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Write balanced equations that describe the formation of the following compounds from elements in their standard states, and use Appendix C to obtain the values of their standard enthalpies of formation: (a) HBr(g)

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Given the data
N2 (g) + O2 (g) → 2 NO (g)            ΔH = +180.7 kJ
2 NO (g) + O2 (g) → 2 NO2 (g)      ΔH = -113.1 kJ
2 N2O (g) → 2 N2 (g) + O2 (g)       ΔH = -163.2 kJ

use Hess’s law to calculate ΔH for the reaction
N2O (g) + NO2 (g) → 3 NO (g)

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From the enthalpies of reaction
H2 (g) + F2 (g) → 2 HF (g)          ΔH = -537 kJ
C (s) + 2 F2 (g) → CF4 (g)          ΔH = -680 kJ
2 C (s) + 2 H2 (g) → C2H4 (g)     ΔH = +52.3 kJ

calculate ΔH for the reaction of ethylene with F 2:
C2H4 (g) + 6 F2 (g) → 2 CF4 (g) + 4 HF (g)

 

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From the enthalpies of reaction

2 H2(g) + O2(g) → 2 H2O(g)          Δ = -483.6 kJ
             3 O2(g) → 2 O3(g)            Δ = +284.6 kJ

calculate the heat of the reaction

3 H2(g) + O3(g) → 3 H2O(g)

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Calculate the enthalpy change for the reaction
P4O6 (s) + 2 O2 (g) → P4O10 (s)
given the following enthalpies of reaction:
P4 (s) + 3 O2 (g) → P4O6 (s)          ΔH = -1640.1 kJ
P4 (s) + 5 O2 (g) → P4O10 (s)        ΔH = -2940.1 kJ

 

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A 2.200-g sample of quinone (C6H4O2) is burned in a bomb calorimeter whose total heat capacity is 7.854 kJ/°C. The temperature of the calorimeter increases from 23.44°C to 30.57°C. What is the heat of combustion per gram of quinone? Per mole of quinone?

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When a 9.55-g sample of solid sodium hydroxide dissolves in 100.0 g of water in a coffee-cup calorimeter (Figure 5.17), the temperature rises from 23.6°C to 47.4°C. Calculate ΔH (in kJ/mol NaOH) for the solution process

NaOH(s) → Na+(aq) + OH(aq)

Assume that the specific heat of the solution is the same as that of pure water.

 

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The specific heat of iron metal is 0.450 J/g-K. How many J of heat are necessary to raise the temperature of a 1.05-kg block of iron from 25.0°C to 88.5°C?

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Two solid objects, A and B, are placed in boiling water and allowed to come to the temperature of the water. Each is then lifted out and placed in separate beakers containing 1000 g water at 10.0°C. Object A increases the water temperature by 3.50°C; B increases the water temperature by 2.60°C. (a) Which object has the larger heat capacity?

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Consider the decomposition of liquid benzene, C6H6 (l), to gaseous acetylene, C2H2 (g):
C6H6 (l) → 3 C2H2 (g)         ΔH = +630 kJ 

(d) If C6H6 (g) were consumed instead of C6H6 (l), would you expect the magnitude of ΔH to increase, decrease, or stay the same? Explain.

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Consider the decomposition of liquid benzene, C6H6 (l), to gaseous acetylene, C2H2 (g):
C6H6 (l) → 3 C2H2 (g)         ΔH = +630 kJ 

(c) Which is more likely to be thermodynamically favored, the forward reaction or the reverse reaction?

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Consider the decomposition of liquid benzene, C6H6 (l), to gaseous acetylene, C2H2 (g):
C6H6 (l) → 3 C2H2 (g)         ΔH = +630 kJ

(b) What is ΔH for the formation of 1 mol of acetylene?

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Consider the decomposition of liquid benzene, C6H6 (l), to gaseous acetylene, C2H2 (g):
C6H6 (l) → 3 C2H2 (g)         ΔH = +630 kJ
(a) What is the enthalpy change for the reverse reaction?

 

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Consider the combustion of liquid methanol, CH 3OH (l):

CH3OH (l) + 3/2 O2 (g) → CO2 (g) + 2 H2O (l)     ΔH = -726.5 kJ

(d) If the reaction were written to produce H2O (g) instead of H2O (l), would you expect the magnitude of ΔH to increase, decrease, or stay the same? Explain.

 

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Consider the combustion of liquid methanol, CH 3OH (l):

CH3OH (l) + 3/2 O2 (g) → CO2 (g) + 2 H2O (l)     ΔH = -726.5 kJ

(c) Which is more likely to be thermodynamically favored, the forward reaction or the reverse reaction? Explain.

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Consider the combustion of liquid methanol, CH 3OH (l):

CH3OH (l) + 3/2 O2 (g) → CO2 (g) + 2 H2O (l)     ΔH = -726.5 kJ

(b) Balance the forward reaction with whole-number coefficients. What is ΔH for the reaction represented by this equation? Explain.

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Consider the combustion of liquid methanol, CH 3OH (l):

CH3OH (l) + 3/2 O2 (g) → CO2 (g) + 2 H2O (l)     ΔH = -726.5 kJ

(a) What is the enthalpy change for the reverse reaction? Explain.

 

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At one time, a common means of forming small quantities of oxygen gas in the laboratory was to heat KClO3:
2 KClO3 (s) → 2 KCl (s) + 3 O2 (g)   ΔH = -89.4 kJ

The decomposition of KClO3 proceeds spontaneously when it is heated. Do you think that the reverse reaction, the formation of KClO3 from KCl and O2, is likely to be feasible under ordinary conditions? Explain your answer.

 

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At one time, a common means of forming small quantities of oxygen gas in the laboratory was to heat KClO3:
2 KClO3(s) → 2 KCl(s) + 3 O 2(g)   ΔH = -89.4 kJ

For this reaction, calculate ΔH for the formation of (b) 8.57 g of KCI.

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At one time, a common means of forming small quantities of oxygen gas in the laboratory was to heat KClO3:
2 KClO3(s) → 2 KCl(s) + 3 O2(g)   ΔH = -89.4 kJ

For this reaction, calculate ΔH for the formation of (a) 0.632 mol of O 2

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Consider the following reaction: 
CH3OH (g) → CO (g) + 2 H 2 (g)       ΔH = +90.7 kJ

(d) How many kilojoules of heat are released when 50.9 g of CO(g) reacts completely with H 2 (g) to form CH3OH(g) at constant pressure?

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Consider the following reaction: 
CH3OH (g) → CO (g) + 2 H 2 (g)       ΔH = +90.7 kJ

(c) For a given sample of CH3OH, the enthalpy change on reaction is 25.8 kJ. How many grams of hydrogen gas are produced? What is the value of ΔH for the reverse of the previous reaction?

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Consider the following reaction: 
CH3OH (g) → CO (g) + 2 H 2 (g)       ΔH = +90.7 kJ

(b) Calculate the amount of heat transferred when 45.0 g of CH 3OH (g) is decomposed by this reaction at constant pressure.

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Consider the following reaction: 
CH3OH (g) → CO (g) + 2 H 2 (g)       ΔH = +90.7 kJ
(a) Is heat absorbed or released in the course of this reaction?

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Without referring to tables, predict which of the following has the higher enthalpy in each case: (d) 1 mol N2 (g) at 100°C or 1 mol N2 (g) at 300°C.

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