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Chapter 11

Chapter 11. Acids, Bases, & Salts Acid/Base Equilibrium. Properties of Acids. sour or tart taste strong acids burn; weak acids feel similar to H 2 O acid solutions are electrolytes acids react with most metals to release H 2 acids cause indicators to change color Acids turn litmus red.

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Chapter 11

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  1. Chapter 11 Acids, Bases, & Salts Acid/Base Equilibrium

  2. Properties of Acids • sour or tart taste • strong acids burn; weak acids feel similar to H2O • acid solutions are electrolytes • acids react with most metals to release H2 • acids cause indicators to change color • Acids turn litmus red

  3. Properties of Bases • bases taste bitter • basic solutions do not burn • basic solutions feel smooth and slippery • basic solutions are also electrolytes • bases usually do not react with metals • bases also cause indicators to change color • Bases turn litmus blue

  4. Three definitions • Arrhenius • Deals with H+ and OH- • Bronsted Lowry • Deals with protons • Lewis • Deals with pairs of electrons

  5. Arrhenius Definition • Arrhenius defined an acid as a substance that ionizes in water to produce H+ ions • Ex. HCl • Any of our acids with H+ in the beginning • Arrhenius defined a base as a substance that ionizes in water to produce OH- ions • Ex. NaOH • Any of the metal hydroxides

  6. Bronsted Lowry Definition • NH3 is known to turn litmus blue, but it does not have hydroxide ion • Needed a new definition • Substances do not need to be in water • A B/L acid is any substance that can donate a proton (H+) • A B/L base is any substance that can accept a proton (H+)

  7. Conjugate pairs HC2H3O2 + H2O → C2H3O2- + H3O+ NH3 + H2O → NH4+ + OH-

  8. Amphoteric Substances • Notice in the examples, that H2O can act as a B/L acid and as a B/L base. • It is know as an amphoteric substance. • Additional examples include • HSO3- • HCO3- • H2PO4- • HPO42- These are the ions that result when polyprotic acids release H+ ions one at a time. Notice that they can accept an H+ or they can release an H+.

  9. Lewis definition • This is the most broad of the definitions • A Lewis acid is an electron pair acceptor. • A Lewis base is an electron pair donor. • Example Lewis acid/base reaction….

  10. Ion product constant for H2O • Remember [H+][OH-] = 1.0 x 10-14 = Kw • So if you know [H+] you can find [OH-] and vice versa • If [H+] = 9.3 x 10-4 M, what is [OH-]

  11. pH or pOH • pH = -log[H+] or pOH = -log[OH-] • pH + pOH = 14 • What is the pH of a solution that contains 4.9 x 10-9 M OH-?

  12. What is the [H+] of a solution that is pH4.82? What is the [OH-] of a solution that is pH8.56?

  13. You must understand that… • When [H+] = [OH-], the solution is neutral, and the pH = 7 • When [H+] > [OH-], the solution is acidic, and the pH <7 • When [H+] < [OH-], the solution is basic, and the pH > 7. • Increasing pH means decreasing [H+]…fewer H+ ions floating in solution….less acidic. • Decreasing pH means increasing [H+]…more H+ ions floating in solution….more acidic.

  14. Weak Acids • When a weak acid is placed in water, a small fraction of its molecules will dissociate into hydrogen ions and conjugate base ions. • Most of the acid molecules will remain in solution as undissociated aqueous particles. • The dissociation constant, Ka, is a measure of the strength of weak acids. • For the reaction…HA ↔ H+ + A- Ka =

  15. Weak Bases • When a weak base is placed in water, a small fraction of its molecules will dissociate into hydroxide ions and conjugate acid ions. • Most of the base molecules will remain in solution as undissociated aqueous particles. • The dissociation constant, Kb, is a measure of the strength of weak bases. • For the reaction…B ↔ HB+ + OH- Kb =

  16. Acid Equilibrium problems • Things to know • All of our acids will be monoprotic, that is, give off only one hydrogen ion. • x will represent the amount of acid that dissociates. • Therefore x also represents the [H+] • We will use ICE to determine concentrations for the Ka expression • We can use HA to represent the acid and A- to represent the conjugate base.

  17. Ka example What is the pH of a 0.20M HC2H3O2, with Ka = 1.8 x 10-5? (also find % ionization)

  18. Polyprotic acids • Polyprotic acids are weak acids (except for H2SO4 which is a strong acid). • They release their H+ ions one at a time. • Each has its own Ka • Ex. H3PO4 H3PO4 → H+ + H2PO4- Ka = 7.5 x 10-3 H2PO4- → H+ + HPO42- Ka = 6.2 x 10-8 HPO42- → H+ + PO43- Ka = 4.8 x 10-13

  19. Kw revisited • Water comes to equilibrium with its ions according to the following reaction… H2O(l) ↔ H+(aq) + OH-(aq) Kw = [H+][OH-] = 1.0 x 10-14 Kw = KaKb pKa + pKb = 14

  20. Strong Acids/Bases Strong Acids Strong Bases HCl LiOH HBr NaOH HI KOH HNO3 RbOH HClO4 CsOH H2SO4 Ba(OH)2 Sr(OH)2

  21. Strong Acids/Base description • Strong acids and bases completely dissociate in water, therefore no Ka or Kb • The dissociations do not reverse. • Oxoacids are acids that contain oxygen. • The greater the number of oxygen atoms attached to the central atom in an oxoacid, the stronger the acid. • That’s because increasing the number of oxygen atoms that are attached to the central atom weakens the attraction that the central atom has for the H+ ion.

  22. Strong Acid/Base calculations • Since these acids and bases completely dissociate in water, the final concentration of H+ ions is the same as the original concentration. • So you can always find the pH of a strong acid solution directly from its concentration. • What is the pH of .20 M HCl? pH = -log(.20) =

  23. Titration • When an acid and a base are mixed, a neutralization reaction occurs. • Acid + base → salt + water • Neutralizations reactions are generally performed by titration, where a base of known concentration is slowly added to an acid (or vice versa) • The progress of a neutralization reaction can be shown in a titration curve. • The equivalence point is the point in the titration when exactly enough base has been added to neutralize all the acid that was initially present. • An indicator will be used to mark the equiv. pt.

  24. Strong acid/strong base titration. • Equivalence point at pH7

  25. Notice the shape of the curve if a strong acid is added to a strong base • The equivalence point is also pH7

  26. Weak acid vs strong base • The equivalence point is > pH7

  27. Weak base vs strong acid • Notice the equivalence point is < pH7

  28. Which indicator to use? • The indicator needs to change color close to the equivalence point.

  29. pH of dissolved salts • If a salt is composed of the conjugates of a strong base and a strong acid, its solution will be neutral. (NaCl) • If a salt is composed of the conjugates of a weak base and a strong acid, its solutions will be acidic. (NH4Cl) • If a salt is composed of the conjugates of a strong base and a weak acid, its solution will be basic. (NaC2H3O2) • If a salt is composed of the conjugates of a weak base and a weak acid, the pH of its solution will depend on the relative strengths of the conjugate acid and base of the specific ions in the salt. (NH4C2H3O2)

  30. Anhydrides • An acid anhydride is a substance that combines with water to form an acid. • Generally nonmetal oxides. • Ex. CO2 + H2O→ H2CO3 • A basic anhydride is a substance that combines with water to form a base. • Generally metal oxides. • Ex. Na2O + H2O→ 2NaOH

  31. What to know for the test? • Find pH from either [H+] or [OH-] • Find [H+] from pH or pOH • Find the # of grams of a base to make a solution of a certain pH • Use titration info to find molar mass acid • Ka problems • Find pH - find Ka from pH • Find concentrations • Find % ionization

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