CHEMISTRY KEYWORDS:
Electrolysis: Decomposition of an ionic compound when
molten or in aqueous solution by an electric current.
Electrolyte: A molten ionic compound or an aqueous
solution of ions that is decomposed during electrolysis.
Electrode: A rod of metal or carbon (graphite) which
conducts electricity to or from an electrolyte.
Cathode: Negative electrode (where reduction
reactions occur)
Anode: Positive electrode (where oxidation reactions
occur).
Faraday: Quantity of electric charge (in coulombs)
carried by 1 mole of electrons or 1 mole of singly charged ions.
Electrode Potential, E: Voltage measured for a
half-cell compared with another half-cell.
Feasible: Likely to take place. (if E standard has a positive
value)
Discharged: Ions changed into atoms or molecules.
CHEMISTRY DEFINITIONS:
Standard Electrode Potential: Voltage produced when a
standard half-cell (ion concentration 1 mol/dm3 at 298 K) is
connected to a standard hydrogen electrode under standard conditions.
Standard Cell Potential: Difference in standard
electrode potential between 2 specified half-cells.
CHEMISTRY IMPORTANT NOTES:
Electrolysis is a redox reaction. Electrons
may be gained or lost in redox reactions.
Species losing
electrons is being oxidised and acts as a reducing agent. Oxidation always
occurs at the anode (negative terminal of the cell).
Species gaining
electrons is being reduced and acts as an oxidising agent. Reduction always
occurs at the cathode (positive terminal of the cell).
Electrolysis is
often used to extract metals that are high in the reactivity series. It is also
used to produce non-metals and to purify some metals. It is generally carried
out in an electrolysis cell.
Mass of a substance
produced at an electrode during electrolysis is proportional to:
·
Time over which a constant electric current
passes.
·
Strength of the electric current.
Mass of a substance
produced at an electrode during electrolysis is proportional to the quantity of
electricity (in coulombs) which passes through the electrolyte.
Avogadro constant,
L, is the number of specified particles in 1 mole.
Redox equilibrium
exists between 2 chemically related species that are in different oxidation
states. It is established when the rate of electron gain equals the rate of
electron loss.
When a metal is
placed into a solution of its ions, we establish an electric potential (voltage)
between the metal and the metal ions in solution. We can measure the difference
in potential between the metal/metal ion system and another system. Electrode
potential is measured in volts. The system used for comparison is the standard
hydrogen electrode.
The standard hydrogen
electrode is one of several types of half-cells used as reference electrodes.
Hydrogen gas at 101 kPa, and a concentration of 1.00 mol/ dm3. A
platinum electrode covered with platinum black is in contact with hydrogen gas
and H+ ions. The platinum electrode is inert, so it does not take part in the
reaction.
Electrode potential
refers to a reduction reaction. So electrons appear on the left-hand side of the
half-equation.
The more positive (or
less negative) the electrode potential, the easier it is to reduce the ions on the
left. Metal on the right is relatively unreactive and is a poor reducing agent.
The more negative
(or less positive) the electrode potential, the more difficult it is to reduce
the ions on the left. Metal on the right is relatively reactive and is a good
reducing agent.
Standard conditions
to make a half-cell:
·
Ions have a concentration of 1.00 mol/dm3.
·
Temperature is 25 °C (298 K).
·
Rod must be pure.
Half-cells are
connected using:
·
Wires connected the metal rods in each half-cell
to a high-resistance voltmeter. Electrons flow around this external circuit
from the metal with the more negative (or less positive) electrode potential to
the metal with the less negative (or more positive) electrode potential.
·
Salt bridge to complete the electrical circuit,
allowing the movement of ions between the 2 half-cells so that ionic balance is
maintained. It does not allow the movement of electrons.
Position of
equilibrium of a reaction may be affected by changes in concentration of reagents,
temperature and pressure of gases. The voltage of an electrochemical cell will
also depend on these factors.
Standard electrode
potential is sometimes referred to as standard reduction potential because it refers
to the reduction reaction (addition of electrons).
The standard
electrode potential for a half-cell is the voltage measured under standard
conditions with a standard hydrogen electrode as the other half-cell.
The 3 main types of
half-cells can be connected to a standard hydrogen electrode:
·
Metal/metal ion half-cell
·
Non-metal/non-metal ion half-cell
·
Ion/ion half-cell
In half-cells that
do not contain a metal, electrical contact with the solution is made by using a
platinum wire or platinum foil as an electrode. Platinum must be in contact
with both the element and the aqueous solution of its ions.
Standard electrode
potential values measure how easy or difficult it is to oxidise/reduce powers
of elements and ions by comparing the E standard values for their half-reactions.
·
The more positive the value of E standard, the greater
the tendency for the half-equation to proceed in the forward direction. Easier
to reduce the species on the left of the half-equation.
·
The less positive the value of E standard, the greater
the tendency for the half-equation to proceed in the reverse direction. Easier
to oxidise species on the left of the half-equation.
A reaction is
feasible if it is likely to occur. If forward reaction is feasible, the reverse
reaction is not feasible.
·
Species on the left of the equation become
weaker oxidising agents, as they accept electrons less readily.
·
Species on the right of the equation become
stronger reducing agents, as they release electrons more readily.
Cations (positive
ions) move towards the cathode, where they gain electrons; the gain of
electrons is reduction. Anions (negative ions) move towards the anode, where
they lose electrons; loss of electrons is oxidation.
SO42-
< NO3- < Cl- < OH- < Br-
< I- (Increasing ease of discharge, and oxidation)
The proportion of
oxygen increases, the more dilute the solution.
SUMMARY:
Q = It (Charge =
Current x Time)
1 F = 96500 C / mol
F = Le (F = Faraday
Constant, L = Avogadro Constant, e = Electron Charge)
L = (charge on 1
mole of electrons) / (1.6 x 10-19 C)
During electrolysis,
reduction occurs at the cathode (negative electrode) because ions gain
electrons from the cathode, and oxidation occurs at the anode (positive
electrode) because ions lose electrons to the anode.
A metal is formed
at the cathode and a non-metal at the anode when molten metal salts are
electrolysed. Hydrogen may be formed at the cathode and oxygen at the anode
when a dilute aqueous solution of metal salts is electrolysed.
A standard hydrogen
electrode is a half-cell in which hydrogen gas at a pressure of 101 kPa bubbles
through a solution of 1.00 mol/dm3 H+ ions.
The standard
electrode potential of a half-cell is the voltage of the half-cell under
standard conditions compared with a standard hydrogen electrode.
The direction of
electron flow in a simple cell is from the half-cell that has the more negative
(or less positive) electrode potential to the half-cell that has the less
negative (or more positive) electrode potential.
A particular redox
reaction will occur if the E standard of the half-equation involving the
species being reduced is more positive than the E standard of the half-equation
of the species being oxidised.
No comments:
Post a Comment