pH and Solubility - AP Chemistry

1. Identify the pH of a neutral solution at 25°C.

   pH = 7. Equal concentrations of $[H^+]$ and $[OH^-]$.

2. Determine the pOH of a solution with $[OH^-] = 1 \times 10^{-4} \text{ M}$.

   pOH = 4. $pOH = -\log(10^{-4}) = 4$

3. What is the solubility product constant for $PbI_2$?

   $K_{sp} = [Pb^{2+}][I^-]^2$. Iodide ion squared due to formula stoichiometry.

4. What is the pH of a $0.1 \text{ M}$ $NaOH$ solution?

   pH = 13. Strong base: $[OH^-] = 0.1$, $pOH = 1$

5. Identify the pH of a neutral solution at 25°C.

   pH = 7. Equal concentrations of $[H^+]$ and $[OH^-]$.

6. State the $K_{sp}$ expression for $ZnS$.

   $K_{sp} = [Zn^{2+}][S^{2-}]$. 1:1 stoichiometry for ion products.

7. What is the solubility product constant for $PbI_2$?

   $K_{sp} = [Pb^{2+}][I^-]^2$. Iodide ion squared due to formula stoichiometry.

8. What is the definition of the solubility product constant ($K_{sp}$)?

   $K_{sp}$ is the equilibrium constant for a solid dissolving in water. Represents dissolution equilibrium of ionic solids.

9. What is the definition of the solubility product constant ($K_{sp}$)?

   $K_{sp}$ is the equilibrium constant for a solid dissolving in water. Represents dissolution equilibrium of ionic solids.

10. State the formula for calculating pOH from $[OH^-]$.

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

11. Determine the pH of a $0.025 \text{ M}$ $HBr$ solution.

    pH = 1.6. $pH = -\log(0.025) = 1.6$

12. Which solution has a higher pH: $0.1 \text{ M}$ $HCl$ or $0.1 \text{ M}$ $NaOH$?

    $0.1 \text{ M}$ $NaOH$. $NaOH$ is basic with higher pH value.

13. Find the pH of a $0.075 \text{ M}$ $Ba(OH)_2$ solution.

    pH = 13.18. Diprotic base: $[OH^-] = 0.15$, $pOH = 0.82$

14. What is the $K_{sp}$ expression for $CaF_2$?

    $K_{sp} = [Ca^{2+}][F^-]^2$. Fluoride ion squared due to formula stoichiometry.

15. State the formula for converting pOH to pH.

    $pH = 14 - pOH$. Derived from $pH + pOH = 14$ relationship.

16. What is the $K_{sp}$ expression for $CuS$?

    $K_{sp} = [Cu^{2+}][S^{2-}]$. 1:1 stoichiometry for ion products.

17. What is the pH of a $0.02 \text{ M}$ $KOH$ solution?

    pH = 12.3. $[OH^-] = 0.02$, $pOH = 1.7$

18. What is the $K_{sp}$ expression for $PbSO_4$?

    $K_{sp} = [Pb^{2+}][SO_4^{2-}]$. 1:1 stoichiometry for ion products.

19. What is the $K_{sp}$ expression for $Mg(OH)_2$?

    $K_{sp} = [Mg^{2+}][OH^-]^2$. Hydroxide ion squared due to formula stoichiometry.

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

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

21. State the ion product constant for water ($K_w$) at 25°C.

    $K_w = 1.0 \times 10^{-14}$. Equilibrium constant for water autoionization.

22. Identify the $K_{sp}$ expression for $BaSO_4$.

    $K_{sp} = [Ba^{2+}][SO_4^{2-}]$. 1:1 stoichiometry for ion products.

23. Identify the $K_{sp}$ expression for $BaSO_4$.

    $K_{sp} = [Ba^{2+}][SO_4^{2-}]$. 1:1 stoichiometry for ion products.

24. What is the pH of a $0.02 \text{ M}$ $KOH$ solution?

    pH = 12.3. $[OH^-] = 0.02$, $pOH = 1.7$

25. What is the $K_{sp}$ expression for $CaCO_3$?

    $K_{sp} = [Ca^{2+}][CO_3^{2-}]$. 1:1 stoichiometry for ion products.

26. What is the $K_{sp}$ expression for $Ag_2CrO_4$?

    $K_{sp} = [Ag^+]^2[CrO_4^{2-}]$. Silver ion squared due to formula stoichiometry.

27. What is the relationship between pH and pOH at 25°C?

    $pH + pOH = 14$. Sum equals 14 at standard temperature.

28. What is the definition of pH?

    pH is the negative logarithm of the hydrogen ion concentration: $pH = -\text{log}[H^+]$. It measures acidity using logarithmic scale.

29. State the formula for calculating pH from $[H^+]$.

    $pH = -\text{log}[H^+]$. Take negative log of hydrogen ion concentration.

30. What is the relationship between pH and pOH at 25°C?

    $pH + pOH = 14$. Sum equals 14 at standard temperature.

31. What is the solubility product constant ($K_{sp}$) expression for $AgCl$?

    $K_{sp} = [Ag^+][Cl^-]$. Products of ion concentrations at equilibrium.

32. What is the $K_{sp}$ expression for $Hg_2I_2$?

    $K_{sp} = [Hg_2^{2+}][I^-]^2$. Iodide ion squared due to formula stoichiometry.

33. What is the solubility product constant for $PbI_2$?

    $K_{sp} = [Pb^{2+}][I^-]^2$. Iodide ion squared due to formula stoichiometry.

34. What is the $K_{sp}$ expression for $Mg(OH)_2$?

    $K_{sp} = [Mg^{2+}][OH^-]^2$. Hydroxide ion squared due to formula stoichiometry.

35. Calculate the pH of a $0.1 \text{ M}$ $CH_3COOH$ solution, $K_a = 1.8 \times 10^{-5}$.

    pH ≈ 2.87. Weak acid requires ICE table calculation.

36. What is the definition of pH?

    pH is the negative logarithm of the hydrogen ion concentration: $pH = -\text{log}[H^+]$. It measures acidity using logarithmic scale.

37. Determine the pOH of a solution with $[OH^-] = 1 \times 10^{-4} \text{ M}$.

    pOH = 4. $pOH = -\log(10^{-4}) = 4$

38. What is the $K_{sp}$ expression for $CuS$?

    $K_{sp} = [Cu^{2+}][S^{2-}]$. 1:1 stoichiometry for ion products.

39. Find the pH of a $0.03 \text{ M}$ $LiOH$ solution.

    pH = 12.5. $[OH^-] = 0.03$, $pOH = 1.52$

40. Determine the pH of a $0.025 \text{ M}$ $HBr$ solution.

    pH = 1.6. $pH = -\log(0.025) = 1.6$

41. What is the $K_{sp}$ expression for $Ag_2CrO_4$?

    $K_{sp} = [Ag^+]^2[CrO_4^{2-}]$. Silver ion squared due to formula stoichiometry.

42. Determine the pOH of a solution with $[OH^-] = 1 \times 10^{-4} \text{ M}$.

    pOH = 4. $pOH = -\log(10^{-4}) = 4$

43. What is the $K_{sp}$ expression for $Ca(OH)_2$?

    $K_{sp} = [Ca^{2+}][OH^-]^2$. Hydroxide ion squared due to formula stoichiometry.

44. What is the $K_{sp}$ expression for $Hg_2I_2$?

    $K_{sp} = [Hg_2^{2+}][I^-]^2$. Iodide ion squared due to formula stoichiometry.

45. Which solution has a higher pH: $0.1 \text{ M}$ $HCl$ or $0.1 \text{ M}$ $NaOH$?

    $0.1 \text{ M}$ $NaOH$. $NaOH$ is basic with higher pH value.

46. Determine the $K_{sp}$ expression for $Fe(OH)_3$.

    $K_{sp} = [Fe^{3+}][OH^-]^3$. Hydroxide ion cubed due to formula stoichiometry.

47. What is the $K_{sp}$ expression for $Ca(OH)_2$?

    $K_{sp} = [Ca^{2+}][OH^-]^2$. Hydroxide ion squared due to formula stoichiometry.

48. What is the pH of a $0.01 \text{ M}$ $HCl$ solution?

    pH = 2. Strong acid completely ionizes: $[H^+] = 0.01$

49. Determine the $K_{sp}$ expression for $Fe(OH)_3$.

    $K_{sp} = [Fe^{3+}][OH^-]^3$. Hydroxide ion cubed due to formula stoichiometry.

50. Find the pH of a $0.03 \text{ M}$ $LiOH$ solution.

    pH = 12.5. $[OH^-] = 0.03$, $pOH = 1.52$

51. State the formula for calculating pOH from $[OH^-]$.

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

52. State the formula for converting pOH to pH.

    $pH = 14 - pOH$. Derived from $pH + pOH = 14$ relationship.

53. Calculate the pH of a $0.005 \text{ M}$ $HCl$ solution.

    pH = 2.3. $pH = -\log(0.005) = 2.3$

54. What is the definition of pH?

    pH is the negative logarithm of the hydrogen ion concentration: $pH = -\text{log}[H^+]$. It measures acidity using logarithmic scale.

55. What is the pH of a $0.1 \text{ M}$ $NaOH$ solution?

    pH = 13. Strong base: $[OH^-] = 0.1$, $pOH = 1$

56. What is the $K_{sp}$ expression for $Ca(OH)_2$?

    $K_{sp} = [Ca^{2+}][OH^-]^2$. Hydroxide ion squared due to formula stoichiometry.

57. Determine the $K_{sp}$ expression for $Fe(OH)_3$.

    $K_{sp} = [Fe^{3+}][OH^-]^3$. Hydroxide ion cubed due to formula stoichiometry.

58. Calculate the pH of pure water at 25°C.

    pH = 7. $[H^+] = 1.0 \times 10^{-7}$ at equilibrium.

59. What is the pH of a $0.02 \text{ M}$ $KOH$ solution?

    pH = 12.3. $[OH^-] = 0.02$, $pOH = 1.7$

60. Calculate the pH of a $0.001 \text{ M}$ $HNO_3$ solution.

    pH = 3. Strong acid: $pH = -\log(0.001) = 3$

61. What is the $K_{sp}$ expression for $Ag_2SO_4$?

    $K_{sp} = [Ag^+]^2[SO_4^{2-}]$. Silver ion squared due to formula stoichiometry.

62. What is the $K_{sp}$ expression for $PbSO_4$?

    $K_{sp} = [Pb^{2+}][SO_4^{2-}]$. 1:1 stoichiometry for ion products.

63. State the formula for calculating pOH from $[OH^-]$.

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

64. Which solution has a higher pH: $0.1 \text{ M}$ $HCl$ or $0.1 \text{ M}$ $NaOH$?

    $0.1 \text{ M}$ $NaOH$. $NaOH$ is basic with higher pH value.

65. State the ion product constant for water ($K_w$) at 25°C.

    $K_w = 1.0 \times 10^{-14}$. Equilibrium constant for water autoionization.

66. Find the pH of a $0.03 \text{ M}$ $LiOH$ solution.

    pH = 12.5. $[OH^-] = 0.03$, $pOH = 1.52$

67. What is the pH of a $0.1 \text{ M}$ $NaOH$ solution?

    pH = 13. Strong base: $[OH^-] = 0.1$, $pOH = 1$

68. Which is more soluble in water: $AgCl$ or $NaCl$?

    $NaCl$. $NaCl$ is highly soluble ionic compound.

69. What is the $K_{sp}$ expression for $Ag_2CrO_4$?

    $K_{sp} = [Ag^+]^2[CrO_4^{2-}]$. Silver ion squared due to formula stoichiometry.

70. Calculate the pH of a $0.1 \text{ M}$ $CH_3COOH$ solution, $K_a = 1.8 \times 10^{-5}$.

    pH ≈ 2.87. Weak acid requires ICE table calculation.

71. State the formula for calculating pH from $[H^+]$.

    $pH = -\text{log}[H^+]$. Take negative log of hydrogen ion concentration.

72. Calculate the pH of a $0.001 \text{ M}$ $HNO_3$ solution.

    pH = 3. Strong acid: $pH = -\log(0.001) = 3$

73. What is the solubility product constant ($K_{sp}$) expression for $AgCl$?

    $K_{sp} = [Ag^+][Cl^-]$. Products of ion concentrations at equilibrium.

74. What is the $K_{sp}$ expression for $CaF_2$?

    $K_{sp} = [Ca^{2+}][F^-]^2$. Fluoride ion squared due to formula stoichiometry.

75. Identify the $K_{sp}$ expression for $BaSO_4$.

    $K_{sp} = [Ba^{2+}][SO_4^{2-}]$. 1:1 stoichiometry for ion products.

76. Calculate the pH of pure water at 25°C.

    pH = 7. $[H^+] = 1.0 \times 10^{-7}$ at equilibrium.

77. Find the pH of a $0.075 \text{ M}$ $Ba(OH)_2$ solution.

    pH = 13.18. Diprotic base: $[OH^-] = 0.15$, $pOH = 0.82$

78. What is the pH of a solution with $[H^+] = 1 \times 10^{-7} \text{ M}$?

    pH = 7. $pH = -\log(10^{-7}) = 7$

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

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

80. What is the pH of a $0.01 \text{ M}$ $HCl$ solution?

    pH = 2. Strong acid completely ionizes: $[H^+] = 0.01$

81. What is the $K_{sp}$ expression for $CuS$?

    $K_{sp} = [Cu^{2+}][S^{2-}]$. 1:1 stoichiometry for ion products.

82. What is the pH of a $0.01 \text{ M}$ $HCl$ solution?

    pH = 2. Strong acid completely ionizes: $[H^+] = 0.01$

83. State the formula for converting pOH to pH.

    $pH = 14 - pOH$. Derived from $pH + pOH = 14$ relationship.

84. What is the $K_{sp}$ expression for $Mg(OH)_2$?

    $K_{sp} = [Mg^{2+}][OH^-]^2$. Hydroxide ion squared due to formula stoichiometry.

85. What is the pH of a solution with $[H^+] = 1 \times 10^{-7} \text{ M}$?

    pH = 7. $pH = -\log(10^{-7}) = 7$

86. Determine the pH of a $0.025 \text{ M}$ $HBr$ solution.

    pH = 1.6. $pH = -\log(0.025) = 1.6$

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

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

88. What is the $K_{sp}$ expression for $Ag_2SO_4$?

    $K_{sp} = [Ag^+]^2[SO_4^{2-}]$. Silver ion squared due to formula stoichiometry.

89. Identify the pH of a neutral solution at 25°C.

    pH = 7. Equal concentrations of $[H^+]$ and $[OH^-]$.

90. If $[OH^-] = 1 \times 10^{-9} \text{ M}$, what is the pH of the solution?

    pH = 5. $pOH = 9$, so $pH = 14 - 9 = 5$

91. Calculate the pH of a $0.1 \text{ M}$ $CH_3COOH$ solution, $K_a = 1.8 \times 10^{-5}$.

    pH ≈ 2.87. Weak acid requires ICE table calculation.

92. What is the $K_{sp}$ expression for $Hg_2I_2$?

    $K_{sp} = [Hg_2^{2+}][I^-]^2$. Iodide ion squared due to formula stoichiometry.

93. What is the $K_{sp}$ expression for $CaF_2$?

    $K_{sp} = [Ca^{2+}][F^-]^2$. Fluoride ion squared due to formula stoichiometry.

94. What is the solubility product constant ($K_{sp}$) expression for $AgCl$?

    $K_{sp} = [Ag^+][Cl^-]$. Products of ion concentrations at equilibrium.

95. State the formula for calculating pH from $[H^+]$.

    $pH = -\text{log}[H^+]$. Take negative log of hydrogen ion concentration.

96. Calculate the pH of pure water at 25°C.

    pH = 7. $[H^+] = 1.0 \times 10^{-7}$ at equilibrium.

97. If $[OH^-] = 1 \times 10^{-9} \text{ M}$, what is the pH of the solution?

    pH = 5. $pOH = 9$, so $pH = 14 - 9 = 5$

98. Find the pH of a $0.075 \text{ M}$ $Ba(OH)_2$ solution.

    pH = 13.18. Diprotic base: $[OH^-] = 0.15$, $pOH = 0.82$

99. Calculate the pH of a $0.001 \text{ M}$ $HNO_3$ solution.

    pH = 3. Strong acid: $pH = -\log(0.001) = 3$

100. If $[OH^-] = 1 \times 10^{-9} \text{ M}$, what is the pH of the solution?

     pH = 5. $pOH = 9$, so $pH = 14 - 9 = 5$

101. What is the $K_{sp}$ expression for $Ag_2SO_4$?

     $K_{sp} = [Ag^+]^2[SO_4^{2-}]$. Silver ion squared due to formula stoichiometry.

102. State the ion product constant for water ($K_w$) at 25°C.

     $K_w = 1.0 \times 10^{-14}$. Equilibrium constant for water autoionization.

103. What is the definition of the solubility product constant ($K_{sp}$)?

     $K_{sp}$ is the equilibrium constant for a solid dissolving in water. Represents dissolution equilibrium of ionic solids.

104. State the $K_{sp}$ expression for $ZnS$.

     $K_{sp} = [Zn^{2+}][S^{2-}]$. 1:1 stoichiometry for ion products.

105. What is the pH of a solution with $[H^+] = 1 \times 10^{-7} \text{ M}$?

     pH = 7. $pH = -\log(10^{-7}) = 7$

106. State the $K_{sp}$ expression for $ZnS$.

     $K_{sp} = [Zn^{2+}][S^{2-}]$. 1:1 stoichiometry for ion products.

107. At 25°C, what is the relationship between $[H^+]$ and $[OH^-]$?

     $[H^+][OH^-] = 1.0 \times 10^{-14}$. Ion product constant for water at 25°C.

108. What is the $K_{sp}$ expression for $PbSO_4$?

     $K_{sp} = [Pb^{2+}][SO_4^{2-}]$. 1:1 stoichiometry for ion products.

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