23. Redox and Electrode Potentials Flashcards
How do you balance half equations in basic solutions?
balance thing being oxidised/reduced
deduce e- needed from oxidation state change
balance on other side w OH-
balance on other side w H2O
What type of reactions happen in batteries (voltaic cells)?
chemical reactions that transfer e- from one species to another
When is a platinum electrode used?
when there is a mixture of 2 different ions of the same element or a gas and a solution of its ion
is used to transport the e-
What does the standard electrode potential represent?
the tendency for the species in a half cell to either gain or lose electrons
Define standard electrode potential.
the electromotive force of a half cell compared to the standard hydrogen electrode measured at 298K, solution conc. 1 moldm-3 and at a pressure of 100kPa
Draw the standard hydrogen half cell connected to a Cu half cell.
Page 3 of electrode potentials leaflet
What is a salt bridge and what is its purpose?
connects the 2 half cells
made of KNO3 as all compounds formed are soluble and it will not react w other solutions
allows ions to transfer between half cells to keep them neutral
What component connects the 2 electrodes in each half cell other than the salt bridge?
high resistance voltmeter: measures potential for electron transfer
In which direction are all half equations written when showing their E(theta)?
reduction half equations
What does decreasing E(theta) mean?
the more negative E is, the greater its tendency to lose e- and go in the reverse direction (i.e. species on the right gets oxidised)
more negative = species on right is strong reducing agent
What does increasing E(theta) mean?
the more positive E is, the greater its tendency to gain e- and go in the forward direction (i.e. as is written, species on left gets reduced)
more positive = species on the left is strong oxidising agent
How do you calculate overall cell potential (E(cell))?
E(most +ve electrode) - E(most -ve electrode)
If E of half cell 1 = 0.34V and E of half cell 2 = -0.11V which will be the forwards reaction?
half cell 1 forwards: species on left is strong oxidising agent
half cell 2 backwards: species on right is strong reducing agent
At which electrode does oxidation take place?
anode: negative electrode: in half cell with most -ve E
At which electrode does reduction take place?
cathode: positive electrode: in half cell with most +ve E
At which electrode are electrons produced?
the electrode where oxidation is taking place: the negative electrode
Which direction do electrons flow?
from the -ve (where they are produced) to the +ve electrode
Zn2+(aq) + 2e- <=> Zn(s)
E(theta) = -0.76V
What will happen if the concentration of Zn2+ is increased?
equilibrium shifts right, removing e- from the system
electrode potential becomes less negative
Zn2+(aq) + 2e- <=> Zn(s)
E(theta) = -0.76V
What will happen if the concentration of Zn2+ is decreased?
equilibrium shifts left, increasing the number of e- in the system
electrode potential becomes more negative
What happens if standard electrode potentials are not measured at 1 moldm-3?
electrode potential will not be accurate: will increase/decrease depending on whether conc has increased/decreased
What happens if the E(cell) is less than 0.06V?
reaction very unlikely to proceed
Why may a reaction not be able to go ahead even if the electrode potential indicate that it could?
is thermodynamically feasible but reaction rate may be very slow i.e. catalyst needed
How do changes in E(theta) due to concentration affect E(cell)?
may increase or decrease the difference between the 2 E(theta)s of the half cells
will alter E(cell) accordingly
What is a primary cell?
non-rechargeable: reactions are not reversible: stops working when the chemical reacting at the electrodes run out
used in clocks
usually Zn and MnO2 reactants with KOH electrolyte
What is a secondary cell?
rechargeable: reaction is reversible and reverse reaction occurs while recharging
used in lithium ion batteries in phones
How does a fuel cell work?
hydrogen rich fuel reacts w oxygen generating e- and a current
fuel and oxygen flow into the cell, electrolyte allows transfer of ions + products flow out
only water produced as product (in pure H2 cell)
What are the advantages and disadvantages of H fuel cells?
AD: clean, quiet, 2/3x more efficient than burning fuel
DIS: production involves toxic chemicals that need to be disposed of, cells are highly flammable
What is the difference between using an alkaline vs acid electrolyte in a fuel cell?
in both cases: O2 in, fuel in, water out, E(cell) = 1.23V, there is a membrane separating half cells
but have different redox systems
What are the hydrogen oxygen fuel cell equations with an acidic electrolyte?
2H+ + 2e- <=> H2 : flip: becomes oxidation equation
1/2 O2 + 2H+ + 2e- <=> H2O : reduction
overall: 1/2 O2 + H2 <=> H2O
H+ travel to and water leaves from the oxygen electrode (cathode)
What are the hydrogen oxygen fuel cell equations with an alkaline electrolyte?
2H2O + 2e- <=> H2 + 2OH- : flip: become oxidation equation
1/2 O2 + H2O + 2e- <=> 2OH- : reduction
overall: 1/2 O2 + H2 <=> H2O
OH- travel to and water leaves from the hydrogen electrode (anode)
Practicals: iodine/thiosulphate redox titrations
What are the reduction, oxidation and overall equations for the titrations?
oxidation: 2(S2O3)2- (aq) => (S4O6)2- (aq) + 2e-
reduction: I2(aq) + 2e- => 2I- (aq)
overall: 2(S2O3)2- + I2 => (S4O6)2- + 2I-
What is the I/THIO reaction used to analyse?
ClO- in bleach
Cu2+ in copper (II) alloys
Cu content in copper alloys (Cu converted to Cu2+ first)
How do to I/THIO reactions work?
the chemical under analysis is reacted with excess KI to produce iodine (yellow-brown solution)
excess = complete reaction
solution titrated against Na2S2O3 and yellow-brown fades as I2 reduced back to I-
when solution is pale straw colour, add starch indicator, turns blue-black
when end point is reached, blue-black just disappears
What happens if you add starch too early in an I/THIO titration?
solid black complex forms, titre is too low
How would you prepare an insoluble copper sample for analysis?
react it with concentrated acid, then neutralise to form Cu2+ ions
(then react w excess KI, titrate, etc.)
What is the reducing agent in the I/THIO reactions?
(S2O3)2- = reducing agent when reacting with I2
Practicals: manganate (VII) redox titrations What does (MnO4)- act as?
oxidising agent
Why does (MnO4)- (and also dichromate) need to be acidified with excess H2SO4?
provides a source of H+ ions which allows the reaction to occur
and provides an excess to prevent other products from forming
What colour is manganate before and after reduction?
manganate: deep purple
Mn2+ : very pale pink
self indicating titration
What is the end point of a manganate titration?
when the first permanent pale pink solution is reached
What is KMnO4 used to analyse?
ethanedioic acid (oxidised to CO2)
What is the colour change of (Cr2O7)2- when reduced?
(Cr2O7)2- : orange
Cr3+ : green
cannot self indicate, green is too strong
Where do you read the meniscus from in manganate titrations?
the top, as the deep purple colour is too difficult to see through
What is the process of a manganate titration?
dissolve compound and make up to 250cm^3 solution in volumetric flask
add 25cm^3 to a conical flask w 10cm^3 H2SO4 ( an excess)
fill burette with standard solution of KMnO4 + titrate
redox: more positive = greater tendency to undergo…
reduction
non-metals tend to be more positive
redox: more negative = greater tendency to undergo…
oxidation
metals tend to be more negative
What are the limitations of electrode potential predictions?
if cell operates under non standard conditions, actual electrode potentials will be different to E theta values
as the cell operates E theta values change
activation energy could be too high or rate of reaction very slow
Why is using a lower temperature for a reaction beneficial to the environment?
less fossils fuels/energy used
reduction in CO2 emissions and global warming
