Friday, August 3, 2012

Notes of Atomic Structure


 Quality Education At Every level

 Chapter #03 
Atomic Structure 
 Prepared by:

Lecturer. Syed Fayyaz Hussain



Atomic Structure

Introduction:

Atom is a smallest particle of an element which can take part in a chemical reaction. Atom consist of subatomic particles such as electron, proton and neutron etc. evidence for the presence of electron, proton and neutrons in the atom is derived through many experiment.

FARADAY’S EXPERIMENT:
         
The substance which can dissociate into their ions can conduct electricity is called electrolyte. This phenomenon was studied in great detail
By Faraday’s .He observed that,
(a)    Electrolyte gave cations and anions on dissociate.
(b)   Cations migrate towards cathode.
(c)    Anions migrate toward anode.
(d)   These ions neutralized their charges by gaining or losing electrons.
(e)    Finally, they get deposited at the electrodes.

CONCLUSIONS:
Accurate determination showed that 96490 coulombs of electricity liberated one gram equivalent of the substance at the electrode. Hence there is some elementary unit of electric charges associated with these ions and it was found that its value is
ē = 1.60 x 10-19 coulombs.
The ions were observed to carry some integral multiply of this charge. The basic unit of electric charge was letter named by stony as “ELECTRONIC CHARGE”.

Discharge Tube Experiment:

All gases air are bad conductors of electricity at normal pressure. But on high voltage and low pressure, they become good conductor. The conduction of electricity was first studied by WILLIAM COOKS. The apparatus used for this purpose is called “DISCHARGE TUBE”.

Construction:

Discharge tube consists of a cylindrical glass tube closed at both ends and fitted with two metallic electrodes. These electrodes are connected to the positive and negative terminals of a battery.
The discharge tube also posses a side tube which is connected to a vacuum pump in order to remove the gas or air from it. The removal of air or gas reduced pressure inside the tube.
                
Working:

When a high voltage is applied to a discharge tube at normal pressure, no phenomenon is observed. But when the vacuum pump is started and most of the gas inside the discharge tube is pumped out and the pressure is reduced to 1 torr, the tube soon begins to emit a soft glow. This gives an indication that the gas in the tube begins to conduct electricity. As the pressure is further reduced, the glowing rays move towards anode. Since these rays are produced at the surface of the cathode therefore these are called “CATHODE
RAYS”.
At still lower pressure about 0.01 torr, the flow from within the tube fades out and dark spaces appears in the discharge tube. At this stage the glass tube at the anode end begins to emit a greenish fluorescence.

Properties of Cathode Rays:

The various experiments performed by different scientists indicate that the cathode rays possess the following characteristics:
1.         These rays travel in a straight lines perpendicular to the cathode surface.
2.         These rays produce sharp shadows if an opaque object is placed in their path.
3.         These rays emerge from cathode and can be focused by using a concave cathode.
4.         These rays can penetrate small thickness of meter e.g. Aluminum or Gold foils
             without producing any perforation in it.
5.         These rays can exert mechanical pressure, showing they possess Kinetic Energy.
6.         These rays are deflected in a magnetic field. This behavior   indicates that they consist
 of charged particles.  
7.         These rays are deflected towards positively charge plate in an electrical field,
            indicating that they consist of negatively charged particles called ELECTRONS.
8.         The charge to mass ratio (e/m) of these rays is 1.76 x 1011 coulomb / Kg.
9.         The charge and mass of these rays are 1.6 x 10-19 coulomb & 9.1 x 10-31 Kg
            respectively.
10.       The rays were seen neither to depend on the material of which the electrodes were
            made nor upon the gas which is filled in the tube.

Conclusion:
On the basis of these properties, it was concluded that cathode rays are negatively charged particles called “ELECTRONS”. Since nature of the cathode rays does not change with the nature of the gas and the cathode used in the tube, hence it can be safely that electrons are the fundamental particles of all atoms.

DISCOVERY OF PROTON:

During the study of the passage of electricity through gases at low pressure, it was observed by Goldstein that, in addition to cathode rays there are also other rays traveling in opposite direction to the cathode rays, and if the cathode is perforated, then these rays can pass through the perforations (canals) and accumulate at the behind of cathode. He named these rays as “Positive Rays” or “Canal Rays”.

Properties of Positive Rays:

1.These rays travel in a straight line.
2.These rays are deflected by electric and magnetic fields opposite to the cathode rays.
   Therefore these are positively charged rays.
3.These rays are not emitted from anode but are produced from       the ionization of gas as a
    result of bombardment of electrons.
4.Their positive charge was found to be equal to that of electron or             simple multiple of it.
5.These particles were found to be much heavier than electrons.       The mass of these
    particles were found to depend on the kind of      gas taken in the discharge tube but it is
    never less than that of an atom or hydrogen.
6. Positive rays unlike cathode rays have different value of e/m       depending on the gas
     present in the discharge tube.

Conclusion:
The charge to mass ratio (e/m) of these rays depends upon the natural of the gas in the discharge tube. The highest (e/m) ratio is obtained in case of hydrogen gas. Therefore hydrogen without electron (H+) is the smallest positive particles of matter. This particle was given the name “PROTON” by Goldstein in 1986. Later on Rutherford showed that proton like electron is also a fundamental particle of matter.

Discovery of Neutron:

The total mass of the protons and electrons in each atom are not sufficient to account for the atomic masses of the different elements. For this reason, scientists believed that there must be another heavy particle inside the atom.
In 1932, James Chadwick discovered the third particle in the nucleus of atom by means of artificial radioactivity.
Chadwick bombarded the nucleus of Beryllium atoms with  a particles and found that it gave highly penetrating radiation. Chadwick put forward the suggestion that these penetrating radiation were due to material particles with mass comparable with that of an atom of hydrogen but carrying no charge. These particles are called “Neutrons” The neutrons must have come out from atoms on disintegration of the bombarded element. This is indicated by the equation.
                        4Be92He4                         C12   +   0n1

Radioactivity:
The natural phenomenon in which certain elements emits invisible radiations is called “Radioactivity” and the substances which give out these radiations are called radioactive substance.
If the radioactive rays are passed through a strong magnetic field as shown in the diagram, they split into three kinds of rays, which are as follows:
(a)        Alpha a rays
(b)        Beta b rays
(c)        Gamma g rays


Properties of a -Rays:

1.         These rays consist of stream of doubly positively charged particles which are four times as heavy as hydrogen atoms. So        they are He nuclei or Helium ion (He++).
2.         The velocity of these particles is  of the velocity of light.
3.         They are deflected by electrical and magnetic field and moves towards –ve plate.
4.         They produce intense ionization of the gas through which they        pass.
5.         Due to heavy mass and low velocity, these particles have low          penetration power. Therefore, these particles can hardly pass through thin layers of solids or a few centimeters of gas.
6.         These particles can effects photographic plate.

Properties of b Rays:

1.         They consist of negatively charged particles which have same e/m ratio as the cathode rays, so they are actually electrons.
2.         Their velocity is about ten times greater than that of a rays.             Hence it is nearly equal to the velocity of light.
3.         These rays are deflected by electric & magnetic fields in the            direction opposite to the a - particles.
4.         Due to lesser mass, the ionization that it produces is small as compared to a - particles.
5.         These rays due to lesser mass and greater velocity have more           penetration power than that of a particles.
6.         The e/m ratio of these rays is equal to that of an electron.
7.         These rays can effects photographic plate.


Properties of g- Rays:

1.         Gamma rays do not consist of particles, actually they are      electromagnetic radiations.
2.         They travel with very high speed which is equal to the velocity of light.
3.         They are not deflected either by electric or by magnetic field. Hence they carry no charge.
4.         They are weak ionizers of gases, due to their non – material nature.
5.         Due to high velocity and non – material nature, they have greater  power of penetration through solid.
6.         Their power to effects a photographic plate is small.
7.         These are more energetic rays than X – rays due to their shorter wave – length.




Plank’s Quantum Theory:

German physicist, Max Plank proposed this theory in 1900 to describe the origin of the radiations from heated bodies. This theory is stated as:
“Radiant energy is emitted or absorbed by a body in the form of small packets, called QUANTA, instead of being continuously. This QUATA of energy are often called PHOTON.
Also the amount of energy given off or absorbed is directly related to the frequency of the light emitted”.

i.e.
E a u
=> E = hu
where  E = energy of quanta
                        u = frequency of radiation
                        h = Plank’s constant
                           = 6.625 x 10-34 J.sec

Spectra:

When an element absorb sufficient amount of energy from a flame or its source its emits radiant energy. When these radiation are passed through a prism, they get separated into different components of colours.
e.g. When a ray of light is passed through the prism, it is dispersed into seven colours. The group of these colours of components is called “SPECTRUM”. Hence.
“A band of rays of different wave – lengths obtained from the decomposition of radiation is called SPECTRUM”.
In a spectrum each colour or component has its own frequency and energy. The light of single wave – length is called “Monochromatic”.
There are two types of spectra.
ü  Continuous spectrum
ü  Discontinuous spectrum

Continuous Spectrum:
“It is spectrum in which different colours are different into each other. There are no dark spaces separating these colours”.
The visible portion of spectrum obtained from ordinary sunlight and lights from incandescent sources are the examples of continuous spectrum.

Line Spectrum:
“In the line spectrum, the bands of colours are separated by dark spaces”.
This type of spectrum may be obtained when light emitted from a gas source pass through a prism. A common way of doing this is to pass an electric current through the gas at low pressure (Crook’s tube). The neon lights used in advertisement make use of this method for producing light and so do Sodium Vapour street lights. If the light from the discharge tube is allowed to pass through the prism, some discrete sharp lines on an otherwise complete dark background are obtained such spectrum is called “LINE SPECTRUM”.

X – Rays:
Professor William Roentgen in 1895 discovered that when cathode rays (electrons) collide with a metal anode, a very penetrating radiation is produced, which were named “X – Rays”. They were also called Roentgen rays”. These rays have the following properties.
Ø  They are invisible rays.
Ø  They can penetrate paper, rubber, glass, metal and human flash.
Ø  They have very short wave length.
Ø  They have high energy and are electromagnetic in nature.
Ø  Different metals produce X – Rays of different wave lengths.

Discovery of Atomic Number:
In 1913, Henry Moseley bombarded different metal anodes with cathode rays in order to det ermine the magnitude of positive charge in atoms. From the study of different wavelengths of the metals, the concluded that wavelengths of the X – rays emitted decreased regularly with the increase in atomic mass. On Careful examination he found that this is due to a fundamental factor in the nucleus, i.e. the number of positive charges in the nucleus. He called the number of positive charges contained in the nucleus as “ATOMIC NUMBER (Z)

Rutherford’s Model of Atom:
            After the discovery of electron, proton & neutron, the attempts were made to see how these particles are arranged in an atom. On the basis of following experiments, Rutherford in 1911 not only discovered nucleus of the atom but also proposed a model of atom.

Experiment:
            Rutherford performed a lot of experiments on the scattering of a particles on thin metal foils like gold. A beam of a - particles                   
(i.e. He2+ ions) was obtained from radioactive element polonium. He bombarded thin foil of gold with a particles and observed the effect on a fluorescent screen.

Observations:
            In his experiment, Rutherford observed that most of the a particles passed through the foil without any deflection, but a few particles were deflected to very great extent. Also, he was found that if 0.0004 cm thick gold foil is used then only one out of 8000 particles suffered a deflection through an angle greater than 90o.

Conclusions:
            On the basis of these observations, Rutherford put forward the following conclusions.
1.         Since majority of the a particles, pass straight and undeflected through foil, it indicates that most of the volume occupied by atom is extraordinary empty.
            2.         The deflection of only few a particles which are positively charged indicates that centre of atom has positive charge.           Rutherford call the centre of atom NUCLEUS.
3.         The total mass of the atom is concentrated in the nucleus, whose dimension is negligible when compared with radius of atom.
4.         The atom as a whole is neutral, therefore it was concluded that        the number of positively charged particles (protons) within the    nucleus must be equal to negatively charged electrons.
5.         The electrons are not stationary but revolving around the nucleus with a very high speed. The centrifugal forces resulting            from the motion of electrons balances the electrostatic attraction of positive nucleus and keeps the electros away from nucleus.

Weakness of Model:
            The Rutherford’s atomic model has two serious defects:
1.         According to Maxwell’s theory a revolving electron will emit energy continuously. Hence due to decrease in energy, the orbit          of revolving electron will become smaller and smaller until it would drop into the nucleus. But it never happen so.
2.         Since revolving electron according to Rutherford emits energy continuously, so it should give continuous spectrum but in actual practice a line spectrum is obtained.

Bohr’s Atomic Theory:
            The weakness in the Rutherford’s model and the formation of line spectrum were improved by Neil Bohr, who proposed a new theory to explain the electronic structure of the atom in 1913. This theory is based upon the following assumptions.

Assumptions:

1.         Electrons in an atom revolve around the nucleus in fixed circular which he called orbits or energy levels.
2.         As long as an electron revolves in a particular energy level it does not emit or absorb energy.
3.         When an electron absorbs energy, it moves to a higher energy level, further away from the nucleus. When it loss energy, it returns to a lower energy level, closer to the nucleus and the            energy is emitted as light.
4.         The electron loses a definite quantity of energy called “Quantum”, when it jumps from an orbit of higher energy level to lower energy level.
5.         The energy is emitted in the form of radiations. The frequency of the energy emitted is directly proportional to the difference in     energy between two levels.

            i.e.       E2 – E1 a u
                        E2 – E1 = hu
            D E = hu
            where, E1 = Energy of electron in 1st orbit.
                                    E2 = Energy of electron in 2nd orbit.
                                    DE = Energy difference b/w two levels.
                                    u   = Frequency of the energy emitted.
                                    h   = Plank’s Constant
6.         The angular momentum (mvr) of an electron in any orbit is integral multiple of .
            where, m         = mass of electron
                           v       = velocity of electron
                           r        = Radius of the orbit
                           n       = Quantum number
                                    = 1, 2, 3, ……………

Application of Bohr Theory:

§  In Hydrogen Atom:
By applying the concept of the quantum of energy and the laws of classical mechanics Bohr worked out a mathematical expression for the derivation of radius, energy, frequency & wave number of an electron moving with a certain orbit in case of Hydrogen atom

Radius of an Orbit:
Suppose an atom of Hydrogen with atomic number ‘Z’ and electro with mass ‘m’ , charge ‘e’ is revolving around its nucleus at the distance ‘r’.
The electrostatic force (centripetal force) between the nucleus and the electron would be
The centrifugal force which keeps the electron away from the nucleus would be . Since the electron is in equilibrium.
\ centrifugal force = centripetal force  
\                         -------------- (i)
According to postulate of Bohr’s Theory:
Þ                    -------------- (ii)
Put this value of ‘v2’ in equation (i)
we have
Þ       
Þ       
Þ       
Þ       
Þ                  -------------- (iii)

Energy of Electron:
An electron possess Kinetic Energy due to its motion around the nucleus and the potential energy due to the coulombic attraction force of the nucleus.
According to equation
Þ       
Þ       
Þ                         -------------- (iv)
Now  P.E. = Fcoul X Distance
Þ       
Þ       
The total energy of electro will be the sum of its K.E. & P.E.
\ E = K.E + P.E
Þ       
Þ                          -------------- (v)
We know that
\        

Þ       
Þ               -------------- (vi)


Frequency:
Suppose and electron with energy E2, jumps from higher energy state n2, to a lower energy state n1, with energy E1. Then the energy released DE by the electron will be
DE = E2 – E1
The energies E2 and E1 of the electron will be
and
\
Þ        22
Þ        -------------- (vii)
According to Bohr’s Postulate,
DE = hu
Þ       
Þ              -------------- (viii)

Wave Number:
The number of waves per unit distance is called “Wave Number” , if “C” is the velocity of light then the wave number is given by
Putting this value in eq. (viii) we have
Þ       
Þ             -------------- (ix)
But
is a constant term which is called Rydberg constant denoted by RH and has value 109678 cm-1.
Þ                       -------------- (x)
The expression is for wave number.
SPECTRUM OF HYDROGEN ATOM:

Although Hydrogen atoms contains only one electron, its spectrum gives a large number of series. Balmer in 1885, studied the spectrum of Hydrogen . he found that , when energy is supplied to the sample of hydrogen gas, individual atoms absorb different amount of energy. The electrons in higher energy levels are unstable & drop back to the lower energy levels &during this process energy is emitted in the from of line spectrum containing various lines of particular frequency & wave length.
                 Balmer observed that a series of lines appeared in  the visible region when the electrons drop from 3rd , 4th ,----nth energy levels to the second orbit . These spectral line are known as “BALMER SERIES”. He proposed an empirical formula to find wave no.(u) of each line.
where n2 = 3,4,5,…………
Lyman later on discovered another series in uv-region . wave no. of each line was sound by similar formula.
where n2 = 2,3,4,…………
Paschen discovered another such series in infrared region. Wave no. of each line was given by:-
Þ             
where n2 = 4,5,6,……………
Bracket found another series infar-infrared region . Wave no. of such lines was given by.     
Þ             
where n2 = 5,6,7,……………
Pfund also found another series far-infrared region. Wave no. of each line was given by:-
Þ             
where n2 = 6,7,8,……………

Heisenberg’s Uncertainty Principle:

Statement:
According to this principle,
“The position & the momentum of an electron cannot be determined accurately simultaneously. One is measured more accurately, the other becomes equally more uncertain at any given instant.”


Mathematical Expression:
If DPx is the uncertainty in the determination of the momentum of a particle & Dx was the uncertainly in the simultaneous determination of it’s position , then the product of these two uncertainties is given by:
DPx . Dx @ h
Thus if one of the two i.e. Px or X was known exactly , then the uncertainty in its other would become infinite . It means that the certainty of determination of one property introduces uncertainty 
For the determination of other.

Explanation:
The uncertainty arise due to the fact that a light with shorter wave length than the located its position. But the momentum of photon (particle of light) increases with decreasing its wave-length.
An electron is so small a particle, that it will be disturbed from its position, on colliding with a high momentum photon particle used to locate it, & hence it is  not possible to say as to what actual position of electron is,

Orbit:
“The fixed circular path on electron  around the nucleus called “Shell”
or “Energy Level” or “Orbit”.
These orbit are designated as K, L, M, N, etc. The maximum number of electron in one orbit is “2n2“, where “n” is the number of orbit. These orbit possess a definite amount of energy increasing outwards from the nucleus.

Orbitals:
If a atomic spectrum is observed, the spectral lines are found to be consist of two or more fine lines closely packed together. Thus the electron In the same orbit may differ in energy by small amount. Thus energy level are further divided into           Sub – Energy Levels or  “Orbitals” & can be defined as          
“Orbitals  are  the  regions  around  the nucleus in which the probability of finding the electron is maximum.”
The Orbitals have been named as s, p, d, f, etc. The maximum no. of electron in s, p, d, f, are 2, 6, 10 & 14 respectively. Each orbit has “Orbital” equal to its quantum number ‘n’. Thus the first orbit contain 1, the second orbit has 2 (i.e. s & p) , the third orbit has 3 (i.e. s, p & d) & fourth has 4 (i.e. s, p, d & f ) Orbitals.
SHAPES OF ORBITAL
S – orbital:
All’s Orbitals are spherical in shape with the nucleus at the centre. Therefore in an ‘S’ orbital, the probability of  finding the electron is uniformly distributed  around the nucleus. It has only one possible orientation in space, because it  spread over all the three axes uniformly . It has no nodal plane.


P – orbital:

The P - orbital are dumb-bell shaped
and they are oriented in the space
along the three mutually perpendicular
axis (x,y,z) and are called Px, Py and Pz
orbitals. All the three P- orbitals are
perpendicular to each other.
These are degenerated orbitals, that are equal energy. Each P – orbital has two lopes, one of which is labeled (+) and the other is labeled (-).

Quantum NumberS:
In 1926 Schrödinger was a mathematician who calculate the probability of location of the electrons in an orbital. Quantum Numbers are the constant used in Schrödinger. Wave equation to describe the energy of an electron, the shape of orbitals an orientation in space around the nucleus of atom. The four quantum numbers are:
ü  Principle quantum number
ü  Azimuthal quantum number
ü  Magnetic quantum number
ü  Spin quantum number

Principle Quantum NumberS:
Principle quantum number describes the size of an orbit and is represent by ‘n’. They have any integral value i.e. n=1,2,3,………………… it never zero. The size and the energy of orbital increase with increase the value of ‘n’.

Azimuthal Quantum NumberS:
Azimuthal quantum number describes the shape of an orbital and its represented by ‘l’. It can have integral values ranging from zero to n-1 i.e. l=0,1,2,…………… (n-1)
If l = o,            the orbital is called s orbital
If l = 1,            the orbital is called p orbital
If l = 2,            the orbital is called d orbital

Magnetic Quantum Numbers:
The magnetic quantum number describes the different orientation of an orbital in the space in applied magnetic field and it is represented by ‘m’. The value of m =- l to + l through zero. i.e. , m = - l ….0….. + l.
If l = 2
Then m = -2 , -1, 0 , +1, +2
i.e. d-orbital has the five orientations.

Spin Quantum Numbers:
It describes the direction of spin of an electron around the nucleus of an atom and is represented by ‘s’. Its values are = + ½ and – ½.
+ ½ spin for clockwise direction
- ½ spin for anti clockwise direction

Principle of Electronic Configuration:

The distribution of electron in the orbitals of the atom is called electronic configuration. Following rules are used in the filing of the electrons in any orbital.
(i)                             Pauli exclusion principle
(ii)                           Aufbau principle
(iii)                         (n+ l) rule
(iv)                         Hund’s rule

Pauli Exclusion Principle:

According to this principle “No two electron of the same atom will have the same value of all its four quantum numbers”.
Therefore in an atom two electrons may have a maximum of three same quantum numbers, same value but the fourth would be different. Thus in an orbital, when the value of n, l and m are same , the two electrons can occupy the same orbital only there spin are opposite.
e.g. for 1st K shell
n=1, l=0, m=0, s=+ ½ 
n=1, l=0, m=0, s=- ½ 

Aufbau principle:

According to this rule the electron are fed in various orbitals in order of increase orbital energy starting with the one ‘s’ orbital”.
Hence electronic configuration of an atom can be built up by placing the electrons to the lowest available orbital until the total number of electrons added is equal to the atomic number ‘Z’. The sequence of increasing orbital energy is below
1s, 2s,2p,3s,3p,4s,3d,4p,……………

(n+ l) Rule:

According to this rule “In builting up the electronic configuration the orbitals with the lowest value of (n+l)fills first, but when the two orbitals have the same value of (n+ l), the orbital with the lower value of ‘n’ fills first”.
Here ‘n’ and ‘l’ stands for principal and azimuthal quantum numbers respectively. Actually this rule is used to determine the energy of any orbital and then we can apply the Aufbau Principle. e.g. 3d – orbital is filled later than 4s – orbital
\         3d – orbital      has (n+l ) = 3+2 = 5
            4s – orbital has (n+l ) = 4+0 = 4
Similarly 4p – orbitals fills before 5s – orbital because 4p orbital has lesser value of ‘n’.
\         4p– orbital       has (n+l ) = 4+1 = 5
            5s – orbital has (n+l ) = 5+0 = 5


Hund’s Rule:
According to this rule “When the orbitals of same energy levels are available, then electrons are distributed in orbitals in such a way as to give the maximum numbers of unpaired electrons . Only when the orbital are separately occupied then the pairing of electrons commences”.
This rule explain the filling of electrons in degenerate (having same energy) orbitals like p,d,f. In simple, it state that the electrons remain unpaired as far as possible. i.e. If there are available orbitals of equal energy to the electrons, the electrons would lie in separate orbitals and have same spin rather than to lie in the same orbital and have paired spin.

e.g. (N(Z=7) = 1s2, 2s2,2Px , 2Py , 2Pz        is true

       and  not                      1s2, 2s2,2Px, 2Py, 2Pz

Atomic Radius:
“It is half distance between two nuclei of two identical atoms covalently bounded together”.
Atomic Radius is measured in Angstrom (Ao) or ‘nm’ units
1 Ao = 10-8cm = 10-10m
1nm = 10-9m

Trend in the Periodic Table:
The atomic radii of elements regularly decrease from left to right along a period. Because increasing atomic number along a period increase the positive charge in the nucleus, resulting in attraction of the outermost shell electrons more and more in word and thus the shell are contracted.
The atomic radii of elements regularly increase from top to bottom in a group. Because as we move from 1 element to the other down a group, a new shell is added. This results in the expansion, of the size of the atom on one hand and the reduction in the attraction of the nucleus for the electron on the other. Therefore the atom radii of the elements increase down the group.

Ionic Radius:
Ions are formed due to loss or gain of electron from the outermost shell. Positive ion is called cation and negative ion is called anion. Ionic Radius is defined as “The radius of spherical ion”.
The ionic radii of cations are smaller from which they are formed because when an atom looses electron to form a positive ion, the outermost shell vanishes away and the next shell becomes the outermost shell, while the quantity of the charge in the nucleus remains the same. As the result of it the attraction of the nucleus for the outermost shell electrons increases causing decrease in the ionic radius.
On the other hand anions are larger than the atoms from which they are formed because in the case of negative ion the electrons are added in the same outer most shell. Therefore the ions tend to expand with the increasing negative charge on them, resulting in the increase in their ionic radii.



Ionization Potential:
“The minimum amount of energy require to remove completely the outer most electrons from an atom in gaseous state resulting in the formation of positive ions is called Ionization Potential” (IP). Simply ionization energy.
When the first electron is removed it is called first ionization energy when the second, third and fourth electron are removed the energies used are called second, third, fourth ionization energy value respectively and son on. The value of I.E. values of element increase from first I.E. to second I.E. to third I.E. and so on.
Al        1st I.E.            Al+ + e-, DHo = +518 Kj/mole
Al+       2nd t I.E.                      Al2+ + e-, DHo =+1817Kj/mole
Al2+      3rd  I.E.                      Al3+ + e-, DHo =+2745Kj/mole

Trends in the Periodic Table:

I.E. decrease down a group:
As we go down a group the addition of new shell increase the atomic radii as well as shielding effect. This result in the decrease in I.E. of the elements down a group in the periodic table.

I.E. increases along a period:
As we go from left to right in the periodic table the atomic number (+charge in the nucleus) of elements increase regularly. This increases the attraction force between the outermost shell electron and the nuclei resulting in the decrease in the atomic radii. Consequently the outermost electrons are held more strongly in the atom and thus more energy is required from this removal. Therefore the ionization of element increase along a period from left to right.

Electron Affinity:
“It is the amount of energy released, when an electron is added to the outermost shell of gaseous atom to form a negative ion”.
Electron Affinity (E.A.) is effected by the factors such as atomic size, nuclear charge and shielding effect of the electrons.
The E.A. increases from left to right in a period due to decrease in the atomic size and increase in the nuclear charge. But it decrease down a group from top to bottom because the atomic size increase.

Electronegativity:
“Electronegativity is the force of attraction of an atom to attract a shared pair of electron towards itself” All elements have different values of electronegativities. Fluorine is the most electronegative element and it used as standard for the comparison of electronegativities of other element. Its value is fixed as 4.

Tend in the Periodic Table:
The electronegativity values of the element increase from left to right in a period because of increase in nuclear charge and the decrease in the size of atom.
On other hand electronegativity values of the elements decrease from top to bottom in the group due to increase in atomic size.

Nature of Chemical Bond:

The nature of a chemical bond can be decided on the basis of electronegativity difference between the two bonded atom i.e.
i.          “When E.N. difference is less than 0.5 it will be a non polar covalent bond”.
ii.         “When the E.N. difference is between 0.6 to 1.7, it will be              Polar – Covalent bond.
iii.        “When E.N. difference is greater than 1.7 then it will be “ionic bond”.
e.g.

(a)        C – H , is a non Polar Covalent bond
\ E.N. difference = 2.2 – 2.1 = 0.4 < 0.7
(b)        H – Cl, is a Polar Covalent bond
            \ E.N. difference = 3.5 – 2.1 = 1.4 > 0.7 but < 1.7
(c)        Na – Cl, is a ionic bond
            \ E.N. difference = 3.5 – 0.9 = 2.4 > 1.7






Table 5.1 Highlights in the Development of Contemporary Atomic Theory 

Year
Name of Scientists
Work
1803

John Dalton (English)

Proposed an atomic theory to explain the laws of definite composition and multiple proportions.

1879

William Crookers (English)

Repeated the observations of earlier workers and concluded that rays produced by an electric current in an evacuated tube consist of a stream of charged particles produced at the cathode.
1885
Johann J. Balmer (Swiss)
Developed a simple empirical equation that could be used to calculate the wavelength of lines in the spectrum of hydrogen atoms.
1886
Eugen Goldstein (Germen)
Characterized “canal rays” (positive ions) produced by an electric current in an evacuated cube.
1895
Wilhelm Roentgen (Germen)
Discovered X rays produced by an electric current in an evacuated cube.

1896
Antoine Becquerel (French)
Discovered natural radioactivity.
1890
Johannes Rydberg (German)
Modified Balmer’s equation to describe the frequencies of the lines in the hydrogen spectrum.
1897
J. J. Thomson (British)
Showed that Crookes’s cathode rays are beams of negativity charged particles (now known to be electrons) and calculated their charge to-mass ratio.
1898
Marie Curie (Polish) and her Husband, Pierre (French)
Isolated polonium and radium from pitchblende (a lead ore) by chemical processes.
1900
Max Planck (German)
Proposed a quantum theory that described the light emitted from a hot object as composed of discrete units called quanta or photons.
1903
Hantaro Nagoka (Japanese)
Postulated a saturn-like atom, a positively charged sphere surrounded by a halo of electrons.
1904
J. J. Thomson
Proposed a model of the atom with electrons embedded in a sea of positive charges, the “plum-pudding model.”
1905
Albert Einstein (German)
Described the relationship of mass and energy.
1909
Robert Milliken (American)
Determined the charge on the electron in his oil drop experiment.
1911
Ernest Rutherford (New Zealander)
Attributed the scattering of a particles by a thin gold foil to an atomic structure in which an atom’s mass is concentrated in a tiny, positively charged nucleus.
1912
J. J. Thomson
Detected the isotopes neon-20 and neon-22.
1913
Niels Bohr (Danish)
Showed that the hydrogen spectrum could be explained by a previously unrecognized property of matter-the energy of electrons in atoms is limited to certain values (quantized).
1913
Henry Moseley (English)
Observed that an element emits characteristics X rays that depend on its nuclear charge (atomic number).
1920
Orme Masson (Australian), William Harkins (American), Rutherford
Independently postulated the existence of an uncharged particle with the mass of proton (the neutron).
1924
Louis de Broglie (French)
Combined equations from Einstein and Planck to suggest that electrons have wave-like properties.
1925
Wolfgang Pauli (German)
Stated that no two electrons in the same atom can have the same set of quantum numbers (Pauli’s exclusion principle).
1926
Erwin Schrödinger (Austrian)
Combined the particle nature of an electron, its wave properties, and quantum restrictions and developed an equation that described the energy and likely location of electron in a probability relationship.
1927
Werner Heisenberg (German)
Showed that we cannot determine simultaneously both the exact position and the momentum of an election (Heisenberg’s uncertainty principle).  
1927
Frederick Hund (German)
Determined that electrons in subshells have maximum unpairing and that unpaired electrons have the same spin (Hund’s rule).
1932
James Chadwick (British)
Characterized neutroms.

Ref# Page no. 144 and 145 General Chemistry 10th edition by Robinson, Odom, Holtzclaw 


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