The atoms were believed to be non -divisible until the discovery of subatomic particles. J. J. Thomson proposed that the atom is a positively charged sphere in which the electrons are embedded.

**SUMMARY**

The atoms were believed to be non -divisible until the
discovery of subatomic particles. J. J. Thomson proposed that the atom is a
positively charged sphere in which the electrons are embedded. However, it
fails to explain the stability of atoms. Rutherford, based on his α-rays
scattering experiment, introduced the term nucleus which is a positively
charged one and the negatively charged electrons are revolving around it, at
high speeds. Bohr modified the Rutherford theory and introduced stationary
orbits by taking into account the quantisation of energy. Louis de Broglie proposed
that all matter possess dual nature. i.e. they behave both as a wave and a
particle. De Broglie wavelength λ = h / mv = h/√(2mev) significant for a
microscopic particle such as an electron. The wave nature of electron is also
proved by Davisson and Germer through electron diffraction. For a microscopic
particle such as an electron, the simultaneous measurement of the conjugate
variables such as position and momentum involves uncertainty which is known as
Heisenberg's uncertainty principle and it is expressed as Δx.Δp ≥ h/4**π.**

De Broglie's concept and Heisenberg's uncertainty
principle lead to the development of quantum mechanical model of atom. Erwin
Schrodinger, developed an equation for an electron wave which is expressed as
Hψ = Eψ. This second order differential equation is exactly solvable for one
electron system such as H, He^{+} etc... but it is quite complex for
multi-electron systems. The SchrÖdinger wave equation is solvable for certain
energy values called eigen values. The wave functions corresponding to these
eigen values are called atomic orbitals. The wave function ψ itself has no
physical meaning. However, |ψ|^{2} is related to the probability of
finding electron around the nucleus. Thus the quantum mechanical model
introduced us the term orbital which is the three dimensional space in which
the probability of finding the electron is maximum. The electron in an orbital
can be described by a set of four quantum numbers namely, principal quantum
number (n) representing the principal energy level, azimuthal quantum number (*l*) representing the sub shell (orbital),
magnetic quantum number (m) representing the different orientation of orbitals
in space and spin quantum number (s) representing the spinning of electron
about its own axis.

The general solution of Schrodinger for a one electron
system can be expressed in spherical polar coordinates (r, θ, φ)

Ψ (r, θ, φ) = R(r).f(θ).g(φ)

(where R(r) is called radial wave function, other two
functions are called angular wave functions).

The plot of 4πr^{2}ψ^{2} versus r gives
the radial distribution curves. The number of radial nodes is given by(n-*l*-1) whereas the angular nodes is equal
to *l*. The angular distribution curve
gives* *the boundary space diagram of
orbital. s orbital is spherical in nature. The shape of p orbital is spherical
and the d orbital has clover leaf shape.

Electrons are filled in various orbitals in the increasing
order of their energies which is known as Aufbau principle. The relative
energies of various orbitals are given by (n+*l*) rule which states that, the lower the value of (n + *l*) for an orbital, the lower is its
energy. If two orbitals have the same value of (n + *l*), the orbital with lower value of n will have the lower energy.
As per Pauli's exclusion principle, the maximum number of electrons that can be
accommodated in an orbital is two. In the case of degenerate orbitals electron
pairing does not take place until all the available degenerate orbitals contain
one electron each. This is known as Hund's rule. Based on these principles, electronic
configurations of atoms can be written. In degenerate orbitals, the completely
filled and half filled configurations are more stable than the partially filled
configurations. This is due to the symmetry and exchange energies.

Tags : Chemistry Chemistry

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