Question

Identify what is oxidized and what is reduced in the following processes. a. \(2 \mathrm{Br}^{-}+\mathrm{Cl}_{2} \rightarrow \mathrm{Br}_{2}+2 \mathrm{Cl}^{-}\) b. \(2 \mathrm{Ce}+3 \mathrm{Cu}^{2+} \rightarrow 3 \mathrm{Cu}+2 \mathrm{Ce}^{3+}\) c. \(2 \mathrm{Zn}+\mathrm{O}_{2} \rightarrow 2 \mathrm{ZnO}\) d. \(2 \mathrm{Na}+2 \mathrm{H}^{+} \rightarrow 2 \mathrm{Na}^{+}+\mathrm{H}_{2}\)

Short answer

a. Br⁻ is oxidized to Br₂; Cl₂ is reduced to Cl⁻. b. Ce is oxidized to Ce³⁺; Cu²⁺ is reduced to Cu. c. Zn is oxidized to ZnO; O₂ is reduced to O²⁻ in ZnO. d. Na is oxidized to Na⁺; H⁺ is reduced to H₂.

Step-by-step solution

Identify the oxidation states of each atom

Before the reaction, Br has an oxidation state of -1, Cl in Cl2 has an oxidation state of 0. After the reaction, Br in Br2 has an oxidation state of 0, and Cl has an oxidation state of -1.

Determine the changes in oxidation states

Comparing their oxidation states before and after the reaction, Br changed from -1 to 0 (loss of electron), and Cl changed from 0 to -1 (gain of electron).

Identify whether oxidation or reduction occurred

Since Br lost an electron, it was oxidized, whereas Cl gained an electron, so it was reduced. For reaction b:

Identify the oxidation states of each atom

Before the reaction, Ce has an oxidation state of 0, Cu has an oxidation state of +2. After the reaction, Ce has an oxidation state of +3, Cu has an oxidation state of 0.

Determine the changes in oxidation states

Comparing their oxidation states before and after the reaction, Ce changed from 0 to +3 (loss of electrons), and Cu changed from +2 to 0 (gain of electrons).

Identify whether oxidation or reduction occurred

Since Ce lost electrons, it was oxidized, whereas Cu gained electrons, so it was reduced. For reaction c:

Identify the oxidation states of each atom

Before the reaction, Zn has an oxidation state of 0, O in O2 has an oxidation state of 0. After the reaction, Zn in ZnO has an oxidation state of +2, O has an oxidation state of -2.

Determine the changes in oxidation states

Comparing their oxidation states before and after the reaction, Zn changed from 0 to +2 (loss of electrons), and O changed from 0 to -2 (gain of electrons).

Identify whether oxidation or reduction occurred

Since Zn lost electrons, it was oxidized, whereas O gained electrons, so it was reduced. For reaction d:

Identify the oxidation states of each atom

Before the reaction, Na has an oxidation state of 0, H has an oxidation state of +1. After the reaction, Na has an oxidation state of +1, H in H2 has an oxidation state of 0.

Determine the changes in oxidation states

Comparing their oxidation states before and after the reaction, Na changed from 0 to +1 (loss of electron), and H changed from +1 to 0 (gain of electron).

Identify whether oxidation or reduction occurred

Since Na lost an electron, it was oxidized, whereas H gained an electron, so it was reduced.

Key concepts

Oxidation

Oxidation is a key concept in redox reactions. It involves the loss of electrons from an atom or molecule. When an element or compound undergoes oxidation, its oxidation state increases. This concept can be seen in different processes such as in the reactions provided. For example, in the reaction of bromide ions with chlorine gas, bromide (\( \text{Br}^- \) ) loses an electron to form bromine (\( \text{Br}_2 \) ), thus it is oxidized. Another example is zinc in the reaction with oxygen. Initially, zinc has an oxidation state of zero, but after it reacts with oxygen to form zinc oxide (\( \text{ZnO} \) ), its oxidation state increases to +2. This indicates a loss of electrons.

• Oxidation involves electron loss.
• The oxidation state increases in the oxidized substance.
• It's essential to identify what gets oxidized in redox reactions.

Reduction

Reduction is the counterpart of oxidation and involves the gain of electrons by an atom or molecule. When a substance undergoes reduction, its oxidation state decreases. For instance, chlorine in its gaseous form (\( \text{Cl}_2 \) ) gains electrons to form chloride ions (\( \text{Cl}^- \) ), showing that it has been reduced as its oxidation state shifts from 0 to -1. Similarly, in the transformation of copper ions (\( \text{Cu}^{2+} \) ) to pure copper, the ions gain electrons and the oxidation state changes from +2 to 0.

• Reduction involves electron gain.
• The oxidation state decreases in the reduced substance.
• Always occurs along with oxidation in a redox reaction.

Oxidation States

Understanding oxidation states is crucial in identifying redox reactions. Oxidation states, or oxidation numbers, help in tracking electron transfer between atoms in a chemical reaction. They serve as indicators of whether atoms have been oxidized or reduced. Normally calculated based on a set of rules, oxidation states dictate the extent of electron loss or gain. For instance, in the process where sodium reacts with hydrogen ions, sodium's oxidation state changes from 0 to +1, indicating oxidation, whereas hydrogen's oxidation state decreases from +1 to 0, pointing to reduction.

• Oxidation states help identify redox reactions.
• Changes in oxidation states indicate electron transfer.
• They are calculated using specific rules and guidelines.

Electron Transfer

Electron transfer is at the heart of redox reactions. It involves the movement of electrons from one species to another, causing changes in oxidation states. This transfer is the key feature that differentiates oxidation and reduction. When a substance is oxidized, it transfers electrons to another species, which is subsequently reduced. For example, electrons move from bromide ions to disperse into chlorine, causing bromide to oxidize and chlorine to reduce. Similarly, in reaction (d), sodium transfers its electrons to hydrogen ions, oxidizing sodium and reducing hydrogen.

• Electron transfer is essential in redox reactions.
• It leads to oxidation and reduction of involved species.
• It entails a shift in oxidation states reflecting electron movement.