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Chapter 18 Acids and Bases

Chapter 18 Acids and Bases. Items from Chapter 17. Reversible Reactions - p. 416 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 18 Acids and Bases

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  1. Chapter 18Acids and Bases

  2. Items from Chapter 17... • Reversible Reactions - p. 416 • 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 17... • Le Chatelier’s Principle - p.421 • 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 17... • Equilibrium Constants (Keq) - p. 418 • 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 17... • Equilibrium Constants - p. 418 • 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 17... • Equilibrium Constants - p. 418 • 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 17... • Equilibrium Constants - p. 418 • 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 17-8, p. 419

  8. Section 18.1Describing Acids and Bases • OBJECTIVES: • List the properties of acids and bases.

  9. Section 18.1Describing 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. Naming Acid Hydrogen _______ide becomes hydro____ic acid Hydrogen_______ate becomes _________ic acid Hydrogen_______ite becomes _______ous acid chlor chlor

  14. Names and Formulas of Acids • 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

  15. 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)

  16. Section 18.2Hydrogen Ions and Acidity • OBJECTIVES: • Given the hydrogen-ion or hydroxide-ion concentration, classify a solution as neutral, acidic, or basic.

  17. Section 18.2Hydrogen Ions and Acidity • OBJECTIVES: • Convert hydrogen-ion concentrations into values of pH, and hydroxide-ion concentrations into values of pOH.

  18. 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 10-14 (mol/L)2 • Kw is called the “ion product constant”

  19. Ion Product Constant • H2O H+ + OH- • Kw is constant in every aqueous solution: [H+] x [OH-] = 1 x 10-14 • If [H+] >1x 10-7 then [OH-] < 1x10-7 • If [H+] < 1x10-7 then [OH-] > 1x10-7 • If we know one,the 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 are also called “alkaline”

  20. Logarithms and the pH concept • Logarithms are powers of ten. • Know how to use the log buttons on your calculator • definition: pH = -log[H+] • in neutral pH = -log(1 x 10-7) = 7 • in acidic solution [H+] > 10-7 • pH < 7 (from 0 to 7 is the acid range) • in base, pH > 7 (7 to 14 is base range)

  21. 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

  22. pH 0 1 3 5 7 9 11 13 14 Acidic Neutral Basic 14 13 11 9 7 5 3 1 0 pOH

  23. 100 10-1 10-3 10-5 10-7 10-9 10-11 10-13 10-14 [H+] Acidic Neutral Basic 10-14 10-13 10-11 10-9 10-7 10-5 10-3 10-1 100 [OH-]

  24. Examples: • Sample 18-1, p.434 • Sample 18-2, p.435 • Sample 18-3, p.436 • Sample 18-4, p.438 • Sample 18-5, p.439

  25. Formula’s to Know Ion Product Constant Kw = [H3O+] [OH-] = 1 x 10-14 Acid Equations pH = - log[H30+] [H3O+] = antilog (-pH) Base Equations pOH = -log [OH-] [OH-] = antilog (-pOH) pH Scale Equation This is from your formula sheet pH + pOH = 14 Note H+ is the same as H3O+

  26. 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

  27. 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) • HIn H1+ + In1- • Examples: litmus, phenolphthalein, bromthymol blue: Fig 18.14, p.445

  28. 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

  29. Acid-Base Indicators • A pH meter(simulation)may give more definitive results • some are large, others portable • works by measuring the voltage between two electrodes • needs to be calibrated • Fig. 18.17, p.446

  30. Section 18.3Acid-Base Theories • OBJECTIVES: • Compare and contrast acids and bases as defined by the theories of Arrhenius, Brønsted-Lowry, and Lewis

  31. Section 18.3Acid-Base Theories • OBJECTIVES: • Identify conjugate acid-base pairs in acid-base reactions.

  32. 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

  33. 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) would not be an Arrhenius base. H2O HCl  H+ + Cl- H2O NaOH  Na+ + OH-

  34. 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!

  35. Polyprotic Acids • However, not all compounds that have hydrogen are acids • Also, not all the hydrogen in an acid may be released as ions CH3COOH CH3COO- + H+ • only those that have very polar bonds are ionizable - this is when the hydrogen is joined to a very electronegative element NH4+ or H2O

  36. Arrhenius examples... • Consider HCl • What about CH4(methane)? • CH3COOH (ethanoic acid, or acetic acid) - it has 4 hydrogens like methane does…? • Table 18.3, p. 442 for bases

  37. What is a Proton • A proton is a positively charged subatomic particle (no neutrons, no electrons) • A hydrogen atom is 1 proton(+) and 1 electron (-) • A hydrogen ion or proton (H+) is a hydrogen atom that has lost its electron leaving only the proton • H  H+(proton) + 1 electron

  38. 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.

  39. 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

  40. 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.

  41. HCl + H2O H3O1+ + Cl1- acid base c.a. c.b. Amphoteric - acts as acid or base (H20, HSO4-, H2PO4-) must be able to give a H+ or accept an H+

  42. 3. Lewis Acids and Bases • Gilbert Lewis focused on the donation or acceptance of a pair of electrons during a reaction • Lewis Acid - electron deficient • Lewis Base - electron rich • Most general of all 3 definitions; acids don’t even need hydrogen! • Sample Problem 18-6, p.447

  43. Lewis Acids/Bases • Several categories of substances can be considered Lewis acids: Electron Poor • 1) positive ions2) having less than a full octet in the valence shell3) polar double bonds (one end)4) expandable valence shells • Several categories of substances can be considered Lewis bases: Electron Rich • 1) negative ions2) one of more unshared pairs in the valence shell3) polar double bonds (the other end)4) the presence of a double bond

  44. Section 18.4Strengths of Acids and Bases • OBJECTIVES: • Define strong acids and weak acids.

  45. Section 18.4Strengths of Acids and Bases • OBJECTIVES: • Calculate an acid dissociation constant (Ka) from concentration and pH measurements.

  46. Section 18.4Strengths of Acids and Bases • OBJECTIVES: • Arrange acids by strength according to their acid dissociation constants (Ka).

  47. Section 18.4Strengths of Acids and Bases • OBJECTIVES: • Arrange bases by strength according to their base dissociation constants (Kb).

  48. 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

  49. Measuring strength • Ionization is reversible. • HA H+ + A- • This makes an equilibrium • Acid dissociation constant = Ka

  50. HA H+ + A- Ka = [H+ ][A- ] [HA] (water is constant so we can ignore it) Stronger acid = more products (ions), thus a larger Ka (Table 18.4, p.448) (Table 18.5 p. 450)

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