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Pure water autoionizes as shown in the equation above. Based on this information, which of the following is correct?
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1
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Pure water autoionizes as shown in the equation above. Based on this information, which of the following is correct?
At 55°C, pH = -SR(log (Kw)) for pure water
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2
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Kw = [H3O+][OH−] = 1.0 × 10−14 at 25°C
Based on the information above, which of the following is true for a sample of pure water at 25°C?
pOH = 7.00
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2 H2O(l) ⇄ H3O+(aq) + OH−(aq)
ΔH°=+56 kJ/molrxn
The endothermic autoionization of pure water is represented by the chemical equation shown above. The pH of pure water is measured to be 7.00 at 25.0°C and 6.02 at 100.0°C. Which of the following statements best explains these observations?
At the higher temperature water dissociates more, [H3O+] = [OH−], and the water remains neutral.
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Which of the following gives the best estimate for the pH of a 5×10−4 M Sr(OH)2(aq) solution at 25°C?
pH ≈ 11.0 because Sr(OH)2 is a strong base.
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5
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Which of the following gives the best estimate for the pH of a 1×10−5 M HClO4(aq) solution at 25°C?
pH ≈ 5.0 because HClO4 is a strong acid.
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Which of the following is the correct mathematical relationship to use to calculate the pH of a 0.10 M aqueous HBr solution?
pH =−log(1.0×10−1)
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The equilibrium for the acid ionization of HCOOH is represented by the equation above and the table gives the percent ionization for HCOOH at different initial concentrations of the weak acid at 25°C. Based on the information, which of the following is true for a 0.125 M aqueous solution of HCOOH?
It has a lower percent ionization, a larger [H3O+]eq, and a higher pOH than a 0.100 M HCOOH solution does.
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NH3(aq) + H2O(l) ⇄ NH4+(aq) + OH−(aq) Kb = 1.8 × 10−5 at 25°C
The reaction between NH3 and water is represented above. A solution that is initially 0.200 M in NH3 has a pH of 11.28. Which of the following correctly predicts the pH of a solution for which [NH3]initial = 0.100 M, and why?
The pH will be lower than 11.28 because the equilibrium concentration of OH− ions decreases when the initial concentration of the base decreases.
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HClO(aq) + H2O(l) ⇄ H3O+(aq) + ClO−(aq)
Ka = 3.98 × 10−8
The acid ionization equilibrium for HClO at 25°C is shown above. A 0.100 M aqueous solution of this acid has a pH of about 4.2. If a solution with an initial concentration of HClO of 0.200 M is allowed to reach equilibrium at the same temperature, which of the following correctly predicts its pH, and why?
The pH will be lower than 4.2 because increasing the concentration of the weak acid produces more H3O+ to establish equilibrium
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The weak acid CH3COOH has a pKa of 4.76. A solution is prepared by mixing 500.mL of 0.150 M CH3COOH(aq) and 0.0200mol of NaOH(s). Which of the following can be used to calculate the pH of the solution?
pH = 4.76 + log (0.02000.0550) = 4.32
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Which of the following provides the correct mathematical expression to calculate the pH of a solution made by mixing 10.0ml of 1.00MHCl and 11.0mL of 1.00MNaOH at 25°C?
pH = 14.00 + log (0.00100.0210) = 12.68
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A buffer solution is formed by mixing equal volumes of 0.12MNH3(aq) and 0.10MHCl(aq), which reduces the concentration of both solutions by one half. Based on the pKa data given in the table, which of the following gives the pH of the buffer solution?
pH = 9.25+log(0.0100.050) = 8.55
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The pH versus volume data for the titration of 0.10 M HNO2(aq) with 0.10 M KOH(aq) is plotted on the graph above. Based on the data, which of the following species is present in the greatest concentration after 6.0 mL of KOH(aq) has been added to the solution of HNO2(aq)?
NO2−(aq)
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A 0.20 M solution of the weak acid potassium hydrogen phthalate (KHP) is titrated with 0.10M NaOH(aq). Based on the titration curve shown in the graph above, the pKa of KHP is closest to which of the following?
5.4
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A 60.mL sample of NaOH(aq) was titrated with 0.10 M HCl(aq). Based on the resulting titration curve shown above, what was the approximate concentration of NaOH in the sample?
0.050M
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The table above provides information on two weak acids. Which of the following explains the difference in their acid strength?
Acid 2 is a stronger acid because it has a more stable conjugate base than acid 1 due to the greater number of electronegative Br atoms.
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The equilibrium reactions for diprotic oxoacids with a general formula H2XO4 are represented by the equations above. The acid ionization constants for H2SeO4 and H2TeO4 are provided in the table. Which of the following best explains the difference in strength for these two acids?
H2TeO4 is weaker because Te is less electronegative than Se, resulting in less stable conjugate bases HTeO4− and TeO42− than those for H2SeO4.
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Lewis diagrams of the weak bases NH3 and NF3 are shown above. Based on these diagrams, which of the following predictions of their relative base strength is correct, and why?
NF3 is a weaker base than NH3 because of the greater electronegativity of F compared with H.
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A student measures the pH of a 0.0100 M buffer solution made with HClO and NaClO, as shown above. The pKa of HClO is 7.40 at 25°C. Based on this information, which of the following best compares the relative concentrations of ClO− and HClO in the buffer solution?
[ClO−] > [HClO]
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HCNO is a weak acid with a pKa value of 3.5. The graph above shows the results of a titration of an aqueous solution of HCNO with 0.100 M NaOH. Based on the results, the concentration of CNO− is greater than the concentration of HCNO at which of the following pH values?
4.0
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C2H3COOH(aq) + H2O(l) ⇄ H3O+(aq) + C2H3COO−(aq)
pKa=4.25
The weak acid ionization equilibrium for C2H3COOH is represented by the equation above. A student measures the pH of C2H3COOH(aq) using a probe and a pH meter in the experimental setup shown. Based on the information given, which of the following is true?
[C2H3COOH] > [C2H3COO−] since the pKa of the weak acid is greater than the pH of the solution.
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CH3CH2COOH(aq) + H2O(l) ⇄ H3O+(aq) + CH3CH2COO−(aq) Ka = 1.4×10−5 at 25°C
CH3CH2COO−(aq) + H2O(l) ⇄ CH3CH2COOH(aq) + OH−(aq) Kb = 7.4×10−10 at 25°C
The acid equilibrium for CH3CH2COOH and the base equilibrium for CH3CH2COO− are represented above. One liter of a buffer solution with pH = 4.85 is made by mixing 0.100 M CH3CH2COOH and 0.100 M NaCH3CH2COO. If 10.0 mL of 0.500 M NaOH is added to the buffer, which of the following is most likely the resulting pH, and why?
The pH will be slightly greater than 4.85, because some CH3CH2COOH will react with the added bases, resulting in a slight decrease in [H3O+].
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A buffer solution that is 0.100 M in both HCOOH and HCOOK has a pH = 3.75. A student says that if a very small amount of 0.100 M HCl is added to the buffer, the pH will decrease by a very small amount. Which of the following best supports the student’s claim?
HCOO− will accept a proton from HCl to produce more HCOOH and H2O.
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CH3COOH(aq) + H2O(l) ⇄ H3O+(aq) + CH3COO−(aq) pKa = 4.76 at 25°C
The equilibrium representing the acid dissociation of CH3COOH is shown above. A buffer solution is prepared by adding 0.10 mol of NaOH(s) to 1.00 L of 0.30 M CH3COOH. Assuming the change in volume is negligible, which of the following expressions will give the pH of the resulting buffer at 25°C?
pH = 4.76 + log(0.100/0.20)
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To prepare a buffer solution for an experiment, a student measured out 53.49 g of NH4Cl(s) (molar mass 53.49 g/mol) and added it to 1.0 L of 1.0 M NH3(aq). However, in the process of adding the NH4Cl(s) to the NH3(aq), the student spilled some of the NH4Cl(s) onto the bench top. As a result, only about 50. g of NH4Cl(s) was actually added to the 1.0 M NH3(aq). Which of the following best describes how the buffer capacity of the solution is affected as a result of the spill?
The solution has a greater buffer capacity for the addition of acid than for base, because [NH3] > [NH4+].
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HC3H5O2(aq) + H2O(l) ⇄ H3O+(aq) + C3H5O2−(aq)
pKa = 4.87
The acid ionization equilibrium for HC3H5O2 is represented by the equation above. A mixture of 1.00L of 0.100MHC3H5O2 and 0.500L of 0.100MNaOH will produce a buffer solution with a pH=4.87. If the NaOH solution was mislabeled and was 1.00M instead of 0.100M, which of the following would be true?
The pH of the resulting solution would be much higher than 4.87 because the weak acid would be completely neutralized by the larger amount of NaOH added.
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The acid ionization equilibrium for HNO2 is represented by the equation above. A 250.0 mL buffer solution is prepared by mixing 125.0 mL of 0.20 M HNO2 and 125.0 mL of 0.1 M NaOH. To test the buffer capacity, the pH is measured and recorded in the table for four samples of the buffer and one sample of a mixture of the buffer and HCl. Which of the following best helps explain why the pH of sample 4 is lower than the pH of the other samples containing only buffer solution?
After measuring the pH of the more acidic sample 3, the pH probe was not rinsed and wiped, resulting in the neutralization of a very small amount of the conjugate base in sample 4.
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Which of the following chemical equilibrium equations best shows what happens in the buffer solutions to minimize the change in pH when a small amount of a strong base is added?
HCO3−(aq) + OH−(aq) ⇄ CO32−(aq) + H2O(l)
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Which of the following mathematical expressions can be used to determine the approximate pH of buffer 1 ?
pH = [14.00+log(2.1×10−4)] + log (0.1000/0.150) = 10.15
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Which mathematical expression can be used to explain why buffer 2 and buffer 3 have the same pH?
log(0.2000/0.200) = log(0.1000/0.100) = log(1)
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