are added: An example of a buffer solution is blood which can absorb acids and bases without changing its pH much. This property is vital since cells can only function over a narrow pH range A buffer solution can be made in two ways: 1. A solution of a weak acid and a salt of its conjugate base 2. A solution of a weak base and a salt of its conjugate acid Buffer solutions are examples of acid-bases solutions containing a common ion By choosing the appropriate components it is possible to create a buffer solution at any chosen pH

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10/11/2009 2 pH = 4.74 Procedure for Calculating the pH Change when an Acid or Base is added to a Buffer Solution Calculate the pH change that occurs when 0.010 mol NaOH is added to 1.0 L of the buffer solution in the previous question. Compare this to the pH change which occurs when 0.010 mol NaOH is added to pure water pH changes from 4.74 to 4.76 pH changes from 7.00 to 12.00 How does a buffer work? Suppose a buffer is made from equal quantities of a weak acid, HA and a salt of its conjugate base, A - The weak acid is present is large amounts since it only dissociates slightly while the salt provides large quantities of its conjugate base When a small amount of OH - ions are added to the solution they are neutralized by the acid producing water: HA(aq) + OH - (aq) A - (aq) + H 2 O(l) There will be a small decrease in HA (present in large amounts) and a small increase in A - (also present in large amounts) but the concentration of H + won’t change much. In other words the OH - ions are replaced by A - and are not allowed to accumulate The pH stability can be understood from the equilibrium expression for the acid: K a = [H + ][A - ]/[HA] or [H + ] = K a x [HA]/[A - ] so the equilibrium concentration of [H + ] (and hence the pH) depends on the ratio [HA]/[A - ]
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