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Acids and Bases

Explore the use of sulfuric acid in industry and learn about the properties and classification of acids and bases. Understand the models of acids and bases, their conjugate pairs, acid strength, and the pH scale.

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Acids and Bases

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  1. Acids and Bases

  2. Acids in Industry • Sulfuric acid, H2SO4, is the chemical manufactured in greatest quantity in the U.S. • Eighty billion pounds of sulfuric acid are used each year to manufacture: • -fertilizers -pharmaceuticals • -detergents -storage batteries • -plastics -metals • -petroleum

  3. 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

  4. 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

  5. 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

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

  7. 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!

  8. Hydronium Ion • Hydronium (H3O+)ion is a hydrated proton -- H+ . H2O. • The H+ ion is simply a proton. It has a very high charge density, so it is strongly attracted to the very electronegative oxygen of the polar water molecule.

  9. Conjugate Acid/Base Pairs • HA(aq) + H2O(l)H3O+(aq) + A(aq) • conj conj conjconj acid 1base 2acid 2base 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-?

  10. Conjugate Acid-Base Pairs • Conjugate Acid Substance Conjugate Base HOH • NH3 • HSO4- • H2PO4- • HPO42- OH- NH2- SO42- HPO42- PO43- H3O+ NH4+ H2SO4 H3PO4 H2PO41-

  11. Conjugate Acid-Base Pairs • Which of the following represent conjugate acid-base pairs? • a) HF, F- e)OH-, HNO3 • b) NH4+, NH3 f) H2O, H3O+ • c) HCl, H2O g) H2SO4, SO42- • d) HC2H3O2, C2H3O2- h) HClO4, ClO4-

  12. Conjugate Bases • Write the conjugate base for each of the following: • a) HClO4 • b) H3PO4 • c) CH3NH3+ ClO4- H2PO4- CH3NH2

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

  14. Acid Strength(continued) • Its equilibrium lies far to the left. (CH3COOH or HC2H3O2) • Yields a much stronger (water is relatively strong) conjugate base than water. (CH3COO or C2H3O2-) Weak Acid:

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

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

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

  18. 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)

  19. Bases(continued) • weak = very little dissociation (or reaction with water) • NH3(aq) + HOH(l)  NH4+(aq) + OH(aq)

  20. 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 • Kw = [H+][OH-] • Only about two molecules in a billion ionize!!

  21. 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-] [H+] < 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.

  22. [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

  23. [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

  24. [H+] & [OH-] Calculations • Calculate the [H+] for a 2.0 x 10-2 M OH-. • Kw = [H+][OH-] • [H+] = Kw/[OH-] • [H+] = 1.0 x 10-14 M2/2.0 x 10-2 M • [H+] = 5.0 x 10-13 M

  25. The pH Scale • pH = log[H+] • pH in water usually ranges from 0 to 14. • Kw = 1.0  1014 M2 = [H+] [OH] • pKw = 14.00 = pH + pOH • As pH rises, pOH falls (sum = 14.00).

  26. Figure 15.5: Indicator paper being used to measure the pH of a solution

  27. Figure 15.4: A pH meter

  28. pOH = 14 pOH = 7 pOH = 0 1 x 10-7 1x 10-14 1 x 100 OH - OH- H3O+ OH- H3O+ H3O+ 1 x 100 1 x 10-7 1 x 10-14 pH = 0 pH = 7 pH = 14

  29. 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.

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

  31. 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

  32. pH Calculations • Calculate the pH value for the following solution at 25 oC. • [H+] = 1.0 x 10-9 M pH = - log [H+] • pH = - log [1.0 x 10-9] • pH = 9.00

  33. pH Calculations • Calculate the pH for the following solution at 25 oC. • [OH-] = 1.0 x 10-6M pH + pOH = 14.00 • pOH = - log [OH-] pH = 14.00 - pOH • pOH = - log [1.0 x 10-6] pH = 14.00 - 6.00 • pOH = 6.00 pH = 8.00

  34. 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

  35. 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.

  36. 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.

  37. 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.

  38. A Buffered Solution • . . . resists change in its pH when either H+ or OH are added. • 1.0 L of 0.50 M H3CCOOH + 0.50 M H3CCOONa • pH = 4.74 • Adding 0.010 mol solid NaOH raises the pH of the solution to 4.76, a very minor change.

  39. Preparation of Buffered Solutions • Buffered solution can be made from: • 1. a weak acid and its salt (e.g. HC2H3O2 & NaC2H3O2). • 2. a weak base and its salt (e.g. NH3 & NH4Cl). • Other examples of buffered pairs are: • H2CO3 & NaHCO3 H3PO4 & NaH2PO4 • NaH2PO4 & Na2HPO4 Na2HPO4 & Na3PO4

  40. Characteristics of a Buffer • 1. The solution contains a weak acid HA and its conjugate base A-. • 2. The buffer resists changes in pH by reacting with any added H+ or OH- so that these ions do not accumulate. • 3. Any added H+ reacts with the base A-. • 4. Any added OH- reacts with the weak acid HA.

  41. Buffered Solutions • Used when need to maintain a certain pH in the system. • Blood • Buffers work by reacting with added H+1 or OH-1 ions so they do not accumulate and change the pH. • Buffers will only work as long as there is sufficient weak acid and conjugate base molecules present.

  42. Buffering Mechanism • HC2H3O2(aq) <---> H+(aq) + C2H3O2-(aq) • The buffering materials dissolved in the solution prevent added H+ or OH- from building up in solution.

  43. Buffering Capacity • . . . represents the amount of H+ or OH the buffer can absorb without a significant change in pH.

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