23. Redox and Electrode Potentials
No description yet. This curated deck is ready for daily study.
Card preview (24)
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