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Acid-Base Titrations

Acid-Base Titrations . Acid-Base Equilibria. Chapter 16. Review of Acids and Bases. An acid is a substance that when dissolved in water increases the concentration of H + A base is a substance that when dissolved in water increases the concentration of OH -. Bronsted -Lowry Acids and Bases.

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Acid-Base Titrations

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  1. Acid-Base Titrations

  2. Acid-Base Equilibria Chapter 16

  3. Review of Acids and Bases • An acid is a substance that when dissolved in water increases the concentration of H+ • A base is a substance that when dissolved in water increases the concentration of OH-

  4. Bronsted-Lowry Acids and Bases • These types of acid base reactions have to do with the transfer of a proton.

  5. Conjugate Acid/Base pairs

  6. The Relative Strengths of Acids and Bases • We can predict the acid or basic properties of conjugate pairs. • A strong acid produces a conjugate base that will be very unwilling to accept a proton. • A weak acid produces a conjugate base that has some tendency to accept a proton. • A substance that contains hydrogen but has negligible acidity will produce a conjugate base that is a very strong base.

  7. The Autoionization of Water • Calculate the concentration of H+ in a solution in which [OH-] is 0.01 M.

  8. The pH scale • Calculate the [H+] in a sample of freshly squeezed apple juice that has a pH of 3.76.

  9. Strong Acids • The seven most common strong acids are:

  10. Strong Bases • The most common strong bases are the hydroxides of the alkali metals.

  11. Weak Acids • Most acidic substances are weak acids and are therefore only partially ionized.

  12. Calculate Ka from pH • A student prepared a 0.10 M solution of formic acid (HCOOH) and measured its pH. The pH at 25oC was found to be 2.38. Calculate Ka for formic acid.

  13. Calculating pH from Ka

  14. Polyprotic Acids

  15. Weak Bases • Many substances behave as weak bases when mixed with water by removing a hydrogen from the water producing OH-.

  16. Relationship Between Ka and Kb

  17. Acid-Base Properties of Salt Solutions • Solutions made from soluble salts (ionic compounds) can be acidic or basic. • We know these solutions produce anions and cations. • Many of these ions are able to react with water to generate H+(aq) or OH-(aq). These types of reactions are often called hydrolysis.

  18. Anions Reacting with Water • Often times an anion is the conjugate base of an acid: • If this anion is the conjugate base of a strong acid it will have negligible tendency to remove a proton from water.

  19. Cations Reacting with Water • Polyatomic cations whose formulas contain one or more protons are usually the conjugate acids of weak bases.

  20. Combined Effect of Cations and Anions in Soluiton • If a solution contains two different ions that can effect the pH differently.

  21. Summary • An anion that is the conjugate base of a strong acid (example Br-) will not affect the pH of a solution. • An anion that is the conjugate base of a weak acid (example CN-) will cause and increase in pH. • An cation that is the conjugate acid of a weak base (example CH3NH3+) will cause a decrease in pH. • Cations of group 1A metals have no affect on pH • other metals will cause a decrease in pH • when a solution contains both the conjugate base of a weak acid and the conjugate acid of a weak base the ion with the larger equilibrium constant will have the greater influence on pH.

  22. Acid-Base Behavior and Chemical Structure • What makes an acid or a base strong or weak? • It’s chemical structure.

  23. Binary Acids • Binary acids are ones that contain only 2 atoms (one being H). • The strength of binary acids where elements X are in the same group decreases as the size of X increases.

  24. Oxyacids • Oxyacids are ones where OH groups are attached to a central atom. • For oxyacids that have the same number of OH groups and the same number of O atoms, acid strength increases with increasing electronegativity of the central atom. • For oxyacids that have the same central atom, acid strength increases as the number of oxygen atoms attached to that central atom increases.

  25. Carboxylic Acids

  26. Lewis Acids and Bases • A Lewis base is a substance that has a free pair of electrons. • A Lewis acid is a substance that can accept that free pair of electrons

  27. Additional Aspects of Aqueous Equilibria Chapter 17

  28. The Common Ion Effect • The common ion effect describes what happens when we add a strong electrolyte to a weak electrolyte solution and both compound share a common ion

  29. Example • What is the pH of a solution made by adding 0.3 mol of acetic acid and 0.3 moles of sodium acetate to enough water to make 1.0 L of solution? (Ka = 1.8 x 10-5)

  30. Buffer Solutions • Buffer solutions are ones that resist drastic pH changes when small amounts of strong acids or bases are added. • These solutions contain significant enough amounts of H+ and OH- that they are able to neutralize any acid or base we might add. • Buffers are usually prepared by mixing a weak acid or base with a salt of that acid or base.

  31. Calculating the pH of a Buffer • What is the pH of a buffer that is 0.12 M lactic acid (HC3H5O3) and 0.1 M sodium lactate (NaC3H5O)? (Ka = 1.4x10-4)

  32. Buffer Capacity and pH Range • Buffer capacity refers to the amount of strong acid or strong base that can be added before the pH begins to change by an appreciable amount. • pH range refers to the range of pH over which the buffer best resists pH change.

  33. Addition of Strong Acids or Bases to Buffers • Adding a strong acid or a strong base cause opposite neutralization reactions to occur in the buffer.

  34. Example • A buffer is made by adding 0.3 mol of acetic acid and 0.3 mol of sodium acetate to enough water to make 1 L of solution. the pH of the buffer is 4.74. (a) Calculate the pH of the solution after 0.02 mol of NaOH are added. (Assume that volume remains constant.)

  35. Titration Calculations • Calculate the pH when the following quantities of 0.1 M NaOH solution have been added to 50.0 mL of 0.1 M HCl • (a) 49.0 mL • (b) 51.0 mL

  36. Weak Acid-Strong Base Titrations • Consider titrating 50.0 mL of 0.1 M acetic acid with 0.1 M NaOH.

  37. Example • Calculate the pH of a solution that is formed when 45.0 mL of 0.1 M NaOH is added to 50.0 mL of 0.1 M CH3COOH. (Ka = 1.8 x 10-5)

  38. Effect of Ka on Titration Curves

  39. Titrations of Polyprotic Acids

  40. Solubility Equilibria • We have seen that some ionic compounds are not soluble in water. • This is only partially true.

  41. Factors that Affect Solubility • Common-Ion Effect

  42. Example • Calculate the solubility of CaF2 at 25oC in 0.01M Ca(NO3)2. (Ksp = 3.9 x 10-11).

  43. Solubility and pH • The pH of a solution will affect the solubility of any substance whose anion is basic. • If a compound contains an anion that is the conjugate base of a weak acid its solubility will increase as the solution becomes more acidic.

  44. Formation of Complex Ions • Metal ions can act as Lewis bases in the presence of a Lewis acid.

  45. Precipitation and Separation of Ions • The equilibrium of BaSO4(s) in water can be achieved in either direction:

  46. Selective Precipitation of ions • Ions can be precipitated out of a solution based on the solubilities of their salts.

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