Chemistry 120

Problem Session 4

1.     The world record for lowest pressure (at sea level) was 658 mm Hg recorded inside typhoon Ida on September 24, 1958, in the Philippine Sea.  Convert this pressure to

a.    torr

           

 

b.    atm

           

 

c.     in Hg

           

 

d.    psi

           

 

2.    A sample of methane gas at 25.0oC and 2.40 atm pressure occupies a volume of 500.0 mL

a.    What would be the gas volume at 25.0oC and 1.70 atm pressure?

 

b.    What would be the volume at 100oC and 2.40 atm pressure?

 

c.     What would be the volume at 50.0oC and 955 torr?

d.    How many grams of methane are in the sample?

 

 

e.    How many molecules of methane are in the sample?

 

 

3.    A 0.595 gram sample of a gas occupies a volume of 250. mL at 750. torr pressure and a temperature of 20.0oC.  What is the molecular mass of the gas?

 

 

4.    What is the density in grams/liter of krypton gas at 25oC and 744 torr pressure?

 

 

5.    If 25.00 mL of a 3.00 M solution of HCl is added to 50.0 g of Ba(HCO3)2, how many mL of dry carbon dioxide gas will be collected at 25.0oC and 745 torr pressure?

2 HCl       +       Ba(HCO3)2  ˆ  2 H2O  +  2 CO2  +  BaCl2

I           0.075 mol         0.193 mol

D         -2x                  -x                          +2x                        +2x     +x

E         0                      0.193-x                                  0.075 mol

 

6.    If 357.0 mL of H2 gas are collected over water at 28.0 oC and 752 torr, how many moles of H2 were collected?

 

mol dry H2           = total pressure – pressure water vapor at 28oC (from pg 4-105 in lab book)

(a)  =752 torr – 28.3 torr = 714 torr

 

 

 

7.     If 25.00 mL of a 3.00 M solution of HCl is added to 50.0 g of Ba(HCO3)2, how many mL of dry carbon dioxide gas will be collected at 25.0oC and 745 torr pressure?

i)      2 HCl       +       Ba(HCO3)2  ˆ  2 H2O  +  2 CO2  +  BaCl2

I           0.075 mol         0.193 mol

D          -2x                   -x                          +2x             +2x      +x

E          0                      0.193-x                                    0.075 mol

 

 

8.     How many grams of Na2SO4 are required to prepare 25.00 mL of 0.200 M Na2SO4 solution?

 

 

9.     How many mL of 0.150 M CaCl2 contain 6.00 g of CaCl2?

 

 

10.  In 33.00 mLof 0.22 M (NH4)3PO4, there are how many

a.     Moles of (NH4)3PO4?

 

 

 

b.     Grams of phosphorous?

 

 

 

c.     Moles of phosphate ions?

d.     NH4+ ions?

11.  A solution is prepared by dissolving 81.15 grams of Na3PO4á12H2O in enough water to prepare 500.0 mL of solution.  The density of the solution is 1.077 g/mL.  Calculate

a.     The % anhydrous salt

 

 

b.     The % hydrated salt

 

 

 

c.     The molarity of Na3PO4

 

12.  If 15.00 g of a mixture of 25.00% CaCl2 and 75.00% inert material are dissolved in 1.00 liter of water, how many mL of 0.500 M silver nitrate solution would be required to precipitate all of the chloride according to the equation: 

2 AgNO3(aq)  +  CaCl2(aq)  ˆ  Ca(NO3)2(aq)  +  2 AgCl(s)

 

 

13.  A 1.00 gram sample of (NH4)2SO4 is added to 50.0 mL of 0.450 M (NH4)3PO4 solution.  Assuming no volume change, what are the resultant molarities of all ions in solution?

 

Mol  (NH4)2SO4 = 1.00 g x 1 mol (NH4)2SO4 /132.15g = 7.56e-3 mol (NH4)2SO4

Mol (NH4)3PO4 = 0.0500 L x 0.450 mol (NH4)3PO4/L = 0.0225 mol (NH4)3PO4

[NH4+1 ] = [2(7.56e-3 mol (NH4)2SO4) + 3(0.0225 mol(NH4)3PO4)]/0.0500 L = 1.65 M NH4+1

[SO4-2 ] = 7.56e-3 mol SO4-2 /0.0500 L = 0.151 M SO4-2

[PO4-3 ] = 0.0225 mol PO4-3/0.0500 L = 0.450 M PO4-3

14.  2.500 g of K3PO4 are dissolved in 100.0 mL of deionized water.  To this solution are added 75.00 mL of 0.300 M AgNO3 solution.

 

                  K3PO4(aq)  +  3 AgNO3(aq)  ˆ  3 KNO3(aq)  +  Ag3PO4(s)

 

            Initial   0.01178 mol   0.0225 mol

            Change   -X                 -3X                      +3X                         +X

X = 0.00750 mol

            Final    0.00428 mol       0 mol            0.0225 mol                  0.00750 mol

 

 

a.     How many grams of Ag3PO4 will precipitate?

 

 

b.     How many grams of which reactant remains unreacted?

 

 

 

c.     What is the final concentration of K+?

 

 

 

15.  When 35.00 mL of 0.100 M iron(III) chloride are added to 20.00 mL of 0.200 M potassium phosphate, a precipitate of iron(III) phosphate is observed to form.  Inventory the concentrations of all ions in solution and calculate the mass of the precipitate.

 

            Fe(Cl)3   +   K3PO4    ˆ      FePO4    +   3   KCl

 

I

3.50 e-3 mol

4.00e-3 mol

 

0 mol

0 mol

x = 3.50 e -3 mol
 
D

-x mol

-x mol

 

+x mol

+3x mol

E

0 mol

5.0e-4 mol

 

3.5 e-3 mol

1.05 e-2 mol

 

[Fe+3] = 0M

[PO4-3] = 5.0 e-4/0.055 = 0.0091 M

[K+] = 3(4.00 e-3)/0.055 L = 0.218 M

[Cl-] = 1.05 e-2/0.055 L = 0.191 M

 

3.5 e-3 mol FePO4 x 150.82 g FePO4/mol = 5.27 g FePO4

16.  A 35.00 mL aliquot of 0.1861 M sulfuric acid is mixed with 20.00 mL of a 0.3016 M barium hydroxide solution.  Write the balanced chemical equation for the reaction and determine [H3O+], [OH-], [Ba+2], [SO4-2], pH, and pOH.  Find the mass of Barium sulfate that precipitates.

 

H2SO4  +  Ba(OH)2  ˆ  2 H2O  +  BaSO4(s)

X = 6.032 e-3 mol
 
 


I

6.514 e -3 mol

6.032 e-3 mol

 

 

0 mol

D

-x mol

-x mol

 

+ 2x mol

+ 1x mol

E

4.82 e-4 mol

0 mol

 

 

6.032 e-3 mol

 

[Ba+2] = 0 M Ba+2

[SO4-1] = 4.82 e-4mol /0.05500 L = 8.18 e-3 M SO4-1

[H3O+1] = 2(4.82 e-4)/0.05500 L = 0.0175 M                                    pH = 1.756

[OH-1] = kw/[H3O+1] = 1.0 e-14/0.0175 = 5.7 e -13                          pOH = 12.24

mass BaSO4 = 6.032 e-3 mol * 233.4 g/mol = 1.408 g BaSO4

 

 

17.  Fill in the table below

Conjugate acid

Conjugate base

HC2H3O2

C2H3O2-1

HCN

CN-1

NH4+1

NH3

H3O+1

H2O

H2O

OH-1

CH3NH3+

CH3NH2

 

 

18.  Calculate the hydroxide ion concentration in a solution if 35.27 mL of it exactly neutralize 24.30 mL of a 0.07243 M solution of sulfuric acid.

 

 

19.  A solution is prepared by dissolving 0.1162 grams of KOH in 250.0 mL of solution.  Calculate

a.     [H+] = 1.2107e-12

b.     [OH-] = 8.284e-3

c.     pH = 11.9182

d.     pOH = 2.0818

20.  Fill in the table below

[H3O+]

[OH-]

pH

pOH

3.582 x 10-6

2.792e-9

5.4459

8.5541

1.539e-6

6.498 x 10-9

5.8128

8.1872

1.705e-4

5.864e-11

3.7682

10.2318

2.986e-5

3.349e-10

4.5249

9.4751

 

21.  A primary standard is made by dissolving 3.7627 grams of oxalic acid dehydrate in enough water to make 500.0 mL of solution.  A 25.00 mL aliquot of the oxalic acid solution requires 41.87 mL of potassium hydroxide solution to titrate it to a phenophalein end point.  A 25.00 mL sample of vinegar is titrated with 15.97 mL of the sodium hydroxide solution.  What is the concentration of acetic acid in the vinegar?

 

Concentration of oxalic acid standard

 

 

Concentration of potassium hydroxide solution

 

Concentration of Acetic acid

 

 

 

22.  Solution A is a solution of HNO3 with a pH of 2.80.  Solution B is a solution of NaOH with a pH of 11.30.  Calculate the pH of the following solutions:

 

Solution A pH = 2.80                                                             [H3O+] = 10-2.80 = 1.6e-3 M

Solution B pH = 11.30            pOH = 2.70                 [OH-1] = 10-2.70 = 2.0e-3 M

 

a.     10.00 mL of A and 5.00 mL of B

Text Box: This one is hard!!            HNO3  +  NaOH    ˆ    H2O    +    NaNO3

I

1.6e-5 mol

1.0e-5 mol

 

 

D

-x mol

-x mol

 

 

E

0.6e-5 mol

0 mol

 

 

[H3O+] = 0.6 e-5mol/0.015 L = 4e-4 M

pH = 3.40

 

b.     10.00 mL of B and 5.00 mL of A

            HNO3  +  NaOH    ˆ    H2O    +    NaNO3

I

0.8e-5 mol

2.0e-5 mol

 

 

D

-x mol

-x mol

 

 

E

0mol

1.2e-5 mol

 

 

[OH-1] = 1.2 e-5mol/0.015 L = 8e-4 M

pOH = 3.10

pH = 10.90

 

c.     How many mL of A must be added to 25.00 mL of B to make a solution of pH = 7.00?

pH = 7 when {H3O+] = {OH-1] at the equivalence point of a titration

 

We need to find out how many mL of A are required to completely react with 25.00 mL of B