chapter 23 Flashcards
carrying out redox titrations- common titrations
- potassium manganate (VII) (KMnO4(aq)) under acidic conditions
- sodium thiosulfate (Na2S2O3(aq)) for determination of iodine
Manganate (VIII) titration procedure
1) a standard solution of KMnO4 added to the burette
2)using a pipette, add a measured volume of the solution being analysed to a conical flask. Excess of dilute sulfuric acid to provide the H+ ions fro reduction of MnO4-
3)manganate solution is decolourised as it is added, end point= first permanent pink colour
4) repeat until you obtain concordant titres
read from top of meniscus
iodine/ thiosulfate titration equations
thiosulfate ions are oxidised and Iodine is reduced
oxidation: 2S2O3 2- (aq) –> S4O6 2- + 2e-
reduction: I2 (aq) + 2e- –> 2I-
overall: 2S2O3 2- (aq) + I2 (aq) + S4O6 2-(aq)
starch can be used to find the end point- blue/black colour disappears
half cells
contains the chemical species present in a redox half equation. voltaic cell can be made by joining two half cells.
Metal/metal ion half-cells
consists of a metal rod dipped into a solution of its aq metal ion- represented by vertical line eg Zn2+(aq)|Zn(s)
Ion/ion half cells
contains ions of the same element in different oxidation states eg Fe3+(aq) + e- –> Fe2+ (aq)
platinum metal electrode
how do you know which electrode has a greater tendency to gain or lose electrons?
in an operating cell:
- the electrode w/ more reactive metal loses electrons and is oxidised- negative electrode
- the electrode w/ the less reactive metal gains electrons and is reduced- positive electrode
standard electrode potential
- the tendency to be reduced and gain electrons is measured as standard electrode potential.
- the standard is a half-cell containing hydrogen gas and a solution containing H+ ions. platinum electrode is used to allow electrons in/out the half cell
- 298K, 100KPa (1 bar), conc 1 moldm-3
- the standard potential of hydrogen electrode is 0V
measuring a standard electrode potential
- to measure a standard electrode potential the half-cell is connected to a standard hydrogen electrode
-2 electrodes are connected by a wire to allow a controlled flow of electrons - 2 solutions connected by a salt bridge which allows ions to flow. typically contains a conc solution of an electrolyte that does not react w/ either solution. eg strip of filter paper soaked in aq KNO3
equilibrium always shows the forward reaction is reduction
the more negative the E value
- greater the tendency to lose electrons and undergo oxidation
- the less the tendency to gain electrons and undergo reduction
- reducing agent
the more positive the E value
- greater the tendency to gain electrons and undergo reduction
- less tendency to lose electrons and undergo oxidation
- oxidising agent
e value summary
- the more negative the E value , the greater the reactivity of a metal in losing electrons
- the more positive the E value, the greater the reactivity of a non-metal on gaining electrons
E cell =
E(positive electrode) - E(negative electrode)
predicting redox reactions
- the most negative system has the greatest tendency to be oxidised and lose electrons
- the most positive system has the greatest tendency to be reduced and gain electrons
you can predict the feasibility of a reaction from E values
limitations of predictions using E values- reaction rate
- reactions w/ v large Ea- slow rate
- electrode potentials give no indication of rate of a reaction
limitations of predictions using E values- concentration
- if the conc of a solution is not 1 moldm-3 then the value will be different from the standard electrode potential
- eg Zn2+(aq) + 2e- –> Zn(s), if the conc of Zn2+ is greater than 1 moldm-3, the equilibrium will shift to the right, removing electrons from the system and making the potential less negative
- any change to electrode potential will affect the value of the overall cell potential
limitations of using E values- other factors
- conditions may be different from standard
- the electrode potentials apply to aq equilibria, many reactions are not aqueous
primary cells
- non-rechargeable, designed to be used once only
- electrical energy is produced by oxidation and reduction at the electrodes
- reactions cannot be reversed, chemicals will be used up & battery will go flat
- eg wall clocks and smoke detectors- low-current, long-storage devices
secondary cells
-rechargeable
-cell reaction producing electrical energy can be reversed during recharging, chemicals are regenerated
eg lead-acid car batteries, nickel-cadmium NiCd cells, lithium ion cells.
fuel cells
- uses the energy from the reaction w/ oxygen to create a voltage
- fuel & oxygen flow in and products flow out. electrolyte remains in the cell
- do not have to be recharged
hydrogen fuel cells- alkali
oxidation: H2(g) + 2OH-(aq) –> 2H2O(l) + 2e-
reduction: 1/2O2(g) + H2O(l) + 2e- –> 2OH-(aq)
overall: H2(g) + 1/2 O2(g) –> H20(l) Ecell= 1.23V
hydrogen fuel cells- acid
oxidation: H2(g) –> 2H+(aq) + 2e-
reduction: 1/2O2(g) + 2H+(aq) + 2e- –> H2O(l)
overall: H2(g) + 1/2O2 –> H2O(l) Ecell= 1.23V
