Complete and balance the following equations, and identify the oxidizing and reducing agents. (Recall that the O atoms in hydrogen peroxide, H 2 O 2 , have an atypical oxidation state.) (a) NO 2 - (aq) + Cr 2 O 7 2- (aq) → Cr 3+ (aq) + NO 3 - (aq) (acidic solution) (b) S(s) + HNO 3 (aq) → H 2 SO 3 (aq) + N 2 O(g) (acidic solution) (c) Cr 2 O 7 2- (aq) + CH 3 OH(aq) → HCOOH(aq) + Cr 3+ (aq) (acidic solution) (d) BrO 3 - (aq) + N 2 H 4 (g) → Br - (aq) + N 2 (g) (acidic solution)

everyone. So ask the balance of following redox reaction under acidic conditions and identify the oxidizing and reducing agents. I need to first separate the whole reaction into two half reactions. F. cr 207 two minus. And this goes to see our three plus. Then we have you four us And this goes to you 02 10 plus. Now we need to balance the non hydrogen and non oxygen elements first. So for the first reaction, you need to balance the chromium. This is already balanced. So for the first we actually need to balance the chromium to chromium on the raft inside. But one on the product side we can put a two in front of cr three plus. Now we need to balance the oxygen by adding H20 liquid to the side. That needs oxygen seven auction on the wrapped inside. And the first equation, We need to add 7 H20 liquid to the product side. For the second reaction have to auction on the product side, we need to add to H20 liquid to the reacting side. Now now I need to balance the hydrogen by adding H plus to decide they need hydrogen 14 hydrogen on the product side, they need to add 14 H plus to the acting side on the first reaction. And then for the second reaction we have four hydrogen on the right hand side. So we need to add four H plus to the product side. Now I need to balance the charges and add the electrons to the more positive side, have a look at the charges plus 14 here. I got to hear six here. You're out here. We have a zero here plus four here. Last two here. Last four here on the first reaction we're gonna need to add electrons to the reacting side. We need to add 14. So we need to add six electrons. For the second reaction we need to add electrons to the product side. I need to add two electrons. Now I need to balance electrons on the two half reactions. We have six on the top and two on the bottom and he's most probably three on the bottom. And for the oxidation reactions were gonna lose electrons. We're gonna have electrons on the product side. For the reduction reaction we're gonna gain electrons. We're gonna have electrons on the react inside. For oxidation half reaction We're gonna have three U four plus six H 2 L. It's gonna be a three year 02 two plus plus 12 h plus six electrons. And for reduction half reaction You see our two 72 minus last 14 H plus plus six electrons. Getting to cr three plus seven H 20. For our loss of electrons. This oxidation and this is our reducing agent, It's gonna be you four plus And then for a gain of electrons is reduction is an oxidizing agent. NBC are to go 7 to -. So now I need to get the overall reaction by adding the two reactions, canceling it out. We cancel out any common species. We're gonna get rid of 6 H20. We're gonna be left with 1 H20. Over here. We're gonna have to H plus over here. We're get we're gonna get rid of our electrons as well. So for the balanced reaction, Get three U 4 plus plus cr two plus, seven to minus plus two H plus. And then she was 302 in plus Plus 2, er 3 plus H 20. Thanks for watching my video and I hope it was helpful.

Ch.20 - Electrochemistry

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Brown 15th Edition

Ch.20 - Electrochemistry

Problem 29a,b,c,d

Chapter 20, Problem 29a,b,c,d

Complete and balance the following equations, and identify the oxidizing and reducing agents. (Recall that the O atoms in hydrogen peroxide, H₂O₂, have an atypical oxidation state.) (a) NO₂^-(aq) + Cr₂O₇^2-(aq) → Cr³⁺(aq) + NO₃^-(aq) (acidic solution) (b) S(s) + HNO₃(aq) → H₂SO₃(aq) + N₂O(g) (acidic solution) (c) Cr₂O₇^2- (aq) + CH₃OH(aq) → HCOOH(aq) + Cr³⁺(aq) (acidic solution) (d) BrO₃^-(aq) + N₂H₄(g) → Br^-(aq) + N₂(g) (acidic solution)

Verified step by step guidance

Identify the oxidation states of each element in the reactants and products. For example, in \( \text{Cr}_2\text{O}_7^{2-} \), chromium has an oxidation state of +6, and in \( \text{Cr}^{3+} \), it is +3.

Determine which elements are oxidized and which are reduced by comparing the changes in oxidation states. Oxidation involves an increase in oxidation state, while reduction involves a decrease.

Write the half-reactions for the oxidation and reduction processes. For example, the reduction half-reaction involves \( \text{Cr}_2\text{O}_7^{2-} \) being reduced to \( \text{Cr}^{3+} \).

Balance each half-reaction for mass and charge. In acidic solutions, add \( \text{H}^+ \) ions to balance hydrogen and \( \text{H}_2\text{O} \) to balance oxygen.

Combine the balanced half-reactions, ensuring that the electrons lost in the oxidation half-reaction equal the electrons gained in the reduction half-reaction. Identify the oxidizing agent (the species that is reduced) and the reducing agent (the species that is oxidized).

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Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Oxidation and Reduction

Oxidation and reduction are chemical processes that involve the transfer of electrons between substances. Oxidation refers to the loss of electrons, resulting in an increase in oxidation state, while reduction involves the gain of electrons, leading to a decrease in oxidation state. In redox reactions, one species is oxidized and another is reduced, which is essential for balancing chemical equations.

Balancing Redox Reactions

Balancing redox reactions requires ensuring that both mass and charge are conserved. This often involves separating the reaction into half-reactions for oxidation and reduction, balancing each half for atoms and charge, and then combining them to form the overall balanced equation. In acidic solutions, hydrogen ions (H+) and water (H2O) are commonly used to achieve balance.

Identifying Oxidizing and Reducing Agents

In a redox reaction, the oxidizing agent is the substance that gains electrons and is reduced, while the reducing agent is the substance that loses electrons and is oxidized. Identifying these agents involves analyzing the changes in oxidation states of the elements involved in the reaction. This understanding is crucial for determining the roles of different reactants in the overall chemical process.