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Acids & Bases
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  1. Acids & Bases

  2. Properties of Acids • Taste sour • Contain H+ ion • pH less than 7 • React with bases to form a salt and water • React with some metals to produce hydrogen gas

  3. Properties of Acids • Turn litmus paper red • Phenolphthalein is colorless in the presence of an acid • Bromothymol blue is yellow in the presence of an acid • Found in citrus fruits in the form of citric acid

  4. Properties of Acids • Found in soured milk and in sore muscles in the form of lactic acid • Found in vitamin C in the form of ascorbic acid • Found in carbonated beverages in the form of carbonic acid…that’s also what you exhale

  5. Properties of Bases • Taste bitter • Contain OH- ion • pH greater than 7 • React with acids to form a salt and water • React with organic material

  6. Properties of Bases • Feel slippery because they immediately begin to dissolve the outer layer of skin tissue • Turn litmus paper blue • Phenolphthalein is fuchsia in the presence of a base • Bromothymol blue is blue in the presence of a base

  7. Properties of Bases • Found in drain cleaners usually in the form of sodium hydroxide • Found in ammonia-based cleaners like Windex • Lye (NaOH) is used to make soaps

  8. Classifying Acids and Bases • Svante Arrhenius—Swedish guy who put forth his definitions of acids and bases in 1884 at the age of 25 • Worked with our buddy van’t Hoff • Received 1903 Nobel Chemistry Prize for electrolytic dissociation discoveries

  9. Arrhenius Acids • are substances that will dissociate in water to yield hydrogen ion (H+)

  10. Arrhenius Bases • are substances that will dissociate in water to yield hydroxide ion (OH-)

  11. Classifying Acids and Bases • Johannes Brønsted (Danish) and Thomas Lowry (English) came up with a new way to classify acids and bases and their conjugates (pairs that have features in common but are opposites) in the 1920’s • never received a Nobel for furthering these acid/base concepts, and Arrhenius never accepted them!

  12. Brønsted-Lowry Acid • A reactant that donates a proton in a chemical reaction • The proton is actually the hydrogen ion…since a hydrogen atom has 1 proton and 1 electron and the ion with a 1+ charge indicates that it has lost an electron

  13. Brønsted-Lowry Base • A reactant that accepts a proton in a chemical reaction

  14. Brønsted-Lowry Conjugate Acid • A product that • is formed when a base accepts a proton • in the reverse reaction, will donate a proton

  15. Brønsted-Lowry Conjugate Base • A product that • is formed when an acid donates a proton (what’s left after the donation occurs) • in the reverse reaction, will accept a proton

  16. Classifying Acids and Bases • Around the same time that Brønsted and Lowry were devising their acid/base scheme, our buddy Gilbert Lewis (yep, the same guy who did the dots) came up with yet another method of classifying them…it’s a broader method than Arrhenius, Brønsted, or Lowry ever postulated

  17. Lewis Acid • A reactant that accepts an electron pair

  18. Lewis Base • A reactant that donates an electron pair

  19. Example #1 HCl (aq) + H2O (l)  H3O+ (aq) + Cl- (aq) Or HCl (aq)  H+ (aq) + Cl- (aq)

  20. Example #1 • HCl is an acid • It dissociates to yield H3O+ (hydronium ion), which is really water with an extra H+. (Arrhenius) • It donates a proton (H+) to water in the first reaction written. (Brønsted-Lowry)

  21. Example #1 • H2O is an base • It accepts a proton (H+) from HCl in the first reaction written. (Brønsted-Lowry)

  22. Example #1 • H3O+ is a conjugate acid • It is produced when the water accepts a proton (H+) from HCl in the first reaction written. (Brønsted-Lowry) • In the reverse reaction, it will donate a proton (H+) to Cl- in the first reaction written. (Brønsted-Lowry)

  23. Example #1 • Cl- is a conjugate base • It is produced when the HCl donates a proton (H+) to water in the first reaction written. (Brønsted-Lowry) • In the reverse reaction, it will accept a proton (H+) from H3O+ in the first reaction written. (Brønsted-Lowry)

  24. Example #2 NH3 (g) + H2O (l)  NH4+ (aq) + OH- (aq)

  25. Example #2 • NH3 is a base • It accepts a proton (H+) from water. (Brønsted-Lowry)

  26. Example #2 • H2O is an acid • It donates a proton (H+) to ammonia. (Brønsted-Lowry)

  27. Example #2 • NH4+ is a conjugate acid • It is produced when the ammonia accepts a proton (H+) from water. (Brønsted-Lowry) • In the reverse reaction, it will donate a proton (H+) to OH-. (Brønsted-Lowry)

  28. Example #2 • OH- is a conjugate base • It is produced when the water donates a proton (H+) to ammonia. (Brønsted-Lowry) • In the reverse reaction, it will accept a proton (H+) from NH4+. (Brønsted-Lowry)

  29. Water—our special friend • Did you notice that it behaved as a base in the first example and as an acid in the second example? • A substance that can behave as either an acid or a base is called, amphoteric or amphiprotic.

  30. Example #3 H H—N—H + H+ H—N—H H H

  31. Example #3 • NH3 is a base • It donates a pair of electrons to H+. (Lewis)

  32. Example #3 • H+ is an acid • It accepts a pair of electrons from NH3. (Lewis)

  33. Autoionization of Water • Every acid and base will dissociate in water…even water (since it’s amphoteric)! 2H2O (l)  H3O+ (aq) + OH- (aq) or H2O (l)  H+ (aq) + OH- (aq)

  34. Autoionization of Water • Usually, the 2nd reaction is the one we will use since H3O+ is just water with an extra H+. • Write the K expression for the 2nd reaction, keeping in mind that we only include gaseous and aqueous phases.

  35. Autoionization of Water • K = [H+][OH-] • note that water is not included because it is a liquid • This expression is known as the Kw, or equilibrium constant for water, expression • K w = [H+][OH-]

  36. Autoionization of Water • The value of Kw is 1 x 10-14 M2 at 25°C. • This is a small K value. • If the temperature changes, so does the value of Kw

  37. Autoionization of Water • Make an equilibrium chart for the dissociation, or autoionization, of water.

  38. Autoionization of Water

  39. Autoionization of Water

  40. Autoionization of Water

  41. Autoionization of Water • Determine the equlibrium concentrations of both the hydrogen ion and the hydroxide ion by plugging into the Kw expression. 1 x 10-14 M2 = [x][x]

  42. Autoionization of Water 1 x 10-14 M2 = x2 1 x 10-7 M = x [H+] = 1 x 10-7 M [OH-] = 1 x 10-7 M

  43. Autoionization of Water • Because the H+ and OH- concentrations are equal, the solution is neutral. • If [H+] > [OH-], then the solution is an acid. • If [H+] < [OH-], then the solution is a base.

  44. Autoionization of Water • Since every acid or base dissociation we will entertain occurs in water, then the Kw expression is applicable to any of these dissociations.

  45. Autoionization of Water • Thus, if you know the [H+] concentration of a solution, you can determine the [OH-] concentration. • And, if you know the [OH-] concentration of a solution, you can determine the [H+] concentration.

  46. pH • Represents the “power of hydrogen” • Calculated by taking the opposite of the logarithm of the hydrogen ion concentration… pH = -log [H+]

  47. pH • Calculate the pH of water at 25°C knowing that the [H+] is 1 x 10-7 M. pH = -log [1 x 10-7] pH = 7

  48. pH • If you already know the pH of a solution, then you can find the [H+] using: [H+] = 10-pH So, [H+] = 10-7 [H+] = 1 x 10-7 M

  49. pOH • Represents the “power of hydroxide” • Calculated by taking the opposite of the logarithm of the hydroxide ion concentration… pOH = -log [OH-]

  50. pOH • Calculate the pOH of water at 25°C knowing that the [OH-] is 1 x 10-7 M. pOH = -log [1 x 10-7] pOH = 7