Balance the following net ionic reactions, and identify...

Balance the following net ionic reactions, and identify which elements are oxidized and which are reduced: a. $\mathrm{MnO}_{2}(s)+\mathrm{HCl}(a q) \rightarrow \mathrm{Mn}^{2+}(a q)+\mathrm{Cl}_{2}(g)$ b. $\mathrm{I}_{2}(s)+\mathrm{S}_{2} \mathrm{O}_{3}^{2-}(a q) \rightarrow \mathrm{S}_{4} \mathrm{O}_{6}^{2-}(a q)+\mathrm{I}^{-}(a q)$ c. $\mathrm{MnO}_{4}^{-}(a q)+\mathrm{Fe}^{2+}(a q) \rightarrow \mathrm{Mn}^{2+}(a q)+\mathrm{Fe}^{3+}(a q)$

Balance the following net ionic reactions, and identify which elements are oxidized and which are reduced:
a. $\mathrm{MnO}_{2}(s)+\mathrm{HCl}(a q) \rightarrow \mathrm{Mn}^{2+}(a q)+\mathrm{Cl}_{2}(g)$
b. $\mathrm{I}_{2}(s)+\mathrm{S}_{2} \mathrm{O}_{3}^{2-}(a q) \rightarrow \mathrm{S}_{4} \mathrm{O}_{6}^{2-}(a q)+\mathrm{I}^{-}(a q)$
c. $\mathrm{MnO}_{4}^{-}(a q)+\mathrm{Fe}^{2+}(a q) \rightarrow \mathrm{Mn}^{2+}(a q)+\mathrm{Fe}^{3+}(a q)$

Key Concepts

Oxidizing and Reducing Agents

In redox reactions, the oxidizing agent is the species that gains electrons and is reduced, while the reducing agent is the one that loses electrons and is oxidized. Recognizing these agents is crucial for predicting the direction of the electron transfer and the overall feasibility of the reaction.

Net Ionic Equations

Net ionic equations focus only on the species that participate directly in the chemical change, omitting the spectator ions that remain unchanged. This simplification allows for a clearer understanding of the main chemical transformation, particularly in aqueous reactions where many ions are present.

Redox Reactions

Redox (reduction-oxidation) reactions involve the transfer of electrons between chemical species. In these reactions, one species loses electrons (is oxidized) while another gains electrons (is reduced). Understanding redox reactions is fundamental for analyzing how chemicals interact in processes ranging from corrosion to metabolism.

Oxidation States

The oxidation state (or oxidation number) of an element in a compound or ion provides a way of tracking electron transfer during chemical reactions. Changes in oxidation states indicate which species have gained or lost electrons. This concept is essential for identifying which components of a reaction are oxidized or reduced.

Half-Reactions

The half-reaction method involves splitting a redox process into two separate reactions—one oxidation and one reduction. Each half-reaction is balanced individually for mass and charge before combining them to yield the overall balanced equation. This approach simplifies the balancing of complex redox reactions.

Transcript

00:01 There are a couple ways that you can balance redox reactions.

00:04 The textbook describes the half -reaction method and the oxidation number method.

00:12 I find it easier to use the half -reaction method.

00:16 It is a little bit more comprehensive and allows you to add water and or hydrogen ions when the reaction is occurring in acidic solution.

00:28 And it turns out that for a few of these in problem 103 and problems 103 and 104, you do need to add in some water and hydrogen ions.

00:41 So all of these will be balanced using the half -reaction method.

00:46 For the first reaction, we have mno2 solid plus hcl, going to mn2 plus and cl2, going to mn2 plus and cl2, gal.

00:58 So for the half reaction method, we will recognize that manganese mno2 solid goes to mn2 plus, and the hcl goes to cl2.

01:11 The next step after writing the half reactions is to balance all of the elements except for hydrogen and oxygen.

01:19 We have mn2 plus, 1mn2 plus, and 1mn2 plus it's balanced.

01:25 We have 2 cls.

01:26 We had 1 cls, so we added 2.

01:28 To here to make sure that it is balanced.

01:32 Then we balance all oxygens with water.

01:36 So because we have two oxygens over here, we need two waters on the right hand side.

01:43 Then we can balance hydrogen ions.

01:45 I'm sorry, we can balance hydrogen atoms with hydrogen ions, assuming that it is an in acidic solution, which i believe for all of these questions, it is assumed that they are occurring in acidic solution.

01:57 So we have four hydrogen hydrogens on the right -hand side.

02:02 We need four hydrogen ions on the left -hand side.

02:06 Then we look at the hydrogens on the left -hand side for the second half -reaction, and we add two hydrogen ions on the right -hand side.

02:16 Then we balance charges with electrons.

02:20 We have a total of a 2 -plus charge on the right -hand side, a 4 -plus charge on the left -hand side for this first -half reaction.

02:27 So we need two electrons to make sure that it is balanced with a 2 plus charge on both sides.

02:33 Then we had no charge down here.

02:36 We had a 2 plus charge here, so we needed two electrons on the right hand side.

02:40 This will always be the case.

02:42 One half reaction will have electrons on the left hand side, the other will have electrons on the right hand side.

02:47 Then we sum things up to make sure our electrons cancel.

02:51 Turns out two hydrogen ions will be canceled.

02:55 And then we get two hydrogen ions, plus two hcls plus mno2 goes to mn2 plus and cl2 and h2 and two h2os.

03:09 Then we see that we have two hydrogens plus two more hydrogens gives us four.

03:16 We have four hydrogens.

03:18 We have one manganese, one manganese.

03:21 We have two cls.

03:23 We have two cls.

03:24 We also need to make sure the charges are balanced...

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