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Chapter 21: Problem 45

Methane and sulfur vapor react to form carbon disulfide and hydrogen sulfide. Carbon disulfide reacts with \(\mathrm{Cl}_{2}(\mathrm{g})\) to form carbon tetrachloride and \(\mathrm{S}_{2} \mathrm{Cl}_{2}\) Further reaction of carbon disulfide and \(\overline{S_{2} C l_{2}}\) produces additional carbon tetrachloride and sulfur. Write a series of equations for the reactions described here.

Short Answer

Expert verified
The balanced reactions are: \( CH4(g) + 2S(g) \rightarrow CS2(g) + 2H2S(g) \) , \( CS2(g) + 3Cl2(g) \rightarrow CCl4(g) +S2Cl2(g) \) , \( CS2(g) + S2Cl2(g) \rightarrow CCl4(g) + 2S(g) \)

Step by step solution

01

Formulate the first reaction

The first reaction is between Methane (CH4) and sulfur (S) to form Carbon disulfide (CS2) and Hydrogen Sulfide (H2S). This is represented by the equation: \(CH4(g) + 2S(g) \rightarrow CS2(g) + 2H2S(g) \)
02

Formulate the second reaction

The second reaction is between Carbon disulfide (CS2) and Chlorine gas (Cl2) to form Carbon tetrachloride (CCl4) and Disulfur dichloride (S2Cl2). This is represented by the equation: \(CS2(g) + 3Cl2(g) \rightarrow CCl4(g) +S2Cl2(g) \)
03

Formulate the third reaction

The third reaction is between Carbon disulfide (CS2) and Disulfur dichloride (S2Cl2) to form Carbon tetrachloride (CCl4) and sulfur (S). This is represented by the equation: \(CS2(g) + S2Cl2(g) \rightarrow CCl4(g) + 2S(g) \)

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Most popular questions from this chapter

When a \(0.200 \mathrm{g}\) sample of \(\mathrm{Mg}\) is heated in air, \(0.315 \mathrm{g}\) of product is obtained. Assume that all the Mg appears in the product. (a) If the product were pure \(\mathrm{MgO}\), what mass should have been obtained? (b) Show that the 0.315 g product could be a mixture of \(\mathrm{Mg} \mathrm{O}\) and \(\mathrm{Mg}_{3} \mathrm{N}_{2}.\) (c) What is the mass percent of \(\mathrm{MgO}\) in the \(\mathrm{MgO}-\mathrm{Mg}_{3} \mathrm{N}_{2}\) mixed product?

The Gibbs energies of formation, \(\Delta G_{f}^{Q}\), for \(\mathrm{KO}_{2}(\mathrm{s})\) and \(\mathrm{K}_{2} \mathrm{O}(\mathrm{s})\) are \(-240.59 \mathrm{kJmol}^{-1}\) and \(-322.09 \mathrm{kJmol}^{-1}\) respectively, at \(298 \mathrm{K}\). Calculate the equilibrium constant for the reaction below at \(298 \mathrm{K}\). Is \(\mathrm{KO}_{2}(\mathrm{s})\) thermodynamically stable with respect to \(\mathrm{K}_{2} \mathrm{O}(\mathrm{s})\) and \(\mathrm{O}_{2}(\mathrm{g})\) at \(298 \mathrm{K} ?\) $$ 2 \mathrm{KO}_{2}(\mathrm{s}) \longrightarrow \mathrm{K}_{2} \mathrm{O}(\mathrm{s})+\frac{3}{2} \mathrm{O}_{2}(\mathrm{g}) $$

Write a chemical equation to represent (a) the reaction of potassium cyanide solution with silver nitrate solution; (b) the combustion of \(\mathrm{Si}_{3} \mathrm{H}_{8}\) in an excess of oxygen; (c) the reaction of dinitrogen with calcium carbide to give calcium cyanamide (CaNCN).

The electrolysis of \(0.250 \mathrm{L}\) of \(0.220 \mathrm{M} \mathrm{MgCl}_{2}\) is conducted until \(104 \mathrm{mL}\) of gas (a mixture of \(\mathrm{H}_{2}\) and water vapor) is collected at \(23^{\circ} \mathrm{C}\) and \(748 \mathrm{mmHg} .\) Will \(\mathrm{Mg}(\mathrm{OH})_{2}(\mathrm{s})\) precipitate if electrolysis is carried to this point? (Use 21 mmHg as the vapor pressure of the solution.)

Lithium superoxide, \(\mathrm{LiO}_{2}(\mathrm{s}),\) has never been isolated. Use ideas from Chapter \(12,\) together with data from this chapter and Appendix \(D\), to estimate \(\Delta H_{f}\) for \(\mathrm{LiO}_{2}(\mathrm{s})\) and assess whether \(\mathrm{LiO}_{2}(\mathrm{s})\) is thermodynamically stable with respect to \(\mathrm{Li}_{2} \mathrm{O}(\mathrm{s})\) and \(\mathrm{O}_{2}(\mathrm{g}).\) (a) Use the Kapustinskii equation, along with appropriate data below, to estimate the lattice energy, \(U,\) for \(\left.\mathrm{LiO}_{2}(\mathrm{s}) . \text { (See exercise } 126 \text { in Chapter } 12 .\right)\) The ionic radii for \(L\) i \(^{+}\) and \(O_{2}^{-}\) are \(73 \mathrm{pm}\) and \(144 \mathrm{pm},\) respectively. (b) Use your result from part (a) in the BornFajans-Haber cycle to estimate \(\Delta H_{\mathrm{f}}^{2}\) for \(\mathrm{LiO}_{2}(\mathrm{s})\) [Hint: For the process \(\mathrm{O}_{2}(\mathrm{g})+\mathrm{e}^{-} \rightarrow \mathrm{O}_{2}^{-}(\mathrm{g}), \Delta H^{\circ}=.\) \(-43 \mathrm{kJ} \mathrm{mol}^{-1} .\) See Table 21.2 and Appendix \(\mathrm{D}\) for the other data that are required.] (c) Use your result from part (b) to calculate the enthalpy change for the decomposition of \(\mathrm{LiO}_{2}(\mathrm{s})\) to \(\mathrm{Li}_{2} \mathrm{O}(\mathrm{s})\) and \(\mathrm{O}_{2}(\mathrm{g}) .\) For \(\mathrm{Li}_{2} \mathrm{O}(\mathrm{s}), \Delta H_{\mathrm{f}}^{\circ}=-598.73\) \(\mathrm{kJmol}^{-1}.\) (d) Use your result from part (c) to decide whether \(\mathrm{LiO}_{2}(\mathrm{s})\) is thermodynamically stable with respect to \(\mathrm{Li}_{2} \mathrm{O}(\mathrm{s})\) and \(\mathrm{O}_{2}(\mathrm{g}) .\) Assume that entropy effects can be neglected.
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