Developing Metals - Electrochemistry

Question 1

draw a diagram of a simple electrochemical cell between copper and zinc

Answer 1

Question 2

what ions are usually present in a salt bridge?

Answer 2

KNO₃

Question 3

what is E_cell?

Answer 3

the maximum voltage produced by a cell

Question 4

how are standard electrode potentials calculated?

Answer 4

hydrogen half cell used as a reference, so other half cells are measured against it

Question 5

how is E_cell calculated?

Answer 5

voltage measured

E_cell = E_red - E_ox

Question 6

how can E_red and E_ox be determined?

Answer 6

E_red is the more positive electrode

E_ox is the more negative electrode

Question 7

if a substance is a good oxidising agent, will it have a high or low electrode potential?

Answer 7

high

Question 8

if a substance is a good reducing agent, will it have a high or low electrode potential?

Answer 8

low

Question 9

why will a cell eventually stop working?

Answer 9

E_red and E_ox grow closer together until they are both the same

Question 10

when might a platinum electrode be used and why?

Answer 10

when non-metals are used

they are inert

Question 11

what is a potential difference?

Answer 11

tendency of an electrode to release or accept electrons

Question 12

what are the conditions for electrochemical cells?

Answer 12

298K

1 mol dm^-3 solution

Question 13

what is the purpose of the salt bridge?

Answer 13

so the circuit can flow, just through ions instead of electrons

Question 14

why might a reaction not occur even with a positive E_cell?

Answer 14

E_cell only tells you whether reaction is possible, not rate

reaction may be so slow that no change is observed

Question 15

why might a negative voltage be recorded?

Answer 15

may be negative at standard conditions

will be positive in different conditions

Question 16

what are some common non-metal half cells?

Answer 16

halogens

MnO₄^-

Question 17

what is the formula of rust?

Answer 17

Fe₂O₃

Question 18

what are the 2 half equations involved in rusting?

Answer 18

Fe_(s) -> Fe²⁺_(aq) + 2e^-

½O₂ + H₂O + 2e^- -> 2OH^-

Question 19

which half reaction is occuring at the positive ‘electrode’ in rusting?

Answer 19

reduction of oxygen to hydroxide ions

Question 20

in which direction are electrons flowing in rusting?

Answer 20

towards where the iron is being oxidised into Fe²⁺

Question 21

explain the process of rusting

Answer 21

at edges of droplet of water, O₂ reduced to OH^-

at centre of water - Fe oxidised

electrons released flow to edges of droplet through metal to reduce more oxygen

iron and hydroxide react to form Fe(OH)₂

reacts with O₂ to form rust - Fe₂O₃

Question 22

why is water reduced at the edges of the water droplet in rusting?

Answer 22

the concentration of water is highest here

Question 23

why is iron oxidised at the centre of a water droplet in rusting?

Answer 23

water concentration is lowest here

Question 24

what are the methods of protecting metals against rusting?

Answer 24

sacrificial protection

impressed currents

Question 25

what is sacrificial protection?

Answer 25

by putting a layer of a different material on top of the metal, so the original layer is sacrificed before the metal can rust

Question 26

how does sacrificial protection help prevent rusting?

Answer 26

the more electronegative metal rusts, shifting its equilibrium to the left to counter, iron equilibrium shifted to the right iron rusting prevented

Question 27

what are some common barriers against rusting?

Answer 27

galvanised steel plastic lining

Question 28

what is galvanised steel and how does it help to prevent rusting?

Answer 28

iron covered in zinc covered in zinc oxide zinc oxide will not rust, and if scratched/broken through, the zinc will rust before the iron can

Question 29

what is the impressed current method in preventing rusting?

Answer 29

making a sacrificial metal an cathode by supplying electrons from an external source reduction will occur at the sacrificial cathode, instead of the anode (iron)