Problem 71 - Chemistry and Chemical Reactivity

Question

Complete and balance the following equations. (a) \(\mathrm{KClO}_{3}+\) heat \(\longrightarrow\) (b) \(\mathrm{H}_{2} \mathrm{S}(\mathrm{g})+\mathrm{O}_{2}(\mathrm{g}) \longrightarrow\) (c) \(\mathrm{Na}(\mathrm{s})+\mathrm{O}_{2}(\mathrm{g}) \longrightarrow\) (d) \(\mathbf{P}_{4}(s)+\mathbf{K O H}(\mathbf{a q})+\mathbf{H}_{2} \mathbf{O}(\ell) \longrightarrow\) (e) \(\mathrm{NH}_{4} \mathrm{NO}_{3}(\mathrm{s})+\) heat \(\longrightarrow\) (f) \(\operatorname{In}(\mathrm{s})+\mathrm{Br}_{2}(\ell) \longrightarrow\) \((g) \operatorname{Sn} \mathrm{Cl}_{4}(\ell)+\mathrm{H}_{2} \mathrm{O}(\ell) \longrightarrow\)

Step-by-Step Solution

Answer

1) 2 KClO\(_3\) → 2 KCl + 3 O\(_2\); 2) 2 H\(_2\)S + 3 O\(_2\) → 2 SO\(_2\) + 2 H\(_2\)O; 3) 4 Na + O\(_2\) → 2 Na\(_2\)O; 4) P\(_4\) + 8 KOH + 8 H\(_2\)O → 4 KH\(_2\)PO\(_4\) + 4 H\(_2\); 5) NH\(_4\)NO\(_3\) → N\(_2\)O + 2 H\(_2\)O; 6) 2 In + 3 Br\(_2\) → 2 InBr\(_3\); 7) SnCl\(_4\) + 4 H\(_2\)O → 4 HCl + SnO\(_2\).

Step 1: Decompose KClO3

The chemical compound potassium chlorate (KClO\(_3\)) decomposes on heating to produce potassium chloride (KCl) and oxygen gas (O\(_2\)). The balanced equation is: \[ 2 \mathrm{KClO}_3 \rightarrow 2 \mathrm{KCl} + 3 \mathrm{O}_2 \]

Step 2: Balance H2S and O2

The reaction between hydrogen sulfide (H\(_2\)S) and oxygen (O\(_2\)) to form sulfur dioxide (SO\(_2\)) and water (H\(_2\)O) can be balanced as:\[ 2 \mathrm{H}_2\mathrm{S} + 3 \mathrm{O}_2 \rightarrow 2 \mathrm{SO}_2 + 2 \mathrm{H}_2\mathrm{O} \]

Step 3: Balance Na and O2

When sodium (Na) reacts with oxygen (O\(_2\)), they form sodium oxide (Na\(_2\)O). The balanced equation is:\[ 4 \mathrm{Na} + \mathrm{O}_2 \rightarrow 2 \mathrm{Na}_2\mathrm{O} \]

Step 4: Balance P4 with KOH and H2O

Phosphorus (P\(_4\)) reacts with potassium hydroxide (KOH) and water (H\(_2\)O) to produce potassium dihydrogen phosphate (KH\(_2\)PO\(_4\)) and hydrogen gas (H\(_2\)). The balanced equation is:\[ \mathrm{P}_4 + 8 \mathrm{KOH} + 8 \mathrm{H}_2\mathrm{O} \rightarrow 4 \mathrm{KH}_2\mathrm{PO}_4 + 4 \mathrm{H}_2 \]

Step 5: Decompose NH4NO3

Ammonium nitrate (NH\(_4\)NO\(_3\)) decomposes when heated to produce dinitrogen monoxide (N\(_2\)O, also known as nitrous oxide) and water (H\(_2\)O). The balanced equation is:\[ \mathrm{NH}_4\mathrm{NO}_3 \rightarrow \mathrm{N}_2\mathrm{O} + 2 \mathrm{H}_2\mathrm{O} \]

Step 6: Balance In and Br2

Indium (In) reacts with bromine (Br\(_2\)) to form indium(III) bromide (InBr\(_3\)). The balanced equation is:\[ 2 \mathrm{In} + 3 \mathrm{Br}_2 \rightarrow 2 \mathrm{InBr}_3 \]

Step 7: Balance SnCl4 with H2O

Tin(IV) chloride (SnCl\(_4\)) reacts with water (H\(_2\)O) to form hydrochloric acid (HCl) and tin(IV) oxide (SnO\(_4\)). The balanced equation is:\[ \mathrm{SnCl}_4 + 4 \mathrm{H}_2\mathrm{O} \rightarrow 4 \mathrm{HCl} + \mathrm{SnO}_2 \]

Key Concepts

Balancing Chemical Equations

Balancing chemical equations is an essential skill in chemistry that ensures the law of conservation of mass is obeyed. This law states that mass cannot be created or destroyed in a chemical reaction. Therefore, the number of atoms for each element must be the same on both sides of the equation. To balance an equation:

• Start by writing the unbalanced equation, showing the reactants and products.
• List each element and tally the number of atoms on each side.
• Use coefficients to adjust the balance without changing the actual substances involved.
• Continuously check your work to ensure each element matches on both sides.

Practice and familiarity help to make balancing equations more intuitive. For example, in the reaction where sodium (\(\mathrm{Na}\)) reacts with oxygen (\(\mathrm{O}_2\)) to form sodium oxide, if we start with \(4 \mathrm{Na} + \mathrm{O}_2 \rightarrow 2 \mathrm{Na}_2\mathrm{O}\), it shows that we successfully equalized four sodium atoms and two oxygen atoms on both sides.

Decomposition Reactions

Decomposition reactions are processes where a single compound breaks down into two or more simpler substances. This type of reaction often involves the application of heat, light, or electricity. Such reactions can be identified by the general form:\(\mathrm{AB} \rightarrow \mathrm{A} + \mathrm{B}\).
For instance, when heating potassium chlorate (\(\mathrm{KClO}_3\)) in Step 1 of our example, it decomposes into potassium chloride (\(\mathrm{KCl}\)) and oxygen gas (\(\mathrm{O}_2\)).

• Thermal decomposition requires energy in the form of heat, leading to the breakdown of compounds like \(\mathrm{KClO}_3\).
• Electrolytic decomposition uses electricity to break down substances, such as the decomposition of water into hydrogen and oxygen gases.
• Photodecomposition requires light; for example, silver chloride decomposes in presence of light.

These reactions are significant in various industrial and laboratory processes.

Redox Reactions

Redox reactions, short for reduction-oxidation reactions, involve the transfer of electrons between two species. These reactions are characterized by one element being oxidized (losing electrons) and another being reduced (gaining electrons). Understanding redox reactions is crucial in fields like battery technology, corrosion, and metabolism. In our exercise:

• The reaction of indium (\(\mathrm{In}\)) with bromine (\(\mathrm{Br}_2\)) to form indium(III) bromide (\(\mathrm{InBr}_3\)) exemplifies a redox reaction.
• Oxidation occurs when indium loses electrons to become indium ions.
• Reduction happens when bromine molecules gain those electrons to form bromide ions.

Redox reactions can be identified by changes in oxidation states. Analyzing these changes helps in balancing complex equations and predicting reaction products.

Combustion Reactions

Combustion reactions are chemical processes that occur when a substance reacts rapidly with oxygen to release energy in the form of light and heat. These reactions are generally exothermic and produce oxidized products. In many cases, hydrocarbons (compounds composed of hydrogen and carbon) are the combustible material. The general equation for a combustion reaction of a hydrocarbon in the presence of oxygen is:\(\mathrm{C}_x\mathrm{H}_y + \mathrm{O}_2 \rightarrow \mathrm{CO}_2 + \mathrm{H}_2\mathrm{O}\).
Within our exercise context, the reaction where hydrogen sulfide (\(\mathrm{H}_2\mathrm{S}\)) reacts with oxygen yields sulfur dioxide and water, showcasing the characteristics of a combustion reaction:

• It transforms potential chemical energy into heat and light energy.
• Complete combustion occurs with a sufficient oxygen supply, producing carbon dioxide and water.
• Incomplete combustion may result in carbon monoxide if oxygen is limited.

Combustion reactions are foundational in industries for producing energy, such as in engines and power generation plants.