Question

Identify each of the following unbalanced reaction equations as belonging to one or more of the following categories: precipitation, acid-base, or oxidation-reduction. a. \(\mathrm{H}_{2} \mathrm{O}_{2}(a q) \rightarrow \mathrm{H}_{2} \mathrm{O}(l)+\mathrm{O}_{2}(g)\) b. \(\mathrm{H}_{2} \mathrm{SO}_{4}(a q)+\mathrm{Zn}(s) \rightarrow \mathrm{ZnSO}_{4}(a q)+\mathrm{H}_{2}(g)\) c. \(\mathrm{H}_{2} \mathrm{SO}_{4}(a q)+\mathrm{NaOH}(a q) \rightarrow \mathrm{Na}_{2} \mathrm{SO}_{4}(a q)+\mathrm{H}_{2} \mathrm{O}(l)\) d. \(\mathrm{H}_{2} \mathrm{SO}_{4}(a q)+\mathrm{Ba}(\mathrm{OH})_{2}(a q) \rightarrow \mathrm{BaSO}_{4}(s)+\mathrm{H}_{2} \mathrm{O}(l)\) e. \(\mathrm{AgNO}_{3}(a q)+\mathrm{CuCl}_{2}(a q) \rightarrow \mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}(a q)+\mathrm{AgCl}(s)\) f. \(\mathrm{KOH}(a q)+\mathrm{CuSO}_{4}(a q) \rightarrow \mathrm{Cu}(\mathrm{OH})_{2}(s)+\mathrm{K}_{2} \mathrm{SO}_{4}(a q)\) g. \(\mathrm{Cl}_{2}(g)+\mathrm{F}_{2}(g) \rightarrow \mathrm{ClF}(g)\) h. \(\mathrm{NO}(g)+\mathrm{O}_{2}(g) \rightarrow \mathrm{NO}_{2}(g)\) i. \(\mathrm{Ca}(\mathrm{OH})_{2}(s)+\mathrm{HNO}_{3}(a q) \rightarrow \mathrm{Ca}\left(\mathrm{NO}_{3}\right)_{2}(a q)+\mathrm{H}_{2} \mathrm{O}(l)\)

Short answer

The classifications for each reaction are as follows: a. Oxidation-reduction b. Oxidation-reduction c. Acid-base d. Acid-base and precipitation e. Oxidation-reduction and precipitation f. Acid-base and precipitation g. Oxidation-reduction h. Oxidation-reduction i. Acid-base

Step-by-step solution

a. $\mathrm{H}_{2} \mathrm{O}_{2}(a q) \rightarrow \mathrm{H}_{2} \mathrm{O}(l)+\mathrm{O}_{2}(g)$

This reaction involves the decomposition of hydrogen peroxide where water and oxygen gas are formed. Since there is only one reactant involved and the oxidation numbers change (oxygen's oxidation number goes from -1 in hydrogen peroxide to -2 in water and 0 in oxygen gas), this is an oxidation-reduction (Redox) reaction.

b. $\mathrm{H}_{2} \mathrm{SO}_{4}(a q)+\mathrm{Zn}(s) \rightarrow \mathrm{ZnSO}_{4}(a q)+\mathrm{H}_{2}(g)$

In this reaction, zinc metal reacts with sulfuric acid to form zinc sulfate (a salt) and hydrogen gas. The oxidation number of zinc changes from 0 to +2, and hydrogen changes from +1 to 0, meaning this is an oxidation-reduction (Redox) reaction.

c. $\mathrm{H}_{2} \mathrm{SO}_{4}(a q)+\mathrm{NaOH}(a q) \rightarrow \mathrm{Na}_{2} \mathrm{SO}_{4}(a q)+\mathrm{H}_{2} \mathrm{O}(l)$

This reaction involves the reaction between sulfuric acid and sodium hydroxide to form a salt (sodium sulfate) and water. As an acid and a base are reacting to form a salt and water, this is an acid-base reaction.

d. $\mathrm{H}_{2} \mathrm{SO}_{4}(a q)+\mathrm{Ba}(\mathrm{OH})_{2}(a q) \rightarrow \mathrm{BaSO}_{4}(s)+\mathrm{H}_{2} \mathrm{O}(l)$

This reaction involves the reaction between sulfuric acid and barium hydroxide, producing barium sulfate (a solid precipitate) and water. This reaction is both an acid-base reaction (since an acid and base react to form water) and a precipitation reaction (since a solid precipitate is formed).

e. $\mathrm{AgNO}_{3}(a q)+\mathrm{CuCl}_{2}(a q) \rightarrow \mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}(a q)+\mathrm{AgCl}(s)$

In this reaction, silver nitrate reacts with copper chloride to form copper nitrate and silver chloride (a solid precipitate). The oxidation numbers of silver and copper change, so this reaction is both an oxidation-reduction (Redox) reaction and a precipitation reaction due to the formation of silver chloride precipitate.

f. $\mathrm{KOH}(a q)+\mathrm{CuSO}_{4}(a q) \rightarrow \mathrm{Cu}(\mathrm{OH})_{2}(s)+\mathrm{K}_{2} \mathrm{SO}_{4}(a q)$

In this reaction, potassium hydroxide reacts with copper sulfate to form copper hydroxide (a solid precipitate) and potassium sulfate. Since a solid precipitate is formed and a reaction occurs between an acid and a base (copper sulfate is considered acidic due to the sulfate ion and potassium hydroxide is a base), this reaction is both a precipitation reaction and an acid-base reaction.

g. \(\mathrm{Cl}_{2}(g)+\mathrm{F}_{2}(g) \rightarrow \mathrm{ClF}(g)\)

This reaction involves the combination of chlorine gas and fluorine gas to form chlorine monofluoride. The oxidation numbers of chlorine and fluorine change during this reaction, making it an oxidation-reduction (Redox) reaction.

h. \(\mathrm{NO}(g)+\mathrm{O}_{2}(g) \rightarrow \mathrm{NO}_{2}(g)\)

In this reaction, nitrogen monoxide reacts with oxygen gas to form nitrogen dioxide. The oxidation numbers of nitrogen and oxygen change during this reaction, making it an oxidation-reduction (Redox) reaction.

i. $\mathrm{Ca}(\mathrm{OH})_{2}(s)+\mathrm{HNO}_{3}(a q) \rightarrow \mathrm{Ca}\left(\mathrm{NO}_{3}\right)_{2}(a q)+\mathrm{H}_{2} \mathrm{O}(l)$

This reaction involves the reaction between calcium hydroxide and nitric acid, producing calcium nitrate (a salt) and water. Since an acid and a base are reacting to form a salt and water, this is an acid-base reaction.

Key concepts

Precipitation Reactions

Precipitation reactions occur when two soluble salts react in solution to form one or more insoluble products, known as precipitates. For example, when solutions of silver nitrate (\text{AgNO}_3) and copper chloride (\text{CuCl}_2) are mixed, the insoluble silver chloride (\text{AgCl}) precipitates out of the solution.

The chemical equation for this reaction is: \[ \text{AgNO}_3(aq) + \text{CuCl}_2(aq) \rightarrow \text{Cu}(\text{NO}_3)_2(aq) + \text{AgCl}(s) \] To identify a precipitation reaction, look for the formation of a solid product from the reaction of two aqueous solutions. These reactions are essential in various fields, including analytical chemistry for performing qualitative analysis of ions.

Acid-Base Reactions

Acid-base reactions, also known as neutralization reactions, involve the reaction of an acid with a base to produce a salt and water. A classic example is the reaction of sulfuric acid (\text{H}_2\text{SO}_4) with sodium hydroxide (\text{NaOH}), which yields sodium sulfate (\text{Na}_2\text{SO}_4) and water (H_2O).

The balanced chemical equation is: \[ \text{H}_2\text{SO}_4(aq) + 2\text{NaOH}(aq) \rightarrow \text{Na}_2\text{SO}_4(aq) + 2\text{H}_2\text{O}(l) \] These reactions are crucial in chemistry as they are used in titrations to determine the concentration of an unknown acid or base solution, and they also play a key role in biological systems and the environment.

Oxidation-Reduction Reactions

Oxidation-reduction (redox) reactions involve the transfer of electrons between reactants, leading to a change in oxidation states. An example is the decomposition of hydrogen peroxide (\text{H}_2\text{O}_2) into water and oxygen.

The equation is: \[ 2\text{H}_2\text{O}_2(aq) \rightarrow 2\text{H}_2\text{O}(l) + \text{O}_2(g) \] Redox reactions are central to many processes such as combustion, metabolism, corrosion, and industrial chemical synthesis. They are characterized by the oxidation number changes seen in the elements involved in the reactions.

Chemical Equation Balancing

Balancing chemical equations is a fundamental skill in chemistry that ensures the principle of conservation of mass is obeyed in every chemical reaction. It involves making sure that there is an equal number of each type of atom on both the reactant and product sides of a chemical equation. This is crucial as it reflects the stoichiometry of the reaction, which is necessary to calculate reactant and product quantities.

For instance, in the oxidation of zinc by sulfuric acid: \[ \text{H}_2\text{SO}_4(aq) + \text{Zn}(s) \rightarrow \text{ZnSO}_4(aq) + \text{H}_2(g) \] This equation is already balanced with one sulfate (\text{SO}_4^{2-}), one zinc atom, two hydrogen atoms, and four oxygen atoms on both sides. Balancing chemical equations often requires practice to identify the correct coefficients that balance the equation.