Question

Balance the following equations and indicate whether they are combination, decomposition, or combustion reactions: $$ \begin{array}{l}{\text { (a) } \mathrm{C}_{3} \mathrm{H}_{6}(g)+\mathrm{O}_{2}(g) \longrightarrow \mathrm{CO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(g)} \\ {\text { (b) } \mathrm{NH}_{4} \mathrm{NO}_{3}(s) \longrightarrow \mathrm{N}_{2} \mathrm{O}(g)+\mathrm{H}_{2} \mathrm{O}(g)} \\ {\text { (c) } \mathrm{C}_{5} \mathrm{H}_{6} \mathrm{O}(l)+\mathrm{O}_{2}(g) \longrightarrow \mathrm{CO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(g)} \\ {\text { (d) } \mathrm{N}_{2}(g)+\mathrm{H}_{2}(g) \longrightarrow \mathrm{NH}_{3}(g)} \\\ {\text { (e) } \mathrm{K}_{2} \mathrm{O}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{KOH}(a q)}\end{array} $$

Short answer

The short answer is: (a) Balanced equation: \(C_3H_6(g) + \frac{9}{2}O_2(g) \longrightarrow 3CO_2(g) + 3H_2O(g)\) - Combustion reaction (b) Balanced equation: \(NH_4NO_3(s) \longrightarrow N_2O(g) + 2H_2O(g)\) - Decomposition reaction (c) Balanced equation: \(C_5H_6O(l) + \frac{15}{2}O_2(g) \longrightarrow 5CO_2(g) + 3H_2O(g)\) - Combustion reaction (d) Balanced equation: \(N_2(g) + 3H_2(g) \longrightarrow 2NH_3(g)\) - Combination reaction (e) Balanced equation: \(K_2O(s) + H_2O(l) \longrightarrow 2KOH(aq)\) - Combination reaction

Step-by-step solution

(a) Balancing reaction and Identifying reaction type

To balance the given reaction, first write down the number of atoms of the elements in the reactants and the products: $$\mathrm{C}_{3} \mathrm{H}_{6}(g)+\mathrm{O}_{2}(g) \longrightarrow \mathrm{CO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(g)$$ C: 3, H: 6, O: 2 (Reactants) and C: 1, H: 2, O: 3 (Products) To balance the reaction, let's start by balancing the carbon atoms by placing the coefficient 3 in front of CO2: $$\mathrm{C}_{3} \mathrm{H}_{6}(g)+\mathrm{O}_{2}(g) \longrightarrow 3 \mathrm{CO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(g)$$ Now, balance the hydrogen atoms by placing the coefficient 3 in front of H2O: $$\mathrm{C}_{3} \mathrm{H}_{6}(g)+\mathrm{O}_{2}(g) \longrightarrow 3 \mathrm{CO}_{2}(g)+3 \mathrm{H}_{2} \mathrm{O}(g)$$ Finally, balance the oxygen atoms by placing the coefficient \(\frac{9}{2}\) in front of O2: $$\mathrm{C}_{3} \mathrm{H}_{6}(g)+\frac{9}{2} \mathrm{O}_{2}(g) \longrightarrow 3 \mathrm{CO}_{2}(g)+3 \mathrm{H}_{2} \mathrm{O}(g)$$ Since the reaction involves a hydrocarbon (C3H6) reacting with oxygen (O2) to form carbon dioxide and water as products, this is a combustion reaction.

(b) Balancing reaction and Identifying reaction type

Balance the given reaction: $$\mathrm{NH}_{4} \mathrm{NO}_{3}(s) \longrightarrow \mathrm{N}_{2} \mathrm{O}(g)+\mathrm{H}_{2} \mathrm{O}(g)$$ Since the reaction involves a single reactant (NH4NO3) breaking down into two products (N2O and H2O), this is a decomposition reaction. The given reaction is already balanced.

(c) Balancing reaction and Identifying reaction type

To balance the given reaction: $$\mathrm{C}_{5} \mathrm{H}_{6} \mathrm{O}(l)+\mathrm{O}_{2}(g) \longrightarrow \mathrm{CO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(g)$$ Follow the same process as in the first reaction. The balanced equation is: $$\mathrm{C}_{5} \mathrm{H}_{6} \mathrm{O}(l)+\frac{15}{2} \mathrm{O}_{2}(g) \longrightarrow 5 \mathrm{CO}_{2}(g)+3 \mathrm{H}_{2} \mathrm{O}(g)$$ This is also a combustion reaction since it involves a hydrocarbon reacting with oxygen to form carbon dioxide and water.

(d) Balancing reaction and Identifying reaction type

Balance the given reaction: $$\mathrm{N}_{2}(g)+\mathrm{H}_{2}(g) \longrightarrow \mathrm{NH}_{3}(g)$$ Place the coefficient 2 in front of NH3, and the coefficient 3 in front of H2. The balanced equation is: $$\mathrm{N}_{2}(g)+3 \mathrm{H}_{2}(g) \longrightarrow 2 \mathrm{NH}_{3}(g)$$ Since the reaction involves two reactants combining to form a single product, this is a combination reaction.

(e) Balancing reaction and Identifying reaction type

Balance the given reaction: $$\mathrm{K}_{2} \mathrm{O}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{KOH}(a q)$$ Place the coefficient 2 in front of KOH. The balanced equation is: $$\mathrm{K}_{2} \mathrm{O}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow 2 \mathrm{KOH}(a q)$$ This is a combination reaction as two reactants combine to form a single product.

Key concepts

Combination Reaction

In chemistry, a combination reaction, also known as a synthesis reaction, occurs when two or more substances combine to form a single product. A typical combination reaction involves two or more reactants usually yielding one product. A classic example of this reaction type is the synthesis of water from hydrogen and oxygen gases.

A simple formula for a combination reaction is:
\[A + B \rightarrow AB\]
The reaction in step 4 of our exercise (\(\mathrm{N}_{2}(g)+3 \mathrm{H}_{2}(g) \longrightarrow 2 \mathrm{NH}_{3}(g)\)) is a perfect illustration of a combination reaction, where nitrogen (\(\mathrm{N}_{2}\)) and hydrogen (\(\mathrm{H}_{2}\)) gases combine to create ammonia (\(\mathrm{NH}_{3}\)). Likewise, in step 5, potassium oxide (\(\mathrm{K}_{2} \mathrm{O}\)) reacts with water (\(\mathrm{H}_{2} \mathrm{O}\)) to yield potassium hydroxide (\(\mathrm{KOH}\)), also exhibiting a combination reaction. When balancing combination reactions, it’s crucial to ensure that the number of atoms for each element is the same on both the reactant and product sides of the equation.

Decomposition Reaction

Opposite to combination reactions are decomposition reactions, where a single compound breaks down into two or more simpler substances. These reactions often require an input of energy in the form of heat, light, or electricity. They are crucial in various processes, including the breakdown of organic matter.

The general formula for a decomposition reaction is:
\[AB \rightarrow A + B\]
In the context of our exercise, step 2 demonstrates a decomposition reaction (\(\mathrm{NH}_{4} \mathrm{NO}_{3}(s) \longrightarrow \mathrm{N}_{2} \mathrm{O}(g)+\mathrm{H}_{2} \mathrm{O}(g)\)), where ammonium nitrate (\(\mathrm{NH}_{4} \mathrm{NO}_{3}\)) decomposes into nitrous oxide (\(\mathrm{N}_{2} \mathrm{O}\)) and water (\(\mathrm{H}_{2} \mathrm{O}\)). Unlike combination reactions, decomposition involves breaking bonds, which may be seen as ‘reversing’ a synthesis process.

Combustion Reaction

A combustion reaction is an exothermic reaction where a substance combines with oxygen, releasing energy in the form of light and heat. Combustion reactions are common in daily life and power the engines of cars and airplanes.

A combustion reaction's basic formula often appears as:
\[Fuel + O_2 \rightarrow CO_2 + H_2O\]
The steps 1 and 3 of our exercise are exemplary combustion reactions. In step 1, propene (\(\mathrm{C}_{3} \mathrm{H}_{6}\)) combusts in the presence of oxygen (\(\mathrm{O}_{2}\)) to form carbon dioxide (\(\mathrm{CO}_{2}\)) and water (\(\mathrm{H}_{2} \mathrm{O}\)). The combustion of benzyl alcohol (\(\mathrm{C}_{5} \mathrm{H}_{6} \mathrm{O}\)) in step 3 follows a similar pattern. Typically, the combustion of hydrocarbons like these is used to generate energy, and the complete combustion always produces carbon dioxide and water.

Chemical Reaction Types

There are many types of chemical reactions, each characterized by the rearrangement of atoms to transform reactants into products. The core types beyond combination, decomposition, and combustion include single replacement, double replacement, and acid-base reactions. Understanding the different reaction types is crucial for studying chemical processes, predicting outcomes of reactions, and balancing chemical equations.

It's important to recognize the reactants and products' states, such as solid (\(s\)), liquid (\(l\)), aqueous (\(aq\)), or gas (\(g\)), as they can provide clues about the reaction type. For example, aqueous solutions often imply a dissolution process or a reaction taking place in water. By identifying patterns and applying laws of conservation of mass and charge, we can systematically approach and solve chemical equations, ensuring that atoms are conserved from reactants to products in any chemical transformation.