# 5.111 Principles of Chemical Science

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2 5.111 Principles of Chemical Science Solutions to selected biology-related questions from problem sets for lectures Acid-base problems A. A pharmaceutical molecule with antifungal properties is only active when deprotonated and negatively charged (A - ). The protonated state (HA) is inactive. If the pk a of this drug is 10.0, (a) calculate the ratio of protonated to deprotonated compound at physiological ph (7.4). (b) Is this drug likely to be a useful pharmaceutical agent? (a) calculate the ratio of protonated to deprotonated compound at physiological ph (7.4). Since we are given both the ph and pk a of the compound, we can use the Henderson- Hasselbalch equation to solve for the ratio of [HA] to [A - ]. ph = pk a - log([ha] / [A - ]) log([ha] / [A - ]) = pk a ph log([ha] / [A - ]) = log([ha] / [A - ]) = 2.6 ([HA] / [A - ]) = ([HA] / [A - ]) = 400 (correct sig figs) The ratio of protonated (inactive) compound to deprotonated (active) compound is 400 to 1 at physiological ph. (b) Is this drug likely to be a useful pharmaceutical agent? Since the vast majority of the compound is in the inactive form at physiological ph, it is unlikely to be a useful pharmaceutical agent.* Ideally, most of the compound would be active in the body. *However, if the active compound is highly potent, it is possible that a small fraction of active compound is sufficient for useful antifungal activity. B. Absorption of aspirin (acetylsalicylic acid, C 9 H 8 O 4,) into the bloodstream occurs only when the molecule is in its conjugate base form. (a) If a patient takes two tablets of aspirin (325 mg each), how many grams of aspirin are available for immediate absorption in the stomach? The ph of the stomach is 1.6, and the pk a of aspirin is 3.5. Since we are given both the ph of the stomach and the pk a of aspirin, we can use the Henderson-Hasselbalch equation to solve for the ratio of [HA] to [A - ]. ph = pk a - log([ha] / [A - ]) log([ha] / [A - ]) = pk a ph log([ha] / [A - ]) = log([ha] / [A - ]) = 1.9 ([HA] / [A - ]) = 79 1

3 The ratio of protonated aspirin to its conjugate base is 79 to 1. So one-eightieth (1/80) of the total aspirin taken will be in the conjugate base form and available for immediate absorption in the stomach: 2 x 325 mg x (1/80) = 8.75 mg 9 mg (b) Would you expect more or less aspirin to be absorbed in the small intestine (ph 7.5) compared to the stomach? Briefly explain your answer (no calculation is required). More aspirin will be absorbed in the small intestine. The higher ph in the intestine means that more aspirin will be in the conjugate base form and therefore available for absorption. C. (Modified from problem in Atkins 3 rd ed) Phosphate buffers are very useful in biochemical experiments. Your research supervisor asks you to make up a phosphate buffer to carry out kinetic assays on an enzyme: (a) Determine the ratio of molarities of PO 4 and HPO 4 ions required to prepare a phosphate buffer with a ph of 12.0? (b) What mass of K 3 PO 4 must be added to 1.00 L of M K 2 HPO 4 (aq) to prepare a buffer solution with a ph of 12.0? (a) We need to think about what happens when you have a solution of PO 4 and HPO 4 in water: HPO 4 (aq) + H 2 O ( ) PO 4 (aq) + H 3 O +1 (aq). We can look up the K a for that reaction in table10.9 (page 391 in the text), recognizing that we are looking at the ionization of the third proton. The K a is 2.1 x Now to determine the ratios of ions we can substitute into the Henderson-Hasselbach equation: ph = pk a - log HPO 4 ( ) log HPO = log( 2.1 x10 13 ) 12.0 = log HPO [ 4 ] 0.68 = log HPO [ 4 ] [ HPO 4 ] = [ HPO 4 ] [ HPO 4 ] = 0.2 PO 4 = 4.79 = 5 (correct sig figs) 2

4 The ratio of HPO 4 to PO 4 is 5, and the ratio of PO 4 to HPO 4 is 0.2 (don t take off points for significant figures) (b) We need to realize that we are adding the potassium salt of HPO 4-2, so the concentration of HPO 4-2 is M, so we can substitute that into the expression we calculated in part (a): HPO 4 = = 4.8 = 2.09 x 10-2 M [ ( )] g mol -1 Molar mass of K 3 PO 4 = 3( ) = g mol -1 mol Mass of K 3 PO 4 = 2.09 x 10-2 ( g mol -1 )1 L L = g ( ) =4 g We need to add 4 g of K 3 PO 4 to 1.0 L. (don t take off points for significant figures) Oxidation-reduction problems D. Hydrogen peroxide (H 2 O 2 ) is a harmful and reactive byproduct of metabolism. To prevent H 2 O 2 from causing extensive oxidative damage to cells, the enzyme catalase catalyzes the conversion of H 2 O 2 to much less reactive molecules, oxygen and water. 2H 2 O 2 2H 2 O + O 2 Use oxidation numbers to identify which molecules are reduced and/or oxidized in this reaction. H 2 O 2 is both reduced to H 2 O (oxidation number of oxygen changes from -1 to -2) and oxidized to O 2 (oxidation number changes from -1 to 0). 3

5 E. Oxidative damage of DNA leads to mutation, which can in turn lead to cancer or genetic defects. The redox potentials of the four nucleotides of DNA are listed below. Which nucleotide (A, G, T, C) is the most likely to undergo oxidation? Explain your answer for full credit. Nucleotide ε (V) G 1.33 A 1.42 C 1.60 T 1.70 G is most likely to undergo oxidation. G has the lowest redox potential of the four bases, which indicates that it is the easiest to oxidize F. The enzyme known as laccase is a multicopper oxidase that is widely distributed in higher plants and fungi. The Fe(CN) 6 4- complex was used to study the structure of the active site of this enzyme. (a) Indicate the oxidation number of the central atom in this complex. Fe is the central atom. Since each of the six cyano (CN - ) groups carries a -1 charge and the total charge of the complex is 4 -, Fe has an oxidation number of 2 +. (b) Calculate the d-electron count The d-electron count of a metal is equal to the group number (from the periodic table) minus the oxidation number of the metal. The oxidation number of Fe is +2 (see part a) The group number of Fe is 8 d-electron count = 8 2 = 6 The d-electron count is d 6 (c )What is the coordination geometry? The coordination number (CN) is 6, so the geometry is octahedral. G. Zinc can act as a Lewis acid, and coordinate amino acids like cysteine by their sulfhydryl (SH) groups. Normally, the sulfhydryl group of cysteine has a pk a of ~8. However, the interaction with zinc can lower the pk a of the sulfhydryl group by as much as 4 ph units. Using your knowledge of acid/base chemistry, estimate the approximate ratio of protonated to deprotonated cysteine at neutral ph for (a) free cysteine (pk a of ~ 8) and 4

6 (b) cysteine coordinated to zinc (pk a of ~ 4). (Since this is an estimate, don t worry about sig figs). O O H 2 N CH C OH H 2 N CH C OH CH 2 CH 2 SH protonated cysteine S deprotonated cysteine We can treat this like a buffer problem and thus take advantage of the Henderson- Hasselbach equation: ph = pk a log [HA/A-] (a) Free cysteine (no zinc present): pk a ~8 ph = pk a log [HA] 7 = 8 log [HA] [HA] = 10 1 When there is no zinc around there is 10 times as much protonated cysteine as there is deprotonated cysteine (activated for nucleophillic attack and therefore the form of cysteine which can partake in chemical reactions). (b) Cysteine Coordinated to Zinc: pk a ~4 ph = pk a log [HA] 7 = 4 log [HA] [HA] = When zinc is coordinated to the cysteine there is 1000 times as much deprotonated cysteine as there is protonated cysteine. In this case a much greater portion of the cysteines are deprotonated and ready to partake in a chemical reaction. 5

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