Chem 1220: Thermodynamics, Electrochemistry, and Nuclear Chemistry
Calculate the standard cell potential for each reaction below, and note whether the reaction is spontaneous under standard state conditions. (a) Mg(s) + Ni2+ (aq) → Mg2+ (aq) + Ni(s) (b) 2Ag+ (aq) + Cu(s) → Cu2+ (aq) + 2Ag(s) (c) Mn(s) + Sn(NO3 )2 (aq) → Mn(NO3 )2 (aq) + Sn(s) (d) 3Fe(NO3 )2 (aq) + Au(NO3 )3 (aq) → 3Fe(NO3 )3 (aq) + Au(s)
(a) +2.115V (spontaneous) (b) +0.4626V (spontaneous) (c) +1.0589V (spontaneous) (d) +0.727V (spontaneous)
Determine whether each of the following reactions is spontaneous under standard conditions. (a) Zn2+ (aq) + Ni (s) --> Zn (s) + Ni2+ (aq) (b) Cl− (aq) + Cu2+ (aq) --> Cu (s) + Cl2 (g) (c) Fe2+ (aq) + Ag+ (aq) --> Fe3+ (aq) + Ag (s)
(a) -0.53; non spontaneous (b) -1.02; non spontaneous (c) 0.35; spontaneous
Using the relevant S° values listed in Appendix G, calculate ∆S° for the following changes: (a) N2(g) + 3H2(g) → 2NH3(g) (b) N2(g) + 5O2(g) → N2O5(g)
(a) -198.1 J/k (b) -348.9 J/K
For each pair of standard cell potential and electron stoichiometry values below, calculate a corresponding standard free energy change (kJ).(a) 0.000 V, n = 2(b) +0.434 V, n = 2 Chemistry 2e 17: Electrochemistry17.4: Potential, Free Energy, and Equilibrium (c) -2.439 V, n = 1
(a) 0 j/mol or 0 KJ/mol (b) -83749 J/mol or -83.7 kJ/mol (c) +235327 J/mol or +235.3 kJ/mol
Calculate ∆S° for the following changes. (a) SnCl4 (l) → SnCl4 (g) (b) CS2 (g) → CS2 (l) (c) Cu(s) → Cu(g) (d) H2O(l) → H2O(g) (e) 2H2 (g) + O2 (g) → 2H2O(l) (f) 2HCl(g) + Pb(s) → PbCl2 (s) + H2 (g) Zn(s) + CuSO4 (s) → Cu(s) + ZnSO4 (s)
(a) 107 J/K (b) -86.4 J/k (c) 133.2 J/k (d) 118.8 J/K (e) -326.6 J/K (f) -171.9 J/K (g) -7.2 J/K
Determine ΔG° for the following reactions. (a) Antimony pentachloride decomposes at 448 °C. The reaction is: SbCl5(g) → SbCl3(g) + Cl2(g) An equilibrium mixture in a 5.00 L flask at 448 °C contains 3.85 g of SbCl5, 9.14 g of SbCl3, and 2.84 g of Cl2. (b) Chlorine molecules dissociate according to this reaction: Cl2(g) → 2Cl(g)1.00% of Cl2 molecules dissociate at 975 K and a pressure of 1.00 atm.
(a) 22.1 (b) 61.5
Determine the oxidation states of the elements in the compounds listed. None of the oxygen- containing compounds are peroxides or superoxides. (a) H3PO4 (b) Al(OH)3 (c) SeO2 (d) KNO2 (e) In2S3 (f) P4O6
(a) H +1, P +5, O -2; (b) Al +3, H +1, O -2; (c) Se +4, O -2; (d) K +1, N +3, O -2; (e) In +3, S -2; (f) P +3, O -2
Calculate the equilibrium constant at the temperature given. (a) O (g) + 2F (g) → 2F O(g) (T = 100 °C) 222 (b) I2 (s) + Br2 (l) → 2IBr(g) (T = 0.0 °C) (c) 2LiOH(s) + CO2 (g) → Li2CO3 (s) + H2O(g) (T = 575 °C) (d)N2O3(g)→NO(g)+NO2(g) (T= -10.0°C) (e) SnCl4 (l) → SnCl4 (g) (T = 200 C) °
(a) K= 1.07x10^-13 (b) K= 2.42x10^-3 (c) K= 2.72x10^4 (d) K= 0.229 (e) K= 16.1
Calculate the equilibrium constant at 25°C for each of the following reactions from the value of ΔG° given. (a) O (g) + 2F (g) → 2OF (g) ∆G° = - 9.2 kJ 222 (b) I2 (s) + Br2 (l) → 2IBr(g) ∆G° = 7.3 kJ (c) 2LiOH(s) + CO2 (g) → Li2CO3 (s) + H2O(g) ∆G° = - 79 kJ (d)N2O3(g)→NO(g)+NO2(g) ∆G°= -1.6kJ (e) SnCl4 (l) → SnCl4 (l) ∆G° = 8.0 kJ
(a) K= 41 (b) K= 0.053 (c) K= 6.9X10^13 (d) K=1.9 (e) K= 0.04
Write cell schematics for the following cell reactions, using platinum as an inert electrode as needed. (a) Mg(s) + Ni2+ (aq) → Mg2+ (aq) + Ni(s)( b) 2Ag+ (aq) + Cu(s) → Cu2+ (aq) + 2Ag(s) (c) Mn(s) + Sn(NO3 )2 (aq) → Mn(NO3 )2 (aq) + Sn(s) (d) 3CuNO3 (aq) + Au(NO3 )3 (aq) → 3Cu(NO3 )2 (aq) + Au(s)
(a) Mg(s) | Mg2+ (aq) Ni+ (aq) | Ni(s) ; (b) Stoichiometric coefficients do not appear in cell notation Cu(s) | Cu2+ (aq) Ag+ (aq) | Ag(s) ; (c) Spectator ions do not appear in cell notation Mn(s) | Mn2+ (aq) Sn2+ (aq) | Sn(s) (d) Neither stoichiometric coefficients nor spectator ions appear in cell notation. Platinum electrode needed Pt(s) | Cu+ (aq), Cu2+ (aq) Au3+ (aq) | Au(s)
Predict the sign of the enthalpy change for the following processes. Give a reason for your prediction. (a) Pb2+ (aq) + S2− (aq) → PbS(s) (b) 2Fe+3O2 →Fe2O2 (c) 2C6H14 (l) + 19O2 (g) → 14H2O(g) + 12CO2 (g)
(a) Negative. The relatively ordered solid precipitating decreases the number of mobile ions in solution. (b) Negative. There is a net loss of three moles of gas from reactants to products. (c) Positive. There is a net increase of seven moles of gas from reactants to products.
One of the important reactions in the biochemical pathway glycolysis is the reaction of glucose-6-phosphate (G6P) to form fructose-6-phosphate (F6P): G6P → F6P ΔG° =1.7 kJ ← (a) Is the reaction spontaneous or nonspontaneous under standard thermodynamic conditions? (b) Standard thermodynamic conditions imply the concentrations of G6P and F6P to be 1 M, however, in a typical cell, they are not even close to these values. Calculate ΔG when the concentrations of G6P and F6P are 120 μM and 28 μM respectively, and discuss the spontaneity of the forward reaction under these conditions. Assume the temperature is 37 °C.
(a) Nonspontaneous as ∆G° > 0 ; (b) ΔG = ΔG° + RT lnQ, ∆G = 1.7× 103+ 8.314× 310× ln 28 =− 2.1 kJ . The forward reaction to produce 120 F6P is spontaneous under these conditions.
Given: P (s) +5O (g) → P O (s) 4 2 410 2H (g) + O (g) → 2H O(g) 222 ∆G° = ∆G° 26−97.0 kJ/mol= 45−7.18 kJ/mol 6H O(g) + P O (s) → 4H PO (l) 2 410 34 ∆G° = 42−8.66 kJ/mol (a) Determine the standard free energy of formation, ∆G° , for phosphoric acid. (b) How does your calculated result compare to the value in Appendix G? Explain.
(a) The standard free energy of formation is the standard free energy change for 1 P4 (s) + 3 H2 (g) + 2O2 (g) → H3PO4 (l) . We can use a Hess's law-like approach. Note 42 that adding the first reaction plus three times the second reaction plus the third reaction gives, after cancelling terms P4 (s) + 6H2 (g) + 8O2 (g) → 4H3PO4 (l)∆G° = (− 2697.0)+ 3−( 457.18)+ −( 428.66) kJ/mol =− 4497.2 kJ/mol rxn Dividing this result by four gives the equation of interest. The standard free energy of formation is -1124.3kJ/mol. (b) The calculation agrees with the value in Appendix G because free energy is a state function (just like the enthalpy and entropy), so its change depends only on the initial and final states, not the path between them.
Consider the decomposition of red mercury(II) oxide under standard state conditions. 2HgO(s, red) → 2Hg(l) + O2 (g) (a) Is the decomposition spontaneous under standard state conditions? (b) Above what temperature does the reaction become spontaneous?
(a) Using the data in Appendix G, determine ∆G° : 298 ∆G° = 2 G∆° (Hg(l)) + G∆° (O (g)) −2 G∆° (HgO(s, red)) ff2f = 2(0) 0 + 2(−58.5−) kJ/mol 117=.0 kJ/mol {} From its value at 298.15 K, the reaction is nonspontaneous; (b) requires the ratio of the standard enthalpy change to the standard entropy change: ∆H° = 2 H∆° (Hg(l)) + H∆° (O (g)) −2 H∆° (HgO(s, red)) ff2f = [2(0) + 0 - 2(-90.83)]kJ/mol = 181.66 kJ/mol ∆S° = 2S° (Hg(l)) S° (O (g+)) 2S° (HgO−(s, red)) 2 = [2(75.9) + 205.0 - 2(70.29)]J/K•mol = 216.42 J/K•mol ∆H° 181.66 × 103 J/molT= = =839K566=C ° ∆S° 216.42 J/KmolAbove 566 °C the process is spontaneous.
Classify the following as acid-base reactions or oxidation-reduction reactions: (a) Na2S(aq) + 2HCl(aq) → 2NaCl(aq) + H2S(g) (b) 2Na(s) + 2HCl(aq) → 2NaCl(aq) + H2 (g) (c) Mg(s) + Cl2 (g) → MgCl2 (s)(d) MgO(s) + 2HCl(aq) → MgCl2 (aq) + H2O(l) (e) K3P(s) + 2O2 (g) → K3PO4 (s) (f) 3KOH(aq) + H3PO4 (aq) → K3PO4 (aq) + 3H2O(l)
(a) acid-base; (b) oxidation-reduction: Na is oxidized, H is reduced; (c) oxidation-reduction: Mg is oxidized, Cl2 is reduced; (d) acid-base; (e) oxidation-reduction: P3- is oxidized, O2 is reduced; (f) acid-base
Predict the conditions (high T, low T, all T, no T) under which each of the following will be spontaneous. (a) C3H8(g) + 5 O2(g) --> 3 CO2(g) + 4 H2O(g) high T (b) 2 H(g) --> H2(g) low T (c) 2 NO2(g) --> 2 NO(g) + O2(g) (endothermic) all T
(a) all T (b) low T (c) All T
Indicate whether the following processes are spontaneous or nonspontaneous. (a) Liquid water freezing at a temperature below its freezing point (b) Liquid water freezing at a temperature above its freezing point (c) The combustion of gasoline (d) A ball thrown into the air (e) A raindrop falling to the ground (f) Iron rusting in a moist atmosphere
(a) spontaneous; (b) nonspontaneous; (c) spontaneous; (d) nonspontaneous; (e) spontaneous; (f) spontaneous
Calculate ΔG° for each of the following reactions from the equilibrium constant at the temperature given. (a)N(g)+O(g)→2NO(g) T=2000°C K =4.1 10-4 × 22P (b) H2(g) + I2(g) → 2HI(g) T = 400 °C KP = 50.0 (c) CO2(g) + H2(g) → CO(g) + H2O(g) T = 980 °C KP = 1.67 (d) CaCO3(s) → CaO(s) + CO2(g) T = 900 °C KP = 1.04 (e)HF(aq)+HO(l)→HO+(aq)+F-(aq) T=25 C ° K =7.2 10-4× (f)AgBr(s)→Ag(aq)+Br(aq) T=25 C ° K =3.3 10 × P
(a) ΔG° = -RT ln KP = -(8.314 J K-1)(2273.15 K)(ln 4.1 × 10-4) = 147 kJ = 1.5 × 102 kJ (b) ΔG° = -(8.314 J K-1)(673.15 K)(ln 50.0) = -21,893 J = -21.9 kJ (c) ΔG° = -(8.314 J K-1)(1253.15 K)(ln 1.67) = -5.34 kJ (d) ΔG° = -(8.314 J K-1)(1173.15 K)(ln 1.04) = -0.383 kJ (e) ΔG° = -(8.314 J K-1)(298.15 K)(ln 7.2 × 10-4) = 17,937 J = 18 kJ (f) ΔG° = -(8.314 J K-1)(298.15 K)(ln 3.3 × 10-13) = 71,240 J = 71 kJ
Use the standard free energy of formation data in Appendix G to determine the free energy change for each of the following reactions, which are run under standard state conditions and 25 °C. Identify each as either spontaneous or nonspontaneous at these conditions. (a) MnO2 (s) → Mn(s) + O2 (g) (b) H2 (g) + Br2 (l) → 2HBr(g) (c) Cu(s) + S(g) → CuS(s) (d) 2LiOH(s) + CO2 (g) → Li2CO3 (s) + H2O(g) (e) CH4 (g) + O2 (g) → C(s, graphite) + 2H2O(g) (f) CS2 (g) + 3Cl2 (g) → CCl4 (g) + S2Cl2 (g)
(a)∆G°= 465.1 kJ (non spontaneous) (b)∆G°= -106.86 kJ (spontaneous) (c)∆G°= -291.9 kJ (spontaneous) (d)∆G°= -83.4 kJ (spontaneous) (e)∆G°= -406.7 kJ (spontaneous) (f)∆G°= -154.3 kJ (spontaneous)
What is the oxidation number of As in the compound K(NH4)2AsO4•6H2O?
+5
Consider the following cell at 298 K: Pt (s)│H2 (g, 1.0 atm) │H+ (aq, 1.0 M)║CH3CO2H (aq, 1.0 M)│H2 (g, 1.0 atm)│Pt (s) The Ka of CH3CO2H is 1.8 × 10−5. Calculate Ecell. Enter your answer to 2 decimal places.
-0.14
Fe2+ (aq) undergoes a disproportionation reaction to give Fe (s) and Fe3+ (aq). In other words, Fe2+ can be oxidized and reduced with the same reaction: 3 Fe2+ (aq) ⇌ Fe (s) + 2 Fe3+ (aq) Use standard reduction potentials for to calculate the overall cell potential (V).
-1.22
Write a balanced chemical equation for the reaction of Al3+ (aq) with Sn (s). The overall standard potential is ________V so the reaction is ____________.
-1.52; non spontaneous
The Trinity test, the first successful test of a nuclear weapon, released the energy equivalent of 20 kilotons (metric) of TNT, C7H5N3O6. The detonation of TNT releases 950 kJ/mol Calculate the energy change (in kJ/mol) for the fission of 235U
-1.789x10^13 J/mol -1.789x10 kJ/mol
Ozone in the lower atmosphere is a pollutant that can be formed by the oxidation of hydrocarbons: CH4(g) + 8 O2(g) --> CO2(g) + 2 H2O(g) + 4 O3(g) Calculate ΔGo (kJ/mol) for this reaction at 36 oC (97 oF). Enter your answer to 1 decimal place.
-144.9 kJ/mol
Ozone in the lower atmosphere is a pollutant that can be formed by the oxidation of hydrocarbons: CH4(g) + 8 O2(g) --> CO2(g) + 2 H2O(g) + 4 O3(g) Calculate ΔGo (kJ/mol) for this reaction at −5 oC (23 oF). Enter your answer to 1 decimal place.
-156.4 kJ/mol
The amount of current a car battery can deliver for 30 s at 0 oC is referred to as the battery's cold cranking amps (CCA). If one were to use a car battery to electrolyze water, what volume of O2 would be produced at 1.00 atm and 273 K if the battery delivered 850 amps for 30 seconds?
0.06607 mol O2 1.48 L
Consider this cell: Fe (s)│FeCl2 (aq)║Co(NO3)2 (aq)│Co (s). The initial concentration of FeCl2 is 0.250 M, but the initial cobalt(II) concentration at the cathode is unknown. Calculate [Fe2+] at equilibrium. Enter your answer to 3 decimal places.
0.251
For this voltaic cell, Eo= _______V.
0.694
An extremely important reaction in biochemistry is the 4-electron reduction of molecular oxygen: O2 (g) + 4 H+ (aq) + 4 e− --> 2 H2O (l) Eo = 1.229 V In calculations involving biological systems, biochemists use Eo' rather than Eo, where Eo' is the cell potential at pH = 7.00. Calculate the cell potential (V) at this pH. Enter your answer to 3 decimal places.
0.815
Consider the following voltaic cell at 298 K: Cd(s) | Cd2+(aq, 0.5 M) || Cu+ (aq, 0.75 M) | Cu(s) Estimate the concentration (M) of Cd2+(aq) when this cell stops running. Enter your answer to 2 decimal places.
0.88
Consider the following voltaic cell at 298 K: Cd(s) | Cd2+(aq, 0.5 M) || Cu+ (aq, 0.75 M) | Cu(s) Calculate the standard cell potential. Enter your answer to 2 decimal places.
0.92
A chunk of solid CO2 is placed in a closed 2.00-L bottle at 360 K. The bottle is capable of withstanding pressures up to 200 atm. Use standard thermodynamic data to determine the pressure (atm) as the dry ice establishes equilibrium with gaseous CO2. Enter your answer as an integer
1,288
A voltaic cell has the following initial state: Zn (s) | Zn2+(aq, 1.0 M) || Cu+ (aq, 1.0 M) | Cu (s) Each half-cell has a volume of 1.00 L.
1.28
A voltaic cell has the following initial state: Zn (s) Zn2+ (aq, 1.0 M) Cu+ (aq, 2.0 M) Cu (s) Each half-cell has a volume of 1.00 L. Calculate the standard cell potential
1.28V
Consider this cell: Pt (s) | H2 (g, 1.0 atm) | HCl (aq, 1.0 M) || Au3+ (aq, 1.0 M) |Fe (s)
1.5
One half-cell contains a silver wire dipped into a 0.25 M solution of AgNO3. Another half-cell contains a zinc electrode in a 0.010 M solution of Zn(NO3)2. Calculate the overall cell potential (V) when the half-cells are connected together. Enter your answer to 2 decimal places.
1.54V (spontaneous)
A voltaic cell has the following initial state: Zn (s) Zn2+ (aq, 1.0 M) Cu+ (aq, 2.0 M) Cu (s) Each half-cell has a volume of 1.00 L. Calculate the equilibrium constant, K, for this reaction
1.82x10^43
Use standard thermodynamic data to calculate the pH of a 0.15 M solution of NH3(aq) at 65 oC. Enter your answer to 2 decimal places.
10.10
Determine the entropy change for the combustion of gaseous propane, C3H8, under the standard conditions to give gaseous carbon dioxide and water.
100.6 J/K
Calculate the number of grams of uranium-235 that must undergo fission in order to release the same energy calculated in Part A
1098 g
An electrochemical cell contains a copper wire, a saturated solution of copper(I) iodide, and an amount of solid copper(I) iodide versus S.H.E.: S.H.E.║CuI (aq)│Cu (s) The potential of this cell was measured to be 0.169 V. Calculate Ksp of CuI. Then, enter the negative log to 2 decimal places.
11.85
The half-life of tritium (3H) is 12.3 yr. How much of a 48.0-mg sample of tritium remains after 21.0 years?
14.7 mg
Carbon monoxide is toxic because it binds much more strongly to the iron in hemoglobin (Hb) than does O2. Consider the following reactions at 310 K: Hb + O2 ⇌ HbO2 ΔGo = -1.51 kJ/mol Hb + CO ⇌ HbCO K = 306 Using these data, calculate the equilibrium constant for the following reaction at 310 K. Enter your answer as an integer. HbO2 + CO ⇌ HbCO + O2
170.28
A voltaic cell has the following initial state: Zn (s) Zn2+ (aq, 1.0 M) Cu+ (aq, 2.0 M) Cu (s) Each half-cell has a volume of 1.00 L. Calculate the number of hours the cell can run at 25 A (1 A= 1 C/s) before reaching equilibrium
2.1 hr
Use standard thermodynamic data to calculate the pOH of a 0.15 M solution of NH3(aq) at 65 oC. Enter your answer to 2 decimal places.
2.7
An oil painting attributed to Rembrandt (1606-1669) is checked by 14C dating. The 14C content of the canvas is 0.975 times that of a living plant. Could the painting have been done by Rembrandt?
209 years
Consider the following reaction occurring at 298 K: BaCO3(s) ⇌ BaO(s) + CO2(g) Show that the reaction is not spontaneous under standard conditions by calculating ΔGrxno (kJ/mol). Enter your answer to 1 decimal place.
219.6
Consider the following reaction: H2 (g) + I2 (g) ⇌ 2 HI (g) The data show the equilibrium constant for this reaction measured at several different temperatures. Use the data to find ΔSo (J/mol*K) for the reaction. Enter your answer as an integer.
225
Consider this cell: Fe (s)│FeCl2 (aq)║Co(NO3)2 (aq)│Co (s). The initial concentration of FeCl2 is 0.250 M, but the initial cobalt(II) concentration at the cathode is unknown. For a cell potential of 0.096 V, calculate [Co2+]. Then, take its negative log and enter your answer to 2 decimal places.
3.10
The half-life of radium-226 is 1600 yr. How long will it take for a 2.5-g sample to decay to 0.60 g?
3300 years
Consider the following reaction occurring at 298 K: BaCO3(s) ⇌ BaO(s) + CO2(g) If BaCO3 is placed in an evacuated flask, what partial pressure of CO2 will be present when the reaction reaches equilibrium? Enter the negative logarithm of the pressure (atm) to 2 decimal places.
38.5
Determine the cell reaction and standard cell potential at 25 °C for a cell made from an anode half-cell containing a cadmium electrode in 1 M cadmium nitrate and an anode half-cell consisting of an aluminum electrode in 1 M aluminum nitrate solution. Is the reaction spontaneous at standard conditions?
3Cd(s) + 2Al3+ (aq) → 3Cd2+ (aq) + 2Al(s) Eo= -1.259V (non spontaneous)
Write the balanced cell reaction for the cell schematic below, calculate the standard cell potential, and note whether the reaction is spontaneous under standard state conditions Cu(s) │Cu2+ (aq) || Au3+ (aq) │Au(s)
3Cu(s) + 2Au3+ (aq) → 3Cu2+ (aq) + 2Au(s) Eo= +1.16V (spontaneous)
A voltaic cell consists of the following: V (s)│V2+ (aq)║saturated Mn(IO3)2 (aq)│Mn (s) The Ksp of Mn(IO3)2 is 4.37 × 10−7, and the concentration of V2+ (aq) is unknown. Determine the concentration of V2+ at which this reaction stops. [Hint: the solution of Mn(IO3)2 is saturated.] Then, take its negative log and enter your answer to 2 decimal places.
4.34
A voltaic cell consists of the following: V (s)│V2+ (aq)║saturated Mn(IO3)2 (aq)│Mn (s) The Ksp of Mn(IO3)2 is 4.37 × 10−7, and the concentration of V2+ (aq) is unknown. The cell potential under these conditions is measured to be 0.0101 V. Determine the concentration of V2+. Then, take its negative log and enter your answer to 2 decimal places.
4.70
Calculate the temperature (Kelvin) at which iodine sublimes spontaneously under standard conditions. Enter your answer to 1 decimal place.
431.8
Fe2+ (aq) undergoes a disproportionation reaction to give Fe (s) and Fe3+ (aq). In other words, Fe2+ can be oxidized and reduced with the same reaction: 3 Fe2+ (aq) ⇌ Fe (s) + 2 Fe3+ (aq) Calculate the equilibrium constant for the disproportionation reaction above. Then, enter the negative log of that constant to 2 decimal places.
5.85x10^-42 or 41.25
Consider the following reaction: H2 (g) + I2 (g) ⇌ 2 HI (g) The data show the equilibrium constant for this reaction measured at several different temperatures. Use the data to find ΔHo (kJ/mol) for the reaction. Enter your answer to 1 decimal place.
50.6
A metal forms the salt MCl3. Electricity is passed through molten MCl3, with a current of 0.700 A for 6.63 hr, producing 3.00 g of the metal, M. Identify the metal
52.0 g/mol Cr
Pt (s) H2 (g, 1.0 atm) CH3COOH (aq, 1.0 M) Au3+ (aq, 1.0 M) Fe (s) Suppose the volume of each half-cell is 1.0 L. If the current is 10.0 A, how much time (in hours) is required for [Au3+] to drop to 0.25 M?
6.03 hr
Above what temperature does the following reaction become nonspontaneous? FeO(s) + CO(g) → CO2(g) + Fe(s) ΔH = −11.0 kJ; ΔS = −17.4 J/K
632 K
The Trinity test, the first successful test of a nuclear weapon, released the energy equivalent of 20 kilotons (metric) of TNT, C7H5N3O6. The detonation of TNT releases 950 kJ/mol How much engird, in kJ, was released by this explosion?
8.36x10^10 kJ
The Trinity test, the first successful test of a nuclear weapon, released the energy equivalent of 20 kilotons (metric) of TNT, C7H5N3O6. The detonation of TNT releases 950 kJ/mol. Calculate the total amount of mass (in g) converted to energy during the Trinity test (part A)
9.30x10^-4 kg 0.930 g
An active (metal) electrode was found to lose mass as the oxidation-reduction reaction was allowed to proceed. Was the electrode an anode or a cathode? Explain.
Active electrodes participate in the oxidation-reduction reaction. Since metals form cations, the electrode would lose mass if metal atoms in the electrode were to oxidize and go into solution. Oxidation occurs at the anode.
Consider the following half-reactions: Half-reaction Eo (V) Cu2+ (aq) + 2 e− --> Cu (s) +0.34 Sn2+ (aq) + 2 e− --> Sn (s) −0.14 Fe2+ (aq) + 2 e− --> Fe (s) −0.44 Zn2+ (aq) + 2 e− --> Zn (s) −0.76 Al3+ (aq) + 3 e− --> Al (s) −1.66 Based on Eo values, which metal is the most easily oxidized?
Al
Which has the larger binding energy, Mg-26 or Al-26
Al: 2.045x10^12 J/mol Mg: 2.091x10^13 J/mol
Many plastic materials are organic polymers that contain carbon and hydrogen. The oxidation of these plastics in air to form carbon dioxide and water is a spontaneous process; however, plastic materials tend to persist in the environment. Explain.
Although the oxidation of plastics is spontaneous, the rate of oxidation is very slow. Plastics are therefore kinetically stable and do not decompose appreciably even over relatively long periods of time.
Which element below is the best reducing agent? Which is the worst? Cu Zn Fe Ag Cr
Best: Zn Worse: Ag
Indicate which substance in the given pairs has the higher entropy value. Explain your choices. (a) C2H5OH(l) or C3H7OH(l) (b) C2H5OH(l) or C2H5OH(g) (c) 2H(g) or H(g)
C3H7OH(l) as it is a larger molecule (more complex and more massive), and so more microstates describing its motions are available at any given temperature. (b) C2H5OH(g) as it is in the gaseous state. (c) 2H(g), since entropy is an extensive property, and so two H atoms (or two moles of H atoms) possess twice as much entropy as one atom (or one mole of atoms).
Write the balanced chemical equation for the combustion of benzene, C6H6(l), to give carbon dioxide and water vapor. Would you expect ΔS to be positive or negative in this process?
C6H6 (l) + 7.5O2 (g) → 3H2O(g) + 6CO2 (g) There are 7.5 moles of gas initially, and 3 + 6 = 9 moles of gas in the end. Therefore, it is likely that the entropy increases as a result of this reaction, and ΔS is positive.
Predict the oxidation state for: CN-
C: +2 N: -3
C6H6
C: -1 H: +1
Predict the oxidation state for: Ch3cH2OH
C: -3 H: +1 c: -1 O: -2
Predict the oxidation state for: CH2NH
C: 0 H: +1 N: -3
In the reaction below, which species undergoes oxidation? CH4 + 2 O2 --> CO2 + 2 H2O
CH4
In the reaction below, ____ is the oxidizing agent. 3 Ce4+(aq) + Cr(s) --> 3 Ce3+(aq) + Cr3+(aq)
Ce^4+ (aq)
In the reaction below, ____ is the reducing agent. 3 Ce4+(aq) + Cr(s) --> 3 Ce3+(aq) + Cr3+(aq)
Cr (s)
Which of the following statements is TRUE?
Endothermic processes decrease the entropy of the surrounds, at constant T and P
Predict the oxidation state for: Fe2O3
Fe: +3 O: -2
One half-cell contains a silver wire dipped into a 0.25 M solution of AgNO3. Another half-cell contains a zinc electrode in a 0.010 M solution of Zn(NO3)2. For a voltaic cell, which half-cell is the cathode? For a voltaic cell, which half-cell is the cathode?
For a voltaic cell, which half-cell is the cathode? silver half-cell For a voltaic cell, which half-cell is the anode? zinc half-cell
Balance each of the following equations according to the half-reaction method: (a) MnO -(aq)+NO -(aq) → MnO (s)+NO -(aq)(in base) (b) MnO 2- (aq) → MnO - (aq) + MnO (s) (in base) (c) Br2 (l) + SO2 (g) → Br- (aq) + SO42- (aq) (in acid)
For an example of the fully worked out solution, see the solution to Exercise 37. (a) 2MnO -(aq)+3NO -(aq)+H O(l) → 2MnO (s)+3NO -(aq)+2OH−(aq);(b) 42223 3MnO 2-(aq)+2H O(l) → 2MnO -(aq)+4OH−(aq)+MnO (s)(inbase);(c) 4242 Br2 (l) + SO2 (g) + 2H2O(l) → 4H+ (aq) + 2Br- (aq) + SO42- (aq)
Balance the half-reaction in acidic and in basic solution, and answer the following questions. H2O2 --> O2
For the reaction in acidic solution, the number of electrons transferred is 2, and they should be on the right side. H+ should be on the right side with a coefficient of 2. H2O should be on neither side with a coefficient of 0. For the reaction in basic solution, OH- should be on the left side with a coefficient of 2. H2O should be on the right side with a coefficient of 2.
Balance the half-reaction in acidic and in basic solution, and answer the following questions. Cr2O72− --> Cr3+
For the reaction in acidic solution, the number of electrons transferred is 6, and they should be on the left side. H+ should be on the left side with a coefficient of 14. H2O should be on the right side with a coefficient of 7. For the reaction in basic solution, OH_ should be on the right side with a coefficient of 14. H2O should be on the left side with a coefficient of 7.
H2C=c=O
H: +1 C:-2 c: +2 O: -2
Predict the oxidation state for: K2O2
K: +1 O: -1
Predict the oxidation state for: Mn2O7
Mn: +7 O: -2
A cylinder contains 3.50 atm NO2(g) and 5.00 atm N2O4(g) at 425 K. What reaction will occur spontaneously?
N2O4(g) ⇌ 2 NO2(g)
Predict the oxidation state for: NO2
N: +4 O: -2
Predict the oxidation state for: N2O5
N: +5 O: -2
Predict the oxidation state for: NaH
Na: +1 H: -1
The decomposition of N2H4(g) to its elements is endothermic. This reaction will be spontaneous at _____.
Only high T
Predict the oxidation state for: PO3^3-
P: +3 O: -2
In the reaction below, which species undergoes reduction? Pb(s) + PbO2(s) + 2 H+(aq) + 2 HSO4-(aq) --> 2 PbSO4(s) + 2 H2O(l)
PbO2
Predict the oxidation state for: SO3
S: +6 O: -2
Identify the oxidant and reductant in each reaction of the previous exercise. (a) Mg(s) + Cu2+ (aq) → Mg2+ (aq) + Cu(s) (b) 2Ag+ (aq) + Ni(s) → Ni2+ (aq) + 2Ag(s)
Species oxidized = reducing agent: (a) Al(s); (b) NO(g); (c) Mg(s) and (d) MnO2(s). Species reduced = oxidizing agent: (a) Zr4+(aq); (b) Ag+(aq); (c) SiO 2- (aq); and (d) ClO − (aq).
Complete and balance each of the following half-reactions (steps 2-5 in half-reaction method): (a) Sn4+ (aq) → Sn2+ (aq) (b) Ag(NH ) + (aq) → Ag(s) + NH (aq) 323 (c) Hg Cl (s) → Hg(l) + Cl− (aq) 22 (d) H2O(l) → O2 (g) (in acidic solution) (e) IO − (aq) → I (s) 32 (f) SO32- (aq) → SO42- (aq) (in acidic solution) (g) MnO - (aq) → Mn 2+ (aq) (in acidic solution) 4 (h) Cl- (aq) → ClO - (aq) (in basic solution)
Stiochiometry of Chemical Reactions Key #39
dentify the strongest oxidant and the strongest reductant. Au Cl2 Cr Ni2+ O2 Sn
Strongest oxidant: Cl2 Strongest reductant: Cr
At room temperature, the entropy of the halogens increases from I2 to Br2 to Cl2. Explain.
The masses of these molecules would suggest the opposite trend in their entropies. The observed trend is a result of the more significant variation of entropy with a physical state. At room temperature, I2 is a solid, Br2 is a liquid, and Cl2 is a gas.
Carbon dioxide decomposes into CO and O2 at elevated temperatures. What is the equilibrium partial pressure of oxygen in a sample at 1000 °C for which the initial pressure of CO2 was 1.15 atm?
The reaction is 2CO2(g) → 2CO(g) + O2(g). x= 1.29x10^-5 atm
Determine the standard free energy change, ∆Gf , for the formation of S (aq) given that the ∆G° for Ag+(aq) and Ag2S(s) are 77.1 kJ/mole and -39.5 kJ/mole respectively, and the fsolubility product for Ag2S(s) is 8 × 10 .
The reaction of interest is: Ag S(s) -----> 2Ag+ (aq) + S2− (aq) K = 8 × 10−51 ΔG° = 286kJ ΔGf S = 90 kJ/mol
In Figure 16.8, all of the possible distributions and microstates are shown for four different particles shared between two boxes. Determine the entropy change, ΔS, for the system when it is converted from distribution (b) to distribution (d)
There are four initial microstates and four final microstates. ∆S=kln wf =1.38×10−23 J/K×ln 4 =0 wi 4
Predict a likely mode of decay for each of the following radionuclides Ba-123
Too few neutrons: positron emission
Predict a likely mode of decay for each of the following radionuclides Se-83
Too many neutrons: Beta Decay
Identify the change in state that does not have an increase in entropy.
Water freezing
When ammonium chloride is added to water and stirred, it dissolves spontaneously and the resulting solution feels cold. Without doing any calculations, deduce the signs of ΔG, ΔH, and ΔS for this process, and justify your choices.
When ammonium chloride is added to water and stirred, it dissolves spontaneously and the resulting solution feels cold. Without doing any calculations, deduce the signs of ΔG, ΔH, and ΔS for this process, and justify your choices.
Predict a likely mode of decay for each of the following radionuclides Fr-223
Z > 83: alpha decay
Consider the following half-reactions: Half-reaction Eo (V) Cu2+ (aq) + 2 e− --> Cu (s) +0.34 Sn2+ (aq) + 2 e− --> Sn (s) −0.14 Fe2+ (aq) + 2 e− --> Fe (s) −0.44 Zn2+ (aq) + 2 e− --> Zn (s) −0.76 Al3+ (aq) + 3 e− --> Al (s) −1.66 Which metals on this list are capable of reducing Fe2+ to Fe? Select all that apply.
Zn and Al
Pt (s) H2 (g, 1.0 atm) CH3COOH (aq, 1.0 M) Au3+ (aq, 1.0 M) Fe (s) Will the iron electro become gold-plated spontaneously? Justify by calculating E. CH3COOH: Ka = 1.8 × 10−5
[H+]= 4.24x10^-3 1.69V (spontaneous)
The reaction below is spontaneous under standard conditions. C2H2 (g) + 4 Cl2 (g) --> 2 CCl4 (l) + H2 (g) In order to get the maximum yield of CCl4, should this reaction be carried out at low T or high T?
low T
Predict a likely mode of decay for each of the following radionuclides Sb-132
too many neutrons: Beta decay
A reactions has ∆H° = 100 kJ/moland ∆S° = 250 J/molK. Is the reaction spontaneous at room temperature? If not, under what temperature conditions will it become spontaneous?
ΔG°= 25.5 kJ/mol T= 400 K
Determine the sign of ΔSsys and ΔSsurr for the reaction below. 2 H2(g) + O2(g) --> 2 H2O(g) ΔHo = −484 kJ/mol
ΔSsys: negative ΔSsurr: positive
Determine the sign of ΔSsys and ΔSsurr for the reaction below. N2(g) + O2(g) --> 2 NO(g) ΔHo = 113 kJ/mol
ΔSsys: positive ΔSsurr: negative
Determine the sign of ΔSsys and ΔSsurr for the reaction below. the isothermal (constant temperature) expansion of an ideal gas
ΔSsys: positive ΔSsurr: negative
Determine the sign of ΔSsys and ΔSsurr for the reaction below. CaS (s) + 2 HCl (aq) --> CaCl2 (aq) + H2S (g) ΔHo = −258 kJ/mol
ΔSsys: positive ΔSsurr: positive
Consider the following reaction at 298 K: NO(g)---->2NO(g) K =0.142 What is the standard free energy change at this temperature? Describe what happens to the initial system, where the reactants and products are in standard states, as it approaches equilibrium.
∆G°= 4.84
By calculating ΔSuniv at each temperature, determine if the melting of 1 mole of NaCl(s) is spontaneous at 500 °C and at 700 °C. S° = 72.11 J S° = 95.06 J ∆H° = 27.95kJ/mol NaCl(s) mol*K NaCl(l) mol*K fusion What assumptions are made about the thermodynamic information (entropy and enthalpy values) used to solve this problem?
∆S uni = -13.2 ∆S uni - 5.8
