CHEM 1128Q Lecture Notes - Lecture 9: Galvanic Cell, Redox, Electric Potential

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In a circuit, the flow of charge is generated by an electrical potential difference between two points in the circuit. Electrical potential is the ability of the electric field to do work on the charge. In a galvanic cell, this electrical potential difference is called the cell"s potential. The cell potential in a galvanic cell results from the difference in the electrical potentials at each electrode. The driving force of an electrochemical reaction is measured by the standard cell potential. Depends on the nature of the redox reaction and of the concentration of the species involved. Temperature: held constant, usually 25 degrees celsius. Units of e standard are v or mv. The standard e of the cell is equal to the standard e of the cathode minus the standard e of the anode. Any redox reaction can be split into two half reactions. The standard e for a half reaction cannot be measured.

Related Questions

I am studying an iron chloride redox battery cell (not a flow battery). It is constructed with an iron negative electrode, a carbon positive electrode with an iron (II) chloride electrolyte. The electrodes are physically isolated from each other in a beaker of iron(II) chloride. This cell can be charged and discharged and will run a motor. The output of the battery is 1.2V and can deliver 0.2A.

Unlike most redox batteries, the positive carbon electrode does not take part in the chemistry. The positive carbon electrode, as far as I understand it, simply provides a reaction site for the positive half reaction.

We are told that the half reaction at the positive electrode converts Fe2+ ions to Fe3+ ions (charge) and back again (discharge).

The half reaction at the negative electrode converts Fe2+ ions to metallic iron (charge) and back again (discharge).

The task is to fully describe the iron chloride redox chemistry at both electrodes during charge and discharge.

Where I am:

I am starting with an assumption that during charge and the application of external electrical energy, the dissociated ions in the electrolyte, Fe2+ and Cl-, must migrate to the oppositely charged electrodes. This would be independent of the electrode material and determined simply by the external EMF applied to the battery during charge. The Fe2+ ions must migrate to the negative electrode while the Cl- ions must migrate to the positive electrode.

The result is that the negative electrode will be surrounded by Fe2+ ions and the positive electrode will be surrounded by Cl- ions.

It seems clear that Fe2+ ions will react with incoming electrons, at the negative electrode, to form metallic iron.

However, I do not understand how any reaction can occur at the positive electrode with respect to any variety of "iron chloride". There should be no Fe2+ ions at the positive electrode. Therefore, no reaction involving Fe2+ ions can occur at the positive electrode.

The positive electrode is carbon surrounded by Cl- ions. It is under the influence of a positive electric potential by the external charge source. The only possible reaction I can see occurring is that the Cl- ions may liberate an electron at the positive electrode to form chlorine gas. However, this is not observed under the charging voltage applied, at 25^OC and 1 atm. There are no bubbles or chlorine smell.

I have found various references such as:

https://courses.lumenlearning.com/boundless-chemistry/chapter/electrolysis/

That state:

Oxidation of ions or neutral molecules occurs at the anode, and reduction of ions or neutral molecules occurs at the cathode. For example, it is possible to oxidize ferrous ions to ferric ions at the anode:

${Fe}^{2 +} (aq) \to {Fe}^{3 +} (aq) + e^{-}$

This is exactly what we are told is happening at the positive electrode but I am unable to see how this can occur.

How can a positive ion be in proximity to the positive electrode in order to react with it?? How can any reaction involving Fe2+ ions happen at the positive electrode??

olivemule889

experment1 Electron Transfer (Redox) Reactions

1. Comparison of Br2/Brâ and Fe3+/Fe2+ Couples

Oxidant-reductant pairs from these couples will he tested, i.e. Br2 with the reductant Fe2+ and

Brâ with the oxidant Fe3+. A reaction will occur in only one of these cases because of the

difference in electrode potential between the two couples.

To test whether reaction has occurred, SCNâ will be used as a probe for Fe3+, with which it

gives a red complex. Note that the air rapidly oxidises Fe2+, so that a faint pink colour is often

obtained when SCNâ is added to Fe2+.

Solutions containing Fe2+ or Fe3+ tend to decompose (although in different ways); both are

stabilised by H+. The iron(II) solution provided is freshly prepared 0.1 M (NH4)2Fe(SO4)2 / 2

M H2SO4.

To 5 drops of the 0.1 M iron(II) solution in a semi-micro tube add 10 drops of 0.1 M KBr and

stir. Place 1 drop of the mixture on a white tile, and test for Fe3+ by adding 1 drop 0.1 M

KSCN. A strong red colour indicates Fe3+.

Repeat using 5 drops of the 0.1 M iron(III) solution and adding 10 drops of 0.1 M KBr.

Repeat using 5 drops of the 0.1 M iron(II) solution and adding 10 drops of 0.1 M Br2.

Repeat using 5 drops of the 0.1 M iron(III) solution and adding 10 drops of 0.1 M Br2.

2.Use of the Br2/Brâ and Fe3+/Fe2+ Couples in an Electrochemical Cell

You are supplied with a pair of carbon rods joined by copper wire. Place two 100 mL beakers

side by side.

In one put 50 mL of the iron(II) solution and 1 mL of 0.1 M KSCN solution. In the other

beaker put 25 mL of 0.2 M Br2 solution and 25 mL of 4 M HCl.

Correct any difference in level between the solutions in the two beakers by adding water to

one. Fold an 11 cm filter paper into a strip about 1 cm wide. Bend the strip in the middle,

moisten with a little 4 M HCl, then insert the strip so that it links the two solutions. Insert one

carbon rod into each solution, then allow to stand undisturbed for 30 minutes while

proceeding with Experiment 3. Carefully observe the apparatus at intervals and record any

change that occurs.

3. Comparison of the Fe3+/Fe2+ and Sn4+/Sn2+ Couples

In this experiment the additional sources of ions are:

tin(IV) ion 0.1 M SnCl4/1 M HCl

tin(II) ion 0.1 M SnCl2/1 M HCl

(Both Sn2+ and Sn4+ solutions are stabilised by H+.)

To 5 drops of the 0.1 M iron(II) solution in a semi-micro tube add 10 drops of the appropriate

0.1 M tin solution and stir. Place 1 drop of the mixture on a white tile, and test for Fe3+.

Record the results.

Similarly, test the 0.1 M iron(III) solution with an excess of the appropriate 0.1 M tin

solution.

You have now enough information to state whether the standard electrode potential of the

Fe3+/Fe2+ couple is more positive or less positive than that of the Sn4+/Sn2+ couple.

please answer following questions

1.Comparison of Br₂/Br^â and Fe³⁺/Fe²⁺ couples

Complete the following table.

Solutions	Observation	Fe³⁺ present (Yes/No)	Did a reaction occur?
Fe(II) and KBr
Fe(III) and KBr
Fe(II) and Br₂
Fe(III) and Br₂

2.Br₂/Br^â and Fe³⁺/Fe²⁺ couples in an Electrochemical Cell

a)What observations did you make over 3 minutes? Did Br₂ or Fe²⁺ react? Write the ion/electron half- equations for this reaction and the complementary reduction. Write the overall redox reaction.

b)Draw/scan/insert a circuit diagram showing the half-reactions that occur in each beaker, the direction of electron flow and the direction of flow of ions.

3.Fe³⁺/Fe²⁺ and Sn⁴⁺/Sn²⁺ couples

Complete the following table.

Solutions	Observation	Fe³⁺ present (Yes/No)	Did a reaction occur?
Fe(II) and Sn(IV)
Fe(II) and Sn(II)

CHEM 1128Q Lecture Notes - Lecture 9: Galvanic Cell, Redox, Electric Potential

CHEM 1128Q Full Course Notes

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