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REDOX II:
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Loss of electrons – oxidation – ON increases
Gain of electrons – reduction – ON decreases
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Electrochemical cells:
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Made from 2 different metals dipped in salt
solutions of their own ions and connected by a
wire
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2 half-cells make a complete cell
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Electrons flow from anode to cathode
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A voltmeter in the external circuit shoes the
voltage between the two half-cells
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They also
measure the flow of electrons
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Half cells can have solutions of two aqueous ions of the same element
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For that electrode must be made
from inert conducting material i
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platinum/graphite
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Reduction on right
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Method of setting up electrochemical cell:
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Clear each strip of metal using sandpaper
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Place each electrode into beaker containing the solution of ions of that metal
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Can be done by
soaking a piece of filter paper in salt solution and draping it between the two
beakers
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The salt bridge usually
contains conc
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It allows
the movement of ions
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Standard Electrode potentials:
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Each half cell has own electrode potential – Measure of

how easily the substance in the half-cell is oxidised
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Reaction with more positive electrode potential goes
forward
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The standard electrode potential – of the half-cell is the voltage

measured under standard conditions when the half-cell is
connected to a standard hydrogen electrode
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00 mol dm-3 conc
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00V
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Ecell = (Ereduction – Eoxidation)
Conditions affect the value of electrode potentials:
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Temperature
Pressure
Concentration

Measuring standard electrode potentials:
Systems involving gases:
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E
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of 1/2Cl2 + e- ↔ ClSet a half-cell where chlorine gas is bubbled into a solution containing chloride
ions
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It has a set up like hydrogen
electrode
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The
conc
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Oxidised forms in centre
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Double vertical lines – salt bridge, single – different physical state
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Inert electrodes shown on the outside
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More reactive non-metal – more easily gains electrons, has more positive
standard electrode potential
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The electrode potentials can be used to predict if disproportionation and other
reactions will happen
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The activation energy is too high
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Bigger cell potential = bigger total entropy change
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TΔS=nFEcell
Cell potential is directly proportional to ln (k) - equilibrium constant
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They are used in:
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Torches, battery operated radios, electronic calculators, mobile phones
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The chemicals are stored separately outside the cell
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o H and O are fed into two separate platinumcontaining electrodes
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The electrolyte is an aqueous alkaline (KOH)
solution
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5O2 + 2e- → 2OHO cell works in acidic conditions:
At the anode the platinum catalyst splits the H2 into protons and
electrons
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Electric current is created which is used to power appliance
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Only waste
product
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5O2 +2H+ + 2e- → H2O – at positive electrode

Some fuel cells can use alcohols: methanol, ethanol
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Advantages:
o Alternative to fossil fuels
o Do not produce pollutants
o Lighter and more efficient
Disadvantages:
o Hydrogen is explosive
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o Adsorbing onto a metal surface – not sure which one is the best
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to release hydrogen
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New sources have to be used
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Titration of potassium manganate (VII) with iron (II) ions
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Ionic equations:
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MnO4- + 8H+ + 5e- → Mn2+ + 4H2O
Fe2+→Fe3+ + e-

Overall equation:
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MnO4- + 8H++ 5Fe2+→ 5Fe3++ Mn2+ + 4H2O

Method:
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Pipette a solution of iron (II) into a conical flask and add small volume of
dilute sulfuric acid
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from burette
and swirl the mixture
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The addition of one drop of potassium manganate (VII) solution will turn the
mixture pale pink
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They form iodine and water
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As colour fades to pale yellow add starch
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Half equations:
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2S2O32- → S4O62- + 2eI2 +2e- → 2I2S2O32- + I2 → 2I- + S4O62-

Source Errors in Titrations:
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The starch indicator has to be added at the right point or blue colour will take
long to disappear
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Precipitate of copper (I) iodide makes seeing the colour of solution hard
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Keeping solution cool helps with that

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