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CHM135H1 Lecture Notes - Lecture 24: Buffer Solution, Alkalosis, Weak Base

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1 Nov 2018

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Gilbert Walker, Eugenia Kumacheva, Kristine Quinlan

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CHM135H1 Lecture 23: Fri Oct 26 2018

CHM135H1 Lecture Notes - Lecture 24: Buffer Solution, Alkalosis, Weak Base

CHM135H1 Lecture Notes - Lecture 25: Prope

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Chm135: lecture 23 applications of equilibria (part 2/3) October 31, 2018 (relevant reading: chemistry 2nd ed. Ha(aq) + h2o(l) a- (aq) + h3o+ (aq) =[3+][ ] [(cid:1853)(cid:1855)(cid:1856)](cid:4679) ph of a buffer is close to pka of the weak acid (plus log [base]/[acid]). For [base] = [acid], log is: ph = pka. Note: the concentrations in the henderson hasselbach equation are equilibrium concentrations of a- and ha. You either need to find concentrations (ice) or assume (if you can) that the initial concentrations are the same as the equilibrium concentrations (i. e. x is small). This is a good assumption if ka is small. Buffer systems maintain a constant ph in blood. The body maintains the ph of blood at ~7. 4; if the ph level changes just a few tenths of a ph unit, serious health consequences can result. A decrease in blood ph is called acidosis, an increase is called alkalosis.

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Experiment 18: Buffers Purpose In this experiment you will prepare several buffers and investigate their properties. Discussion A solution of a weak acid and one of its salts, in approximately equal concentrations, resists changes in pH when acid or base is added to it. (Note: The pH of a solution is a measure of the concentration of free H3Ot ion; it does not indicate concentration of combined acid such as HA.) The same is true for a solution of a weak base and one of its salts. A solution with this property is called a buffer (see illustration next page). Bufter systems are extremely important in many chemical and biological systems. For example, your blood contains a buffer system utilizing carbonic acid and bicarbonate ion. The important equilibrium in an acidic buffer is HA (aq) H20 ) A (a)0 () acid conjugate base in which the concentrations of acid and conjugate base are approximately equal. If a small amount of free acid (H30) is added to this solution, it reacts with the conjugate base A (reaction 1 in reverse). The equilibrium constant for this reaction is very large, so the reaction proceeds to near completion, absorbing the added acid. Similarly, added base (OH-) reacts with the acid HA as shown below: 2. HA (aq) + OH-(aq) ATaq) + H2O (1) This reaction also proceeds to completion, effectively absorbing any added base. Since the H30 and OH- concentrations are kept constant, the solution thus resists pH changes when an acid or base is added to it. A similar argument can be made for a basic buffer The actual pH of a buffer can be determined from the value of the Ka and the concentrations of the acid and conjugate base obtained from an initial-change- equilibrium table. However, a shortcut for buffers is available, and is shown below. We begin by writing and rearranging the equilibrium constant expression for the acid dissociation. If the -log of both sides of equation (3) is taken, a relationship involving pH is then obtained. H30IA] HA] 3. HA HAI -loglH30'1- -log Ka A- 141

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