Question

Indicate whether the following balanced equations involve oxidation-reduction. If they do, identify the elements that undergo changes in oxidation number. (a) \(2 \mathrm{AgNO}_{3}(a q)+\mathrm{CoCl}_{2}(a q) \longrightarrow 2 \mathrm{AgCl}(s)+ \mathrm{Co}\left(\mathrm{NO}_{3}\right)_{2}(a q)\) (b) \(2 \mathrm{PbO}_{2}(s) \longrightarrow 2 \mathrm{PbO}(s)+\mathrm{O}_{2}(g)\) (c) \(2 \mathrm{H}_{2} \mathrm{SO}_{4}(a q)+2 \mathrm{NaBr}(s) \longrightarrow \mathrm{Br}_{2}(l)+\mathrm{SO}_{2}(g)+ \mathrm{Na}_{2} \mathrm{SO}_{4}(a q)+2 \mathrm{H}_{2} \mathrm{O}(l)\)

Short answer

(a) The reaction does not involve oxidation-reduction, as no elements have changes in oxidation numbers. (b) The reaction involves oxidation-reduction. Pb changes from +4 to +2, and O changes from -2 to 0. (c) The reaction involves oxidation-reduction. S changes from +6 to +4, and Br changes from -1 to 0.

Step-by-step solution

Assign oxidation numbers to all elements

We must assign oxidation numbers to all elements involved in each balanced equation. We will use the following rules: 1. The oxidation number of an element in its free state is 0. 2. The oxidation number of a monoatomic ion is equal to its charge. 3. The oxidation number of hydrogen is +1 if it is bonded to a non-metal and -1 if bonded to a metal. 4. The oxidation number of oxygen is -2, except in peroxides (-1) and when bonded to fluorine (+2). 5. The oxidation number of fluorine is -1; all other halogens have an oxidation number of -1 in their compounds, except when they are bonded to a more electronegative element. 6. The sum of the oxidation numbers of all atoms in a compound must be zero; the sum of the oxidation numbers in a polyatomic ion must be equal to its charge.

Compare the oxidation numbers

We will compare the oxidation numbers before and after the reaction for each element in each balanced equation to determine if any changes occurred. (a) \(2\mathrm{AgNO}_{3}(a q)+\mathrm{CoCl}_{2}(a q) \longrightarrow 2\mathrm{AgCl}(s)+ \mathrm{Co}\left(\mathrm{NO}_{3}\right)_{2}(a q)\) Ag: +1 -> +1 (no change) N: +5 -> +5 (no change) O: -2 -> -2 (no change) Co: +2 -> +2 (no change) Cl: -1 -> -1 (no change) (b) \(2\mathrm{PbO}_{2}(s) \longrightarrow 2\mathrm{PbO}(s)+\mathrm{O}_{2}(g)\) Pb: +4 -> +2 (change) O: -2 -> 0 (change) (c) \(2\mathrm{H}_{2}\mathrm{SO}_{4}(a q)+2\mathrm{NaBr}(s) \longrightarrow \mathrm{Br}_{2}(l)+\mathrm{SO}_{2}(g)+ \mathrm{Na}_{2}\mathrm{SO}_{4}(a q)+2\mathrm{H}_{2}\mathrm{O}(l)\) H: +1 -> +1 (no change) S: +6 -> +4 (change) O: -2 -> -2 (no change) Na: +1 -> +1 (no change) Br: -1 -> 0 (change)

Identify elements with changes in oxidation numbers

We can now identify the elements that have undergone changes in oxidation numbers in the given balanced equations: (a) No elements underwent any changes in oxidation numbers, so the reaction does not involve oxidation-reduction. (b) In equation (b), Pb and O have changes in oxidation numbers. Pb changes from +4 to +2 and O changes from -2 to 0. Therefore, this is an oxidation-reduction reaction. (c) In equation (c), S and Br have changes in oxidation numbers. S changes from +6 to +4, and Br changes from -1 to 0. Therefore, this is an oxidation-reduction reaction.

Key concepts

Oxidation Numbers

Oxidation numbers help us track how many electrons are lost or gained by an element during a chemical reaction. They are fundamental in determining whether a reaction involves redox processes. An oxidation number is like a tool that indicates the degree of oxidation (loss of electrons) or reduction (gain of electrons) of an element in a compound.

• For an element in its free state, the oxidation number is 0. For example, O₂ or Fe has an oxidation number of 0.
• In a monoatomic ion, the oxidation number is equal to the charge of the ion. For instance, Na⁺ would have an oxidation number of +1.

Common rules are applied for assigning oxidation numbers:

• Hydrogen generally has an oxidation number of +1, except when bonded to metals in which it is -1.
• Oxygen typically is assigned an oxidation number of -2, except in peroxides (-1) or when bonded with fluorine.
• The sum of the oxidation numbers in a neutral compound must be zero, and for a polyatomic ion, it must be equal to the ion’s charge.

These rules are crucial to identifying changes in oxidation numbers, which in turn help us understand the happening of redox reactions in chemical equations.

Redox Chemistry

Redox chemistry refers to oxidation-reduction reactions where the oxidation state of atoms is altered. These reactions are vital because they involve the transfer of electrons between substances.

An element is oxidized if it loses electrons, which increases its oxidation number. Conversely, an element is reduced if it gains electrons, thereby decreasing its oxidation number.

Redox reactions can often be split into two half-reactions: oxidation and reduction. In this way, you can clearly see the electron transfer.

• In the oxidation half, the substance loses electrons. For example, in reaction (c) from the exercise, Br changes from -1 to 0, indicating oxidation as it loses electrons.
• In the reduction half, the substance gains electrons. In the same reaction, S changes from +6 to +4 when it is reduced, signifying a gain of electrons.

Determining these halves is very important because it helps chemists both understand and predict the outcomes of reactions, as well as identify the flow of electrons during the chemical process.

Chemical Equations

Chemical equations represent how reactants transform into products in a chemical reaction. For redox reactions, correctly balanced equations are essential in identifying changes in oxidation numbers. A balanced chemical equation must obey the law of conservation of mass, where the number of atoms for each element is the same on both sides of the equation.

In redox reactions, it's not just about balancing atoms, but also about balancing electron transfer. An equation is balanced when:

• The atoms for each element are equal on both sides.
• The charge is balanced, meaning the total charge is the same on both sides, indicating that electrons lost = electrons gained.

For example, in exercise (b), the balanced equation shows changes in lead ( ext{Pb}) and oxygen ( ext{O}) oxidation numbers, highlighting a clear redox reaction.
Overall, chemical equations provide a detailed depiction of chemical reactions and their dynamics, especially in determining the nature and extent of oxidation and reduction activities involved.