MODULE 5: Redox and Electrode Potentials Flashcards
Balancing equations by oxidation numbers:
- Assign oxidation numbers to each species in an equation
- Balance species that have a changed oxidation number
- Balance any remaining atoms
Reagents/products likely added to an equation when predicting:
H+, OH- or H2O- usually to balance atoms or charges on each side
Two examples of redox titrations:
Fe(2+)/MnO4(-) and I2/S2O3(2-)
Procedure for MnO4(-) titration:
- Add KMnO4 to burette
- Add sulfuric acid to conical flask to reduce KMnO4 to MnO4(-)
- Add MnO4(-) in small increments- it’ll decolourise as it is added
- No indicator is needed as the reaction is self indicating, the end point is when a permanent pink colour remains
- Repeat titration until 2 concordant results are obtained
Use of MnO4(-) for percentage purity of Fe(II) compounds:
- Prepare 250 cm3 standard solution of impure iron sulphate solution to volumetric flask
- Using a pipette, measure 25 cm3 of the impure solution and to a conical flask, add 10 cm3 sulphuric acid to the solution
- Titrate against MnO4(-) solution
- Calculate moles of MnO4(-) that reacted from mean titre
- Calculate moles of Fe2+ that reacted in 25 cm3
- Multiply by 10 to get moles in the whole 250 cm3 solution
- Calculate mass of unknown sample
- Calculate percentage purity using mass of solution reacted/mass of impure sample x 100
Uses of iodine/thiosulfate:
- ClO- concentration in bleach
- Cu2+ concentration in copper (II) compunds
- Cu concentration in copper (II) alloys
What indicator is used in an iodine thiosulfate titration?
Starch (colourless to black)
Standard electrode potential definition:
The tendency to be reduced and gain electrons compared to hydrogen half cell under standard conditions
Measuring a standard electrode potential:
- Half cell must be attached to a hydrogen half cell
- Connected via a salt bridge in each solution to allow the movement of ions (soaked in a solution that wil now react with either reactant)
- Connected by a wire between the electrodes to allow the movement of electrons
Metal/metal half-cells:
An electrode made of a metal is placed into a solution of the same metal (but aqueous.) Where the ions come into contact with the electrode, an equilibrium is established:
X 2+ + 2e- X(s)
Ion/ion half-cells:
A solution made containing ions of the same species in different oxidation states. An inert electrode of platinum is used to establish this equilibirum:
X3+ + e- –> X2+
Meanings of standard electrode potential values:
More positive: more likely to be reduced
More negative: more likely to be oxidised
Cell potential equation:
Cell potential = positive electrode standard pot. - negative electrode standard pot.
Predicting feasibility of a reaction:
Using standard electrode potentials to calculate the standard cell potential (positive = feasible)
Limitations of predicting feasibility of a reaction using standard electrode potentials:
- Activation energy may be too large for the reaction to start, causing a slow rate of reaction/no reaction at all
- Standard electrode potential is measured at 1 mol dm-3, so if the solution is less concentrated it’ll be less feasible, but if it is more concentrated it’ll be more feasible
- Similar applies to other conditions (temperature and pressure)
