16. Using the information below, determine the enthalpy change (in kJ) for the reaction of 50.0 g of oxy gas (MW 32.0) with excess diborane gas (B2H6; MW 27.6). B2H6(g) + 302(g) B2O3(s) + 3H20() AH--2147.5 kJ a. -3360 kJ b. -2150 kJ c. -1300 kJ d. -1120 kJ

16. Using the information below, determine the enthalpy change (in kJ) for the reaction of 50.0 g of oxy gas (MW 32.0) with excess diborane gas (B2H6; MW 27.6). B2H6(g) + 302(g) B2O3(s) + 3H20() AH--2147.5 kJ a. -3360 kJ b. -2150 kJ c. -1300 kJ d. -1120 kJ

Chemistry & Chemical Reactivity

9th Edition

ISBN:9781133949640

Author:John C. Kotz, Paul M. Treichel, John Townsend, David Treichel

Publisher:John C. Kotz, Paul M. Treichel, John Townsend, David Treichel

Chapter5: Principles Of Chemical Reactivity: Energy And Chemical Reactions

Section: Chapter Questions

Problem 42PS: You should use care when dissolving H2SO4 in water because the process is highly exothermic. To...

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Benzoic acid, C6H5COOH, is a common standard used in bomb calorimeters, which maintain a constant volume. If 1.20 g of benzoic acid gives off 31, 723 J of energy when burned in the presence of excess oxygen and in a water bath having a temperature of 24.6 C, calculate q, w, H, and U for the reaction.
Before the introduction of chlorofluorocarbons, sulfur dioxide (enthalpy of vaporization, 6.00 kcal/mol) was used in household refrigerators. What mass of SO2 must be evaporated to remove as much heat as evaporation of 1.00 kg of CC12F2 (enthalpy of vaporization is 17.4 kJ/mol)? The vaporization reactions for SO2 and are SO2(l)SO2(g) and CCl2F(l)CCl2F2(g), respectively.
Magnesium sulfate is often used in first-aid hot packs, giving off heat when dissolved in water. A coffee-cup calorimeter at 25C contains 15.0 mL of water at 25C. A 2.00-g sample of MgSO4 is dissolved in the water and 1.51 kJ of heat are evolved. (You can make the following assumptions about the solution: volume=15.0 mL, density=1.00 g/mL, specific heat=4.18J/gC.) (a) Write a balanced equation for the solution process. (b) Is the process exothermic? (c) What is qH2O? (d) What is the final temperature of the solution? (e) What are the initial and final temperatures in F?
In a coffee-cup calorimeter, 1.60 g NH4NO3 is mixed with 75.0 g water at an initial temperature of 25.00C. After dissolution of the salt, the final temperature of the calorimeter contents is 23.34C. Assuming the solution has a heat capacity of 4.18 J/C g and assuming no heat loss to the calorimeter, calculate the enthalpy change for the dissolution of NH4NO3 in units of kJ/mol.
The enthalpy of vaporization for water is 40.65 kJ mol-1. As a design engineer for a project in a desert climate, you are exploring the option of using evaporative cooling. (a) If the air has an average volumetric heat capacity of 0.00130 J cm-3 K-1, what is the minimum mass of water that would need to evaporate in order to cool a 5 m? 5 m room with a 3 m ceiling by 5°F using this method? (b) Is this a spontaneous or nonspontaneous process?
In the 1880s, Frederick Trouton noted that the enthalpy of vaporization of 1 mol pure liquid is approximately 88 times the boiling point, Tb, of the liquid on the Kelvin scale. This relationship is called Troutons rule and is represented by the thermochemical equation liquid gas H = 88 Tb, joules Combined with an empirical formula from chemical analysis, Troutons rule can be used to find the molecular formula of a compound, as illustrated here. A compound that contains only carbon and hydrogen is 85.6% C and 14.4% H. Its enthalpy of vaporization is 389 J/g, and it boils at a temperature of 322 K. (a) What is the empirical formula of this compound? (b) Use Troutons rule to calculate the approximate enthalpy or vaporization or one mole of the compound. Combine the enthalpy of vaporization per mole with that same quantity per gram to obtain an approximate molar mass of the compound. (c) Use the results of parts (a) and (b) to find the molecular formula of this compound. Remember that the molecular mass must be exactly a whole-number multiple of the empirical formula mass, so considerable rounding may be needed.
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You should use care when dissolving H2SO4 in water because the process is highly exothermic. To measure the enthalpy change, 5.2 g of concentrated H2SO4() was added (with stirring) to 135 g of water in a coffee-cup calorimeter. This resulted in an increase in temperature from 20.2 C to 28.8 C. Calculate the enthalpy change for the process H2SO4() H2SO4(aq), in kf/mol.
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