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3.1.11 electrode potentials and electrochemical cells

Terms in this set (121)

what happens when rod of metal is dipped into a solution of its own ions
an equilibrium is set up between the solid metal and the aqueous metal ions
write a half equation for zinc (s) to zinc (II)
Zn (s) ⇌ Zn2+(aq) + 2e-
write a half equation for copper (II) to copper (III)
Cu2+ (aq) ⇌ Cu3+ (aq) + e-
what is the simplest salt bridge made of
Filter paper soaked in saturated solution of KNO3 (potassium nitrate)
why is a salt bridge used rather than a piece of wire
to avoid further metal/ion potentials in the circuit
why are salt bridges necessary
Complete the circuit, but avoid further metal/ion potentials as does not perform electrochemistry. Allows ion movement to balance the charge. Do not react with electrodes
what symbol is used to represent a salt bridge in standard notation
II
what does the potential of a half cell tell us
how readily electrons are released by the metal - how good a reducing agent the metal is
what does a more negative electrode tell us
that it loses its electrons more readily than one higher
what type of species goes on the outside (furthest from the salt bridge) in standard cell notation
the most reduced species
what does I indicate
Phase boundary (solid/liquid/gas)
How would an aluminium/copper cell be represented
Al(s) | Al3+(aq) || Cu2+(aq) | Cu(s)
what happens at the left-hand electrode
land hand electrode is where oxidation occurs
left hand electrode is the half cell with the most negative E value
What happens at the right hand electrode
right hand electrode is where reduction occurs
right hand electrode is the half cell with the most positive E value
which side of the cell has the most negative E value
what happens to the metal with the most negative E value
oxidation - left hand electrode
how to calculate emf
in terms of electrons, explain the meaning of the term oxidising agent
electron acceptor
what is used to compare the tendency of different metals to release electrons
a standard hydrogen electrode
describe the standard hydrogen electrode
1. Hydrogen gas bubbled into a solution of H+ (aq) ions
2. 1 moldm-3
3.298K and 100kPa
4. Platinum as electrode
draw the hydrogen electrode
what is the potential of the standard hydrogen electrode
zero
why might you use other standard electrode occasionally
they are cheaper/easier/quicker to use and can provide just as good a reference
platinum is expensive
if an E value is more negative, what does it mean in terms of oxidising/reducing power?
better reducing agent (easier to oxidise)
if an E value is more positive, what odes it mean in terms of oxidising/reducing power
better oxidising agent (easier to reduce)
what factors will change E values
concentration of ions
temperature
what happens if you reduce the concentration of ions in the left hand half cell
equilibrium moves to the left to oppose the change of removing ions; this releases more electrons.
The E of the LH cell becomes more negative, so the emf of the cell increases
when would you use a platinum electrode
When both the oxidised and reduced forms of the metal are in aqueous solution
why is platinum chosen
inert so does not take part in electrochemistry
good conductor to complete circuit
what is the electrochemical series
A measure of how easily a metal forms ions in solution by losing electrons
how to calculate the voltage of the cell
subtract the lesser positive value from the other value
how would you predict if a reaction would occur
take the 2 half equations
find the species that is being reduced (the right hand electrode)
calculate its E value minus the E value of the species that is being oxidised (the left hand cell)
if E > 0 the reaction will occur
give examples of non-rechargeable cells
zinc/copper cells
zinc/carbon cell
what is the zinc/copper cell
Daniell cell
emf of 1.1
why is the zinc/copper cell not practical for portable devices
because of the liquids it contains
what was the first commercial cell made from
zinc/copper(II)
What are zinc/carbon cells more commonly known as?
Disposable batteries
what are the two reactions that take place in zinc/carbon cells
Zn reduced to ZN2+
NH4+ reduced to NH3 at carbon electrode
what is the electrolyte in a long life alkaline battery
potassium hydroxide
what are rechargeable batteries
If voltage applied is greater than the cell provides on discharging, the reaction is reversed & the cell recharges.
what are the reactions that occur in a lead/acid battery
Pb +SO4 2- -> PbSO4 (s) + 2e-
PbO2 + 4H+ + SO4 2- + 2e- -> PbSO4 + 2H2O
what are lead-acid batteries
rechargeable batteries used to operate the starter motors of cars
Nickel/cadmium cells are rechargeable AA batteries. what reactions occur at the electrodes
Cd(OH)2 (s) + 2e- -> Cd(s) + 2OH-
NiO(OH) (s) + H2O (l) + e- -> Ni(OH)2 (s) + OH- (aq)
what are the electrode reactions in a lithium cell
POSITIVE ELECTRODE = Li+ + CoO2 + e- → Li+[CoO2 ] -
NEGATIVE ELECTRODE = Li → Li+ + e-
when is the lithium cell used
in laptops and mobile phones
what is a fuel cell
a cell that is used to generate electric current; does not require electrical recharging
what are the reaction that take place at the two electrons in an alkaline hydrogen fuel cell
2H2 + 4OH- -> 4H2O + 4e-
O2 + 2H2O + 4e- -> 4OH-
draw a diagram of a hydrogen fuel cell
what is the weakest oxidising agent off electrode potential values
the most negative electrode potential
why is it better to use a fuel cell than to burn H2 in air, even though the same overall reaction occurs
in combustion, sulfur containing compounds (SO2,SO3) and nitrogen containing compounds (NO2, NOx) are produced due to the high temperatures an the S and N in air
these are bad for the environment
this does not occur in a fuel cell; the only product is water
more efficient
what are the disadvantages of fuel cells
hydrogen is a flammable gas with low b.p. -> hard and dangerous to store and transport -> expensive to buy
fuel cells gave a limited lifetime and use toxic chemicals in their manufacture
How do you find the weakest reducing agent from a table of electrode potentials
most positive E value
then it is the PRODUCT of the reduction equation i.e. imagine equation going from right to left
what is the reason that some cells cannot be recharged
reaction of the cell is not reversible - a product is produced that either dissipates or cannot be converted back into the reactants
why might the emf of a cell change after a period of time
concentrations of the ions change - the reagents are used up
how can the emf of a cell be kept constant
reagents are supplied constantly, so the concentrations of the ions are constant; E remains constant
what are some limitations of predicting feasibility
E does not include the kinetics of a reaction
some cells have a positive E value but the reaction might not take place if its required activation energy is high
in such cases, the rate of reaction is so slow that it is considered to be negligible
what are the advantages of hydrogen fuel cells
reduction in carbon dioxide emissions as only water is released
more efficient than diesel or petrol powered vehicles
what are the disadvantages of hydrogen fuel cells
expensive to store and transport hydrogen
less lifetime compared to conventional cells
toxic chemicals are used to produce this fuel cell
what is an alternative to the hydrogen fuel cell
methanol and ethanol
ethanol is produced by fermentation of sugars
ethanol is less expensive and easier to store when compared to that of hydrogen
what is the equation of the methanol and ethanol fuel cells
OXYGEN ELECTRODE - 4H+ + O2 + 4e- -> 2H2O
ETHANOL ELECTRODE - C2H5OH + 3H2O -> 2CO2 + 12H+ + 12e-
OVERALL
C2H5OH + 3O2 -> CO2 + 2H2O
write the conventional representation of the cell used to measure the standard electrode potetial for the Ag+/Ag electrode
Pt(s)/H2(g)/H+(aq)//Ag+(aq)/Ag(s)
state the conditions necessary when measuring the standard electrode potential
298K
100kPa
both solutions of 1 moldm-3
zero current
use data from the table to explain, in terms of redox, what happens when a soluble gold(I) compound containing Au+ ions is added to water
emf = 1.68-1.23 = 0.45
the Au+ ions oxidise water
you would observe bubbles of oxygen
equation:
4Au+ + 2H2O -> 4Au + O2 + 4H+
suggest why potassium chloride would not be suitable for use in the salt bridge of this cell
chloride ions form a precipitate with silver ions
use data in the table to explain what happens when a solution of iron(II) chloride is exposed to the air
E O2 > E Fe3+
therefore the iron(II) ions are oxidised into iron (III) ions by oxygen
suggest two reasons why the concentration of chlorine in the vase water decreases with tie
biocide reacts with bacteria
chlorine is give off
chlorine has reacted with water to form HCl and O2
suggest why this sampling techniques has no effect on the rate at which the concentration of chlorine in the vase water decreases
because the concentration of the remaining solution remains unchanged
why was it important to use an excess of potassium iodide solution
so that all the chlorine was reduced
use the standard electrode potential data to explain why I2 oxidises S2O3 2- under standard condition
E I2 > E S4O6 2-
0.54
explain the function of the salt bridge
the ions in the ionic substance move through the salt bridge
to maintain charge balance
completes the circuit
how to deduce the weakest reducing agent
species on RHS has most positive E (highest potential)
use data from the table to justify why sulphate ions should not be capable of oxidising bromide ions
E SO2 < E Br2
suggest the major advantage of using the fuel cell
a fuel cell converts more of the available energy from combustion of hydrogen into kinetic energy of the car
with reference to electrons, give the meaning of the term reducing agent
electron donor
explain why Cl is reduced when HOCl acts as an oxidising agent in terms of the oxidation state of this atom
chlorine goes from +1 to 0
how to identify which is the positive electrode
which has the more positive E
suggest one reason why a zinc oxygen cell is not recharged
the cell reactions cannot be reduced
deduce half equations for the electrode reactions in a hydrogen-oxygen fuel cell.
H2 -> 2H+ + 2e-
O2 + 4H+ + 4e- -> 2H2O
use the half equations of fuel cells to explain how an electric current can be generated
hydrogen electrode produces electrons and the oxygen electrode accepts eletrons
explain why a fuel cell does not need to be recharged
hydrogen is supplied continuously
suggest the main advantage of using hydrogen in a fuel cell rather than in an internal combustion engine
in a fuel cell, a greater proportion of the energy available from the hydrogen-oxygen reaction is converted into useful energy `
identify one major hazard associated with the use of a hydrogen-oxygen fuel cell in a vehicle
hydrogen is flammable
Draw a clearly labelled diagram to show the components and reagents, including their concentrations, in this Fe(III)/Fe(II) electrode.
platinum electrode
solution in beaker is a mixture of named soluble iron(II) compound and named soluble iron (III) compunds
concentration of Fe(II) and Fe(III) ions are both 1 mol dm-3
what is the purpose of a salt bridge
allow movement of ions between electrodes
state one essential requirement of the soluble ionic compound used to make the salt bridge
must not react with the electrolyte
A Daniell cell was set up using 100 cm3 of a 1.0 mol dm-3 copper(II) sulfate solution. The cell was allowed to produce electricity until the concentration of the copper(II) ions had decreased to 0.50 mol dm-3 .
Calculate the decrease in mass of the zinc electrode
moles of Cu(II) reacted = volume x concentration = (100/1000) x 0.5 = 0.05
1:1 ratio = moles of zinc reacted = 0.05
mass lost = moles x Mr = 0.05 x 65.4 = 3.27g
you are provided with a daniell cell, including a zinc electrode of known mass. briefly outline how you would carry out an experiment to confirm the decrease in mass of the zinc electrode
allow the cell to discharge until the concentration of the Cu(II) ions is 0.5. this is confirmed by colorimetric measurements. then weigh the Zn electrode before and after the experiment
explain why rechargeable cells are connected to solar cells
solar cells do not supply enough energy all the time. therefore rechargeable cells can store electrical energy for use when the solar cells are not working
suggest one reason why many waste disposal centres contain separate sections for cells and batteries
to prevent pollution of the environment by toxic or dangerous substances
how does a salt bridge provide and electrical connection between the two electrodes
it maintains and ionic balance by having ions which can move through it by completing the circuit
suggest why potassium chloride would not be a suitable salt for the salt bridge in this cell
the chloride ions would react with the copper ions
in the external circuit of this cell, the electrons flow through the ammeter from right to left. suggest why the electrons move in this direction
because the Cu(II) ions in the left hand electrode are more concentrated
so the reaction of Cu(II) ions with 2e- will occur in preference at the left hand electrode
explain why the current in the external circuit o this cell falls to zero after the cell has operated for some time
eventually the copper ions in each electrode will be at the same concentration
suggest why the recharging of a lithium cell may lead to release of carbon dioxide into the atmosphere
electricity for recharging the cell may come from power stations burning fossil fuels
suggest a reason why the aluminium used as the electrode is rubbed with sandpaper prior to use
to remove the oxide layer on the aluminium
a simple salt bridge can be prepared by dipping a piece of filter paper into potassium carbonate solution. explain why a such a salt bridge would nto be suitable for use in this cell
there would be a reaction between the carbonate and the aluminium (III) ions
use data from the table to explain why Au+ ions are not normally found in aqueous solution. write an equation to show how Au+ ions would react with water
E Au+ > E O2
so Au+ ions will oxidise water
2Au+ + H2O -> 2Au + 1/2O2 + 2H+
in terms of electrons, state the meaning of the term oxidising agent
an electron acceptor
explain how E values can be used to deduce the identity of the weakest oxidising agent in the table
species on the LHS with the most negative electrode potential
deduce one essential property of the non-reactive porous separator labelled in the diagram
allows ions to pass through it
suggest the function of the carbon rod in the cell
allows electrons to flow
the zinc electrode acts as a container for the cell and is protected from external damage. suggest why a cell often leaks after being used for a long time
zinc gets oxidised
state one environmental advantage of rechargeable cells compared with a non-rechargeable cell
metal is reused
deduce the equation for the overall reaction that occurs in the ethanol=oxygen fuel cell
C2H6O + 3O2 -> 2CO2 + 3H2O
deduce a half equation for the reaction at the ethanol electrode. in this half equation ethanol reacts with water to form carbon dioxide and hydrogen ions
C2H6O + 3H2O -> 2CO2 +12H+ + 12e-
suggest why ethanol can be considered to be carbon neutral
CO2 released by combustion is equal to that of the atmospheric CO2 taken up in photosynthesis
use data from the table below to explain why dilute hydrochloric acid cannot be used to acidify potassium manganate (VII) in a titration
manganate would oxidise Cl-
because E for MnO4- is more positive than that for the Cl2
give two reasons why platinum is a suitable electrode for an Fe(II) and Fe(III) electrolyte
it is inert
conducts electricity
use data from the table to explain why chlorine should undergo a redox reaction with water
ECl2 > E O2
suggest one reason why the redox reaction between chlorine and water does not normally occur in the absence of light
the activation energy is high and light provides the activation energy
H2O2 + 2H+ + 2e- -> 2H20
explain in terms of oxidation state what happens to hydrogen peroxide when it is reduced
O goes from oxidation state -1 in H2O2 to oxidation to -2 in H2O
in terms of electrons, state what happens to a reducing agent in a redox reaction
it donates electrons
how to identify the strongest reducing agent form the species in the table
the strongest reducing agent is the most negative E value
use the data from the table to explain why fluorine reacts with water
E F2 > E O2
fluorine is a more powerful oxidising agent than oxygen
suggest one reason why the cell cannot be electrically recharged
reactions not reversible
give one reason why the emf of the lead-acid cell changed after several hours
the reagents are used up
A student determined the concentration of iron(II) ions in a solution of iron(II) chloride by titration with acidified potassium dichromate(VI) solution. A second student titrated the same solution of iron(II) chloride with acidified potassium manganate(VII) solution. By reference to the table, explain why the second student obtained a greater value for the concentration of iron(II) ions.
MnO4- will oxidise the chloride ion so a larger volume is needed
suggest two properties of platinum that make it suitable for use as an external electrical contact in the cell
platinum is a conductor
platinum is inert
suggest one reason why water is not used as a solvent in this cell
because lithium reacts with water
describe a standard hydrogen electrode
1. Hydrogen gas
2. 1 moldm-3 H+
3.298K and 100kPa
3. Platinum as electrode
suggest what reactions occur, if any, when hydrogen gas is bubbled into a solution containing a mixture of iron(II) and iron(III) ions. explain your answer
Fe3+ ions reduced to Fe2+ because E Fe3+ > E H+
State why the electrode potential for the standard hydrogen electrode is equal to 0.00V.
by definition as it is set to this value
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