Molecular Structure of Acids and Bases - AP Chemistry

1. Which is the conjugate base of H₂SO₄?

   HSO₄⁻. $H_2SO_4$ loses one proton to form $HSO_4^-$.

2. List one example of a strong acid.

   Hydrochloric acid (HCl). One of the seven common strong acids.

3. What is the Lewis definition of an acid?

   A species that accepts an electron pair. Lewis theory focuses on electron pair interactions, not protons.

4. What is the general formula for calculating pH?

   $pH = -\log[H^+]$. Negative log of hydrogen ion concentration.

5. What is the definition of a Bronsted-Lowry acid?

   A species that donates a proton (H⁺). This is the proton transfer definition used in acid-base chemistry.

6. Identify the conjugate base of HCl.

   Cl⁻. Formed when HCl loses its proton ($H^+$).

7. What is the general formula for calculating pH?

   $pH = -\log[H^+]$. Negative log of hydrogen ion concentration.

8. What is the primary component of stomach acid?

   Hydrochloric acid (HCl). Gastric acid is primarily HCl with pH around 1-2.

9. What is the Henderson-Hasselbalch equation for pOH?

   $pOH = pK_b + \log\left(\frac{[BH⁺]}{[B]}\right)$. Relates buffer pOH to base/acid ratio.

10. What is the definition of a Bronsted-Lowry base?

    A species that accepts a proton (H⁺). This is the proton acceptor definition in Bronsted-Lowry theory.

11. What does a small $K_a$ value indicate about an acid?

    The acid is weak, dissociating less in solution. Small $K_a$ means lower ionization tendency.

12. What is the effect of dilution on the pH of a buffer?

    Minimal change in pH. Buffer capacity maintains pH despite concentration changes.

13. Identify the Lewis acid in AlCl₃ + Cl⁻ → AlCl₄⁻.

    AlCl₃. $AlCl_3$ accepts the electron pair from $Cl^-$.

14. Define the term 'amphoteric'.

    A species that can act as both an acid and a base. Can donate or accept protons depending on conditions.

15. Which is stronger: HCl or HF?

    HCl. HCl is a strong acid; HF is weak.

16. State the relationship between $K_a$ and $K_b$ for conjugate pairs.

    $K_a \times K_b = K_w$. This relationship connects conjugate acid-base pairs.

17. What does a large $K_a$ value indicate about an acid?

    The acid is strong, dissociating more in solution. Large $K_a$ means greater ionization tendency.

18. What does a small $K_a$ value indicate about an acid?

    The acid is weak, dissociating less in solution. Small $K_a$ means lower ionization tendency.

19. What is the Lewis definition of an acid?

    A species that accepts an electron pair. Lewis theory focuses on electron pair interactions, not protons.

20. Identify the conjugate base of HCl.

    Cl⁻. Formed when HCl loses its proton ($H^+$).

21. Identify the Lewis base in NH₃ + BF₃ → NH₃BF₃.

    NH₃. $NH_3$ donates its lone pair to $BF_3$.

22. What is a buffer solution?

    A solution that resists changes in pH when acids or bases are added. Contains weak acid/base and conjugate pairs.

23. What is the definition of a Bronsted-Lowry base?

    A species that accepts a proton (H⁺). This is the proton acceptor definition in Bronsted-Lowry theory.

24. List one example of a strong base.

    Sodium hydroxide (NaOH). One of the common strong bases.

25. Which is the conjugate base of H₂SO₄?

    HSO₄⁻. $H_2SO_4$ loses one proton to form $HSO_4^-$.

26. What is the Lewis definition of an acid?

    A species that accepts an electron pair. Lewis theory focuses on electron pair interactions, not protons.

27. List one example of a weak base.

    Ammonia (NH₃). A common weak base that accepts protons from water.

28. List one example of a weak base.

    Ammonia (NH₃). A common weak base that accepts protons from water.

29. What is the formula for calculating pOH?

    $pOH = -\log[OH⁻]$. Negative log of hydroxide ion concentration.

30. Which is stronger: HCl or HF?

    HCl. HCl is a strong acid; HF is weak.

31. What is the Henderson-Hasselbalch equation for pH?

    $pH = pK_a + \log\left(\frac{[A⁻]}{[HA]}\right)$. Relates buffer pH to acid/base ratio.

32. What is the common ion effect?

    The shift in equilibrium due to the addition of a common ion. Suppresses ionization of weak acids or bases.

33. What is the definition of a Bronsted-Lowry base?

    A species that accepts a proton (H⁺). This is the proton acceptor definition in Bronsted-Lowry theory.

34. Define the term 'weak acid'.

    An acid that partially dissociates in solution. Only a small percentage ionizes in aqueous solution.

35. Define the term 'weak acid'.

    An acid that partially dissociates in solution. Only a small percentage ionizes in aqueous solution.

36. What does a small $K_a$ value indicate about an acid?

    The acid is weak, dissociating less in solution. Small $K_a$ means lower ionization tendency.

37. What is the Henderson-Hasselbalch equation for pOH?

    $pOH = pK_b + \log\left(\frac{[BH^+]}{[B]}\right)$. Relates buffer pOH to base/acid ratio.

38. Identify the Lewis acid in AlCl₃ + Cl⁻ → AlCl₄⁻.

    AlCl₃. $AlCl_3$ accepts the electron pair from $Cl^-$.

39. Identify an example of an amphoteric substance.

    Water (H₂O). Can act as acid or base in different reactions.

40. Identify the Lewis base in NH₃ + BF₃ → NH₃BF₃.

    NH₃. $NH_3$ donates its lone pair to $BF_3$.

41. Identify the conjugate base of HCl.

    Cl⁻. Formed when HCl loses its proton ($H^+$).

42. What does a large $K_a$ value indicate about an acid?

    The acid is strong, dissociating more in solution. Large $K_a$ means greater ionization tendency.

43. Define the term 'strong base'.

    A base that completely dissociates in solution. Nearly 100% ionization occurs in aqueous solution.

44. Which is stronger: HCl or HF?

    HCl. HCl is a strong acid; HF is weak.

45. Identify the common ion in a mixture of HCl and NaCl.

    Cl⁻. Both compounds contain chloride ions.

46. What is the pH of a solution with $[H⁺] = 1 \times 10^{-3}$ M?

    3. Using $pH = -\log(1 \times 10^{-3}) = 3$.

47. Identify an example of an amphoteric substance.

    Water (H₂O). Can act as acid or base in different reactions.

48. State the relationship between pH and pOH.

    $pH + pOH = 14$. This relationship holds at 25°C due to $K_w$.

49. What is the Lewis definition of a base?

    A species that donates an electron pair. Lewis bases are electron pair donors in chemical reactions.

50. Define the term 'amphoteric'.

    A species that can act as both an acid and a base. Can donate or accept protons depending on conditions.

51. List one example of a weak base.

    Ammonia (NH₃). A common weak base that accepts protons from water.

52. Define the term 'strong acid'.

    An acid that completely dissociates in solution. Nearly 100% ionization occurs in aqueous solution.

53. Identify the Lewis base in NH₃ + BF₃ → NH₃BF₃.

    NH₃. $NH_3$ donates its lone pair to $BF_3$.

54. Identify the Lewis acid in AlCl₃ + Cl⁻ → AlCl₄⁻.

    AlCl₃. $AlCl_3$ accepts the electron pair from $Cl^-$.

55. What is the Lewis definition of a base?

    A species that donates an electron pair. Lewis bases are electron pair donors in chemical reactions.

56. What does a large $K_a$ value indicate about an acid?

    The acid is strong, dissociating more in solution. Large $K_a$ means greater ionization tendency.

57. What is the pH of a buffer with equal concentrations of acid and base?

    pH = pK_a. When $\log(1) = 0$, so pH equals $pK_a$.

58. What is the pOH of a solution with $[OH⁻] = 1 \times 10^{-5}$ M?

    5. Using $pOH = -\log(1 \times 10^{-5}) = 5$.

59. What is the primary component of stomach acid?

    Hydrochloric acid (HCl). Gastric acid is primarily HCl with pH around 1-2.

60. What is the Henderson-Hasselbalch equation for pH?

    $pH = pK_a + \log\left(\frac{[A⁻]}{[HA]}\right)$. Relates buffer pH to acid/base ratio.

61. What is the effect of dilution on the pH of a buffer?

    Minimal change in pH. Buffer capacity maintains pH despite concentration changes.

62. What is a buffer solution?

    A solution that resists changes in pH when acids or bases are added. Contains weak acid/base and conjugate pairs.

63. What is the formula for the base dissociation constant, $K_b$?

    $K_b = \frac{[BH^+][OH^-]}{[B]}$. Measures the extent of base dissociation in solution.

64. Define the term 'amphoteric'.

    A species that can act as both an acid and a base. Can donate or accept protons depending on conditions.

65. What is the pH of a buffer with equal concentrations of acid and base?

    pH = pK_a. When $\log(1) = 0$, so pH equals $pK_a$.

66. List one example of a strong acid.

    Hydrochloric acid (HCl). One of the seven common strong acids.

67. What is the formula for calculating pOH?

    $pOH = -\log[OH^{-}]$. Negative log of hydroxide ion concentration.

68. What is the primary component of stomach acid?

    Hydrochloric acid (HCl). Gastric acid is primarily HCl with pH around 1-2.

69. State the relationship between $K_a$ and $K_b$ for conjugate pairs.

    $K_a \times K_b = K_w$. This relationship connects conjugate acid-base pairs.

70. What is the formula for the base dissociation constant, $K_b$?

    $K_b = \frac{[BH⁺][OH⁻]}{[B]}$. Measures the extent of base dissociation in solution.

71. What is the Henderson-Hasselbalch equation for pOH?

    $pOH = pK_b + \log\left(\frac{[BH⁺]}{[B]}\right)$. Relates buffer pOH to base/acid ratio.

72. What is the general formula for calculating pH?

    $pH = -\log[H⁺]$. Negative log of hydrogen ion concentration.

73. What is a buffer solution?

    A solution that resists changes in pH when acids or bases are added. Contains weak acid/base and conjugate pairs.

74. Define the term 'strong base'.

    A base that completely dissociates in solution. Nearly 100% ionization occurs in aqueous solution.

75. State the relationship between $K_a$ and $K_b$ for conjugate pairs.

    $K_a \times K_b = K_w$. This relationship connects conjugate acid-base pairs.

76. Define the term 'strong acid'.

    An acid that completely dissociates in solution. Nearly 100% ionization occurs in aqueous solution.

77. What is the effect of dilution on the pH of a buffer?

    Minimal change in pH. Buffer capacity maintains pH despite concentration changes.

78. What is the Lewis definition of a base?

    A species that donates an electron pair. Lewis bases are electron pair donors in chemical reactions.

79. What is the pH of a solution with $[H⁺] = 1 \times 10^{-3}$ M?

    3. Using $pH = -\log(1 \times 10^{-3}) = 3$.

80. Identify an example of an amphoteric substance.

    Water (H₂O). Can act as acid or base in different reactions.

81. Identify the common ion in a mixture of HCl and NaCl.

    Cl⁻. Both compounds contain chloride ions.

82. What is the common ion effect?

    The shift in equilibrium due to the addition of a common ion. Suppresses ionization of weak acids or bases.

83. Identify the conjugate acid of NH₃.

    NH₄⁺. Formed when $NH_3$ gains a proton ($H^+$).

84. Define the term 'strong acid'.

    An acid that completely dissociates in solution. Nearly 100% ionization occurs in aqueous solution.

85. List one example of a weak acid.

    Acetic acid (CH₃COOH). A common weak acid found in vinegar.

86. Define the term 'strong base'.

    A base that completely dissociates in solution. Nearly 100% ionization occurs in aqueous solution.

87. What is the pH of a buffer with equal concentrations of acid and base?

    pH = pK_a. When $\log(1) = 0$, so pH equals $pK_a$.

88. State the relationship between pH and pOH.

    $pH + pOH = 14$. This relationship holds at 25°C due to $K_w$.

89. Define the term 'weak base'.

    A base that partially dissociates in solution. Only a small percentage ionizes in aqueous solution.

90. State the relationship between pH and pOH.

    $pH + pOH = 14$. This relationship holds at 25°C due to $K_w$.

91. Define the term 'weak base'.

    A base that partially dissociates in solution. Only a small percentage ionizes in aqueous solution.

92. Identify the common ion in a mixture of HCl and NaCl.

    Cl⁻. Both compounds contain chloride ions.

93. Which is the conjugate base of H₂SO₄?

    HSO₄⁻. $H_2SO_4$ loses one proton to form $HSO_4^-$.

94. What is the definition of a Bronsted-Lowry acid?

    A species that donates a proton (H⁺). This is the proton transfer definition used in acid-base chemistry.

95. Define the term 'weak acid'.

    An acid that partially dissociates in solution. Only a small percentage ionizes in aqueous solution.

96. What is the pH of a solution with $[H⁺] = 1 \times 10^{-3}$ M?

    3. Using $pH = -\log(1 \times 10^{-3}) = 3$.

97. What is the formula for the base dissociation constant, $K_b$?

    $K_b = \frac{[BH⁺][OH⁻]}{[B]}$. Measures the extent of base dissociation in solution.

98. Define the term 'weak base'.

    A base that partially dissociates in solution. Only a small percentage ionizes in aqueous solution.

99. What is the value of $K_w$ at 25°C?

    $1.0 \times 10^{-14}$. The ion product constant for water at standard temperature.

100. What is the pOH of a solution with $[OH⁻] = 1 \times 10^{-5}$ M?

     5. Using $pOH = -\log(1 \times 10^{-5}) = 5$.

101. List one example of a strong acid.

     Hydrochloric acid (HCl). One of the seven common strong acids.

102. What is the Henderson-Hasselbalch equation for pH?

     $pH = pK_a + \log\left(\frac{[A^{-}]}{[HA]}\right)$. Relates buffer pH to acid/base ratio.

103. What is the pOH of a solution with $[OH⁻] = 1 \times 10^{-5}$ M?

     5. Using $pOH = -\log(1 \times 10^{-5}) = 5$.

104. What is the formula for calculating pOH?

     $pOH = -\log[OH⁻]$. Negative log of hydroxide ion concentration.

105. What is the definition of a Bronsted-Lowry acid?

     A species that donates a proton (H⁺). This is the proton transfer definition used in acid-base chemistry.

106. What is the formula for the acid dissociation constant, $K_a$?

     $K_a = \frac{[H⁺][A⁻]}{[HA]}$. Measures the extent of acid dissociation in solution.

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