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COMMON ION EFFECT. COMMON ION an ion common with one in a system at equilibrium which places a stress on the equilibrium Common Ion Common Ion.

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Presentation on theme: "COMMON ION EFFECT. COMMON ION an ion common with one in a system at equilibrium which places a stress on the equilibrium Common Ion Common Ion."— Presentation transcript:

1 COMMON ION EFFECT

2 COMMON ION an ion common with one in a system at equilibrium which places a stress on the equilibrium Common Ion Common Ion

3 Uses of Common Ion Effect 1. control pH of a weak acid or base 2. control formation of a precipitate

4 BUFFER Example Non-example Example Non-example Example Non-example A solution which resists a change in pH when an acid or base is added consists of a weak acid or base and a salt containing a common ion of its conjugate

5 How does LeChateliers Principle explain the operation of a buffer?

6 Example of a buffer system CH 3 COOH + HOH CH 3 COO - + H 3 O + NaCH 3 COO (aq) Na + + CH 3 COO -

7 Characteristics of a Good Buffer

8 1. operates over a narrow pH range (< 1 pH unit) 2. no reactions between buffers in a multiple buffer system 3. range can be extended using more than one buffer

9 Henderson-Hasselbalch Equation

10 Maximum buffering will occur when ratio is close to 1, or when pH = pK a

11 1. What is the pH of a 0.20 M acetic acid solution?

12 Add 10.0 mL of 0.20 M NaOH to 50.0 mL of the preceding solution. What is the pH?

13 Add 5.0g sodium acetate (MM 82.05) to 500. mL of the 0.20 M acetic acid solution. What is the pH?

14 Add 10.0 mL of 0.20 M NaOH to 50.0 mL of the preceding solution. What is the pH?

15 2. Calculate the mass of ammonium chloride (MM 43.6) needed to buffer 250. mL of 2.0 M ammonia to a pH of 10.

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17 TITRATION CURVES

18 Titration Curve A graphical history of a titration typically a plot of the pH (dependent variable) and volume titrant (independent variable)

19 Uses of Titration Curves 1. determine equivalence point 2. determine number of ionization reactions 3. determine optimum buffer region 4. determine possible indicators

20 Shape of Titration Curve Strong acid - strong base Weak acid - strong base

21 Shape of Titration Curve

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23 Equivalence Point 1. Midpoint between points of inflection 2. Plot of the slope of each point of the curve against volume titrant ( pH/ V vs V avg )

24 Number of Ionization Reactions CH 3 COOH - NaOH H 2 C 2 O 4 - NaOH

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27 Optimum Buffer Region Area where the concentration of molecules and their conjugate ions are relatively high

28 Indicators Need to choose for each titration system Dependent on pH at equivalence point

29 ACID-BASE INDICATORS

30 Acid-base indicators are weak Bronsted- Lowry compounds that are different colors in acid and base form.

31 Acid-base indicators are all large organic molecules. HIn H + + In - Color 1 Color 2

32 Phenolphthalein Colorless acid form, HIn

33 Phenolphthalein Pink base form, In -

34 The color change occurs at a different pH for different indicators. The pH at which the indicator changes color is dependent on the K a of the indicator as a weak acid.

35 HIn H + + In -

36 Experiments have shown that the minimum amount of change of HIn In - that can be detected visually is

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38 Thus, from the Henderson-Hasselbalch equation, one can select an appropriate indicator for a titration based upon the K a of the indicator and the pH at the equivalence point.

39 What is the pH at the equivalence point of a titration of 25.0 mL each of 0.10 M HCl and 0.10 M NaOH?

40 What is the pH at the equivalence point of a titration of 25.0 mL each of 0.10 M CH 3 COOH and 0.10 M NaOH?

41 Phenolphthalein K a = 1 x pH of perceptible color change?

42 SOLUBILITY EQUILIBRIA

43 Saturated Solution Maximum amount of solute dissolved in a specific volume of solvent at a specific temperature

44 Saturated Solution Equilibrium is established between a solid solute and ions from the solute

45 Super-Saturated Solution More than the normal maximum amount of solute is dissolved in a solution.

46 Question at a constant temperature, what is the difference in concentration of a saturated solution: (1 mL vs 1 ML solution) (1 mg vs 1 kg solid)

47 The concentration of a saturated solution remains the same, no matter how much solid is present, as long as the temperature remains constant.

48 The concentration of a solid remains the same at a constant temperature.

49 By convention, equations for the formation of saturated solutions are written in the format solid solution AgCl(s) Ag + + Cl -

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52 3. What is the solubility of silver chloride in water at 25 o C? (K sp = 1.6 x )

53 4. What is the solubility of lead(II) bromide at 25 o C? (K sp = 4.6 x )

54 6. What mass of nickel is dissolved in 100. mL of saturated nickel(II) hydroxide? (K sp = 1.6 x ) What is the pH of this solution?

55 Which is more soluble? Ag 2 CO 3 [K sp = 8.5 x ] or CaCO 3 [K sp = 3.4 x ]

56 SOLUBILITY ---- ACIDITY ---- PRECIPITATION

57 8. If gram of magnesium hydroxide (MM 58.3) is added to 1.00L of water, will it all dissolve? (K sp = 8.9 x ) Below what pH would the solution be buffered so that it does all dissolve?

58 9. Calculate the concentration of NH 4 + from ammonium chloride required to prevent the precipitation of Ca(OH) 2 in a liter of solution that contains 0.10 mole of ammonia and 0.10 mole of calcium ion.

59 10.If 50. mL of 0.012M barium chloride are mixed with 25 mL of 1.0 x M sulfuric acid, will a precipitate form? HINT: use the concentration quotient Q as we used it before

60 11.You have a aqueous solution of Zn 2+ and Pb 2+ both M. Both form insoluble sulfides. Approximately what pH will allow maximum precipitation of one ion and leave the other in solution? [K sp ZnS = 2.5 x ] [K sp PbS = 7 x ]

61 SOLUBILITY ---- COMMON IONS ---- COMPLEX IONS

62 12. Calculate the molar solubility of silver thiocyanate, AgSCN, in pure water and in 0.010M NaSCN.

63 Complex Ion A charged species consisting of a metal ion surrounded by ligands

64 LIGAND An ion or molecule, acting as a Lewis base, attached to the central metal ion using the d-orbitals of the metal

65 Coordination Number The number of ligands attached to the central metal ion. 2, 4, or 6 are most common CN

66 Metal ions add ligands one step at a time. Ag + + NH 3 Ag(NH 3 ) + K f1 = 2.1 x 10 3 K f1 = 2.1 x 10 3 Ag(NH 3 ) + + NH 3 Ag(NH 3 ) 2 + K f2 = 8.2 x 10 3 K f2 = 8.2 x 10 3 where K f = formation constant

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68 You need to familiarize yourself with typical complex ions, Appendix K

69 Note that a formation constant reflects the stability of the complex.

70 13. Calculate the equilibrium constant for AgI(s) + 2NH 3 (aq) [Ag(NH 3 ) 2 ] + (aq) + I - (aq)

71 14. Will 5.0 mL of 2.5 M NH 3 dissolve mole AgCl?

72 15.A solution is prepared by adding 0.10 mole Ni(NH 3 ) 6 Cl 2 to 0.50 L of 3.0 M NH 3. Calculate the [Ni(NH 3 ) 6 2+ ] and [Ni 2+ ] in the solution.


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