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Chemistry Module 5 > 5.2.3 Electrode potential and fuel cells > Flashcards

5.2.3 Electrode potential and fuel cells Flashcards

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1
Q

oxidation

A

loss of electrons

increase in oxidation number

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2
Q

reduction

A

gain of electrons

decrease in oxidation number

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3
Q

in an oxidation half equation where are the electrons

A

on the right

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4
Q

in a reduction half equation where are the electrons

A

on the left

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5
Q

reducing agent

A

electron donors

get oxidised

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6
Q

oxidising agents

A

electron acceptors

get reduced

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7
Q

how do you combine half equations

A

multiply half equations to get equal electrons
add half equations together and cancel electrons
multiply half equations to get equal electrons
add half equations together and cancel electrons

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8
Q

common redox titrations

A

thiosulfate

manganate

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9
Q

thiosulfate redox reaction

A

2(S2O3)^2- (aq) + I2 (aq) -> 2I- (aq) +S4O6 2- (aq)

yellow/brown to colourless
starch indicator is added near the end point when iodine fades a pale yellow, colour change is blue/black to colourless

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10
Q

manganate redox titration

A

self-indicating

MnO4- (aq) + 8H+ (aq) + 5Fe 2+ (aq) -> Mn 2+ (aq) + 4H2O (l) + 5Fe 3+ (aq)

purple to colourless
if manganate is in the burette then colourless to purple will be change

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11
Q

choosing the correct acid for a manganate titration

A

need to supply 8H+ ions
use dilute sulfuric acid
in excess

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12
Q

why can’t you use ethanoic acid in manganate titration

A

can’t supply the large amount of H+ ions needed

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13
Q

why can’t you use conc HCl in manganate titrations

A

as Cl- ions would be oxidised to Cl2 by MnO4-
as E0 value of Mn is greater than the Cl
leads to greater amount of manganate being used and toxic Cl2 gas being produced

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14
Q

why can’t nitric acid be used in manganate titration

A

as it is an oxidising agent
will oxidise Fe2+ to Fe3+
as E0 of NO3 is greater than the fe2 and fe3

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15
Q

other useful manganate titrations

A 
with hydrogen peroxide 
with ethanedioate (slower so heat to 60 degrees)
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16
Q

electrochemical cells

A

a cell contains 2 half-cells
these are connected via a salt bridge
simple half cells contain a metal (acts as an electrode) and a 1.0 moldm-3 solution containing that metal
these will produce a small voltage if connected into a circuit

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17
Q

drawing the half-cells

A

2 beakers
1 with 1 metal electrode and a 1M solution of the metal
the other with another metal but exactly the same
connect them up through a voltmeter
electron flow needs to be drawn in the correct direction
MUST DRAW AND LABEL SALT BRIDGE BETWEEN BEAKERS

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18
Q

why does a voltage form in half-cells

A

when connected, one of the metals will have a greater tendency to oxidise than the other
place the one with the most positive E0 value on the bottom and follow the anti-clockwise rule
more electrons will build up on one electrode than the other so potential difference is established

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19
Q

why use a high resistance voltmeter

A

to stop the current flowing in a circuit
to measure the maximum possible E value
reactions will not be occurring as voltmeter stops current from flowing

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20
Q

salt bridge

A

used to connect the circuit
free moving ions conduct the charge
made of filter paper soaked in salt solution e.g. potassium nitrate
salt should be unreactive with electrodes and electrode solutions

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21
Q

why isn’t a wire used as a salt bridge

A

wire isn’t used as it would establish its own electrode system with the solutions

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22
Q

what happens if current is allowed to flow

A

reactions then occur separately at each electrode
voltage will fall to zero and reactants are used up
most positive will always undergo reduction
most negative will always undergo oxidation

23
Q

what are standard electrode potentials always measured as

A

reduction values
most positive will always undergo reduction
most negative will always undergo oxidation

24
Q

measuring the electrode potential of the cell

A

can’t measure the absolute potential of a half electrode on its own but can measure the potential difference between the two half-cells measured
connect to another half cell of a known potential and calculate difference
connect to standard hydrogen cell

25
Q

what is the reference electrode we use

A

hydrogen electrode

potential of 0 volts

26
Q

what is the standard hydrogen equilibrium

A

H2 (g) (reversible) 2H+ (aq) + 2e-

27
Q

components of a standard hydrogen electrode

A
  1. hydrogen gas at pressure 100kPa
  2. solution containing the hydrogen ion at 1moldm-3
  3. temperature at 298K
28
Q

draw the standard reference electrode

A

draw the two beakers
platinum as 1 electrode, 1M solution of HCl
hydrogen gas is collected at 100kPa
KNO3 salt bridge
metal electrode and metal ion solution at 1M
voltmeter

29
Q

secondary standards

A

calibrated against the standard hydrogen cell
standard electrode that has been calibrated against the primary standard
common are:
silver/silver chloride E0= +0.22V
calomel electrode E0= +0.27V

30
Q

standard electrode potential

A

when an electrode system is connected to the hydrogen electrode system and standard conditions apply, the potential difference measured is called the standard electrode potential

31
Q

what are the standard conditions for electrode potentials

A

all ion solutions at 1M
temperature 298K
gases at 100kPa
no current flowing

32
Q

what is the most useful application of electrode potentials

A

to show direction of spontaneous change for redox reactions

33
Q

how do you work out the E value of a cell

A

E reduction- E oxidation
subtract the most positive value from the most negative value
a spontaneous change will always have a positive E value

34
Q

more positive the E0

A

greater tendency for the species on the left to reduce and act as an oxidsing agent

35
Q

more negative the E0

A

greater tendency for the species on the right to oxidise and act as reducing agents

36
Q

most powerful reducing agents

A

found at most negative end of the series on the right

one with lower oxidation number

37
Q

most powerful oxidising agents

A

found at the most positive end of the series on the left

one with higher oxidation number

38
Q

what is cell EMF

A

measure of how far from equilibrium the cell reaction lies

more positive the more likely the reaction is to occur

39
Q

effects of changing conditions on cell can be explained by what

A

Le Chatelier’s principle

40
Q

effect of concentration on cell emf

A

increasing concentration of “reactants” would increase emf and decreasing them would cause emf to decrease

41
Q

effect of temperature on cell emf

A

most cells are exothermic in the spontaneous direction so this would result in a decrease in E cell

42
Q

even if E value is positive what may not occur

A

the reaction
may occur so slowly or not occur at all
if reaction has a high activation energy it’ll not occur

43
Q

explaining homogeneous catalysis using E values

A

reaction between I- and S2O82- is catalysed by Fe2+
uncatalyzed is very slow as reaction needs a collision between 2 negative ions, but the repulsion results in a high activation energy
for a substance to act as a homogeneous catalyst its electrode potential value must lie between the electrode potentials of the 2 reactants so it can oxidise one and reduce the other

44
Q

cells

A

electrochemical cells can be used as a commercial source of electrical energy
can be non-rechargeable, rechargeable and fuel cells

45
Q

when are cells non-rechargeable

A

when reactions that occur within them are non-reversible

46
Q

fuel cells

A

use the energy from the reaction of a fuel with oxygen to create a voltage

47
Q

what are fuel cells vehicles fuelled by

A

hydrogen gas

hydrogen-rich fuels

48
Q

hydrogen fuel cell

A

uses a potassium hydroxide electrolyte

overall reaction: 2H2 + O2 -> 2H2O

49
Q

voltage of fuel cells

A

remains constant as they are continuously fed with fresh O2 and H2 so they maintain a constant conc of reactants
ordinary cells the voltage drops over time as the reactant conc drops

50
Q

fuel cells in acidic conditions

A

e cell is the same as alkaline conditions as the overall equation is the same

51
Q

using standard conditions fuel cells

A

rate is too slow to produce an appreciable current
higher temps are therefore used to increase the rate but reaction is then exothermic so emf falls
higher pressure counteracts this

52
Q

advantages of fuel cells over conventional petrol or diesel-powered vehicles

A

less pollution and CO2 (pure hydrogen emits only water whilst hydrogen-rich fuels produce only small amounts of air pollutants and CO2)
greater efficiency

53
Q

limitations of hydrogen fuel cells

A

storing and transporting hydrogen in terms of safety, feasibility of a pressurised liquid and a limited life cycle of a solid “adsorber” or “absorber”
limited lifetime (regular replacement) high production costs
use of toxic chemicals in their production

Chemistry Module 5 (7 decks)

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