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Final Exam

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    Name: ______________________________

    Section: _____________________________

    Student ID#:__________________________

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


    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.)

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


    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?


    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?


    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)?


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


    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. \(\ce{HSO4-}\)


    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).


    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


    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)


    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


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

    1. addition of NaOH
    2. addition of HCl
    3. addition of NaF


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

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


    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.


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

    Final Exam is shared under a not declared license and was authored, remixed, and/or curated by LibreTexts.

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