Question

Specify which of the following equations represent oxidationreduction reactions, and indicate the oxidizing agent, the reducing agent, the species being oxidized, and the species being reduced. a. \(\mathrm{CH}_{4}(g)+\mathrm{H}_{2} \mathrm{O}(g) \rightarrow \mathrm{CO}(g)+3 \mathrm{H}_{2}(g)\) b. \(2 \mathrm{AgNO}_{3}(a q)+\mathrm{Cu}(s) \rightarrow \mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}(a q)+2 \mathrm{Ag}(s)\) c. \(\mathrm{Zn}(s)+2 \mathrm{HCl}(a q) \rightarrow \mathrm{ZnCl}_{2}(a q)+\mathrm{H}_{2}(g)\) d. \(2 \mathrm{H}^{+}(a q)+2 \mathrm{CrO}_{4}^{2-}(a q) \rightarrow \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q)+\mathrm{H}_{2} \mathrm{O}(i)\)

Short answer

The short answer: a. Redox reaction. Oxidizing Agent: H\(_2\)O Reducing Agent: CH\(_4\) Species oxidized: C in CH\(_4\) Species reduced: O in H\(_2\)O b. Redox reaction. Oxidizing Agent: AgNO\(_3\) Reducing Agent: Cu Species oxidized: Cu Species reduced: Ag in AgNO\(_3\) c. Redox reaction. Oxidizing Agent: HCl Reducing Agent: Zn Species oxidized: Zn Species reduced: H in HCl d. Not a redox reaction.

Step-by-step solution

Determine Oxidation States

For each equation, we will analyze the elements' oxidation states in the reactants and products. a. CH\(_4\) + H\(_2\)O → CO + 3H\(_2\) In CH\(_4\), C has -4, H has +1 In H\(_2\)O, O has -2, H has +1 In CO; C has +2, O has -2 In H\(_2\), H has +1 b. 2AgNO\(_3\) + Cu → Cu(NO\(_3\))\(_2\) + 2Ag In AgNO\(_3\), Ag has +1, N has +5, O has -2 In Cu, Cu has 0 In Cu(NO\(_3\))\(_2\), Cu has +2, N has +5, O has -2 In Ag, Ag has 0 c. Zn + 2HCl → ZnCl\(_2\) + H\(_2\) In Zn, Zn has 0 In HCl, H has +1, Cl has -1 In ZnCl\(_2\), Zn has +2, Cl has -1 In H\(_2\), H has +1 d. 2H\(^+\)(aq) + 2CrO\(_4^{2-}\)(aq) → Cr\(_2\)O\(_7^{2-}\)(aq) + H\(_2\)O(l) In H\(^+\), H has +1 In CrO\(_4^{2-}\), Cr has +6, O has -2 In Cr\(_2\)O\(_7^{2-}\), Cr has +6, O has -2 In H\(_2\)O, H has +1, O has -2

Identify Redox Reactions

Now, we examine the oxidation states to see if any changed during the reaction. a. C: -4 to +2 (change), other elements' oxidation states remain the same. Redox reaction. b. Ag: +1 to 0 (change), Cu: 0 to +2 (change), other elements' oxidation states remain the same. Redox reaction. c. Zn: 0 to +2 (change), H: +1 to 0 (change). Redox reaction. d. The oxidation states of the elements remain the same. Not a redox reaction.

Identify Oxidizing and Reducing Agents and Species

Now we identify the oxidizing agents, reducing agents, and the species being oxidized and reduced for the redox reactions (a, b, and c). a. Oxidizing Agent: H\(_2\)O Reducing Agent: CH\(_4\) Species oxidized: C in CH\(_4\) Species reduced: O in H\(_2\)O b. Oxidizing Agent: AgNO\(_3\) Reducing Agent: Cu Species oxidized: Cu Species reduced: Ag in AgNO\(_3\) c. Oxidizing Agent: HCl Reducing Agent: Zn Species oxidized: Zn Species reduced: H in HCl

Key concepts

Oxidizing Agent

An oxidizing agent is a substance that causes the oxidation of another species by accepting electrons from it.
In simpler terms, it's the element or compound that gets reduced in a chemical reaction. It gains electrons and undergoes a decrease in oxidation state.

For example, in the reaction between methane ( CH_4 ) and water (as seen in equation a), water acts as the oxidizing agent. Here, water ( H_2O ) accepts electrons from carbon in CH_4 , allowing carbon to be oxidized from an oxidation state of -4 to +2.
The oxygen within water gains electrons, reducing its effective oxidation state.

• Oxidizing agents are critical for the progress of chemical reactions such as combustion, respiration, and rusting.
• Without the oxidizing agent, the oxidation process, which is crucial for energy production and material transformations, cannot occur.

Understanding the role of oxidizing agents is essential for mastering redox chemistry and anticipating the products of chemical reactions.

Reducing Agent

A reducing agent, also known as the reductant, is the substance that donates electrons to another species.
It causes the reduction of other entities while it itself is oxidized. This means its oxidation state increases as it loses electrons.

Looking at the given reactions, copper ( Cu ) in equation b acts as the reducing agent. Copper loses electrons, which allows the silver ( Ag^+ ) in silver nitrate ( AgNO_3 ) to be reduced to metallic silver ( Ag ).
This process also changes the copper's oxidation state from 0 to +2 as it forms copper nitrate ( Cu(NO_3) _2 ).

• Reducing agents play a vital role in various industrial and biological processes, such as metal extraction and cellular respiration.
• By understanding reducing agents, students can predict which substances will be oxidized and which will be reduced in a reaction.

Mastering the behavior of reducing agents helps clarify the flow of electrons in redox reactions.

Oxidation States

The concept of oxidation states (or oxidation numbers) is a fundamental part of redox reactions, used to keep track of electron transfer.
An element's oxidation state indicates its degree of oxidation, or loss of electrons.

In a molecule or compound, each atom is assigned an oxidation state based on a set of rules:

• For elemental substances (like O_2 or metallic Cu ), the oxidation state is 0.
• The oxidation state for a simple ion is equal to its charge ( Na^+ has an oxidation state of +1).
• In compounds, oxygen typically has an oxidation state of -2, and hydrogen +1.

By examining changes in oxidation states, chemists can identify which species are oxidized and which are reduced.
For instance, in the reaction of zinc ( Zn ) with hydrochloric acid ( HCl ) from equation c, the oxidation state of zinc changes from 0 to +2, while hydrogen's changes from +1 to 0, marking this as a redox process.
Recognizing these shifts allows for the identification of reactants and products that partake in electron transfer.

Chemical Equations

Chemical equations are symbolic representations of chemical reactions, where reactants are transmuted into products.
These equations detail which molecules take part in the reaction and how they interact with each other.
Understanding and balancing chemical equations is crucial for accurately describing chemical processes.

For instance, in equation b, 2AgNO_3 + Cu → Cu(NO_3)_2 + 2Ag , the equation reveals how two moles of silver nitrate react with copper to produce copper nitrate and silver.
The coefficients in the equation represent the proportions of substances involved, ensuring mass conservation in accordance with the law of conservation of mass.

Redox reactions specifically involve the balancing of electron transfer, which might not be immediately obvious in a simple chemical equation.
Balancing such equations involves equalizing the number of electrons lost in oxidation and gained in reduction, ensuring the sum of charges remains constant.

• Balanced chemical equations provide insight into molecular changes occurring during a reaction.
• They form the basis for quantitative studies of chemical processes including calculating yields and determining reactant amounts needed.

Mastering chemical equations is a core competency in chemistry, essential for both theoretical understanding and practical application.