23) Redox and electrode potentials Flashcards
What two things will there always be in a redox reaction?
oxidising agent
reducing agent
Define oxidising agent
a reagent that oxidises (takes e- from) another species
Define reducing agent
a reagent that reduces (adds e- to) another species
What should you do when writing an overall equation for redox reactions?
balance the changes in oxidation number and then balance any remaining atoms afterwards
If one species in an overall redox equation has an oxidation number change of +6, and a second species has an oxidation number change of -1, what should be done?
multiply second species by 6 to give -6
Describe manganate (VII) titrations
a standard solution of KMnO4 is added to burette with excess of dilute sulfuric acid
as it reacts it is decolourised from deep purple colour until end point = permanent pink colour - indicating an excess of MnO4 - ions (self-indicating)
What is the purpose of excess of dilute sulfuric acid in a manganate (VII) titration?
provides H+ for the reduction of MnO4 - ions
Why are burette readings taken from the top of the meniscus in a manganate (VII) titration?
manganate (VII) solution is a deep purple colour
What can manganate (VII) titrations be used for?
- analysis of reducing agents e.g. Fe 2+ (analysing the purity of an iron (II) compound) or ethanedioic acid (COOH)2
- reduce MnO4 - to Mn 2+
What can iodine/thiosulfate titrations be used to determine?
- chlorate ions, ClO- content in household bleach (HCl added to acidify solution)
- Cu 2+ content in copper (II) compounds
- Cu content in copper alloys (dissolved in conc. nitric acid + neutralisation to produce Cu 2+)
Describe iodine/thiosulfate titrations
- a solution of oxidising agent to be analysed is added to a conical flask with an excess of potassium iodide - iodine produced resulting in a yellow-brown colour
- titrated with sodium thiosulfate solution (S2O3 2- oxidised and I2 reduced)
- when end point is approached, sol. has faded to a pale straw colour, a small amount of starch indicator is added -> blue-black solution which turns colourless at end point
Define voltaic cell
a type of electrochemical cell which converts chemical energy into electrical energy
What does a half cell contain?
the chemical species present in a redox half-equation
How is a voltaic cell made?
by connecting together 2 different half-cells allowing e- flow
Describe metal/metal ion half-cells
a metal rod dipped into a solution of its aqueous metal ion
at the phase boundary where the metal is in contact with its ions, an equilibrium is set up e.g. Zn 2+(aq) + 2e- <=> Zn(s)
Describe ion/ion half-cells
contains ions of the same element in different oxidation states
there is no metal to transport e- into/out of the half-cell so an inert metal electrode made out of platinum is used
What happens in an operating cell with two metal/metal ion half-cells connected?
the electrode with more reactive metal loses e- and is oxidised
the electrode with the less reactive metal gains e- and is reduced
Define standard electrode potential E⦵
the emf of a half-cell compared with a standard hydrogen half-cell under standard conditions
What is the standard electrode potential of a standard hydrogen half-cell under standard conditions?
0V
How is standard electrode potential measured?
the half-cell is connected to a standard hydrogen electrode by a wire to allow controlled flow of e-
recorded by voltmeter
Describe a standard hydrogen half-cell
platinum electrode
H2(g) pumped in
acid sol. 1moldm-3 H+(aq)
What does a salt bridge typically contain?
a concentrated solution of electrolyte that does not react with either solution e.g. a strip of filter paper soaked in KNO3(aq)
How is the equilibrium for standard electrode potentials shown?
with the forward reaction being the reduction
How are electrode potentials sorted in a data reference table?
in order of E⦵ - most negative at top
The more negative E⦵ value, the greater the tendency to _ e- and undergo _?
lose
oxidation
The more positive E⦵ value, the greater the tendency to _ e- and undergo _?
gain
reduction
E⦵cell = ?
E⦵(positive electrode) - E⦵(negative electrode)
What should you remember when predicting redox reactions from a data reference table?
top-right, bottom-left rule
the strongest reducing agent is at the top right
the strongest oxidising agent is at the bottom left
What should you remember when writing an overall equation from half-equations in a data reference table?
the oxidation half-equation is obtained by reversing the equilibrium
What are the limitations of predicting redox reactions from a data reference table? (4)
- gives no indication of rate of a reaction
- if concentration is not equal to 1moldm-3 electrode potential will be different to standard value
- may not be under standard conditions
- many reactions are not aqueous (aqueous equilibria)
Describe primary cells
- non-rechargeable
- electrical energy is produced by reduction and oxidation at electrodes (irreversible reactions)
- eventually chemicals will be used up, voltage decreases and battery goes flat
- alkaline (Zn/MnO4 and potassium hydroxide alkaline electrolyte, Zn oxidised -> ZnO and Mn in MnO2 reduced -> Mn2O3)
Describe secondary cells
- rechargeable - cell reaction can be reversed during recharging so chemicals are regenerated
- lithium-ion and lithium ion polymer cells (when in use e- flow from negative -> positive electrode and Li+ from anode to cathode to keep electrical charge balance)
Describe fuel cells
- uses the energy from the reaction of a fuel with oxygen to create voltage
- fuel and oxygen flow into cell and products flow out, electrolyte remains
- operate continually provided fuel and O2 supplied
- do not have to be recharged
Give the overall equation for hydrogen fuel cells and the E⦵cell
H2(g) + 1/2 O2(g) -> H2O(l)
E⦵cell = 1.23V
Suggest two advantages of using methanoic acid as the fuel in a fuel cell rather than hydrogen
liquid - easier to store and transport
greater cell potential / voltage
hydrogen is flammable
