ATOMIC STRUCTURE
Atoms
• The word “Atom” is derived from the Greek word “Atomio or Atomos” meaning indivisible or uncut-able.
• Atoms are the smallest particles until the 19th century.
• “John Dalton” conducted various experiments & stated Dalton’s Atomic Model (discarded now).
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Dalton’s Atomic Model
•
Atoms are small indivisible particles which makes
up matter.
•
Matter can neither be created nor destroyed.
•
Atoms of the same element are similar in all aspects.
•
Conversely, atoms of different elements are different
in all aspects.
•
Atoms combine in small whole numbers to form compounds.
•
Atom is the smallest unit of matter that takes
part in chemical reactions.
Defects of Dalton’s
Atomic Model
•
Atoms can be further divided into subatomic particles
like electrons, protons and neutrons.
•
Nuclear fission and fusion reactions show that
atoms can be created and destroyed.
• The Discovery of isotopes proves that atoms of same the element need not be similar in all aspects.
• The Discovery of isobars proves that atoms of different
elements need not be different in all aspects.
• The Discovery of polymers and macromolecules proves
that atoms can combine in large numbers.
Some terms of Chemistry
•
Isotopes: Atoms of the same element having the same atomic number but a different mass number.
•
Isobars: Atoms of different elements having the same
mass number with different atomic numbers.
•
Isotones: Atoms of different elements having the same number of neutrons.
•
Isoelectronic species: The species having the same number of electrons.
•
Sub-Atomic Particles: Particles that are smaller than an atom.
Discharge tube
•
The discharge tube consists of a cylindrical glass
tube about 50 cm long & 4cm in diameter
with two electrodes at two ends.
•
The pressure of the gas inside can be decreased and can be controlled by a gauge.
•
A high potential difference of 10,000 to 15,000 volts is set between the electrodes.
•
This apparatus helped to find some sub-atomic particles.
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Discovery of Electron
•
When sufficient high voltage is applied inside the discharge tube, current flows like a stream of particles from cathode to anode.
•
These were called cathode rays or cathode ray particles.
•
By making holes in the anode and by coating the tube behind the anode with a phosphorescent material (zinc sulphide), a bright spot on the coating is developed due to the striking of these rays on that material. (the same thing happens on a television set)
•
The name Electron was given by J.J.Thomson.
Properties of Cathode Ray Particles
•
The cathode rays start from the cathode and move towards
the anode.
•
These rays are
not visible but their behaviour can be observed
with the help of certain kinds of materials (fluorescent or phosphorescent) which glow when hit by them.
•
In the absence of an electrical or magnetic field,
these rays travel in straight lines.
•
In presence of an electric or magnetic field, the behaviour of cathode rays are similar to
negatively charged particles, electrons.
•
The characteristics of cathode rays (electrons)
do not depend upon the material of electrodes
or the nature of the gas present in the cathode
ray tube.
•
Thus electrons are the basic constituent of all the
atoms.
Charge to Mass Ratio of Electron
In 1897, British physicist J.J.Thomson measured the ratio
of charge of an electron to the ratio of mass
of an electron by placing electrical and magnetic fields perpendicular to the path of electrons in
cathode-ray tube.
Thomson argued that the rate of deviation of a particle depends
on: Magnitude of charge of the electron( Directly
Proportional)
Mass of the electron(Indirectly Proportional)
Strength of Magnetic Field & electric Field(Directly Proportional)
•
Thomson’s apparatus :
•
-e/m = 1.758820 × 10^11 C/kg
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Charge of an electron
•
Millikan’s Oil Drop Experiment:
•
Oil droplets in the form of mist entered through
a tiny hole in the upper plate of electrical
condenser.
•
The downward motion of these droplets was
viewed through the telescope with a
micrometre eyepiece.
•
By measuring the rate of fall of these droplets,
Millikan was able to measure the charge of
oil droplets.
•
The air inside the chamber was ionized by passing
a beam of X-rays through it. The electrical
charge on these oil droplets was acquired by collisions with gaseous ions.
• The
fall of these charged oil droplets can be changed depending on the charge on the droplets, the polarity and strength of
the voltage applied to the plate.
•
By measuring the effects of electrical field strength
on the motion of oil droplets, Millikan concluded
that the magnitude of electrical charge, q, on the droplets is always an integral multiple of the electrical
charge(e), q = n*e, where n = 1, 2, 3....
Millikan’s apparatus :
He found that the charge
on an electron is
– 1.6 × 10^–19 C.
The present accepted
value of the electrical charge is – 1.6022
× 10^–19 C.
The mass of the electron
( me) was deter mined by combining these results with Thomson’s value of e/m ratio.
e/( e/m ) = 1.6022*(10^-19)C/
1.758820 *(10^11)C/kg =
m = 9.1094×10^–31 kg
Discovery of the proton:
•
Electrical discharge in the modified cathode ray
tube led to the discovery of particles carrying
positive charge, also known as canal rays.
•
The characteristics of these positively charged
particles are
(i)
The positively charged particles depend upon
the nature of gas present in the cathode
ray tube.
These are simply the positively charged gaseous ions. These rays travel from anode to cathode.
(i)
The charge to mass ratio of the particles
was depending on the gas from which these
originate.
(ii)
The behaviour of these particles in the magnetic
or electrical field is opposite to that observed
for electrons or cathode rays.
•
The smallest and lightest positive ion was obtained
from hydrogen and was called a proton. This
positively charged particle was characterised in 1919.
Discovery of Neutron
·
Indirect interference: The mass of a helium atom
is approximately 4 amu. But, only there are
2 proton atoms each of mass of 1 AMU.
·
The mass of electrons is very low(negligible).
A need was felt for the presence of electrically neutral particle
as one of the constituents of an atom.
·
Chadwick (1932) discovered and named neutron by
bombarding a thin sheet of beryllium by α-particles.
· When an alpha particle collided with a beryllium atom, it released a carbon atom & a neutron.
Atomic Models
•
The three main atoms models which were proposed are :
Thomson’s model. Rutherford’s
model. Bohr’s model.
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Thomson’s atomic model
·
J. J. Thomson(1898) proposed that an atom possesses
a spherical shape (radius approximately 10^-10
m) in which the positive charge is uniformly
distributed & electrons embedded in it in
the most electrostatically stable manner.
·
It looks like a watermelon.
·
Hence, it is called plum-pudding, watermelon
or raisin model.
Alpha- Particle scattering experiment:
In the Gold foil experiment, they placed a very thin gold foil in the centre of an alpha ray detector.
They used Polonium as an alpha particle source.
Whenever an alpha
particle hits a nucleus of an atom, it
gets deflected.
Results
(i) Most of the space in the atom is empty
as most of the α– particles passed through
the foil undeflected.
(ii) A few positively charged α– particles
were deflected. The deflection must be due
to enormous repulsive force showing that
the positive charge of the atom is not spread throughout the atom as Thomson had presumed. The positive charge has to be concentrated in a very
small volume that repelled and deflected
the positively charged α– particles.
(iii) Calculations by Rutherford showed that
the volume occupied by the nucleus is negligibly
small as compared to the total volume of
the atom. The radius of the atom is about
10^–10 m, while that of the nucleus is 10^–15 m.
Rutherford’s Atomic Model
(i) The
positive charge and most of the mass of the atom was densely concentrated in an extremely small region. This very small portion of the atom was called the nucleus by Rutherford.
(ii) The nucleus is surrounded by electrons that move around the nucleus with a very high speed in circular paths called orbits
. Thus, it resembles the solar system in
which the nucleus plays the role of the sun and the
electrons of revolving planets.
(iii) iii) Electrons and the nucleus are held
together by electrostatic forces of attraction.
Defects of Rutherford’s model
•
There is a theory (by Maxwell) stating that whenever
a charged particle accelerates, it emits
radiation.
•
As the electron moves in a circular path, it is
accelerating. So, it must emit radiation.
•
As it emits radiation, it must lose some of its
energy.
•
As it loses some of its energy its orbit must
continuously shrink; until the electron collides
with the nucleus.
•
Mathematical equations show that it will take
only 10^-8 seconds to collide with the nucleus.
•
But, this does not happen in the atom, the atom
is stable.
•
So, Rutherford’s Atomic model is not completely
correct.
•
Another drawback is that his model did not explain
how the electrons were distributed and what
is the energies of these electrons.
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Properties of electromagnetic wave motion
•
The oscillating electric and magnetic fields produced
by oscillating charged particles are perpendicular
to each other and both are perpendicular
to the direction of propagation of the wave.
•
Unlike sound waves or
water waves, electromagnetic waves do not require a medium and can move in a vacuum.
It is now well established that there are many types of electromagnetic radiations, which differ from one another in wavelength (or frequency). These constitute what is called the electromagnetic spectrum.
Finding the nature of light
•
Some of the experimental phenomena such as diffraction
and interference can be explained by the
wave nature of electromagnetic radiation.
•
The things which couldn’t be explained by the wave
nature of electromagnetic radiation:
Black body radiation(The nature of emission
of radiation from hot bodies).
Photoelectric effect(Ejection of electrons
from a metal surface when radiation strikes
it).
Variation of heat capacity of solids as
a function of temperature. Line spectra of
atoms with special reference to hydrogen.
Blackbody radiation
Planck suggested that atoms and molecules
could emit (or absorb) energy only in small
quantities and not in a continuous manner, a belief popular at that time. Planck gave the name quantum to the smallest quantity of energy that can be emitted
or absorbed in the form of electromagnetic radiation.
The energy (E ) of a quantum of radiation
is proportional to its frequency (ν
) and is expressed by the equation:
E = hν
h=6.626×10^–34 J s.
{Planck’s constant}
The frequency must be a whole number.
Presentation Done by Tharun A
Photo-electric effect
•
Hertz’ s experiment: In 1887, H. Hertz performed
an experiment. He used a discharge tube;
but used a little lower voltages (not enough to travel through the tube). He focused
lights of different frequencies on it.
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Results
•
(i) The
electrons are ejected from the metal surface as
soon as the beam of light strikes the surface, i.e., there is no time lag between the striking of the light beam
and the ejection of electrons from the metal
surface.
•
(ii) The
number of electrons ejected is proportional to
the intensity or brightness of the light.
•
(iii) For
each metal, there is a characteristic minimum frequency,ν0(also known as threshold frequency) below which the photoelectric effect is not observed. At
a frequency ν >ν0, the ejected electrons
come out with certain kinetic energy. The
kinetic energies of these electrons increase with the increase of the frequency of the light used.
Explanation
•
Shining a beam of light onto a metal surface
can be viewed as shooting a beam of particles,
the photons.
•
When
a photon of sufficient energy strikes an electron in the atom of the metal, it transfers its energy instantly to the electron
during the collision and the electron is
ejected without any delay.
Greater the energy possessed by the photon,
greater will be a transfer of energy to the
electron and greater the kinetic energy of the ejected electron.
The kinetic energy of the ejected electron is
proportional to the frequency of the electromagnetic
radiation.
•
The striking photon has energy equal to hν and
the minimum energy required to eject the
electron is hν0 (also called work function, W0), then the difference in energy (hν – hν0 ) is transferred as the kinetic energy of the electron.
•
The kinetic energy of the ejected electron is
given:
•
E = hν = hν0 +1/2m(v^2)
Spectrum
•
Spectrum: A type of categorisation of different
waves according to their wavelengths. (Singular
- Spectra)
•
Absorption spectrum: The spectrum that is got
after absorption of certain wavelengths by
a particle.
•
Emission spectrum: The spectrum that is got during
release of photons by an exciting object.
Spectroscopy: A branch of the science that deals with spectrum.
Excited sample: The sample has many atoms in it that have absorbed a lot of radiation.
Uses
•
These became the evidence for quantised electronic levels.
•
Every element has a unique absorption & emission spectrum.
•
This method was used to find elements like caesium & rubidium.
•
Spectroscopy has developed a lot nowadays
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Bohr’s model
•
Neils Bohr (1913) explained the general features
of the structure of hydrogen atom and its
spectrum.
•
Bohr’s model for the hydrogen atom is based on the
following postulates:
•
i) The electron in the hydrogen atom can move
around the nucleus in a circular path (orbits)
of fixed radius and energy.
Arranged in form of concentric circles around
nucleus
•
ii) The energy of an electron in the orbit does
not change with time.
The electron will move from a lower stationery state to a higher stationary state when required amount of energy is absorbed by the electron & vice versa.
iii) The frequency of radiation absorbed
or emitted occurs between two stationary
states that differ in energy by ∆E is given by:
Where E1 and E2
are the energies of the lower and higher allowed energy states respectively. This expression is commonly known
as Bohr’s frequency rule.
iv)In a given stationary state it can be
expressed as in equation:
Where me is the mass of
electron, v is the velocity and r is the radius of the orbit in which the electron is moving.
Thus an electron can move only in those
orbits for which its angular momentum is an integral multiple of h/2p. That means angular momentum is quantised(restricted).
Radiation is emitted or absorbed only
when the transition of electrons takes place from one quantised value of angular momentum to another. Therefore, Maxwell’s electromagnetic theory does not apply here
that is why only certain fixed orbits are allowed.
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