# Final Exam


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Section: _____________________________

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This "final exam" is due Wednesday, March 18th at 3:00 via online/Canvas or emial to dlarsen@ucdavis.edu

## Q9.1

Calculate the equilibrium concentrations of NO, O2, and NO2 in a mixture at 250 °C that results from the reaction of 0.20 M NO and 0.10 M O2. (Hint: K is large; assume the reaction goes to completion then comes back to equilibrium.)

$\ce{2NO}(g)+\ce{O2}(g)⇌\ce{2NO2}(g) \hspace{20px} K_c=2.3×10^5\textrm{ at 250 °C} \nonumber$

## Q9.2

Carbon reacts with water vapor at elevated temperatures.

$$\ce{C}(s)+\ce{H2O}(g)⇌\ce{CO}(g)+\ce{H2}(g) \hspace{20px} K_c=\textrm{0.2 at 1000 °C}$$

What is the concentration of CO in an equilibrium mixture with [H2O] = 0.500 M at 1000 °C?

## Q9.3

At 773 K the equilibrium constant for the Haber-Bosch process is 1.45x10-5. The chemical reaction involved in the Haber-Bosch process is:

$N_{2\, (g)} + 3H_{2\, (g)} \rightleftharpoons 2NH_{3\: (g)} \nonumber$

1. What is $$\Delta G°(773)$$ for this reaction?
2. What is $$\Delta G$$ at 773 K for transforming 1 mole of N2 and 3 moles H2 held at 30 atm to 2 moles of NH3 held at 1 atm?
3. In which direction would the previous reaction run spontaneously?

## Q9.4

At a temperature of 60 °C, the vapor pressure of water is 0.196 atm. What is the value of the equilibrium constant KP for the transformation at 60 °C?

$\ce{H2O}(l)⇌\ce{H2O}(g) \nonumber$

There are two different equations that you were introduced to address this problem. What are they and how do they differ (if at all)?

## Q9.5

Compare and contrast the three definitions of Acids and Bases (from readings). Give two examples of each.

## Q9.6

Show by suitable net ionic equations that each of the following species can act as a Brønsted-Lowry acid:

1. $$\ce{H3O+}$$
2. HCl
3. NH3
4. CH3CO2H
5. $$\ce{NH4+}$$
6. $$\ce{HSO4-}$$

## Q9.7

Acetic acid gives vinegar a sour taste and strong aroma. Its $$K_a$$ value is $$1.75 \times 10^{-5}$$. What is the pH of the solution if 0.59 grams of acetic acid is dissolved in 40 mL of water? (Hint: This is an ICE table, but you need to start with a balance reaction of course).

## Q9.8

Calculate the pH and the pOH of each of the following solutions at 25 °C for which the substances ionize completely:

1. 0.000259 M HClO4
2. 0.21 M NaOH
3. 0.000071 M Ba(OH)2
4. 2.5 M KOH

## Q9.9

Calculate the ionization constant for each of the following acids or bases from the ionization constant of its conjugate base or conjugate acid:

1. $$\ce{HTe^{−}}$$ (as a base)
2. $$\ce{(CH3)3NH+}$$ (as an acid)
3. $$\ce{HAsO4^2-}$$ (as a base)
4. $$\ce{HO2-}$$ (as a base)
5. $$\ce{C6H5NH3+}$$ (as an acid)
6. $$\ce{HSO3-}$$ (as a base)

## Q9.10

What is the effect on the concentrations of $$\ce{NO2-}$$, $$\ce{HNO2}$$, and $$\ce{OH^{−}}$$ when the following are added to a solution of $$\ce{KNO2}$$ in water:

1. HCl
2. HNO2
3. NaOH
4. NaCl
5. KNO

## Q9.11

Which of the following will increase the percent of HF that is converted to the fluoride ion in water?

## Q9.12

Determine whether aqueous solutions of the following salts are acidic, basic, or neutral:

1. Al(NO3)3
2. RbI
3. KHCO2
4. CH3NH3Br

## Q9.13

Which of the following concentrations would be practically equal in a calculation of the equilibrium concentrations in a 0.134-M solution of H2CO3, a diprotic acid:

• $$\ce{[H3O+]}$$,
• $$[OH^−]$$
• $$[H_2CO_3]$$
• $$\ce{[HCO3- ]}$$
• $$\ce{[CO3^2- ]}$$

No calculations are needed to answer this question.

## Q9.14

Explain why a buffer can be prepared from a mixture of NH4Cl and NaOH but not from NH3 and NaOH.

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