Enthalpy and calorimetry

$A 298\,\text{K}$, the combustion of carbon monoxide gas to give carbon dioxide proceeds with release of energy according to:

$$2\,\mathrm{CO(g)} + \mathrm{O_2(g)} \rightarrow 2\,\mathrm{CO_2(g)} \quad \Delta H_1^\circ = -566.0\,\text{kJ}$$

1. a) Determine the value of the standard enthalpy of the oxidation reaction of carbon monoxide, expressed per mole of $\mathrm{CO_2(g)}$ formed.
2. b) Calculate the standard enthalpy change, $\Delta H_2^\circ$, for the decomposition reaction:

$$\mathrm{CO_2(g)} \rightarrow \mathrm{CO(g)} + \frac{1}{2}\,\mathrm{O_2(g)}$$

The industrial synthesis of ammonia is described by the Haber-Bosch process. The following enthalpy diagram shows the energy associated with this thermochemical transformation at $298\,\text{K}$:

The figure outlines the enthalpy balance of the reaction:

The gaseous reactants are at a relative enthalpy of $0\,\text{kJ}$ and the products at $-92.0\,\text{kJ}$ per mole of $\mathrm{N_2(g)}$ that reacts.

1. a) Write the balanced thermochemical equation with the smallest integer stoichiometric coefficients and deduce, with justification, the sign and numerical value indicated by the arrow with the question mark.
2. b) Calculate the amount of heat exchanged at constant pressure when $0.50\,\text{mol}$ of $\mathrm{N_2(g)}$ react completely with excess hydrogen, and determine the heat required to decompose $68.0\,\text{g}$ of gaseous ammonia into its elements at the same temperature.

Data: Molar mass: $M(\mathrm{NH_3}) = 17.0\,\text{g}\times\text{mol}^{-1}$.

Calcium carbonate decomposes thermally according to the following thermochemical equation at $298\,\text{K}$ and $1\,\text{bar}$:

$$\mathrm{CaCO_3(s)} \rightarrow \mathrm{CaO(s)} + \mathrm{CO_2(g)} \quad \Delta H^\circ = +178.0\,\text{kJ}\times\text{mol}^{-1}$$

1. a) Justify whether the process is exothermic or endothermic and which physical quantity this enthalpy change represents if the reaction is carried out in an open container at constant pressure.
2. b) Calculate the heat exchanged, expressed in kilojoules, when $25.0\,\text{g}$ of pure calcium carbonate decompose completely.

Data: Atomic molar masses: $M(\mathrm{Ca}) = 40.1\,\text{g}\times\text{mol}^{-1}$, $M(\mathrm{C}) = 12.0\,\text{g}\times\text{mol}^{-1}$, $M(\mathrm{O}) = 16.0\,\text{g}\times\text{mol}^{-1}$.

Two fuels are available for home heating: methane, $\mathrm{CH_4(g)}$, and propane, $\mathrm{C_3H_8(g)}$. The standard thermochemical equations for complete combustion at $298\,\text{K}$ are:

$$\mathrm{CH_4(g)} + 2\,\mathrm{O_2(g)} \rightarrow \mathrm{CO_2(g)} + 2\,\mathrm{H_2O(l)} \quad \Delta H^\circ = -890.4\,\text{kJ}\times\text{mol}^{-1}$$ $$\mathrm{C_3H_8(g)} + 5\,\mathrm{O_2(g)} \rightarrow 3\,\mathrm{CO_2(g)} + 4\,\mathrm{H_2O(l)} \quad \Delta H^\circ = -2220.0\,\text{kJ}\times\text{mol}^{-1}$$

1. a) Calculate the energy released per gram of fuel consumed in each case.
2. b) State, justifying your answer with the results obtained, which of the two hydrocarbons provides the greater energy yield per unit of mass.

Data: Atomic molar masses: $M(\mathrm{C}) = 12.0\,\text{g}\times\text{mol}^{-1}$, $M(\mathrm{H}) = 1.0\,\text{g}\times\text{mol}^{-1}$.

In a constant-pressure calorimeter of $1\,\text{bar}$ (whose heat capacity is considered negligible), $100.0\,\text{mL}$ of an aqueous solution of $\mathrm{HCl} 0.50\,\text{M}$ are mixed with $100.0\,\text{mL}$ of an aqueous solution of $\mathrm{NaOH} 0.50\,\text{M}$, both initially at $21.0\,^\circ\text{C}$. After neutralization is complete, the final temperature of the homogeneous mixture stabilizes at $24.4\,^\circ\text{C}$.

1. a) Write the net ionic equation of the process and determine the amount of substance, in moles, of water formed.
2. b) Calculate the experimental molar enthalpy of neutralization, expressed in $\text{kJ}\times\text{mol}^{-1}$ of water produced.

Data: Assume that the density of all aqueous solutions is $1.00\,\text{g}\times\text{mL}^{-1}$ and that their mean specific heat capacity is equal to that of pure water, c_e $= 4.18\,\text{J}\times\text{g}^{-1}\times\text{K}^{-1}$. Assume there are no heat losses to the surroundings.

Gaseous hydrogen, $\mathrm{H_2(g)}$, and liquid hydrazine, $\mathrm{N_2H_4(l)}$, are compared as aerospace fuels. Under standard conditions at $298\,\text{K}$, the molar enthalpies of formation are $\Delta H_f^\circ[\mathrm{H_2O(l)}] = -285.8\,\text{kJ}\times\text{mol}^{-1}$ and $\Delta H_f^\circ[\mathrm{N_2H_4(l)}] = +50.6\,\text{kJ}\times\text{mol}^{-1}$.

The complete combustion reactions with oxygen gas form, respectively, liquid water in the case of hydrogen, and nitrogen gas together with liquid water in the case of hydrazine.

1. a) Write both balanced thermochemical equations per mole of fuel consumed and determine the standard enthalpy of combustion of each in $\text{kJ}\times\text{mol}^{-1}$.
2. b) If a rocket has a tank with a payload limit of $500\,\text{kg}$ of fuel, work out, with justification, which substance will supply the greater total energy and determine the volume of dry gaseous product that would remain after cooling and completely condensing the water (at $1\,\text{bar}$ and $298\,\text{K}$) for each fuel. Consider only the stoichiometric products: just enough oxygen, with no air or inert gas. This cold volume does not represent the hot exhaust of an engine.

Data: $M(\mathrm{H}) = 1.0\,\text{g}\times\text{mol}^{-1}$, $M(\mathrm{N}) = 14.0\,\text{g}\times\text{mol}^{-1}$, $R = 0.0831\,\text{bar}\times\text{L}\times\text{mol}^{-1}\times\text{K}^{-1}$.

In a lab experiment to determine the enthalpy of hydration and subsequent neutralization of calcium oxide, a student carries out two experiments in a constant-pressure calorimeter with negligible heat capacity:

1. $2.80\,\text{g}$ of $\mathrm{CaO(s)}$ are reacted in $200.0\,\text{g}$ of liquid water, producing the reaction $\mathrm{CaO(s)} + \mathrm{H_2O(l)} \rightarrow \mathrm{Ca(OH)_2(s)}$, and the water is observed to heat up, absorbing $3.26\,\text{kJ}$.

2. $3.70\,\text{g}$ of $\mathrm{Ca(OH)_2(s)}$ are reacted with an excess of aqueous solution of $\mathrm{HCl}$, releasing $5.84\,\text{kJ}$ according to: $\mathrm{Ca(OH)_2(s)} + 2\,\mathrm{HCl(aq)} \rightarrow \mathrm{CaCl_2(aq)} + 2\,\mathrm{H_2O(l)}$.

1. a) Calculate the standard molar enthalpy change for each of the two reactions above.
2. b) Applying Hess's law, determine the standard enthalpy for the direct reaction:

$$\mathrm{CaO(s)} + 2\,\mathrm{HCl(aq)} \rightarrow \mathrm{CaCl_2(aq)} + \mathrm{H_2O(l)}$$

Data: Atomic molar masses: $M(\mathrm{Ca}) = 40.1\,\text{g}\times\text{mol}^{-1}$, $M(\mathrm{O}) = 16.0\,\text{g}\times\text{mol}^{-1}$, $M(\mathrm{H}) = 1.0\,\text{g}\times\text{mol}^{-1}$. Consider the heats corrected to 298 K and standard pressure; neglect dilution effects and heat losses.