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Chapter 8: Problem 46

Use the appropriate table to calculate \(\Delta H^{\circ}\) for (a) the reaction between copper(II) oxide and carbon monoxide to give copper metal and carbon dioxide. (b) the decomposition of one mole of methyl alcohol (CH \(_{3} \mathrm{OH}\) ) to methane and oxygen gases.

Short Answer

Expert verified
Question: Calculate the standard enthalpy change (ΔH°) for the following reactions: (a) CuO (s) + CO (g) -> Cu (s) + CO2 (g) (b) 2 CH3OH (l) -> CH4 (g) + O2 (g) + H2O (l) Answer: (a) ΔH° = -125.7 kJ/mol (b) ΔH° = 116.8 kJ/mol

Step by step solution

01

Write the balanced chemical equation

The balanced chemical equation for this reaction is: CuO (s) + CO (g) -> Cu (s) + CO2 (g)
02

Find the standard enthalpies of formation

According to a table of standard enthalpies of formation, we have: ΔHf°(CuO) = -157.3 kJ/mol ΔHf°(CO) = -110.5 kJ/mol ΔHf°(Cu) = 0 kJ/mol (Standard enthalpy of formation for elements in their standard states is always zero) ΔHf°(CO2) = -393.5 kJ/mol
03

Calculate ΔH° for the reaction

Use the equation mentioned in the analysis: ΔH° = Σ ΔHf°(products) - Σ ΔHf°(reactants) ΔH° = [ΔHf°(Cu) + ΔHf°(CO2)] - [ΔHf°(CuO) + ΔHf°(CO)] ΔH° = [(0) + (-393.5)] - [(-157.3) + (-110.5)] ΔH° = -393.5 - (-267.8) ΔH° = -125.7 kJ/mol So, the standard enthalpy change for this reaction is ΔH° = -125.7 kJ/mol. (b) Decomposition of one mole of methyl alcohol to methane and oxygen gases:
04

Write the balanced chemical equation

The balanced chemical equation for this reaction is: 2 CH3OH (l) -> CH4 (g) + O2 (g) + H2O (l)
05

Find the standard enthalpies of formation

According to a table of standard enthalpies of formation, we have: ΔHf°(CH3OH) = -238.7 kJ/mol ΔHf°(CH4) = -74.8 kJ/mol ΔHf°(O2) = 0 kJ/mol (Standard enthalpy of formation for elements in their standard states is always zero) ΔHf°(H2O) = -285.8 kJ/mol
06

Calculate ΔH° for the reaction

Use the equation mentioned in the analysis: ΔH° = Σ ΔHf°(products) - Σ ΔHf°(reactants) ΔH° = [ΔHf°(CH4) + ΔHf°(O2) + ΔHf°(H2O)] - [2 × ΔHf°(CH3OH)] ΔH° = [(-74.8) + (0) + (-285.8)] - [2 × (-238.7)] ΔH° = -360.6 - (-477.4) ΔH° = 116.8 kJ/mol So, the standard enthalpy change for this reaction is ΔH° = 116.8 kJ/mol.

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Most popular questions from this chapter

On a hot day, you take a six-pack of soda on a picnic, cooling it with ice. Each empty (aluminum) can weighs \(12.5 \mathrm{~g}\) and contains \(12.0\) oz of soda. The specific heat of aluminum is \(0.902 \mathrm{~J} / \mathrm{g} \cdot{ }^{\circ} \mathrm{C}\); take that of soda to be \(4.10 \mathrm{~J} / \mathrm{g} \cdot{ }^{\circ} \mathrm{C}\) (a) How much heat must be absorbed from the six-pack to lower the temperature from \(25.0^{\circ}\) to \(5.0^{\circ} \mathrm{C}\) ? (b) How much ice must be melted to absorb this amount of heat? \(\left(\Delta H_{\mathrm{fus}}\right.\) of ise is given in Table 8.2.)

Determine whether the statements given below are true or false. Consider enthalpy \((H)\) (a) It is a state property. (b) \(q_{\text {reaction }}\) (at constant \(\left.P\right)=\Delta H=H_{\text {products }}-H_{\text {reactants }}\) (c) The magnitude of \(\Delta H\) is independent of the amount of reactant. (d) In an exothermic process, the enthalpy of the system remains unchanged.

In the late eighteenth century Priestley prepared ammonia by reacting \(\mathrm{HNO}_{3}(g)\) with hydrogen gas. The thermodynamic equation for the reaction is $$ \mathrm{HNO}_{3}(g)+4 \mathrm{H}_{2}(g) \longrightarrow \mathrm{NH}_{3}(g)+3 \mathrm{H}_{2} \mathrm{O}(g) \quad \Delta H=-637 \mathrm{~kJ} $$ (a) Calculate \(\Delta H\) when one mole of hydrogen gas reacts. (b) What is \(\Delta H\) when \(10.00 \mathrm{~g}\) of \(\mathrm{NH}_{3}(g)\) is made to react with an excess of steam to form \(\mathrm{HNO}_{3}\) and \(\mathrm{H}_{2}\) gases?

To produce silicon, used in semiconductors, from sand \(\left(\mathrm{SiO}_{2}\right)\), a reaction is used that can be broken down into three steps: $$ \begin{aligned} \mathrm{SiO}_{2}(s)+2 \mathrm{C}(s) \longrightarrow \mathrm{Si}(s)+2 \mathrm{CO}(g) & & \Delta H=689.9 \mathrm{~kJ} \\ \mathrm{Si}(s)+2 \mathrm{Cl}_{2}(g) \longrightarrow \mathrm{SiCl}_{4}(g) & & \Delta H=-657.0 \mathrm{~kJ} \\ \mathrm{SiCl}_{4}(g)+2 \mathrm{Mg}(s) \longrightarrow 2 \mathrm{MgCl}_{2}(s)+\mathrm{Si}(s) & & \Delta H=-625.6 \mathrm{~kJ} \end{aligned} $$ (a) Write the thermochemical equation for the overall reaction for the formation of silicon from silicon dioxide; \(\mathrm{CO}\) and \(\mathrm{MgCl}_{2}\) are byproducts. (b) What is \(\Delta H\) for the formation of one mole of silicon? (c) Is the overall reaction exothermic?

When one mole of calcium carbonate reacts with ammonia, solid calcium cyanamide, \(\mathrm{CaCN}_{2}\), and liquid water are formed. The reaction absorbs \(90.1 \mathrm{~kJ}\) of heat. (a) Write a balanced thermochemical equation for the reaction. (b) Using Table 8.3, calculate \(\Delta H_{\mathrm{f}}^{\circ}\) for calcium cyanamide.
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