Exp. 8 Acid-Base Equilibria - Experiment 8 8 AcidBase Equilibria Determination of Acid Ionization Constants PURPOSE AND LEARNING OBJECTIVES To obtain

Exp. 8 Acid-Base Equilibria - Experiment 8 8 AcidBase...

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Experiment 8 8 93 Acid–Base Equilibria: Determination of Acid Ionization Constants PURPOSE AND LEARNING OBJECTIVES To obtain quantitative values of acid ionization constants by measur- ing the pH. To prepare buffer solutions and observe their resistance to change in pH. PRINCIPLES According to the Brønsted-Lowry definition of acids and bases, an acid is a proton (H ) donor and a base is a proton acceptor. The stronger the acid, the greater its ability to donate protons. Consider, for example, the reaction of HCl with water. HCl (aq) H 2 O (l) W H 3 O (aq) Cl (aq) or HCl (aq) W H (aq) Cl (aq) HCl is observed to dissociate essentially 100% in water. The K a value is large, K a ~10 7 , indicating the equilibrium position lies far to the right. HCl is therefore considered to be a strong acid. All acids with K a > 1 are considered strong acids. In this experiment the reaction of acetic acid, CH 3 COOH, with water will be studied. Unlike HCl, CH 3 COOH is a weak acid and dissociates only slightly in water. CH 3 COOH (aq) W H (aq) CH 3 COO (aq)
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EXPERIMENT 8 94 For CH 3 COOH, the value of K a is fairly small, K a 1.76 10 5 , indicating the equilibrium lies far to the left. In a 1.0 M CH 3 COOH solution, 99% of the acetic acid molecules remain undissociated. Buffer Solutions and the Determination of K a Acid–base reactions belong to a very important class of reactions. For example, in living systems the control of pH by acid–base equilibria is crucial for survival. Our average blood pH is 7.4. A 0.2 unit shift in pH results in serious changes in blood chemistry causing severe illness. A 0.4 unit shift in pH is generally fatal. An understanding of acid–base equilibria is necessary to un- derstand how blood pH is controlled. Buffer solutions are remarkably resistant to pH changes. All living systems contain buffer solutions. Buffer solutions generally consist of a weak acid and its conjugate base or a weak base and its conjugate acid. According to the Brønsted-Lowry theory, acid–base reactions produce a conjugate acid–base pair. Consider, for example, the reac- tion of an acid, HA, with water. HA (aq) H 2 O (l) W H 3 O (aq) A (aq) K HA H O A a 3 = + - 6 6 6 @ @ @ acid base conjugate acid conjugate base In this reaction, the HA gives up a proton to produce A , its conjugate base, and H 2 O acts as a base by accepting a proton to produce H 3 O , its conjugate acid. A weak acid produces a strong conjugate base and a strong acid produces a weak conjugate base. A buffered solution must contain appreciable amounts of both a weak acid and its conjugate base to resist a change in pH when a small amount of a strong acid or base is added to the solution. When a small amount of strong acid is added to a buffered solution, it reacts completely with the base, A . H (aq) A (aq) HA (aq) When a small amount of strong base is added to a buffered solution, it reacts completely with the weak acid, HA.
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