Complete and balance each combustion reaction equation: $$\mathrm {a} .\mathrm{C}_{4} \mathrm{H}_{6}(g)+\mathrm{O}_{2}(g) \longrightarrow \quad \mathrm{b} . \mathrm{C}(s)+\mathrm{O}_{2}(g) \longrightarrow$$ $$ \mathrm{c} . \mathrm{CS}_{2}(s)+\mathrm{O}_{2}(g) \longrightarrow \quad \text { d. } \mathrm{C}_{3} \mathrm{H}_{8} \mathrm{O}(l)+\mathrm{O}_{2}(g) \longrightarrow $$

Balancing chemical equations is crucial because it reflects the law of conservation of mass, stating that matter cannot be created or destroyed. When we write a combustion reaction equation, each element must have the same number of atoms on both the reactant and product sides. For example, balancing the combustion of \text{C}_4\text{H}_6 involves making sure the carbons (Cs) and hydrogens (Hs) from the hydrocarbon end up as carbon dioxide (CO\(_2\)) and water (H\(_2\)O) molecules.

Each step in the balancing process addresses a specific element, and we use coefficients to equalize atom counts for the elements involved. In combustion reactions, it's common to finalize the balancing process by adjusting the dioxygen (O\(_2\)) molecules, which can sometimes result in fractional coefficients, signaling that a full molecule of O\(_2\) is not wholly consumed during the reaction.

Understanding Reaction Proportions

Stoichiometry is the area of chemistry that pertains to the quantitative relationships between the substances involved in chemical reactions. It involves the calculation of reactants and products in chemical equations and is fundamental in determining the amounts of substances consumed and generated. For instance, the stoichiometry of the reaction \text{C}_4\text{H}_6(g) + \frac{11}{2} \text{O}_2(g) \rightarrow 4 \text{CO}_2(g) + 3 \text{H}_2\text{O}(g) tells us that for every molecule of \text{C}_4\text{H}_6 combusted, 4 molecules of CO\(_2\) and 3 molecules of H\(_2\)O are produced, requiring 5.5 molecules of O\(_2\).

Effective stoichiometry depends upon a balanced chemical equation, which allows the accurate prediction of product yields from given quantities of reactants. This is essential not only for laboratory chemical reactions but also for industrial chemical processes where efficiency and cost-effectiveness are key.

Outcome of Combustion Reactions

The products of combustion reactions typically include carbon dioxide (CO\(_2\)) and water (H\(_2\)O) when a hydrocarbon burns in the presence of oxygen, as observed in the equations provided. If the burning compound contains other elements, such as sulfur in carbon disulfide (CS\(_2\)), additional products like sulfur dioxide (SO\(_2\)) can be formed.

With complete combustion, ample oxygen leads to the formation of CO\(_2\), but with insufficient oxygen, carbon monoxide (CO) or even elemental carbon, in the form of soot, may result. Understanding the potential products of a combustion reaction is vital in predicting the environmental impact of burning different substances, such as the generation of pollutants like SO\(_2\), which contributes to acid rain.

Chemistry Behind Combustion

Oxidation reactions are chemical processes where an atom or compound loses electrons, often involving the gain of oxygen or the loss of hydrogen. Combustion is a specific type of rapid oxidation reaction that releases energy, typically in the form of heat and light. The oxidation here occurs when oxygen from the air reacts with the other elements in the compound being burned.

In the context of our reactions, the carbon (C) and hydrogen (H) atoms in the hydrocarbons and organics like alcohol (\text{C}_3\text{H}_8\text{O}) undergo oxidation to form CO\(_2\) and H\(_2\)O. These reactions not only sustain fire but also are the principle behind internal combustion engines and other technologies that leverage chemical energy.