5.3 Enthalpy

Explain how the heat measured in the reference differs from the enthalpy change for the exothermic reaction described by the following equation:
$\text{HCl} \left(a q\right) + \text{NaOH} \left(a q\right) \rightarrow \text{NaCl} \left(a q\right) + \text{H}_{2} \text{O} \left(l\right)$

Calculate the enthalpy of solution (ΔH for the dissolution) per mole of NH₄NO₃ under the conditions described in the reference.

Calculate the enthalpy of solution (ΔH for the dissolution) per mole of CaCl₂ (refer to the reference).

How much heat is produced by burning 4.00 moles of acetylene under standard state conditions?

How many moles of isooctane must be burned to produce 100 kJ of heat under standard state conditions?

When 2.50 g of methane burns in oxygen, 125 kJ of heat is produced. What is the enthalpy of combustion per mole of methane under these conditions?

A sample of 0.562 g of carbon is burned in oxygen in a bomb calorimeter, producing carbon dioxide. Assume both the reactants and products are under standard state conditions, and that the heat released is directly proportional to the enthalpy of combustion of graphite. The temperature of the calorimeter increases from 26.74 °C to 27.93 °C. What is the heat capacity of the calorimeter and its conte…

Homes may be heated by pumping hot water through radiators. What mass of water will provide the same amount of heat when cooled from 95.0 to 35.0 °C, as the heat provided when 100 g of steam is cooled from 110 °C to 100 °C?

Does the standard enthalpy of formation of H₂O(g) differ from ΔH° for the reaction $\text{2H}_{2} \left(g\right) + \text{O}_{2} \left(g\right) \rightarrow 2 \text{H}_{2} \text{O} \left(g\right) ?$

How many kilojoules of heat will be released when exactly 1 mole of manganese, Mn, is burned to form Mn₃O₄(s) at standard state conditions?

The following sequence of reactions occurs in the commercial production of aqueous nitric acid:
$4 \text{NH}_{3} \left(g\right) + 5 \text{O}_{2} \left(g\right) \rightarrow 4 \text{NO} \left(g\right) + 6 \text{H}_{2} \text{O} \left(l\right) \Delta \text{H} = −907 \text{kJ}$
$2 \text{NO} \left(g\right) + \text{O}_{2} \left(g\right) \rightarrow 2 \text{NO}_{2} \left(g\right) \Delta \text{H} = −113 \text{kJ}$
$3 \text{NO}_{2} + \text{H}_{2} \text{O} \left(l\right) \rightarrow 2 \text{HNO}_{3} \left(a q\right) + \text{NO} \left(g\right) \Delta \text{H} = −139 \text{kJ}$

Determine the total enthalpy change for the production of one mole of aqueous nitric acid by this process. Coproducts of the net reaction include water and nitrogen monoxide.

From the molar heats of formation in Appendix G, determine how much heat is required to evaporate one mole of water: $\text{H}_{2} \text{O} \left(l\right) \rightarrow \text{H}_{2} \text{O} \left(g\right)$

Calculate $\Delta H °$ for the process
$\text{Sb} \left(s\right) + \frac{5}{2} \text{Cl}_{2} \left(g\right) \rightarrow \text{SbCl}_{5} \left(s\right)$

from the following information:
$\begin{matrix} \\ \text{Sb} \left(s\right) + \frac{3}{2} \text{Cl}_{2} \left(g\right) \rightarrow \text{SbCl}_{3} \left(s\right) \Delta H ° = −314 \text{kJ} \\ \text{SbCl}_{3} \left(s\right) + \text{Cl}_{2} \left(g\right) \rightarrow \text{SbCl}_{5} \left(s\right) \Delta H ° = −80 \text{kJ}\end{matrix}$

Calculate ΔH for the process $\text{Hg}_{2} \text{Cl}_{2} \left(s\right) \rightarrow 2 \text{Hg} \left(l\right) + \text{Cl}_{2} \left(g\right)$

from the following information:
$\begin{matrix}\text{Hg} \left(l\right) + \text{Cl}_{2} \left(g\right) \rightarrow \text{HgCl}_{2} \left(s\right) \Delta H = −224 \text{kJ} \\ \text{Hg} \left(l\right) + \text{HgCl}_{2} \left(s\right) \rightarrow \text{Hg}_{2} \text{Cl}_{2} \left(s\right) \Delta H = −41.2 \text{kJ}\end{matrix}$

Calculate the standard molar enthalpy of formation of NO(g) from the following data:
$\begin{matrix}\text{N}_{2} \left(g\right) + 2 \text{O}_{2} \left(g\right) \rightarrow 2 \text{NO}_{2} \left(g\right) \Delta H ° = 66.4 \text{kJ} \\ \text{2NO} \left(g\right) + \text{O}_{2} \left(g\right) \rightarrow 2 \text{NO}_{2} \left(g\right) \Delta H ° = −114.1 \text{kJ}\end{matrix}$

Using the data in Appendix G, calculate the standard enthalpy change for each of the following reactions:

(a) $\text{Si} \left(s\right) + 2 \text{F}_{2} \left(g\right) \rightarrow \text{SiF}_{4} \left(g\right)$

(b) $2 \text{C} \left(s\right) + 2 \text{H}_{2} \left(g\right) + \text{O}_{2} \left(g\right) \rightarrow \text{CH}_{3} \text{CO}_{2} \text{H} \left(l\right)$

(c) $\text{CH}_{4} \left(g\right) + \text{N}_{2} \left(g\right) \rightarrow \text{HCN} \left(g\right) + \text{NH}_{3} \left(g\right) ;$

(d) $\text{CS}_{2} \left(g\right) + 3 \text{Cl}_{2} \left(g\right) \rightarrow \text{CCl}_{4} \left(g\right) + \text{S}_{2} \text{Cl}_{2} \left(g\right)$

The decomposition of hydrogen peroxide, H₂O₂, has been used to provide thrust in the control jets of various space vehicles. Using the data in Appendix G, determine how much heat is produced by the decomposition of exactly 1 mole of H₂O₂ under standard conditions.
$2 \text{H}_{2} \text{O}_{2} \left(l\right) \rightarrow 2 \text{H}_{2} \text{O} \left(g\right) + \text{O}_{2} \left(g\right)$

Calculate the enthalpy of combustion of butane, C₄H₁₀(g) for the formation of H₂O(g) and CO₂(g). The enthalpy of formation of butane is −126 kJ/mol.

The white pigment TiO₂ is prepared by the reaction of titanium tetrachloride, TiCl₄, with water vapor in the gas phase: $\text{TiCl}_{4} \left(g\right) + 2 \text{H}_{2} \text{O} \left(g\right) \rightarrow \text{TiO}_{2} \left(s\right) + 4 \text{HCl} \left(g\right) .$

How much heat is evolved in the production of exactly 1 mole of TiO₂(s) under standard state conditions?

In the early days of automobiles, illumination at night was provided by burning acetylene, C₂H₂. Though no longer used as auto headlamps, acetylene is still used as a source of light by some cave explorers. The acetylene is (was) prepared in the lamp by the reaction of water with calcium carbide, CaC₂:
$\text{CaC}_{2} \left(s\right) + 2 \text{H}_{2} \text{O} \left(l\right) \rightarrow \text{Ca} \left(\text{OH}\right)_{2} \left(s\right) + \text{C}_{2} \text{H}_{2} \left(g\right) .$
Calculate the standard enthalpy of the reaction. The $\Delta H_{\text{f}}^{\circ}$ of CaC₂ is −15.14 kcal/mol.

The enthalpy of combustion of hard coal averages −35 kJ/g, that of gasoline, 1.28 $\times$ 10⁵ kJ/gal. How many kilograms of hard coal provide the same amount of heat as is available from 1.0 gallon of gasoline? Assume that the density of gasoline is 0.692 g/mL (the same as the density of isooctane).

Among the substances that react with oxygen and that have been considered as potential rocket fuels are diborane [B₂H₆, produces B₂O₃(s) and H₂O(g)], methane [CH₄, produces CO₂(g) and H₂O(g)], and hydrazine [N₂H₄, produces N₂(g) and H₂O(g)]. On the basis of the heat released by 1.00 g of each substance in its reaction with oxygen, which of these compounds offers the best possibility as a rocket fu…

Ethylene, C₂H₄, a byproduct from the fractional distillation of petroleum, is fourth among the 50 chemical compounds produced commercially in the largest quantities. About 80% of synthetic ethanol is manufactured from ethylene by its reaction with water in the presence of a suitable catalyst. $\text{C}_{2} \text{H}_{4} \left(g\right) + \text{H}_{2} \text{O} \left(g\right) \rightarrow \text{C}_{2} \text{H}_{5} \text{OH} \left(l\right)$

Using the data in the table in Appendix G, calculate ΔH° for the reaction.

Propane, C₃H₈, is a hydrocarbon that is commonly used as a fuel.

(a) Write a balanced equation for the complete combustion of propane gas.

(b) Calculate the volume of air at 25 °C and 1.00 atmosphere that is needed to completely combust 25.0 grams of propane. Assume that air is 21.0 percent O₂ by volume. (Hint: We will see how to do this calculation in a later chapter on gases—for now use the info…