Question

In the following reactions, label the oxidizing agent and the reducing agent. (a) $$ \mathrm{P}_{4}(s)+5 \mathrm{O}_{2}(g) \longrightarrow \mathrm{P}_{4} \mathrm{O}_{10}(s) $$ (b) $$ \mathrm{Co}(s)+\mathrm{Cl}_{2}(g) \longrightarrow \mathrm{CoCl}_{2}(s) $$

Short answer

(a) Oxidizing agent: \( \mathrm{O}_2 \); Reducing agent: \( \mathrm{P}_4 \). (b) Oxidizing agent: \( \mathrm{Cl}_2 \); Reducing agent: \( \mathrm{Co} \).

Step-by-step solution

Determine Oxidation States in Reaction (a)

In reaction (a) \( \mathrm{P}_{4}(s) + 5 \mathrm{O}_{2}(g) \longrightarrow \mathrm{P}_{4}\mathrm{O}_{10}(s) \), phosphorus (P) starts with an oxidation state of 0 because it is in its elemental form \( \mathrm{P}_4 \). Oxygen (O) also starts with an oxidation state of 0, being in its diatomic elemental form \( \mathrm{O}_2 \). In the product \( \mathrm{P}_{4}\mathrm{O}_{10} \), each oxygen is typically \(-2\) in compounds, so phosphorus must be \(+5\) to balance the charges in \( \mathrm{P}_4\mathrm{O}_{10} \). Phosphorus is oxidized (its oxidation state increases from 0 to +5), and oxygen is reduced (its oxidation state decreases from 0 to -2).

Identify Agents in Reaction (a)

Since oxygen (\( \mathrm{O}_2 \)) is being reduced, it is the oxidizing agent because it causes the oxidation of phosphorus. Meanwhile, phosphorus (\( \mathrm{P}_4 \)) is the reducing agent because it causes the reduction of oxygen.

Determine Oxidation States in Reaction (b)

In reaction (b) \( \mathrm{Co}(s) + \mathrm{Cl}_2(g) \longrightarrow \mathrm{CoCl}_2(s) \), cobalt (Co) begins with an oxidation state of 0 as a free element, and chlorine (Cl) starts with a state of 0 as \( \mathrm{Cl}_2 \). In the product \( \mathrm{CoCl}_2 \), each chlorine is \(-1\) (common oxidation state for Cl), making cobalt \(+2\) to balance out the charges. Cobalt is oxidized (0 to +2), and chlorine is reduced (0 to -1).

Identify Agents in Reaction (b)

Chlorine (\( \mathrm{Cl}_2 \)) is being reduced, so it acts as the oxidizing agent because it causes the oxidation of cobalt. Cobalt is the reducing agent since it donates electrons for the reduction of chlorine.

Key concepts

Oxidation States

In the realm of chemistry, oxidation states, also known as oxidation numbers, help us keep track of electron transfers in chemical reactions. They indicate the hypothetical charge an atom would have if all bonds were completely ionic.

Some key points about oxidation states are:

• In their elemental form, atoms are always assigned an oxidation state of zero.
• When forming compounds, atoms may lose, gain, or share electrons, leading to changes in their oxidation states.
• An increase in oxidation state indicates oxidation, while a decrease indicates reduction.

Let’s apply this to reaction (a) from our original exercise:
- Phosphorus (\(\mathrm{P}_4\)) and oxygen (\(\mathrm{O}_2\)) both start with a 0 oxidation state.- In \(\mathrm{P}_4\mathrm{O}_{10}\), oxygen has an oxidation state of \(-2\), making phosphorus \(+5\) to keep the molecule neutral.

Using this tool, we can understand how electrons shift and interact during reactions.

Chemical Reactions

Chemical reactions are transformative processes where bonds between atoms break, and new ones form, altering the substances involved. Such transformations hinge on understanding reactants and products:

• Reactants are the starting materials.
• Products are new materials formed from the reaction.

To illustrate, let's ponder over reaction (b) from the initial exercise:
- The reactants, cobalt (\(\mathrm{Co}\)) and chlorine (\(\mathrm{Cl}_2\)), engage to form a new compound, \(\mathrm{CoCl}_2\).- Here, both reactants swap electrons to build a stable compound.

Such intricate electron exchanges underscore how compounds can form new chemical entities. Understanding these processes helps us predict product outcomes using reaction equations.

Redox Processes

Redox, short for reduction-oxidation, describes all chemical reactions where one atom undergoes reduction while another experiences oxidation. Reduction involves gaining electrons, whereas oxidation involves losing electrons. These reactions are significant for understanding the nature of oxidizing and reducing agents.

Looking closer at the example from the exercise:
- In reaction (a), phosphorus is oxidized with its oxidation state moving from 0 to +5, while oxygen is reduced from 0 to -2. - In reaction (b), cobalt is oxidized going from 0 to +2, and chlorine is reduced moving from 0 to -1.

Each reaction includes agents:

• The oxidizing agent in a redox process gains electrons, thus undergoing reduction itself.
• The reducing agent loses electrons, causing another species to reduce.

By grasping these processes and agents, we can appreciate the cyclical exchange of electrons shaping chemical transformations.