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Chapter 16 Aqueous Ionic Equilibrium 2008, Prentice Hall Chemistry: A Molecular Approach, 1 st Ed. Nivaldo Tro Roy Kennedy Massachusetts Bay Community.

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Presentation on theme: "Chapter 16 Aqueous Ionic Equilibrium 2008, Prentice Hall Chemistry: A Molecular Approach, 1 st Ed. Nivaldo Tro Roy Kennedy Massachusetts Bay Community."— Presentation transcript:

1 Chapter 16 Aqueous Ionic Equilibrium 2008, Prentice Hall Chemistry: A Molecular Approach, 1 st Ed. Nivaldo Tro Roy Kennedy Massachusetts Bay Community College Wellesley Hills, MA

2 2 The Danger of Antifreeze each year, thousands of pets and wildlife die from consuming antifreeze most brands of antifreeze contain ethylene glycol sweet taste initial effect drunkenness metabolized in the liver to glycolic acid HOCH 2 COOH if present in high enough concentration in the bloodstream, it overwhelms the buffering ability of HCO 3, causing the blood pH to drop when the blood pH is low, it ability to carry O 2 is compromised acidosis the treatment is to give the patient ethyl alcohol, which has a higher affinity for the enzyme that catalyzes the metabolism of ethylene glycol ethylene glycol (aka 1,2–ethandiol)

3 Tro, Chemistry: A Molecular Approach3 Buffers buffers are solutions that resist changes in pH when an acid or base is added they act by neutralizing the added acid or base but just like everything else, there is a limit to what they can do, eventually the pH changes many buffers are made by mixing a solution of a weak acid with a solution of soluble salt containing its conjugate base anion

4 Tro, Chemistry: A Molecular Approach4 Making an Acid Buffer

5 Tro, Chemistry: A Molecular Approach5 How Acid Buffers Work HA (aq) + H 2 O (l) A (aq) + H 3 O + (aq) buffers work by applying Le Châteliers Principle to weak acid equilibrium buffer solutions contain significant amounts of the weak acid molecules, HA – these molecules react with added base to neutralize it you can also think of the H 3 O + combining with the OH to make H 2 O; the H 3 O + is then replaced by the shifting equilibrium the buffer solutions also contain significant amounts of the conjugate base anion, A - these ions combine with added acid to make more HA and keep the H 3 O + constant

6 Tro, Chemistry: A Molecular Approach6 H2OH2O How Buffers Work HA + H3O+H3O+ A A Added H 3 O + new HA

7 Tro, Chemistry: A Molecular Approach7 H2OH2O HA How Buffers Work HA + H3O+H3O+ A Added HO new A A

8 Tro, Chemistry: A Molecular Approach8 Common Ion Effect HA (aq) + H 2 O (l) A (aq) + H 3 O + (aq) adding a salt containing the anion, NaA, that is the conjugate base of the acid (the common ion) shifts the position of equilibrium to the left this causes the pH to be higher than the pH of the acid solution lowering the H 3 O + ion concentration

9 Tro, Chemistry: A Molecular Approach9 Common Ion Effect

10 Tro, Chemistry: A Molecular Approach10 Ex What is the pH of a buffer that is M HC 2 H 3 O 2 and M NaC 2 H 3 O 2 ? Write the reaction for the acid with water Construct an ICE table for the reaction Enter the initial concentrations – assuming the [H 3 O + ] from water is 0 HC 2 H 3 O 2 + H 2 O C 2 H 3 O 2 + H 3 O + [HA][A - ][H 3 O + ] initial change equilibrium

11 Tro, Chemistry: A Molecular Approach11 [HA][A - ][H 3 O + ] initial change equilibrium Ex What is the pH of a buffer that is M HC 2 H 3 O 2 and M NaC 2 H 3 O 2 ? represent the change in the concentrations in terms of x sum the columns to find the equilibrium concentrations in terms of x substitute into the equilibrium constant expression +x+x+x+x x x x x HC 2 H 3 O 2 + H 2 O C 2 H 3 O 2 + H 3 O +

12 Tro, Chemistry: A Molecular Approach12 determine the value of K a since K a is very small, approximate the [HA] eq = [HA] init and [A ] eq = [A ] init solve for x [HA][A - ][H 3 O + ] initial change -x-x+x+x+x+x equilibrium x x Ex What is the pH of a buffer that is M HC 2 H 3 O 2 and M NaC 2 H 3 O 2 ? x K a for HC 2 H 3 O 2 = 1.8 x 10 -5

13 Tro, Chemistry: A Molecular Approach13 Ex What is the pH of a buffer that is M HC 2 H 3 O 2 and M NaC 2 H 3 O 2 ? K a for HC 2 H 3 O 2 = 1.8 x check if the approximation is valid by seeing if x < 5% of [HC 2 H 3 O 2 ] init the approximation is valid x = 1.8 x [HA][A - ][H 3 O + ] initial change -x-x+x+x+x+x equilibrium x

14 Tro, Chemistry: A Molecular Approach14 Ex What is the pH of a buffer that is M HC 2 H 3 O 2 and M NaC 2 H 3 O 2 ? x = 1.8 x substitute x into the equilibrium concentration definitions and solve [HA][A - ][H 3 O + ] initial change -x-x+x+x+x+x equilibrium E x x x

15 Tro, Chemistry: A Molecular Approach15 Ex What is the pH of a buffer that is M HC 2 H 3 O 2 and M NaC 2 H 3 O 2 ? substitute [H 3 O + ] into the formula for pH and solve [HA][A - ][H 3 O + ] initial change -x-x+x+x+x+x equilibrium E-5

16 Tro, Chemistry: A Molecular Approach16 Ex What is the pH of a buffer that is M HC 2 H 3 O 2 and M NaC 2 H 3 O 2 ? check by substituting the equilibrium concentrations back into the equilibrium constant expression and comparing the calculated K a to the given K a the values match [HA][A - ][H 3 O + ] initial change -x-x+x+x+x+x equilibrium E-5 K a for HC 2 H 3 O 2 = 1.8 x 10 -5

17 Tro, Chemistry: A Molecular Approach17 Practice - What is the pH of a buffer that is 0.14 M HF (pK a = 3.15) and M KF?

18 Tro, Chemistry: A Molecular Approach18 Practice - What is the pH of a buffer that is 0.14 M HF (pK a = 3.15) and M KF? Write the reaction for the acid with water Construct an ICE table for the reaction Enter the initial concentrations – assuming the [H 3 O + ] from water is 0 HF + H 2 O F + H 3 O + [HA][A - ][H 3 O + ] initial change equilibrium

19 Tro, Chemistry: A Molecular Approach19 [HA][A - ][H 3 O + ] initial change equilibrium Practice - What is the pH of a buffer that is 0.14 M HF (pK a = 3.15) and M KF? represent the change in the concentrations in terms of x sum the columns to find the equilibrium concentrations in terms of x substitute into the equilibrium constant expression +x+x+x+x x 0.14 x x x HF + H 2 O F + H 3 O +

20 Tro, Chemistry: A Molecular Approach20 determine the value of K a since K a is very small, approximate the [HA] eq = [HA] init and [A ] eq = [A ] init solve for x [HA][A - ][H 3 O + ] initial change -x-x+x+x+x+x equilibrium x 0.14 x Practice - What is the pH of a buffer that is 0.14 M HF (pK a = 3.15) and M KF? x K a for HF = 7.0 x 10 -4

21 Tro, Chemistry: A Molecular Approach21 Practice - What is the pH of a buffer that is 0.14 M HF (pK a = 3.15) and M KF? K a for HF = 7.0 x check if the approximation is valid by seeing if x < 5% of [HC 2 H 3 O 2 ] init the approximation is valid x = 1.4 x [HA][A 2 - ][H 3 O + ] initial change -x-x+x+x+x+x equilibrium x

22 Tro, Chemistry: A Molecular Approach22 Practice - What is the pH of a buffer that is 0.14 M HF (pK a = 3.15) and M KF? x = 1.4 x substitute x into the equilibrium concentration definitions and solve [HA][A 2 - ][H 3 O + ] initial change -x-x+x+x+x+x equilibrium E x x 0.14 x

23 Tro, Chemistry: A Molecular Approach23 Practice - What is the pH of a buffer that is 0.14 M HF (pK a = 3.15) and M KF? substitute [H 3 O + ] into the formula for pH and solve [HA][A - ][H 3 O + ] initial change -x-x+x+x+x+x equilibrium E-3

24 Tro, Chemistry: A Molecular Approach24 Practice - What is the pH of a buffer that is 0.14 M HF (pK a = 3.15) and M KF? check by substituting the equilibrium concentrations back into the equilibrium constant expression and comparing the calculated K a to the given K a the values are close enough [HA][A - ][H 3 O + ] initial change -x-x+x+x+x+x equilibrium E-3 K a for HF = 7.0 x 10 -4

25 Tro, Chemistry: A Molecular Approach25 Henderson-Hasselbalch Equation calculating the pH of a buffer solution can be simplified by using an equation derived from the K a expression called the Henderson- Hasselbalch Equation the equation calculates the pH of a buffer from the K a and initial concentrations of the weak acid and salt of the conjugate base as long as the x is small approximation is valid

26 Tro, Chemistry: A Molecular Approach26 Deriving the Henderson-Hasselbalch Equation

27 Tro, Chemistry: A Molecular Approach27 Ex What is the pH of a buffer that is M HC 7 H 5 O 2 and M NaC 7 H 5 O 2 ? Assume the [HA] and [A - ] equilibrium concentrations are the same as the initial Substitute into the Henderson-Hasselbalch Equation Check the x is small approximation HC 7 H 5 O 2 + H 2 O C 7 H 5 O 2 + H 3 O + K a for HC 7 H 5 O 2 = 6.5 x 10 -5

28 Tro, Chemistry: A Molecular Approach28 Practice - What is the pH of a buffer that is 0.14 M HF (pK a = 3.15) and M KF?

29 Tro, Chemistry: A Molecular Approach29 Practice - What is the pH of a buffer that is 0.14 M HF (pK a = 3.15) and M KF? find the pK a from the given K a Assume the [HA] and [A - ] equilibrium concentrations are the same as the initial Substitute into the Henderson-Hasselbalch Equation Check the x is small approximation HF + H 2 O F + H 3 O +

30 Tro, Chemistry: A Molecular Approach30 Do I Use the Full Equilibrium Analysis or the Henderson-Hasselbalch Equation? the Henderson-Hasselbalch equation is generally good enough when the x is small approximation is applicable generally, the x is small approximation will work when both of the following are true: a) the initial concentrations of acid and salt are not very dilute b) the K a is fairly small for most problems, this means that the initial acid and salt concentrations should be over 1000x larger than the value of K a

31 Tro, Chemistry: A Molecular Approach31 How Much Does the pH of a Buffer Change When an Acid or Base Is Added? though buffers do resist change in pH when acid or base are added to them, their pH does change calculating the new pH after adding acid or base requires breaking the problem into 2 parts 1. a stoichiometry calculation for the reaction of the added chemical with one of the ingredients of the buffer to reduce its initial concentration and increase the concentration of the other added acid reacts with the A to make more HA added base reacts with the HA to make more A 2. an equilibrium calculation of [H 3 O + ] using the new initial values of [HA] and [A ]

32 Tro, Chemistry: A Molecular Approach32 Ex What is the pH of a buffer that has mol HC 2 H 3 O 2 and mol NaC 2 H 3 O 2 in 1.00 L that has mol NaOH added to it? If the added chemical is a base, write a reaction for OH with HA. If the added chemical is an acid, write a reaction for it with A. Construct a stoichiometry table for the reaction HC 2 H 3 O 2 + OH C 2 H 3 O 2 + H 2 O HAA-A- OH mols Before mols added mols After

33 Tro, Chemistry: A Molecular Approach33 Ex What is the pH of a buffer that has mol HC 2 H 3 O 2 and mol NaC 2 H 3 O 2 in 1.00 L that has mol NaOH added to it? Fill in the table – tracking the changes in the number of moles for each component HC 2 H 3 O 2 + OH C 2 H 3 O 2 + H 2 O HAA-A- OH mols Before mols added mols After

34 Tro, Chemistry: A Molecular Approach34 Ex What is the pH of a buffer that has mol HC 2 H 3 O 2 and mol NaC 2 H 3 O 2 in 1.00 L that has mol NaOH added to it? If the added chemical is a base, write a reaction for OH with HA. If the added chemical is an acid, write a reaction for it with A. Construct a stoichiometry table for the reaction Enter the initial number of moles for each HC 2 H 3 O 2 + OH C 2 H 3 O 2 + H 2 O HAA-A- OH mols Before mols change mols End

35 Tro, Chemistry: A Molecular Approach35 Ex What is the pH of a buffer that has mol HC 2 H 3 O 2 and mol NaC 2 H 3 O 2 in 1.00 L that has mol NaOH added to it? using the added chemical as the limiting reactant, determine how the moles of the other chemicals change add the change to the initial number of moles to find the moles after reaction divide by the liters of solution to find the new molarities HC 2 H 3 O 2 + OH C 2 H 3 O 2 + H 2 O HAA-A- OH mols Before mols change mols End new Molarity

36 Tro, Chemistry: A Molecular Approach36 Ex What is the pH of a buffer that has mol HC 2 H 3 O 2 and mol NaC 2 H 3 O 2 in 1.00 L that has mol NaOH added to it? Write the reaction for the acid with water Construct an ICE table for the reaction Enter the initial concentrations – assuming the [H 3 O + ] from water is 0, and using the new molarities of the [HA] and [A ] HC 2 H 3 O 2 + H 2 O C 2 H 3 O 2 + H 3 O + [HA][A - ][H 3 O + ] initial change equilibrium

37 Tro, Chemistry: A Molecular Approach37 [HA][A - ][H 3 O + ] initial change equilibrium Ex What is the pH of a buffer that has mol HC 2 H 3 O 2 and mol NaC 2 H 3 O 2 in 1.00 L that has mol NaOH added to it? represent the change in the concentrations in terms of x sum the columns to find the equilibrium concentrations in terms of x substitute into the equilibrium constant expression +x+x+x+x x x x x HC 2 H 3 O 2 + H 2 O C 2 H 3 O 2 + H 3 O +

38 Tro, Chemistry: A Molecular Approach38 determine the value of K a since K a is very small, approximate the [HA] eq = [HA] init and [A ] eq = [A ] init solve for x [HA][A - ][H 3 O + ] initial change -x-x+x+x+x+x equilibrium x x Ex What is the pH of a buffer that has mol HC 2 H 3 O 2 and mol NaC 2 H 3 O 2 in 1.00 L that has mol NaOH added to it? x K a for HC 2 H 3 O 2 = 1.8 x 10 -5

39 Tro, Chemistry: A Molecular Approach39 Ex What is the pH of a buffer that has mol HC 2 H 3 O 2 and mol NaC 2 H 3 O 2 in 1.00 L that has mol NaOH added to it? K a for HC 2 H 3 O 2 = 1.8 x check if the approximation is valid by seeing if x < 5% of [HC 2 H 3 O 2 ] init the approximation is valid x = 1.47 x [HA][A - ][H 3 O + ] initial change -x-x+x+x+x+x equilibrium x

40 Tro, Chemistry: A Molecular Approach40 Ex What is the pH of a buffer that has mol HC 2 H 3 O 2 and mol NaC 2 H 3 O 2 in 1.00 L that has mol NaOH added to it? x = 1.47 x substitute x into the equilibrium concentration definitions and solve [HA][A - ][H 3 O + ] initial change -x-x+x+x+x+x equilibrium E x x x

41 Tro, Chemistry: A Molecular Approach41 Ex What is the pH of a buffer that has mol HC 2 H 3 O 2 and mol NaC 2 H 3 O 2 in 1.00 L that has mol NaOH added to it? substitute [H 3 O + ] into the formula for pH and solve [HA][A - ][H 3 O + ] initial change -x-x+x+x+x+x equilibrium E-5

42 Tro, Chemistry: A Molecular Approach42 Ex What is the pH of a buffer that has mol HC 2 H 3 O 2 and mol NaC 2 H 3 O 2 in 1.00 L that has mol NaOH added to it? check by substituting the equilibrium concentrations back into the equilibrium constant expression and comparing the calculated K a to the given K a the values match [HA][A - ][H 3 O + ] initial change -x-x+x+x+x+x equilibrium E-5 K a for HC 2 H 3 O 2 = 1.8 x 10 -5

43 Tro, Chemistry: A Molecular Approach43 Ex What is the pH of a buffer that has mol HC 2 H 3 O 2 and mol NaC 2 H 3 O 2 in 1.00 L that has mol NaOH added to it? find the pK a from the given K a Assume the [HA] and [A - ] equilibrium concentrations are the same as the initial HC 2 H 3 O 2 + H 2 O C 2 H 3 O 2 + H 3 O + K a for HC 2 H 3 O 2 = 1.8 x [HA][A - ][H 3 O + ] initial change -x-x+x+x+x+x equilibrium x

44 Tro, Chemistry: A Molecular Approach44 Ex What is the pH of a buffer that has mol HC 2 H 3 O 2 and mol NaC 2 H 3 O 2 in 1.00 L that has mol NaOH added to it? Substitute into the Henderson-Hasselbalch Equation Check the x is small approximation HC 2 H 3 O 2 + H 2 O C 2 H 3 O 2 + H 3 O + pK a for HC 2 H 3 O 2 = 4.745

45 Tro, Chemistry: A Molecular Approach45 Ex 16.3 – Compare the effect on pH of adding mol NaOH to a buffer that has mol HC 2 H 3 O 2 and mol NaC 2 H 3 O 2 in 1.00 L to adding mol NaOH to 1.00 L of pure water? HC 2 H 3 O 2 + H 2 O C 2 H 3 O 2 + H 3 O + pK a for HC 2 H 3 O 2 = 4.745

46 Tro, Chemistry: A Molecular Approach46 Basic Buffers B: (aq) + H 2 O (l) H:B + (aq) + OH (aq) buffers can also be made by mixing a weak base, (B:), with a soluble salt of its conjugate acid, H:B + Cl H 2 O (l) + NH 3 (aq) NH 4 + (aq) + OH (aq)

47 Tro, Chemistry: A Molecular Approach47 Ex What is the pH of a buffer that is 0.50 M NH 3 (pK b = 4.75) and 0.20 M NH 4 Cl? find the pK a of the conjugate acid (NH 4 + ) from the given K b Assume the [B] and [HB + ] equilibrium concentrations are the same as the initial Substitute into the Henderson-Hasselbalch Equation Check the x is small approximation NH 3 + H 2 O NH OH

48 Tro, Chemistry: A Molecular Approach48 Ex What is the pH of a buffer that is 0.50 M NH 3 (pK b = 4.75) and 0.20 M NH 4 Cl? find the pK b if given K b Assume the [B] and [HB + ] equilibrium concentrations are the same as the initial Substitute into the Henderson- Hasselbalch Equation Base Form, find pOH Check the x is small approximation Calculate pH from pOH NH 3 + H 2 O NH OH

49 Tro, Chemistry: A Molecular Approach49 Buffering Effectiveness a good buffer should be able to neutralize moderate amounts of added acid or base however, there is a limit to how much can be added before the pH changes significantly the buffering capacity is the amount of acid or base a buffer can neutralize the buffering range is the pH range the buffer can be effective the effectiveness of a buffer depends on two factors (1) the relative amounts of acid and base, and (2) the absolute concentrations of acid and base

50 HAA-A- OH mols Before mols added mols After Effect of Relative Amounts of Acid and Conjugate Base Buffer mol HA & mol A - Initial pH = 5.00 Buffer mol HA & mol A - Initial pH = 4.05 pK a (HA) = 5.00 after adding mol NaOH pH = 5.09 HA + OH A + H 2 O HAA-A- OH mols Before mols added mols After after adding mol NaOH pH = 4.25 a buffer is most effective with equal concentrations of acid and base

51 HAA-A- OH mols Before mols added mols After HAA-A- OH mols Before mols added mols After Effect of Absolute Concentrations of Acid and Conjugate Base Buffer mol HA & 0.50 mol A - Initial pH = 5.00 Buffer mol HA & mol A - Initial pH = 5.00 pK a (HA) = 5.00 after adding mol NaOH pH = 5.02 HA + OH A + H 2 O after adding mol NaOH pH = 5.18 a buffer is most effective when the concentrations of acid and base are largest

52 Tro, Chemistry: A Molecular Approach52 Effectiveness of Buffers a buffer will be most effective when the [base]:[acid] = 1 equal concentrations of acid and base effective when 0.1 < [base]:[acid] < 10 a buffer will be most effective when the [acid] and the [base] are large

53 53 Buffering Range we have said that a buffer will be effective when 0.1 < [base]:[acid] < 10 substituting into the Henderson-Hasselbalch we can calculate the maximum and minimum pH at which the buffer will be effective Lowest pHHighest pH therefore, the effective pH range of a buffer is pK a ± 1 when choosing an acid to make a buffer, choose one whose is pK a is closest to the pH of the buffer

54 Tro, Chemistry: A Molecular Approach54 Ex. 16.5a – Which of the following acids would be the best choice to combine with its sodium salt to make a buffer with pH 4.25? Chlorous Acid, HClO 2 pK a = 1.95 Nitrous Acid, HNO 2 pK a = 3.34 Formic Acid, HCHO 2 pK a = 3.74 Hypochlorous Acid, HClOpK a = 7.54

55 Tro, Chemistry: A Molecular Approach55 Ex. 16.5a – Which of the following acids would be the best choice to combine with its sodium salt to make a buffer with pH 4.25? Chlorous Acid, HClO 2 pK a = 1.95 Nitrous Acid, HNO 2 pK a = 3.34 Formic Acid, HCHO 2 pK a = 3.74 Hypochlorous Acid, HClOpK a = 7.54 The pK a of HCHO 2 is closest to the desired pH of the buffer, so it would give the most effective buffering range.

56 Tro, Chemistry: A Molecular Approach56 Ex. 16.5b – What ratio of NaCHO 2 : HCHO 2 would be required to make a buffer with pH 4.25? Formic Acid, HCHO 2, pK a = 3.74 to make the buffer with pH 4.25, you would use 3.24 times as much NaCHO 2 as HCHO 2

57 Tro, Chemistry: A Molecular Approach57 Buffering Capacity buffering capacity is the amount of acid or base that can be added to a buffer without destroying its effectiveness the buffering capacity increases with increasing absolute concentration of the buffer components as the [base]:[acid] ratio approaches 1, the ability of the buffer to neutralize both added acid and base improves buffers that need to work mainly with added acid generally have [base] > [acid] buffers that need to work mainly with added base generally have [acid] > [base]

58 Tro, Chemistry: A Molecular Approach58 Buffering Capacity a concentrated buffer can neutralize more added acid or base than a dilute buffer

59 Tro, Chemistry: A Molecular Approach59 Titration in an acid-base titration, a solution of unknown concentration (titrant) is slowly added to a solution of known concentration from a burette until the reaction is complete when the reaction is complete we have reached the endpoint of the titration an indicator may be added to determine the endpoint an indicator is a chemical that changes color when the pH changes when the moles of H 3 O + = moles of OH, the titration has reached its equivalence point

60 Tro, Chemistry: A Molecular Approach60 Titration

61 Tro, Chemistry: A Molecular Approach61 Titration Curve a plot of pH vs. amount of added titrant the inflection point of the curve is the equivalence point of the titration prior to the equivalence point, the known solution in the flask is in excess, so the pH is closest to its pH the pH of the equivalence point depends on the pH of the salt solution equivalence point of neutral salt, pH = 7 equivalence point of acidic salt, pH < 7 equivalence point of basic salt, pH > 7 beyond the equivalence point, the unknown solution in the burette is in excess, so the pH approaches its pH

62 Tro, Chemistry: A Molecular Approach62 Titration Curve: Unknown Strong Base Added to Strong Acid

63 Tro, Chemistry: A Molecular Approach63

64 Tro, Chemistry: A Molecular Approach64 Titration of 25 mL of M HCl with M NaOH HCl (aq) + NaOH (aq) NaCl (aq) + H 2 O (aq) initial pH = -log(0.100) = 1.00 initial mol of HCl = L x mol/L = 2.50 x before equivalence point added 5.0 mL NaOH 5.0 x mol NaOH 2.00 x mol HCl

65 Tro, Chemistry: A Molecular Approach65 Titration of 25 mL of M HCl with M NaOH HCl (aq) + NaOH (aq) NaCl (aq) + H 2 O (aq) at equivalence, 0.00 mol HCl and 0.00 mol NaOH pH at equivalence = 7.00 after equivalence point added 30.0 mL NaOH 5.0 x mol NaOH xs

66 Tro, Chemistry: A Molecular Approach66 Titration of 25 mL of M HCl with M NaOH HCl (aq) + NaOH (aq) NaCl (aq) + H 2 O (aq) at equivalence, 0.00 mol HCl and 0.00 mol NaOH pH at equivalence = 7.00 after equivalence point added 30.0 mL NaOH 5.0 x mol NaOH xs

67 Tro, Chemistry: A Molecular Approach67 added 30.0 mL NaOH mol NaOH pH = added 35.0 mL NaOH mol NaOH pH = Adding NaOH to HCl 25.0 mL M HCl mol HCl pH = 1.00 added 5.0 mL NaOH mol HCl pH = 1.18 added 10.0 mL NaOH mol HCl pH = 1.37 added 15.0 mL NaOH mol HCl pH = 1.60 added 20.0 mL NaOH mol HCl pH = 1.95 added 25.0 mL NaOH equivalence point pH = 7.00 added 40.0 mL NaOH mol NaOH pH = added 50.0 mL NaOH mol NaOH pH = 12.52

68 Tro, Chemistry: A Molecular Approach68 Titration of 25.0 mL of M HCl with M NaOH The 1st derivative of the curve is maximum at the equivalence point Since the solutions are equal concentration, the equivalence point is at equal volumes

69 Tro, Chemistry: A Molecular Approach69 After about pH 3, there is practically no HCl left, it has all been reacted and become NaCl + H 2 O

70 Tro, Chemistry: A Molecular Approach70 Titration of 25 mL of M HCHO 2 with M NaOH HCHO 2(aq) + NaOH (aq) NaCHO 2(aq) + H 2 O (aq) Initial pH: [HCHO 2 ][CHO 2 - ][H 3 O + ] initial change -x-x+x+x+x+x equilibrium xxx K a = 1.8 x 10 -4

71 Tro, Chemistry: A Molecular Approach71 Titration of 25 mL of M HCHO 2 with M NaOH HCHO 2(aq) + NaOH (aq) NaCHO 2 (aq) + H 2 O (aq) initial mol of HCHO 2 = L x mol/L = 2.50 x before equivalence added 5.0 mL NaOH HAA-A- OH mols Before2.50E-300 mols added --5.0E-4 mols After 2.00E-35.0E-4 0

72 72 Titration of 25 mL of M HCHO 2 with M NaOH HCHO 2(aq) + NaOH (aq) NaCHO 2 (aq) + H 2 O (aq) initial mol of HCHO 2 = L x mol/L = 2.50 x at equivalence added 25.0 mL NaOH HAA-A- OH mols Before2.50E-300 mols added E-3 mols After 02.50E-3 0 [HCHO 2 ][CHO 2 - ][OH ] initial change +x+x-x-x+x+x equilibrium x5.00E-2-xx CHO 2 (aq) + H 2 O (l) HCHO 2(aq) + OH (aq) K b = 5.6 x [OH - ] = 1.7 x M

73 Tro, Chemistry: A Molecular Approach73 Titration of 25 mL of M HCl with M NaOH HCHO 2(aq) + NaOH (aq) NaCHO 2 (aq) + H 2 O (aq) after equivalence point added 30.0 mL NaOH 5.0 x mol NaOH xs

74 Tro, Chemistry: A Molecular Approach74 Titration of 25 mL of M HCl with M NaOH HCHO 2(aq) + NaOH (aq) NaCHO 2 (aq) + H 2 O (aq) after equivalence point added 30.0 mL NaOH 5.0 x mol NaOH xs

75 Tro, Chemistry: A Molecular Approach75 added 30.0 mL NaOH mol NaOH xs pH = added 35.0 mL NaOH mol NaOH xs pH = Adding NaOH to HCHO 2 added 12.5 mL NaOH mol HCHO 2 pH = 3.74 = pK a half-neutralization initial HCHO 2 solution mol HCHO 2 pH = 2.37 added 5.0 mL NaOH mol HCHO 2 pH = 3.14 added 10.0 mL NaOH mol HCHO 2 pH = 3.56 added 15.0 mL NaOH mol HCHO 2 pH = 3.92 added 20.0 mL NaOH mol HCHO 2 pH = 4.34 added 40.0 mL NaOH mol NaOH xs pH = added 25.0 mL NaOH equivalence point mol CHO 2 [CHO 2 ] init = M [OH ] eq = 1.7 x pH = 8.23 added 50.0 mL NaOH mol NaOH xs pH = 12.52

76 Tro, Chemistry: A Molecular Approach76

77 Tro, Chemistry: A Molecular Approach77 Titration of 25.0 mL of M HCHO 2 with M NaOH The 1st derivative of the curve is maximum at the equivalence point Since the solutions are equal concentration, the equivalence point is at equal volumes pH at equivalence = 8.23

78 Tro, Chemistry: A Molecular Approach78 at pH 3.74, the [HCHO 2 ] = [CHO 2 ]; the acid is half neutralized half-neutralization occurs when pH = pK a

79 Tro, Chemistry: A Molecular Approach79 Titrating Weak Acid with a Strong Base the initial pH is that of the weak acid solution calculate like a weak acid equilibrium problem e.g., 15.5 and 15.6 before the equivalence point, the solution becomes a buffer calculate mol HA init and mol A init using reaction stoichiometry calculate pH with Henderson-Hasselbalch using mol HA init and mol A init half-neutralization pH = pK a

80 Tro, Chemistry: A Molecular Approach80 Titrating Weak Acid with a Strong Base at the equivalence point, the mole HA = mol Base, so the resulting solution has only the conjugate base anion in it before equilibrium is established mol A = original mole HA calculate the volume of added base like Ex 4.8 [A ] init = mol A /total liters calculate like a weak base equilibrium problem e.g., beyond equivalence point, the OH is in excess [OH ] = mol MOH xs/total liters [H 3 O + ][OH ]=1 x

81 Tro, Chemistry: A Molecular Approach81 Ex 16.7a – A 40.0 mL sample of M HNO 2 is titrated with M KOH. Calculate the volume of KOH at the equivalence point Write an equation for the reaction for B with HA. Use Stoichiometry to determine the volume of added B HNO 2 + KOH NO 2 + H 2 O

82 Tro, Chemistry: A Molecular Approach82 Ex 16.7b – A 40.0 mL sample of M HNO 2 is titrated with M KOH. Calculate the pH after adding 5.00 mL KOH Write an equation for the reaction for B with HA. Determine the moles of HA before & moles of added B Make a stoichiometry table and determine the moles of HA in excess and moles A made HNO 2 + KOH NO 2 + H 2 O HNO 2 NO 2 - OH mols Before mols added mols After

83 Tro, Chemistry: A Molecular Approach83 Ex 16.7b – A 40.0 mL sample of M HNO 2 is titrated with M KOH. Calculate the pH after adding 5.00 mL KOH. Write an equation for the reaction of HA with H 2 O Determine K a and pK a for HA Use the Henderson- Hasselbalch Equation to determine the pH HNO 2 + H 2 O NO 2 + H 3 O + HNO 2 NO 2 - OH mols Before mols added mols After Table 15.5 K a = 4.6 x 10 -4

84 Tro, Chemistry: A Molecular Approach84 Ex 16.7b – A 40.0 mL sample of M HNO 2 is titrated with M KOH. Calculate the pH at the half-equivalence point Write an equation for the reaction for B with HA. Determine the moles of HA before & moles of added B Make a stoichiometry table and determine the moles of HA in excess and moles A made HNO 2 + KOH NO 2 + H 2 O HNO 2 NO 2 - OH mols Before mols added mols After at half-equivalence, moles KOH = ½ mole HNO 2

85 Tro, Chemistry: A Molecular Approach85 Ex 16.7b – A 40.0 mL sample of M HNO 2 is titrated with M KOH. Calculate the pH at the half-equivalence point. Write an equation for the reaction of HA with H 2 O Determine K a and pK a for HA Use the Henderson- Hasselbalch Equation to determine the pH HNO 2 + H 2 O NO 2 + H 3 O + HNO 2 NO 2 - OH mols Before mols added mols After Table 15.5 K a = 4.6 x 10 -4

86 Tro, Chemistry: A Molecular Approach86 Titration Curve of a Weak Base with a Strong Acid

87 Tro, Chemistry: A Molecular Approach87 Titration of a Polyprotic Acid if K a1 >> K a2, there will be two equivalence points in the titration the closer the K a s are to each other, the less distinguishable the equivalence points are titration of 25.0 mL of M H 2 SO 3 with M NaOH

88 Tro, Chemistry: A Molecular Approach88 Monitoring pH During a Titration the general method for monitoring the pH during the course of a titration is to measure the conductivity of the solution due to the [H 3 O + ] using a probe that specifically measures just H 3 O + the endpoint of the titration is reached at the equivalence point in the titration – at the inflection point of the titration curve if you just need to know the amount of titrant added to reach the endpoint, we often monitor the titration with an indicator

89 Tro, Chemistry: A Molecular Approach89 Monitoring pH During a Titration

90 Tro, Chemistry: A Molecular Approach90 Indicators many dyes change color depending on the pH of the solution these dyes are weak acids, establishing an equilibrium with the H 2 O and H 3 O + in the solution HInd (aq) + H 2 O (l) Ind (aq) + H 3 O + (aq) the color of the solution depends on the relative concentrations of Ind :HInd when Ind :HInd 1, the color will be mix of the colors of Ind and HInd when Ind :HInd > 10, the color will be mix of the colors of Ind when Ind :HInd < 0.1, the color will be mix of the colors of HInd

91 91 Phenolphthalein

92 Tro, Chemistry: A Molecular Approach92 Methyl Red

93 Tro, Chemistry: A Molecular Approach93 Monitoring a Titration with an Indicator for most titrations, the titration curve shows a very large change in pH for very small additions of base near the equivalence point an indicator can therefore be used to determine the endpoint of the titration if it changes color within the same range as the rapid change in pH pK a of HInd pH at equivalence point

94 94 Acid-Base Indicators

95 Tro, Chemistry: A Molecular Approach95 Solubility Equilibria all ionic compounds dissolve in water to some degree however, many compounds have such low solubility in water that we classify them as insoluble we can apply the concepts of equilibrium to salts dissolving, and use the equilibrium constant for the process to measure relative solubilities in water

96 Tro, Chemistry: A Molecular Approach96 Solubility Product the equilibrium constant for the dissociation of a solid salt into its aqueous ions is called the solubility product, K sp for an ionic solid M n X m, the dissociation reaction is: M n X m (s) nM m+ (aq) + mX n (aq) the solubility product would be K sp = [M m+ ] n [X n ] m for example, the dissociation reaction for PbCl 2 is PbCl 2 (s) Pb 2+ (aq) + 2 Cl (aq) and its equilibrium constant is K sp = [Pb 2+ ][Cl ] 2

97 Tro, Chemistry: A Molecular Approach97

98 Tro, Chemistry: A Molecular Approach98 Molar Solubility solubility is the amount of solute that will dissolve in a given amount of solution at a particular temperature the molar solubility is the number of moles of solute that will dissolve in a liter of solution the molarity of the dissolved solute in a saturated solution for the general reaction M n X m (s) nM m+ (aq) + mX n (aq)

99 Tro, Chemistry: A Molecular Approach99 Ex 16.8 – Calculate the molar solubility of PbCl 2 in pure water at 25 C Write the dissociation reaction and K sp expression Create an ICE table defining the change in terms of the solubility of the solid [Pb 2+ ][Cl ] Initial00 Change+S+2S EquilibriumS2S PbCl 2 (s) Pb 2+ (aq) + 2 Cl (aq) K sp = [Pb 2+ ][Cl ] 2

100 Tro, Chemistry: A Molecular Approach100 Ex 16.8 – Calculate the molar solubility of PbCl 2 in pure water at 25 C Substitute into the K sp expression Find the value of K sp from Table 16.2, plug into the equation and solve for S [Pb 2+ ][Cl ] Initial00 Change+S+2S EquilibriumS2S K sp = [Pb 2+ ][Cl ] 2 K sp = (S)(2S) 2

101 Tro, Chemistry: A Molecular Approach101 Practice – Determine the K sp of PbBr 2 if its molar solubility in water at 25 C is 1.05 x M

102 Tro, Chemistry: A Molecular Approach102 Practice – Determine the K sp of PbBr 2 if its molar solubility in water at 25 C is 1.05 x M Write the dissociation reaction and K sp expression Create an ICE table defining the change in terms of the solubility of the solid [Pb 2+ ][Br ] Initial00 Change+(1.05 x )+2(1.05 x ) Equilibrium(1.05 x )(2.10 x ) PbBr 2 (s) Pb 2+ (aq) + 2 Br (aq) K sp = [Pb 2+ ][Br ] 2

103 Tro, Chemistry: A Molecular Approach103 Practice – Determine the K sp of PbBr 2 if its molar solubility in water at 25 C is 1.05 x M Substitute into the K sp expression plug into the equation and solve K sp = [Pb 2+ ][Br ] 2 K sp = (1.05 x )(2.10 x ) 2 [Pb 2+ ][Br ] Initial00 Change+(1.05 x )+2(1.05 x ) Equilibrium(1.05 x )(2.10 x )

104 Tro, Chemistry: A Molecular Approach104 K sp and Relative Solubility molar solubility is related to K sp but you cannot always compare solubilities of compounds by comparing their K sp s in order to compare K sp s, the compounds must have the same dissociation stoichiometry

105 Tro, Chemistry: A Molecular Approach105 The Effect of Common Ion on Solubility addition of a soluble salt that contains one of the ions of the insoluble salt, decreases the solubility of the insoluble salt for example, addition of NaCl to the solubility equilibrium of solid PbCl 2 decreases the solubility of PbCl 2 PbCl 2 (s) Pb 2+ (aq) + 2 Cl (aq) addition of Cl shifts the equilibrium to the left

106 Tro, Chemistry: A Molecular Approach106 Ex – Calculate the molar solubility of CaF 2 in M NaF at 25 C Write the dissociation reaction and K sp expression Create an ICE table defining the change in terms of the solubility of the solid [Ca 2+ ][F ] Initial Change+S+2S EquilibriumS S CaF 2 (s) Ca 2+ (aq) + 2 F (aq) K sp = [Ca 2+ ][F ] 2

107 Tro, Chemistry: A Molecular Approach107 Ex – Calculate the molar solubility of CaF 2 in M NaF at 25 C Substitute into the K sp expression assume S is small Find the value of K sp from Table 16.2, plug into the equation and solve for S [Ca 2+ ][F ] Initial Change+S+2S EquilibriumS S K sp = [Ca 2+ ][F ] 2 K sp = (S)( S) 2 K sp = (S)(0.100) 2

108 Tro, Chemistry: A Molecular Approach108 The Effect of pH on Solubility for insoluble ionic hydroxides, the higher the pH, the lower the solubility of the ionic hydroxide and the lower the pH, the higher the solubility higher pH = increased [OH ] M(OH) n (s) M n+ (aq) + nOH (aq) for insoluble ionic compounds that contain anions of weak acids, the lower the pH, the higher the solubility M 2 (CO 3 ) n (s) 2 M n+ (aq) + nCO 3 2 (aq) H 3 O + (aq) + CO 3 2 (aq) HCO 3 (aq) + H 2 O(l)

109 Tro, Chemistry: A Molecular Approach109 Precipitation precipitation will occur when the concentrations of the ions exceed the solubility of the ionic compound if we compare the reaction quotient, Q, for the current solution concentrations to the value of K sp, we can determine if precipitation will occur Q = K sp, the solution is saturated, no precipitation Q < K sp, the solution is unsaturated, no precipitation Q > K sp, the solution would be above saturation, the salt above saturation will precipitate some solutions with Q > K sp will not precipitate unless disturbed – these are called supersaturated solutions

110 Tro, Chemistry: A Molecular Approach110 precipitation occurs if Q > K sp a supersaturated solution will precipitate if a seed crystal is added

111 Tro, Chemistry: A Molecular Approach111 Selective Precipitation a solution containing several different cations can often be separated by addition of a reagent that will form an insoluble salt with one of the ions, but not the others a successful reagent can precipitate with more than one of the cations, as long as their K sp values are significantly different

112 Tro, Chemistry: A Molecular Approach112 Ex What is the minimum [OH ] necessary to just begin to precipitate Mg 2+ (with [0.059]) from seawater? precipitating may just occur when Q = K sp

113 Tro, Chemistry: A Molecular Approach113 Ex What is the [Mg 2+ ] when Ca 2+ (with [0.011]) just begins to precipitate from seawater? precipitating Mg 2+ begins when [OH ] = 1.9 x M

114 Tro, Chemistry: A Molecular Approach114 Ex What is the [Mg 2+ ] when Ca 2+ (with [0.011]) just begins to precipitate from seawater? precipitating Mg 2+ begins when [OH ] = 1.9 x M precipitating Ca 2+ begins when [OH ] = 2.06 x M when Ca 2+ just begins to precipitate out, the [Mg 2+ ] has dropped from M to 4.8 x M

115 Tro, Chemistry: A Molecular Approach115 Qualitative Analysis an analytical scheme that utilizes selective precipitation to identify the ions present in a solution is called a qualitative analysis scheme wet chemistry a sample containing several ions is subjected to the addition of several precipitating agents addition of each reagent causes one of the ions present to precipitate out

116 Tro, Chemistry: A Molecular Approach116 Qualitative Analysis

117 117

118 Tro, Chemistry: A Molecular Approach118 Group 1 group one cations are Ag +, Pb 2+, and Hg 2 2+ all these cations form compounds with Cl that are insoluble in water as long as the concentration is large enough PbCl 2 may be borderline molar solubility of PbCl 2 = 1.43 x M precipitated by the addition of HCl

119 119 Group 2 group two cations are Cd 2+, Cu 2+, Bi 3+, Sn 4+, As 3+, Pb 2+, Sb 3+, and Hg 2+ all these cations form compounds with HS and S 2 that are insoluble in water at low pH precipitated by the addition of H 2 S in HCl

120 120 Group 3 group three cations are Fe 2+, Co 2+, Zn 2+, Mn 2+, Ni 2+ precipitated as sulfides; as well as Cr 3+, Fe 3+, and Al 3+ precipitated as hydroxides all these cations form compounds with S 2 that are insoluble in water at high pH precipitated by the addition of H 2 S in NaOH

121 Tro, Chemistry: A Molecular Approach121 Group 4 group four cations are Mg 2+, Ca 2+, Ba 2+ all these cations form compounds with PO 4 3 that are insoluble in water at high pH precipitated by the addition of (NH 4 ) 2 HPO 4

122 122 Group 5 group five cations are Na +, K +, NH 4 + all these cations form compounds that are soluble in water – they do not precipitate identified by the color of their flame

123 Tro, Chemistry: A Molecular Approach123 Complex Ion Formation transition metals tend to be good Lewis acids they often bond to one or more H 2 O molecules to form a hydrated ion H 2 O is the Lewis base, donating electron pairs to form coordinate covalent bonds Ag + (aq) + 2 H 2 O(l) Ag(H 2 O) 2 + (aq) ions that form by combining a cation with several anions or neutral molecules are called complex ions e.g., Ag(H 2 O) 2 + the attached ions or molecules are called ligands e.g., H 2 O

124 Tro, Chemistry: A Molecular Approach124 Complex Ion Equilibria if a ligand is added to a solution that forms a stronger bond than the current ligand, it will replace the current ligand Ag(H 2 O) 2 + (aq) + 2 NH 3(aq) Ag(NH 3 ) 2 + (aq) + 2 H 2 O (l) generally H 2 O is not included, since its complex ion is always present in aqueous solution Ag + (aq) + 2 NH 3(aq) Ag(NH 3 ) 2 + (aq)

125 Tro, Chemistry: A Molecular Approach125 Formation Constant the reaction between an ion and ligands to form a complex ion is called a complex ion formation reaction Ag + (aq) + 2 NH 3(aq) Ag(NH 3 ) 2 + (aq) the equilibrium constant for the formation reaction is called the formation constant, K f

126 Tro, Chemistry: A Molecular Approach126 Formation Constants

127 Tro, Chemistry: A Molecular Approach127 Ex – mL of 1.5 x M Cu(NO 3 ) 2 is mixed with mL of 0.20 M NH 3. What is the [Cu 2+ ] at equilibrium? Write the formation reaction and K f expression. Look up K f value Determine the concentration of ions in the diluted solutions Cu 2+ (aq) + 4 NH 3 (aq) Cu(NH 3 ) 2 2+ (aq)

128 128 Ex – mL of 1.5 x M Cu(NO 3 ) 2 is mixed with mL of 0.20 M NH 3. What is the [Cu 2+ ] at equilibrium? Create an ICE table. Since K f is large, assume all the Cu 2+ is converted into complex ion, then the system returns to equilibrium [Cu 2+ ][NH 3 ][Cu(NH 3 ) 2 2+ ] Initial6.7E Change-6.7E-4-4(6.7E-4)+ 6.7E-4 Equilibriumx E-4 Cu 2+ (aq) + 4 NH 3 (aq) Cu(NH 3 ) 2 2+ (aq)

129 Tro, Chemistry: A Molecular Approach129 Ex – mL of 1.5 x M Cu(NO 3 ) 2 is mixed with mL of 0.20 M NH 3. What is the [Cu 2+ ] at equilibrium? Cu 2+ (aq) + 4 NH 3 (aq) Cu(NH 3 ) 2 2+ (aq) Substitute in and solve for x confirm thex is small approximation [Cu 2+ ][NH 3 ][Cu(NH 3 ) 2 2+ ] Initial6.7E Change-6.7E-4-4(6.7E-4)+ 6.7E-4 Equilibriumx E-4 since 2.7 x << 6.7 x 10 -4, the approximation is valid

130 Tro, Chemistry: A Molecular Approach130 The Effect of Complex Ion Formation on Solubility the solubility of an ionic compound that contains a metal cation that forms a complex ion increases in the presence of aqueous ligands AgCl (s) Ag + (aq) + Cl (aq) K sp = 1.77 x Ag + (aq) + 2 NH 3(aq) Ag(NH 3 ) 2 + (aq) K f = 1.7 x 10 7 adding NH 3 to a solution in equilibrium with AgCl (s) increases the solubility of Ag +

131 131

132 Tro, Chemistry: A Molecular Approach132 Solubility of Amphoteric Metal Hydroxides many metal hydroxides are insoluble all metal hydroxides become more soluble in acidic solution shifting the equilibrium to the right by removing OH some metal hydroxides also become more soluble in basic solution acting as a Lewis base forming a complex ion substances that behave as both an acid and base are said to be amphoteric some cations that form amphoteric hydroxides include Al 3+, Cr 3+, Zn 2+, Pb 2+, and Sb 2+

133 Tro, Chemistry: A Molecular Approach133 Al 3+ Al 3+ is hydrated in water to form an acidic solution Al(H 2 O) 6 3+ (aq) + H 2 O (l) Al(H 2 O) 5 (OH) 2+ (aq) + H 3 O + (aq) addition of OH drives the equilibrium to the right and continues to remove H from the molecules Al(H 2 O) 5 (OH) 2+ (aq) + OH (aq) Al(H 2 O) 4 (OH) 2 + (aq) + H 2 O (l) Al(H 2 O) 4 (OH) 2 + (aq) + OH (aq) Al(H 2 O) 3 (OH) 3(s) + H 2 O (l)

134 Tro, Chemistry: A Molecular Approach134


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