Химия на английском. 2

6.30.Hydrosulfuric acid, H2S(aq), is a weak diprotic acid that is sometimes used in analytical work. Calculate the pH and [HS−(aq)] of a 7.5 10–3 mol/L solution.

6.31.A 0.10 mol/L solution of a weak monoprotic acid was found to be 5.0% dissociated. Calculate Ka.

6.32.Oxalic acid, HOOCCOOH, is a weak diprotic acid that occurs naturally in some foods, including rhubarb. Calculate the pH of a solution of oxalic acid that is prepared by dissolving 2.5 g in 1.0 L of water. What is the concentration of hydrogen oxalate, HOOCCOO−, in the solution?

6.33.A sample of blood was taken from a patient and sent to a laboratory for testing. Chemists found that the blood pH was 7.40. They also found that the hydrogen

carbonate ion concentration was 2.6 10–2 mol/L. What was the concentration of carbonic acid in the blood?

6.34.An aqueous solution of household ammonia has a molar concentration of 0.105 M. Calculate the pH of the solution.

6.35.Hydrazine, N2H4, has been used as a rocket fuel. The concentration of an aqueous solution of hydrazine is 5.9 10–2 mol/L. Calculate the pH of the solution.

6.36.Morphine, C17H19NO3, is a naturally occurring base that is used to control pain. A 4.5 10–3 mol/L solution has a pH of 9.93. Calculate Kb for morphine.

6.37.Methylamine, CH3NH2, is used to manufacture several prescription drugs. Calculate [OH−] and pOH of a 0.25 mol/L aqueous solution of methylamine.

6.38.At room temperature, trimethylamine, (CH3)3N, is a gas with a strong ammonialike odour. Calculate [OH−] and the percent of trimethylamine molecules that react with water in a 0.22 mol/L aqueous solution.

6.39.An aqueous solution of ammonia has a pH of 10.85. What is the concentration of the solution?

6.40.Use the table of Ka values in Appendix 4 to list the conjugate bases of the following acids in order of increasing base strength: formic acid, HCOOH; hydrofluoric acid, HF(aq); benzoic acid, C6H5COOH; phenol, C6H5OH.

6.41.Compare Kb for ammonia, NH3, and for trimethylamine, (CH3)3N. Which is the stronger acid, NH4+ or (CH3)3NH+?

6.42.A buffer solution is made by mixing 250 mL of 0.200 mol/L aqueous ammonia and 400 mL of 0.150 mol/L ammonium chloride. Calculate the pH of solution.

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6.43.A buffer solution is made by mixing 200 mL of 0.200 mol/L aqueous ammonia and 450 mL of 0.150 mol/L ammonium chloride. Calculate the pH of solution.

6.44.A buffer solution contains 0.200 mol/L nitrous acid, HNO2(aq), and 0.140 mol/L potassium nitrite, KNO2(aq). What is the pH of the buffer solution?

6.45.A buffer solution is prepared by dissolving 1.80 g of benzoic acid, C6H5COOH, and 1.95 g of sodium benzoate, NaC6H5COO, in 800 mL of water. Calculate the pH of the buffer solution.

6.46.Predict whether an aqueous solution of each salt is neutral, acidic, or basic.

6.47. Is the solution of each salt acidic, basic, or neutral? For solutions that are not neutral, write equations that support your predictions.

6.48.Compare Ka for benzoic acid, C6H5COOH, and for phenol, C6H5OH. Which is the stronger base, C6H5COO−(aq) or C6H5O−(aq)? Explain your answer.

6.49.Sodium hydrogen sulfite, NaHSO3, is a preservative that is used to prevent the discolouration of dried fruit. In aqueous solution, the hydrogen sulfite ion can act as

either an acid or a base. Predict whether NaHSO3 dissolves to form an acidic solution or a basic solution. (Refer to Appendix 4 for ionization data.)

6.50.Sodium carbonate and sodium hydrogen carbonate both dissolve to form basic solutions. Comparing solutions with the same concentration, which of these salts forms the more basic solution? Explain.

6.51.Potassium phosphate and potassium dihydrogen phosphate both dissolve to form basic solutions. Comparing solutions with the same concentration, which of these salts forms the more basic solution? Explain.

6.52.Determine whether or not each ion reacts with water. If the ion does react, write the chemical equation for the reaction. Then predict whether the ion forms an acidic

solution or a basic solution.

6.53. A chemist measures the pH of aqueous solutions of Ca(OH)2, CaF2, NH4NO3, KNO3, and HNO3. Each solution has the same concentration. Arrange the solutions from most basic to most acidic.

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6.54. Write the balanced chemical equation that represents the dissociation of each compound in water. Then write the corresponding solubility product expression.

6.55.The maximum solubility of silver cyanide, AgCN, is 1.5 10–8 mol/L at 25°C. Calculate Ksp for silver cyanide.

6.56.A saturated solution of copper(II) phosphate, Cu3(PO4)2 , has a concentration of 6.1 10–7 g Cu3(PO4)2 per 100 mL of solution at 25 °C. What is Ksp for Cu3(PO4)2?

6.57.A saturated solution of CaF2 contains 1.2 1020 molecules of calcium fluoride per liter of solution. Calculate Ksp for CaF2.

6.58.Ksp for silver chloride, AgCl, is 1.8 10–10 at 25 oC.

a)Calculate the molar solubility of AgCl in a saturated solution at 25 °C.

b)How many molecules of AgCl are dissolved in 1.0 L of saturated silver chloride solution?

c)What is the percent (m/v) of AgCl in a saturated solution at 25 °C?

6.59.Calculate the molar solubility of Fe(OH)3 at 25 °C (ref. Appendix 6).

6.60.Calculate the solubility (in mol/L and in g/L) of Zn(IO3)2 in a saturated solution.

6.61.Determine the molar solubility of AgCl:

a) in pure water; b) in 0.15 mol/L NaCl.

6.62.Determine the molar solubility of lead(II) iodide, PbI2, in 0.050 mol/L NaI.

6.63.Calculate the molar solubility of calcium sulfate, CaSO4:

a) in pure water; b) in 0.25 mol/L Na2SO4.

6.64. Calculate the molar solubility of lead(II) chloride, PbCl2: a) in pure water; b) in 0.10 mol/L CaCl2.

6.65.The maximum solubility of barium fluoride, BaF2, at 25 °C, is 1.3 g/L.

a)Calculate Ksp for BaF2 at 25 °C.

b)Calculate the solubility of BaF2 in molecules of BaF2/L.

6.66.A solution of BaCl2 is added to a solution of Na2SO4.

a)Calculate the solubility (in mol/L and in g/L) of BaSO4 in pure water.

b)Calculate the solubility (in mol/L and in g/L) of BaSO4 in 0.085 M Na2SO4.

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6.67.A solution contains 0.15 mol/L of NaCl and 0.0034 mol/L Pb(NO3)2. Does a precipitate form? Include a balanced chemical equation for the formation of the possible precipitate.

6.68.One drop (0.050 mL) of 1.5 mol/L potassium chromate, K2CrO4, is added to 250 mL of 0.10 mol/L AgNO3. Does a precipitate form?

6.69.A chemist adds 0.010 g of CaCl2 to 500 mL of 0.0015 mol/L sodium carbonate, Na2CO3. Does a precipitate of calcium carbonate form?

6.70.0.10 mg of magnesium chloride, MgCl2, is added to 250 mL of 0.0010 mol/L NaOH. Does a precipitate of magnesium hydroxide form?

6.71.100 mL of 1.0 10–3 mol/L Pb(NO3)2 is added to 40 mL of 0.040 mol/L NaCl. Does a precipitate form?

6.72.25 mL of 0.10 mol/L NaOH is added to 500 mL of 0.00010 mol/L cobalt(II) chloride, CoCl2. Does a precipitate form?

6.73.250 mL of 0.0011 mol/L Al2(SO4)3 is added to 50 mL of 0.022 mol/L BaCl2. Does a precipitate form? Include a balanced chemical equation for the formation of the possible precipitate.

6.74.A chemist adds 1.0 mg of NaI to 50 mL of a 0.010 mol/L solution of Pb(NO3)2. Does a precipitate form?

6.75.How many milligrams of Na2SO4 will just begin to precipitate calcium sulfate, CaSO4, from 500 mL of a 0.10 mol/L solution of CaCl2?

6.76.How many drops of 0.0010 mol/L silver nitrate solution will just begin to precipitate AgCl from 5000 mL of a 0.90% (m/v) solution of NaCl? (Assume that one drop equals 0.050 mL.)

6.77.Compare the values of solubility products constants for two salts with the same

begin to add solid Na2SO4.

a)Explain why SrSO4 precipitates first.

b)How many milligrams of Al2(SO4)3 will just begin to precipitate SrSO4 from the solution?

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7.ELECTROCHEMISTRY

7.1.Write a net ionic equation for a reaction in which:

a)Fe2+ acts as an oxidizing agent;

b)Al acts as a reducing agent;

c)Au3+ acts as an oxidizing agent;

d)Cu acts as a reducing agent;

e)Sn2+ acts as an oxidizing agent and as a reducing agent.

7.2.Write the oxidation half-reaction, the reduction half-reaction, and the overall cell reaction for each of the following galvanic cells. Identify the anode and the cathode in

each case. (In part (b), platinum is present as an inert electrode.)

a)Sn(s) | Sn2+(aq) || Tl+(aq) | Tl(s);

b)Cd(s) | Cd2+(aq) || H+(aq) | H2(g) | Pt(s).

7.3.A galvanic cell involves the overall reaction of iodide ions with acidified permanganate ions to form manganese(II) ions and iodine. The salt bridge contains potassium nitrate.

a)Write the half-reactions, and the overall cell reaction.

b)Identify the oxidizing agent and the reducing agent.

c)The inert anode and cathode are both made of graphite. Solid iodine forms on one of them. Which one?

7.4.Look up the standard reduction potentials of the following half-reactions (ref. Appendix 7). Predict whether acidified nitrate ions will oxidize manganese(II) ions to manganese(IV) oxide under standard conditions.

MnO2(s) + 4 H+(aq) + 2e− → Mn2+(aq) + 2 H2O(l);

NO3–(aq) + 4 H+(aq) + 3e− → NO(g) + 2 H2O(l).

7.5. Predict whether each reaction is spontaneous or non-spontaneous under standard conditions.

a) 2 Cr(s) + 3 Cl2(g) → 2 Cr3+(aq) + 6 Cl–(aq).

b)Zn2+(aq) + Fe(s) → Zn(s) + Fe2+(aq).

c)5 Ag(s) + MnO4–(aq) + 8 H+(aq) → 5 Ag+(aq) + Mn2+(aq) + 4 H2O(l).

7.6.Predict whether each reaction is spontaneous or non-spontaneous under standard conditions in an acidic solution.

a)H2O2(aq) → H2(g) + O2(g).

b)3 H2(g) + Cr2O72–(aq) + 8 H+(aq) → 2 Cr3+(aq) + 7 H2O(l).

7.7.Determine the standard cell potential for each of the following redox reactions.

a)CuSO4(aq) + Ni(s) → NiSO4(aq) + Cu(s).

b)Fe(s) + 4 HNO3(aq) → Fe(NO3)3(aq) + NO(g) + 2 H2O(l).

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7.8. Determine if each of the following balanced redox reactions is spontaneous as written, calculate the cell potential.

a) Sn(s) + 2 Cu+(aq) → Sn2+(aq) + 2 Cu(s).

b)Mg(s) + Pb2+(aq) → Pb(s) + Mg2+(aq).

c)2Mn2+(aq) + 8H2O(l) + 10Hg2+(aq) → 2MnO4–(aq) + 16H+(aq) + 5Hg22+(aq).

7.9.Write the two half-reactions for the following redox reactions. Subtract the two reduction potentials to find the standard cell potential for a galvanic cell in which this reaction occurs.

a) Cl2(g) + 2 Br–(aq) → 2 Cl–(aq) + Br2(l).

b)2 Cu+(aq) + 2 H+(aq) + O2(g) → 2 Cu2+(aq) + H2O2(aq).

7.10.Determine the standard cell potential for each of the following redox reactions.

a) 3 Mg(s) + 2 Al3+(aq) → 3 Mg2+(aq) + 2 Al(s). b) 2 K(s) + F2(g) → 2 K+(aq) + 2 F–(aq).

c) Cr2O72–(aq) + 14 H+(aq) + 6 Ag(s) → 2 Cr3+(aq) + 6 Ag+(aq) + 7 H2O(l).

7.11. The cell potential for the following galvanic cell is given: Eocell = 1.750 V. Zn | Zn2+ (1 mol/L) || Pd2+ (1 mol/L) | Pd.

Determine the standard reduction potential for the following half-reaction: Pd2+(aq) + 2e− → Pd(s).

7.12. These equations represent overall cell reactions. Determine the standard potential for each cell and identify the reactions as spontaneous or nonspontaneous as written.

a) 2 Al3+(aq) + 3 Cu(s) → 2 Cu2+(aq) + 2 Al(s). b) Hg2+(aq) + 2 Cu+(aq) → 3 Cu2+(aq) + Hg(l).

c)Cd(s) + 2 NO3–(aq) + 4 H+(aq) → 2 Cd2+(aq) + 2 NO2(g) + 2 H2O(l).

7.13.Write the standard cell notation for the following cells in which the half-cell listed is connected to the standard hydrogen electrode. An example is Na | Na+ || H+ | H2.

Determine the voltage of the cells formed.

7.14.Calculate the cell potential of voltaic cells that contain the following pairs of half-

cells.

a)Chromium in a solution of Cr3+ ions; copper in a solution of Cu2+ ions.

b)Zinc in a solution of Zn2+ ions; platinum in a solution of Pt2+ ions.

c)A half-cell containing both HgCl2 and Hg2Cl2; lead in a solution of Pb2+ ions.

d)Tin in a solution of Sn2+ ions; iodine in a solution of I– ions.

7.15.Predict the products of the electrolysis of a 1 mol/L solution of sodium chloride.

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7.16.The electrolysis of molten calcium chloride produces calcium and chlorine. Write

a)the half-reaction that occurs at the anode;

b)the half-reaction that occurs at the cathode;

c)the chemical equation for the overall cell reaction.

7.17.For the electrolysis of molten lithium bromide, write

a)the half-reaction that occurs at the negative electrode;

b)the half-reaction that occurs at the positive electrode;

c)the net ionic equation for the overall cell reaction.

7.18.Explain why calcium can be produced by the electrolysis of molten calcium chloride, but not by the electrolysis of aqueous calcium chloride.

7.19.One half-cell of a galvanic cell has a nickel electrode in a 1 mol/L nickel(II) chloride solution. The other half-cell has a cadmium electrode in a 1 mol/L cadmium chloride solution.

a)Find the cell potential.

b)Identify the anode and the cathode.

c)Write the oxidation half-reaction, the reduction half-reaction, and the overall cell reaction.

7.20.An external voltage is applied to change the galvanic cell in question 7.21 into an electrolytic cell. Repeat parts (a) to (c) for the electrolytic cell.

7.21.Calculate the mass of zinc plated onto the cathode of an electrolytic cell by a current of 750 mA in 3.25 h.

7.22.How many minutes does it take to plate 0.925 g of silver onto the cathode of an electrolytic cell using a current of 1.55 A?

7.23.The nickel anode in an electrolytic cell decreases in mass by 1.20 g in 35.5 min. The oxidation half-reaction converts nickel atoms to nickel(II) ions. What is the constant current?

7.24.The following two half-reactions take place in an electrolytic cell with an iron anode and a chromium cathode.

Oxidation: Fe(s) → Fe2+(aq) + 2e−.

Reduction: Cr3+(aq) + 3e− → Cr(s).

During the process, the mass of the iron anode decreases by 1.75 g.

a)Find the change in mass of the chromium cathode.

b)Explain why you do not need to know the electric current or the time to complete part (a).

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PERIODIC TABLE

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