Assistant Lecture Aayad Amaar

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1 Assistant Lecture Aayad Amaar Acid-Base Buffers A buffer solution contains components that enable the solution to resist large changes in ph when either acids or bases are added. Buffer solutions may be prepared in the laboratory to maintain optimum conditions for a chemical reaction. Buffers are routinely used in commercial products to maintain optimum conditions for product behavior. Buffer solutions also occur naturally. Blood, for example, is a complex natural buffer solution maintaining a ph of approximately 7.4, optimum for oxygen transport. The major buffering agent in blood is the mixture of carbonic acid (H 2 CO 3 ) and bicarbonate ions (HCO 3 - ). The Buffer Process The basis of buffer action is the establishment of equilibrium between either a weak acid and its conjugate base or a weak base and its conjugate acid. Let s consider the case of a weak acid and its salt. A common buffer solution may be prepared from acetic acid (CH3COOH) and sodium acetate (CH3COONa). Sodium acetate is a salt that is the source of the conjugate base CH3COO -. An equilibrium is established in solution between the weak acid and the conjugate base. A buffer solution functions in accordance with LeChatelier s principle, which states that an equilibrium system, when stressed, will shift its equilibrium to relieve that stress. This principle is illustrated by the following examples. Addition of Base (OH - ) to a Buffer Solution Addition of a basic substance to a buffer solution causes the following changes. OH - from the base reacts with H 3 O + producing water.

2 Molecular acetic acid dissociates to replace the H 3 O + consumed by the base, maintaining the ph close to the initial level. This is an example of LeChatelier s principle, because the loss of H 3 O + (the stress) is compensated by the dissociation of acetic acid to produce more H 3 O +. Addition of Acid (H 3 O + ) to a Buffer Solution Addition of an acidic solution to a buffer results in the following changes. H3O + from the acid increases the overall [H 3 O + ]. The system reacts to this stress, in accordance with LeChatelier s principle, to form more molecular acetic acid; the acetate ion combines with H 3 O +. Thus, the H 3 O + concentration and therefore, the ph, remain close to the initial level. These effects may be summarized as follows: Buffer capacity is a measure of the ability of a solution to resist large changes in ph when a strong acid or strong base is added. More specifically, buffer capacity is described as the amount of strong acid or strong base that a buffer can neutralize without significantly changing its ph. Buffering capacity against base is a function of the concentration of the weak acid (in this case CH 3 COOH). Buffering capacity against acid is dependent on the concentration of the anion of the salt, the conjugate base (CH 3 COO - ). Preparation of a Buffer Solution It is useful to understand how to prepare a buffer solution and how to determine the ph of the resulting solution. Many chemical reactions produce the largest amount of product only when they are run at a known, constant ph. The study of biologically important processes in the laboratory often requires conditions that approximate the composition of biological fluids. A constant

3 ph would certainly be essential. The buffer process is an equilibrium reaction and is described by an equilibrium constant expression. For acids, the equilibrium constant is represented as Ka, the subscript a implying an acid equilibrium. For example, the acetic acid/sodium acetate system is described by Using a few mathematical maneuvers we can turn this equilibrium-constant expression into one that will allow us to calculate the ph of the buffer if we know how much acid (acetic acid) and salt (sodium acetate) are present in a known volume of the solution. First, multiply both sides of the equation by the concentration of acetic acid, [CH 3 COOH]. This will eliminate the denominator on the right side of the equation. Now, dividing both sides of the equation by the acetate ion concentration [CH 3 COO - ] will give us an expression for the hydronium ion concentration [H 3 O + ] Once we know the value for [H 3 O + ], we can easily find the ph. To use this equation: assume that [CH 3 COOH] represents the concentration of the acid component of the buffer. assume that [CH 3 COO - ] represents the concentration of the conjugate base (principally from the dissociation of the salt, sodium acetate) component of the buffer.

4 Example The Henderson-Hasselbalch Equation The solution of the equilibrium-constant expression and the ph are sometimes combined into one operation. The combined expression is termed the Henderson-Hasselbalch equation. For the acetic acid/sodium acetate buffer system, Where pk a = - log K a The form of this equation is especially amenable to buffer problem calculations. In this expression, [CH 3 COOH] represents

5 the molar concentration of the weak acid and [CH 3 COO - ] is the molar concentration of the conjugate base of the weak acid. The generalized expression is:

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