Introduction to Organic and Biochemistry (CHE 124)

1. Introduction to Organic and Biochemistry(CHE 124) Reading Assignment General, Organic, and Biological Chemistry: An Integrated Approach 3rd. 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 + H2O ⇌ CN- + H3O+ 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) H30+ • Hydrogen atom H· • Hydro (hydrogen) group H • Hydride H:or H¯ • Hydrogen gas or molecule H:H or H2 • Hydroxide HO or OH¯ • Hydroxyl group OH

5. Acid and Conjugate Base • Acid and conjugate base differ by presence or absence of a proton. Conjugate Acid Base HCN + H2O ⇌ CN- + H3O+ 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 N2O4(g)⇌2 NO2(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. • Keq = [NO2]2= 4.6 x 10-3 [ ] = molarity [N2O2] Keq = Products Reactants

8. Writing Equilibrium Equation Keq • To write an equilibrium constant (Keq) 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 Keq = [C]c [D]d [A]a [B]b

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

10. Ka • 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 H2O(l) + CO2 (g) ⇌ H2CO3(aq) water carbon dioxide carbonic acid • Describe where / why this reaction occurs? • Describe what happens if you increase [CO2] • Reaction proceeds in the forward direction (to the right) • Describe what happens if you decrease [CO2] • In which direction does the reaction proceed. • Describe what happens if you increase [H2CO3]

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 + H2O ⇌ Cl- + H3O+ Acid Base Base Acid NH3 + H2O ⇌ HN4+ + OH- Base Acid Acid Base

15. Water Can Ionize H2O (l) + H2O (l) ↔ H3O+ (aq) + OH(aq)hydronium ion hydroxide ion Acid Base Acid Base • Kw = [H3O+][OH-] = 1 X 10-14 Kw is water equilibrium constant.

16. pH pH = -log [H3O+] • Measure of [H3O+] • 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 HClO4 perchloric acid HNO3 nitric acid H2SO4 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- LiOH Lithium hydroxide NaOH Sodium hydroxide KOH Potassium 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) + H2O(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 H2CO3 ↔H3O+ + HCO3- • 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 [H3O+] (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. H2CO3 ↔H3O+ + HCO3- • carbonic acid hydronium ion bicarbonate + + H+ OH- ↔ H2O

25. Acidosis • Acidosis - low blood pH • Leads to light headedness, coma, death. • Respiratory Acidosis • Characteristics • Low blood pH; high blood PCO2; normal or high (if compensating) blood HCO3- • 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 PCO2; low blood HCO3- • 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 [H3O+] • 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, CO2 is exhaled at a faster rate thus reducing the partial pressure of carbon dioxide (PCO2). 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. PCO2 increase leads to increase [carbonic acid] and thus [hydronium] and a drop in pH.

27. Reducing [H3O+] 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 CO2 + H20 ↔ H2CO3 ↔H3O+ + HCO3- kidneys release H3O+ in urine. shifts equation kidneys generate / release bicarbonate shifts equation ↑ respiration rate (breathing becomes more rapid) causes ↓ pCO2 (lungs remove carbon dioxide from blood and release it into atmosphere) shifts equation. Think about exercise

29. Alkalosis • Alkalosis - high blood pH. • Leads to headaches, nervousness, cramps, and convulsions and death. • Respiratory alkalosis • Characteristics • high blood pH; low blood PCO2; normal or lower (if kidneys are compensating) blood HCO3- • 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 PCO2; high blood HCO3- • Causes • Excessive use of antacids and constipation. • See Table 7.7 p. 244 and Figure 7.12 p 245

30. Correcting Alkalosis CO2 + H20 ↔ H2CO3 ↔H3O+ + HCO3- kidneys generate and release acid into blood shifts equation kidneys remove HCO3- from blood and release it into urine. shifts equation ↓ respiration rate (breathing slows) causes ↑ pCO2 shifts equation.