Electrochemical Cell and Its Types, Galvanic cell
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| Electrochemical Cell and Its Types, Galvanic cell |
Here you will learn about,
Electrochemistry
Electrochemical cell and its types
Gibbs Free Energy Calculation using EMF
Equilibrium
Constant Calculation Using EMF
Nernst Equation
Finding
Concentration cell potential using Nernst Equation
Definition of Electrochemistry
The field of study known as "Electrochemistry"
combines the study of ionic solutions with that of solid-state electrons. Any
material that will be used in electrochemistry requires essential measurements,
depending on the uses, to confirm its susceptibility, conductance,
responsiveness, interaction, consistency, and lifespan in a given medium.
The study of the correlation between electrical energy
and chemical changes is the focus of the branch of chemistry known as
electrochemistry. Electrochemical reactions are those in which electric
currents are either generated or input. These responses can be roughly divided
into two categories:
Electrical energy produces chemical change i.e., the
electrolysis phenomenon
Chemical energy to electrical energy conversion. i.e.,
the production of electricity using redox reactions that occur spontaneously.
An oxidation or reduction reaction at a polarized electrode surface is the subject of electrochemistry, which studies the movement of electrons in such reactions. At a particular potential, each analyte is oxidized or reduced, and the current measured is proportional to concentration. This method is an effective approach to bioanalysis.
Galvanic cell
Galvanic, also known as Voltaic, and electrolytic
cells are the two varieties of electrochemical cells. While electrolytic cells
utilize non-spontaneous reactions and therefore need an external electron
source, such as a DC battery or an AC power source, galvanic cells get their
energy from spontaneous redox reactions. Anode and cathode, which can be formed
of the same metal or two distinct metals, as well as an electrolyte, in which
the two electrodes are submerged, make up both galvanic and electrolytic cells.
DC electrical power is usually generated by galvanic cells. A straightforward galvanic cell would just have one electrolyte separated from it by a semi-permeable membrane, or a more complicated one would have two distinct half-cells joined by a salt bridge. In order to balance the developing charges at the electrodes, the salt bridge contains an inert electrolyte like potassium sulphate, whose ions will diffuse into the half-cells.
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| Galvanic cell Diagram |
The anode is where oxidation happens, and the cathode
is where reduction happens. The anode is the negative terminal for the galvanic
cell because the anode's reaction serves as the source of electrons for the
current.
Voltage is an intense attribute, meaning it is
independent of the system's size and material content. Since galvanic cells
contain a positive EMF, we want to rearrange the equation so that it will
result in a positive value when the other EMF is added.
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| Galvanic cell Example |
The two EMF readings for the zinc half-reaction are
+0.382 V and +1.221 V. We simply sum them all together to obtain a rough
estimate of 1.5 V, which represents the EMF of an alkaline AA battery.
Gibbs Free Energy Calculation using EMF
Let's say someone asks us to express the energy in
additional thermodynamic terms. Let's apply the following equation, where n
represents the number of electrons exchanged, E represents the EMF in its
standard condition, and F represents the Faraday constant, which is 96,485
C/mol.
Instead of joules, Gibbs free energy is typically
stated in kilojoules. We can determine from the sign which way the reaction
must change to achieve equilibrium. Accordingly, a system operating under
normal circumstances would have to move to the right, transforming some
reactants into products before coming to equilibrium. The magnitude shows us
how far away from equilibrium the standard state is.
Equilibrium Constant Calculation Using EMF
Assume
that in order to determine how favorable this reaction is; it is necessary to
determine the equilibrium constant K under standard conditions. The high K
value suggests that the reaction will proceed fully to completion and is
particularly beneficial to the products. For the batteries, the reaction will
proceed until Go
=0, or equilibrium, has been reached.
The
value of ΔG equals zero when the reactants and products of the electrochemical
cell are in equilibrium. The reaction quotient and the equilibrium constant
(Kc) are the same at this point. Because Δ G = -nFE, the equilibrium cell
potential is also 0.
The
following equation is generated by substituting the values of Q and E into the
Nernst equation.
0 = E0cell – (RT/nF) ln Kc
The equation is changed by converting the
natural logarithm into base-10 logarithm and replacing T=298K (standard
temperature).
E0cell = (0.0592V/n)
log Kc
The following equation created by rearranging this one.
log Kc = (nE0cell)/0.0592V
As a result, the equilibrium constant's link to the standard cell potential is found. The value of E0cell will be greater than 0 when Kc is greater than 1 (you know the value of Kc is directly related to E0 because value of Kc present in log) , indicating that the equilibrium supports the forward reaction. Similarly, E0cell will have a negative value when Kc is less than 1, indicating that the opposite reaction will likely be preferred.
Nernst Equation
“Nernst equation is an
equation relating the capacity of an atom/ion to take up one or more electrons
(reduction potential) measured at any conditions to that measured at standard
conditions (standard reduction potentials) of 298K and one molar or one atmospheric
pressure.”
Walther
Hermann Nernst, a German chemist, developed the equation. The cell potential of
an electrochemical cell at any given temperature, pressure, and reactant
concentration is frequently determined using the Nernst equation.
The
standard cell potential, temperature, reaction quotient, and the cell potential
of an electrochemical cell are all related by the Nernst equation. The Nernst
equation can be used to calculate the cell potentials of electrochemical cells
even in unusual circumstances.
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| Nernst Equation |
Ecell
= Cell Potential Of The Cell
F
= Faraday Constant
E0 = Cell Potential Under Standard Conditions
Product / Reactant =Q = Reaction Quotient
R
= Universal Gas Constant
T
= Temperature
N
= Number Of Electrons Transferred In The Redox Reaction
Finding Concentration cell potential using Nernst Equation
Consider
a concentration cell, a particular type of galvanic cell that consists of two
identical half-cells of the same material that differ only in concentration.
The sodium ion, potassium ion, or Calcium ion pumps in our cell membranes,
the ATP synthase employed in energy production, and the concentration gradients
in our nerve cells are all examples of concentration cells.
In
addition to the Henderson-Hasselbalch equation, the thermodynamics equation, is where the Nernst equation originates. When a concentration cell tries
to reach equilibrium, a little voltage is generated. The Nernst Equation can be
used to determine the potential created by a concentration cell and is as
follows:
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| Concentration cells Nernst equation |
The standard state EMF is 0 for any concentration cell because the two half-cells have identical half-reactions.
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