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Acids and Bases

Acids and Bases. Operational definitions are based on observed properties. Compounds can be Classified as acid or base by observing these sets of properties. Properties of Acids. Taste sour (acere – Latin for sour) (Lemons, vinegar)

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Acids and Bases

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  1. Acids and Bases Operational definitions are based on observed properties. Compounds can be Classified as acid or base by observing these sets of properties.

  2. Properties of Acids • Taste sour (acere – Latin for sour) (Lemons, vinegar) • Cause certain organic dyes to change colour (Turns blue litmus paper to red – BAR) • Acid properties are destroyed by Bases (React with bases to form a salt and water) • Acid solutions are Electrolytes (substance in solution that conduct an electric current – Acids can be strong or weak electrolytes) • Acids react (corrode) with active metals (Group I and II as well as Zn and Aluminum) (Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)) • Acids react with carbonates (CO32-) and hydrogen carbonates (HCO31-) to produce carbon dioxide gas {2HCl(aq) + Na2CO3(s) → 2NaCl(aq) + H2O(l) + CO2(g)} • Certain nonmetal oxides will dissolve to produce acid solutions. (SO3(g) + H2O → H2SO4(aq) (SO3(g) is the acid anhydride – without water)

  3. Properties of Bases • Bases taste bitter; mustard and soap • Bases cause weak organic acids (dyes) to change colour (red litmus paper to blue {BB} Basic Blue • Acids destroy base properties - react with acids to form salts and water • Bases are electrolytes {strong or weak} • Feel soapy, slippery • Bases are formed when the oxide of some metals dissolve in water (CaO(s) + H2O → Ca(OH)2(aq) {CaO is the base anhydride}

  4. Acid/Base definitions • Definition #1: Arrhenius (traditional) • Acids are compounds with ionizable hydrogen– produce H+ ions (or hydronium ions H3O+) in solution • Bases are compounds that produce OH- ions in solution (problem: some bases don’t have hydroxide ions!) The reaction between an acid and a base:H+(aq) + OH-(aq) → H2O (l)

  5. Arrhenius acid is a substance that produces H+ (H3O+) in water. The HCl molecule is ionized. (ionization) Arrhenius base is a substance that produces OH- in water. The ions are dissociated. (dissociation)

  6. Some acids have more than one ionizable hydrogen • H2SO4 → H+(aq) + HSO41-(aq • HSO41- → H+(aq) + SO42-(aq) H2SO4 is diprotic • H3PO4(aq) → H+(aq) + H2PO41-(aq) • H2PO41-(aq) → H+(aq) + HPO42-(aq) • HPO42-(aq) → H+(aq) + PO43-(aq) Phosphoric acid is a triprotic acid.

  7. Water self-ionization H2O ↔ H+(aq) + OH-(aq) [H+] = [OH-] = 10-7Mat SATP Keq = [H+][ OH-] [H2O(l)] Kw = [H+][ OH-] = 10-7 x 10-7 (at 25ºC) Kw = 10-14 at SATP

  8. H2O ↔ H+(aq) + OH-(aq) What happens to this equilibrium if HCl(g) dissolves in the water? HCl(g) + H2O(l) → H3O+(aq) + Cl-(aq) • Increasing Decreasing • H2O ↔ H3O+(aq) + OH-(aq) • [H+] > [OH-] = acidic • What happens when sodium hydroxide dissolves? NaOH(s) + H2O → Na+(aq) + OH-(aq) • Decreasing Increasing • H2O ↔ H3O+(aq) + OH-(aq) • [H+] < [OH-] = basic (alkaline solution) If [H+] = 10-7 then [OH-] = 10-7 solution is neutral (SATP)

  9. pH and logs • [H+] is important in the study of acid-base chemistry. pH is the widely used scale to show [H+]. • pH = -log[H+] or pH = 1 . log[H+] [H+] = 10 – pH (the antilog) A logarithm is the power to which ten must be raised to get a number. log1000 = log(103) = 3

  10. pH calculations • For a neutral solution • pH = -log[H+] • pH = -log [10-7] • pH = - [-7] • pH = 7 at SATP • Example: • [H+] = 5 x 10-3 • pH = -log [5 x 10-3] • pH = -log [0.005] • pH = - (-2.3) = 2.3

  11. pH and pOH • pOH = - log [OH-] or [OH-] = 10 - pOH • Kw = [H+] x [OH-] = 1 x 10-14 (at 25ºC) • pKw= pH + pOH • 14= pH + pOH • Example: • If pH = (2.3) what is the [OH-]? • pH + pOH = 14 • pOH = 14 – pH • pOH = 14 – 2.3 • pOH = 11.7 • pOH = -log [OH-] • [OH-] = inverse log -11.7 or (10 - 11.7) • [OH-] = 2.0 x 10-12

  12. [H3O+], [OH-] and pH • What is the pH of the 0.0010 M NaOH solution? • [OH-] = 0.0010 (or 1.0 X 10-3 M) • pOH = - log 0.0010 • pOH = 3 • pH + pOH = 14 • pH = 14 – 3 = 11 • OR Kw = [H3O+] [OH-] • 1.0 x10-14 = [H3O+] x 1.0 X 10-3 • [H3O+] = 1.0 x 10-11 M • pH = - log (1.0 x 10-11) = 11.00

  13. Problem 1:The pH of rainwater collected in a certain region of the northeastern New Brunswick on a particular day was 4.82. What is the H+ ion concentration of the rainwater? [H+] = 1.51 x 10-5 Problem 2:The OH- ion concentration of a blood sample is 2.5 x 10-7M. What is the pH of the blood? pOH = 6.6 pH = 7.4 Problem 3: A chemist dilutes concentrated hydrochloric acid to make two solutions: (a) 3.0 M and (b) 0.0024 M. Calculate the [H3O+], pH, [OH-], and pOH of the two solutions at 25°C. b) [H3O+] = [2.4x10-3], pH = 2.62, pOH = 11.38, [OH-] = 4.2 x 10-12 a) [H3O+] = [3.0], pH = - 0.48, pOH = 14.48, [OH-] = 3.3 x 10-15 Problem 4: What is the [H3O+], [OH-], and pOH of a solution with pH = 3.67? Is this an acid, base, or neutral? [H3O+] = 2.14 x10-4, pOH = 10.33, [OH-] = 4.68x 10-11 It is an acid. Problem 5: Problem #4 with pH = 8.05? [H3O+] = 8.92 x10-9, pOH = 5.95, [OH-] = 1.12x 10-6 It is an acid.

  14. Definition #2: Brønsted – Lowry Acids – proton donor A “proton” is a hydrogen ion (the atom lost it’s electron) Bases – proton acceptor (accepts a hydrogen ion) No longer needs to contain the OH- ion Acid/Base Definitions

  15. A Brønsted-Lowry acid is a proton donor A Brønsted-Lowry base is a proton acceptor acid conjugateacid conjugatebase base

  16. + + + H H O H O Cl Cl H H H The Bronsted-Lowry concept base conjugate acid conjugate base acid conjugate acid-base pairs • Acids and bases are identified based on whether they donate or accept H+. • “Conjugate” acids and bases are found on the products side of the equation. A conjugate base is the same as the starting acid minus H+.

  17. Practice problems Identify the acid, base, conjugate acid, conjugate base, and conjugate acid-base pairs: CH3OOH(aq) + H2O(l)  CH3COO–(aq) + H3O+(aq) conjugate base conjugate acid acid base conjugate acid-base pairs OH–(aq) + HCO3–(aq)  CO32–(aq) + H2O(l) base acid conjugate base conjugate acid conjugate acid-base pairs

  18. Base Conjugate acid \ \ NH3(g) + H2O(l) ↔ NH4+(aq) + OH-(aq) / / Acid Conjugate Base HCl(aq) + H2O(l) ↔ H3O+(aq) + Cl-(aq) Acid Base Conjugate Conjugate Acid Base The water has acted as both an acid and a base, depending on what it is mixed with. Substances that can act as both an acid and a base are amphoteric (also called amphiproteric).

  19. Strong acid and base : HA(aq) + H2O(l) ↔ H3O+(aq) + A-(aq) B(aq) + H2O(l) ↔ BH+(aq) + OH-(aq) At equilibrium the ionic form is favored Weak acid and base : HA(aq) + H2O(l) ↔ H3O+(aq) + A-(aq) B(aq) + H2O(l) ↔ BH+(aq) + OH-(aq) At equilibrium the molecular form is favored

  20. CH3COOH(aq) + H2O(l) ↔ H+(aq) + CH3COO-(aq) Keq= [H+] [CH3COO-] . [CH3COOH] [H2O(l)] [H2O] is a constant, so collect the constants (Keq)[H2O(l)] = [H+] [CH3COO-] [CH3COOH]  (Keq)[H2O] is represented Ka(ionization constant for an acid) Ka = [H+] [CH3COO-] = 1.8 x 10-5 [CH3COOH] Ka < 1 weak acid General Formula for the ionization constant of a weak acid.

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