Transcript electrochemical cell

electrochemistry
 redox
reactions ( cells )
 electrolysis
M.Morelli
I.T.I. Malignani - Udine
redox
oxidation is :
reduction is :
• addition of oxygen
• loss of hydrogen
• loss of electrons
• an increase in
oxidation state
• loss of oxygen
• addition of hydrogen
• addition of electrons
• a decrease in
oxidation state
M.Morelli
I.T.I. Malignani - Udine
a simple reaction will serve to demonstrate:
oxidation and reduction
if we immerse a piece of solid copper in a colourless silver
nitrate solution, we can tell that a chemical reaction occurs
because the solution turns blue and we can notice a silvery
coating that forms on the piece of copper
after 5 minutes
M.Morelli
I.T.I. Malignani - Udine
what is the nature of this chemical change ?
you known that copper (II ) ions in solution are blue, so copper ions (II)
must be forming and the silver nitrate solution contains silver ions ;
these ions form the solid silver-coloured material ;
the unbalanced reaction between solid copper and silver ions is this :
Ag+ (aq) + Cu (s)  Cu2+ (aq) + Ag (s)
This equation tells you that solid copper atoms are changing to copper
ions at the same time that silver ions are changing to solid silver atoms;
The reaction opposite can be split into half equations which show what
is being reduced and what is being oxidised:
Cu (s)  Cu2+ (aq) + 2 eAg+(aq) + e-  Ag (s)
loss of electrons = oxidation
gain of electrons = reduction
The global reaction is called redox
M.Morelli
I.T.I. Malignani - Udine
summary
Redox are the transfer of electrons from one
reactant to another
• when there is an oxidation, there is also a
reduction.
• the substance which loses electrons is oxidised.
• the substance which gains electrons is reduced
M.Morelli
I.T.I. Malignani - Udine
Electrochemical cells
an electrochemical cell is :
- a chemical system in which an oxidation-reduction reaction
occurs spontaneously;
-the transfer of electrons is used to produce an electric current;
the cell voltage is a measure of the tendency of electrons to flow;
if two substances have a great difference in their tendency of lose
and acquire electrons, the voltage of the cell will be high ; cell
voltage can be used to predict whether a particular redox reaction
will occur spontaneously ;
-Electrochemical cells that use an oxidation-reduction
reaction to generate an electric current are known as galvanic
or voltaic cells.
M.Morelli
I.T.I. Malignani - Udine
an example of voltaic cell is shown below:
H+ ions flow toward the cathode, where they are reduced to H2 gas ( cathode ).
On the other side of the cell, the zinc metal is oxidized becoming Zn2+ (anode ).
The electrode at
which reduction
occurs is called
cathode.
The electrode at
which oxidation
occurs is called
anode.
M.Morelli
I.T.I. Malignani - Udine
To provide a basis for comparing the results of one experiment
with another, the following set of standard-state conditions for
electrochemical measurements has been defined:
- solutions are 1 M.
- gases have a partial pressure of 0.1 MPa ( 1atm).
- standard-state measurements are taken at 25oC.
Cell potentials measured under standard-state conditions are
represented by the symbol Eo.
The standard-state cell potential Eo, measures the strength of
the driving force behind the chemical reaction.
The larger is the difference between the oxidizing and reducing
strengths of the reactants and products, the larger is the cell
potential.
M.Morelli
I.T.I. Malignani - Udine
Example: the experimental value for the standard-state cell potential for the
reaction between zinc metal and acid is 0.76 volts.
Zn(s) + 2 H+(aq)
Zn2+(aq) + H2(g)
Eo = 0.76 V
the cell potential for this reaction measures the relative reducing power of zinc
metal compared with hydrogen gas. But it doesn't tell us anything about the
absolute value of the reducing power for either zinc metal or H2.
We therefore arbitrarily define the standard-state potential for the reduction of
H+ ions to H2 gas as exactly zero volts.
2 H+ + 2 e-
Eo = 0.00 V
H2
this reference point is used to calibrate the potential of any other halfreaction
M.Morelli
I.T.I. Malignani - Udine
the overall cell potential for a reaction must be the difference of
the potentials for the reduction and oxidation half-reactions
Eocell = Eored – Eoox
If the overall potential for the reaction between zinc and acid is
0.76 V, and the half-cell potential for the reduction of H+ ions is 0
volts, then the half-cell potential for the oxidation of zinc metal
must be 0.76 V .
Zn
Zn2+ + 2 e-
2 H+ + 2 e-
H2
Eoox = 0.76 V
Eored = 0.00 V
¯¯¯¯¯¯¯¯¯¯¯¯¯¯¯¯¯¯¯¯¯¯¯¯¯¯¯¯¯¯¯¯¯¯¯¯¯¯¯¯¯
Zn + 2 H+
Zn2+ + H2
M.Morelli
Eocell = Eored - Eoox = 0.76 V
I.T.I. Malignani - Udine
a Daniell cell is an example of a voltaic cell used as a battery.
the model of Daniell cell is :
the reaction that occurs in this cell is:
Zn(s) + Cu2+(aq)
Zn2+(aq) + Cu(s)
Electrons flow from the anode to the
cathode in a voltaic cell.
They flow from the electrode at which
they are given off to the electrode at
which they are consumed.
the line notation for a standard-state Daniell cell is written as follows:
anode (-) Zn | Zn2+(1.0 M) || Cu2+(1.0 M) | Cu(+) catode
Reading from left to right, this line notation therefore corresponds to the direction in
which electrons flow.
M.Morelli
I.T.I. Malignani - Udine
classwork
can you describe the problem in your own words ?
•draw a concept map for the term electrochemical cell
include the following terms : anode, cathode, electrode,
oxidation ,reduction, salt bridge, voltage and cell
•using the table of standard-state potentials answer to these
questions :
1- can aluminium foil immersed in an electrolyte solution be
used to restore the lustre to silverware? justify your answer
by using the cell potential ( assume 1 M solutions and 25°C)
2- can a zinc strip immersed in an electrolyte solution be
used to restore the lustre to silverware?
3- calculate the Daniell’s cell potential
M.Morelli
I.T.I. Malignani - Udine
laboratory experience
1 - build three different electrochemical cells
2 - complete a detailed diagram of each
3 - determine the half reactions and complete reactions occurring in each cell
4 - calculate the cell potential for each cell.
Material
50 mL beaker (3); zinc strip,copper
strip, lead strip
wire and
voltmeter
clip
(2),
salt
bridge,
1.0 M copper (II) nitrate, 1.0 M zinc
nitrate, 1.0 M lead nitrate
Fill the lines with the correct terms
for every cell
M.Morelli
I.T.I. Malignani - Udine
Electrolytic Cells
Voltaic cells use a spontaneous chemical reaction to
drive an electric current through an external circuit.
These cells are important because they are the basis for
the batteries that fuel modern society. But they aren't the
only kind of electrochemical cell.
It is also possible to construct a cell that does work on a
chemical system by driving an electric current through
the system. These cells are called electrolytic cells.
Electrolysis is used to drive an oxidation-reduction
reaction in a direction in which it does not occur
spontaneously.
M.Morelli
I.T.I. Malignani - Udine
The Electrolysis of Molten NaCl
An idealized cell for the electrolysis of sodium chloride is
shown in the figure below. A source of direct current is
connected to a pair of inert electrodes immersed in
molten sodium chloride.
The net effect of passing an
electric current through the
molten salt in this cell is to
decompose sodium chloride
into its elements, sodium metal
and chlorine gas.
Reactions of NaCl (s):
Cathode (-):
Na+ + e-  Na
Anode (+):
2 Cl- Cl2 + 2 e-
M.Morelli
I.T.I. Malignani - Udine
Difference between
electrolytic cells:
voltaic
cells
and
Voltaic cells use the energy given off in a
spontaneous reaction to do electrical work.
Electrolytic cells use electrical work as source of
energy to drive the reaction in the opposite
direction.
M.Morelli
I.T.I. Malignani - Udine
The commercial Downs cell used to electrolyse
sodium chloride
Chlorine gas that forms on the
graphite anode inserted into the
bottom of this cell bubbles
through the molten sodium
chloride into a funnel at the top of
the cell. Sodium metal that forms
at the cathode floats up through
the molten sodium chloride into a
sodium-collecting ring, from which
it is periodically drained.
The diaphragm that separates the two electrodes is a screen of iron gauze, which
prevents the explosive reaction that would occur if the products of the electrolysis
reaction came in contact.
M.Morelli
I.T.I. Malignani - Udine
The Electrolysis of Aqueous NaCl
The figure below shows a cell in which an aqueous solution of
sodium chloride is electrolysed.
Electrolysis of aqueous
NaCl solutions gives a
mixture of hydrogen and
chlorine gas and an
aqueous
sodium
hydroxide solution.
2 NaCl(aq) + 2 H2O(l) 2 Na+(aq) + 2 OH-(aq) + H2(g) + Cl2(g)
Electrolysis of aqueous sodium chloride is a more important process commercially,
because the demand for chlorine is much larger than the demand for sodium.
Electrolysis of an aqueous NaCl solution has two other advantages. It produces H2 gas
at the cathode, which can be collected and sold. It also produces NaOH, which can be
drained from the bottom of the electrolytic cell and sold.
M.Morelli
I.T.I. Malignani - Udine
Electrolysis of Water
Another example of electrolysis
A pair of inert electrodes are sealed in opposite ends of a container
designed to collect the H2 and O2 gas given off in this reaction.
The electrodes are then connected to a battery or another source of
electric current.
By itself, water is a very poor conductor of electricity. We therefore
add an electrolyte to water to provide ions that can flow through the
solution, thereby completing the electric circuit. The electrolyte must
be soluble in water. It should also be relatively inexpensive. Most
importantly, it must contain ions that are harder to oxidize or reduce
than water.
Reaction :
2 H2O(l)  2 H2(g) + O2(g)
M.Morelli
I.T.I. Malignani - Udine
Corrosion :
corrosion occurs by oxidation-reduction reactions.
When salt is spread on the roads to melt ice, the salt
accelerates the formation of rust on car bodies,
bridges, and other steel structures; near the oceans,
the salt suspended in the air assists in corroding
metal objects;
in each of these cases the oxidation of the metal to
its oxide is central to the corrosion process.
M.Morelli
I.T.I. Malignani - Udine
Electroplating:
electrochemical process for depositing a thin layer of
metal on a metallic base. To electroplate an object with
a metal “X “, use:
- the object as cathode
- a strip of X as anode
- a solution of a compound of X as electrolyte
Objects are electroplated to prevent corrosion, to
obtain a hard surface or attractive finish, to purify
metals (as in the electro refining of copper),
Typical products of electroplating are silver-plated
tableware, chromium-plated car accessories.
M.Morelli
I.T.I. Malignani - Udine