CHM 2046C Module 12 Homework Packet Name: __________

 Jespersen 7th editin Module 12i:  Acid-Base Equilibria in Aquous Solutions Jespersen Chapter 16 Possible Answers 1 1 B1. Preparing a Buffer with a desired pH Section 16.7 2 1 3 3 3 Module 12i Total: 14 Module 12ii: Solubility and Simultaneous Equilibria Jespersen Chapter 17 1 1 2 2 2 2 1 Module 12ii Total: 11

Part A.  The Common-Ion and the Common-Ion Effect      1 point

1.     What is a common-ion?

2.  What would be the common ion for the following weak acids and weak bases and in addition to its formula give its name:

(a) Lactic Acid:  CH3CH(OH)COOH

(b) Acetic Acid:  CH3COOH

(c) Ammonia dissolved in Water:  NH3 in H2O

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3. What is the Common-Ion Effect for Weak acids, Weak base, or Salt Solubility?

Discuss Examples in General:

a)

b)

4. Does the pH of the solution increase, decrease, or stay the same when you:

(a) Add Ammonium nitrate to a dilute aqueous solution of ammonia?

(b) Add solid potassium acetate to a dilute aqueous solution of acetic acid?

(c) Add solid sodium lactate to a dilute aqueous solution of lactic acid?

(d) Add solid Potassium nitrate to a dilute solution of Potassium hydroxide?

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Part B.  Controlling pH: Buffer Solutions     1 points

1.     What is a buffer solution?

2. (a)  State the Henderson-Hasselbalch equation for an acid:

2. (b) State the Henderson-Hasselbalch equation for a base:

3. According to the Henderson-Hasselbalch equation, there are two terms that establish the pH of a buffer solution.

(a)What are they?

The first:                      The Second :

(b) What are the relative magnitudes of their contribution to the value of the pH?

4.  What are the two requirements for a buffer solution?

(a)

(b)

5. Explain how it is that the pH of a buffer solution does not change if the solution is diluted.

6. Explain what happens if the concentrations of the buffering components are increased so that their proportion maintains the same ratio.

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Part C1.  Derivation of Henderson-Hasselbalch Equations     1 point

1.     Given the ionization of weak acid, HA in water, derive the Henderson-Hasselbalch equation for a weak acid.

Start with writing the ionization reaction of the Weak Acid and setup the equilibrium constant expression for Ka:

2.     Given the ionization of base acid, BOH in water, derive the Henderson-Hasselbalch equation for a weak base.

Start with writing the ionization reaction of the weak Base and setup the equilibrium constant expression for Kb:

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Part D  Acid-Base/pH Titrations Curves     3 points

1. How do titration curves differ for strong acid-strong base, strong acid-weak base; weak acid-strong base, and weak acid-weak base titrations? Sketch curves for each and describe their pH equivalence points.

 (a) Strong Acid-Strong Base titrations                 (b) Strong Acid-Weak Base titrations                   (c) Weak Acid-Strong Base titrations:                   (d) Weak Acid-Weak Base titrations Page 6   (2) Titrations against polyprotic acids can be divided into separate titrations for each hydrogen ion removed. Explain and demonstrate with a titration curve.                               (3)  What are acid-base indicators. Choose appropriate indicators for the four titrations listed in question #1.

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Part E.  Acid-Base/pH Titration Points of Interests     3 points

 1. When a weak acid is titrated against a strong base,  there are five regions or possible points of interest. Describe/Explain each for both titrations:   Strong Base (BOH)-Weak Acid (HA):   First: Before base has been added,       Second: Anywhere between the start and the equivalence point,         Third: Exactly halfway to the equivalence point,         Fourth: At the equivalence point,         Fifth: Beyond the equivalence point,         Strong Acid-Weak Base:   2.  When a weak base is titrated against a strong acid,  there are five regions or possible points of interest. Describe/Explain each for both titrations:   First: Before acid has been added,             Page 8: Second: Anywhere between the start and the equivalence point,       Third: Exactly halfway to the equivalence point,             Fourth: At the equivalence point,             Fifth: Beyond the equivalence point,             3.   What are acid-base indicators. Choose appropriate indicators for the four titrations listed in question #1.

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Part F.  Acid-Base/pH Titration Calculations     3 points

 You titrate 25.0 mL of 0.10 M NH3 with 0.10 M HCl.   a.  What is the pH of the solution Before acid has been added?             b.  What is the pH at the equivalence point?             c. What is the pH at the halfway point of the titration?                 If you can do Part (d) successfully it will be worth a 5 point bonus: Bonus: d.  Calculate the pH of the solution after adding 5.00, 15.0, 20.0, 22.0, and 30.0 mL of acid.   Combine this information with that in parts a., b., and c. and plot the titration curve..

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Part G: Writing Ksp Expressions   1 point

Write the dissociation reactions for the following insoluble salts, then setup the Ksp expressions:

1.   Cu(OH)2  (s)   +   HOH  (l)    ß à

Ksp =                                                                                                                   Ksp = 2.2 x 10 -20

2.   MgCO3  (s)   +   HOH  (l)    ß à

Ksp =                                                                                                                   Ksp = 6.8 x 10 -6

3.   AlPO4  (s)   +   HOH  (l)    ß à

Ksp =                                                                                                                   Ksp = 6.3 x 10 -19

4.   Cu2S      (s)   +   HOH  (l)      ß à

Ksp =                                                                                                                   Ksp = 2.2 x 10 -48

5.   Mg(OH)2  (s)   +   HOH  (l)    ß à

Ksp =                                                                                                                   Ksp = 5.6 x 10 -12

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H. Comparing Solubility of Salts        1 point

An insoluble substance is defined as one that has solubility less than 0.01 moles/liter of water.

Using  Ksp  values, tell which salt in each pair is more soluble in water:

1.  AgCl   (Ksp = 1.8 x 10-10) or   AgCN  (Ksp = 6.0 x 10-17)

2.  Mg(OH)2  (Ksp = 5.6 x 10-12) or   Ca(OH)2  (Ksp = 5.5 x 10-5)

3.  MgCO3   (Ksp = 6.8 x 10-6) or   CaCO3  (Ksp = 3.4 x 10-9)

4.  FeCO3   (Ksp = 3.1 x 10-11) or   Ag2C2O4  (Ksp = 5.4 x 10-12)

Which compound in each pair is more soluble in water than is predicted by a calculation from Ksp :

5.    AgBr  (Ksp = 5.4 x 10-13) or   AgCN  (Ksp = 6.0 x 10-17)

6.    PbCl2  (Ksp = 1.7 x 10-5) or   PbCO3  (Ksp = 7.4 x 10-14)

7.    PbCl2  (Ksp = 1.7 x 10-5) or   PbF2  (Ksp = 3.3 x 10-8)

Which insoluble compound in each pair should be more soluble in nitric acid than in pure water:

8.    PbCl2  (Ksp = 1.7 x 10-5) or   PbS  (Ksp = 9.0 x 10-24)

9.  AgCl   (Ksp = 1.8 x 10-10) or   Al(OH)3  (Ksp = not known)

10.    AgI  (Ksp = 8.5 x 10-17) or   Ag2CO3  (Ksp = 8.5 x 10-12)

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I. Estimating Salt Solubility from Ksp       2 points

Given the following insoluble salts, estimate the salt’s molar solubility from the Ksp :

1.   Cu(OH)2  (s)   +   HOH  (l)    ß à

Ksp = 2.2 x 10 -20

2.   MgCO3  (s)   +   HOH  (l)    ß à

Ksp = 6.8 x 10 -6

3.   Al PO4  (s)   +   HOH  (l)    ß à

Ksp = 9.8 x 10 -21

4.   Cu2S      (s)   +   HOH  (l)      ß à

Ksp = 2.2 x 10 -48

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J. Determining Ksp from a Salt’s Solubility        2 points

Given the following insoluble salts, estimate the salt’s molar solubility from the Ksp :

1.   Calcium hydroxide, Ca(OH)2 (s)    dissolves in water to the extent of 1.04 g/L. What is the value of the Ksp ?

2.   You place 1.834 g of  Ca(OH)2 (s)     in 1.00 L of pure water at 25oC. The pH of the solution is found to be 12.68. What is the value of the Ksp ?

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Part L.  Common Ion Effect on Solubility      2 points

What is the solubility, in milligrams per milliliter, of BaCO3,  Ksp =1.2 x 10 -10, in

(a) pure water and

(b) in water containing 10.0 mg/ml of Na2CO3?

(c) What is the pH of the solution?  Ka2HCO3 =  4.8 x 10 -11

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Part N:  Precipitation Reactions                  2  points

If the concentration of the strontium ion, Sr 2+, is 2.5 x 10 -4 M does precipitation of SrSO4 occur when enough of the soluble salt  Na2SO4 is added to make the solution,  2.5 x 10 -4 M? Ksp for SrSO4 is 3.4 x 10 -7

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Part K: Chapter 17 Key Terms            1 point

1. ___________________ – a measure of how much hydronium or hydroxide can be added to a buffer solution before the buffer can no longer control the pH

2. .________________a material that is one color below a particular pH and another color above it; typically these materials are themselves weak acids or bases.

3. ___________________ – a solution that resists a change in pH when hydroxide or hydronium ions are added.

4. ___________________  - The limiting of acid (or base) ionization caused by the addition of its conjugate base (or conjugate acid)

5. ___________________ – the point in a titration at which one reactant has been exactly consumed by the solution of the other reactant.

6. ___________________ – the equilibrium constant for the formation of a complex ion.

7. ___________________ – an equation that allows us to calculate the pOH of a buffer:

 [conjugate acid] pOH = pKb + log ------------------------ [base]

8. ___________________ – the equilibrium constant for the dissociation of a sparingly soluble salt. For the salt AxBy, it has the form:

Ksp = [A y+]x [B x-]y

9. ___________________ – the substance added during a titration.

10. __________________ – the addition of a base to an acid (or vice versa) until the equivalence point is reached