Skip to main content
Chemistry LibreTexts

12.9: Acids and Bases (Exercises)

  • Page ID
    172974
  • \( \newcommand{\vecs}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)

    \( \newcommand{\vecd}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash {#1}}} \)

    \( \newcommand{\id}{\mathrm{id}}\) \( \newcommand{\Span}{\mathrm{span}}\)

    ( \newcommand{\kernel}{\mathrm{null}\,}\) \( \newcommand{\range}{\mathrm{range}\,}\)

    \( \newcommand{\RealPart}{\mathrm{Re}}\) \( \newcommand{\ImaginaryPart}{\mathrm{Im}}\)

    \( \newcommand{\Argument}{\mathrm{Arg}}\) \( \newcommand{\norm}[1]{\| #1 \|}\)

    \( \newcommand{\inner}[2]{\langle #1, #2 \rangle}\)

    \( \newcommand{\Span}{\mathrm{span}}\)

    \( \newcommand{\id}{\mathrm{id}}\)

    \( \newcommand{\Span}{\mathrm{span}}\)

    \( \newcommand{\kernel}{\mathrm{null}\,}\)

    \( \newcommand{\range}{\mathrm{range}\,}\)

    \( \newcommand{\RealPart}{\mathrm{Re}}\)

    \( \newcommand{\ImaginaryPart}{\mathrm{Im}}\)

    \( \newcommand{\Argument}{\mathrm{Arg}}\)

    \( \newcommand{\norm}[1]{\| #1 \|}\)

    \( \newcommand{\inner}[2]{\langle #1, #2 \rangle}\)

    \( \newcommand{\Span}{\mathrm{span}}\) \( \newcommand{\AA}{\unicode[.8,0]{x212B}}\)

    \( \newcommand{\vectorA}[1]{\vec{#1}}      % arrow\)

    \( \newcommand{\vectorAt}[1]{\vec{\text{#1}}}      % arrow\)

    \( \newcommand{\vectorB}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)

    \( \newcommand{\vectorC}[1]{\textbf{#1}} \)

    \( \newcommand{\vectorD}[1]{\overrightarrow{#1}} \)

    \( \newcommand{\vectorDt}[1]{\overrightarrow{\text{#1}}} \)

    \( \newcommand{\vectE}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash{\mathbf {#1}}}} \)

    \( \newcommand{\vecs}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)

    \( \newcommand{\vecd}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash {#1}}} \)

    Exercises (Arrhenius Acids and Bases)

    1. Define Arrhenius acid.
    2. Define Arrhenius base.

    3. What are some general properties of Arrhenius acids?

    1. What are some general properties of Arrhenius bases?
       
    2. Identify each substance as an Arrhenius acid, an Arrhenius base, or neither.
      1. NaOH
      2. C2H5OH
      3. H3PO4
         
    3. Identify each substance as an Arrhenius acid, an Arrhenius base, or neither.
      1. C6H12O6
      2. HNO2
      3. Ba(OH)2
         
    4. Write the balanced chemical equation for the neutralization reaction between KOH and H2C2O4. What is the salt?
       
    5. Write the balanced chemical equation for the neutralization reaction between Sr(OH)2 and H3PO4. What is the salt?
       
    6. Write the balanced chemical equation for the neutralization reaction between HCl and Fe(OH)3. What is the salt?
       
    7. Write the balanced chemical equation for the neutralization reaction between H2SO4 and Cr(OH)3. What is the salt?
       
    8. CaCl2 would be the product of the reaction of what acid and what base?
       
    9. Zn(NO3)2 would be product of the reaction of what acid and what base?
       
    10. BaSO4 would be product of the reaction of what acid and what base?
       
    11. Na3PO4 would be product of the reaction of what acid and what base?

    Answers

    1.   a compound that increases the H+ concentration in water

    3.   sour taste, react with metals, and turn litmus red

    5. 
    1. Arrhenius base
    2. neither
    3. Arrhenius acid

    7.   2KOH + H2C2O4 → 2H2O + K2C2O4; K2C2O4

    9.   3HCl + Fe(OH)3 → 3H2O + FeCl3; FeCl3

    11.   HCl and Ca(OH)2

    13.   H2SO4 and Ba(OH)2

    Exercises (Brønsted-Lowry Acids and Bases)

    1. Define Brønsted-Lowry acid. How does it differ from an Arrhenius acid?
    2. Define Brønsted-Lowry base. How does it differ from an Arrhenius base?

    3. Write the dissociation of hydrogen bromide in water as a Brønsted-Lowry acid-base reaction and identify the proton donor and proton acceptor.

    4. Write the dissociation of nitric acid in water as a Brønsted-Lowry acid-base reaction and identify the proton donor and proton acceptor.

    5. Pyridine (C5H5N) acts as a Brønsted-Lowry base in water. Write the hydrolysis reaction for pyridine and identify the Brønsted-Lowry acid and Brønsted-Lowry base.

    6. The methoxide ion (CH3O) acts as a Brønsted-Lowry base in water. Write the hydrolysis reaction for the methoxide ion and identify the Brønsted-Lowry acid and Brønsted-Lowry base.

    7. Identify the Brønsted-Lowry acid and Brønsted-Lowry base in this chemical equation.

      H3PO4 + OH → H2PO4 + H2O

    8. Identify the Brønsted-Lowry acid and Brønsted-Lowry base in this chemical equation.

      H2C2O4 + 2F → 2HF + C2O42

    9. Predict the products of this reaction, assuming it undergoes a Brønsted-Lowry acid-base reaction.

      HC2H3O2 + C5H5N → ?

    10. Predict the products of this reaction, assuming it undergoes a Brønsted-Lowry acid-base reaction.

      (C2H5)3N + H2O → ?

    11. What is the conjugate acid of H2O? of NH3?

    12. What is the conjugate acid of H2PO4? of NO3?

    13. What is the conjugate base of HSO4? of H2O?

    14. What is the conjugate base of H3O+? of H2SO4?

    15. Identify the conjugate acid-base pairs in this reaction.

      HSO4 + PO43 → SO42 + HPO42

    16. Identify the conjugate acid-base pairs in this reaction.

      HClO3 + (C2H5)3N → ClO3 + (C2H5)3NH+

    17. Identify the conjugate acid-base pairs in this reaction.

      NH3 + C6H5O → C6H5OH + NH2

    18. Identify the conjugate acid-base pairs in this reaction.

      C5H5NH+ + C2O42 → C5H5N + HC2O4suming it undergoes a Brønsted-Lowry acid-base reaction.

    Answers

    1.   A Brønsted-Lowry acid is a proton donor. It does not necessarily increase the H+ concentration in water.

    3.   HBr + H2O → H3O+ + Br; PD: HBr; PA: H2O

    5.   C5H5N + H2O → C5H5NH+ + OH; PD: H2O; PA: C5H5N

    7.   BL acid: H3PO4; BL base: OH

    9.   C2H3O2 and C5H5NH+

    11. H3O+; NH4+

    13. SO42; OH

    15. HSO4 and SO42; PO43 and HPO42

    17. NH3 and NH2; C6H5O and C6H5OH

    Exercises (Acid-Base Titrations)

    1. Define titration.
    2. What is the difference between the titrant and the analyte?

    3. True or false: An acid is always the titrant. Explain your answer.

    4. True or false: An analyte is always dissolved before reaction. Explain your answer.
    5. If 55.60 mL of 0.2221 M HCl was needed to titrate a sample of NaOH to its equivalence point, what mass of NaOH was present?

    6. If 16.33 mL of 0.6664 M KOH was needed to titrate a sample of HC2H3O2 to its equivalence point, what mass of HC2H3O2 was present?
       
    7. It takes 45.66 mL of 0.1126 M HBr to titrate 25.00 mL of Ca(OH)2 to its equivalence point. What is the original concentration of the Ca(OH)2 solution?
       
    8. It takes 9.77 mL of 0.883 M H2SO4 to titrate 15.00 mL of KOH to its equivalence point. What is the original concentration of the KOH solution?

    Answers

    1. a chemical reaction performed in a quantitative fashion

    3. False; a base can be a titrant, or the reaction being performed may not even be an acid-base reaction.

    5. 0.494 g

    7. 0.1028 M

    Exercises (Strong and Weak Acids and Bases and their Salts)

    1. Differentiate between a strong acid and a weak acid.
    1. Differentiate between a strong base and a weak base.
       
    2. Identify each as a strong acid or a weak acid. Assume aqueous solutions.
      1. HF
      2. HCl
      3. HC2O4
         
    3. Identify each as a strong base or a weak base. Assume aqueous solutions.
      1. NaOH
      2. Al(OH)3
      3. C4H9NH2
         
    4. Write a chemical equation for the ionization of each acid and indicate whether it proceeds 100% to products or not.
      1. HNO3
      2. HNO2
      3. HI3
         
    5. Write a chemical equation for the ionization of each base and indicate whether it proceeds 100% to products or not.
      1. NH3
      2. (CH3)3N
      3. Mg(OH)2
         
    6. Write the balanced chemical equation for the reaction of each acid and base pair.
      1. HCl + C5H5N
      2. H2C2O4 + NH3
      3. HNO2 + C7H9N
         
    7. Write the balanced chemical equation for the reaction of each acid and base pair.
      1. H3C5H5O7 + Mg(OH)2
      2. HC3H3O3 + (CH3)3N
      3. HBr + Fe(OH)3
         
    8. Identify each salt as neutral, acidic, or basic.
      1. NaBr
      2. Fe(NO3)2
      3. Fe(NO3)3
         
    9. Identify each salt as neutral, acidic, or basic.
      1. NH4I
      2. C2H5NH3Cl
      3. KI
         
    10. Identify each salt as neutral, acidic, or basic.
      1. NaNO2
      2. NaNO3
      3. NH4NO3
         
    11. Identify each salt as neutral, acidic, or basic.
      1. KC2H3O2
      2. KHSO4
      3. KClO3
         
    12. Write the hydrolysis reaction that occurs, if any, when each salt dissolves in water.
      1. K2SO3
      2. KI
      3. NH4ClO3
         
    13. Write the hydrolysis reaction that occurs, if any, when each salt dissolves in water.
      1. NaNO3
      2. CaC2O4
      3. C5H5NHCl
         
    14. When NH4NO2 dissolves in H2O, both ions hydrolyze. Write chemical equations for both reactions. Can you tell if the solution will be acidic or basic overall?
       
    15. When pyridinium acetate (C5H5NHC2H3O2) dissolves in H2O, both ions hydrolyze. Write chemical equations for both reactions. Can you tell if the solution will be acidic or basic overall?
       
    16. A lab technician mixes a solution of 0.015 M Mg(OH)2. Is the resulting OH concentration greater than, equal to, or less than 0.015 M? Explain your answer.
       
    17. A lab technician mixes a solution of 0.55 M HNO3. Is the resulting H+ concentration greater than, equal to, or less than 0.55 M? Explain your answer.

    Answers

    1. A strong acid is 100% ionized in aqueous solution, whereas a weak acid is not 100% ionized.

    3. 

    1. weak acid
    2. strong acid
    3. weak acid
    5. 
    1. HNO3(aq) → H+(aq) + NO3(aq); proceeds 100%
    2. HNO2(aq) → H+(aq) + NO2(aq); does not proceed 100%
    3. HI3(aq) → H+(aq) + I3(aq); does not proceed 100%
    7.
    1. HCl + C5H5N → Cl + C5H5NH+
    2. H2C2O4 + 2NH3 → C2O42 + 2NH4+
    3. HNO2 + C7H9N → NO2 + C7H9NH+
    9.
    1. neutral
    2. acidic
    3. acidic
    11.
    1. basic
    2. neutral
    3. acidic
    13.
     
    a. \(\mathrm{SO}_3{ }^{2-}+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{HSO}_3^{-}+\mathrm{OH}^{-}\)
    b. no reaction
    c. \(\mathrm{NH}_4{ }^{+}+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{NH}_3+\mathrm{H}_3 \mathrm{O}^{+}\)
     
    15. \(\mathrm{NH}_4{ }^{+}+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{NH}_3+\mathrm{H}_3 \mathrm{O}^{+} ; \mathrm{NO}_2^{-}+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{HNO}_2+\mathrm{OH}^{-}\); it is not possible to determine whether the solution will be acidic or basic.
     
    17. greater than \(0.015 \mathrm{M}\) because there are two \(\mathrm{OH}^{-}\)ions per formula unit of \(\operatorname{Mg}(\mathrm{OH})_2\)
     

    Exercise (Autoionization of Water)

     
    1. Does [H+] remain constant in all aqueous solutions? Why or why not?
    2. Does [OH] remain constant in all aqueous solutions? Why or why not?

    3. What is the relationship between [H+] and Kw? Write a mathematical expression that relates them.

    4. What is the relationship between [OH] and Kw? Write a mathematical expression that relates them.

    5. Write the chemical equation for the autoionization of water and label the conjugate acid-base pairs.

    6. Write the reverse of the reaction for the autoionization of water. It is still an acid-base reaction? If so, label the acid and base.

    7. For a given aqueous solution, if [H+] = 1.0 × 10−3 M, what is [OH]?

    8. For a given aqueous solution, if [H+] = 1.0 × 10−9 M, what is [OH]?

    9. For a given aqueous solution, if [H+] = 7.92 × 10−5 M, what is [OH]?

    10. For a given aqueous solution, if [H+] = 2.07 × 10−11 M, what is [H+]?

    11. For a given aqueous solution, if [OH] = 1.0 × 10−5 M, what is [H+]?

    12. For a given aqueous solution, if [OH] = 1.0 × 10−12 M, what is [H+]?

    13. For a given aqueous solution, if [OH] = 3.77 × 10−4 M, what is [H+]?

    14. For a given aqueous solution, if [OH] = 7.11 × 10−10 M, what is [H+]?

    15. What are [H+] and [OH] in a 0.344 M solution of HNO3?

    16. What are [H+] and [OH] in a 2.86 M solution of HBr?

    17. What are [H+] and [OH] in a 0.00338 M solution of KOH?

    18. What are [H+] and [OH] in a 6.02 × 10−4 M solution of Ca(OH)2?

    19. If HNO2 is dissociated only to an extent of 0.445%, what are [H+] and [OH] in a 0.307 M solution of HNO2?

    20. If (C2H5)2NH is dissociated only to an extent of 0.077%, what are [H+] and [OH] in a 0.0955 M solution of (C2H5)2NH?

    Answers

    1.   [H+] varies with the amount of acid or base in a solution.

    3.   \[\left [ H^{+} \right ]=\frac{K_{W}}{\left [ OH^{-} \right ]}\]

    5.   H2O + H2O → H3O+ + OH; H2O/H3O+ and H2O/OH

    7.   1.0 × 10−11 M

    9.   1.26 × 10−10 M

    11. 1.0 × 10−9 M

    13. 2.65 × 10−11 M

    15. [H+] = 0.344 M; [OH] = 2.91 × 10−14 M

    17. [OH] = 0.00338 M; [H+] = 2.96 × 10−12 M

    19. [H+] = 0.00137 M; [OH] = 7.32 × 10−12 M

    Exercises  (The pH Scale)

    1. Define pH. How is it related to pOH?
    2. Define pOH. How is it related to pH?

    3. What is the pH range for an acidic solution?

    4. What is the pH range for a basic solution?

    5. What is [H+] for a neutral solution?

    6. What is [OH] for a neutral solution? Compare your answer to Exercise 6. Does this make sense?

    1. Which substances in Table 12.7.1 are acidic?
       
    2. Which substances in Table 12.7.1 are basic?
       
    3. What is the pH of a solution when [H+] is 3.44 × 10−4 M?
    4. What is the pH of a solution when [H+] is 9.04 × 10−13 M?

    5. What is the pH of a solution when [OH] is 6.22 × 10−7 M?

    6. What is the pH of a solution when [OH] is 0.0222 M?

    7. What is the pOH of a solution when [H+] is 3.44 × 10−4 M?

    8. What is the pOH of a solution when [H+] is 9.04 × 10−13 M?

    9. What is the pOH of a solution when [OH] is 6.22 × 10−7 M?

    10. What is the pOH of a solution when [OH] is 0.0222 M?

    11. If a solution has a pH of 0.77, what is its pOH, [H+], and [OH]?

    12. If a solution has a pOH of 13.09, what is its pH, [H+], and [OH]?

    Answers

    1.   pH is the negative logarithm of [H+] and is equal to 14 − pOH.

    3.   pH < 7

    5.   1.0 × 10−7 M

    7.   Every entry above pure water is acidic.

    9.   3.46

    11. 7.79

    13. 10.546.21

    15. pOH = 13.23; [H+] = 1.70 × 10−1 M; [OH] = 5.89 × 10−14 M

    Exercises (Buffers)

    1. Define buffer. What two related chemical components are required to make a buffer?
    1. Can a buffer be made by combining a strong acid with a strong base? Why or why not?
       
    2. Which combinations of compounds can make a buffer? Assume aqueous solutions.
      1. HCl and NaCl
      2. HNO2 and NaNO2
      3. NH4NO3 and HNO3
      4. NH4NO3 and NH3
         
    3. Which combinations of compounds can make a buffer? Assume aqueous solutions.
      1. H3PO4 and Na3PO4
      2. NaHCO3 and Na2CO3
      3. NaNO3 and Ca(NO3)2
      4. HN3 and NH3
         
    4. For each combination in Exercise 3 that is a buffer, write the chemical equations for the reactions of the buffer components when a strong acid and a strong base is added.
       
    5. For each combination in Exercise 4 that is a buffer, write the chemical equations for the reactions of the buffer components when a strong acid and a strong base is added.
    6. The complete phosphate buffer system is based on four substances: H3PO4, H2PO4, HPO42, and PO43. What different buffer solutions can be made from these substances?
       
    7. Explain why NaBr cannot be a component in either an acidic or a basic buffer.
       
    8. Two solutions are made containing the same concentrations of solutes. One solution is composed of H3PO4 and Na3PO4, while the other is composed of HCN and NaCN. Which solution should have the larger capacity as a buffer?
       
    9. Two solutions are made containing the same concentrations of solutes. One solution is composed of NH3 and NH4NO3, while the other is composed of H2SO4 and Na2SO4. Which solution should have the larger capacity as a buffer?

    Answers

    1. A buffer is the combination of a weak acid or base and a salt of that weak acid or base.

    3.

    a. no
    b. yes
    c. no
    d. yes

    5. 3b: strong acid: \(\mathrm{NO}_2^{-}+\mathrm{H}^{+} \rightarrow \mathrm{HNO}_2\); strong base: \(\mathrm{HNO}_2+\mathrm{OH}^{-} \rightarrow \mathrm{NO}_2^{-}+\mathrm{H}_2 \mathrm{O}\); 3d: strong base: \(\mathrm{NH}_4{ }^{+}+\mathrm{OH}^{-} \rightarrow \mathrm{NH}_3+\mathrm{H}_2 \mathrm{O}\); strong acid: \(\mathrm{NH}_3+\mathrm{H}^{+} \rightarrow \mathrm{NH}_4{ }^{+}\)

    7. Buffers can be made from three combinations: (1) \(\mathrm{H}_3 \mathrm{PO}_4\) and \(\mathrm{H}_2 \mathrm{PO}_4^{-}\), (2) \(\mathrm{H}_2 \mathrm{PO}_4{ }^{-}\)and \(\mathrm{HPO}_4{ }^{2-}\), and (3) \(\mathrm{HPO}_4{ }^{2-}\) and \(\mathrm{PO}_4{ }^{3-}\). (Technically, a buffer can be made from any two components.)

    9. The phosphate buffer should have the larger capacity.

    Additional Exercises

    1. Write the balanced chemical equation between \(\mathrm{Zn}\) metal and \(\mathrm{HCl}(\mathrm{aq})\). The other product is \(\mathrm{ZnCl}_2\).

    2. Write the neutralization reaction in which \(\mathrm{ZnCl}_2\), also found in Exercise 1, is the salt product.

    3. Why isn't an oxide compound like CaO considered a salt? (Hint: what acid-base combination would be needed to make it if it were a salt?)

    4. Metal oxides are considered basic because they react with \(\mathrm{H}_2 \mathrm{O}\) to form \(\mathrm{OH}\) compounds. Write the chemical equation for a reaction that forms a base when \(\mathrm{CaO}\) is combined with \(\mathrm{H}_2 \mathrm{O}\).

    5. Write the balanced chemical equation between aluminum hydroxide and sulfuric acid.

    6. Write the balanced chemical equation between phosphoric acid and barium hydroxide.

    7. Write the equation for the chemical reaction that occurs when caffeine \(\left(\mathrm{C}_8 \mathrm{H}_{10} \mathrm{~N}_4 \mathrm{O}_2\right)\) acts as a Brønsted-Lowry base.

    8. Citric acid \(\left(\mathrm{C}_6 \mathrm{H}_8 \mathrm{O}_7\right)\) is the acid found in citrus fruits. It can lose a maximum of three \(\mathrm{H}^{+}\)ions in the presence of a base. Write the chemical equations for citric acid acting stepwise as a Brønsted-Lowry acid.

    9. Can an amphiprotic substance be a strong acid and a strong base at the same time? Explain your answer.

    10. Can an amphiprotic substance be a weak acid and a weak base at the same time? If so, explain why and give an example.

    11. Under what conditions will the equivalence point of a titration be slightly acidic?

    12. Under what conditions will the equivalence point of a titration be slightly basic?

    13. Write the chemical equation for the autoionization of \(\mathrm{NH}_3\).

    14. Write the chemical equation for the autoionization of HF.

    15. What is the \(\mathrm{pOH}\) range for an acidic solution?

    16. What is the pOH range for a basic solution?

    17. The concentration of commercial \(\mathrm{HCl}\) is about \(12 \mathrm{M}\). What is its \(\mathrm{pH}\) and \(\mathrm{pOH}\) ?

    18. The concentration of concentrated \(\mathrm{H}_2 \mathrm{SO}_4\) is about 18 M. Assuming only one \(\mathrm{H}^{+}\) comes off the \(\mathrm{H}_2 \mathrm{SO}_4\) molecule, what is its \(\mathrm{pH}\) and \(\mathrm{pOH}\) ? What would the \(\mathrm{pH}\) and pOH be if the second \(\mathrm{H}^{+}\)were also ionized?

    Answers

    1. \(\mathrm{Zn}+2 \mathrm{HCl} \rightarrow \mathrm{ZnCl}_2+\mathrm{H}_2\)

    3. The \(\mathrm{O}^{2-}\) ion would come from \(\mathrm{H}_2 \mathrm{O}\), which is not considered a classic acid in the Arrhenius sense.

    5. \(2 \mathrm{Al}(\mathrm{OH})_3+3 \mathrm{H}_2 \mathrm{SO}_4 \rightarrow \mathrm{Al}_2\left(\mathrm{SO}_4\right)_3+6 \mathrm{H}_2 \mathrm{O}\)

    7. \(\mathrm{C}_8 \mathrm{H}_{10} \mathrm{~N}_4 \mathrm{O}_2+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{C}_8 \mathrm{H}_{10} \mathrm{~N}_4 \mathrm{O}_2 \mathrm{H}^{+}+\mathrm{OH}^{-}\); the \(\mathrm{H}^{+}\)ion attaches to one of the \(\mathrm{N}\) atoms in the caffeine molecule.

    9. As a strong acid or base, an amphiprotic substance reacts \(100 \%\) as an acid or a base, so it cannot be a base or an acid at the same time.

    11. if the salt produced is an acidic salt

    13. \(\mathrm{NH}_3+\mathrm{NH}_3 \rightarrow \mathrm{NH}_4^{+}+\mathrm{NH}_2^{-}\)

    15. \(\mathrm{pOH}>7\)

    17. \(\mathrm{pH}=-1.08 ; \mathrm{pOH}=15.08\)


    12.9: Acids and Bases (Exercises) is shared under a CC BY-NC-SA 4.0 license and was authored, remixed, and/or curated by LibreTexts.

    • Was this article helpful?