This document provides a summary of a presentation on conductometry. It discusses electrochemical cells, types of electrodes including reference and indicator electrodes. It also describes the Nernst equation and its applications in determining solubility products and for analytical chemistry purposes such as measuring ion concentrations using cell potentials. Electrode types including electrodes of the first, second and third kind are explained along with examples like the silver/silver chloride electrode.

2. CONDUCTOMETRY
LAB
PRESENTATION
 Electrochemical cells
 Types of electrodes
 Nernst Equation & its
Application

3. Prepared & Presented by:
Sidra Safdar
Durrani
M.Sc. Final year
Presented to:
Miss Atiya Firdous
For the Course of:
Conductometry
lab

4. ELECTROCHEMICAL CELLS
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5. ELECTROCHEMICAL CELLS
Electrochemical cells are Batteries
Alkaline Batteries
Car Batteries
Cell Phone batteries

6. ELECTROCHEMICAL CELLS
Electrochemical cells are Batteries
Lithium Coin Cell
Space Ship Batteries
Ni-Metal Hydride

7. ELECTROCHEMICAL CELLS
An electrochemical cell – a system
of
salt
electrodes ,
bridge
that
electrolytes ,
allow
and
oxidation
and reduction reactions to occur
and electrons to flow through an
external circuit .

8. FOR ANY CELL
 Oxidation
a l w a ys
occurs
at
the
anode and reduction at the cathode.
 Electrons flow through the wire and
go from anode to cathode.
 Anions (- ions) migrate to the anode
and cations (+ions) migrate towards
the cathode usually through the salt
bridge.

9. ANALYZING THE
ELECTROCHEMICAL CELLS
 The reaction that is higher on the reduction
chart is the reduction and the lower is
oxidation and is written in reverse.
 The salt bridge allows ions to migrate from
one half-cell to the other without allowing
the solutions to mix.

10. A Cu/Sn ELECTROCEMICAL
CELL

11. A
HYDROGEN HALF-CELL

12. A H 2/ A g E L E C T R O C H E M I C A L C E L L
WITH A KNO3 SALT BRIDGE

13. The term ―electrochemical cell‖ is
often used to refer to a:
Voltaic Cell – one with a spontaneous reaction
SOA over SRA on the activity series
Eocell greater than zero = spontaneous
Electrolytic cell – one with a nonspontaneous
reaction
SOA below SRA – i.e. zinc sulfate and lead
solid cell
Eocell less than zero= nonspontaneous

14. VOLTAIC CELLS (AKA
GALVANIC CELL)
A device that spontaneously produces electricity by
redox
Composed of two half-cells; which each consist of a
metal rod or strip immersed in a solution of its own
ions or an inert electrolyte.
The electrons flow from the anode to the cathode (―a
before c‖) through an electrical circuit rather than
passing directly from one substance to another
A porous boundary to maintain cell neutrality
(mostly is a salt bridge containing an inert aqueous
electrolyte ; such as Na2SO4(aq) or KNO3(aq)).

15. ELECTROLYTIC CELLS
Electrolytic Cell – a cell in which a nonspontaneous
redox reaction is forced to occur; a combination of
two electrodes, an electrolyte and an external power
source.
Electrolysis – the process of supplying electrical
energy to force a nonspontaneous redox reaction to
occur
The external power source acts as an “electron
pump”; the electric energy is used to do work on
the electrons to cause an electron transfer .
Electrons are pulled from the
anode and pushed to the
cathode by the battery or
power supply

16. TYPES OF ELECTRODES
TYPES OF ELECTRODES ON THE BASIS OF USES:-
Electrodes are classified into two types on the
basis of their uses:•Reference Electrode
•Indicator Electrode
TYPES OF ELECTRODES ON THE BASIS OF THEIR
WORKING:-
Whether the electrode is indicator or reference, it is
classified into the following types:
•Electrode of First kind
•Electrode of Second kind
•Electrode of Third kind

17. REFERENCE ELECTRODE:A reference electrode is an electrode which has a stable
and well-known electrode potential. A half-cell with an
accurately known electrode potential, Eref, that is
independent of the concentration of the analyte or any
other ions in the solution and always treated as the lefthand electrode. Common examples are:
Normal hydrogen electrode (NHE
Saturated calomel electrode (SCE)
Copper-copper (II) sulfate electrode (CSE)
Silver chloride electrode
pH-electrode

18. INDICATOR ELECTRODE:-
It can serve either as an anode or a cathode,
depending on the applied polarity. One of the
electrodes in some ""classical two-electrode""
cells can also be considered a ""working""
(""measuring,"" ""indicator,"" or ""sensing"")
electrode). Indicator electrode, responds
rapidly and reproducibly to changes in the
concentration of an analyte ion (or groups of
analyte ions).

19. ELECTRODE OF FIRST KIND:This type of electrode always contains a
single chemical element in contact with an
electrolyte solution containing its own ions,
or charged species originating from this
element (e.g., silver immersed in a silver
nitrate solution,or hydrogen gas and
oxonium ions, etc.). The equilibrium
electrode potential of this electrode is a
function of the activity of the ions in the
solution and the activity of the chemical
element.
Subtypes:
metal/metal
ion
electrode, gas electrode, amalgam electrode.

20. Silver chloride electrode:Silver/ chloride electrode is a type of reference
electrode, commonly used in electrochemical
measurements. For example, it is usually the
internal reference electrode in pH meters. As
another example, the silver chloride electrode is
the most commonly used reference electrode for
testing cathodic protection corrosion control
systems in sea water environments. The electrode
functions as a redox electrode and the reaction is
between the silver metal (Ag) and its salt — silver
chloride (AgCl, also called silver (I) chloride).
.

21. ELECTRODE OF SECOND KIND:A metal electrode assembly with the equilibrium
potential being a function of the concentration of an
anion in the solution.Typical examples are the
silver/silver-chloride electrode and the calomel
electrode. Contrast with electrode of the first kind
and electrode of the third kind, the assembly consists
of a metal, in contact with a slightly soluble salt of
this metal (or metal - oxide), immersed in a solution
containing the same anion as that of the metal salt
(e.g.,
silver/silver
chloride/potassium
chloride
solution). The potential of the metal is controlled by
the concentration of its cation in the solution, but
this, in turn, is controlled by the anion concentration
in the solution through the solubility product of the
slightly soluble metal salt.

22. ELECTRODE OF THIRD KIND:A metal electrode assembly with the equilibrium
potential being a function of the concentration of a
cation, other than the cation of the electrode metal, in
the solution. These have been used, with limited
success, in sensors for metal ions for metals that are
not stable in aqueous solutions, e.g., calcium and
magnesium.
Contrast with electrode of the first kind and electrode of
the second kind, the assembly consists of a metal in
contact with two slightly soluble salts (one containing
the cation of the solid metal, the other the cation to be
determined, with both salts having a common anion)
immersed in a solution containing a salt of the second
metal,
(e.g.
zinc
metal/zinc
oxalate/calcium
oxalate/calcium salt solution).

23. NERNST EQUATION
In electrochemistry, the Nernst equation is an equation
that relates the equilibrium reduction potential of a halfcell in an electrochemical cell (or the total voltage
(electromotive force) for a full cell) to the standard
electrode potential, temperature, activity, and reaction
quotient of the underlying reactions and species used.
It is named after the German physical chemist who first
formulated it, Walther Nernst.

24. APPLICATIONS OF EQUATION
1. Oxygen and the aquatic environment
The presence of oxygen in the atmosphere has a
profound effect on the redox properties of the aquatic
environment— that is, on natural waters exposed
directly or indirectly to the atmosphere, and by
extension, on organisms that live in an aerobic
environment. This is due, of course, to its being an
exceptionally strong oxidizing agent and thus a lowlying sink for electrons from most of the elements and
all organic compounds. Those parts of the environment
that are protected from atmospheric oxygen are equally
important because it is only here that electrons are
sufficiently available to produce the "reducing"
conditions that are essential for processes varying
from photosynthesis to nitrogen fixation.

25. APPLICATIONS OF EQUATION
2. Analytical chemistry application
A very large part of Chemistry is concerned, either
directly or indirectly, with determining the
concentrations of ions in solution. Any method that
can accomplish such measurements using relatively
simple physical techniques is bound to be widely
exploited. Cell potentials are fairly easy to measure,
and although the Nernst equation relates them to
ionic activities rather than to concentrations, the
difference between them becomes negligible in
solutions where the total ionic concentration is less
than about
10–3 M.

26. APPLICATIONS OF EQUATION
3. Determination of solubility products
The concentrations of ions in equilibrium with a
sparingly soluble salt are sufficiently low that
their direct determination can be quite difficult.
A far simpler and common procedure is to set
up a cell in which one of the electrode reactions
involves the insoluble salt, and whose net cell
reaction corresponds to the dissolution of the
Salt. Nernst equation can be modified in the
given form in order to determine Ksp
Ecell = E - (0.0591 / n) log Ksp

28. THANKS
EVERYON
E