Chapter 17

What is the pH range of human blood? How is human blood maintained in this pH range?

What is a buffer? How does a buffer work? How does it neutralize added acid? Added base?

What is the common ion effect?

What is the pH of a buffer when the concentrations of both buffer components (the weak acid and its conjugate base) are equal? What happens to the pH when the buffer contains more of the weak acid than the conjugate base? More of the conjugate base than the weak acid?

Suppose that a buffer contains equal amounts of a weak acid and its conjugate base. What happens to the relative amounts of the weak acid and conjugate base when a small amount of strong acid is added to the buffer? What happens when a small amount of strong base is added?

What factors influence the effectiveness of a buffer? What are the characteristics of an effective buffer?

What is the effective pH range of a buffer (relative to the pKa of the weak acid component)?

Describe acid-base titration. What is the equivalence point?

The pH at the equivalence point of the titration of a strong acid with a strong base is 7.0. However, the pH at the equivalence point of the titration of a weak acid with a strong base is above7.0. Explain.

The titration of a diprotic acid with sufficiently different pKa's displays two equivalence points.Why?

What is the difference between the endpoint and the equivalence point in a titration?

What is an indicator? How can an indicator signal the equivalence point of a titration?

What is the solubility-product constant? Write a general expression for the solubility constant of a compound with the general formula AmXn.

What is molar solubility? How do you obtain the molar solubility of a compound from Ksp?

How does a common ion affect the solubility of a compound? More specifically, how is the solubility of a compound with the general formula AX different in a solution containing one of the common ions (A+ or X-) than it is in pure water? Explain.

How is the solubility of an ionic compound with a basic anion affected by pH? Explain.

For a given solution containing an ionic compound, what is the relationship between Q, Ksp, and the relative saturation of the solution?

What is selective precipitation? Under which conditions does selective precipitation occur?

Solve an equilibrium problem (using an ICE table) to calculate the pH of each solution. a. a solution that is 0.20 M in HCHO2 and 0.15 M in NaCHO2 b. a solution that is 0.16 M in NH3 and 0.22 M in NH4Cl

Solve an equilibrium problem (using an ICE table) to calculate the pH of each solution. a. a solution that is 0.195 M in HC2H3O2 and 0.125 M in KC2H3O2 b. a solution that is 0.255 M in CH3NH2 and 0.135 M in CH3NH3Br

Solve an equilibrium problem (using an ICE table) to calculate the pH of each solution. a. 0.15 M HF b. 0.15 M NaF c. a mixture that is 0.15 M in HF and 0.15 M in NaF

Solve an equilibrium problem (using an ICE table) to calculate the pH of each solution. a. 0.18 M CH3NH2 b. 0.18 M CH3NH3Cl c. a mixture that is 0.18 M in CH3NH2 and 0.18 M in CH3NH3Cl

A buffer contains significant amounts of ammonia and ammonium chloride. Write equations that demonstrate how this buffer neutralizes added acid and added base.

Calculate the pH of the solution that results from each mixture. a. 50.0 mL of 0.15 M HCHO2 with 75.0 mL of 0.13 M NaCHO2 b. 125.0 mL of 0.10 M NH3 with 250.0 mL of 0.10 M NH4Cl

Calculate the pH of the solution that results from each mixture. a. 150.0 mL of 0.25 M HF with 225.0 mL of 0.30 M NaF b. 175.0 mL of 0.10 M C2H5NH2 with 275.0 mL of 0.20 M C2H5NH3Cl

A 250.0-mL buffer solution is 0.250 M in acetic acid and 0.250 M in sodium acetate. a. What is the initial pH of this solution? b. What is the pH after addition of 0.0050 mol of HCl? c. What is the pH after addition of 0.0050 mol of NaOH?

For each solution, calculate the initial and final pH after the addition of 0.010 mol of HCl. a. 500.0 mL of pure water b. 500.0 mL of a buffer solution that is 0.125 M in HC2H3O2 and 0.115 M in NaC2H3O2 c. 500.0 mL of a buffer solution that is 0.155 M in C2H5NH2 and 0.145 M in C2H5NH3Cl

For each solution, calculate the initial and final pH after the addition of 0.010 mol of NaOH. a. 250.0 mL of pure water b. 250.0 mL of a buffer solution that is 0.195 M in HCHO2 and 0.275 M in KCHO2 c. 250.0 mL of a buffer solution that is 0.255 M in CH3CH2NH2 and 0.235 M in CH3CH2NH3Cl

Determine whether the mixing of each pair of solutions results in a buffer. a. 75.0 mL of 0.10 M HF; 55.0 mL of 0.15 M NaF b. 150.0 mL of 0.10 M HF; 135.0 mL of 0.175 M HCl c. 165.0 mL of 0.10 M HF; 135.0 mL of 0.050 M KOH d. 125.0 mL of 0.15 M CH3NH2; 120.0 mL of 0.25 M CH3NH3Cl e. 105.0 mL of 0.15 M CH3NH2; 95.0 mL of 0.10 M HCl

Blood is buffered by carbonic acid and the bicarbonate ion.Normal blood plasma is 0.024 M in HCO3- and 0.0012 M H2CO3(pKa1 for H2CO3 at body temperature is 6.1). a. What is the pH of blood plasma? b. If the volume of blood in a normal adult is 5.0 L, what mass of HCl can be neutralized by the buffering system in blood before the pH falls below 7.0 (which would result in death)? c. Given the volume from part b, what mass of NaOH can be neutralized before the pH rises above 7.8?

The fluids within cells are buffered by H2PO4- and HPO42-. a. Calculate the ratio of HPO42- to H2PO4- required to maintain a pH of 7.1 within a cell. b. Could a buffer system employing H3PO4 as the weak acid and H2PO4- as the weak base be used as a buffer system within cells? Explain.

Two 25.0-mL samples, one 0.100 M HCl and the other 0.100 M HF, are titrated with 0.200 M KOH. a. What is the volume of added base at the equivalence point for each titration? b. Is the pH at the equivalence point for each titration acidic, basic, or neutral? c. Which titration curve has the lower initial pH? d. Sketch each titration curve.

Consider the titration of a 35.0-mL sample of 0.175 M HBr with 0.200 M KOH. Determine each quantity. a. the initial pH b. the volume of added base required to reach the equivalence point c. the pH at 10.0 mL of added base d. the pH at the equivalence point e. the pH after adding 5.0 mL of base beyond the equivalence point

A 20.0-mL sample of 0.125 M HNO3 is titrated with 0.150 M NaOH. Calculate the pH for at least five different points on the titration curve and sketch the curve. Indicate the volume at the equivalence point on your graph.

Consider the titration of a 25.0-mL sample of 0.115 M RbOH with 0.100 M HCl. Determine each quantity. a. the initial pH b. the volume of added acid required to reach the equivalence point c. the pH at 5.0 mL of added acid d. the pH at the equivalence point e. the pH after adding 5.0 mL of acid beyond the equivalence point

Consider the titration of a 25.0-mL sample of 0.175 M CH3NH2 with 0.150 M HBr. Determine each quantity. a. the initial pH b. the volume of added acid required to reach the equivalence point c. the pH at 5.0 mL of added acid d. the pH at one-half of the equivalence point e. the pH at the equivalence point f. the pH after adding 5.0 mL of acid beyond the equivalence point

Methyl red has a pKa of 5.0 and is red in its acid form and yellow in its basic form. If several drops of this indicator are placed in a 25.0-mL sample of 0.100 M HCl, what color does the solution appear? If 0.100 M NaOH is slowly added to the HCl sample, in what pH range will the indicator change color?

Phenolphthalein has a pKa of 9.7. It is colorless in its acid form and pink in its basic form. For each of the pH values, calculate [In-]/[HIn] and predict the color of a phenolphthalein solution. a. pH = 2.0 b. pH = 5.0 c. pH = 8.0 d. pH = 11.0

Write balanced equations and expressions for Ksp for the dissolution of each ionic compound. a. BaSO4 b. PbBr2 c. Ag2CrO4

Write balanced equations and expressions for Ksp for the dissolution of each ionic compound. a. CaCO3 b. PbCl2 c. AgI

Use the given molar solubilities in pure water to calculate Ksp for each compound. a. MX; molar solubility = 3.27 * 10-11 M b. PbF2; molar solubility = 5.63 * 10-3 M c. MgF2; molar solubility = 2.65 * 10-4 M

Use the given molar solubilities in pure water to calculate Ksp for each compound. a. BaCrO4; molar solubility = 1.08 * 10-5 M b. Ag2SO3; molar solubility = 1.55 * 10-5 M c. Pd(SCN)2; molar solubility = 2.22 * 10-8 M

Consider the compounds with the generic formulas listed and their corresponding molar solubilities in pure water. Which compound has the smallest value of Ksp? a. AX; molar solubility = 1.35 * 10-4 M b. AX2; molar solubility = 2.25 * 10-4 M c. A2X; molar solubility = 1.75 * 10-4 M

The solubility of copper(I) chloride is 3.91 mg per 100.0 mL of solution. Calculate Ksp for CuCl.

Calculate the molar solubility of MX (Ksp = 1.27 * 10-36) in each liquid or solution. a. pure water b. 0.25 M MCl2 c. 0.20 M Na2X

Predict whether a precipitate forms if you mix 75.0 mL of a NaOH solution with pOH = 2.58 with 125.0 mL of a 0.018 M MgCl2 solution. Identify the precipitate, if any.

Predict whether a precipitate forms if you mix 175.0 mL of a 0.0055 M KCl solution with 145.0 mL of a 0.0015 M AgNO3 solution. Identify the precipitate, if any.

A buffer is created by combining 150.0 mL of 0.25 M HCHO2 with 75.0 mL of 0.20 M NaOH. Determine the pH of the buffer.

In analytical chemistry, bases used for titrations must often be standardized; that is, their concentration must be precisely determined. Standardization of sodium hydroxide solutions can be accomplished by titrating potassium hydrogen phthalate (KHC8H4O4), also known as KHP, with the NaOH solution to be standardized. a. Write an equation for the reaction between NaOH and KHP. b. The titration of 0.5527 g of KHP required 25.87 mL of an NaOH solution to reach the equivalence point. What is the concentration of the NaOH solution?

A 0.5224-g sample of an unknown monoprotic acid was titrated with 0.0998 M NaOH. The equivalence point of the titration occurs at 23.82 mL. Determine the molar mass of the unknown acid.