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Acid-Base Equilibria

Acid-Base Equilibria. Arrhenius. acids increase [H + ] when dissolved in water acids can be classified as monoprotic, diprotic or triprotic bases increase [OH - ] when dissolved in water bases can be classified as monobasic, dibasic, or tribasic. A-B Strength.

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Acid-Base Equilibria

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  1. Acid-Base Equilibria

  2. Arrhenius • acids increase [H+] when dissolved in water • acids can be classified as monoprotic, diprotic or triprotic • bases increase [OH-] when dissolved in water • bases can be classified as monobasic, dibasic, or tribasic

  3. A-B Strength • strong acids & bases ionize completely and are strong electrolytes • 7 acids & 8 bases (memorize them!!!) • diprotic acids & dibasic bases do NOT ionize completely, only their first H+ or OH- ionizes completely. • Strong acid + strong base = neutral salt • strong A-B are not equilibrium expressions, but all other A-B are reversible

  4. Bronsted-Lowry A-B • restricts definition to H+ • acids donate H+ • bases accept H+ • allows the classification of less traditional A-B • conjugate A-B pairs = 2 formulas in an equation whose formulas differ by a H+

  5. What is the acid, base, and the conjugates? HClO +H2O H3O+ + ClO- CO32- + H2O OH- + HCO3-

  6. Wait, water can go both ways? • amphoteric substances can behave as either an acid or base depending on what they react with. • water and anions with protons (H+) attached are most common amphoterics

  7. Weak A-B • only partially ionize in water and are weak electrolytes • can be written as equilibrium expressions with a Ka or Kb value • K value indicates how much the acid or base will ionize (high K = higher ionization) • larger K values indicate a stronger acid or base • For di- and tri- protic/basic, there will be 2 K values (one for the first ionization and one for the second)

  8. Autoionization of Water H2O + H2O OH- + H3O+ • reversible equilibrium where water can donate a proton to itself • Kw = 1.0 x 10-14 at room temp. • For any conjugate A-B pair, Kw = Ka x Kb • What is the Ka value for NH4+?

  9. Example Is an aqueous solution of Na2HPO4 acidic or basic?

  10. pH scale pH = -log[H+] • works for pH ranges from 2-12 and approximates pH outside that window • The exponent on the [H+] is an indicator of approximate pH.

  11. Strong A-B • no equilibrium b/c all acid/base ionizes • use original acid concentration to calculate pH • Calculate the [H+] and pH in a solution of 0.37M hydrochloric acid. • Calculate the [OH-] and pH in a 0.58M solution of NaOH

  12. Weak Acids • use RICE to find equilibrium concentrations: R HA + H2O H3O+ + A- I I 0 0 C -x +x +x E I-x x x • b/c Ka for most weak acids is less than 10-3, I-x is about equal to I, so Ka = x2/I

  13. Example What is the pH of a 0.20M solution of acetic acid (CH3COOH)?

  14. Weak Bases • similar calculations as acids (replace H3O+ with OH-) R B + H2O OH- + HB+ I I 0 0 C -y +y +y E I-y y y • b/c Kb for most weak bases is less than 10-3, I-y is about equal to I, so Kb = y2/I

  15. Example What is the pH of a 0.68M solution of aqueous ammonia?

  16. Classify the following as weak/strong acids/bases: • chloric acid • ammonium chloride • calcium hydroxide • ethyl amine • sodium cyanide

  17. Example Measurements show that the pH of a 0.10M solution of acetic acid is 2.87. What is the Kb of potassium acetate?

  18. A-B properties of salt solutions • for the most part anions are slightly basic (because they attract protons) and cations are slightly acidic (because they can donate protons) • ions from strong A-B are the only neutral ions • To determine if a salt is acidic or basic, look at the ions it forms: • ignore any neutral ions • if anion is left, salt is basic • if cation is left, salt is acidic • if both cation & anion are neutral, salt is neutral • if both cation & anion are not neutral, the A-B-ness can’t be determined from the formula

  19. Classify the following salts as acidic, basic, or neutral: • NaNO2 • CH3NH3Cl • NaCl • MgSO4 • Al2(SO3)3

  20. A-B-ness & chemical structure 3 factors affect attraction of electrons (acidity increases with stronger attraction of electrons) • ionic charge – when comparing similar atoms, more positive ions are stronger acids • oxidation # on central atom – when comparing similar formulas with the same central atom, the higher the ox#, the stronger the acid • electronegativity – when comparing similar formulas with different central atoms, the higher the EN, the stronger the acid

  21. Common Base Reactions • Strong bases also include hydrides (H-), nitrides (N3-), and carbides (C22-) • NaH + H2O  H2 + Na+ + OH- • Mg3N2 + 6H2O  2NH3 + 3Mg2+ + 6OH- • Ca2C2 + 2H2O  C2H2 + Ca2+ + 2OH- • strong bases also include oxides of groups 1&2 metals • Li2O + H2O  2Li+ + 2OH- • CaO + H2O  Ca2+ + 2OH-

  22. Common Acid Reactions • nonmetal oxides (aka. acid anhydrides) turn into acids when placed in water • SO2 + H2O H2SO3 • CO2 + H2O H2CO3 • Cl2O7 + H2O  2H+ + 2ClO4-

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