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Chapter 18: Problem 14

Indicate the oxidizing and reducing agents in the following redox reactions: (a) \(\mathrm{Ca}(\mathrm{s})+2 \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{Ca}(\mathrm{OH})_{2}(a q)+\mathrm{H}_{2}(g)\) (b) \(\mathrm{Mg}(s)+2 \mathrm{HCl}(a q) \longrightarrow \mathrm{MgCl}_{2}(a q)+\mathrm{H}_{2}(g)\)

Short Answer

Expert verified
(a) Oxidizing agent: H鈧侽, Reducing agent: Ca; (b) Oxidizing agent: HCl, Reducing agent: Mg.

Step by step solution

01

Identify the Initial and Final Oxidation States

For each species in the reactions, determine the initial and final oxidation states. In reaction (a), calcium (Ca) starts in the elemental state with an oxidation number of 0, and changes to +2 in Ca(OH)鈧�, while hydrogen (H) shifts from +1 in H鈧侽 to 0 in H鈧�. In reaction (b), magnesium (Mg) starts at 0 and changes to +2 in MgCl鈧�, and hydrogen shifts from +1 in HCl to 0 in H鈧�.
02

Determine the Oxidizing Agent

The oxidizing agent is the species that gets reduced by gaining electrons. In reaction (a), hydrogen in water (H鈧侽) gains electrons and is reduced to H鈧�. Thus, water (H鈧侽) acts as the oxidizing agent. In reaction (b), hydrogen in hydrochloric acid (HCl) gains electrons and is reduced to H鈧�. Therefore, HCl is the oxidizing agent.
03

Determine the Reducing Agent

The reducing agent is the species that gets oxidized by losing electrons. In reaction (a), calcium (Ca) loses electrons and is oxidized to Ca虏鈦�. Therefore, calcium (Ca) is the reducing agent. In reaction (b), magnesium (Mg) loses electrons and is oxidized to Mg虏鈦�, making magnesium (Mg) the reducing agent.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Oxidizing Agent
In a redox reaction, the oxidizing agent is the substance that undergoes reduction. In simpler terms, it is the species that gains electrons. This process happens since the oxidizing agent causes another substance to lose electrons, therefore enabling their own reduction. It's like a friend who lets you borrow their crayons: you gain color variety, while they part with a few crayons.

For example, in the reactions provided:
  • In reaction (a): Water (\(\mathrm{H}_{2} \mathrm{O}\)) acts as the oxidizing agent. It accepts electrons from calcium, resulting in the reduction of hydrogen from \(+1\) to \(0\) in \(\mathrm{H}_{2}\).
  • In reaction (b): Hydrochloric acid (\(\mathrm{HCl}\)) serves as the oxidizing agent. The hydrogen ion in \(\mathrm{HCl}\) accepts electrons from magnesium, and reduces from \(+1\) in \(\mathrm{HCl}\) to \(0\) in \(\mathrm{H}_{2}\).
Recognizing the oxidizing agent helps in understanding the electron transfer that supports the reaction dynamics.
Reducing Agent
The reducing agent in a redox reaction is the species that donates electrons and becomes oxidized. Essentially, it is the electron giver of the reaction group, helping another substance receive the electrons it donates. Think of it like sharing your packed lunch with a friend, lessening your own but making them satisfied.

For the given examples:
  • In reaction (a): Calcium (\(\mathrm{Ca}\)) is the reducing agent. It loses electrons, seeing its oxidation state change from \(0\) to \(+2\) in \(\mathrm{Ca(OH)}_{2}\).
  • In reaction (b): Magnesium (\(\mathrm{Mg}\)) functions as the reducing agent. It is oxidized as its electrons are given away, leading to a shift from \(0\) to \(+2\) in \(\mathrm{MgCl}_{2}\).
Grasping the concept of the reducing agent helps you see the role of substances in driving the redox reaction.
Oxidation States
Oxidation states, often called oxidation numbers, are positive or negative values assigned to elements in a chemical species. They help indicate the degree of oxidation, or loss of electrons, of an element in a compound.

Using oxidation states, you can track how electrons are distributed in molecules or ions involved in a reaction. Here's how they shift in our examples:
  • In reaction (a):
    • Calcium (\(\mathrm{Ca}\)) shifts from oxidation state \(0\) to \(+2\).
    • Hydrogen (\(\mathrm{H}\)), in water, moves from \(+1\) to \(0\) when forming \(\mathrm{H}_{2}\).
  • In reaction (b):
    • Magnesium (\(\mathrm{Mg}\)) transitions from \(0\) to \(+2\).
    • Hydrogen in HCl changes from \(+1\) to \(0\) in \(\mathrm{H}_{2}\).
By understanding oxidation states, we can better analyze how atoms in a reaction gain or lose electrons. This is essential for understanding the entirety of redox processes.

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Most popular questions from this chapter

Indicate the oxidizing and reducing agents in the following redox reactions: (a) \(\mathrm{Cr}^{2+}(a q)+\mathrm{AgI}(s) \longrightarrow \mathrm{Cr}^{3+}(a q)+\mathrm{Ag}(s)+\Gamma(a q)\) (b) \(\operatorname{Sn}^{2+}(a q)+\mathrm{Hg}^{2+}(a q) \longrightarrow \mathrm{Sn}^{4+}(a q)+\mathrm{Hg}_{2}^{2+}(a q)\)

Calculate the oxidation number for silicon in the following compounds: (a) \(\mathrm{Si}_{2} \mathrm{H}_{6}\) (b) \(\mathrm{SiO}_{2}\) (c) \(\mathrm{Si}_{2} \mathrm{~N}_{4}\) (d) \(\mathrm{CaSiO}_{3}\)

Indicate the oxidizing and reducing agents in the following redox reactions: (a) \(\mathrm{CuO}(s)+\mathrm{H}_{2}(g) \longrightarrow \mathrm{Cu}(s)+\mathrm{H}_{2} \mathrm{O}(l)\) (b) \(\mathrm{PbO}(s)+\mathrm{CO}(g) \longrightarrow \mathrm{Pb}(s)+\mathrm{CO}_{2}(g)\)

The nonspontaneous redox reaction of chromium metal and aqueous aluminum nitrate is an electrolytic process according to the following ionic equation: $$ \mathrm{Cr}(s)+\mathrm{Al}^{3+}(a q) \longrightarrow \mathrm{Al}(s)+\mathrm{Cr}^{3+}(a q) $$ The half-reactions are separated into two compartments. A Cr electrode is placed in \(1.00 \mathrm{M} \mathrm{Cr}\left(\mathrm{NO}_{3}\right)_{3},\) and an \(\mathrm{Al}\) electrode in \(1.00 \mathrm{M} \mathrm{Al}\left(\mathrm{NO}_{3}\right)_{3} .\) Indicate each of the following for the electrolytic cell: (a) oxidation half-cell reaction (b) reduction half-cell reaction (c) anode and cathode (d) direction of electron flow (e) direction of \(\mathrm{NO}_{3}^{-}\) in the salt bridge

Write a balanced equation for each of the following redox reactions in a basic solution using the half-reaction method: (a) \(\mathrm{Cl}_{2}(g)+\mathrm{BrO}_{2}^{-}(a q) \longrightarrow \mathrm{Cl}^{-}(a q)+\mathrm{BrO}_{3}^{-}(a q)\) (b) \(\mathrm{MnO}_{2}(s)+\mathrm{O}_{2}(g) \longrightarrow \mathrm{MnO}_{4}^{-}(a q)+\mathrm{H}_{2} \mathrm{O}(l)\)
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