Question

Sulfuric acid, \(\mathrm{H}_{2} \mathrm{SO}_{4}\), oxidizes many metallic elements. One of the effects of acid rain is that it produces sulfuric acid in the atmosphere, which then reacts with metals used in construction. Write balanced oxidation-reduction equations for the reaction of sulfuric acid with Fe, \(\mathrm{Zn}, \mathrm{Mg}, \mathrm{Co},\) and \(\mathrm{Ni}\).

Short answer

The balanced oxidation-reduction equations for the reaction of sulfuric acid with the given metals are: 1. Fe: \(H_2SO_4 + Fe → FeSO_4 + H_2\) 2. Zn: \(H_2SO_4 + Zn → ZnSO_4 + H_2\) 3. Mg: \(H_2SO_4 + Mg → MgSO_4 + H_2\) 4. Co: \(H_2SO_4 + Co → CoSO_4 + H_2\) 5. Ni: \(H_2SO_4 + Ni → NiSO_4 + H_2\) In each reaction, the metal changes its oxidation state from 0 to +2, while hydrogen changes its oxidation state from +1 to 0, resulting in the formation of a metal sulfate and hydrogen gas.

Step-by-step solution

Identify the Oxidation States

Before starting with the individual reactions, let's identify the oxidation states of sulfuric acid (H2SO4) and the five metals. In sulfuric acid, hydrogen has a +1 oxidation state, sulfur has a +6 oxidation state, and oxygen has a -2 oxidation state. The oxidation states of the five metals are as follows: Fe (0), Zn (0), Mg (0), Co(0), and Ni(0), as they are in their elemental form.

Sulfuric Acid and Iron (Fe)

First, let's write the reaction of sulfuric acid with iron: H2SO4 + Fe → FeSO4 + H2 Now, let's balance the equation. In this reaction, the oxidation state of iron changes from 0 to +2, while the oxidation state of hydrogen changes from +1 to 0. The balanced equation is: 1( H2SO4) + 1(Fe) → 1(FeSO4) + 1(H2)

Sulfuric Acid and Zinc (Zn)

Now, let's write the reaction of sulfuric acid with zinc: H2SO4 + Zn → ZnSO4 + H2 In this reaction, the oxidation state of zinc changes from 0 to +2, while the oxidation state of hydrogen changes from +1 to 0. The balanced equation is: 1( H2SO4) + 1(Zn) → 1(ZnSO4) + 1(H2)

Sulfuric Acid and Magnesium (Mg)

Now, let's write the reaction of sulfuric acid with magnesium: H2SO4 + Mg → MgSO4 + H2 In this reaction, the oxidation state of magnesium changes from 0 to +2, and the oxidation state of hydrogen changes from +1 to 0. The balanced equation is: 1( H2SO4) + 1(Mg) → 1(MgSO4) + 1(H2)

Sulfuric Acid and Cobalt (Co)

Now, let's write the reaction of sulfuric acid with cobalt: H2SO4 + Co → CoSO4 + H2 In this reaction, the oxidation state of cobalt changes from 0 to +2, and the oxidation state of hydrogen changes from +1 to 0. The balanced equation is: 1( H2SO4) + 1(Co) → 1(CoSO4) + 1(H2)

Sulfuric Acid and Nickel (Ni)

Finally, let's write the reaction of sulfuric acid with nickel: H2SO4 + Ni → NiSO4 + H2 In this reaction, the oxidation state of nickel changes from 0 to +2, and the oxidation state of hydrogen changes from +1 to 0. The balanced equation is: 1( H2SO4) + 1(Ni) → 1(NiSO4) + 1(H2) In all the reactions discussed above, sulfuric acid reacts with each metal to form a metal sulfate and hydrogen gas. These balanced equations illustrate the process of oxidation and reduction in the context of reacting sulfuric acid with different metals.

Key concepts

Sulfuric Acid Reactivity

Sulfuric acid, often denoted as \( \mathrm{H}_2\mathrm{SO}_4 \), is a highly reactive, strong acid that is commonly used industrially and is also a component of acid rain. It has the ability to oxidize metals, which means it can accept electrons from metal atoms during chemical reactions.

Sulfuric acid reacts with metals such as iron (Fe), zinc (Zn), magnesium (Mg), cobalt (Co), and nickel (Ni) to produce metal sulfates. In these reactions, the metals lose electrons (are oxidized) and the hydrogen ions in the acid gain electrons (are reduced), forming hydrogen gas (H2). The sulfate ions (\( \mathrm{SO}_4^{2-} \) remain unchanged. This behavior exemplifies the dual nature of \( \mathrm{H}_2\mathrm{SO}_4 \) reacting as an oxidizing agent and an acid, emphasizing its versatile reactivity in different chemical contexts.

Balancing Chemical Equations

Balancing chemical equations is a fundamental skill in chemistry that ensures the law of conservation of mass is fulfilled; the quantity of each element must remain the same before and after a reaction. In the reactions of \( \mathrm{H}_2\mathrm{SO}_4 \) with various metals, each balanced equation must have the same number of atoms of each element on both sides.

The process involves adjusting coefficients to get equal numbers of each type of atom on both the reactant and product sides. As shown in the exercise, for each metal reaction with sulfuric acid, coefficients are assigned so that the number of iron, zinc, magnesium, cobalt, and nickel atoms and the number of sulfur, hydrogen, and oxygen atoms balance. This fundamental process guarantees the equations accurately depict the physical reality of the reactions occurring.

Metal Oxidation

Metal oxidation is the loss of electrons by a metal atom, typically leading to the formation of an oxide or, in the case of reactions with sulfuric acid, a metal sulfate. In each reaction between \( \mathrm{H}_2\mathrm{SO}_4 \) and a metal, the metal atom starts with an oxidation state of zero, as it is in its elemental form. After reacting with sulfuric acid, the metal's oxidation state increases (usually becoming positive), signifying that it has been oxidized.

Examples of Metal Oxidation with Sulfuric Acid

• Fe becomes FeSO4, and its oxidation state changes from 0 to +2.
• Zn becomes ZnSO4, and its oxidation state changes from 0 to +2.
• Mg becomes MgSO4, and its oxidation state changes from 0 to +2.

These changes confirm that the metals have been oxidized since they have lost electrons to the acid, which reduces the hydrogen ions to hydrogen gas.

Acid Rain Chemistry

Acid rain refers to precipitation that has a lower pH than normal rainwater, due to the presence of dissolved acidic compounds like sulfuric acid and nitric acid. These acids form in the atmosphere when sulfur dioxide (SO2) and nitrogen oxides (NOx) react with water, oxygen, and other chemicals.

When sulfuric acid formed in the atmosphere falls as acid rain, it can have harmful effects on buildings, infrastructure, ecosystems, and human health. The reactivity of \( \mathrm{H}_2\mathrm{SO}_4 \) with metals used in construction, as shown in the exercise, leads to corrosion and weakening of metal structures. Understanding the chemistry of acid rain is crucial for developing strategies to mitigate its impacts and protect the environment and constructed materials.