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Chapter 20 Acids and Bases

Chapter 20 Acids and Bases. Items from Chapter 19. Reversible Reactions - p. 539 In a reversible reaction, the reactions occur simultaneously in both directions double arrows used to indicate this 2SO 2(g) + O 2(g)  2SO 3(g)

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Chapter 20 Acids and Bases

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  1. Chapter 20Acids and Bases

  2. Items from Chapter 19... • Reversible Reactions - p. 539 • In a reversible reaction, the reactions occur simultaneously in both directions • double arrows used to indicate this 2SO2(g) + O2(g) 2SO3(g) • In principle, almost all reactions are reversible to some extent

  3. Items from Chapter 19... • Le Chatelier’s Principle - p.541 • If a stress is applied to a system in dynamic equilibrium, the system changes to relieve the stress. • Stresses that upset the equilibrium in a chemical system include: changes in concentration, changes in temperature, and changes in pressure

  4. Items from Chapter 19... • Equilibrium Constants (Keq) - p. 545 • Chemists generally express the position of equilibrium in terms of numerical values • These values relate to the amounts of reactants and products at equilibrium

  5. Items from Chapter 19... • Equilibrium Constants - p. 545 • consider this reaction: aA + bB  cC + dD • The equilibrium constant (Keq) is the ratio of product concentration to the reactant concentration at equilibrium, with each concentration raised to a power (= the coefficient)

  6. Items from Chapter 19... • Equilibrium Constants - p. 545 • consider this reaction: aA + bB  cC + dD • Thus, the “equilibrium constantexpression” has the general form: [C]c x [D]d [A]a x [B]b ( [ ] = molarity ) Keq =

  7. Items from Chapter 19... • Equilibrium Constants - p. 545 • the equilibrium constants provide valuable information, such as whether products or reactants are favored: Keq > 1, products favored at equilibrium Keq < 1, reactants favored at equilibrium • Sample Problem 19-2, p. 545

  8. Section 20.1Describing Acids and Bases • OBJECTIVES: • List the properties of acids and bases.

  9. Section 20.1Describing Acids and Bases • OBJECTIVES: • Name an acid or base, when given the formula.

  10. Properties of acids • Taste sour (don’t try this at home). • Conduct electricity. • Some are strong, others are weak electrolytes. • React with metals to form hydrogen gas. • Change indicators (blue litmus to red). • React with hydroxides to form water and a salt.

  11. Properties of bases • React with acids to form water and a salt. • Taste bitter. • Feel slippery (don’t try this either). • Can be strong or weak electrolytes. • Change indicators (red litmus turns blue).

  12. Names and Formulas of Acids • An acid is a chemical that produces hydrogen ions (H1+) when dissolved in water • Thus, general formula = HX, where X is a monatomic or polyatomic anion • HCl(g) named hydrogen chloride • HCl(aq) is named as an acid • Name focuses on the anion present

  13. Names and Formulas of Acids 1. When anion ends with -ide, the acid starts with hydro-, and the stem of the anion has the suffix -ic followed by the word acid 2. When anion ends with -ite, the anion has the suffix -ous, then acid 3. When anion ends with -ate, the anion suffix is -ic and then acid • Table 20.1, page 578 for examples

  14. Names and Formulas of Bases • A base produces hydroxide ions (OH1-) when dissolved in water. • Named the same way as any other ionic compound • name the cation, followed by anion • To write the formula: write symbols; write charges; then cross (if needed) • Sample Problem 20-1, p. 579

  15. Section 20.2Hydrogen Ions and Acidity • OBJECTIVES: • Given the hydrogen-ion or hydroxide-ion concentration, classify a solution as neutral, acidic, or basic.

  16. Section 20.2Hydrogen Ions and Acidity • OBJECTIVES: • Convert hydrogen-ion concentrations into values of pH, and hydroxide-ion concentrations into values of pOH.

  17. Hydrogen Ions from Water • Water ionizes, or falls apart into ions: H2O ® H1+ + OH1- • Called the “self ionization” of water • Occurs to a very small extent: [H1+ ] = [OH1-] = 1 x 10-7 M • Since they are equal, a neutral solution results from water • Kw = [H1+ ] x [OH1-] = 1 x 10-14M2 • Kw is called the “ion product constant”

  18. Ion Product Constant • H2O H+ + OH- • Kw is constant in every aqueous solution: [H+] x [OH-] = 1 x 10-14 M2 • If [H+] > 10-7 then [OH-] < 10-7 • If [H+] < 10-7 then [OH-] > 10-7 • If we know one, other can be determined • If [H+] > 10-7, it is acidic and [OH-] < 10-7 • If [H+] < 10-7, it is basic and [OH-] > 10-7 • Basic solutions also called “alkaline” • Sample problem 20-2, p. 582

  19. Logarithms and the pH concept • Logarithms are powers of ten. • Review from earlier lessons, and p. 585 • definition: pH = -log[H+] • in neutral pH = -log(1 x 10-7) = 7 • in acidic solution [H+] > 10-7 • pH < -log(10-7) • pH < 7 (from 0 to 7 is the acid range) • in base, pH > 7 (7 to 14 is base range)

  20. pH and pOH • pOH = -log [OH-] • [H+] x [OH-] = 1 x 10-14 M2 • pH + pOH = 14 • Thus, a solution with a pOH less than 7 is basic; with a pOH greater than 7 is an acid

  21. 100 10-1 10-3 10-5 10-7 10-9 10-11 10-13 10-14 14 13 11 9 7 5 3 1 0 10-14 10-13 10-11 10-9 Basic 10-7 10-5 10-3 10-1 100 [H+] pH 0 1 3 5 7 9 11 13 14 Acidic Neutral Basic pOH [OH-]

  22. Examples: • Sample 20-3, p.586 • Sample 20-4, p.586 • Sample 20-5, p.587 • Sample 20-6, p.588

  23. Measuring pH • Why measure pH? • Everything from swimming pools, soil conditions for plants, medical diagnosis, soaps and shampoos, etc. • Sometimes we can use indicators, other times we might need a pH meter

  24. Acid-Base Indicators • An indicator is an acid or base that undergoes dissociation in a known pH range, and has different colors in solution (more later in chapter) • Examples: litmus, phenolphthalein, bromthymol blue: Fig 20.8, p.590

  25. Acid-Base Indicators • Although useful, there are limitations to indicators: • usually given for a certain temperature (25 oC), thus may change at different temperatures • what if the solution already has color? • ability of human eye to distinguish colors

  26. Acid-Base Indicators • A pH meter may give more definitive results • some are large, others portable • works by measuring the voltage between two electrodes • needs to be calibrated • Fig. 20.10, p.591

  27. Section 20.3Acid-Base Theories • OBJECTIVES: • Compare and contrast acids and bases as defined by the theories of Arrhenius, Brønsted-Lowry, and Lewis

  28. Section 20.3Acid-Base Theories • OBJECTIVES: • Identify conjugate acid-base pairs in acid-base reactions.

  29. Svante Arrhenius • Swedish chemist (1859-1927) - Nobel prize winner in chemistry (1903) • one of the first chemists to explain the chemical theory of the behavior of acids and bases • Dr. Hubert Alyea-last graduate student of Arrhenius. (link below) http://www.woodrow.org/teachers/ci/1992/Arrhenius.html

  30. Hubert N. Alyea (1903-1996)

  31. 1. Arrhenius Definition • Acids produce hydrogen ions (H1+) in aqueous solution. • Bases produce hydroxide ions (OH1-) when dissolved in water. • Limited to aqueous solutions. • Only one kind of base (hydroxides) • NH3 (ammonia) could not be an Arrhenius base.

  32. Svante Arrhenius (1859-1927)

  33. Polyprotic Acids • Some compounds have more than 1 ionizable hydrogen. • HNO3 nitric acid - monoprotic • H2SO4 sulfuric acid - diprotic - 2 H+ • H3PO4 phosphoric acid - triprotic - 3 H+ • Having more than one ionizable hydrogen does not mean stronger!

  34. Polyprotic Acids • However, not all compounds that have hydrogen are acids • Also, not all the hydrogen in an acid may be released as ions • only those that have very polar bonds are ionizable - this is when the hydrogen is joined to a very electronegative element

  35. Arrhenius examples... • Consider HCl • What about CH4 (methane)? • CH3COOH (ethanoic acid, or acetic acid) - it has 4 hydrogens like methane does…? • Table 20.4, p. 595 for bases

  36. 2. Brønsted-Lowry Definitions • Broader definition than Arrhenius • Acid is hydrogen-ion donor (H+ or proton); base is hydrogen-ion acceptor. • Acids and bases always come in pairs. • HCl is an acid. • When it dissolves in water, it gives it’s proton to water. • HCl(g) + H2O(l) H3O+ + Cl- • Water is a base; makes hydronium ion.

  37. Johannes Bronsted / Thomas Lowry (1879-1947) (1874-1936)

  38. Acids and bases come in pairs... • A conjugate base is the remainder of the original acid, after it donates it’s hydrogen ion • A conjugate acid is the particle formed when the original base gains a hydrogen ion • Indicators are weak acids or bases that have a different color from their original acid and base

  39. Acids and bases come in pairs... • General equation is: • HA(aq) + H2O(l) H3O+(aq) + A-(aq) • Acid + Base Conjugate acid + Conjugate base • NH3 + H2O NH41+ + OH1- base acid c.a. c.b. • HCl + H2O H3O1++ Cl1- • acid base c.a. c.b. • Amphoteric - acts as acid or base

  40. 3. Lewis Acids and Bases • Gilbert Lewis focused on the donation or acceptance of a pair of electrons during a reaction • Lewis Acid - electron pair acceptor • Lewis Base - electron pair donor • Most general of all 3 definitions; acids don’t even need hydrogen! • Sample Problem 20-7, p.599

  41. Gilbert Lewis (1875-1946)

  42. Section 20.4Strengths of Acids and Bases • OBJECTIVES: • Define strong acids and weak acids.

  43. Section 20.4Strengths of Acids and Bases • OBJECTIVES: • Calculate an acid dissociation constant (Ka) from concentration and pH measurements.

  44. Section 20.4Strengths of Acids and Bases • OBJECTIVES: • Arrange acids by strength according to their acid dissociation constants (Ka).

  45. Section 20.4Strengths of Acids and Bases • OBJECTIVES: • Arrange bases by strength according to their base dissociation constants (Kb).

  46. Strength • Strong acids and bases are strong electrolytes • They fall apart (ionize) completely. • Weak acids don’t completely ionize. • Strength different from concentration • Strong-forms many ions when dissolved • Mg(OH)2 is a strong base- it falls completely apart when dissolved. • But, not much dissolves- not concentrated

  47. Measuring strength • Ionization is reversible. • HA H+ + A- • This makes an equilibrium • Acid dissociation constant = Ka • Ka = [H+ ][A- ] (water is constant) [HA] • Stronger acid = more products (ions), thus a larger Ka (Table 20.8, p.602)

  48. What about bases? • Strong bases dissociate completely. • B + H2O BH+ + OH- • Base dissociation constant = Kb • Kb = [BH+ ][OH-] [B] (we ignore the water) • Stronger base = more dissociated, thus a larger Kb.

  49. Strength vs. Concentration • The words concentrated and dilute tell how much of an acid or base is dissolved in solution - refers to the number of moles of acid or base in a given volume • The words strong and weak refer to the extent of ionization of an acid or base • Is concentrated weak acid possible?

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