Name:  Professor Key                                                                         

 

                                                                

Chemistry 3600

Second Examination

November 3, 1999

 

Directions: There are 4 questions, each worth equal points.  Some are more difficult than others.  Read over all questions before you start to answer.  Answering the harder questions last will insure that you obtain the maximum score.  You must show your work and pertinent formula's to get full credit.  Final numerical results are not as important as the methods used to find them.

 

Useful formula:

Useful information:

 


1.      Review of Chemical Equilibrium

 

(a)    Calculate the mean activity coefficients for CaCO3 in 0.1 F NaOH using the Debye-Hükel limiting law.

First, calculate ionic strength for 0.1 F NaOH

Next, calculate f+/- for Ca2+ and CO32- using

(b)   Using these activity coefficients, calculate the effective equilibrium constant, in terms of molarity, for dissolution of CaCO3. CaCO3 has Ksp=5´10-9 and

 

(c)    Calculate grams of CaCO3 (FW 100.09 g/mole) that will dissolve in 1.00 L of 0.1 F NaOH using the results from part (b). Hint: Ignore base reactions of carbonate because the solution is basic.

Since the pH is about 13, almost all carbonate is in the CO32- form

 

CaCO3 (s)

Ca2+(aq)

CO32- (aq)

Init

S

0

0

Change

-X

+X

+X

Final

S

X

X

 

 

2. Systematic Treatment of Chemical Equilibrium

 

A chemist is surprised to find that the dissolution of CaCO3 changes the pH of solution. Demonstrate that you could solve this problem for them using a systematic approach to chemical equilibrium. You may assume a closed system (no CO2 (g) in atmosphere).

 

(a)    Write the chemical equations

 

(b)   Give the charge balance equation

 

 

(c)    Give the mass balance equations

 

 

(d)   Give the chemical equilibrium expressions

 

 

 

(e)    Count equations and species. Are there enough equations to solve for concentrations of all species?

 

The 6 " unknowns" are Ca2+, H+, OH-, H2CO3, HCO3-, CO32-

There are 8 equations

Yes I can solve it!

 

3. Acid-Base Equilibria

(a)    Answer the following questions

1.      What is "carbonate error" in an acid/base titration?

An error in determination of end point due to carbonic acid formation in solutions.

 

2.      Where does the carbonate come from?

Air

 

3.      What is the carbonate error more important in the base than the acid solution?

Because carbonic acid reacts with base and CO2 solubility increases with pH

 

(b)   A chemist needs to prepare about 400 mL of a buffer solution of pH 5.0 using acetic acid. What should the weight of sodium acetate (FW 82.034 g/mole) should be added to 400 mL of 0.1 F acetic acid (FW 60.052 g/mole) solution to prepare the buffer? (Ka=1.75´10-5 for acetic acid.)

 

 

4. Acid-Base Titrations/Advanced Acid-Base Chemistry

20 mL of a diprotic acid, H2A, is titrated with 0.1 F NaOH titrant. The resulting titration curve is shown below.

(a)    Label the two points where the solution has the maximum buffer capacity by A and C, and label first and second equivalence points as B and D.

(b)   Estimate the formal concentration of H2A from the graph.

 

 

(c)    There are simple expressions for the pH at points A, B, and C. Give the equations below.

 

(d)   Estimate Ka1 and Ka2 for this acid from your estimate of pH at points A, B, and C.

 

 

(e)    At what indicator pH would you want your color-change indicator to have in order to titrate to the second equivalence point?

At pH=10