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chemistry a molecular approach

Chemistry A Molecular Approach 5th Edition Nivaldo Tro - Solutions

Does a large positive electrode potential indicate a strong oxidizing agent or a strong reducing agent? What about a large negative electrode potential?
How can Table 20.1 be used to predict whether or not a metal will dissolve in HCl? In HNO3? TABLE 20.1 Standard Electrode Potentials at 25 °C Reduction Half-Reaction High tendency towards reduction F₂(g) + 2 e H₂O₂(aq) + 2 H+(aq) + 2 e PbO₂ (s) + 4 H*(aq) + SO² (aq) + 2 e MnO4 (aq) + 4 H
Which metal can be used as a sacrificial electrode to prevent the rusting of an iron pipe?a) Au b) Ag c) Cu d) Mn
Explain why E°cell, ΔG°rxn, and K are all interrelated.
Does a redox reaction with a small equilibrium constant (K < 1) have a positive or a negative E°cell? Does it have a positive or a negative ΔG°rxn?
How does Ecell depend on the concentrations of the reactants and products in the redox reaction occurring in the cell? What effect does increasing the concentration of a reactant have on Ecell? Increasing the concentration of a product?
Use the Nernst equation to show that Ecell = E°cell under standard conditions.
What are the anode and cathode reactions in a common drycell battery? In an alkaline battery?
What are the anode and cathode reactions in a lead–acid storage battery? What happens when the battery is recharged?
What are the three common types of portable rechargeable batteries, and how does each one work?
What is a fuel cell? What is the most common type of fuel cell, and what reactions occur at its anode and cathode?
Explain how a fuel-cell breathalyzer works.
List some applications of electrolysis.
The anode of an electrolytic cell must be connected to which terminal—positive or negative—of the power source?
What species is oxidized, and what species is reduced in the electrolysis of a pure molten salt?
If an electrolytic cell contains a mixture of species that can be oxidized, how do you determine which species will actually be oxidized? If it contains a mixture of species that can be reduced, how do you determine which one will actually be reduced?
Why does the electrolysis of an aqueous sodium chloride solution produce hydrogen gas at the cathode?
What is overvoltage in an electrochemical cell? Why is it important?
How is the amount of current flowing through an electrolytic cell related to the amount of product produced in the redox reaction?
What is corrosion? Why is corrosion only a problem for some metals (such as iron)?
Explain the role of each of the following in promoting corrosion: moisture, electrolytes, and acids.
Balance each redox reaction occurring in acidic aqueous solution. a. K(s) + Cr³+ (aq). Cr(s) + K+ (aq) b. Al(s) + Fe²+ (aq) Al³+ (aq) + Fe(s) c. BrO3(aq) + N₂H₂(g) →→→ Br¯(aq) + N₂(g)
How can the corrosion of iron be prevented?
Balance each redox reaction occurring in acidic aqueous solution. a. Zn(s) + Sn²+ (aq) b. Mg(s) + Cr³+ (aq) — Zn²+ (aq) + Sn(s) Mg²+ (aq) + Cr(s) c. MnO4 (aq) + Al(s) →→→ Mn²+ (aq) + A1³+ (aq)
Balance each redox reaction occurring in acidic aqueous solution. a. PbO₂(s) + I (aq) →→→ Pb²+ (aq) + 1₂(s) 2+ b. SO32 (aq) + MnO4 (aq) → SO42² (aq) + Mn²+ (aq) 2- c. S₂03² (aq) + Cl₂(g) → SO4² (aq) + Cl¯ (aq)
Balance each redox reaction occurring in acidic aqueous solution. a. I (aq) + NO₂ (aq) b. CIO4 (aq) + Cl¯(aq) c. NO3(aq) + Sn²+ (aq) 1₂(s) + NO(g) CIO3(aq) + Cl₂(g) Sn+ (aq) + NO(g)
Balance each redox reaction occurring in basic aqueous solution. a. H₂O₂(aq) + ClO₂(aq) CIO₂ (aq) + O₂(g) b. Al(s) + MnO4 (aq) → MnO₂(s) + Al(OH)4 (aq) c. Cl₂(g) →→→ Cl(aq) + CIO (aq)
Balance each redox reaction occurring in basic aqueous solution. a. MnO4 (aq) + Br (aq) → MnO₂(s) + BrO3(aq) b. Ag(s) + CN- (aq) + O₂(g) → Ag(CN)₂ (aq) c. NO₂ (aq) + Al(s) - NH3(g) + AlO₂ (aq)
Sketch a voltaic cell for each redox reaction. Label the anode and cathode and indicate the half-reaction that occurs at each electrode and the species present in each solution. Also indicate the direction of electron flow. a. 2 Ag+ (aq) + Pb(s) → 2 Ag(s) + Pb²+ (aq) b. 2 C1O₂(g) + 21 (aq) 2
Sketch a voltaic cell for each redox reaction. Label the anode and cathode and indicate the half-reaction that occurs at each electrode and the species present in each solution. Also indicate the direction of electron flow. 2+ Ni(s) + Mg²+ (aq) H₂(g) + Fe²+ (aq) a. Ni²+ (aq) + Mg(s) b. 2 H+
Calculate the standard cell potential for each of the electrochemical cells in Problem 43.Problem 43Sketch a voltaic cell for each redox reaction. Label the anode and cathode and indicate the half-reaction that occurs at each electrode and the species present in each solution. Also indicate the
Calculate the standard cell potential for each of the electrochemical cells in Problem 44.Problem 44Sketch a voltaic cell for each redox reaction. Label the anode and cathode and indicate the half-reaction that occurs at each electrode and the species present in each solution. Also indicate the
Consider the voltaic cell: a. Determine the direction of electron flow and label the anode and the cathode.b. Write a balanced equation for the overall reaction and calculate E°cell.c. Label each electrode as negative or positive.d. Indicate the direction of anion and cation flow in the salt
Consider the voltaic cell:a. Determine the direction of electron flow and label the anode and the cathode.b. Write a balanced equation for the overall reaction and calculate E°cell.c. Label each electrode as negative or positive.d. Indicate the direction of anion and cation flow in the salt
Use line notation to represent each electrochemical cell in Problem 43.Problem 43Sketch a voltaic cell for each redox reaction. Label the anode and cathode and indicate the half-reaction that occurs at each electrode and the species present in each solution. Also indicate the direction of electron
Use line notation to represent each electrochemical cell in Problem 44.Problem 44Sketch a voltaic cell for each redox reaction. Label the anode and cathode and indicate the half-reaction that occurs at each electrode and the species present in each solution. Also indicate the direction of electron
Make a sketch of the voltaic cell represented by the line notation.Write the overall balanced equation for the reaction and calculate E°cell. Sn(s) Sn²+ (aq) || NO(g) |NO3(aq), H+ (aq) | Pt(s)
Make a sketch of the voltaic cell represented by the line notation.Write the overall balanced equation for the reaction and calculate E°cell. Mn(s) | Mn²+ (aq)||CIO₂ (aq) | CIO2(g) | Pt(s)
Which metal could you use to reduce Mn2+ ions but not Mg2+ ions?
Which metal can be oxidized with an Sn2+ solution but not with an Fe2+ solution?
Determine whether or not each metal dissolves in 1 M HCl. For those metals that do dissolve, write a balanced redox reaction showing what happens when the metal dissolves.a. Al b. Ag c. Pb
Determine whether or not each metal dissolves in 1 M HCl. For those metals that do dissolve, write a balanced redox reaction showing what happens when the metal dissolves.a. Cu b. Fe c. Au
Determine whether or not each metal dissolves in 1 M HNO3.For those metals that do dissolve, write a balanced redox reaction showing what happens when the metal dissolves.a. Cub. Au
Calculate E°cell for each balanced redox reaction and determine if the reaction is spontaneous as written. 2+ a. 2 Cu(s) + Mn²+ (aq) b. MnO₂(aq) + 4H+ (aq) + Zn(s) c. Cl₂(g) + 2 F (aq) 2 Cut (aq) + Mn(s) 2+ Mn²+ (aq) + 2 H₂O(1) + Zn²+ (aq) F₂(g) + 2 CI (aq)
Determine whether or not each metal dissolves in 1 M HIO3. For those metals that do dissolve, write a balanced redox equation for the reaction that occurs.a. Au b. Cr
Calculate E°cell for each balanced redox reaction and determine if the reaction is spontaneous as written. a. O₂(g) + 2 H₂O(1) + 4 Ag(s) b. Br₂(1) 21 (aq) c. PbO₂ (s) + 4H+ (aq) + Sn(s) 4 OH(aq) + 4 Ag+ (aq) 2 Br (aq) + 1₂(s) 2+ Pb²+ (aq) + 2 H₂O(1) + Sn²+ (aq)
Which metal cation is the best oxidizing agent? a. Pb²+ b. Cr³+ c. Fe²+ d. Sn²+
Which metal is the best reducing agent? a. Mn b. Al c. Ni d. Cr
Use tabulated electrode potentials to calculate ΔG°rxn for each reaction at 25 °C. a. Pb²+ (aq) + Mg(s) b. Br₂() + 2 CI (aq) c. MnO₂(s) + 4H+ (aq) + Cu(s) Pb(s) + Mg2+ (aq) 2 Br (aq) + Cl₂(8) 2+ Mn²+ (aq) + 2 H₂O(l) + Cu²+ (aq)
Use tabulated electrode potentials to calculate ΔG°rxn for each reaction at 25 °C. 2+ a. 2 Fe³+ (aq) + 3 Sn(s) 2 Fe(s) + 3 Sn²+ (aq) b. O₂(g) + 2 H₂O(l) + 2 Cu(s) →→→ 4 OH(aq) + 2 Cu²+ (aq) c. Br₂(1) + 21 (aq) 2 Br (aq) + 1₂(s)
Calculate the equilibrium constant for each of the reactions in Problem 65.Problem 65Use tabulated electrode potentials to calculate ΔG°rxn for each reaction at 25 °C. a. Pb²+ (aq) + Mg(s) b. Br₂() + 2 CI (aq) c. MnO₂(s) + 4H+ (aq) + Cu(s) Pb(s) + Mg2+ (aq) 2 Br (aq) + Cl₂(8) 2+ Mn²+
Calculate the equilibrium constant for each of the reactions in Problem 66.Problem 66Use tabulated electrode potentials to calculate ΔG°rxn for each reaction at 25 °C. 2+ a. 2 Fe³+ (aq) + 3 Sn(s) 2 Fe(s) + 3 Sn²+ (aq) b. O₂(g) + 2 H₂O(l) + 2 Cu(s) →→→ 4 OH(aq) + 2 Cu²+ (aq) c.
Calculate the equilibrium constant for the reaction between Ni2+(aq) and Cd(s) (at 25 °C)
Calculate the equilibrium constant for the reaction between Fe2+(aq) and Zn(s) (at 25 °C).
Calculate ΔG∘rxn and E°cell for a redox reaction with n = 2 that has an equilibrium constant of K = 25 (at 25 °C).
Calculate ΔG°rxn and E°cell for a redox reaction with n = 3 that has an equilibrium constant of K = 0.050 (at 25 °C).
A voltaic cell employs the redox reaction:Calculate the cell potential at 25 °C under each set of conditions. 2 Fe³+ (aq) + 3 Mg(s) - 2 Fe(s) + 3 Mg²+ (aq)
A voltaic cell employs the following redox reaction:Calculate the cell potential at 25 °C under each set of conditions. Sn²+ (aq) + Mn(s) 2+ Sn(s) + Mn²+ (aq)
An electrochemical cell is based on these two half-reactions:Calculate the cell potential at 25 °C. Ox: Pb(s) Pb²+ (aq, 0.10 M) + 2 e Red: MnO4 (aq, 1.50 M) + 4H* (aq, 2.0 M) + 3 e MnO₂ (s) + 2 H₂O(1)
Why do exothermic processes tend to be spontaneous at low temperatures? Why does their tendency toward spontaneity decrease with increasing temperature?
Under which set of conditions is ΔGrxn for the reaction A(g) → B(g) most likely to be negative?a) PA = 10.0 atm; PB = 10.0 atmb) PA = 10.0 atm; PB = 0.010 atmc) PA = 0.010 atm; PB = 10.0 atmd) PA = 0.010 atm; PB = 0.010 atm
What is the significance of the change in Gibbs free energy (ΔG) for a reaction?
Which statement is true for the freezing of liquid water below 0°C?a) ΔH is positive; ΔS is negative; ΔG is negativeb) ΔH is negative; ΔS is negative; ΔG is negativec) ΔH is positive; ΔS is positive; ΔG is positived) ΔH is positive; ΔS is negative; ΔG is positive
Predict the spontaneity of a reaction (and the temperature dependence of the spontaneity) for each possible combination of signs for ΔH and ΔS (for the system).a. ΔH negative, ΔS positive b. ΔH positive, ΔS negativec. ΔH negative, ΔS negative d. ΔH positive, ΔS positi
State the third law of thermodynamics and explain its significance.
Why is the standard entropy of a substance in the gas state greater than its standard entropy in the liquid state?
How does the standard entropy of a substance depend on its molar mass? On its molecular complexity?
How can you calculate the standard entropy change for a reaction from tables of standard entropies?
What are three different methods to calculate ΔG° for a reaction?Which method would you choose to calculate ΔG° for a reaction at a temperature other than 25 °C?
Why is free energy “free”?
Explain the difference between ΔG° and ΔG.
Why does water spilled on the floor evaporate even though ΔG° for the evaporation process is positive at room temperature?
How do you calculate the change in free energy for a reaction under nonstandard conditions?
How does the value of ΔG° for a reaction relate to the equilibrium constant for the reaction? What does a negative ΔG° for a reaction imply about K for the reaction? A positive ΔG°?
Which of these processes is spontaneous?a. The combustion of natural gasb. The extraction of iron metal from iron orec. A hot drink cooling to room temperatured. Drawing heat energy from the ocean’s surface to power a ship
Two systems, each composed of two particles represented by circles, have 20 J of total energy. Which system, A or B, has the greater entropy? Why? System A 10 J System B 12 J 8 J
Which of these processes are nonspontaneous? Are the nonspontaneous processes impossible?a. A bike going up a hillb. A meteor falling to Earthc. Obtaining hydrogen gas from liquid waterd. A ball rolling down a hill
Two systems, each composed of three particles represented by circles, have 30 J of total energy. In how many energetically equivalent ways can you distribute the particles in each system?Which system has greater entropy? System A 10 J System B 12 J 10 J 8 J
Calculate the change in entropy that occurs in the system when 1.00 mole of isopropyl alcohol (C3H8O) melts at its melting point (-89.5 °C). See Table 12.9 for heats of fusion. TABLE 12.9 Heats of Fusion of Several Substances Liquid Chemical Formula Melting Point (°C) Water 0.00 Isopropyl
Calculate the change in entropy that occurs in the system when 1.00 mole of diethyl ether (C4H10O) condenses from a gas to a liquid at its normal boiling point (34.6 °C). See Table 12.7 for heats of vaporization. TABLE 12.7 Heats of Vaporization of Several Liquids at Their Boiling Points and at 25
Calculate the change in entropy that occurs in the system when 45.0 g of acetone (C3H6O) freezes at its melting point (-94.8 °C).See Table 12.9 for heats of fusion. TABLE 12.9 Heats of Fusion of Several Substances Liquid Chemical Formula Melting Point (°C) Water 0.00 Isopropyl
Calculate the change in entropy that occurs in the system when 55.0 g of water vaporizes from a liquid to a gas at its boiling point (100.0 °C). See Table 12.7 for heats of vaporization. TABLE 12.7 Heats of Vaporization of Several Liquids at Their Boiling Points and at 25 °C A Hvap (kJ/mol) at
Without doing any calculations, determine the sign of ΔSsys for each chemical reaction. a. Mg(s) + Cl₂(g) →→→ MgCl₂(s) b. 2 H₂S(g) + 3 O₂(g) → 2 H₂O(g) + 2 SO₂(g) c. 203(g) d. HCI(g) + NH3(g) NH4Cl(s) 3 O₂(g)
Without doing any calculations, determine the signs of ΔSsys and ΔSsurr for each chemical reaction. In addition, predict under what temperatures (all temperatures, low temperatures, or high temperatures), if any, the reaction is spontaneous. a. 2 CO(g) + O(g) 2 CO(g) b. 2 NO(g) 2NO(g) + O(g) AH
Without doing any calculations, determine the sign of ΔSsys for each chemical reaction.a. 2 KClO3(s) → 2 KCl(s) + 3 O2( g)b. CH2 = CH2( g) + H2( g) → CH3CH3(g)c. Na(s) + 1/2 Cl2(g) → NaCl(s)d. N2(g) + 3 H2(g) → 2 NH3(g)
Calculate ΔSsurr at the indicated temperature for each reaction. a. ΔΗxn = -385 kJ; 298 K c. ΔHxn +114 kJ; 298 K b. ΔΗan d. ΔΗ, xn −385 kJ; 77 K +114 kJ; 77 K
Given the values of ΔH °rxn, ΔS °rxn, and T, determine ΔSuniv and predict whether or not each reaction is spontaneous. a. AH n = +115 kJ; ASxn=-263 J/K; T = 298 K b. AH n = -115 kJ; ASn = +263 J/K; T = 298 K c. AH n = -115 kJ; ASn=-263 J/K; T = 298 K d. AH = -115 kJ; ASn=-263 J/K; T = 615
A reaction has ΔH°rxn = -112 kJ and ΔS °rxn = 354 J/K. At what temperature is the change in entropy for the reaction equal to the change in entropy for the surroundings?
Given the values of ΔHrxn, and T, determine ΔSrxn, and predict whether or not each reaction is spontaneous. a. AH n = -95 kJ; ASxn-157J/K; T = 298 K b. AH = -95 kJ; ASn = -157 J/K; T = 855 K c. AHxn = -157 J/K; T = 298 K +95 kJ; ASn = -95 kJ; ASixn -95 kJ; ASn = d. AH = +157J/K; T = 398 K
Calculate the change in Gibbs free energy for each of the sets of ΔHrxn, ΔSrxn, and T given in Problem 41. Predict whether or not each reaction is spontaneous at the temperature indicated.Problem 41Given the values of ΔH°rxn, ΔS°rxn, and T, determine ΔSuniv and predict whether or not each
Calculate the change in Gibbs free energy for each of the sets of ΔHrxn, ΔSrxn, and T given in Problem 42. Predict whether or not each reaction is spontaneous at the temperature indicated.Problem 42Given the values of ΔHrxn, and T, determine ΔSrxn, and predict whether or not each reaction is
Calculate the free energy change for this reaction at 25 °C. Is the reaction spontaneous? C3H8(g) + 5O₂(g) 5 O₂(g) → 3 CO₂(g) + 4H₂O(g) AHin = -2217 kJ; ASixn 101.1 J/K rxn
Calculate the free energy change for this reaction at 25 °C. Is the reaction spontaneous? 2 Ca(s) + O2(g) — 2 CaO(s) −1269.8 kJ; Δ.Sixn -364.6J/K ΔΗ, xn
Fill in the blanks in the table. Both ΔH and ΔS refer to the system. ΔΗ T AS + AG Temperature dependent Low Temperature Spontaneous High Temperature Spontaneous Nonspontaneous Nonspontaneous
Predict the conditions (high temperature, low temperature, all temperatures, or no temperatures) under which each reaction is spontaneous.a. H2O( g) → H2O(l)b. CO2(s) → CO2(g)c. H2(g) → 2 H(g)d. 2 NO2( g) → 2 NO(g) + O2(g) (endothermic)
How does the molar entropy of a substance change with increasing temperature?
For each pair of substances, choose the one that you expect to have the higher standard molar entropy (S°) at 25 °C. Explain your choices. a. CO(g); CO₂(g) c. Ar(g); CO₂(g) e. NO₂(g); CH3CH₂CH3(g) b. CH₂OH(1); CH3OH(g) d. CH₂(g); SiH4(g) f. NaBr(s); NaBr(aq)
For each pair of substances, choose the one that you expect to have the higher standard molar entropy (S°) at 25 °C. Explain your choices. a. NaNO3(s); NaNO3(aq) c. Br₂(1); Br₂(g) e. PC13(g); PC15(g) b. CH4(g); CH3CH3(g) d. Br₂(g); F₂(g) f. CH3CH₂CH₂CH3(g); SO₂(g)
What is the molar entropy of a pure crystal at 0 K? What is the significance of the answer to this question?
Rank each set of substances in order of increasing standard molar entropy (S°). Explain your reasoning. a. NH3(g); Ne(g); SO₂(g); CH3CH₂OH(g); He(g) b. H₂O(s); H₂O(1); H₂O(g) C. CH4(g); CF4(g); CC14(g)

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