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11.E: Acids and Bases

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    The following questions are related to the material covered in this chapter. For additional discussion on each topic, also check the links included in each heading.

    12.2: 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

      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

    12.3: Brønsted-Lowry Acids and Bases

    Exercise \(\PageIndex{1}\)
    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 + HC2O4−

    nswers

    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

    12.4: Acid-Base Titrations

    Exercise \(\PageIndex{1}\)
    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

    12.5: Strong and Weak Acids and Bases and their Salts

    Exercise \(\PageIndex{1}\)
    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. Write the hydrolysis reaction that occurs, if any, when each salt dissolves in water.
      1. K2SO3
      2. KI
      3. NH4ClO3
    9. Write the hydrolysis reaction that occurs, if any, when each salt dissolves in water.
      1. NaNO3
      2. CaC2O4
      3. C5H5NHCl
    10. 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?
    11. 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?
    12. 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.
    13. 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.
      1. weak acid
      2. strong acid
      3. weak acid
      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%
      1. HCl + C5H5N → Cl− + C5H5NH+
      2. H2C2O4 + 2NH3 → C2O42− + 2NH4+
      3. HNO2 + C7H9N → NO2− + C7H9NH+
      1. SO32− + H2O → HSO3− + OH−
      2. no reaction
      3. NH4+ + H2O → NH3 + H3O+
    11. NH4+ + H2O → NH3 + H3O+; NO2− + H2O → HNO2 + OH−; it is not possible to determine whether the solution will be acidic or basic.
    13. greater than 0.015 M because there are two OH− ions per formula unit of Mg(OH)2

    12.6: Autoionization of Water

    Exercise \(\PageIndex{4}\)
    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

    12.7: The pH Scale

    Exercise \(\PageIndex{1}\)
    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. 10.54

    15. 6.21

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

    12.8: Buffers

    Exercise \(\PageIndex{1}\)
    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.
      1. no
      2. yes
      3. no
      4. yes
    5. 3b: strong acid: NO2− + H+ → HNO2; strong base: HNO2 + OH− → NO2− + H2O; 3d: strong base: NH4+ + OH− → NH3 + H2O; strong acid: NH3 + H+ → NH4+

    7. Buffers can be made from three combinations: (1) H3PO4 and H2PO4−, (2) H2PO4− and HPO42−, and (3) HPO42− and PO43−. (Technically, a buffer can be made from any two components.)

    9. The phosphate buffer should have the larger capacity.

    This page titled 11.E: Acids and Bases is shared under a mixed license and was authored, remixed, and/or curated by Anonymous.