# Redox Reactions

- Page ID
- 283118

Learning Objectives – Oxidation and Reduction Reactions

After completing this unit the student will be able to:

- Define electrochemical process.
- Differentiate between oxidation and reduction using a simple equation or daily products.
- Determine oxidation numbers and balance redox equations.
- Show relationship between electric current from redox reaction to the rate of reaction.
- Show relationship between voltage and free energy change of the reaction.
- Describe the relationship between work and potential difference of a reaction.
- Show how free energy is related to electrochemical reaction.
- Show relationship between current, voltage and resistance.
- Differentiate between Galvanic and Electrolytic cells.
- Calculate the voltage and Gibbs free energy for a chemical reaction.
- Assemble a redox reaction using two half equations.
- Determine the spontaneity of a chemical reaction.
- Write free hand notation of a redox reaction.
- Use the Nernst equation in solving redox reactions.
- Illustrate redox reaction using market examples of dry-cell batteries, lead-acid storage batteries, a fuel-cell car, Hydrogen–oxygen fuel cell etc.
- Solve any redox reaction problems using the electromotive force values in a standard reduction potentials.

## Background Material

Please read Chapter 10 on Redox Reaction on Page 230 in the textbook “Analytical Chemistry and Quantitative Analysis” by David S. Hage and James D. Carr.

## Questions before class

- What is electrochemical process? Define oxidation and reduction using electrons.
**Solution:**

- Using the following equation, 2Mg (
*s*) + O_{2}(*g*) → 2MgO (*s*); write two half equations for oxidation and reduction.**Solution:**

- Make a summary of rules governing oxidation numbers and balancing equations using oxidation numbers.
**Solution:**

** **

- Make a summary of rules governing redox reactions and how they are balanced in basics and acidic media.

- Write an equation relating electric charge (
*q*) with measured coulombs (C).**Solution:**

** **

- Give an equation for current relating it to quantity of charge (q) and time (t).
**Solution:**

** **

- Write a relationship equation between work (w) resulting from charge (
*q)*moving through a potential difference (*E)*:**Solution:**

- Relate free energy change (
**Δ***G*) and maximum possible electrical work.**Solution:**

- Using Ohm’s law give the relationship between current (I) and potential (V) or resistance (
**Ω**)**Solution:**

- What is the relationship between Power and work?
**Solution:**

- Using an example define a Galvanic cell, draw a sketch of the cell, name the electrodes, anode, cathode, salt bridge, write down the half reactions and a balanced overall cell equation.
**Solution:**

- Write the cell reaction-using notation.
\[\ce{Cd}(s) \mid \ce{Cd(NO3)2}(aq) \mid\mid \ce{AgNO3}(aq) \mid \ce{Ag}(s)\nonumber\]

**Solution:**

- Draw a sketch representing the standard hydrogen half equation at 1-atm and 1M acidic solution.
**Solution:**

- How can you use standard electrode potentials to calculate the cell potential, illustrate using SHE and a silver half cell (Ag
^{+}(*aq*)+e^{-}↔Ag(*s*) E° = 0.799) - Write down and equation for calculating the overall electrode potential of the cell.
**Solution:**

** **

- Write down Nernst equation.
**Solution:**

- Use Nernst equation calculate the overall cell potential for the following reaction:
\[\ce{Ag} (s) \mid \ce{Ag+} (aq)\: \mathrm{(0.001 = a)} \mid\mid \ce{Ag+} (aq)\: \mathrm{(0.010 = a)} \mid \ce{Ag} (s)\nonumber\]

**Solution:**

- What is the relationship between Δ
*G,*Δ*H, T and*Δ*S?***Solution:**

__Questions in class Group 1:__

Given a cell with the following half equations:

\[\ce{C} (s) + \ce{2O^2-} (aq) → \ce{CO2} (g) + \ce{4e-}\nonumber\]

\[\ce{Ga^3+} (aq) + \ce{3e-} → \ce{Ga} (s)\nonumber\]

This cell was 100% efficient and produces 17.26 kilogram of gallium metal in exactly 1260 minutes.

- Write a balanced equation for the reaction;
__Solution:__

- Calculate the quantity of gallium formed in moles;
__Solution:__

- Calculate the number of electron moles passed through the cell in forming gallium metal?
__Solution:__

- Calculate the charge (C) carried by the electrons that passed through the cell;
__Solution:__

- Calculate the currents (A) using the charge (C) passed through the cell in a period of 1260 minutes.
__Solution:__

** **

__Questions in class Group 2__:

Given the cell: *Cu(s) | Cu ^{2+} (0.50 M) || I^{-} (0.30 M) | I_{3}^{-} (0.15 M) | Pt*; and information below,

\[\ce{I3- + 2e- → 3I-} \hspace{30px} E^0 = \mathrm{0.535\: V}\nonumber\]

\[\ce{Cu^2+} + \ce{2e-} → \ce{Cu}(s) \hspace{30px} E^0 = \mathrm{0.339\: V}\nonumber\]

- What does
**|**stands for in the notation?__Solution:__

- What does
**||**stands for in the notation?__Solution:__

- Write the ionic half-cell reactions for the cathode and anode.
__Solution:__

- Write the overall ionic cell equation for the reaction and its E
^{0 }value.__Solution:__

- Calculate the voltage of the cell at 298K.
__Solution:__

__Questions in class Group 3__:

As researchers look for alternative fuels, the production of H_{2} is being extensively studied. The following example shows a reduction reaction that can be used to produce H_{2} under basic conditions:

\[\ce{2H2O} (l) + \ce{2e-} → \ce{H2} (g) + \ce{2OH-} (aq) \hspace{30px} E^o = \mathrm{- 0.83\: V}\nonumber\]

- What is the Nernst expression for this half-cell?
__Solution:__

- Write the Nernst expression so that the influence of the [OH
^{-}] appears as a function of pH assuming that activity of H_{2}= 4.0 x 10^{-5}.__Solution:__

- What is the value of E for this reaction when the pH of the solution is maintained at 7.00?
__Solution:__

- What is the value of E for this reaction when the pH of the solution is maintained at 10.00?
__Solution:__

__Questions in class Group 4__:

Use the following cell between Cu/Zn and the electrode potentials for the half cells to answer the question:

\[\ce{Zn^2+}(aq) + \ce{2e-} → \ce{Zn}(s)\: \textrm{------------------------} \: E^o = \mathrm{-0.76\:V}\nonumber\]

\[\ce{Cu^2+}(aq) + \ce{2e-} → \ce{Cu}(s)\: \textrm{------------------------} \: E^o = \mathrm{0.34\:V}\nonumber\]

- What is the overall cell equation for the reaction?
__Solution:__

- What is the overall cell potential for the reaction?
__Solution:__

- Calculate the Gibbs free energy (Δ
*G*) for the reaction in kilojoules?__Solution:__

- What is the equilibrium constant for the reaction?
__Solution:__

- What is the role of the salt bridge?
__Solution:__

__Questions in class Group 5__:

Using standard reduction potentials, calculate the potential for the following cell at 298 K:

\[\mathrm{Zn \mid Zn^{2+}(\textit{aq}),\: 0.1\: M \mid\mid Cu^{2+}(\textit{aq}),\: 1.0\: M \mid Cu}\nonumber\]

- Why is it not possible to generate electrical current by simply placing strips of copper and zinc in a beaker containing an aqueous solution of Zn(NO
_{3})_{2}and Cu(NO_{3})_{2}?__Solution:__

- Using shorthand notation, Zn | Zn
^{2+}(aq), 0.1 M || Cu^{2+}(aq), 1.0 M | Cu, write the half equation for anode, cathode and the overall reaction equation:__Solution:__

- What is the expression for Q and calculate the Q value?
__Solution:__

- Calculate the E
^{o}value for the reaction?__Solution:__

- What is the overall potential of the cell?
__Solution:__

- Is the reaction spontaneous or not?
__Solution:__

## Contributors and Attributions

- Grant Wangila, University of Arkansas Pine Bluff (wangilag@uapb.edu)
- Sourced from the Analytical Sciences Digital Library