5.2.3 - Redox & electrode potentials Flashcards
Constructing half equations when the oxidation no. INCREASES
Delocalised e- goes on the products side
Constructing half equations when the oxidation no. DECREASES
Delocalised e- goes on the reactants side
Half eqns. in acidic conditions
Add H2Os to balance oxygens
H+ to balance H’s
Half eqns. in alkaline conditions
Add H2Os to balance oxygens
OH- to balance H’s
Combining half eqns to get the overall eqn
Balance delocalised e- then combine the two
Delocalised e- should cancel
Redox titrations
Involves transfer of e- from one species to another
Titrations of an oxidising agent against a reducing agent
Use of acidified KMnO4
Purple potassium manganate is in the burette
Sample analysed is in the flask with an excess of dilute sulfuric acid
As the MnO4- ions react, they form Mn 2+ ions which are colourless
Why can’t HCl be used w/ MnO4-
MnO4- would oxidise Cl- to Cl2 and then affect the vol of KMnO4 required in the titration
Why can’t conc. H2SO4/HNO3 be used w/ MnO4-
They are oxidising agents themselves so affect the vol of KMnO4 required
Why can’t ethanoic acid be used w/ MnO4-
It’s a weak acid and would not provide enough H+ ions
Reacting ratio of Fe 2+ : MnO4-
5:1
How is Fe (0) used in the redox titration
Oxidised w/ H2SO4 to Fe (2+) ready for analysis
How is Fe(+3) used in the redox titration
Reacted w/ Zn to reduce it to Fe (2+)
Reacting ratio of C2O4 2-: MnO4-
2.5 : 1
C2O4 2-
Ethanedioate ion
Self indicating
Titration that does not need an indicator
Autocatalysis
A reaction where one of the products acts as a catalyst
Reaction w/ MnO4- and C2O4 2-
As both ions are -ve they repel each other and the reaction is slow and needs warming at the start
However Mn2+ acts as a catalyst and speeds up reaction
Cu^2+/ I-
2 Cu2+ + 4I- —-> 2CuI + I2
CuI - white ppt
ClO-/I-
2 I- + ClO- + 2H+ —> Cl- + I2 + H2O
I2/S2O3 2-
2 S2O3 + I2 —> 2 I- + S4O6 2-
Reacting ratio Cu2+:I2
2:1
Reacting ratio of ClO- : I2
1:1
Reacting ratio of S2O3 2- : I2
2:1
Electrochemistry
Control of transfer of e- to produce electrical energy
Standard electrode potential
The emf of a half cell compared to the standard hydrogen electrode, measured at 298 K w/ sol. conc. of 1 moldm3 and gas pressure of 101 kPa
Strong reducing agents have …
more -ve electrode potentials and so get oxidised
Strong oxidising agents have …
More +ve electrode potential and so get reduced
Calculating cell potential
More +ve electrode potential - more -ve electrode potential
Where does oxidation occur
Anode (+ve)
Where des reduction occur
Cathode (-ve)
How are half cells eqns written
Showing the reduction so e- are on the lhs
Species on the lhs are oxidising agents and species on the rhs are reducng agents
When is a cell reaction feasible
When the cell potential is +ve
Why may a feasible reaction not occur
HIgh Ea or conditions aren’t standard/ diff conc
Explaining changes in a cell
….. affects which electrode
The eqm shifts to the … in order to ..
… accepts/ donates more e-
The electrode potential is more +ve/-ve and the cell potential is more/less +ve
What is the salt bridge soaked in and why
KNO3
Highly soluble
Easily enters solutions to balance charges
Unlikely to form ppts
How do electrons travel in a cell
In the wire
How do ions travel in a cell
Through the salt bridge
Metal used in half cells
Platinum (s)
Features of primary cells
Non-rechargebale
Can only be used once as the redox reactions arent reversible
Chemicals get used up –> pd falls –> battery goes flat —> discarded
Used in low-current, long-storage devices e.g. wall clocks
Redox system in most primary cells
Zn/MnO2 (alkaline)
Both involve H2O and OH- (cancel out)
MnO2/ Mn2O3 - more +ve, OA, cathode
Zn/ZnO - more -ve, RA, anode
Feautures of secondary cells
Rechargable - each reaction is reversed during recharging, regenerating chemicals
Li + in modern devices e.g. phones, tablets
Nickel/cadmium and nickel metal hydride in torches
Fuel cells
Uses energy from reactions w/ O2 to create a pd (combustion)
Fuel and O2 flow into fuel cells and produts out. Electrolyte remains in cell
Can operate continously and dont have to be recharged
Hydrogen fuel cells
Can be acid or alkali (generate same potential difference - same overall reaction)
Form no CO2 during combustion, only H2O
Acid hydrogen fuel cells
H2 enters at anode
H+ moves to cathode
O2 enters at cathode
H2O leaves at cathode
Alkali hydrogen fuel cells
H2 enters at anode
H2O leaves at anode
OH- moves to cathode
O2 enters at cathode
Negative of hydrogen fuel cell
H2 is a gas and difficult to store
What’s the purpose of a salt bridge
Separates the solution
Complete the electrochemical circuit
Balance charges by releasing K+ and NO3- into the different solutions
Identifying the +ve electrode using voltmeter readings
If +ve, electrode connected to +ve terminal (reduction eqn)
If -ve, electrode connected to -ve terminal (oxidation eqn)
Why does the emf of a cell change once connected to electrodes w/ a flowing current
Conc of ions change
Function of the platinum electrode
Allow transfer of electrons
Reaction surface
Why may a redox reaction not occur in the absence of light
High Ea
Light breaks bonds
Properties of platinum that makes it suitable as an electrode
Inert
Conductor
Reduction of H2O2
2H+ + H2O2 +2e- —-> 2H2O
pH of SHE
0
Difference between fuel cell and modern storage cell
Fuel reacts w oxygen to give electrical energy
Risks with Li based cells
Toxicity
Fire
