Acids and Bases—Definitions Strong Acids and Bases Chemical Equilibrium Weak Acids and Bases

1. Chapter 8Acids, Bases and Buffers in the Body Acids and Bases—Definitions Strong Acids and Bases Chemical Equilibrium Weak Acids and Bases pH and the pH Scale pKa Amino Acids—Common Biological Weak Acids Buffers and Blood—The Bicarbonate Buffer System

2. Brønsted Acid – Base Theory H+ = Proton = hydrogen cation H3O+ = hydronium ion (a majority of protons exist as hydronium ions in water) but this is often represented simply as H+ inequations. H+ + H2O → H3O+ Lewis dot and VSEPR structures for H3O+ Acid = proton donor Base = proton acceptor Acid properties sour taste (vinegar – pickles – citrus fruits) corrosive to metals Base properties bitter taste (many herbs are bases) slippery feel (makes soap out of fats/oils on skin)

3. Strong Acids Dissociate completely when added to water. HCl→ H+ + Cl- H2O or …… HCl + H2O → H3O+ + Cl- An HCl solution has no HCl molecules floating around!! Write the dissociation reaction for perchloric acid.

4. Strong Bases e.g. LiOH, NaOH, KOH, or Ca(OH)2. Dissociate completely when added to water to produce hydroxide ions (OH-). NaOH → Na+ + OH- H2O or …… OH- + H3O+ → 2 H2O or simply OH- + H+ → H2O Write the dissociation reaction for Ca(OH)2.

5. Neutralization Reactions Acid + Base → salt (ionic compound) + water HCl + NaOH → NaCl + H2O Write the neutralization reaction for any two other strong acid base pairs. Strong Bases - LiOH, NaOH, KOH, or Ca(OH)2. Strong Acids – HClO4, H2SO4, HI, HBr, HCl, HNO3. What salt is formed form the reaction of nitric acid with calcium hydroxide? a) CaNO3 b) Ca2NO3 c) Ca(NO3)2 d) CaNO2 2 HNO3 + Ca(OH)2 → Ca(NO3)2 + 2 H2O

6. Kw = [H+] [OH-] = 1.0 x 10-14 Equilibrium H2O ↔ H+ + OH- 2H2O ↔ H3O+ + OH- A small # of the molecules in pure water will dissociate. The concentration of H+ in pure water at 25ºC is always 1.0 x 10-7 M. What is the [OH-] in pure water? a) 0 b) 1.0 x 10-7 M c) < 1.0 x 10-7 M d) > 1.0 x 10-7 M The Equilibrium Constant (K) K = [H+] [OH-] [H2O] ――― liquid water is not included in equilibrium expressions

7. Equilibrium K = __[NH3]2_ [N2] [H2]3 N2(g) + 3H2(g)↔ 2NH3(g) K = _[CO][H2]3_ [CH4] [H2O] H2O(g) + CH4(g)↔ CO(g) +3H2(g) K = ??? Include water when it’s a gas Equilibrium expressions do not include pure liquids or solids. Always include gases or solutes if the reaction is in a solution.

8. Le Châtelier’s Principle If the equilibrium of a reaction is disturbed ( adding more of one reagent, heat, or pressure) the concentrations of reagents will readjust until K is the same (at constant T). K = __[NH3]2_ [N2] [H2]3 N2(g) + 3H2(g)↔ 2NH3(g) + heat Add reagent that is reactant – Equilibrium shifts to make more product (right) Add reagent that is product – Equilibrium shifts to make more reactants (left) Remove reagent? – Equilibrium shifts to replace the reagent removed. Add/remove heat (increase/decrease T) – Treat heat same as reactant or product For the reaction above increasing T will cause equilibrium to shift to make more reactants. The value of K for any reaction changes with temperature. Increase Pressure – Equilibrium shifts to reduce P by shifting to side with less gas. For the reaction above increasing P will cause equilibrium to shift to make more products. The Haber process used to make ammonia is carried out at 10 atm. P to increase yield.

9. Weak Acids e.g. Acetic Acid (CH3 – COOH) Only partially dissociate when added to water. Write the dissociation reaction for acetic acid. (Use H+ rather than H3O+) Use the ↔ symbol to indicate an equilibrium is established. CH3 – COOH(aq) ↔ H+(aq) + CH3-COO- . Write the equilibrium expression for acetic acid. Use Ka for the constant (the a indicates it is an acid) Ka = [H+] [CH3-COO-] = 1.75 x 10-5 [CH3-COOH] Write the dissociation reaction and equilibrium expression for HCN and HF. Ka = 6.2 x 10-10 for HCN and 6.5 x 10-4 for HF. Which acid is stronger? a) HCN b) HF

10. Weak Acids e.g. Acetic Acid (CH3 – COOH) Only partially dissociate when added to water. CH3 – COOH(aq) ↔ H+(aq) + CH3-COO- . Ka = [H+] [CH3-COO-] = 1.75 x 10-5 [CH3-COOH] Le Châtelier’s Principle – weak acids and their salts – buffers What happens when you add sodium acetate to a solution containing acetic acid? What is the formula for sodium acetate? NaCH3-COO Is the compound …. a) ionic b) covalent Is this compound soluble in water …. a) yes b) no Which way will the reaction above shift due to adding sodium acetate?…. a) right b) left

11. Weak Acids Only partially dissociate when added to water. Ka = [H+] [CH3-COO-] = 1.75 x 10-5 [CH3-COOH] CH3 – COOH(aq) ↔ H+(aq) + CH3-COO-(aq) . Conjugate base. Le Châtelier’s Principle – weak acids and their salts – buffers NaCH3-COO(aq) → Na+(aq) + CH3-COO-(aq). Will this reaction occur …. Na+(aq) + OH-(aq) ↔ NaOH(aq)? a) yes b) no Will this reaction occur …. CH3-COO-(aq) + H2O ↔ OH-(aq) + CH3-COOH(aq) a) yes b) no The soluble salt of a weak acid is a/an a) acid b) base

12. Polyprotic acids – more than one acidic hydrogen atom. H2CO3 and H3PO4 A separate dissociation equation and equilibrium expression can be written for each acidic H. H2CO3(aq) → H2O(ℓ) + CO2(aq) ↔ H+(aq) + HCO3-(aq) Ka = [H+][HCO3-]=4.5 x 10-7 [H2CO3] pKa1 = -log Ka = 6.35 HCO3-(aq) → H+(aq) + CO32-(aq) Ka = [H+][CO32-]=4.8 x 10-11 [HCO3-] pKa2 = -log Ka = 10.32

13. pH scale pOH = - log [OH-] pH = - log [H+] Kw = [H+] [OH-] = 1.0 x 10-14 H2O ↔ H+ + OH- In pure water [H+] = [OH-] = 1.0 x 10-7 M pH = 7.0 and pOH = 7.0note that pH + pOH = 14 (this is always true in aqueous solution) This follows from [H+] [OH-] = 1.0 x 10-14 (log x + log y = log xy) pOH is rarely used but rather converted into pH by ….. pH = 14 - pOH

14. pH + pOH = 14 pOH = - log [OH-] pH = - log [H+] Strong Acids Dissociate completely when added to water. HCl(aq)→ H+(aq)+ Cl-(aq) a) 1.0 b) 1.50 c) 2.00 d) 2.50 M What is the pH of a 0.020 M HCl solution? What is the [H+] in this solution? a) 0.0 M b) 0.010 M c) 0.020 M d) 1.00 M This is the approximate composition and pH of stomach acid What is the pH of a 0.020 M CH3COOH (acetic acid) solution? a) < 2.0 b) = 2.0 c) > 2.0 but less than 7.0 d) > 7.0

15. pH + pOH = 14 pOH = - log [OH-] pH = - log [H+] Strong Bases Dissociate completely when added to water. KOH(aq)→ K+(aq)+ OH-(aq) a) 2.0 b) 7.0 c) 12.0 d) 14.0 What is the pH of a 0.020 M KOH solution? The [OH-] is 0.02 M since a strong base will completely dissociate in water. Strategy: find the pOH --- then calculate the pH What is the pH of a 0.020 M NH3 (ammonia) solution? NH3 + H2O ↔ NH4+ + OH-. a) < 2.0 b) = 2.0 c) > 2.0 but less than 7.0 d) > 7.0

16. pH scale

17. Weak Acids Only partially dissociate when added to water. Ka = [H+] [CH3-COO-] = 1.75 x 10-5 [CH3-COOH] pKa = 4.76 CH3 – COOH(aq) ↔ H+(aq) + CH3-COO-(aq) . Conjugate base. Le Châtelier’s Principle – weak acids and their salts – buffers What is a buffer solution? A solution that will resist changes in pH when either an acid or base are added A solution containing the combination of a weak acid and its conjugate base. Often made by mixing the acid with the salt of the acid. e.g. acetic acid + sodium acetate CH3 – COOH(aq) ↔ H+(aq) + CH3-COO-(aq). Add acid …. ↑ reaction shifts a) right b) left pH decreases slightly (you are trading strong acid for weak acid) Add base …. reaction shifts a) right b) left OH- + H+ → H2O & ……. ↓ pH increases slightly

18. Buffer Demo pKa = 7.2 or … pH ~ 7.2 when [H2PO4-(aq)] = [ HPO42-(aq)] H2PO4-(aq) → H+(aq) + HPO42-(aq) Test water buffering capacity Prepare phosphate buffer – addHPO42-(aq) first will this a) add or b) remove H+? Test phosphate buffering capacity What is a buffer solution? A solution that will resist changes in pH when either an acid or base are added A solution containing the combination of a weak acid and its conjugate base. Often made by mixing the acid with the salt of the acid. e.g. acetic acid + sodium acetate

19. What is a buffer solution? A solution that will resist changes in pH when either an acid or base are added A solution containing the combination of a weak acid and its conjugate base. Often made by mixing the acid with the salt of the acid. e.g. acetic acid + sodium acetate What is the Blood’s buffer system? H2CO3(aq) ↔H2O(ℓ) + CO2(aq)↔ H+(aq) + HCO3-(aq) Ka = [H+][HCO3-]=4.5 x 10-7 [H2CO3] pKa1 = -log Ka = 6.35

20. What is the Blood’s buffer system? H2CO3(aq) ↔H2O(ℓ) + CO2(aq)↔ H+(aq) + HCO3-(aq) Blood pH is maintained in a narrow range of 7.35-7.45. If the blood pH drops below this range, a condition called acidosis occurs. If the blood pH becomes elevated, a condition called alkalosis exists. How does cellular activity affect blood pH?

21. 9.8 Buffers and Blood—The Bicarbonate Buffer System

22. 9.8 Buffers and Blood—The Bicarbonate Buffer System • Metabolic alkalosis can occur with excessive vomiting. • To lower the pH, ammonium chloride can be given.

23. Amino Acids H O | || H2N – C – C – OH | R H O | || ↔ H2N – C – C – O- | R + H+ H O | || H3N+ – C – C – OH ↔ | R + H+ The amino group is a base and can accept a proton. The carboxylic acid group is an acid and can donate a proton Given that ‘R’ for the amino acid alanine is –CH3, Draw the structure of alanine. Given that the pK’s are 2.3 for the -COOH group and 9.7 for the –NH2 group, describe it’s ionic form at various pH values. What is it’s ionic form in blood at pH 7.4?

24. Chapter 8Acids, Bases and Buffers in the Body Acids and Bases—Definitions Strong Acids and Bases Chemical Equilibrium Weak Acids and Bases pH and the pH Scale pKa Amino Acids—Common Biological Weak Acids Buffers and Blood—The Bicarbonate Buffer System

25. Define acids and bases using the Brønsted Model. Distinguish between the descriptive characteristics of acids and bases related to taste and physical properties. Distinguish between strong and weak acids/bases. Know the strong acids and the strong bases produced by hydroxides of Ca and any Group 1 metal. Write equations representing the dissociation of strong acids and bases in aqueous solution. Describe chemical equilibrium. Construct the equilibrium expression for water. Write the expression for Kw. Construct equilibrium equations for weak acids and bases. Write equilibrium expressions for weak acids and bases based on these equations. Describe LeChâtelier’s Principle. For equilibrium equations be able to indicate the direction of the reaction change for adding/removing reactants, adding/removing products, changing the temperature (for reactions with energy listed as a reactant or product), and changes in pressure (for reactions with gas phase reactants/products). Define buffers. Relate buffering activity to acid-base reactions and LeChâtelier’s Principle. Describe the general components of a buffer system. Describe the buffering system of the blood. Construct equations for blood the bicarbonate/CO2 buffer system. Draw the general structure of an amino acid where R represents the variable side chain. Explain how these compound can act as both an acid and a base. Given that ‘R’ for the amino acid alanine is –CH3, Draw the structure of alanine. Given that the pK’s are 2.3 for the -COOH group and 9.7 for the –NH2 group, describe it’s ionic form at various pH values. What is it’s ionic form in blood at pH 7.4?