Redox

Question 1

Single Replacement Reaction

Answer 1

When one element replaces another element in a compound

A metal will replace a less reactive metal ion

Element + Compound —> New Element + New Compound

Question 2

Single Replacement Reaction Example:
Mg (s) + CuCl(2) (aq)
(Balanced Equation, Complete Ionic Equation, Net Ionic Equation)

Answer 2

Balanced Equation:
Mg (s) + CuCl(2) (aq) —> MgCl(2) (aq) + Cu (s)

Complete Ionic Equation:
Mg (s) + Cu(2+) (aq) + 2Cl(-) (aq) —> Mg(2+) (aq) + 2Cl(-) (aq) + Cu (s)

Net Ionic Equation:
Mg (s) + Cu(2+) (aq) —> Mg(2+) (aq) + Cu (s)

Question 3

Redox Reaction

Answer 3

A redox reaction involves the transfer of electrons from one chemical species to another

Question 4

Reduction

Answer 4

When a chemical species gains electrons

Question 5

Oxidation

Answer 5

When a chemical species loses electrons

Question 6

Half Equation

Answer 6

A half equation represents either an oxidation or reduction half of a redox equation including a loss or gain of electrons

Question 7

Reduction Half Equation Example:
O(2) (g)

Answer 7

O(2) (g) + 4e(-) —> 2O(2-) (s)

Question 8

Oxidation Half Equation: Mg(2+) (s)

Answer 8

Mg (s) —> Mg(2+) (s) + 2e(-)
To balance with reduction half equation:
2Mg (s) —> 2Mg(2+) (s) + 4e(-)

Question 9

Complete Half Equation Example:
2Li (s) + Br(2) (l) —> 2LiBr (s)

Answer 9

Oxidation Half Equation:
Li(s) —> Li+ (s) + e(-)

Reduction Half Equation:
Br(2) (l) + 2e(-) —> 2Br(-) (s)

To Balance Half Equations:
2Li(s) —> 2Li+ (s) + 2e(-)

Full Equation:
Br(2) (l) + 2Li(s) —>2BrLi (s)

Question 10

Oxidising Agent or Oxidant

Answer 10

A chemical species that causes another chemical species to be oxidised

Question 11

Reducing Agent or Reagent

Answer 11

A chemical species that causes another chemical species to be reduced

Question 12

SEP Table

Answer 12

Metals at the top are most reactive , most likely to oxidise, strongest reducing agents

Question 13

Reactivity for a Redox Reaction to Occur

Answer 13

A more reactive metal will be oxidised by a less reactive metal cation (the metal donates its electrons and the cation receives the electrons). Therefore, for a spontaneous redox reaction to occur, the metal ions must be less reactive than the solid metal

Question 14

Example of Reactivity for a Redox Reaction to Occur:
Copper wire is placed in silver nitrate solution

Answer 14

Cu is more reactive and therefore more likely to oxidise than Ag, so the redox reaction will take place.

Reduction Half Equation:
Ag(+) (aq) + e(-) —> Ag (s)

Oxidation Half Equation:
Cu (s) —> Cu(2+) (aq) + 2e(-)

Full Equation:
2Ag(+) (aq) + Cu (s) —> 2Ag (s) + Cu(2+) (aq)

Question 15

Potential Difference

Answer 15

Potential difference exists between two half-cells connected by an external wire.
Potential difference has the symbol E and unit of volts. It is measured using a voltmeter under standard conditions

Question 16

Standard Conditions

Answer 16

Pressure 100k Pa
1M concentration for solutions
Temperature 298K (25 degrees)

Question 17

Oxidation States

Answer 17

Oxidation states represent the charge that an atom would have if it was an ion

Question 18

Determining whether a redox reaction has occurred using oxidation states

Answer 18

If there is no change in oxidation numbers for all atoms in a reaction it is not a redox reaction

Question 19

Increase in oxidation state

Answer 19

An increase in oxidation state means an atom has been oxidised

Question 20

Decrease in oxidation state

Answer 20

A decrease in oxidation state means an atom has been reduced

Question 21

Oxidation rule for a free element

Answer 21

Oxidation state of a free element is 0

Question 22

Oxidation rule for a simple ion. E.g. Na(+)

Answer 22

The oxidation number of a simple ion is equal to the charge of the ion

Question 23

Oxidation rule for compounds (main group metals, hydrogen, oxygen)

Answer 23

Main group metals have an oxidation number equal to the charge of their ions

Hydrogen has an oxidation number of +1 in compounds with non-metals

Oxygen has on oxygen number of -2

Question 24

Balancing half equations using the half reaction method

Answer 24

1. Write the skeleton oxidation and reduction half equation
2. Balance all the elements except for hydrogen and oxygen in the half equations
3. Balance the oxygen atoms using H(2)O
4. Balance the hydrogen atoms using H(+)
5. Balance the charges by using electrons. The total charge on the left-hand side should equal the total charge on the right-hand side
6. Add the reactions together, cancelling the electrons and any other elements that are present on both sides (e.g. H(2)O and H(+))

Question 25

Galvanic Cell

Answer 25

Also known as voltaic cell
Electrochemical cell in which chemical energy is converted into electrical energy
Electrical energy produced via spontaneous redox reactions within the cell

Question 26

Parts of a galvanic cell

Answer 26

Two half cells: half reactions occur in separate half cells. Each half cell contains an electrode in contact with a solution (electrolyte).

Electrodes (anode and cathode): metal strips. Anode is where oxidation occurs, negative electrode. Cathode is where reduction occurs, positive electrode.

Solution (electrolyte): Contains ions. Conducts electricity as charged ions freely move. Soluble solutions that match electrodes ensure no competing reactions.

External circuit: Wire allows electron movement from anode to cathode.

Salt bridge: Contains ions that move to balance the charges formed in the half cells. Cations move towards cathode, and anions towards anode. Usually KNO(3) or NaNO(3).

Question 27

Electrolytic Cell

Answer 27

Electrolysis occurs, producing a non-spontaneous redox reaction by the passage of electrical energy from a power supply through a conducting liquid.
Convert electrical energy into chemical energy (operate in reverse compared to galvanic cells)

Question 28

Parts of an electrolytic cell

Answer 28

Porous barrier: half reactions occur in one container, porous barrier separates products

Solution (electrolyte): Contains a salt (made of positive and negative ions) that conducts electricity

Cathode: Reduction occurs, negatively charged

Anode: Oxidation occurs, positively charged

Voltage: Required voltage is higher than the voltage produced in an equivalent galvanic cell

Question 29

! Understand that the ability of an atom to gain or lose electrons can be predicted from the atom’s position in the periodic table, and explained with reference to valence electrons, consideration of energy and the overall stability of the atom !

Question 30

Redox conjugate pair

Answer 30

An electron donor and acceptor

Question 31

Inert electrode

Answer 31

an electrode that serves only as a source or sink for electrons without playing a chemical role in the electrode reaction
E.g. carbon