9.2 Flashcards
electrochemical cells are applications of
the fact that redox reactions involve transfers of electrons that immediately suggests a link between this type of chemical reactivity and eelectricity
types of electrochemical cells
Voltaic (galvanic) cells - generate electricity from chemical reactions
Electrolytic cells - drive chemical reactions using chemical energy
Difference between carrying out a redox reaction normally and in a voltaic cell
When the reaction is carried out in a single test tube the electrons flow spontaneously from the reducing agent to the oxidizing agent in the solution and energy is released in the form of heat. (specific case of Zn and Cu)
There is however a different way of organizing this reaction so that the energy released in the redox reaction instead of being lost as heat is available as electrical energy. It is really just a case of separating the two half equations into so called half-cells and allowing the electrons to flow between them only through an external circuit.
SImplest type of cell
Putting a strip of metal into a solution of its ions
What happens in the half-cell of a reducing agent
reducing agent’s atoms will form ions by releasing electrons that will make the surface of the metal negatively charged with respect to the solution causing a charge separation, known as an electrode potential, between the metal and its ions in solution. At the same time, ions in the solution gain electrons to form reducing agent’s atoms, so an equilibrium exists.
position of this equilibrium determines the size of the electrode potential in the half cell and depends on the reactivity of the metal
in general how are metal reactivity and electrode potential in its half cell linked
The more reactive a metal, the more negative its electrode potential in its half-cell
electrodes in Voltaic cell
Cathode - positive; Reduction
Anode - negative; oxidation
*metals higher on activity series at anode ( better reducing agent is oxidized)
electrons flow from the anode to cathode
external electronic circuit in voltaic cell
electrons flow from the anode to cathode through the external electronic circuit connected to the metal electrodes - a voltmeter can also be attached to record the voltage generated
salt bridge in voltaic cell
ion movement through salt bridge - glass tube or strip of absorptive paper that contains an aqueous solutions of ions
movement of ions neutralises any build up charge and maintains the potential difference
anions move from the cathode to the anode
cations move from the anode to the cathode
solution chosen in the voltaic cell is often
aqueous NaNO3 or KNO3 as these ions will not interfere with the reactions at the electrodes
no salt bridge =
no voltage generated
cell diagram convention
*a single vertical line represents a phase boundary such as that between a solid electrode and an aqueous solution within a half-cell
*a double vertical line represents the salt bridge
*the aqueous solutions of each electrode are placed next to the salt bridge
*the anode is generally put on the left and the cathode on the right, so electrons flow from left to right
*spectator ions are usually omitted from the diagram
*If a half cell includes two ions, they are separated by a comma because they are in the same phase
E.g.
ZN(s) | Zn2+ (aq) ||Cu2+ (aq) | Cu(s)
anode cathode
oxidation reduction
———————>
electrons flow this way via external circuit
Zn(s) –> Zn2+ (aq) + 2e-
Cu2+ (aq) + 2e- –> Cu(s)
Electromotive force (EMF)
a potential difference generated by a voltaic cell
magnitude of this voltage depends on the difference in the tendencies of these two half-cells to undergo reduction
why do we need standard hydrogen electrode (SHE)
clearly, the electrode potential of a single half-cell cannot be measured in isolation, but only when electrons flow as it is linked in this way to another half-cell. Therefore, in order to draw up a list of the relative reducing power of different half-cells, it is necessary to compare them all with some fixed reference point that acts as a standard for measurement.
In electrochemistry, the reference standard is the standard hydrogen electrode
The standard hydrogen electrode
sometimes called the standard hydrogen half-cell
half cell that has an electrode potential value of 0 V
therefore while connecting it to another half cell it will give us a single electrode potential of that half cell
platinum electrode is put into a glass tube with holes (to allow bubbles of H2(g) to escape) containing acid solution with 1.0 mol dm-3 H+(aq)
platinum is used as the conducting metal in the electrode because it is fairly inert and will not ionize - and it also acts as a catalyst for the reaction of proton reduction
concentration of H+ (aq) is 1.0 mol dm-3
pH=0
Pressure of H2(g) is 100kPa
298K
2H+ (aq) + 2e- ⇌ H2 (g)
reaction is reversible occurring as reduction of H+ ( forward reaction) or as oxidation of H2 ( backward), depending on the electrode potential of the half cell to which it is linked
standard half-cells
half cells under standard conditions
standard electrode potential
EMF generated when it is connected to standard hydrogen electrode by an external circuit and a salt bridge measured under standard conditions
standard electrode potentials are given for
the reduction reaction
points regarding standard electrode potentials
*all E values refer to reduction reaction
*the E values do not depend on the total number of electrons, so do not have to be scaled up or down according to the stoichiometry of the equation
*the more positive the E value for a half cell, the more readily it is reduced
In a voltaic cell the half-cell with the higher (more positive) electrode potential is the
cathode +
In a voltaic cell the half cell with the lower (more negative) electrode potential is the
anode -
electrons always flow towards the half-cell with
the highest E value
calculating the cell potential E cell
E cell = E half cell where reduction occurs - E half-cell where oxidation occurs
Note:
*The E values used in this expression must be the reduction potentials as supplied in data tables. Do not invert them before substituting into equation.
*The E values do not have to be multiplied according to the stoichiometry of the redox equation. This is because they are intensive quantities and do not depend on the total number of electrons shown in the equation
How to determine the spontaneity of a reaction
If E cell is positive, the reaction is spontaneous as written
Electrode potential and free energy change ( E cell and Delta G)
Delta G = -nFE
n = number of moles of electrons transferred in the reaction
F = the charge carried by 1 mole of electrons, known as the Faraday constant ( value of approximately 96500)
when E = 0 and delta G = 0 –> reaction is at equilibrium
trends of strength and E values
In the table of standard electrode potentials
up to down –> increasing strength as oxidizing agent
down to up –> increasing strength as reducing agent
The higher the value of E of a half cell
the stronger the oxidizing agent
The lower the value of E
the stronger the reducing agent
A little caution about interpreting E data
E data does not give any info about the rate
so a reaction that is predicted to be spontaneous may give no observable sign of reaction because the activation energy may be to high for the reaction to occur at any appreciable rate
+ E values discussed relate to only standard conditions and any changes in temp concentration or pressure will alter the values, possibly even changing whether a reaction is spontaneous or not
Electrolytic cell
*uses external source of electrical energy to bring about a redox reaction that would otherwise be non-spontaneous
*source of electric power is a battery or a DC power source
*the electrodes are immersed in the electrolyte and connected to the power supply; they must not touch each other; Electrodes are made from a conducting substance - generally metal or graphite; They are described as inert when they do not take part in the redox reactions
*Electric wires connect the electrodes to the power supply
electrolyte
where the reactant is present in the process of electrolysis
liquid, usually molten ionic compound or solution of an ionic compound
as the electric current passes through the electrolyte, redox reactions occur at the electrodes, removing the charges on the ions and forming products that are electrically neutral
the ions are therefore said to be discharged during this process
difference between electrodes and electrolytes
electrodes are electronically conducting
electrolytes are ionically conducting
electrodes in electrolytic cells
Cathode - negative; reduction; ions gain electrons
Anode - positive; oxidation; ions lose electrons
determining the products in electrolytic cells
- Identify all the ions present in the electrolyte and determine which will migrate to which electrode : anions to anode and cations to cathode
- Where there is more than one possible reaction at each electrode, determine which will occur. Write the half equation for the reaction at each electrode, showing electrons released at the anode in oxidation and taken up at the cathode in reduction.
- Balance the electrons lost and gained at the anode and the cathode, then add the two half-equations to write the equation for the net reaction
- Consider what changes would be observed in the cell as a result of the redox processes occurring. These may include colour changes in the electrolyte, precipitation of solid, gas discharge, or pH changes.
The electrolysis of molten salts
When the electrolyte is a molten salt, the only ions present are those from the compound itself as there is no solvent. So, usually, only one ion migrates to each electrode, and it is straightforward to predict the reactions that will occur
Electrolysis of Aqueous solutions
- when electrolysis is carried out on an aqueous solution, predicting the products at the electrodes is more difficult because water itself can be oxidized or reduced
- So when a solute M+A- is in an aqueous solution, there is more than one redox reaction possible at each electrode ——> specifically at the anode: either A- or H2O can be oxidized; At the cathode: either M+ or H2O can be reduced.
*the discharge of an ion at the electrode in these cases is known as selective discharge.
*The outcome is determined by the following factors:
1. relative E values of the ions
2. the relative concentrations of the ions in the electrolyte
3. the nature of the electrode
Charge =
current x time
number of faradays (F) =
moles of electrons
The amount of product formed in electrolysis is determined by
the current, duration, and the charge on the ion.
equations that link amount of product and charge
Q= It
C= A x s [current - A; Time - s]
F = ~Q/96500 = number of moles
use mole ratio in equation
m = n x M
electroplating
the process of using electrolysis to deposit a layer of a metal on top of another metal or other conductive object
an electrolytic cell used for electroplating has the following features:
- an electrolyte containing the metal ions which are to be deposited
- the cathode made of the object to be plated
- sometimes the anode is made of the same metal which is to be coated because it may be oxidized to replenish the supply of ions in the electrolyte
Voltaic cell vs electrolytic cell
types of cell - Ecell - delta G - type of reaction
voltaic - >0 - <0 - spontaneous
electrolytic - <0 - >0 - non-spontaneous
equilibirum - 0 - 0 - dead battery
