Question

Specify which of the following are oxidation-reduction reactions, and identify the oxidizing agent, the reducing agent, the substance being oxidized, and the substance being reduced. a. \(\mathrm{Cu}(s)+2 \mathrm{Ag}^{+}(a q) \rightarrow 2 \mathrm{Ag}(s)+\mathrm{Cu}^{2+}(a q)\) b. \(\mathrm{HCl}(g)+\mathrm{NH}_{3}(\mathrm{~g}) \rightarrow \mathrm{NH}_{4} \mathrm{Cl}(s)\) c. \(\mathrm{SiCl}_{4}(l)+2 \mathrm{H}_{2} \mathrm{O}(l) \rightarrow 4 \mathrm{HCl}(a q)+\mathrm{SiO}_{2}(s)\) d. \(\mathrm{SiCl}_{4}(l)+2 \mathrm{Mg}(s) \rightarrow 2 \mathrm{MgCl}_{2}(s)+\operatorname{Si}(s)\) e. \(\mathrm{Al}(\mathrm{OH})_{4}^{-}(a q) \rightarrow \mathrm{AlO}_{2}^{-}(a q)+2 \mathrm{H}_{2} \mathrm{O}(l)\)

Short answer

The redox reactions are: a. \(Cu(s)+2Ag^+(aq) \rightarrow 2Ag(s)+Cu^{2+}(aq)\): Oxidizing agent: \(Ag^+\), Reducing agent: \(Cu\), Substance being oxidized: \(Cu\), Substance being reduced: \(Ag^+\) d. \(SiCl_4(l)+2Mg(s) \rightarrow 2MgCl_2(s)+Si(s)\): Oxidizing agent: \(SiCl_4\), Reducing agent: \(Mg\), Substance being oxidized: \(Mg\), Substance being reduced: \(SiCl_4\)

Step-by-step solution

Oxidation numbers for reaction a

\(Cu(0), Ag^+(+1)\) for the reactants and \(Ag(0), Cu^{2+}(+2)\) for the products. #b.

Oxidation numbers for reaction b

\(H(+1), Cl(-1), N(-3), H(+1)\) for the reactants and \(N(-3), H(+1), Cl(-1)\) for the product. #c.

Oxidation numbers for reaction c

\(Si(+4), Cl(-1), H(+1), O(-2)\) for the reactants and \(H(+1), Cl(-1), Si(+4), O(-2)\) for the products. #d.

Oxidation numbers for reaction d

\(Si(+4), Cl(-1), Mg(0)\) for the reactants and \(Mg(+2), Cl(-1), Si(0)\) for the products. #e.

Oxidation numbers for reaction e

\(Al(+3), O(-2), H(+1)\) for the reactant and \(Al(+3), O(-2), H(+1)\) for the products. Step 2: Compare the oxidation states of reactants and products for each reaction Check if there is any change in the oxidation states between reactants and products in each reaction. #a.

Comparing oxidation states for reaction a

Cu goes from 0 to +2, and Ag goes from +1 to 0, indicating a redox reaction. #b.

Comparing oxidation states for reaction b

No change in oxidation states, thus it's not a redox reaction. #c.

Comparing oxidation states for reaction c

No change in oxidation states, thus it's not a redox reaction. #d.

Comparing oxidation states for reaction d

Si goes from +4 to 0, and Mg goes from 0 to +2, indicating a redox reaction. #e.

Comparing oxidation states for reaction e

No change in oxidation states, thus it's not a redox reaction. Step 3: Identify the redox agents and the substances being oxidized/reduced in each redox reaction #a.

Redox agents and substances for reaction a

Oxidizing agent: \(Ag^+\), Reducing agent: \(Cu\), Substance being oxidized: \(Cu\), Substance being reduced: \(Ag^+\) #d.

Redox agents and substances for reaction d

Oxidizing agent: \(SiCl_4\), Reducing agent: \(Mg\), Substance being oxidized: \(Mg\), Substance being reduced: \(SiCl_4\)

Key concepts

Oxidation States

Understanding oxidation states is crucial when studying redox reactions. These states represent the hypothetical charge that an atom would have if all bonds to different atoms were fully ionic. To determine oxidation numbers, we follow specific rules:

• Elements in their standard state have an oxidation number of zero, like \(Cu\) in reaction a.
• For ions, the oxidation number is equal to the charge of the ion, such as \(Ag^+\) which has an oxidation state of +1.
• Common oxidation states: Oxygen is usually -2, and hydrogen is usually +1.

By comparing the oxidation states of elements in the reactants and products of a reaction, we can identify if a change in oxidation state has occurred, which indicates a possible redox reaction. For example, in reaction a, the change of copper from 0 to +2 and silver from +1 to 0 signifies a redox event.

Oxidizing Agent

An oxidizing agent, also known as an oxidant, is a substance that gains electrons during a chemical reaction. By accepting electrons, it helps another substance to be oxidized. Identifying the oxidizing agent requires understanding the flow of electrons within the reaction.
For instance, in reaction a (\(Cu(s) + 2Ag^+(aq) \rightarrow 2Ag(s) + Cu^{2+}(aq)\)), \(Ag^+\) acts as the oxidizing agent. This is because it gains electrons from copper, thus getting reduced itself while oxidizing the copper.

Reducing Agent

The reducing agent in a redox reaction is the component that loses electrons. By losing electrons, it causes another substance to gain them, thereby getting oxidized in the process.

• It's crucial to note who donates the electrons, as that defines the reducing agent.
• In reaction a, copper (\(Cu\)) is the reducing agent. It loses its electrons to \(Ag^+\), allowing silver to be reduced to \(Ag\).

Understanding the reducing agent reveals who is giving up electrons in the reaction, an essential aspect of redox processes.

Redox Reaction

Redox reactions, short for oxidation-reduction reactions, involve the transfer of electrons between two substances. These reactions encompass both the process of oxidation (loss of electrons) and reduction (gain of electrons).
A key indicator of a redox reaction is the change in oxidation states of the involved elements. Let's take reaction a as an example: the copper's oxidation state increases from 0 to +2 as it loses electrons, while silver ions (\(Ag^+\)) decrease from +1 to 0, reflecting electron gain.

• It's essential to look for paired changes in oxidation states to identify redox reactions.

Not all reactions qualify as redox; some, like reactions b, c, and e, show no changes in oxidation states and thus are ruled out. Recognizing redox reactions is fundamental for understanding how substances interact at the electron-transfer level.