Redox And Electrode Potential Flashcards
Reducing agents
Donates electrons and decreases oxidation number of another species.
Oxidising agent
Accepts electrons and increase oxidation number of another species.
Standard electrode potential (E•)
E.m.f of a half-cell compared with a standard hydrogen half cell
- 1 mol dm-3
- 298K
- 100kpA
- Platinum electrode
Salt bridge
Allows ions to move
Wire
Allows electrons to flow
What is the process taking place at the positive and negative electrodes?
Electrons flow away from the least positive electrochemical cell and flows to the most positive.
- Least positive loses electrons, most positive gains
What does a more negative electrode potential tell us ?
- Greater tendency to lose electrons and undergo oxidation
- Less tendency to gain electrons and undergo reduction
- Equilibrium shifts to the left
Cell potential equation
(most positive)-(least positive)
What value do all feasible electrode potential reactions have?
+ V
When will a reaction occur
If the oxidising agent has a more positive value than the redox system of the reducing agent
Steps of comparing feasibility in equations
- identify which is being oxidised
- Reverse oxidised equation
- Combine 2 equations to obtain feasible reaction
- Compare equation to one stated in question
Limitations of predicting feasibility (concentration)
Non standard conditions alter value for E, half equations are at equilibria so changes in concentration will shift the position which affects electron transfer
Limitations of predicting feasibility (rate of reaction)
Reaction rate may be very slow due to a high activation energy and electrode potential give no indication of rate of reaction.
Limitations of predicting feasibility: Aqueous solutions
Standard electrode potentials apply to aqueous equilibria but some reactions that may take place are not aqueous.
Three types of electrochemical cells
- Non rechargeable cells (1• non-reservable)
- Rechargeable cells (2• reversible)
- Fuel cells