Electrode Potentials + Electrochemical Cells Flashcards
Basics
- when electrons move , you get electricity
- electrons move from more reactive metal to less reactive metal
- in an electrochemical cell : ALWAYS 2 REACTIONS - REDOX
- anode = positive = oxidation
- cathode = negative = reduction
- electrode potential + electromotive forever are the same thing - E°
Half cells / electrodes
- when a rod of metal is dipped into a solution / its own ions , an equilibrium is set up
E.g. dipping zinc into zinc sulfate solution sets up the following equilibrium:
Zn (s) = Zn2+ (aq) + 2e-
What makes an electrochemical cell
HALF - CELL + HALF - CELL = ELECTROCHEMICAL CELL
An electrochemical cell is made from …
- 2 different metals
- both dipped in salt so,it ions of their own ions
- connected by a wire ( the external circuit )
- each one is a half - cell
- circuit completed by a salt bridge
- use a voltmeter in the external circuit to measure the voltage between 2 half - cells
What are salt bridges
- complete circuit
- piece of filter paper soaked in a solution of salt ( usually saturated potassium nitrate KNO3)
What are salt bridges used for ?
- it allows ions to flow through + balance the charges to maintain neutrality
- stops both going into equilibrium
Why are salt bridges used rather than a piece of wire ?
- to avoid further metal / ion potentials in the circuit
- they must be inert + not react with the solutions in the half - cells
What happens in the zinc / copper electrochemical cell ( zinc side ) ?
- Zinc loses electrons more easily than copper
- ZN (S)= ZN2+ (aq) + 2e-
- releases electrons into the external circuit
- more negative electrode potential , showing it loses electrons more easily
- better reducing agent
- is oxidised
What happens in the zinc / copper electrochemical cell ( copper side ) ?
- Copper loses electrons less easily than zinc
- Cu(s) + 2e- =Cu2+ (aq)
- gains electrons from external circuit (same number as zinc lost )
- more positive electrode potential , harder to lose electrons
- is reduced
Overall equation in the cell
Cu 2+ (aq) + Zn (s) = Cu (s) + Zn2+ (aq)
What happens in electrochemical cell when voltmeter removed ?
Electrons allowed to flow , they would do so from zinc to copper
What following changes would take place ?
- Zinc would dissolve to form Zn2+ (aq) in solution -> increasing the concentration of Zn2+ (aq)
- The electrons would flow through the wire to the copper rod
- They would combine with Cu2+ (aq) ions from the copper sulfate solution -> depositing fresh copper on the rod + decreasing the conc of Cu2+ (aq) in solution
Daniel Cell
An electrochemical cell when two half cells are connected they generate electricity
Electrode potentials
These measure how easily a meta is oxidised
- how good of a reducing agent it is
- how easily it loses electrons
- if it’s more negative when connected to hydrogen , it is better reducing agent
What is a standard electrode potential
THE STANDARD ELECTRODE POTENTIAL, E°, OF A HALF - CELL IS THE VOLTAGE MEASURED UNDER STANDARD CONDITIONS WHEN THE HALF - CELL IS CONNECTED TO A STANDARD HYDROGEN ELECTRODE
Easy to oxidised = very negative potential
Harder to oxidise = less negative / positive potential
Use if the SHE
- you measure the electrode potential of a half - cell against the SHE
- this is to COMPARE THE TENDENCY OF DIFFERENT METALS TO RELEASE ELECTRONS
What is SHE defined as
The potential of SHE is defined as zero
- so if it’s connected to another electrode , the measured voltage ( the electromotive force) is the electrode potential of that cell if it’s at standard conditions
How do they work
- hydrogen gas is bubbled into a solution of H+ (aq) ions
- since the hydrogen doesn’t conduct , electrical contact is made via a piece of unreactive platinum metal
- metal is coated with finely divided platinum to increase the surface area + allow any reaction to proceed rapidly
What are electrodes with negative E° values better at + show cell notation for a SHE
- releasing electrons ( better reducing agents) than hydrogen
Pt | H2 | 2H+ || (whatever is being measured against )
Standard conditions
- Solutions have conc of 1.00 moldm-3
- Temperature if solution is 298k (25°C)
- Pressure is 100 kPa
Changing these conditions will change the electrode potential
Representing cells
- more -ve electrode potential goes left
- when giving the value of the emf state the polarity (I.e. wether it is positive/ negative ) of the right - hand electrode , as the cell representation is written
- charges go closest to salt bridge
Electrochemical series
- a list of standard electrode potentials for electrochemical half - cells
- arranged in this order with the most negative values at the top
- the number of electrons involved in the reaction has no effect on the value of E°
More +ve electrode potentials show that ….
- left hand substance is more easily reduced
- right hand substance is more stable
- best oxidising agent (best at being reduced)
- Ag+ is the best oxidising agent
More -ve electrode potentials show that ….
- right hand substance more easily oxidised
- left hand substance is more stable
- best reducing agent (best at being oxidised)
- Li is the best reducing agent
What is it written as and what do you need to do to it
Written as reduction eq , when working out any potentials , the more -ve value equation needs to be flipped
Calculating standard cell potentials
The voltage obtained by connecting two standard electrodes together is found by the difference between the two E° values .
E° cell = E° reduced - E° oxidised Always more +ve - more -ve
Predicting direction of redox rxns
Using standard electrode potentials we can predict if a rxn is feasible
- electrons move from more -ve to more +ve electrode
Using electrode potentials :
- the species OXIDISED is the REDUCING AGENT , this needs to have a more -ve potential value for it to be feasible
- the species REDUCED is OXIDISING AGENT , needs to have a more +ve potential for it to be feasible
To work out if feasible
- work out which species is being oxidised (loses electrons) - this is positive cell
- work out which species is being reduced (gains electrons) - this is negative cell
- look on electrochemical series at the emf values
- more negative cell electrode potential value ( may have more + value) minus the more positive cell electrode potential ( may have more -ve value)
How to tell feasibility after calculation
Feasible = +ve value
Not feasible = -ve value
Example :
Fe3+ (aq) + Cl- (aq) -> Fe2+ (aq) + 1/2Cl2 (aq)
Iron being reduced so gaining electrons so this is the -ve electrode
Chlorine is being oxidised so losing electrons so is the +ve electrode
Fe3+ (aq) + e- -> Fe2+ (aq) = +0.77
Cl2 (aq) + 2e- -> 2Cl- (aq) = + 1.36
Negative electrode - positive electrode
0.77 - (+1.36) =-0.59
SO NOT FEASIBLE
Types of electrochemical cells
Batteries are a type of electrochemical cell which provide the electricity we use to power things like watches + phones . A battery refers to a number of cells connected together.
RECHARGEABLE
- these can be recharged by reversing the cell rxns
- this is done by applying an external voltage greater than the voltage of the cell to drive the electrons in the opposite direction
NON-RECHARGEABLE
- non -rechargeable batteries are cheaper than rechargeable ones ,however, are worse is the long run as the don’t last longer
FUELL CELLS
Lithium ion rechargeable battery
- positive electrode is made of lithium cobalt oxide ,LiCoO2
- negative electrode is carbon
- these are arranged in layers with a sandwich of solid electrolyte in between
- on charging , electrons are forced through the external circuit from positive to negative electrode + at the same time lithium ions move through the electrolyte towards the positive electrode to maintain the balance if charge
- on discharging, electrons + Li+ ions move from the negative electrode back to the positive electrode
- a single cell gives a voltage of between 3.5 v and 4.0 v , compared with around 1.5 v for most other cells , sed in laptops, tablets , smartphones
Strength
Light ( lithium is a dense material) , won’t leak (electrode is solid polymer)
Weaknesses
Limited life span , loses capacity,expensive , fire / explosion risk
Eq @ +ve electrode
Li+ + CoO2 + e- -> Li+[CoO2]-
Eq @ -ve electrode
Li -> Li+ + e-
Alkaline hydrogen- oxygen fuel cell
-the cell has two electrodes if a porous platinum- based material
- hydrogen + oxygen are fed into 2 separate electrodes
- they are separated by a semipermeable anion membrane which allows OH- and water to pass through but not hydrogen or oxygen gas
- the electrolyte is sodium hydroxide solution
Hydrogen enters at the -ve electrode and the following half reaction takes place :
2H2(g) + 4OH- -> 4H2O (l) + 4e-
EMF = -0.83V
This releases electrons which flow through the circuit from the -ve hydrogen electrode to the +ve oxygen electrode . Oxygen enters + the following half rxn takes place :
O2 (g) + 2H20 (g) + 4e- -> 4OH-
EMF = +0.40 V
What does this do
This accepts electrons from the other electrode + releases OH- ions which travel through the semi - permeable membrane
Final overall effect / eq
2H2(g) + O2(g) -> 2H20 (g)
EMF = 1.23 V
Benefits of fuel cells
- more efficient - convert more available energy to kinetic energy
- takes place at lower temp so no nitrogen produced
- only waste product is water
- no nasty toxic chemicals
- no CO2 emissions from cell itself
- don’t need to be recharged as long as hydrogen and oxygen are supplied
Problems with fuel cells
- need energy to produce hydrogen and oxygen
- can be supplied by electrolysis but this requires energy from burning fossil fuels - co2 emissions
- hydrogen us highly flammable
- hard to store + transport hydrogen gas
- burning 1g of hydrogen gives out around three times as much energy as burning 1g of petrol but the hydrogen takes up around 8000 times as much space
