Download Structure of Atom

STRUCTURE OF ATOM
Atom :
It is one of the large no. of particle that make up a given element together. An atom is the smallest particle
of chemical element that can take part in a chemical reaction.
Atom are made up of protones, electrones & neutrons .
The atomic radius is of the order of 10–10 metre but almost all its mass is concentrated in small region
called nucleus.
Size of nucleus = 10–15 m.
Sub-Atomic particles :
Dalton’s concept of the indivisibility of atom was disproved by experimental evidence obtained by
scientists.
It was found that atoms, although tiny are composed of particles called electron, proton & neutron.
Cathode Rays : Discovery of electrons :The nature and existance of electron was established by experiment
on conduction of electricity through gases.
Negatively charge rays emitted from cathode in discharge tube are known as cathode rays
Properties of cathode rays :
  They travel in straight line.
  They carry negative charge.
  They are made up of material particles.
  They produce heating effect
  They cause ionization of the gas through which they pass.
  They produce X–ray when they stripe against the surface of hard metals like w, mo etc.
  They produce green fluorescence on the glass walls of the discharge tube as well as on certain other
phosphorescent or fluorescent substance such as zinc sulphide.
  They affect the photo graphic plates.
  They possess penetrating effect.
Electron :
It is negatively charged particles with mass 9.1 × 10–31 Kg and charge equal to 1.6022 × 10–19C. It was
discovered by J.J. Thomson. Its relative charge is –1 & mass 0.000544 Approximately taken as O. They
are denoted by -1e0
Anode rays :
On passing high voltage between the electrodes, It was found that some rays were coming from the side of
the anode which passed through the holes in the calhode & produced green fluorescence on the opposite
glass wall coated with Zinc sulphide screen. These ray were called Anode Rays,Canal Rays.
Properties of anode rays :
(i) They travel in straight time.
(ii) They are made up of meterial particle.
(iii) They are positively charged
(iv) Unlike cathode rays, the ratio charges/mass is found to be different for the particles constituting
anode rays when different gases are taken inside the discharge tube.
(v) The value of charge e on the particle constituting the anode rays in also found to depend upon the
nature of the gas taken in the discharge tube
(vi) The mass of the particles constituting the anode rays is also found to be different for different gases
taken in the dischange tube. However its value is found to be nearly equal to equal to that of the
atom of the gas.
Protons :
They are positively charged particles with mass 1.67262 × 10–27 kg & charge equal to 1.6022 × 10–19C.
Proton was discovered by Goldstein (1866). They are present in highly dense central part of atom called
nucleus of atom. Their mass is 1.007274. They are .
Neutron :
It is neutral particle with mass equal to 1.67493 × 10–27 kg Neutron was discovered by Chadwick
The neutron is also present in nucleus of an atom. Except hydrogen,the atoms of all other elements
including isotopes of hydorgen contain all the three particles neutron, proton & electron.
mass of neutron is 1.0078764.
Denoted by 10 n
THOMSON MODEL OF ATOM
 J.J. Thomson proposed a model in which atom was assumed to be uniform sphere
with radius 10–8 cm of positive electricity (positive charge) with electrons embedded into
it in such a way as to give the most stable electrostatic arrangement.
 In this model, the atom is visualized as a pudding or cake of +ve charge with raisin
embedded into it.
Drawback of Thomson Model of atom :
This model of atom could account for the electrical neutrality of atom, but it could not explain the results
of gold foil experiments carried out by Rutherford.
RUTHERFORD’S -SCATTERING EXPERIMENTS :
Rutherford bombarded high energy -particles from radioactive source on thin foil(10-4cm) of gold metal.
The thin gold foil had a circular fluorescent ZnS screen around it, whenever an -particle struck the
screen a tiny flash of light was produced at that point. He observed that
 Most of  - particles passed(nearly99%) through the gold foil undeflected.
 A small fraction of  - particles was deflected by small angles.
 A very few  - particles (1 in 20000) bounced back.
Conclusion of  - scattering Experiment :
  An atom consists of tiny positively charged centre called nucleus.
  Nucleus consists of proton & neutrons called nucleons.
  Nucleons are much heavier than the electron indicating that mass of atom lies in its nucleus. The
total no. of nucleons is termed as mass number.
  The electrons are outside the nucleus revolving around nucleus at high speed.
  The number of electrons in an atom is equal to no. of protons in it, making it Electrically neutral.
  Electrons & nucleus are held together by electrostatic force of attraction.
Draw backs of Rutherford's Model of Atom :
 When a body is moving in an orbit, it undergoes acceleration so an electron moving around nucleus in an
orbit in under acceleration.
According to Maxwell’s electromagnetic theory charged particles when accelerated must emit
electromagnetic radiations. Therefore an electron in an orbit will emit radiations; the energy carried by
radiations comes from electronic motion. Its path will become closer to nucleus and ultimately should
spiral into nucleus within 10-8s. But actually this doesn’t happen.
Thus Rutherford’s model cannot explain the stability of atom if the motion of e– is described of the basis
of classical mechanics & electromagnetic theory.
 It doesn’t give any idea about distribution of electron around the nucleus and about their energies.
 It was not able to explain Hydrogen spectrum.
Developments Leading to the Bohr’s Model of Atom :
(i) Dual nature of radiations:
Light radiations are regarded as having a dual nature i.e. wave nature as well as particle nature.
(ii) Experimental results regarding spectra can be explained only by assuming Quantized electronic energy
level in atoms.
Wave Nature :
A wave is spread out in space. Two or more waves can exist in same region or space. When two waves
are present together, their resultant wave can be larger or smaller than individual waves In simple words,
wave is delocalized and two waves may interfere. Their interference may be constructive or destructive.
Particle Nature :
A particle occupies a well defined position in space which cannot be simultaneously occupied by another
particle. i.e. a particle is localized in space. If there is more than one particle in a given region or space,
their sum is equal to number of individual particles i.e. two particles do not interfere.
Electromagnetic waves :
Those waves which consist of oscillating electric & magnetic field are called electromagnetic waves.
These radiation which are associated with electric & magnetic field are called electro magnetic radiations.
All of them move with same speed regardless of their wavelengths. They travel at the speed of 3×10 8
ms–1, which is called speed of light. It is denoted by C. These do not require any medium for propagation.
Frequency :
It is defined as number of waves passing through a point in one second. It is denoted by  It’s unit is
s–1 or Hertz. It is related to velocity & wavelength by formula
c
υ=
λ
c is the velocity of light (2.9979 × 108 ms–1) taken of 3×108 ms–1  is wave length.
Wavelength :
It is the distance between centre of two adjacent crests or troughs. It is denoted by  Its unit is metre or
nm (nanometer)
Wave number :
It is the reciprocal of wave length . It is defined as no. of waves in 1cm or 1m length i.e. per unit length. It
is denoted by υ
1
υ=
-1
-1
λ
Unit cm or m
Amplitude :
The height of crest or depth of trough is called amplitude of wave.
Electromagnetic Spectrum :
When these electromagnetic radiations are arranged in order of their increasing wavelengths or decreasing
frequencies, the complete spectrum obtained is called electromagnetic spectrum.
Plank's Quantum Theory :
 The energy is radiated or a absorbed by a body not continuously but discontinuously in form of small
packets.
 Each packet is called quantum. In case of light the quantum is called 'Photon'.
 The energy of quantum is directly proportional to the frequency () of the radiation
E υ
E= nhυ (for n photons)
E = hυ (for one photon)
E=
hc
λ
h = 6.626 × 10 –34 Js
According to planck's theory, energy is always emitted in intergal multiple of .hυ
Limitation of Plank's Quantum Theory :
Planck could not explain why energies should be quantized in this manner.
He could not explain the distribution of intensity in radiation from a black body as a function of frequency
at different with this assumption that energy are quantized.
Photo electric Effect :
It is the phenomenon is which the surface of alkali metals like potassium & calsium emit electrons, when
a beam of light with high frequency is made to fall on them. The ejected electrons are called photo
electrons.
According to the wave theory, both number of electrons ejected and their energies should depend upon the
intensity of light. In practice, it is found that while the no.of electrons ejected depends upon intensity of
incident light but their energy do not.
Black body radiations :
An ideal body which which can emit & absorb radiations of all frequencies is called black body, The
radiation emitted is called black body radiation.
Particle Nature of Light :
The observation in photo electric effect that the number of photoelectrons ejected is directly proportional
to intensity of incident light showed that light consist of particles called photons moving with speed light.
They strike the electrons & photo electrons are ejected.
E = h
using planck’s Quantum theory
Energy of photon = work function + kinetic energy of ejected electrons
hυ = hυ0 +
1 2
mv  hυ = w 0 + K.E
2
Where W0 is work function, 0 is the threshold frequency.
hc hc 1 2
=
+ mv
λ
λ0
2
Where λ 0 = Thershold wave length
n = Frequency
 K.E. = hυ - hυo = h(υ - υo )
As K.E. =
1 2
mv
2

me = mass of electron
Work Function :
It is the minimum energy required to eject an electron from a metal. It is equal to h0. It is also equal to
Ionization Energy.
Spectrum :
It is a combination of light of different wavelength or frequency is called spectrum if it is in continuous
manner is called continuous spectrum.
Cause of spectrum :
When electromagnetic radiations interact with matter energy is exchanged and atoms & molecules may
absorb this energy and electron gets excited to higher energy levels. when they come back to lower energy
level, they emit radiation belonging to different parts electromagnetic spectrum.
Types of spectrum :
There are two types of spectrum
(i) Emission spectrum.
(ii) Absorption spectrum.
(i) Emission spectrum :
The spectrum of radiations emitted by a substance that has absorbed energy by heating it or by
irradiating it. It involves bright lines separated by dark bonds.
(ii) Absorption spectrum :
It is like photographic negative of emission spectra. radiation is passed through a sample which
absorbs radiation of certain wavelengths. The missing wave length which correspond to the
radiation absorbed by matter leave dark spaces in bright continuous spectrum. It is called absorption
spectrum.
Atomic Spectra :
The emission spectrum of elements consisting of discrete sharp lines having different wavelength
separated by dark bands is called atomic spectra. Such spectra is also called line spectra.
The characteristic lines in atomic spectra can be used in chemical analysis to identify unknown atoms as
much as finger prints are used to identify people.
Hydrogen spectra :
If a discharge is passed through hydrogen gas at a low pressure, some hydrogen atoms are formed. which
emit light in the visible region. By using spectroscope it is found to comprise of series of lines of different
wavelength. The four line can be seen by eyes, but many more obseved photographically in ultraviolet
region.
Lyman Series :
When excited electrons in hydrogen atoms fall from higher energy levels to first energy level(n=1), the
series of lines observed are called lyman series. They are observed in ultraviolet region.
1 1 
υ  R  2  2  , n  2,3, 4,5
1 n 
Balmer Series :
When electrons fall form higher Energy level to second energy level(n=2) the series of lines observed is
called balmer series.
1 1
They are observed in visible region. υ  R  2  2  , n  3, 4,5.........
2
n 
Paschen Series :
In this series excited electrons of hydrogen atoms fall from higher energy levels to third energy
level(n=3). They are observed in infra-red region.
Paschen.
1 1
υ  R  2  2  , n  4,5,6.........
3 n 
Brackett Series :
In this seried excited electrons of Hydrogen atoms fall from higher energy levels to fourth energy
levels(n=4). They are also observed in infra-red region.
Brackett
1 1
υ  R  2  2  , n  5,6,7.........
4 n 
P fund Series :
In this series excited electrons of hydrogen atoms fall from higher energy levels to fifth energy levels.
They are also found is infra red region.
Ryd berg formula :
υ
1 1 2
1
 RH  2  2  Z

 n1 n 2 
υ = wave number
RH = Rydberg constant = 109677cm–1
n1 = Lower energy level
n2 = Higher Energy level
BOHR’S THEORY :
 An atom consist of a small heavy positively charged nucleus in the centre surrounded by electrons. The
electrons in an atom revolve around the nucleus only in certain selected circular paths which have fixed
value of radius and energy.These circular paths are known as orbits. These orbits are associated with
definite energies and are also called energy shells or energy levels. Represented as K,L,M,N…… and
numbered as n=1,2,3,4,5…….
 The energy of an electron in the orbit does not change with time .So the energy of the an electron in a
particular energy shell remains constant. So these orbits are also called stationary states or allowed
energy states.

When an one electron moved from one orbit to another, it either radiates or absorbs energy. It is moved
towards nucleus, energy is radiated.
E  E 2  E1  hυ 
hc
λ
1312
kJ mol1 for H  atom
2
n
2 2 m Z2 e2
En 
for other atoms
n2 h2
n 2h 2
1
rn  2
 rn  n 2 or rn 
2
4  me Z
Z
En 
Where, z is atomic no.
rn is radius of nth orbit.
Radius of various nuclei = 1.4 × 10–13 cm × A1/3
Where A  mass no. of given nucleus.
 The angular momentum of orbits is quantized in which of
mvr = nh/2π
where n = 1, 2, 3..........
.
 Energy of orbit closest to the nucleus is lowest where as energy of orbit far away from the nucleus is
highest.
 As we go higher, difference in energy of energy levels goes on decreasing.
Achievements of Bohr’s Theory :
 It explains the stability of the atoms.
 It explains the line spectrum of hydrogen.
 It could help to calculate energy & velocity of electrons in various energy levels
 It helped to calculate radius of various energy levels in hydrogen & hydrogen like species.
Drawback of Bohr’s Theory :
 It could not explain spectum of multi electron atoms.
 It could not explain zeeman effect (Splitting of spectral lines in strong magnetic field).
 It could not explain stark effect (Splitting of spectral lines in electric field).
 It could not explain shape of molecules.
 It was not in accordance with Heisenberg’s uncertainty principle.
Quantum Mechanical Model :
A mathematical theory that developed from quantum theory and is used to explain the behaviour of atoms,
molecules & elementary particles.
It is developed on the basis of Heisenberg’s uncertainty. Principle & dual behavior of matter.
Dual Behavior of Matter :
Light radiations are associated with dual behavior, matter should also be associated with dual behavior. It
means matter is associated with both wave as well as particles.
The diffraction pattern obtained when beam of electrons is made to fall on Ni or Al crystal is similar to
that of x-rays which is associated with wave like behavior. Particle nature of electron is proved with the
help of photo electric effect.
de Broglie Equation :
Every particle can be considered to be associated with a wavelength & wave properties. The wavelength
of moving particles or object can be calculated with the help of de Broglie equation.
de- Broglie relation states that the wavelength associated with moving object or an electron is inversely
proportional to the momentum of a particle.
h
h
λ


where P is momentum of particle = mv
mv P
Heisen berg’s uncertainty principle :
It is not possible to determine the exact position & velocity simultaneously for a sub-atomic particle like
electron at any given instant to an absolute degree of accuracy.
If x is uncertainty in position & V is uncertainty in velocity , then.
x.v 
h
4πm
Important features of Quantum Mechanical Model of atom:
 The energy of electrons in atom is quantized ie can have certain values.
 The existence of quantized electronic energy levels is a direct result of the wave like properties of
electrons.
 Both the exact position & exact velocity of an electron an atom cannot be determined simultaneously.
 An atomic orbital have wave function . There are many oribital in an atom. An orbital cannot have more
than two electron.
 The orbitals are filled in increasing order of energy. All the information about the electron in an atom
stored in orbital wave function .
 The probability of finding electron at a point with in an atom is proportional to square of orbital wave
function i.e., |2| at that point. It is known as probility density & is always positive from the value of 2 at
different points with in atom, it is possible to predict the region around the nucleus where electron most
probably will be found.
Orbital :
It is a region or space where there is maximum probaility of finding electron. There are four types of
orbitals s, p, d, f.
Shape of orbitals :
Quantum numbers :
Atomic orbital can be specified by giving their corresponding energies and angular momenta which are
quantised (i.e. they can have some specific values). The quantized values can be expressed in terms of
quantum no. They are used to get complete information about electron i.e. location, energy, spin etc.
Principal Quantum Number :
It specifies the location and energy of an electron. It is the measare of the effective volume of the electron
clound. It is denoted by 'n'.
Its possible values are 1, 2, 3, 4, for K, L, M, N.n accordance with Heisenberg’s
Angular Momentum Quantum Number :
It is also called azimuthal quantum no. The quantised values of orbital angular momentum can be
specified in terms of angular momentum quantum no.
 It determines the shape of the orbital. It is denoted by . The permitted values of L are 0, 1, 2 etc. up to n–
1
.
 can have value equal to zero.
 If value of n is 4,  can have 0, 1, 2, 3, represent s, p, d, f orbitals.
where mVr is angular moment.
Magnetic quantum number :
A electron having angular momentum produces a magnetic field which canbe observed.
 Denoted by mL
 Its value depends on value of 'L'.
 It determines magnetic orientation of an orbital.
 Its permitted values are – to + l inculding zero.
 It has total number of values equal to 2 + 1.
 It  = 1 then m = –1, 0, +1
Where, is magnetic moment
and
n is no. of unpaired electrons.
Spin quantum number :
It indicates the direction in which electron revolves. spin is magnetic property and is also quantissed. It
has two permitted values .
Spin moment =
Pauli exclusion principle :
No two electron in an atom can have the same four quantum no.
It can also be stated as – An orbital can have maximum two electrons & they must be opposite spin
quantum numbers.
Afbau Principle :
Electron are filled in various orbitals in the increasing order of their energies i.e., orbital having lowest
energy will be filled first & the orbital having highest energy will be filled last.
Rules for getting sequence of filling the orbital with electrons :
 Orbitals fill in the order of increasing value of (n + ) This means that between 3d, and 4s orbitals. 4s
(n + = 4 + 0 = 4) will fill before 3d (n + l = 3 + 2 = 5)
 If two orbital have same value of (n + ), the one with lower 'n' will be filled first . Thus between 2p and
3s. 2p (n + = 2 + 1 = 3). 3s (n +  = 3 + 0 = 3)
2p will be filled before 3s.
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s 4f < 5d < 6p < 7s < 6d < 7p is sequence of
orbital in which electones are filled.
Hund’s rule of maximum multiplicity :
No electron pairing takes place in p,d, & f orbitals untill each orbital in the given sub-shell contains one
electron having paralled spin e.g., N(7) has electronic configuration 1s2, 2s2 2px12py1 is correct according
to Hund’s rule. whereas.
1s2 2s2 2px2 2py1 is wrong.
Stability of Half filled & completely filled orbitals :
(i) Symmetrical distribution of electrons :
We know that symmetry leads to stability. Half filled & completeily filled orbitals are symmetrical,
therefore more stable.
(ii) Exchange Energy :
When two or more electrons with same spin in the degenerate orbitals, they can exchange their positions
& energy released is called excharge energy.
Greater the no.of exchanges, greater will be exchange energy & more will be stability. In case of half
filled & completely filled orbitals, maximum no. of exchanges are possible, therefore, it leads to
maximum stability.