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Acids and Bases

Acids and Bases. Chapter 14. Properties of Acids. Acids: taste sour (citrus fruits & vinegar) affect indicators (e.g. turn blue litmus red) produce H + ions in aqueous solution corrosive to metals pH < 7. Classifying Acids.

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Acids and Bases

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  1. Acids and Bases • Chapter 14

  2. Properties of Acids • Acids: • taste sour (citrus fruits & vinegar) • affect indicators (e.g. turn blue litmus red) • produce H+ ions in aqueous solution • corrosive to metals • pH < 7

  3. Classifying Acids • Organic acids contain a carboxyl group or -COOH -- HC2H3O2 & citric acid. • Inorganic acids -- HCl, H2SO4, HNO3. • Oxyacids -- acid proton attached to oxygen -- H3PO4. • Monoprotic -- HCl & HC2H3O2 • Diprotic -- H2SO4 • Triprotic -- H3PO4

  4. Properties of Bases • Bases: • taste bitter • feel slippery • affect indicators (e.g. turn red litmus blue) • produce OH- ions in aqueous solution • pH > 7 • caustic

  5. Models of Acids and Bases • Arrhenius Concept: Acids produce H+ in solution, bases produce OH ion. • Brønsted-Lowry: Acids are H+ donors, bases are proton acceptors. • HCl + H2O  Cl + H3O+ • acid base

  6. Hydronium Ion • Hydronium ion is a hydrated proton -- H+.H2O. • The H+ ion is simply a proton. It has a very high charge density, so it strongly is attracted to the very electronegative oxygen of the polar water molecule.

  7. Conjugate Acid/Base Pairs • HA(aq) + H2O(l)  H3O+(aq) + A(aq) • conj conj conj conj acid 1 base 2 acid 2 base 1 • conjugate base: everything that remains of the acid molecule after a proton is lost. • conjugate acid: formed when the proton is transferred to the base. • Which is the stronger base--H2O or A-?

  8. Acid Dissociation Constant (Ka) • HA(aq) + H2O(l)  H3O+(aq) + A(aq) • Ka values for common acids are found in Table 14.2 on page 663.

  9. Bronsted-Lowry Model • The Bronsted-Lowry Model is not limited to aqueous solutions like the Arrhenius Model. • NH3(g) + HCl(g) ----> NH4Cl(s) • This is an acid-base reaction according to Bronsted-Lowry, but not according to Arrhenius!

  10. Acid Strength • Its equilibrium position lies far to the right. (HNO3) • Yields a weak conjugate base. (NO3) Strong Acid:

  11. Acid Strength(continued) • Its equilibrium lies far to the left. (CH3COOH) • Yields a much stronger (water is relatively strong) conjugate base than water. (CH3COO) Weak Acid:

  12. A strong acid is nearly 100 % ionized, while a weak acid is only slightly ionized.

  13. Diagram a represents a strong acid, while b represents a weak acid which remains mostly in the molecular form.

  14. The relationship of acid strength and conjugate base strength for acid-base reactions.

  15. Arranging Species According to Increasing Basic Strength • H2O, F-, Cl-, NO2-, & CN- • Use Table 14.2 on page 663. • Cl- is weakest since it is conjugate base of strong acid and weaker than water. Use Ka values to arrange the remaining bases. • Cl- < H2O < F- < NO2- < CN-

  16. Water as an Acid and a Base • Water is amphoteric (it can behave either as an acid or a base). • H2O + H2O  H3O+ + OH • conj conj acid 1 base 2 acid 2 base 1 • Kw = 1  1014 M2 at 25°C

  17. Ion product Constant, Kw • Kw is called the ion-product constant or dissociation constant. • neutral solution [H+] = [OH-] = 1.0 x 10 -7 M • acidic solution [H+] > [OH-] [H+] > 1.0 x 10-7 M • basic solution [H+] < [OH-] [OH-] > 1.0 x 10-7 M • No matter what the concentration of H+ or OH- in an aqueous solution, the product,Kw, will remain the same.

  18. [H+] & [OH-] Calculations • Calculate the [H+] for a 1.0 x 10-5 M OH-. • Kw = [H+][OH-] • [H+] = Kw/[OH-] • [H+] = 1.0 x 10-14 M2/1.0 x 10-5 M • [H+] = 1.0 x 10-9 M

  19. [H+] & [OH-] CalculationsContinued • Calculate the [OH-] for a 10.0 M H+. • Kw = [H+][OH-] • [OH-] = Kw/[H+] • [OH-] = 1.0 x 10-14 M2/10.0 M • [OH-] = 1.0 x 10-15 M

  20. Kw & H • At 60oC, the value of Kw is 1 x 10-13 for the dissociation of water: • 2 H2O(l) <---> H3O+(aq) + OH-(aq) • Is this reaction exothermic or endothermic? • Endothermic -- Kw increased with temperature.

  21. The pH Scale • pH = log[H+] • pH in water usually ranges from 0 to 14. • Kw = 1.00  1014 = [H+] [OH] • pKw = 14.00 = pH + pOH • As pH rises, pOH falls (sum = 14.00).

  22. pH & [H+] pH = 0 pH = 7 pH = 14 1x 10-14 1 x 10-7 1 x 100 OH - H3O+ OH- OH- H3O+ H3O+ 1 x 100 1 x 10-7 1 x 10-14

  23. Logarithms • -log 1.00 x 10-7 = 7.000 • 7.000 • characteristic mantissa • The number of significant digits in 1.00 x 10-7 is three, therefore, the log has three decimal places. The mantissa represents the log of 1.00 and the characteristic represents the exponent 7.

  24. pH scale and pH values for common substances. A pH of 1 is 100 times more acidic than a pH of 3.

  25. pH Calculations • What is the pOH, [H+], & [OH-] for human blood with a pH of 7.41? • pH + pOH = 14.00 • pOH = 14.00 - pH • pOH = 14.00 - 7.41 • pOH = 6.59

  26. pH CalculationsContinued • What is the pOH, [H+], & [OH-] for human blood with a pH of 7.41? • pH = - log [H+] • [H+] = antilog (-pH) • [H+] = antilog (-7.41) • [H+] = 3.9 x 10-8 M Note: The number of significant figures in the antilog is equal to the number of decimal places in the pH.

  27. pH CalculationsContinued • What is the pOH, [H+], & [OH-] for human blood with a pH of 7.41? • pOH = - log [OH-] • [OH-] = antilog (-pOH) • [OH-] = antilog (-6.59) • [OH-] = 2.6 x 10-7 M Note: The number of significant figures in the antilog is equal to the number of decimal places in the pOH.

  28. pH of Strong Acid Solutions • Calculate the pH of a 0.10 M HNO3 solution. • Major species are: H+, NO3-, and H2O • Sources of H+ are from HNO3 and H2O -- amount from water is insignificant. • [H+] = 0.10 M pH = - log [H+] • pH = - log [0.10] • pH = 1.00 Note: The number of significant figures in the [H+] is the same as the decimal places in the pH.

  29. pH & Significant Figures • log • # Significant Figures -------> # decimal places • <------- • inv log • pH = - log [H+] [H+] = inv log (-pH) • [H+] = 1.0 x 10-5 M pH = 5.00

  30. Solving Weak Acid Equilibrium Problems • List major species in solution. • Choose species that can produce H+ and write reactions. • Based on K values, decide on dominant equilibrium. • Write equilibrium expression for dominant equilibrium. • List initial concentrations in dominant equilibrium.

  31. Solving Weak Acid Equilibrium Problems (continued) • Define change at equilibrium (as “x”). • Write equilibrium concentrations in terms of x. • Substitute equilibrium concentrations into equilibrium expression. • Solve for x the “easy way.” x can be neglected when concentration is 2 powers of 10 (100x) greater than Ka or Kb. • Verify assumptions using 5% rule. • Calculate [H+] and pH.

  32. pH of Weak Acid Solutions • Calculate the pH of a 0.100 M HOCl solution. • Major species: HOCl and HOH • Ka HOCl = 3.5 x 10-8 & Ka HOH = 1.0 x 10-14 •  HOCl will be only significant source of [H+]. • Ka = 3.5 x 10-8 = [H+][OCl-]/[HOCl]

  33. pH of Weak Acid SolutionsContinued • ICE • [HOCl] [OCl-] [H+] • Initial (mol/L) 0.100 0 0 • Change (mol/L) - x + x + x • Equil. (mol/L) 0.100 - x 0 + x 0 + x

  34. pH of Weak Acid SolutionsContinued • Ka = 3.5 x 10-8 = [H+][OCl-]/[HOCl] • 3.5 x 10-8 = [x][x]/[0.100 - x] • Ka is more than 100 x smaller than concentration, x can be neglected in the denominator. • Ka = 3.5 x 10-8 = [x][x]/[0.100] • x2 = 3.5 x 10-9 • x= 5.9 x 10-5 M

  35. pH of Weak Acid SolutionsContinued • Approximation check: • % dissociation = (x/[HA]o) (100%) • % dissociation = (x/[HOCl]o) (100%) • % dissociation = (5.9 x 10-5/0.100)(100%) • % dissociation = 0.059 % • This is much less than 5 % and therefore the approximation was valid.

  36. Percent Dissociation (Ionization) The percent dissociation calculation is exactly the same as the one to check the 5 % approximation. See Sample Exercise 14.10 on pages 678 and 679.

  37. % Dissociation Calculations • In a 0.100 M lactic acid solution (HC3H5O3), lactic acid is 3.7 % dissociated. Calculate the Ka for this acid. • Major species: HC3H5O3 & HOH • HC3H5O3(aq) <---> H+(aq)+ C3H5O3-(aq) • Ka = [H+][C3H5O3-]/ [HC3H5O3]

  38. % Dissociation CalculationsContinued • ICE • [HC3H5O3] [C3H5O3-] [H+] • Initial (M) 0.10 0 0 • Change (M) - 3.7 x 10-3 + 3.7 x 10-3 + 3.7 x 10-3 • Equil. (M) 0.10 + 3.7 x 10-3 + 3.7 x 10-3

  39. % Dissociation CalculationsContinued • Ka = [H+][C3H5O3-]/ [HC3H5O3] • Ka = [3.7 x 10-3]2/ [0.10] • Ka = 1.4 x 10-4

  40. The effect of dilution on the % dissociation and [H+] of a weak acid solution.

  41. Bases • Bases are often called alkalis because they often contain alkali or alkaline earth metals. • “Strong” and “weak” are used in the same sense for bases as for acids. • strong = complete dissociation (hydroxide ion supplied to solution) • NaOH(s)  Na+(aq) + OH(aq)

  42. Bases(continued) • weak = very little dissociation (or reaction with water) • H3CNH2(aq) + H2O(l)  H3CNH3+(aq) + OH(aq) • See Table 14.3 on page 685 for Kb values of common bases. • Kb calculations are identical to Ka calculations.

  43. Polyprotic Acids • . . . can furnish more than one proton (H+) to the solution. See Table 14.4 on page 689 for Ka values for common polyprotic acids. Know Sample Exercises 14.15 & 14.16 on pages 689-692.

  44. Acid-Base Properties of Salts

  45. Structure and Acid-Base Properties • Two factors for acidity in binary compounds: • Bond Polarity (high is good) • Bond Strength (low is good)

  46. The effect of the number of attached oxygen on the H-O bond in a series of chlorine oxyacids.

  47. Oxides • Acidic Oxides (Acid Anhydrides): • OX bond is strong and covalent. • SO2, NO2, CrO3 • Basic Oxides (Basic Anhydrides): • OX bond is ionic. • K2O, CaO

  48. Lewis Acids and Bases • Lewis Acid: electron pair acceptor • Lewis Base: electron pair donor

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