Complete and balance the following equations, and identify...

Complete and balance the following equations, and identify the oxidizing and reducing agents: $$ \begin{array}{l}{\text { (a) } \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q)+\mathrm{I}(a q) \longrightarrow \mathrm{Cr}^{3+}(a q)+\mathrm{IO}_{3}^{-}(a q)} \\ \quad {\text { (acidic solution) }} \\ {\text { (b) } \mathrm{MnO}_{4}^{-}(a q)+\mathrm{CH}_{3} \mathrm{OH}(a q) \longrightarrow \mathrm{Mn}^{2+}(a q)+} \\ \quad {\mathrm{HCOOH}(a q) \text { (acidic solution) }}\end{array} \\ {\text {(c) } \mathrm{I}_{2}(s)+\mathrm{OCl}^{-}(a q) \longrightarrow \mathrm{IO}_{3}^{-}(a q)+\mathrm{Cl}^{-}(a q)} \\ {\text { (acidic solution) }} \\ {\text { (d) } \mathrm{As}_{2} \mathrm{O}_{3}(s)+\mathrm{NO}_{3}(a q) \longrightarrow \mathrm{H}_{3} \mathrm{AsO}_{4}(a q)+\mathrm{N}_{2} \mathrm{O}_{3}(a q)} \\ {(\text { acidic solution })} \\ {\text { (e) } \operatorname{MnO}_{4}^{-}(a q)+\operatorname{Br}^{-}(a q) \longrightarrow \mathrm{MnO}_{2}(s)+\mathrm{BrO}_{3}^{-}(a q)} \\ {\text { (basic solution) }} \\ {\text { (f) } \mathrm{Pb}(\mathrm{OH})_{4}^{2-}(a q)+\mathrm{ClO}^{-}(a q) \longrightarrow \mathrm{PbO}_{2}(s)+\mathrm{Cl}^{-}(a q)} \\ {\text { (basic solution) }} $$

Complete and balance the following equations, and identify the oxidizing and reducing agents:
$$
\begin{array}{l}{\text { (a) } \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q)+\mathrm{I}(a q) \longrightarrow \mathrm{Cr}^{3+}(a q)+\mathrm{IO}_{3}^{-}(a q)} \\ \quad {\text { (acidic solution) }} \\ {\text { (b) } \mathrm{MnO}_{4}^{-}(a q)+\mathrm{CH}_{3} \mathrm{OH}(a q) \longrightarrow \mathrm{Mn}^{2+}(a q)+} \\ \quad {\mathrm{HCOOH}(a q) \text { (acidic solution) }}\end{array} \\ {\text {(c) } \mathrm{I}_{2}(s)+\mathrm{OCl}^{-}(a q) \longrightarrow \mathrm{IO}_{3}^{-}(a q)+\mathrm{Cl}^{-}(a q)} \\ {\text { (acidic solution) }} \\ {\text { (d) } \mathrm{As}_{2} \mathrm{O}_{3}(s)+\mathrm{NO}_{3}(a q) \longrightarrow \mathrm{H}_{3} \mathrm{AsO}_{4}(a q)+\mathrm{N}_{2} \mathrm{O}_{3}(a q)} \\ {(\text { acidic solution })} \\ {\text { (e) } \operatorname{MnO}_{4}^{-}(a q)+\operatorname{Br}^{-}(a q) \longrightarrow \mathrm{MnO}_{2}(s)+\mathrm{BrO}_{3}^{-}(a q)} \\ {\text { (basic solution) }} \\ {\text { (f) } \mathrm{Pb}(\mathrm{OH})_{4}^{2-}(a q)+\mathrm{ClO}^{-}(a q) \longrightarrow \mathrm{PbO}_{2}(s)+\mathrm{Cl}^{-}(a q)} \\ {\text { (basic solution) }}
$$

Key Concepts

Acidic and Basic Reaction Media

Redox reactions can occur under acidic or basic conditions, and the medium affects the balancing process. In acidic solutions, H+ and H2O are often used to balance hydrogen and oxygen atoms, whereas OH- and H2O are employed in basic solutions. Understanding the role of the reaction medium is key to properly balancing redox equations under different conditions.

Half-Reaction Method

The half-reaction method involves breaking down an overall redox reaction into two separate reactions—one for oxidation and one for reduction. Each half-reaction is balanced for mass and charge independently, and then combined to give a balanced overall equation. This method is especially useful in reactions that occur under different pH conditions.

Oxidizing and Reducing Agents

An oxidizing agent is a species that causes another substance to oxidize by accepting electrons, while itself being reduced. Conversely, a reducing agent donates electrons to another species, causing reduction, while being oxidized. Identifying these agents is crucial for understanding the electron flow and the overall chemical transformation in redox reactions.

Redox Reactions

Redox (reduction-oxidation) reactions involve the transfer of electrons between chemical species, leading to changes in their oxidation states. Recognizing the species that are oxidized (losing electrons) and reduced (gaining electrons) is fundamental to understanding and balancing these reactions.

Oxidation States

Determining oxidation states is essential in redox chemistry as it helps identify which atoms undergo changes during a reaction. By assigning oxidation numbers, one can ascertain how many electrons are transferred, forming the basis of the half-reaction method used for balancing redox processes.

Transcript

00:01 Podcast we're looking to balance complete equations, identify the oxidizing and the reducing agent.

00:07 So in our first part here we have cr207 to minus.

00:15 That's in the aqueous state.

00:17 I won't add on the states just to be able to run through this a bit quicker, but i'll read them out.

00:21 That's aqueous state add i minus.

00:24 Acqueous state at 3h2o in the liquid state.

00:32 That is in equilibrium with 2cr3 plus aqueous add 7h2o liquid state add i 03 minus aqueous add 6h plus aqueous where we have cr 207 to minus aqueous add a h plus add i minus all aqueous that is in equilibrium with 2r3 plus add 4 4h2o, liquid add i .o3 minus aqueous, where cr072 minus is our oxidizing agent, i minus is our reducing agent.

01:31 So next what we have is 4mno4 minus, aqueous add 12h plus add 5, c .h3oh, that is in equilibrium, 4mn.

01:48 2 plus add 11 h2o 5hc02h changing the oxidation state of mn and c we get mn 04 minus that is x add 4 multiplied by minus 2 equals minus 1 where x is plus 7 equals positive 2 for mn 2 plus and so m n 2 plus and so m n 04 minus is the oxidizing agent and ch3oh is the reducing agent.

02:44 So moving on to the next example.

02:55 So our overall reaction here is i2, add h2o, add 5 ocl minus, cine equilibrium of 2, i -03 minus, add 2h plus, add 5cl minus.

03:12 So we change our oxidation state of i.

03:16 So we have i2 is 2x, that's equal 0, x equals 0, i03 minus x add 3, multiplied by minus 2 equals minus 1, x equals positive 5.

03:33 So i 2 is the oxidizing agent is oxidized, so therefore it is the reducing agent...

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