1. Introduction to Organic and Biochemistry (CHE 124) Reading Assignment General, Organic, and Biological Chemistry: An Integrated Approach 3 rd. Ed. Ramond Chapter 7 Acids, Bases, and Equilibrium Gasses, Solutions, Colloids, and Suspensions Work Problems 7. 8, 12, 24, 26, 29, 30, 32, 36, 40, 44, 52, 56, 60, 66

2. Acid / Base Acid –Sour taste (never taste lab chemicals!) –Dissolves metals –Turns litmus pink Base –Bitter taste (never taste lab chemicals!) –Feel slippery (soapy) –Turns litmus blue See common acids and bases Table 7.1 p. 224.

3. Acid / Definitions Arrhenius definition –Acid – compound that produces H + (protons) in aqueous solution. HCl → H + + Cl - –Base – compound that produces OH - in aqueous solution. NaOH → Na + + OH - Bronsted-Lowery definition –Acid – releases H + (protons). –Base - H + (proton) acceptor HCN + H 2 O ⇌ CN - + H 3 O + arrows mean reversible Acid Base Base Acid Reversible - means products can be converted into products and products can be converted into reactants.

4. Hydrogen and Related Species Proton (hydron)H + Hydronium ion (interchangable with proton) H 3 0 + Hydrogen atom H· Hydro (hydrogen) groupH HydrideH:orH ¯ Hydrogen gas or moleculeH:H or H 2 HydroxideHO orOH ¯ Hydroxyl groupOH

5. Acid and Conjugate Base Acid and conjugate base differ by presence or absence of a proton. Conjugate Acid Base HCN + H 2 O ⇌ CN - + H 3 O + Conjugate Acid Base Amphoteric – compound that can act as an acid or a base.

6. Equilibrium Consider the reversible reaction of decomposition of dinitrogen tetroxide to form nitrogen dioxide. –See Fig. 7.2 p. 226 N 2 O 4 (g) ⇌ 2 NO 2(g) colorless brown Eventually the color stops changing (getting browner). This is equilibrium – the rate of the forward and reverse reaction are equal. The concentration of each species remains constant. Note the double arrow.

7. Equilibrium Constant If the concentration of the reactant and product of an equilibrium equation are determined then the following equation is true. K eq = [NO 2 ] 2 = 4.6 x 10 -3 [ ] = molarity [N 2 O 2 ] K eq = Products Reactants

8. Writing Equilibrium Equation K eq To write an equilibrium constant (K eq ) equation. –Before you start, BALANCE the EQUATION! –ONLY SPECIES WHOSE CONCENTRATION CAN CHANGE ARE INCLUDED. –Do NOT include solvents or solids in the equation. aA + bB ⇌ cC + dD A,and B are reactants C and D are products a,b,c,and d are coefficients K eq = [C] c [D] d [A] a [B] b

9. What does K eq Tell us? K eq > 1 (larger number) –[reactant] < [product] –Favors products K eq < 1 (small number) –[reactant] > [product] –Favors reactants Keq = 1 –[reactant] = [product]

10. KaKa Ka = acidity constant –Special name of Keq for acid base reactions. pKa = -log Ka.

11. Le Chatelier’s Principle Le Chatelier’s Principle states that when a reversible reaction is pushed out of equilibrium, the reaction responds to reestablish equilibrium. –Vary [reactant] or [product] by adding (or removing) reactants or products.

12. Example of Le Chatelier’s Principle carbonic anhydrase H 2 O(l) + CO 2 (g) ⇌ H 2 CO 3 (aq) water carbon dioxide carbonic acid –Describe where / why this reaction occurs? –Describe what happens if you increase [CO 2 ] Reaction proceeds in the forward direction (to the right) –Describe what happens if you decrease [CO 2 ] In which direction does the reaction proceed. –Describe what happens if you increase [H 2 CO 3 ]

13. Catalysts Catalyst increase the rate of the reaction by lowering the activation energy. –Catalyst Do not alter the equilibrium Do not alter the Keq.

14. Water is Amphoteric Amphoteric a compound that can act as an acid or a base. HCl + H 2 O ⇌ Cl - + H 3 O + Acid Base Base Acid NH 3 + H 2 O ⇌ HN 4 + + OH - Base Acid Acid Base

15. Water Can Ionize H 2 O (l) + H 2 O (l) ↔ H 3 O + (aq) + OH(aq) hydronium ion hydroxide ion Acid BaseAcid Base K w = [H 3 O + ][OH - ] = 1 X 10 -14 K w is water equilibrium constant.

16. pH pH = -log [H 3 O+] Measure of [H 3 O + ] scale is continuum from 0 - 14 –7 is neutral; Neutral - neither acidic or basic –0 - 6.99 is acidic –7.01 - 14 is basic (alkaline) one pH unit change represents 10 fold change in [H + ] See Fig. 7.6 p 233 See Table 7.3 p. 233

17. pH of Strong Acids Strong Acid – dissociates 100% in water. HCl → H + + Cl - HCl hydrochloric acid (muriatic acid) HBr hydrobromic acid HI hydriodic acid HClO 4 perchloric acid HNO 3 nitric acid H 2 SO 4 sulfuric acid [H + ] is equal to the [H + ] of the acid Weak Acid – dissociates less than 100% in water. –All other acids

18. pH of Strong Bases Strong Base – dissociate 100% in water. NaOH → Na + + OH - LiOHLithium hydroxide NaOHSodium hydroxide KOHPotassium hydroxide Ca(OH) 2 Calcium hydroxide Sr(OH) 2 strontium hydroxide Ba(OH) 2 barium hydroxide pOH is dependent on the concentration of the strong base Weak Base – dissociates less than 100% in water.

19. Neutralization Neutralization reaction of an acid and base to form water and a salt. HCl(aq) + NaOH(aq) → NaCl(aq) + H 2 O(l) Acid Base Salt Water –If equal amounts of acid and base are added, the pH will equal 7. Titration –A technique used to determine the concentration of an acid or base solutions. Uses Buret and and an indicator See p. 238 Fig. 7.9

20. pH effects the Concentration of the Acid and Conjugate Base A few points to understand: –When pH = pKa [acid] = [conjugate base] –When pH < pKa [acid] > [conjugate base] –When pH > pKa [acid] < [conjugate base] This alters the charge on many biological molecules by changing them form the acid form to the base form (carboxylate ion) –Use the fatty acid example.

21. Buffer Buffer - substance that resists changes in pH thus stabilizing its relative pH. – Buffers are often a solution containing a weak acid and its conjugate base –See example on next slide. –Buffers work within 1 pH unit either side of the pKa of the weak acid.

22. Buffer Carbonic acid is a weak acid. It dissociates in aqueous solution to form hydronium ion and bicarbonate H 2 CO 3 ↔ H3O + + HCO 3 - carbonic acid hydronium ion bicarbonate

23. Buffering Blood pH of blood = 7.35 – 7.45 Blood carries many acids which can alter it’s pH. –Fatty acids, lactic acid, phosphoric acid, carbonic acid. Body uses two approaches to control pH (p.245 Fig 7.12) –Use of Buffers (see next slide) Carbonic Acid / Bicarbonate buffer system –Reduce [H 3 O + ] (see following slides) Action of lungs Filtering by kidneys

24. Use of Buffers Carbonic acid is a weak acid that buffers blood. It dissociates in aqueous solution to form hydronium ion (acid) and bicarbonate. H 2 CO 3 ↔ H 3 O + + HCO 3 - carbonic acid hydronium ion bicarbonate OH - + +H++H+ ↔ H2OH2O

25. Acidosis Acidosis - low blood pH Leads to light headedness, coma, death. –Respiratory Acidosis Characteristics –Low blood pH; high blood P CO2 ; normal or high (if compensating) blood HCO 3 - Causes –Diseases / conditions that limit carbon dioxide exchange by lungs such as ppneumonia, emphysema, cystic fibrosis, shallow breathing or holding your breath. –Metabolic Acidosis Characteristics –low blood pH; normal or low (if lungs are compensating) blood P CO2 ; low blood HCO 3 - Causes –Presence of ketone bodies ( acetone, acetoacetic acid, beta hydroxybuyteric acid) due to starvation or poorly controlled diabetes. See Table 7.7 p. 244 and Figure 7.12 p. 245

26. Reducing [H 3 O + ] Lungs remove excess acid through increase in respiration rate –As the blood becomes more acidic, the respiratory center of the brain signals for faster breathing. –With faster breathing, CO 2 is exhaled at a faster rate thus reducing the partial pressure of carbon dioxide (P CO2 ). This reduces the [carbonic acid] thus reducing the [hydronium] producing an increase in pH. –This happens when you exercise. Lungs remove excess base by reducing rate of respiration. –Breathing becomes slower and more shallow. P CO2 increase leads to increase [carbonic acid] and thus [hydronium] and a drop in pH.

27. Reducing [H 3 O + ] Cont’ Kidneys remove excess acid by releasing bicarbonate into the blood. –The increase in [bicarbonate] shifts the equilibrium toward carbonic acid. This reduces the [hydronium].

28. Correcting Acidosis CO 2 + H 2 0 ↔ H 2 CO 3 ↔ H 3 O + + HCO 3 - ↑ respiration rate (breathing becomes more rapid) causes ↓ pCO 2 (lungs remove carbon dioxide from blood and release it into atmosphere) shifts equation. Think about exercise kidneys generate / release bicarbonate shifts equation kidneys release H 3 O + in urine. shifts equation

29. Alkalosis Alkalosis - high blood pH. Leads to headaches, nervousness, cramps, and convulsions and death. –Respiratory alkalosis Characteristics –high blood pH; low blood P CO2 ; normal or lower (if kidneys are compensating) blood HCO 3 - Causes –Occurs when CO2 is exhaled from the body more quickly than it is produced by cells. –Hyperventilation brought on by anxiety, CNS damage, aspirin poisoning, fever, etc –Metabolic alkalosis Characteristics –high blood pH; normal or high (if lungs are compensating) blood P CO2 ; high blood HCO 3 - Causes –Excessive use of antacids and constipation. See Table 7.7 p. 244 and Figure 7.12 p 245

30. Correcting Alkalosis CO 2 + H 2 0 ↔ H 2 CO 3 ↔ H 3 O + + HCO 3 - ↓ respiration rate (breathing slows) causes ↑ pCO 2 shifts equation. kidneys generate and release acid into blood shifts equation kidneys remove HCO 3 - from blood and release it into urine. shifts equation