# 2.3 Oxidizing and Reducing Agents ## 2.3 Defining Oxidizing and Reducing Agents In a redox (reduction-oxidation) reaction, electrons are transferred between species. The substances involved are categorized as oxidizing or reducing agents based on their role in this electron transfer. **Oxidizing Agent (Oxidant):** - A substance that *accepts electrons*. - It *causes oxidation* in another substance. - Its own oxidation state *decreases* (it gets reduced). - Contains an element that is *gaining electrons*. **Reducing Agent (Reductant):** - A substance that *donates electrons*. - It *causes reduction* in another substance. - Its own oxidation state *increases* (it gets oxidized). - Contains an element that is *losing electrons*. An oxidizing agent facilitates the oxidation of another substance, while a reducing agent facilitates the reduction of another substance. ## Identifying Oxidizing and Reducing Agents To identify oxidizing and reducing agents in a chemical reaction, it is crucial to determine the oxidation states of all elements involved in the reactants and products. 1. **Assign Oxidation States:** Determine the oxidation number for each atom in the reactants and products. 2. **Identify Changes:** Look for elements whose oxidation states change. - If an element's oxidation state *increases*, it has been *oxidized*. The substance containing this element is the **reducing agent**. - If an element's oxidation state *decreases*, it has been *reduced*. The substance containing this element is the **oxidizing agent**. **Important Note:** The oxidizing or reducing agent refers to the *entire molecule or formula unit*, not just the atom that undergoes a change in oxidation number. For related concepts on assigning oxidation states, refer to . ### Examples of Identification **1. Phosphorus and Nitric Acid Reaction** Consider the reaction where phosphorus reacts with nitric acid: $$ 3 \stackrel{0}{\mathrm{P}} + 5 \mathrm{H}\stackrel{+5}{\mathrm{N}}\stackrel{-2}{\mathrm{O}}_{3} + 2 \mathrm{H}_{2}\mathrm{O} \longrightarrow 3 \mathrm{H}_{3}\stackrel{+5}{\mathrm{P}}\stackrel{-2}{\mathrm{O}}_{4} + 5 \stackrel{+2}{\mathrm{N}}\stackrel{-2}{\mathrm{O}} $$ - **Phosphorus (P):** Its oxidation state changes from $0$ in $\mathrm{P}$ to $+5$ in $\mathrm{H}_{3}\mathrm{PO}_{4}$. Since its oxidation state *increases*, $\mathrm{P}$ is undergoing **oxidation**. Therefore, elemental $\mathrm{P}$ is the **reducing agent**. - **Nitrogen (N):** Its oxidation state changes from $+5$ in $\mathrm{HNO}_{3}$ to $+2$ in $\mathrm{NO}$. Since its oxidation state *decreases*, $\mathrm{N}$ is undergoing **reduction**. Therefore, $\mathrm{HNO}_{3}$ (the molecule containing N) is the **oxidizing agent**. **2. Manganese Dioxide and Hydrochloric Acid Reaction** Consider the reaction of manganese dioxide with hydrochloric acid: $$ \stackrel{+4}{\mathrm{Mn}}\stackrel{-2}{\mathrm{O}}_{2} + 4 \mathrm{H}\stackrel{-1}{\mathrm{Cl}} \longrightarrow \stackrel{+2}{\mathrm{Mn}}\stackrel{-1}{\mathrm{Cl}}_{2} + 2 \mathrm{H}_{2}\mathrm{O} + \stackrel{0}{\mathrm{Cl}}_{2} $$ - **Manganese (Mn):** Its oxidation state changes from $+4$ in $\mathrm{MnO}_{2}$ to $+2$ in $\mathrm{MnCl}_{2}$. Since its oxidation state *decreases*, $\mathrm{Mn}$ is undergoing **reduction**. Therefore, $\mathrm{MnO}_{2}$ is the **oxidizing agent**. - **Chlorine (Cl):** Its oxidation state changes from $-1$ in $\mathrm{HCl}$ to $0$ in $\mathrm{Cl}_{2}$. Since its oxidation state *increases*, $\mathrm{Cl}$ is undergoing **oxidation**. Therefore, $\mathrm{HCl}$ is the **reducing agent**. ## Worked Examples ### Example 1: Reaction in Steel Manufacture In the manufacture of steel, iron(III) oxide reacts with carbon monoxide. Let's identify the oxidizing and reducing agents. **Reaction:** $$ \mathrm{Fe}_{2}\mathrm{O}_{3} + 3\mathrm{CO} \longrightarrow 2\mathrm{Fe} + 3\mathrm{CO}_{2} $$ **Step-by-step Identification:** 1. **Assign Oxidation States:** $$ \stackrel{+3}{\mathrm{Fe}}_{2}\stackrel{-2}{\mathrm{O}}_{3} + 3\stackrel{+2}{\mathrm{C}}\stackrel{-2}{\mathrm{O}} \longrightarrow 2\stackrel{0}{\mathrm{Fe}} + 3\stackrel{+4}{\mathrm{C}}\stackrel{-2}{\mathrm{O}}_{2} $$ 2. **Identify Changes in Oxidation States:** - **Carbon (C):** Changes from $+2$ in $\mathrm{CO}$ to $+4$ in $\mathrm{CO}_{2}$. The oxidation state *increases*. - **Iron (Fe):** Changes from $+3$ in $\mathrm{Fe}_{2}\mathrm{O}_{3}$ to $0$ in $\mathrm{Fe}$. The oxidation state *decreases*. 3. **Determine Oxidized and Reduced Species:** - Carbon is being oxidized (increase in oxidation state). - Iron is being reduced (decrease in oxidation state). 4. **Identify Reducing Agent:** The reactant $\mathrm{CO}$ contains the carbon atom that is being oxidized. Therefore, $\mathrm{CO}$ is the **reducing agent**. 5. **Identify Oxidizing Agent:** The reactant $\mathrm{Fe}_{2}\mathrm{O}_{3}$ contains the iron atom that is being reduced. Therefore, $\mathrm{Fe}_{2}\mathrm{O}_{3}$ is the **oxidizing agent**. ### Example 2: Concept Assessment Exercise 2.5 Identify the oxidizing and reducing agents in the following reaction: $$ \mathrm{CuO} + \mathrm{H}_{2} \longrightarrow \mathrm{Cu} + \mathrm{H}_{2}\mathrm{O} $$ **Solution:** 1. **Assign Oxidation States:** $$ \stackrel{+2}{\mathrm{Cu}}\stackrel{-2}{\mathrm{O}} + \stackrel{0}{\mathrm{H}}_{2} \longrightarrow \stackrel{0}{\mathrm{Cu}} + \stackrel{+1}{\mathrm{H}}_{2}\stackrel{-2}{\mathrm{O}} $$ 2. **Identify Changes:** - **Copper (Cu):** From $+2$ in $\mathrm{CuO}$ to $0$ in $\mathrm{Cu}$. (Decrease, *reduction*) - **Hydrogen (H):** From $0$ in $\mathrm{H}_{2}$ to $+1$ in $\mathrm{H}_{2}\mathrm{O}$. (Increase, *oxidation*) 3. **Identify Agents:** - Since $\mathrm{Cu}$ in $\mathrm{CuO}$ is reduced, $\mathrm{CuO}$ is the **oxidizing agent**. - Since $\mathrm{H}$ in $\mathrm{H}_{2}$ is oxidized, $\mathrm{H}_{2}$ is the **reducing agent**. ### Example 3: Activity 2.1 - Identifying Agents Experimentally **Procedure:** 1. Prepare solutions of ferrous sulphate ($\mathrm{FeSO}_{4}$) and potassium permanganate ($\mathrm{KMnO}_{4}$). 2. Transfer $10\,\mathrm{cm}^{3}$ of $\mathrm{FeSO}_{4}$ solution into a test tube. 3. Add $10\,\mathrm{cm}^{3}$ of dilute $\mathrm{H}_{2}\mathrm{SO}_{4}$. 4. Add a few drops of $\mathrm{KMnO}_{4}$ solution to the test tube. **Observation and Explanation:** - **Observation:** Upon adding $\mathrm{KMnO}_{4}$ (which is purple) to the acidified $\mathrm{FeSO}_{4}$ solution, the purple color of $\mathrm{KMnO}_{4}$ is discharged (disappears). - **Explanation:** - The disappearance of the purple color indicates that the $\mathrm{MnO}_{4^{-}}$ ion (where $\mathrm{Mn}$ is in the $+7$ oxidation state) has been reduced to a colorless $\mathrm{Mn}^{2+}$ ion (where $\mathrm{Mn}$ is in the $+2$ oxidation state). - This reduction of $\mathrm{KMnO}_{4}$ is caused by $\mathrm{FeSO}_{4}$. In this reaction, the $\mathrm{Fe}^{2+}$ ions from $\mathrm{FeSO}_{4}$ are oxidized to $\mathrm{Fe}^{3+}$ ions. - Therefore, $\mathrm{FeSO}_{4}$ reduces $\mathrm{KMnO}_{4}$, making $\mathrm{FeSO}_{4}$ the **reducing agent**. - Conversely, $\mathrm{KMnO}_{4}$ oxidizes $\mathrm{FeSO}_{4}$, making $\mathrm{KMnO}_{4}$ the **oxidizing agent**. For understanding the relative strength of oxidizing and reducing agents, see and . ## Possible Questions and Answers **Q:** What is the fundamental difference in electron transfer between an oxidizing agent and a reducing agent? **A:** An oxidizing agent *accepts electrons* (gets reduced), while a reducing agent *donates electrons* (gets oxidized). **Q:** In the reaction $2\mathrm{Na} + \mathrm{Cl}_{2} \rightarrow 2\mathrm{NaCl}$, identify the oxidizing agent. **A:** $\mathrm{Cl}_{2}$ is the oxidizing agent. Chlorine's oxidation state changes from $0$ to $-1$, meaning it is reduced. **Q:** Why is the *entire molecule* considered the oxidizing or reducing agent, and not just the atom that changes oxidation state? **A:** Chemical agents act as a whole species, providing electrons or accepting them as part of their molecular structure. For example, $\mathrm{H}_{2}\mathrm{SO}_{4}$ as an oxidizing agent, despite only sulfur changing oxidation state, functions as a complete entity in the reaction. | Property | Oxidizing Agent | Reducing Agent | | :------------------------ | :------------------------------------------------ | :---------------------------------------------- | | **Electron Transfer** | Accepts electrons (electron acceptor) | Donates electrons (electron donor) | | **Effect on Other Subs.** | Causes oxidation in another substance | Causes reduction in another substance | | **Effect on Self** | Gets reduced (oxidation state decreases) | Gets oxidized (oxidation state increases) | | **Contains Element** | Element undergoing reduction | Element undergoing oxidation | **Significance:** Understanding oxidizing and reducing agents is fundamental to comprehending a vast array of chemical reactions, from biological processes like respiration and photosynthesis to industrial applications such as metal extraction (for example, steel manufacture using $\mathrm{CO}$ to reduce iron oxides) and electrochemistry (batteries). These concepts are crucial for predicting reaction outcomes and designing chemical processes. For applications in electrochemical cells, see .