Use information from the chapter to write chemical equations to represent each of the following: (a) reaction of cesium metal with chlorine gas (b) formation of sodium peroxide \(\left(\mathrm{Na}_{2} \mathrm{O}_{2}\right)\) (c) thermal decomposition of lithium carbonate (d) reduction of sodium sulfate to sodium sulfide (e) combustion of potassium to form potassium superoxide

The reaction between metals and chlorine gas is an example of a synthesis reaction, where two reactants combine to form a single product. In the case of cesium metal reacting with chlorine gas, cesium (\text{Cs}) donates an electron to chlorine (\text{Cl}_2), forming cesium chloride (\text{CsCl}). The balanced chemical equation is: \[\begin{equation} 2 \text{Cs} + \text{Cl}_2 \rightarrow 2 \text{CsCl} \end{equation}\].This reaction is highly exothermic, releasing a large amount of energy. This illustrates the general trend of alkali metals reacting vigorously with halogens to form ionic compounds.

Sodium peroxide (\text{Na}_2\text{O}_2) is formed when sodium (\text{Na}), an alkali metal, reacts with oxygen (\text{O}_2). The reaction is a form of oxidation where the sodium loses electrons and the oxygen gains them. The equation representing this redox reaction is:\[\begin{equation} 2 \text{Na} + \text{O}_2 \rightarrow \text{Na}_2\text{O}_2 \end{equation}\].This reaction can illuminate the broader concept of peroxide formation, which involves the combination of a metal with oxygen in higher oxidation states. Since sodium is in Group 1 of the periodic table, it easily loses one electron, leading to the peroxide anion where the oxygen is in a \text{-1} oxidation state.

Thermal decomposition reactions involve a compound breaking down into two or more products when heated. Lithium carbonate (\text{Li}_2\text{CO}_3) decomposes to lithium oxide (\text{LiO}) and carbon dioxide (\text{CO}_2) upon heating. The chemical equation is:\[\begin{equation}\text{Li}_2\text{CO}_3 \rightarrow 2\text{LiO} + \text{CO}_2 \end{equation}\].The decomposition of lithium carbonate is an endothermic reaction, meaning it requires heat to proceed. It highlights the stability of carbonates which, when heated, generally decompose to release \text{CO}_2 gas.

Reduction reactions entail the gain of electrons by a substance. In this case, sodium sulfate (\text{Na}_2\text{SO}_4) is reduced to sodium sulfide (\text{Na}_2\text{S}) using carbon as a reducing agent. This is shown in the balanced equation:\[\begin{equation}\text{Na}_2\text{SO}_4 + 4\text{C} \rightarrow \text{Na}_2\text{S} + 4\text{CO}_2 \end{equation}\].This exemplifies the classic redox reaction where carbon is oxidized to carbon dioxide (\text{CO}_2), and sulfate (\text{SO}_4^{2-}) is reduced to sulfide (\text{S}^{2-}). These types of redox reactions are fundamental in metallurgy and chemical industries.

The combustion of potassium (\text{K}) in the presence of oxygen leads to the formation of potassium superoxide (\text{KO}_2). The balanced chemical equation is:\[\begin{equation}2 \text{K} + \text{O}_2 \rightarrow 2 \text{KO}_2 \end{equation}\].Combustion reactions are a type of redox reaction where a substance combines with oxygen, often releasing light and heat as energy. In this particular reaction, the product, potassium superoxide, is a strong oxidizing agent and it is quite unique, as most alkali metals form normal oxides or peroxides upon combustion. This emphasizes the highly reactive nature of alkali metals and their tendency to form various types of oxides with oxygen.