Redox and Electrode Potentails A2 Flashcards
Why do you not need to add an indicator to a Manganate (VII) titration?
Manganate (VII) solution is purple. It changes to colourless when reduced to Mn2+.
What colour change do you look for when adding manganate solution?
Colourless to permanent pink.
What is different when reading the burette solution with a manganate titration?
Read from top of meniscus because dark purple colour means you can’t read the bottom of the miniscus.
In a thiosulphate-iodine titration, what is the initial colour change?
Orange-brown fades to yellow as iodine is used up.
What indicator is used at the end of a thiosulphate-iodine titration?
When do you add it?
Colour change?
Starch solution
Add only when faded to pale yellow colour.
Solution turns blue/black in I2 which disappears at end point in I-.
What are the 3 electrochemical cells?
- Metal / metal ion half cells
- non-metal / non-metal ion half cells
- Metal ion / metal ion half cells
What does a half cell compromise of?
An element in 2 different oxidation states.
How do you set up a non-metal/non-metal ion half cell? (specifically hydrogen)
A glass tube with holes is used to allow bubbles of gas (non-metal - e.g. Hydrogen) to escape into the solution, either side of the platinum electrode.
Solution has 1 moldm^3 H+ from HCl
Which electrode is used to measure standard electrode potentials of other half cells?
What are the conditions for standard electrodes?
Standard hydrogen electrode
(1 mol dm-3 H+ conc - HCl solution)
Hydrogen at 298K (25 C), 100kPa, Platinum electrode for inert electrodes, 1 mol dm-3 conc solution
Which electrochemical cell do you not use a platinum electrode?
In the metal/metal ion half cells:
The metal is used as the rod and the solution contains the metal ions.
How are two half cells connected to make a cell?
- Wires with a voltmeter between electrodes
- A salt bridge between solutions
What is a salt bridge?
Concentrated solution of an electrolyte soaked into filter paper to connect two half cells.
Electrolyte = a susbstance that dissociates in water to form ions.
What occurs between metal and metal ions/ non-metal and non-metal ions / metal ions in half cells?
Equilibrium
If a standard electrode potential is very negative/positive, what does that mean for the likelihood of gaining/loosing electrons? Why is this?
Very positive = Likely to gain electrons / be reduced - equilibrium shifts right.
BECAUSE stronger oxidising agent.
Very negative = Likely to lose electrons / be oxidised - equilibrium shifts left.
BECAUSE stronger reducing agent.
How do you work out the cell potential from electrode potentials of half cells? When is the reaction feasible?
Eθ of more positive electrode - Eθ of more negative electrode
Feasible if cell potential is positive (opposite to ΔG).
Give 2 limitations of predictions of feasibility from electrode potentials.
1) Reaction may not take place because activation energy too high
2) Reaction may not take place with 1 mol dm-3 solutions
What are the limitations to standard electrode potentials (Eθ)? (4 limitations)
- Non-standard conditions alter the value of Eθ.
- Half equations are equilibia so changes in conc will shift the position, which alters electron transfer.
- Reaction rates may be slow due to high activation energy.
- Not all reactions take place in aqueous solution.
What is the difference between primary and secondary cells?
Primary cells = non-rechargable - chemical used up
Secondary cells = rechargable
What does a fuel cell do? Describe them. What are the limitations?
Use energy from reactions with fuel and oxygen to produce a voltage.
- Reactants flow in and out - electrolyte remains in the cell.
- Can operate continuously as long as fuel and O2 provided.
- Do not have to be recharged.
LIMITATIONS
- Reactions will generally only take place if the difference between Eθ values is large (over 0.4v).
How do you work out half equations from electrochemical series?
Write the equation with the more positive Eθ on the top and the equation with the more negative Eθ backwards, on the bottom.
(Reaction is clockwise)
How do you construct a redox equation from half equations?
Identify which is reduction and which is oxidation.
Multiply equations so number of electrons are equal.
Add/merge them together.
How do you construct a balanced redox equation from a reaction (reactants and products)?
e.g. 2HI + H2SO4 –> H2S + I2
Identify changing oxidation numbers and total changes.
I (2x-1 –> 2x0) total = +2 S (+6–>-2) total = -8
Balance the total increase and decrease in oxidation number.
I x 4 - total = +8
8HI + H2SO4 –> H2S + 4I2
Balance the H and O
8H and 4O = 4H2O
8HI + H2SO4 –> H2S + 4I2 + 4H2O
What will many redox equations require the addition of to balance them and when?
Acidic conditions = H+
Alkaline conditions = OH-
If H and O don’t balance = H2O
