BASICS OF ELECTROCHEMISTRY

J.Daniell and M.Faraday – English scientists who gave basics for electrochemistry (18-19 century)

Electrochemistry is the study of interchange of chemical and electrical energy Electrochemical processes – redox processes

1. GALVANIC processes 2. ELECTROLYTIC processes COROSSIVE processes

1. GALVANIC processes

Electrochemical cell – derives electrical energy from spontaneous redox reactions taking place within the cell In an electrochemical cell – an

electric potential

between

two dissimilar metals

is created It is also called a Galvanic cell or a Voltaic cell , named after

Luigi Galvani

, or

Alessandro Volta

respectively

in Galvanic cell:

Chemical energy

is converted to

Electrical energy

Chemical

reaction is

spontaneous

and produces

electricity

METAL ELECTRODE

Rod of metal

immersed in a salt solution is called an

electrode Metal electrode potential

will be either

positive

or

negative

HALF-CELL

Metal

electrode potential

is measured in comparison with

Hydrogen

electrode potential

HYDROGEN ELECTRODE POTENTIAL

Standard hydrogen electrode

allows H 2 gas molecules to interact directly with H + dissolved in water and with the electrons from the external circuit simultaneously

Pt, H 2 /H + - such special electrode is assigned a potential of ZERO volts

METAL ELECTRODE POTENTIAL Standard metal electrode potential

is determined at solute concentrations standard temperature of 1 Molar , gas pressure which is usually 25°C of 1 atm , and a

Standard half-cell potential

a superscript



o

Me

is denoted by a degree sign as 1.Measured against 2.Concentration – –

standard hydrogen electrode 1 Molar

3.Pressure –

1 atmosphere

4.Temperature –

25 ° C

METAL ACTIVITY ROW

Li-K-Ca-Mg-Al-Zn-Cr-Fe-Cd-Pb-

H

-Cu-Hg-Ag-Pt-Au -3.0 -2.4 -0.76 -0.4

0.0V

+0.8 +1.7

GALVANIC CELLS

Galvanic cell different metals connected with a conductor – two half-cells of Such galvanic cell formulated: can be ( )Zn│ZnSO 4 ││ CuSO 4 │Cu(+) Anode ( negative electrode, active metal) is written on the left side, and Cathode ( positive electrode, less active metal) – on the right side

GALVANIC PROCESSES

It is customary to visualize the cell reaction in terms of two half reactions: an oxidation half-reaction and a reduction half-reaction

A (-)

More active

Me

0

– ne → Me

n+ (oxidation)

C (+)

Less active

Me

n+

+ ne → Me

0 (reduction)

• • By definition:

Anode

– the electrode where oxidation (loss of electrons) takes place; the anode attracts anions (Zinc electrode is the anode)

Cathode

– the electrode where reduction (gain of electrons) takes place; the cathode attracts cations (Copper electrode is the cathode)

GALVANIC PROCESSES

Anode (-) Zn 0 – 2e → Zn 2+ (oxidation) Cathode (+) Cu 2+ + 2e → Cu 0 (reduction) Total reaction: Zn 0 + CuSO 4 → Cu 0 + ZnSO 4 (redox)

ELECTROMOTIVE FORCE Cell potential

(also called – E cell electromotive force or EMF) has a contribution :

•

from the anode

which is a measure of its ability to lose electrons (called as

oxidation potential

), and

•

from the cathode

which has a contribution based on its ability to gain electrons (called as

reduction potential

) Cell potential can then expressed: E cell = reduction potential – oxidation potential

ELECTROMOTIVE FORCE

E cell =



o reduction –



o oxidation



0 reduction



0 oxidation – – cathode standard potential, V anode standard potential, V

Standard potential for the cell is equal:

to more

positive



o value minus more

negative



o value

In our case:

E cell =



o

(Cu

2+

/Cu)

–



o

(Zn

2+

/Zn)

E cell =

0.34 – (- 0.76) ═ 1.1 V

TYPES of GALVANIC CELLS

There are different number of battery cell types which include:

• • • • • •

wet cell dry cell battery battery molten-salt cell reserve cell fuel cell battery concentration cell battery battery battery Batteries can also be made up for different types of materials :

• • • • •

zinc-carbon alkaline battery nickel-zinc battery lithium-ion lead-acid battery baterry battery Wet cell

battery has a

liquid electrolyte

, while

dry cell

has an electrolyte that is

immobilized as a paste

, which is only liquid enough to allow the flow of electrons

Molten salt

batteries use

molten salt

as an electrolyte – used in electric vehicles

Reserve

battery – a battery that is often stored in an

unassembled form

and only activates when the inner parts are assembled. Can be stored for long periods of time and are often a part of emergency kits. Provide power for a few minutes

CONCENTRATION CELL

Is a kind of a wet galvanic cell that has two equivalent half-cells of the same material differing only in concentrations Rods of Cu immersed in salt solution (CuSO 4 ) of different concentration (A) Cu | CuSO 4 (0.05 M) || CuSO 4 (2.0 M) | Cu (C) One can calculate the potential developed by such a cell using the Nernst Equation

Potential of half-cell:

 (Me) =  o (Me) – (0.0592/n)·log 10 [Me n+ ]

(A) (C) Cu 0 – 2e → Cu 2+ ( 0.05 M ) Cu 2+ ( 2.0 M ) + 2e → Cu 0

DRY HERMETIC CELL

Zinc-carbon

cell battery Electrolyte – immobilized as a paste (

mixture of NH 4 Cl

and

starch or C powder

)

( –) Anode – zinc box (

cell is encased in Zinc casing)

(+) Cathode – carbon (graphite) rod, surrounded by a layer of MnO 2 (-A) Zn



NH 4 Cl



MnO 2



C (+C)

Cells are often encased in a plastic or metal casing, with two points of a negative (-) and a positive (+) sides

ALKALINE CELL

REDOX REACTIONS: (-A) Zn + 2OH − 2e → ZnO + H 2 O

(oxidation)

(+C) 2MnO 2 + H 2 O + 2e− →Mn 2 O 3 + 2OH −

(reduction)

Anode – zinc powder Cathode – MnO 2 Electrolyte – KOH

The usual voltage for an alkaline battery is 1.5 V

FUEL CELL

Fuel cell is an electrochemical energy conversion in use today use hydrogen and oxygen device as the chemicals – most fuel cells Hydrogen fuel cell converts the chemicals hydrogen water , and in the process it produces electricity and oxygen into Battery has all of its chemicals stored inside , and converts those chemicals into electricity. This means that a battery eventually “goes dead” With a fuel cell , chemicals constantly flow into the cell so it never “goes dead” As long as there is a flow of chemicals into the cell, the electricity flows out…

FUEL CELL

Hydrogen fuel cell operates similar to a battery – it has two electrodes, an anode and a cathode , separated by a membrane. Oxygen electrode and hydrogen over the other passes over one Hydrogen reacts to a catalyst on the anode that converts the hydrogen gas into negatively charged electrons (e-) and positively charged ions (H+).

Electrons flow out of the cell to be used as electrical energy Hydrogen ions move through the electrolyte membrane to the cathode where they combine with oxygen and the electrons to produce water

Reactions involved in a fuel cell are as follows:

(-) Anode side:

2H 2 - 4e => 4H +

(+)Cathode side:

O 2 + 4H + + 4e => 2H 2 O

Net "redox" reaction:

2H 2 + O 2 => 2H 2 O

Unlike batteries, fuel cells never run out

ALKALINE FUEL CELL

Alkaline fuel cells (AFCs) were one of the first fuel cell technologies developed, and they were the first type widely used in the U.S. space program to produce electrical energy and water on-board spacecrafts These fuel cells use a solution of potassium hydroxide in water as the electrolyte and can use a variety of metals as a catalyst at the anode and cathode The reactions involved in a fuel cell are as follows:

Anode Reaction (Oxidation): Cathode Reaction (Reduction):

2H 2 O 2 + 2O + 4e – 2−

Overall Cell Reaction (Redox)

: 2H 2 + O 2 → 2H 2 O + 4e − → 2O 2− → 2H 2 O

FUEL CELLS –

PROBLEMS TO BE SOLVED PRODUCTION: Hydrogen can be produced using diverse, domestic resources including fossil fuels, such as natural gas and coal (with carbon sequestration); nuclear; biomass; and other renewable energy technologies, such as wind, solar, geothermal, and hydro-electric power.

The overall challenge to hydrogen production is cost reduction STORAGE: On-board hydrogen storage for transportation applications continues to be one of the most technically challenging barriers to the widespread commercialization of hydrogen-fueled vehicles

.

Specific hydrogen storage material classes: on-board reversible metal hydrides, hydrogen adsorbents, and chemical hydrogen storage materials Regenerative Fuel Cells: produce electricity from H 2 and O 2 and generate heat and water as byproducts, just like other fuel cells. However, regenerative fuel cell systems can also use electricity from solar power or some other source to divide the excess water into oxygen and hydrogen fuel by "electrolysis"

Difference between Cell and Battery

A battery stack or is basically nothing but a pile of galvanic cells – battery consists of multiple cells Volta was the inventor of the voltaic pile , the first electrical battery Usual voltage cell is 1.5 V for a single alkaline and the voltage can be increased by adding on more cells Depending on application of the battery, the cells are combined to provide a higher voltage, for example a 9-Volt battery would have 6 alkaline cells with a 1.5 V charge

2. ELECTROLYTIC processes

Electrolytic cell is an electrochemical cell that undergoes a redox reaction when electrical energy is applied It is most often used to decompose chemical compounds, in a process called electrolysis (the Greek word

lysis

means to

break up

) Important examples of electrolysis are the decomposition of water into hydrogen and oxygen , and bauxite into aluminium and other chemicals

in Electrolytic cell:

Electrical energy

is converted to

Chemical energy

ELECTROLYSIS

Electrolytic cell has three component parts: an electrodes (a cathode and an anode ) electrolyte and two The electrolyte is usually a solution of water in which ions are dissolved Molten salts such as sodium chloride are also used as electrolytes Anode Cathode is positive electrode +( A ) – negative electrode - ( C )

Molten salt:

NaCl → Na + + Cl Electrolyte provides ions that flow to the electrodes , where charge-transferring reactions take place (-C)

Na

+

+ e → Na

o (+A)

Cl

-

e → Cl

o (reduction) (oxidation) (Cl + Cl → Cl 2 )

ELECTROLYSIS – application

1.

2.

3.

Rechargeable battery – cell battery, which acts as a galvanic cell when discharging (converting chemical energy to electrical energy), and an electrolytic cell when being charged (converting electrical energy to chemical energy Electrorefining – process of electrolytic refining of metals is used to extract impurities from crude metals. Here in this process a block of crude metal is used as anode, a diluted salt of that metal is used as electrolyte and plates of that pure metal is used as cathode Electrolysis – used to produce acids, alkalis, non-metals (H 2 , Cl 2, O 2 , As)

ELECTROLYSIS – application

4.

5.

Electroforming – reproduction (making copies) of objects by electro deposition. The surface of the wax mould which bears exact impression of the object, is coated with graphite powder in order to make it conducting. Then the mould is dipped into the electrolyte solution as cathode and metal will be deposited on the graphite coated surface of the mould Electroplating – cathode is an object on which the electroplating to be done (e.g. iron being copper-platted by using copper electroplating)

RECHARGEABLE BATTERY

A rechargeable battery (e.g. lead-acid battery) acts as:

• •

Galvanic cell – when discharging (converting chemical energy to electrical energy), and Electrolytic cell – when being charged electrical energy to chemical energy) (converting

Discharging

In the discharged state both the positive and negative plates become lead(II) sulfate (PbSO 4 ) and the electrolyte loses much of its dissolved sulfuric acid and becomes primarily water. The discharge process is driven by the conduction of electrons from the negative plate back into the cell at the positive plate in the external circuit

Negative plate

reaction (Anode Reaction): Pb(s) + HSO 4 − (aq ) → PbSO 4 (s) + H + (aq) + 2e -

Positive plate

reaction (Cathode Reaction): PbO 2 (s) + HSO 4 − (aq) + 3H + (aq) + 2e → PbSO 4 (s) + 2H 2 O(l)

Total reaction

can be written: Pb(s) + PbO 2 (s) + 2H 2 SO 4 (aq ) → 2PbSO 4 (s) + 2H 2 O(l)

Charging

In the charged state, each cell contains negative plates of elemental lead (Pb) and positive plates of lead(IV) oxide (PbO 2 ) in an electrolyte of approximately 33.5% v/v sulfuric acid (H 2 SO 4 ). The charging process is driven by the forcible removal of electrons from the positive plate and the forcible introduction of them to the negative plate by the charging source

Negative plate

reaction: PbSO 4 (s) + H + (aq) + 2e → Pb(s) + HSO 4 − (aq)

Positive plate

reaction: PbSO 4 (s) + 2H 2 O(l) → PbO 2 (s) + HSO 4 − (aq) + 3H + (aq) + 2e Overcharging with high charging voltages generates oxygen and hydrogen gas by electrolysis of water, which is lost to the cell. Periodic maintenance of lead acid batteries requires inspection of the electrolyte level and replacement of any water that has been lost

Ion motion

During discharge, H+ produced at the negative plates and from the electrolyte solution moves to the positive plates where it is consumed, while HSO 4 − is consumed at both plates. The reverse occurs during charge Due to the freezing-point depression of the electrolyte, as the battery discharges and the concentration of sulfuric acid decreases, the electrolyte is more likely to freeze during winter weather when discharged