SNS - General Chemistry - Redox Reactions and Electrochemistry

Question 1

Redox Reactions

Oxidation Numbers

Monoatomic ions

Answer 1

The oxidation number of any monoatomic ion (an ion of only one atom) is equal to the charge on the ion

Question 2

Redox Reactions

Oxidation Numbers

Free elements

Answer 2

Free elements (eg Na, H2, P4, N2) have oxidation numbers of 0

Question 3

Redox Reactions Oxidation Numbers Oxygen

Answer 3

The oxidation numbe of oxygenis -2 except in peroxides, where it is -1

Question 4

Redox Reactions Oxidation Numbers Hydrogen

Answer 4

+1 except whenit occurs after a metal ion, where it is -1

Question 5

Redox Reactions Oxidation Numbers Fluorine

Answer 5

-1

Question 6

Redox Reactions Oxidation Numbers Group IA Elements

Answer 6

+1

Question 7

Redox Reactions Oxidation Numbers Group IIA Elements

Answer 7

+2

Question 8

Redox Reactions Oxidation Numbers Compounds

Answer 8

The sum of oxidation numbers in a compound is 0 except for polyatomic ions where the sum is equal to the charge of the ion

Question 9

Redox Reactions Balancing Cr₂O₇²- (aq) + Cl- (aq) → Cr³+ (aq) + Cl₂ (g) Involves an acidic solution

Answer 9

Oxidation: Cr₂O₇²- (aq) → Cr³+ (aq) Reduction: Cl- (aq) → Cl₂ (g) Since this invlve an acidic solution, H+ are used to balance the half equations: Cr₂O₇²- (aq) + 14H+ (aq) → Cr³+ (aq) + 7H₂O (l) Cl- (aq) → Cl₂ (g) Equalising the charges: Cr₂O₇²- (aq) + 14H+ (aq) + 6e- → Cr³+ (aq) + 7H₂O (l) Cl- (aq) → Cl₂ (g) +2e- The second half equation must be multiplied by 3 to equalise the lectrons on both sides and the two balanced half equations can be added: Cr₂O₇²- (aq) + 14H+ (aq) + 6Cl- (aq) → Cr³+ (aq) + 7H₂O (l) +3Cl₂ (g)

Question 10

Electrochemical Concepts Electrolysis

Answer 10

Non-spontaneous reactions that are driven by an outside source of electrical energy. Occur in electrolytic cells Redox reactions take place at the electrodes - oxidation at the anode, reduction at the cathode

Question 11

Electrochemical Concepts

Faraday’s Law

Answer 11

Theorises that the amount of chemical charge induced in an electrolytic cell is directly proportional to the number of moles of electrons exchanged during a redox reaction.

For a reaction which involves the transfer of n electrons per atom, Mⁿ⁺ + ne- → M (s) one mole of M will be produced for every n moles e- supplied. The number of moles e- needed to produce a given amount of M can now be related to a measurable electrical property. One electron carries a charge of 1.6 x 10-19 C. Thus the charge carried by one mole of electrons = (1.6 x 10-19)(avogadro number, 6 x 10²³) = 96487 C/mol e- = 1 Faraday

Question 12

Electrochemical Concepts

Galvanic Cell

Answer 12

Has negative ∆G, therefore spontaneous. Reactions supply energy and are used to do work - harnessed by placing the two electrodes in separate half-cells connected by an apparatus that permits the flow of electrons - a salt bridge which permits ion exchange

Question 13

Electrochemical Concepts

Electrode Potential

Answer 13

The cell’s potential power is dependent on the spontaneity of the oxidation reduction reaction. The greater the spontaneity, the more electrons produced per starting molecule

Question 14

Electrochemical Concepts Electromotive Force

Answer 14

The driving force that pushes the electrons through the circuit. It is the potential difference between the electrode of the cell A positive emf value indicates a spontaneous reaction, a negative one indicates a non-spontaneous reaction

Question 15

Electrochemical Concepts

Standard Reduction Potential, E⁰

Answer 15

Reduction potential measured under standard conditions.

Relative reactivities of different half-cells can be compared to predict the direction of electron flow. A higher E⁰ means a greater tendency for reduction to occur

Question 16

Electrochemical Concepts

Standard Electrode Potential,

E⁰ Conditions

Answer 16

Standard Conditions:

1. Temp - 20⁰C
2. Concentration of each ion - 1M
3. Partial pressure of each gas - 1atm
4. Metals - in their pure states

Question 17

Assign oxidation numbers to the atoms in the following reaction in order to determine the oxidised and reduced species and the oxidising and reducing agents:

SnCl2 + PbCl4 → SnCl4 + PbCl2

Answer 17

SnCl2 : Sn = +2, Cl = -1

PbCl4 : Pb = +4, Cl = -1

SnCl4 : Sn = +4, Cl = -1

PbCl2 : Pb = +2, Cl = -1

Sn = oxidised = reducing agent

Pb = reduced = oxidising agent

Question 18

Balance the redox equation:

MnO₄^- + I₂ + Mn²⁺

Answer 18

1. Balance with H+ and e-:

MnO₄^- + 8H+ 5e- → Mn²⁺ + 4H2O

2I^- → I₂ + 2e-

2. Combine:

2MnO4- + 16H+ 10I^- → 2Mn2+ + 8H2O + 5I₂

Question 19

Electrochemical Cells

Answer 19

Contained systems in which a redox reaction occurs

Two types: Galvanic and Electrolytic in which spontaneous and non-spontaneous reactions occur respectively. Both contain electrodes at which oxidation and reduction occur (anode = oxidation, cathode = reduction)

Question 20

Electrochemical Concepts

Daniell Cell

Answer 20

Galvanic cell

A zinc bar placed in aqueous ZnSO4 solution and a copper bar is placed in an aqueous CuSO4 solution. The anode is the zinc bar where Zn (s) is oxidised to Zn2+ (aq). The cathode is the copper bar, where Cu2+ (aq) is reduced to Cu (s). If the two half-cells were not separated, the Cu2+ would react directly with the zinc bar and no useful electrical work would be done. To complete the circuit, the half-cells must be connected - without this the electrons from the zinc oxidation would not be able to get to the Cu2+

Question 21

Electrochemical Concepts

Salt Bridge

Answer 21

Connects half cells in a Galvanic cell

Contains an inert electrolyte such as KCl or NH4NO3 whose ions don’t react with the electrodes or the ions in solution. At the same time, its anions (eg Cl-) diffuse from the salt bridge into the solution within the half-cell containing the anode to balance the charge from the newly oxidised species. The cations (eg K+) flow into the solution of the cathode half-cell to balance the charge created by the newly reduced species

As this flow proceeds, the salt bridge is depleted acounting for the relatively short lifespan of the cell

Question 22

Electrochemical Concepts

Cell Diagram

Answer 22

Shorthand notation for representing the reactions in an electrochemical cell.

anode | anode solution || cathode solution | cathode

where || indicates a salt bridge or some other form of barrier

For example, the Daniell reaction:

Zn (s) | Zn²⁺(xM SO4^2-) || Cu2+(yM SO4^2-) | Cu (s)

Question 23

Electrochemical Concepts

Electrolytic Cells

Answer 23

Non-spontaneous - electrical energy is required to induce reaction. Oxidation and reduction half-reactions are usually placed in a single container

eg cell in which molten NaCl electrolysed to form Cl2 (g) and Na (l). Na+ migrate towards the cathode where they are reduced to Na (l). Cl- migrate towards the anode where they are oxidised to Cl2 (g). Used in sodium and chlorine production

Question 24

Electrochemical Concepts

Electrochemical Cells

Electrode Charge Designations

Answer 24

1. Galvanic cell - anode considered to be negative because the spontaneous oxidation reaction that takes place there is the original source of the cell’s negative charge (ie of electrons).
2. Electrolytic cell - anode is considered positive since it is attached to the positive pole of the battery and so it attracts anions from the solution

In both, however, oxidation takes place at the anode and reduction at the cathode

Question 25

Electrochemical Concepts

Reduction Potentials

Answer 25

The tendency of a species to acquire electrons and be reduced

Used to determine the species oxidised or reduced

Each species has its own reduction potential, the more positive, the greater the tendency to be reduced

Question 26

Given the following half-reactions and Eº values, determine the species to be oxidised and that to be reduced:

Ag+ + e- → Ag (s) Eº = +0.80V

Tl+ + e- → Tl (s) Eº = -0.34V

Answer 26

Ag reduced, Tl oxidised:

Ag+ + Tl → Ag + Tl+

Question 27

Electrochemical Concepts

Electromotive Force

Answer 27

The difference in potential between two half-cells

Standard reduction potentials are used to calculate the standard electromotive force (emf) of a reaction - determined by adding the standard reduction potential of the reduced species and the standard oxidisation potential of the oxidised species

Question 28

Given that the standard reduction potentials for Sm³⁺ and [RhCl6]^3- are -2.41V and +0.44V respectively, calculate the emf of the reaction:

Sm³⁺ + Rh +6Cl^- → Sm (s) + [RhCl6]^3-

Answer 28

= 2.41 -0.44 = -2.85V

Therefore this reaction would procedd spontaneously to the left, in which case Sm would be oxidised and [RhCl6]3- reduced

Question 29

Electrochemical Concepts

Emf and Gibbs Free Energy

Answer 29

The thermodynamic criterion for determining the spontaneity of a reaction is Gibbs free energy - the maximum amount of useful work produced by a chemical reaction

In an electrochemical cell the work done is dependent on the number of coulombs and the energy available. Thus ∆G and emf are related as follows:

∆G = -nFE_cell

where n is the number of moles electrons exchanged, F is faraday’s constant, and Ecell is the emf in the cell

nb if Faraday’s constant is expressed in C (J/V) then ∆G must be expressed in J not kJ

Question 30

Emf and Gibbs Free Energy

Answer 30

If a reaction takes place under standard conditions, the equation becomes:

∆G = -nFEºcell

Question 31

Nernst Equation

Answer 31

Used to determine the effect of species concentration on emf

Ecell = Eºcell - (RT/nF)(ln Q)

where Q is the reaction quotient for a given reaction. For example in the reaction aA + bB → cC +dD,

Q = ([C]^c [D]^d) / ([A]a [B]^b)

Question 32

Emf and Keq

Answer 32

For reactants in solution, ∆G⁰ can be determined in another manner.

∆G⁰ = -RT ln Keq

where R = gas constant (8.314 J/K · mol), T = temp in K, Keq = equilibrium constant for the reaction

∆G⁰ = -nFEºcell = -RT ln Keq

If the values of n, T and Keq are known, Eºcell for the redox reaction can be calculated