electrode potentials Flashcards
what is the process of oxidation and reduction expressed in terms of
electron transfer
oxidation
loss of electrons
reduction
gain of electrons
oxidation is an increase in
oxidation state
reduction is a decrease in
oxidation state
what can we work out by assigning oxidation states
which species have been oxidised and which have been reduced without writing half equations
oxidising agent
species that accepts electrons
reducing agent
species that donates electrons
example of an oxidising agent
halogens
example of a reducing agent
reactive metals
what does an electrochemical cell use the electron transfer which occurs during a redox reaction to do
produce electrochemical energy
what does each electrochemical cell consist of
2 half cells corresponding to the two half equations occurring in the redox reaction
what does a half cell consist of
two species of the same element in different oxidation states
what is the electrode in a copper half cell
solid copper
what is the solution in copper half cell
1 mol dm3 copper 2+ ions
what is the IUPAC convention for writing half equations for electrode reactions
to give them as an equilibrium reaction where the forward reaction is reduction
what must the half cells be connected via in order to make use of the electrical energy produced in a redox reaction
connected via an external circuit and a salt bridge
what is the circuit usually
a conducting wire
what flow through the wire
electrons
what does the salt bridge consst of
a strip of filter paper soaked in saturated potassium nitrate or potassium chloride solution
what flows through the salt bridge
ions
what are the electrodes connected to
a high resistance voltmeter
what does the voltmeter easure
electromotive force (emf) in volts
what is the emf
potential difference
what is the other name for emf
cell potential
what are the standard conditions are there for a half cell
- 100kPa
- 298K
- solution of ions at 1moldm-3
which metal goes on the rhs
the more positive metal
which metal goes on the rhs in the standard copper-zinc cell
copper
why is the copper the positive electrode in the copper-zinc cell
because it is connected to the positive terminal of the voltmeter and so is the positive electrode
what is the IUPAC convention for cell diagrams
to draw the positive electrode as the half cell on the RHS (except when measuring standard electrode potential against hydrogen electrode when hydrogen always on left)
what occurs at the negative electrode
oxidation
what occurs at the positive electrode
reduction
where do electrons flow in cell
from negative electrode to positive electrode
what does the standard electrode potential of a half cell indicate
its tendency to lose/gain electrons
is it possible to measure the standard electrode potential of one half cell alone
no
what is used as a reference to measure standard electrode potential against
a primary standard
`what is the primary standard for measuring electrode potentials
the standard hydrogen electrode
what emf does the standard hydrogen electrode have by definition
0.00V
standard hydrogen electrode
- H2 gas at 1 00kPa
- 1 moldm-3H+ (aq)
- 298K
- Pt electrode
equation for standard hydrogen electrode
2H+ + 2e->< H2(g)
standard electrode potential
emf of a half cell compared with a standard hydrogen half-cell
measured at 298K, all solutions having conc of 1moldm-3 and all gases at a pressure of 100kPa
what is the standard electrode potential found by measuring
the voltage of a cell with the standard hydrogen half cell as one half cell and the half cell under investigation as the other
what gives the sign of the electrode potential
the polarity of the half cell relative to the hydrogen half cell
conventional representation of a cell: what does a vertical solid line indicate
a phase boundary eg between a solid and a solution
conventional representation of a cell: what does a double vertical line show
a salt bridge
conventional representation of a cell: which species for each half cell is written next to the salt bridge
the species with the highest oxidation state
conventional representation of a cell: where is the standard hydrogen half cell shown when measuring the standard electrode potential of a half cell
LHS
conventional representation of a cell: where is the positive electrode shown when showing any other electrochemical cell
RHS
conventional representation of a cell: where is the negative electrode shown when showing any other electrochemical cell
LHS
what is an ion/ion half cell
half cells in which both species are aqueous ions
what must an ion/ion half cell contain
both the oxidised and reduced species
what electrode is used in an ion/ion half cell
platinum
what are standard electrode potential values determined by
measurement against the standard hydrogen electrode
what are all reactions shown as in the electrochemical series
reductions
what can we determine looking at the electrochemical series
the relative reactivity of different species
metals react by losing what to form positive ions
electrons
what do reactive metals undergo more readily
oxidation
the most reactive metals will have the most negative
E0 values
what are the most reactive metals good at
being reducing agents
what is the correlation between negativity of E0 value of metal and tendency for species on RHS to lose electrons and be oxidised
more negative E0 value, greater tendency for species on RHS to lose electrons and be oxidised
how to non metals react
by gaining electrons to form negative ions
reactive metals undergo what more readily
reduction
most reactive non metals will have most positive what
E0 values
most reactive non metals are good at what
being oxidising agents
correlation between positivity of E0 value of non metal and tendency for species of LHS to gain electrons and be reduced
more positive E0 value, greater tendency for species on LHS to gain electrons and be reduced
what must a half cell be connected to to determine the electrode potential of it
the standard hydrogen electrode
what can be connected to generate an emf
any 2 half cells
what is the cell potential the difference between
the standard electrode potentials of the two half cells
what is chemical energy converted into in an electrochemical cell
electrical energy
what does the emf drop to once the chemicals are used up
0.00v
which half equation goes forwards
the one with the more positive e0 value- gains electrons
which half equation goes backwards
the one with the more negative e0 value- loses electrons
how to determine which reaction will occur in the cell or in a test tube
add the two half equations, making sure that the number of electrons lost equals number of electrons gained
measuring electrode potential for a metal/metal ion half cell eg copper: equation
Cu2+ +2e- <> Cu(s)
the conventional representation of a cell: what does a vertical solid line indicate
a phase boundary eg between a solid and a solution
the conventional representation of a cell: a double vertical line shows the
salt briddge
the conventional representation of a cell: the species with the highest ? for each half cell is written next to the salt bridge
oxidation state
the conventional representation of a cell: when measuring the standard electrode potential of a half cell the standard hydrogen half cell is always shown on the
left hand side
The conventional representation of a cell: on which side is the positive electrode shown
The RHS
The conventional representation of a cell: on which side is the negative electrode
Left hand side
The conventional representation of a cell: using what equation can the cell potential be calculated
E0RHS-E0LHS
To measure E0 for ion/ion half cells: what are ion/ion half cells
Half cells in which both species are aqueous ions
To measure E0 for ion/ion half cells: what must the half cells contain
Both the oxidised and reduced species
To measure E0 for ion/ion half cells: what electrode is used
Platinum
The electrochemical series: in what order can the standard electrode potential values determined by measurement against standard hydrogen electrode be placed in
Numerical order
The electrochemical series: what are all reactions shown as
Reductions
The electrochemical series: what can we determine by looking at electrochemical series
Relative reactivity of different species
The electrochemical series: metals react by losing _____ to form ______ ions
Losing electrons to form positive ions
The electrochemical series: reactive metals undergo what more readily
Oxidation
The electrochemical series: the most reactive metals will have the most _______ E0 values
Negative
The electrochemical series: the most reactive metals are good _______ agents
Reducing
The electrochemical series: the more negative the E0 values the ________ the tendency for the species on the _____ to _____ electrons and be _______
The more negative the E0 the greater the tendency for the species on the RHS to lose electrons and he oxidised
The electrochemical series: non metals react by ______ electrons to form ______ ions
Gaining electrons to form negative ions
The electrochemical series: reactive non metals undergo ______ more readily
Reduction
The electrochemical series: the most reactive non metals will have the most _____ E0 values
Positive
The electrochemical series: the most reactive non metals are good ______ agents
Oxidising
The electrochemical series: the more positive the E0 the _______ the tendency for the species on the _____ to _____ electrons and be _____
The more positive the E0 the greater the tendency for the species on the LHS to gain electrons and be reduced
Measuring the emf of an electrochemical cell, required prac: how else can emf be determined other than using hydrogen electrode
Any two half cells can be connected to generate an emf
Cell potential = difference between standard electrode potentials of 2 half cells
Using electrode potentials to predict chemical reactions: what is chemical energy converted to in an electrochemical cell
Electrical energy
Using electrode potentials to predict chemical reactions: when did the emf drop to 0.00V in an electrochemical cell
Once the chemicals are used up
Using electrode potentials to predict chemical reactions: how can we predict what reaction occurs
By looking at the E0 values of the two half cells
Using electrode potentials to predict chemical reactions: which half equation gains electrons and what does this mean
The half equation with the more positive E0 value gains electrons, so goes forwards
Using electrode potentials to predict chemical reactions: which half equation loses electrons and what does this mean
The half equation with the more negative E0 value loses electrons, so goes backwards
Using electrode potentials to predict chemical reactions: what must be certified when adding the two half equations together
Number of electrons lost=number of electrons gained
Change in mass of the electrode: why does the mass of the negative electrode decrease when a current is drawn
The metal will be oxidised to metal ions
Change in mass of the electrode: what happens to the mass of the positive electrode and why
Increases as metal ions turn to metal
The effect of concentration on the emf of a cell: what are the values of E0 measured under standard conditions of
100kPa, 298K, 1 mol dm-3
The effect of concentration on the emf of a cell: what can be used to predict how concentration changes affect EMF
Le chateliers lrinciple
The effect of concentration on the emf of a cell (eg Zn|Zn2+||Cu2+|Cu): what will happen if [Cu2+]>1moldm3
- equilibrium shifts right
- ECu more negative
- Cu2+ gains more electrons
- Ecell will increase- bigger difference in E
The effect of concentration on the emf of a cell (eg Zn|Zn2+||Cu2+|Cu): what will happen if [Cu2+]<1moldm-3
- equilibrium shift LHS
- ECu less positive
- Cu2+ will gain feevwr electrons
- Ecell will decrease- cell has smaller difference in E
The effect of concentration on the emf of a cell (eg Zn|Zn2+||Cu2+|Cu): what will happen if [Zn2+]>1mol dm-3
- equilibrium will shift RHS
- EZn less negative
- Zn will release fewer electrons
- Ecell will decrease as cell has smaller difference in E
The effect of concentration on the emf of a cell (eg Zn|Zn2+||Cu2+|Cu): what will happen if [Zn2+] < 1moldm-3
- equilibrium shift LHS
- EZn become more negative
- Zn will release more electrons
- Ecell will increase as cell has a bigger difference in E
if conditions are no longer standard what is cell potential written as as opposed to E0Cell
ECell
Limitations of using electrode potentials tocpredict redox reactions
- predictions using electrode potentials tell us about equilibrium but not RoR
- many reactions in laboratory not done under standard conditions- will affect electrode potential vales
Commercial applications of electrochemical cells: what 3 main types can cells be divided into
- non- rechargeable cells
- rechargeable cells
- fuel cells
Non rechargeable cells: how are they designed to be used only once
Reactions occurring in cells cannot he reversed
Non rechargeable cells: what will happen when the chemicals are used up
The battery will go flat and the emf will fall to 0.00V
Non rechargeable cells: used in
Smoke detectors and clocks
Non rechargeable cells: overall equation of zinc and manganese dioxide cell when it discharged
2MnO2 + 2H2O + Zn > 2MnO(OH) + 2OH- + Zn2+
Non rechargeable cells: function of porous separator
Allows ions to flow
Non rechargeable cells: function of carbon rod
Allows electrons to flow
Non rechargeable cells: why cell often leaks after being used for a long time
Zn used up as reaction proceeds
Rechargeable cells: why can they be recharged
Reaction occurring in cell can be reversed and chemicals in cell regenerated
Rechargeable cells: examples
- lead/acid- cars
- Ni/Cd- torches/radios
- lithium/ion-phones, tablets, cameras etc
Lithium ion cells: why are they light
Lithium is the least dense mental
Lithium ion cells: reaction occurring at positive electrode
Co(IV) reduced to Co(III)
Lithium ion cells: reaction occurring at negative electrode
Li(0) oxidised to Li(+1)
Lithium ion cells: equation for discharge of cell
CoO2 + Li > Li+[CoO2]-
Lithium ion cells: equation for recharging cell
Discharge reaction reversed
Li+[CoO2]- > CoO2 + Li
The effect of concentration on the emf of a cell (eg Zn|Zn2+||Cu2+|Cu): what will happen if [Cu2+]<1moldm-3
- equilibrium shift LHS
- ECu less positive
- Cu2+ will gain feevwr electrons
- Ecell will decrease- cell has smaller difference in E
The effect of concentration on the emf of a cell (eg Zn|Zn2+||Cu2+|Cu): what will happen if [Zn2+]>1mol dm-3
- equilibrium will shift RHS
- EZn less negative
- Zn will release fewer electrons
- Ecell will decrease as cell has smaller difference in E
The effect of concentration on the emf of a cell (eg Zn|Zn2+||Cu2+|Cu): what will happen if [Zn2+] < 1moldm-3
- equilibrium shift LHS
- EZn become more negative
- Zn will release more electrons
- Ecell will increase as cell has a bigger difference in E
if conditions are no longer standard what is cell potential written as as opposed to E0Cell
ECell
Limitations of using electrode potentials tocpredict redox reactions
- predictions using electrode potentials tell us about equilibrium but not RoR
- many reactions in laboratory not done under standard conditions- will affect electrode potential vales
Commercial applications of electrochemical cells: what 3 main types can cells be divided into
- non- rechargeable cells
- rechargeable cells
- fuel cells
Non rechargeable cells: how are they designed to be used only once
Reactions occurring in cells cannot he reversed
Non rechargeable cells: what will happen when the chemicals are used up
The battery will go flat and the emf will fall to 0.00V
Non rechargeable cells: used in
Smoke detectors and clocks
Non rechargeable cells: overall equation of zinc and manganese dioxide cell when it discharged
2MnO2 + 2H2O + Zn > 2MnO(OH) + 2OH- + Zn2+
Non rechargeable cells: function of porous separator
Allows ions to flow
Non rechargeable cells: function of carbon rod
Allows electrons to flow
Non rechargeable cells: why cell often leaks after being used for a long time
Zn used up as reaction proceeds
Rechargeable cells: why can they be recharged
Reaction occurring in cell can be reversed and chemicals in cell regenerated
Rechargeable cells: examples
- lead/acid- cars
- Ni/Cd- torches/radios
- lithium/ion-phones, tablets, cameras etc
Lithium ion cells: why are they light
Lithium is the least dense mental
Lithium ion cells: reaction occurring at positive electrode
Co(IV) reduced to Co(III)
Lithium ion cells: reaction occurring at negative electrode
Li(0) oxidised to Li(+1)
Lithium ion cells: equation for discharge of cell
CoO2 + Li > Li+[CoO2]-
Lithium ion cells: equation for recharging cell
Discharge reaction reversed
Li+[CoO2]- > CoO2 + Li
Fuel cells: what does it use to create a voltage
Energy from the reaction of a fuel with oxygen
Fuel cells: what flows into the fuel cell and what flows out
Fuel and oxygen flow into fuel cell and products flow out
Fuel cells: what remains in the cell
The electrolyte
Fuel cells: what can fuel cells operate continuously depending on
Provided the fuel and oxygen are supplied into cell
Fuel cells: do fuel cells have to be recharged
No
The alkaline hydrogen oxygen fuel cell: what are the cells two electrodes made out of
Porous Pt based material
The alkaline hydrogen oxygen fuel cell: what are the two electrodes separated by
A semi permeable membrane
The alkaline hydrogen oxygen fuel cell: what is the electrolyte
Sodium hydroxide solution
The alkaline hydrogen oxygen fuel cell: which reaction takes place when hydrogen enters at the negative electrode
H2 + 2OH- > 2H2O + 2e-
The alkaline hydrogen oxygen fuel cell: reaction taking place at positive electrode
O2 + 2H2O + 4e- > 4OH-
The alkaline hydrogen oxygen fuel cell: overall equation
2H2 + O2 > 2H2O
Advantages of fuel cells: why are they more efficient
They convert more of their available energy into kinetic energy
Advantages of fuel cells: what is the only waste product
Water
Advantages of fuel cells: why do they not need to be recharged
They produce a continuous supply of energy so long as hydrogen and oxygen are supplied so they do not need to be recharged
Disadvantages of fuel cells: what is the disadvantage relating to hydrogen being produced
Can be done by electrolysis of water but requires electricity- burning fossil fuels
Disadvantages of fuel cells: why must hydrogen be handled carefully when stored and transported
It’s explosive and flammable
Disadvantages of fuel cells: what is there nothing of relating to hydrogen fuel
No infrastructure to provide hydrogen fuel for cars
