Topic 14: Further Redox

Question 1

what is meant by the term ‘standard electrode potential’, EӨ?

Answer 1

The potential difference of a cell when the electrode is connected to the standard hydrogen electrode under standard conditions (298K, 100 kPa, 1.00 mol dm-3)

Question 2

Answer 2

Zn or Fe, as they have electrode potentials LESS than -0.28, which is the elctrode potential to form Co metal

Question 3

What might prevent a reducing agent from being as effective as the electrode potentials might seem to suggest?

Answer 3

Rate too slow / activation energy too high. The reaction is under non-standard conditions

Question 4

Answer 4

Zn can reduce Fe3+ to Fe2+: +0.77 + 0.76 = +1.53

Zn can reduce Fe2+ to Fe: -0.44 + 0.76 = +0.32

Both these values are positive so the reaction is feasible

Question 5

Explain the need for a standard reference electrode in electrochemical measurements.

Answer 5

A standard reference electrode is so that comparisons can be made between any half cells.

Can only measure a potential difference.

Cannot measure the potential difference between a metal and a solution of its ions

Question 6

Answer 6

Fe3+ and Fe2+ has a more positive electrode potential than I2 and I-. This means I- would reduce Fe3+ / Fe3+ would oxidise I-

OR 2Fe3+ + 2I- → 2Fe2+ + I2 has an E value of 0.23V. This means I- would reduce Fe3+ / Fe3+ would oxidise I-

Question 7

When a metal is placed in a solution of its ions, the electrical potential set up between the metal and the solution cannot be measured without a reference electrode. Explain why.

Answer 7

Introducing another metal wire would set up its own potential difference.

You can only measure a potential difference, so you need the ref electrode.

Question 8

The amount of of sodium chlorate in a sample is found by reacting it w excess acidified potassium iodide: ClO– +2I– +2H+ →I2 +Cl– +H2O.

I2 produced is then titrated with sodium thiosulphate solution.

10.0 cm3 of bleach was pipetted into a 250 cm3 flask and made up to the mark with distilled water. A 25cm3 portion of the solution was added to excess acidified potassium iodide in a conical flask.

This mixture was titrated with 0.100 mol dm–3 sodium thiosulphate solution, using starch indicator. The mean titre was 12.5cm3.

The equation for the reaction between iodine and thiosulphate ions: 2S2O32– + I2 → S4O62– + 2I–

Calculate the moles of chlorate(I) ions in 1.00 dm3 of the original bleach.

Answer 8

Question 9

The standard electrode potentials are shown below.

Pb2+(aq) + 2e —> Pb(s) (–0.13V)

PbO2(s) + 4H+(aq) + 2e —> Pb2+(aq) + 2H2O(l) (+1.46V)

Calculate the standard e.m.f. of the cell.

Answer 9

Eθ = + 1.46 – ( – 0.13) = + 1.59 (V)

Question 10

State why E values cannot predict that a reaction will occur, only that it is possible.

Answer 10

High Ea so slow reaction / reactants are kinetically stable
IGNORE any mention of non-standard conditions

Question 11

Answer 11

A= salt bridge containing saturated solution of NaCl

B= Pt electrode

C= solution containing V2+ and V3+ ions

Question 12

Answer 12

Question 13

Using half eqns, write an overall eqn for the reaction which occurs in an alkaline and acidic hydrogen-oxygen fuel cell.

Answer 13

O2(g) + 2H2O(l) + 4e- → 4OH-(aq) and 4OH-(aq) + 2H2 → 4H20 + 4e-

ACID: H2 → 2H+ + 2e- and O2 + 4H+ + 4e- → 2H20

Add the above half-equations for each acid and alkali, cancelling 4e— to get: 2H2(g) + O2(g) → 2H2O(l)

Question 14

The acid/ alkali in a hydrogen fuel cell is for what?

Answer 14

The acid/ alkali in a hydrogen fuel cell is to allow the movement of ions between electrodes.