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Acids and Bases

This chapter delves into the essential properties of acids and bases, including their tastes, effects on indicators, and ion production. It explores the classification of acids (organic, inorganic, and oxyacids) and the models used to understand acid-base behavior, including the Arrhenius and Brønsted-Lowry concepts. The role of hydronium ions and conjugate acid-base pairs is explained, along with acid strength and calculations involving pH and ion concentrations. The pH scale is discussed, highlighting the importance of Kw and its variation with temperature.

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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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