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chemistry principles and practice
Chemistry Principles And Practice 3rd Edition Daniel L. Reger, Scott R. Goode, David W. Ball - Solutions
Calculate ΔH °, ΔS °, and ΔG ° for each of the following reactions at 298 K. State whether the direction of spontaneous reaction is consistent with the sign of the enthalpy change, the entropy change, or both. Use Appendix G for data.Appendix G
Calculate ΔH °, ΔS °, and ΔG ° for each of the following reactions. State whether the direction of spontaneous reaction is consistent with the sign of the enthalpy change, the entropy change, or both. Use Appendix G for data.
Use standard entropies and heats of formation to calculate ΔG°f at 25°C for
What is the sign of the standard Gibbs free-energy change at low temperatures and at high temperatures for the combustion of acetaldehyde?
What is the sign of the standard Gibbs free-energy change at low temperatures and at high temperatures for the formation of hydrogen sulfide from the elements?
What is the sign of the standard Gibbs free-energy change at low temperatures and at high temperatures for the synthesis of ammonia?
What is the sign of the standard Gibbs free-energy change at low temperatures and at high temperatures for the decomposition of phosgene?
Predict the temperature at which the reaction in Exercise 17.67 comes to equilibrium. Consider the equation ΔG = ΔH - TΔS. At some value of T, ΔG equals zero and the reaction is at equilibrium. Set ΔG = 0, substitute for ΔH ° and ΔS ° (assuming that they do not vary much with temperature)
Predict the temperature at which the reaction in Exercise 17.68 comes to equilibrium.Exercise 17.68What is the sign of the standard Gibbs free-energy change at low temperatures and at high temperatures for the formation of hydrogen sulfide from the elements?
Predict the temperature at which the reaction in Exercise 17.69 comes to equilibrium.Exercise 17.69What is the sign of the standard Gibbs free-energy change at low temperatures and at high temperatures for the synthesis of ammonia?
Predict the temperature at which the reaction in Exercise 17.70 comes to equilibrium.Exercise 17.70What is the sign of the standard Gibbs free-energy change at low temperatures and at high temperatures for the decomposition of phosgene?
Calculate ΔG ° at 400 and 600 K for the following reactions.
Calculate ΔG ° at 300 and 390 K for the following reactions.
Suppose you are looking for a chemical reaction that is spontaneous at low temperatures but proceeds in the reverse direction at high temperatures. What are the signs of ΔH ° and ΔS ° for such a reaction?
Suppose you are looking for a chemical reaction that is spontaneous at high temperatures but proceeds in the reverse direction at low temperatures. What are the signs of ΔH ° and ΔS ° for such a reaction?
Identify which of the following statements are incorrect and change them so that they are true. The statements refer to the formation of 1 mol methanol (CH3OH) from carbon monoxide and hydrogen (all at 1 atm pressure and in the gas phase):(a) The direction of spontaneous reaction depends entirely
Decide whether each of the following statements is true or false. If false, rewrite it to make it true.(a) The entropy of a substance increases on going from the liquid to the vapor state at any temperature.(b) An exothermic reaction will always be spontaneous.(c) Reactions with a positive ΔH °
Determine whether the vaporization of methanol is spontaneous at 80 °C and 1 atm. Use the thermodynamic data in Appendix G. State any assumptions you make.
Determine whether the condensation of nitromethane is spontaneous at 40 °C and 1 atm. Use the thermodynamic data in Appendix G. State any assumptions you make.Thermodynamic Data From Appendix G:
At 298 K, ΔG ° = -70.52 kJ for the reaction(a) Calculate ΔG at the same temperature when PNO = 1.0 × 10-4 atm, PO2 = 2.0 × 10-3 atm, and PNO2 = 0.30 atm.(b) Under the conditions in part a, in which direction is the reaction spontaneous?
At 298 K, ΔG ° = -6.36 kJ for the reaction(a) Calculate ΔG at the same temperature when PN2O = 4.0 × 10-2 atm, PO2 = 4.2 × 10-3 atm, and PN2O4 = 0.40 atm.(b) Under the conditions in part a, in which direction is the reaction spontaneous?
At 298 K, ΔG ° = +27.4 kJ for the reaction(a) Calculate ΔG at the same temperature when [Pb2+] = 4.0 × 10-4 M and [Cl-] = 2.5 × 10-3 M.(b) Under the conditions in part a, in which direction is the reaction spontaneous?
At 298 K, ΔG ° = +11.51 kJ for the reaction(a) Calculate ΔG at the same temperature when [Ca2+] = 3.5 × 10-2 M and [F-] = 2.3 × 10-3 M.(b) Under the conditions in part a, in which direction is the reaction spontaneous?
Calculate the normal boiling point of methanol (CH3OH). Use the thermodynamic data in Appendix G. Compare your answer with the experimentally measured boiling point.Aassume that ΔH ° and ΔS ° do not change with temperature.
Calculate the normal boiling point of nitromethane (CH3NO2). Use the thermodynamic data in Appendix G. Compare your answer with the experimentally measured boiling point.Aassume that ΔH ° and ΔS ° do not change with temperature.
For each reaction, an equilibrium constant at 298 K is given. Calculate ΔG ° for each reaction.Aassume that ΔH ° and ΔS ° do not change with temperature.
Aassume that ΔH ° and ΔS ° do not change with temperature.
Use the standard Gibbs free-energy change to calculate the value of the equilibrium constant for the reactionFor each reaction, an equilibrium constant at 298 K is given. Calculate ΔG ° for each reaction.
Use the data in Appendix G to calculate the value of the equilibrium constant for the reaction 2SO2(g)For each reaction, an equilibrium constant at 298 K is given. Calculate ΔG ° for each reaction.
Suppose you have an endothermic reaction with ΔH ° = +15 kJ and a ΔS ° of +150 J/K. Calculate ΔG ° and Keq at 10, 100, and 1000 K.For each reaction, an equilibrium constant at 298 K is given. Calculate ΔG ° for each reaction.
Suppose you have an endothermic reaction with ΔH ° = +15 kJ and a ΔS ° of -150 J/K. Calculate ΔG ° and Keq at 10, 100, and 1000 K.For each reaction, an equilibrium constant at 298 K is given. Calculate ΔG ° for each reaction.
Suppose you have an exothermic reaction with ΔH ° = -15 kJ and a ΔS ° of -150 J/K. Calculate ΔG ° and Keq at 10, 100, and 1000 K.For each reaction, an equilibrium constant at 298 K is given. Calculate ΔG ° for each reaction.
Suppose you have an exothermic reaction with ΔH ° = -15 kJ and a ΔS ° of -150 J/K. Calculate ΔG ° and Keq at 10, 100, and 1000 K.For each reaction, an equilibrium constant at 298 K is given. Calculate ΔG ° for each reaction.
Calculate ΔG ° and ΔG at 303 °C for the following equation.For each reaction, an equilibrium constant at 298 K is given. Calculate ΔG ° for each reaction.
Calculate ΔG ° and ΔG at 37 °C for the following equation.For each reaction, an equilibrium constant at 298 K is given. Calculate ΔG ° for each reaction.
State whether increasing temperature increases or decreases the value of the equilibrium constant for the following reactions.For each reaction, an equilibrium constant at 298 K is given. Calculate ΔG ° for each reaction.
State whether increasing temperature increases or decreases the value of the equilibrium constant for the following reactions.For each reaction, an equilibrium constant at 298 K is given. Calculate ΔG ° for each reaction.
Calculate the vapor pressure of each of the following at the given temperature.For each reaction, an equilibrium constant at 298 K is given. Calculate ΔG ° for each reaction.
Calculate the vapor pressure of each of the following at the given temperature.For each reaction, an equilibrium constant at 298 K is given. Calculate ΔG ° for each reaction.
A 220-ft3 sample of gas at standard temperature and pressure is compressed into a cylinder, where it exerts pressure of 2000 psi. Calculate the work (inJ) performed when this gas expands isothermally against an opposing pressure of 1.0 atm. (The amount of work that can be done is equivalent to the
What is the sign of the standard Gibbs free-energy change at low temperatures and at high temperatures for the explosive decomposition of TNT? Use your knowledge of TNT and the chemical equation, particularly the phases, to answer this question.
The equilibrium constant for the formation of phosgene is measured at two different temperatures.At 506 °C, Keq = 1.3; at 530 °C, Keq = 0.78. Calculate ΔH ° and ΔS ° for this reaction. Under standard-state conditions, over what temperature range is the reaction spontaneous?
Elemental boron, in the form of thin fibers, can be made by reducing a boron halide with H2.Calculate ΔH °, ΔS °, and ΔG ° at 25 °C for this reaction. Is the reaction predicted to be product favored at equilibrium at 25 °C? If so, is it enthalpy driven or entropy driven?
Calculate the standard Gibbs free-energy change when SO3 forms from SO2 and O2 at 298 K. Why is sulfur trioxide an important substance to study?
The thermite reaction is(a) Calculate ΔG ° for this reaction.(b) Calculate Keq for this reaction.Assume T = 298 K. You may have to do some mathematical manipulations to get your final numerical answer.
Chemists and engineers who design nuclear power plants have to worry about high-temperature reactions because it is possible for water to decompose.(a) Under what conditions does this reaction occur spontaneously?(b) Under conditions where the decomposition of water is spontaneous, do nuclear
Another type of thermite reaction uses Cr2O3 instead of Fe2O3:(a) Calculate ΔH °, ΔS °, and ΔG ° for this reaction.(b) On the assumption that the energy released goes to warming up the products, which reaction generates the greatest temperature, this one or the thermite reaction in Exercise
The reaction of carbon with metal oxide ores is used to isolate metals whenever possible, because carbon is an inexpensive reactant. You are asked to determine the feasibility of using carbon to prepare Ca from lime, which has the formula CaO. The reaction would beAn analysis of the economics shows
A cycloalkane is a hydrocarbon that contains a ring of carbon atoms with two hydrogen atoms bonded to each carbon. Th e standard enthalpy changes for the combustion of several gaseous cycloalkanes to form gaseous water and carbon dioxide have been measured and are summarized in the table
Assign the oxidation numbers to all elements in(a) K2S,(b) NH3,(c) BaO2,(d) Cr2O2-7(e) Br2.Strategy Apply the rules listed earlier, in order, for each element in the formula of the substance.
Determine the oxidation numbers of the elements in the following reaction, and determine whether it is a redox reaction. If so, what is being oxidized, and what is being reduced?Strategy Apply the rules for assigning oxidation numbers and look for elements that are changing oxidation numbers. Any
Complete and balance the following oxidation-reduction reactions.Strategy Use the half-reaction method to balance redox reactions.
Balance the following redox reaction that occurs in basic solution.Strategy Follow the steps for balancing a redox reaction, but add OH-(aq) ions to the balanced half-reactions to convert it to a basic solution.
Consider the cell shown in Figure 18.3. One half-cell consists of silver metal in a silver nitrate solution, and the other half-cell has a piece of copper metal immersed in a copper(II) nitrate solution. A salt bridge that contains a sodium nitrate solution connects the two half-cells. Th e
Calculate the standard potential and state the direction in which the reaction proceeds spontaneously forStrategy Find the two half-reactions in Table 18.1, reversing one of them to make it an oxidation process. Combine the E°s for the two half-reactions to determine the voltage of a voltaic cell
Use Table 18.1 to(a) List metals that are and are not oxidized by H+(aq) under standard conditions.(b) Find an oxidizing agent that will oxidize copper metal.Strategy Consult Table 18.1 as an activity series: The higher half-reaction will be the reduction process of a spontaneous chemical
Calculate the standard free energy change for the reactionStrategy Use Table 18.1 to determine the cell potential; then use Equation 18.2 to calculate the ΔG°. You will also have to determine the number of moles of electrons, n, that are transferred in this reaction.Equation 18.2Table 18.1
What is the equilibrium constant for the following reaction?Strategy Determine the number of electrons transferred, then use Equation 18.4.Equation 18.4
A voltaic cell consists of a half-cell of iron ions in a solution with [Fe2+] = 2.0 M and[Fe3+] = 0.75 M , and a half-cell of copper metal immersed in a solution containing Cu2+ at a concentration of 3.6 × 10-4 M . What is the voltage of this cell?Strategy First, determine E° for the spontaneous
Using standard potentials, predict the electrolysis reaction that occurs, and the standard potential of the electrolysis reaction, for a solution of nickel perchlorate using inert electrodes.Strategy Determine what oxidation and reduction half-reactions might occur, including those involving the
A constant current of 0.500 A passes through a silver nitrate solution for 90.0 minutes .What mass of silver metal is deposited at the anode?Strategy First, determine the half-reaction involved. Then, determine the total charge of the process and relate that to the stoichiometry of the
Assign the oxidation numbers to the elements in ClO-3.
Is the following reaction a redox reaction? If so, what is being oxidized, and what is being reduced?
Balance the following redox reaction.
Balance the following reaction, which occurs in basic solution.
Draw and label the parts of a voltaic cell that is based on the following reaction.Determine which half-cell is the positive electrode, which is the negative electrode, and indicate the direction of electron flow.
Calculate the standard potential of the following reaction, and state whether the reaction is spontaneous as written.
Most disinfectants kill bacteria by oxidizing them. Which substance is the better oxidant,from the point of standard potentials: Cl2 or I2?
From the standard reduction potentials in Table 18.1, find the standard free energy change for the following reaction:Table 18.1
Find the equilibrium constant for the following reaction.
What is the voltage of the iron ion-copper cell when the [Fe2+] = 1.55 M, [Fe3+] 0.066 M, and [Cu2+] = 0.500 M?
Predict the oxidation-reduction reaction that occurs in the electrolysis of a solution of ZnBr2.
The anode reaction in this electrolysis cell is the oxidation of water to O2(g). What volume of O2(g), measured at standard temperature and pressure was produced?
Describe oxidation and reduction. Compare the electron transfer in a redox reaction with the electron donation in a Lewis acid-base reaction.
List the halogens in order of increasing oxidizing power.
Which is a better reducing agent: zinc or mercury?
List four species that can oxidize Fe2+ to Fe3+.
List three species that can reduce Al3+ to Al.
In a “dead” battery, the chemical reaction has come to equilibrium. What are the values of ΔG and E for a dead battery?
What is the difference between a battery and a fuel cell?
What are the differences between anodic and cathodic protection from corrosion?
Assign the oxidation numbers of all atoms in the following species.
Assign the oxidation numbers of all atoms in the following species.
Assign the oxidation numbers of all atoms in the following ions.
Assign the oxidation numbers of all atoms in the following species.
Assign the oxidation numbers of all atoms in the following compounds.
Assign the oxidation numbers of all atoms in the following species.
Assign the oxidation numbers of all atoms in the following species.
Assign the oxidation numbers of all atoms in the following species.(a) KMnO4 (b) H2O (c) Cl2
Assign the oxidation numbers of all atoms in the following species.
Assign the oxidation numbers of all atoms in the following species.
Assign the oxidation numbers of all atoms in the following compounds.(a) KHF2(b) H2Se(c) NaO2(d) C2H6
Assign the oxidation numbers of all atoms in the following compounds.
Assign the oxidation numbers of all atoms in the following species.
Balance the following reactions, and specify which species is oxidized and which is reduced.(a) H2 + O2 → H2O(b) Fe + O2 → Fe2O3(c) Al2O3 + C → Al + CO2
Balance the following reactions, and specify which species is oxidized and which is reduced.(a) Fe2O3 + H2 → Fe + H2O(b) CuCl2 + Na → NaCl + Cu(c) C + O2 → CO2
Balance the following reactions, and specify which species is oxidized and which is reduced.(a) Na + FeCl3 → Fe + NaCl(b) SnCl2 + FeCl3 → SnCl4 + FeCl2(c) CO + Cr2O3 → Cr + CO2
Balance the following reactions, and specify which species is oxidized and which is reduced.(a) Na + Hg2Cl2 → NaCl + Hg(b) HCl + Zn → ZnCl2 + H2(c) H2 + CO2 → CO + H2O
Complete and balance each half-reaction in acid solution,and identify it as an oxidation or a reduction.
Write balanced equations for the following half reactions.Specify whether each is an oxidation or reduction.
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