8: Redox Reactions
- Page ID
- 516592
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\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)- To investigate the principles of redox reactions by observing the chemical conversion of iron(II) ions to iron(III) ions using various oxidizing agents (such as hydrogen peroxide, potassium permanganate, and sodium hypochlorite) and the reduction of iron(III) ions to iron(II) ions using various reducing agents (such as sodium sulfite, sodium bromide, and sodium iodide).
- To utilize chemical indicators (potassium thiocyanate and potassium ferricyanide) to test for and confirm the presence of iron(III) ions and iron(II) ions, respectively, after redox reactions have occurred.
- To observe and document the characteristic colorful states of manganese (such as Mn2+, MnO2, and MnO4−) corresponding to various oxidation numbers (+2 through +7) and use these visual changes to determine the products of redox reactions and write corresponding balanced half-reactions and overall equations.
INTRODUCTION
- Iron Ion Redox Chemistry (Parts A and B): These experiments involve observing the conversion between iron(II) ions and iron(III) ions. Part A examines the oxidation of Fe2+ to Fe3+ using various oxidizing agents, including hydrogen peroxide (H2O2), potassium permanganate (KMnO4), and sodium hypochlorite (NaClO). Conversely, Part B investigates the reduction of Fe3+ to Fe2+ using reducing agents like sodium sulfite (Na2SO3), sodium bromide (NaBr), and sodium iodide (NaI). The presence of the respective iron ions is confirmed using specific chemical indicators: potassium thiocyanate (KSCN) is utilized to test for Fe3+, and potassium ferricyanide (K3[Fe(CN)6]) is used to test for Fe2+.
- Colorful States of Manganese (Part C): This section focuses on the relationship between oxidation number and visible color for manganese. Reactions are performed involving manganese compounds (such as manganese(II) sulfate and potassium permanganate) to observe and document the characteristic colors corresponding to various oxidation numbers of manganese, which range from +2 (e.g., Mn2+, colorless or light pink) up to +7 (e.g., MnO4−, very dark purple). These observations are then used to analyze and write balanced half-reactions and overall equations for the redox processes.
Manganese occurs in the following oxidation numbers.
|
Oxidation number |
Species |
Color |
|---|---|---|
|
+2 |
\( \ce{Mn^{2+}}\) (aq) |
Colorless or light pink |
|
+3 |
\( \ce{Mn^{3+}}\) (aq) |
Rose or red |
|
+4 |
\( \ce{MnO_{2}}\) (s) |
Dark brown |
|
+5 |
\( \ce{MnO^{3-}_{4}}\) (aq) |
Blue |
|
+6 |
\( \ce{MnO^{2-}_{4}}\) (aq) |
Green |
|
+7 |
\( \ce{MnO^{-}_{4}}\) (aq) |
Very dark purple |
In redox chemistry, color is your data. Manganese is unique because its color tells you its oxidation state instantly.
- The Prediction: Before you mix anything in Part C, look at Table 1. If you start with purple \(\ce{MnO4^-}\) (+7) and reduce it slightly to green \(\ce{MnO4^{2-}}\) (+6), you know exactly one electron was transferred.
- The Data: If you see brown precipitate, you have hit the +4 state (\(\ce{MnO2}\)). This "visual voltmeter" allows you to determine the number of electrons transferred without doing any math.
Oxidation vs. Reduction:
- Oxidation: Loss of electrons (Oxidation Number increases).
- Reduction: Gain of electrons (Oxidation Number decreases).
Oxidizing vs. Reducing Agents:
- Oxidizing Agent: The species that takes electrons (it is reduced).
- Reducing Agent: The species that gives electrons (it is oxidized).
- 8.1: Redox Reactions - Experiment
- This page provides safety precautions for handling caustic chemicals, along with a list of necessary equipment for a lab experiment involving iron and manganese reactions. The experimental procedure is divided into three parts: reactions of iron ions with oxidizing agents, iron(III) with reducing agents, and observing manganese's color changes. It concludes with a focus on proper chemical disposal methods.
- 8.2: Redox Reactions - Pre-lab
- This page explores the confirmation of iron oxidation using potassium thiocyanate and the role of oxidizing agents. It also discusses potassium ferricyanide for indicating the reduction of iron(III) to iron(II) and the oxidation of sulfite to sulfate, requiring oxidation number assignments. Furthermore, it includes a logic check with halogens to predict oxidative capabilities based on redox reaction properties.
- 8.3: Redox Reactions - Data and Report
- This page outlines a laboratory experiment focused on redox reactions involving iron(II) and iron(III) with various oxidizing and reducing agents. It provides tables for observation recording, instructions for calculating oxidation states, and questions on balancing half-reactions and writing overall equations. Theoretical questions aid in comparing halides to identify stronger reducing agents, emphasizing the principles of reaction dynamics and redox chemistry.