Chem Exam 3

Calculate the pH of the solution given the following.  Answer to two decimal places.

Initial concentration of 0.30 M NH4Cl.         

(NH3: Kb = 1.76x10-5)

4\.88 (with margin: 0.1)

Calculate the pH of the solution given the following.  Answer to two decimal places.

Initial concentration of 0.20 M HCN.       

(HCN: Ka = 4.9x10-10)

5 (with margin: 0.1)

Which of the following is strictly a Lewis acid

AlCl3

Which of the following salt would result in a slightly more basic solution?

KF

Which of the following bases is the strongest?

(CH3CH2)3N         Kb = 5.2 × 10^-4

Which of the following compound is the most basic?

 d. HCN (aq) pKa = 9.31

If a certain solution required a steady pH of 4, which of the following acids would make a good candidate for this job?

HCHO2 (aq) pKa = 3.74

Waters self-ionization constant is always 1x10^-14 

False

All Bronsted-Lowery acids are proton donors

True

Lewis bases are electron pair acceptors

False

Lewis acids are electron pair acceptors

True

Which of the following describes an basic solution?

 pOH = 2.2

Which of the following compounds can be an amphiprotic compound?

H20

Which of the following solids would have the best solubility

 PbBr2 Ksp = 4.67 x 10^-6

In regards to solubility.  When do you expect to see a precipitate form during a reaction?

When Q > Ksp

The end point of a titration is __________.

When you see the signal to stop the titration by pH indicator.

Which of the following describes the correct conjugate acid/base pairs for the equation below?

HC2H3O2 + NH3 ↔ NH4+ + C2H3O2-

 

(HC2H3O2/ C2H3O2-) and (NH3/ NH4+)

Which of the following titration curves best represents a weak acid titrated with a strong base?

C

With the common ion effect, concerning molar solubility of a solid;

 Molar solubility decrease with the common ion effect.

For a weak acid, you would expect its Ka value to be:

Smaller than 1

Rank #1 through #3 each of the following in terms of **increasing** acidity (#1 = most acidic)

Hi       1             

**HCl**  2             

**HF**     3              

Rank #1 through #3 each of the following in terms of **increasing** acidity (#1 = most acidic)

**HClO3**         1             

**HBrO3**        2             

**HIO2**     3             

Rank #1 through #3 each of the following in terms of **increasing** acidity (#1 = most acidic)

HBr      1             

**HBrO**    3             

**HBrO3**  2             

For a buffer solution that is made of 200.0 mL of 0.2000 M HClO with 100.0 mL of 0.300 M KClO. The Ka for HClO is 2.9 x 10-8.

a. What is the pH of the buffer? (6 pts)

 

b. To this point, if 10.00 mL of 0.120 M NaOH is added, what is the new pH? (4pts)

a) pH = 7.41

b) pH = 7.44

Calculate the molar solubility of AgCl in a solution containing 0.18 M AgNO3 (Ksp = 1.8 x 10^-10).

Molar Solubility = 1.0 x 10^-9

Consider the titration of a 35.0 mL sample of 0.175 M HBr with 0.200 M KOH.

Determine each quantity. 

(a) The initial pH.

(b) The volume of base added required to reach the equivalence point.

(c) The pH at 10.0 mL of added base.

(d) The pH at the equivalence point.

a) 0.76

b) 31.0 mL or 0.031 L 

c) 1.04

d) 7

Starting with 250 mL of a 0.250 M solution of acetic acid (CH3COOH, Ka = 1.8 x 10-5)

(a) Calculate the initial pH of the solution.

(b) Calculate the pH after adding 250 mL of 0.125M NaOH

(c) Calculate the pH after adding 500 mL of 0.125M NaOH

a) 2.68

b) 4.74

c) 8.83

A solution containing sodium hydroxide is mixed with one containing barium nitrate to form a solution that is 0.240 M in NaOH and 0.120 M in Ba(NO3)2.

Calculate the reaction quotient and determine if a precipitate will form or not.

(Ba(OH)2, Ksp = 5.0 x 10-3)

a) Reaction quotient: Q = 6.912 x 10^-3

b) Since Q > Ksp or 6.912 x 10^-3 > 5.0 x 10^-3 precipitate will form.