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Acid -Base Chemistry An acid is a H + (proton) donor . General formula of acid = H-A .

Acid -Base Chemistry An acid is a H + (proton) donor . General formula of acid = H-A . H-A H + + A - acid proton conjugate base. When H-A dissociates to H + and A – in aqueous (water) solution, the “hydronium ion” or H 3 O + forms :

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Acid -Base Chemistry An acid is a H + (proton) donor . General formula of acid = H-A .

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  1. Acid-Base Chemistry An acid is a H+ (proton) donor. General formula of acid = H-A. H-A H+ + A- acid proton conjugate base

  2. When H-A dissociates to H+ and A– in aqueous (water) solution, the “hydronium ion” or H3O+ forms: H-A+ H2O H3O+ + A– means •• •• H-A + H-O-H H-O-H + A– + •• H

  3. A Base is a H+(proton) acceptor. A base has an unshared pair of electrons to share with H+. we can represent a base by the general formula :B or :Base. :Base + H+ H-Base+ from an conjugate acid acid

  4. When a Base accepts protons from water, it increases the concentration of HO– (hydroxide ions) in solution: :B + H-O-H H-B+ + – O-H OR :B + H2O HB+ +–OH •• •• •• •• ••

  5. Examples HNO3 + H2O H3O+ + NO3 – nitric acid nitrate ion acidconjugate base HCl + H2O H3O+ + Cl– hydrochloric acid chloride ion acidconjugate base

  6. :NH3+ H2O +NH4 + – OH ammonia ammonium ion baseconjugate acid

  7. Strengths of Acids & Bases A strong acid HA completely dissociates to H+ (= H3O+) and A– in H2O: H-A + H2O H3O+ + A– . 1 mole H3O+ + 1 mole A– 1 mole to start gives

  8. The common strong acids are: HCl hydrochloric acid HBr hydrobromic acid HI hydroiodic acid HNO3 nitric acid H2SO4 sulfuric acid [HClO4 perchloric acid]don’t memorize [ ]

  9. A strong base is one which completely dissociates to provide HO– in water. These bases are ionic metal hydroxides: Metal cation + hydroxide anion MOH M+ + –OH 1 mole 1 mole + 1 mole

  10. There are 4 common strong bases: LiOH lithium hydroxide NaOH sodium hydroxide KOH potassium hydroxide Ba(OH)2 barium hydroxide

  11. Some Advice... It is best to memorizethelistsof strong acids & strong bases -- assume all others are weak.

  12. But Also, LEARN HOW to write the equations for the reaction of each chemical as a strong acid or strong base.

  13. What about WEAK acids & bases? A weak acid or weak base is incompletely ionized or dissociated in water solution. The ionization reaction of a weak acid or weak base is a reversible equilibrium!!

  14. Example CH3CO-H + H2O CH3CO– + H3O+ O O acetic acid acetate before equilibrium: 1 mole 0 0 after equilibrium: 0.98 mole 0.02 0.02

  15. The position of equilibrium is different for each weak acid and weak base. The equilibrium constant describes how much ionization has occurred.

  16. For acetic acid: O O CH3CO-H + H2O CH3CO– + H3O+ [CH3CO2–]•[H3O+] OR Keq = [CH3CO2H]•[H2O] [CH3CO2–]•[H3O+] Keq•[H2O] = Ka= [CH3CO2H]

  17. We use Ka as a measure of the strength of an acid because [H2O] is very large and does not effectively change as a result of the ionization reaction. Typical Ka values for weak acids are in the range of 1 x 10-1 to 1 x 10-10

  18. Sample Ka values: do NOT memorize!!!! WhatKa= CH3CO2H acetic acid 1.8x10-5 H2O water 1 x 10-16 CH3CH2OH alcohol 1 x 10-17

  19. What do these small Ka values mean? * Solutions of weak acids and bases at equilibrium contain both the unionized (H-A) and ionized forms (H+ and A–). * The largest species present is H-A (unionized acid.)

  20. And... * Any compound having Ka smaller than H2O is considered “neutral” = “not acidic in water”. Thus, alcohols are not acidic compounds. They are neutral substances.

  21. Monoprotic acids are those which can release one H+ per molecule of acid. Polyprotic acids can release more than one H+per molecule of acid.

  22. Examples Monoprotic HCl Hydrochloric Acid DiproticH2SO4 Sulfuric acid TriproticH3PO4 Phosphoric Acid

  23. Amphiprotic (Amphoteric) Compounds can act as acids or bases. Examples HCO3– bicarbonate ion as acid HCO3– CO3-2 + H+ as base HCO3– + H+H2CO3

  24. Dihydrogen phosphate H2PO4–is another important amphoteric ion. as acid H2PO4–HPO4-2 + H+ as base H2PO4–+ H+H3PO4

  25. Water is the most important amphoteric compound of all. H-O-H + H-O-H HO– + H3O+ “acid” “base” Keq = [HO–]•[H3O+] OR [H2O]•[H2O] Keq•[H2O]2 = Kw = [HO–]•[H3O+]

  26. In pure H2O at room temperature, Kw has a value of 1 x 10-14. Since Kw = [HO–]•[H3O+], this means [HO–]=[H3O+] = 1x10-7 moles/liter in neutral water.

  27. pH is defined as pH = -log ([H3O+]) of a solution. A logarithm tells what power of 10 = a certain number. That is, if y = 10x, then log(y) = x. Thus, 100 = 102 , so log(100) = 2.

  28. We saw that for pure H2O, [H3O+] = 1 x 10-7. Since -log (1 x 10-7) = 7, we say neutral water has pH = 7.

  29. Recall: [HO–] and [H3O+] are related through the equilibrium 2 H2O HO– + H3O+ . Acids are H+ donors and cause an increase in [H3O+]. So: pH decreases. Also: [HO–] decreases. Acidic solutions have pH <7. (in the range 0 to 6.99)

  30. Again: 2 H2O HO– + H3O+ . A Base increases [HO-] or decreases [H3O+]. So pH increases: a basic or “alkaline” solution has a pH >7 (in the range 7+ to 14).

  31. 14 Basic [HO–] > [H3O+] • Neutral [HO–] = [H3O+] pH Acidic [HO–] < [H3O+] 0

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