Redox

Question 1

Define Standard electrode potential

Answer 1

The standard electrode potential, E*, of a half- cell is the EMF generated when it is connected to the standard hydrogen electrode by an external circuit and a salt bridge, measured under standard conditions.

Question 2

Explain the function of a salt bridge in voltalc cells

Answer 2

The purpose of the salt bridge is to keep the solutions electrically neutral and allow the free flow of ions from one cell to another. Without the salt bridge, positive and negative charges will build up around the electrodes causing the reaction to stop.

Question 3

Outline the difference between an electrolytic cell and a voltaic cell

Answer 3

Electrolyic cell copnverts electrical energy to chemical energy where as voltalic cells do the oppisite
Electroyliuc cells require an energy supply, while voltalic do not

Electrolytic cells have non-spontaenous redox reactions where as volatlic have spontaneous redox reaction

In an elecotryltic cell the the cathod is negative and anonde posiitve where as in volatlic it is opposite

Electrolytic have one solution, where as volatlic have two
Elecytrolytic has no salt bridge, and voltalic must have a salt bridge
Electoryltic, oxidation occurs at positive elctorde na dopposite for volatlic

Question 4

Explain why solid sodium chloride does not conduct electricity but its molten form does.

Answer 4

In solid state the ions are in a lattice and thus not mobile, however, in molten state the ions are mobile ie free to move and carry a charge

Question 5

Why may Aluminium be preferred to Iron?

Answer 5

Al does not corrode, rust,
Al is less dense/ better conductor, more malleable

Question 6

What is Galvanising and why may it be done?

Answer 6

Galvanisation is a process wherein a thin layer of zinc metal gets deposited on iron objects. This is done to prevent rusting of iron by protecting it to come in contact with air and moisture. Zinc metal, being more reactive, reacts with air to form hard layer of zinc oxide, which prevents air from passing through it.

Question 7

Outline sacrificial protection

Answer 7

Sacrificial protection is a corrosion protection method in which a more electrochemically active metal is electrically attached to a less active metal. The highly active metal donates electrons to replace those which may have been lost during oxidation of the protected metal

Question 8

What is Voltaic cell? List the components

Answer 8

A voltaic cell consists of two different half-cells, connected together to enable the electrons transferred during the redox reaction to produce energy in the form of electricity. The cells are connected by an external wire and by a salt bridge, which allows the free movement of ions.

Question 9

What is an Electrolytic cell? List the components

Answer 9

In an electrolytic cell electricity is passed through an electrolyte and electrical energy is converted into chemical energy. An electrolyte is a substance which does not conduct electricity when solid, but does conduct electricity when molten or in aqueous solution and is chemically decomposed in the process

Question 10

Formula for Charge

Answer 10

Current (A) x Time (s) = Charge

Question 11

What is a Faraday (F)?

Answer 11

A faraday is equal to 96500 Columbs which is the charge carried by one mole of electrons

Question 12

What factors effect ion discharge at an electrode in Electrolysis?

Answer 12

-position in electrochemical series (the lower the more readily it is reduced)

-The relative concentration of the ions in the electrolyte

-The nature of the electrode, if not inert will increase the conc of ion

Question 13

Where do Oxidation and Reduction occur in an electrolytic cell?

Answer 13

Oxidation occurs at the positive anode and reduction occurs at the negative cathode

Question 14

Where do Oxidation and Reduction occur in a voltaic cell?

Answer 14

Oxidation occurs at the negative anode and reduction occurs at the positive cathode

Question 15

What does the E* value of a half cell indicate?

Answer 15

The higher the E* value (MORE POSITIVE) the stronger the Oxidising AGENT

The lower the E* value (MORE NEGATIVE) the stronger the Reducing AGENT

Question 16

Formula for Free Gibs using Faraday

Answer 16

G = -n(no. of moles of electrons transferred) x F (96500 C/mol) x E* (J)

Question 17

How do we indicate the spontaneity using the E* of a cell?

Answer 17

If the E* of a cell is positive then G is negative and thus reaction is spontaneous
If the E* of a cell is negative then G is positive and thus reaction is not spontaneous
If E* is 0 then G is 0 and thus the reaction is at equilibrium

Question 18

Formula of E* of cell?

Answer 18

E* of half cell where reduction occurs - E* Half cell where oxidation occurs

(More positive - less positive)

Question 19

What direction do electrons flow in voltaic cells?

Answer 19

Electrons flow from the anode to the cathode, from the half cell with the smaller (more negative) E* value to the larger (more positive) E* value via the external circuit

Question 20

Which half cell in a voltaic cell becomes the cathode and which the anode?

Answer 20

The half-cell with the higher E* value is the cathode (+), and
the half-cell with the lower E* value is the anode (–).

Question 21

What are standard conditions for a standard half cell?

Answer 21

-all solutions must have a concentration of 1.0 mol dm–3;
-all gases must be at a pressure of 100 kPa;
-all substances used must be pure;
-temperature is 298 K;
-if the half-cell does not include a solid metal, platinum is used as the electrode.

Question 22

Define Electromotive force

Answer 22

The electromotive force (EMF) of a cell is the greatest potential difference that it can generate. It is measured in volts.

Question 23

How do voltaic cells generate a flow of current?

Answer 23

A voltaic cell converts the energy released from a spontaneous, exothermic reaction into electrical energy.

Question 24

What are the rules for the cell diagram convention?

Answer 24

-a single vertical line represents a phase boundary such as that between a solid electrode and an aqueous solution within a half-cell;
-a double vertical line represents the salt bridge;
-the aqueous solutions of each electrode are placed next to the salt bridge;
-the anode is generally put on the left and the cathode on the right, so electrons flow from left to right;
-spectator ions are usually omitted from the diagram;
-if a half-cell includes two ions, they are separated by a comma because they are in the same phase

Question 25

How does the reactivity of a metal affect electrode potential?

Answer 25

In general the more reactive a metal, the more negative its electrode potential in its half-cell.

Question 26

Describe the Wrinkler method

Answer 26

The principle is based on a sequence of redox reactions as follows.

1 The dissolved oxygen, O2(g), in the water is ‘fixed’ by the addition of a manganese(II) salt such as MnSO4. Reaction of this salt with O2 in basic solution causes oxidation of Mn(II) to higher oxidation states, such as Mn(IV):
2Mn(II)(aq) + O2(g) + 4OH–(aq) → 2MnO2(s) + 2H2O(l) +2 +4

2 Acidified iodide ions, I–, are added to the solution, and are oxidized by the Mn(IV) to I2

MnO2(s) + 2I–(aq) + 4H+(aq) → Mn2+(aq) + I2(aq) + 2H2O(l) +4 +2

3 The iodine produced is then titrated with sodium thiosulfate, as described earlier: 2S2O32–(aq) + I2(aq) → 2I–(aq) + S4O62–(aq)
So we can see that in the overall sequence, for every 1 mole of O2 in the water, 4 mol of S2O32– are used.

Question 27

What is BOD?

Answer 27

Biological Oxygen Demand is the quantity of oxygen needed to oxidize organic matter in a sample of water over a five-day period at a specified temperature.

Question 28

How do the reactivity of NON-metals correlate to its reducing ability?

Answer 28

More reactive non- metals are stronger oxidizing AGENTS than less reactive non-metals.

Question 29

Summary of steps in writing redox equations

Answer 29

Summary of steps in writing redox equations.
1 Assign oxidation states to determine which atoms are being oxidized and which are being reduced.
2 Write half-equations for oxidation and reduction as follows:
(a) balance the atoms other than H and O;
(b) balance each half-equation for O by adding H2O as needed;
(c) balance each half-equation for H by adding H+ as needed;
(d) balance each half-equation for charge by adding electrons to the sides with the more positive charge.
(e) check that each half-equation is balanced for atoms and for charge.
3 Equalize the number of electrons in the two half-equations by multiplying
each appropriately.
4 Add the two half-equations together, cancelling out anything that is the same on both sides.

Question 30

What factor affects the qualntity of products discharged at an electrode in electrolysis?

Answer 30

• The more electrons flowing through the circuit the more products formed
• the charge of the ion (Na+ will need half as many electrons as Mg2+, thus Na will discharge faster)