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Chapter 8

Chapter 8. Acids and Bases. Acids. Have a sour taste. Vinegar owes its taste to acetic acid. Citrus fruits contain citric acid. React with certain metals to produce hydrogen gas. React with carbonates and bicarbonates to produce carbon dioxide gas. Bases. Have a bitter taste.

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Chapter 8

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  1. Chapter 8 Acids and Bases

  2. Acids Have a sour taste. Vinegar owes its taste to acetic acid. Citrus fruits contain citric acid. React with certain metals to produce hydrogen gas. React with carbonates and bicarbonates to produce carbon dioxide gas Bases Have a bitter taste. Feel slippery. Many soaps contain bases.

  3. Acids • Acids are a class of chemical compounds. Originally, they were defined as substances that tasted sour. • Vinegar • Citrus fruits • Arrhenius defined acids as substances that produced the hydrogen ion when dissolved in water.

  4. Bases • Bases are a class of chemical compounds, originally defined as substances that tasted bitter and left the skin with a slippery feeling. • Antacids • Oven cleaner • Arrhenius defined bases as substances that produced the OH- ion when dissolved in water.

  5. A Brønsted acid is a proton donor A Brønsted base is a proton acceptor base acid acid base A Brønsted acid must contain at least one ionizable proton!

  6. HI (aq) H+ (aq) + I- (aq) CH3COO- (aq) + H+ (aq) CH3COOH (aq) H2PO4- (aq) H+ (aq) + HPO42- (aq) H2PO4- (aq) + H+ (aq) H3PO4 (aq) Identify each of the following species as a Brønsted acid, base, or both. (a) HI, (b) CH3COO-, (c) H2PO4- Brønsted acid Brønsted base Brønsted acid Brønsted base

  7. Strength of Acids and Bases • Strong acids and strong bases react completely with water. • Weak acids and weak bases exist in equilibrium in water. • See Table 8.1 pg 242 • You need to memorize the strong acids and bases

  8. H2O NaCl (s)Na+ (aq) + Cl- (aq) HCl (aq) + H2O (l) H3O+ (aq) + Cl- (aq) HNO3 (aq) + H2O (l) H3O+ (aq) + NO3- (aq) CH3COOHCH3COO- (aq) + H+ (aq) HClO4 (aq) + H2O (l) H3O+ (aq) + ClO4- (aq) H2SO4 (aq) + H2O (l) H3O+ (aq) + HSO4- (aq) Strong Electrolyte – 100% dissociation Weak Electrolyte – not completely dissociated Strong Acids are strong electrolytes

  9. HF (aq) + H2O (l) H3O+ (aq) + F- (aq) HNO2 (aq) + H2O (l) H3O+ (aq) + NO2- (aq) HSO4- (aq) + H2O (l) H3O+ (aq) + SO42- (aq) H2O (l) + H2O (l) H3O+ (aq) + OH- (aq) H2O NaOH (s) Na+ (aq) + OH- (aq) H2O KOH (s) K+ (aq) + OH- (aq) H2O Ba(OH)2 (s) Ba2+ (aq) + 2OH- (aq) Weak Acids are weak electrolytes Strong Bases are strong electrolytes

  10. F- (aq) + H2O (l) OH- (aq) + HF (aq) NO2- (aq) + H2O (l) OH- (aq) + HNO2 (aq) Weak Bases are weak electrolytes • Conjugate acid-base pairs: • The conjugate base of a strong acid has no measurable strength. • H3O+ is the strongest acid that can exist in aqueous solution. • The OH- ion is the strongest base that can exist in aqeous solution.

  11. Conjugate Acid/Base • Conjugate acid:When a Bronsted-Lowry base accepts a hydrogen ion. • Conjugate base: When a Bronsted-Lowry acid donates a hydrogen ion. • Conjugate acid-base pair: These two molecules together.

  12. base acid acid base conjugate base conjugate acid base acid Table 8.2 page 245

  13. Amphiprotic • When a substance can act as an acid or a base • Carbonic acid and water • Sulfuric acid and water • Problem 8.1 page 246

  14. Strong acids have weak conjugate bases • Weak acids have strong conjugate bases • Strong bases have weak conjugate acids • Weak bases have strong conjugate acids • Problem 8.2 page 249

  15. HA (aq) + H2O (l) H3O+ (aq) + A- (aq) HA (aq) H+ (aq) + A- (aq) [H+][A-] Ka = [HA] Weak Acids (HA) and Acid Ionization Constants Ka is the acid ionization constant pKa is the negative log Ka pKa = - log Ka • Table 8.3 page 249 • Problem 8.3 and 8.4 page 250

  16. Equilibrium Expressions • Because weak acids and weak bases exist in equilibrium, we can write equilibrium expressions for their reactions. • Write Ka for H2CO3, NH4+ and H3PO4

  17. + H H O O [ ] + + - H H H O H O H H H2O (l) H+ (aq) + OH- (aq) H2O + H2O H3O+ + OH- Acid-Base Properties of Water autoionization of water conjugate acid base acid conjugate base

  18. [H+][OH-] Kc = [H2O] H2O (l) H+ (aq) + OH- (aq) The Ion Product of Water [H2O] = constant Kc[H2O] = Kw = [H+][OH-] The ion-product constant (Kw) is the product of the molar concentrations of H+ and OH- ions at a particular temperature. Solution Is [H+] = [OH-] neutral At 250C Kw = [H+][OH-] = 1.0 x 10-14 [H+] > [OH-] acidic [H+] < [OH-] basic • Problem 8.5 page 254

  19. = [OH-] = 1 x 10-14 Kw 1.3 [H+] What is the concentration of OH- ions in a HCl solution whose hydrogen ion concentration is 1.3 M? Kw = [H+][OH-] = 1.0 x 10-14 [H+] = 1.3 M = 7.7 x 10-15M

  20. pH [H+] pH – A Measure of Acidity pH = -log [H+] Solution Is At 250C neutral [H+] = [OH-] [H+] = 1 x 10-7 pH = 7 [H+] > 1 x 10-7 pH < 7 acidic [H+] > [OH-] basic [H+] < [OH-] [H+] < 1 x 10-7 pH > 7

  21. Other important relationships pH = -log [H+] [H+][OH-] = Kw = 1.0 x 10-14 [H+] = 10-pH pH Meter Table 8.4 page 258

  22. [H+] = Kw [OH-] The pH of rainwater collected in a certain region of the northeastern United States on a particular day was 4.82. What is the H+ ion concentration of the rainwater? pH = -log [H+] = 10-4.82 = 1.5 x 10-5M [H+] = 10-pH The OH- ion concentration of a blood sample is 2.5 x 10-7 M. What is the pH of the blood? pH = -log [H+] • Problem 8.6 page 256

  23. Reactions of Acids and Bases • When an acid reacts directly with a base, the reaction is called neutralization. • In general, neutralization reactions produce water and an ionic compound. • There are some examples of neutralization reactions that do not produce water, but we will not focus on those reactions in this class. • Look at the reaction between hydrochloric acid and sodium hydroxide

  24. Titrations In a titration a solution of accurately known concentration is added gradually added to another solution of unknown concentration until the chemical reaction between the two solutions is complete. Equivalence point – the point at which the reaction is complete Indicator – substance that changes color at (or near) the equivalence point Slowly add base to unknown acid UNTIL the indicator changes color

  25. H2SO4 + 2NaOH 2H2O + Na2SO4 M M rxn volume acid moles red moles base volume base base acid coef. 4.50 mol H2SO4 2 mol NaOH 1000 ml soln x x x 1000 mL soln 1 mol H2SO4 1.420 mol NaOH What volume of a 1.420 M NaOH solution is required to titrate 25.00 mL of a 4.50 M H2SO4 solution? WRITE THE CHEMICAL EQUATION! 25.00 mL = 158 mL

  26. M V rxn volume red moles red moles oxid M oxid oxid red coef. 5 mol Fe2+ 1 0.1327 mol KMnO4 x x 5Fe2+ + MnO4- + 8H+ Mn2+ + 5Fe3+ + 4H2O x 1 mol KMnO4 0.02500 L Fe2+ 1 L 16.42 mL of 0.1327 M KMnO4 solution is needed to oxidize 25.00 mL of an acidic FeSO4 solution. What is the molarity of the iron solution? WRITE THE CHEMICAL EQUATION! 16.42 mL = 0.01642 L 25.00 mL = 0.02500 L 0.01642 L = 0.4358 M Problem 8.8 page 261

  27. Consider an equal molar mixture of CH3COOH and CH3COONa H+ (aq) + CH3COO- (aq) CH3COOH (aq) OH- (aq) + CH3COOH (aq) CH3COO- (aq) + H2O (l) • A buffer solution is a solution of: • A weak acid or a weak base and • The salt of the weak acid or weak base • Both must be present! A buffer solution has the ability to resist changes in pH upon the addition of small amounts of either acid or base. Add strong acid Add strong base

  28. HCl H+ + Cl- HCl + CH3COO- CH3COOH + Cl-

  29. Acetic acid has a Ka of 1.8 x 10-5 . What is the pH of a buffer solutions containing 0.15 M Acetic Acid and 0.10 M acetate? Nitrous acid has a Ka of 4.5 x 10-4 . What is the pH of a buffer solutions containing 0.10 M HNO2 and 0.50 M NO2- ? Calculate the pH of the 0.30 M NH3/0.36 M NH4Cl buffer system if the Ka is 5.6 x 10-10 ? Problem 8.9 page 264 Problem 8.10 page 266

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