1styear CHEMISTRY Notes Chapter-4

Chapter-4
CHEMICAL BOND

Chemical Bond
Introduction
Atoms of all the elements except noble gases have incomplete outermost orbits and tends to complete them by chemical combination with the other atoms.
In 1916, W Kossel described the ionic bond which is formed by the transfer of electron from one atom to another and also in 1916 G.N Lewis described about the formation of covalent bond which is formed by the mutual sharing of electrons between two atoms.
Both these scientists based their ideas on the fact that atoms greatest stability when they acquire an inert gas electronic configuration.

Definition
When two or more than two atoms are combined with each other in order to complete their octet a link between them is produced which is known as chemical bond.
OR
The force of attraction which holds atoms together in the molecule of a compound is called chemical bond.

Types of Chemical Bond
There are three main types of chemical bond.
1. Ionic bond or electrovalent bond
2. Covalent bond
3. Co-ordinate covalent bond or Dative covalent bond

Ionic Bond OR Electrovalent Bond
Definition
A chemical bond which is formed by the complete shifting of electron between two atoms is called ionic bond or electrovalent bond.
OR
The electrostatic attraction between positive and negative ions is called ionic bond.

Conditions for the Ionic Bond Formation
1. Electronegativity
Ionic bond is formed between the element having a difference of electronegativity more than 1.7 or equal to 1.7 eV.
Therefore ionic bond is generally formed between metals (low electronegative) and non-metal (high electronegative) elements.

2. Ionization Potential
We know that ionic bond is formed by the transference of electron from one atom to another, so in the formation of ionic bond an element is required which can lose its electrons from the outer most shell. It is possible to remove electron from the outermost shell of metals because of their low ionization potential values.

3. Electron Affinity
In the formation of ionic bond an element is also required which can gain an element is also required which can gain electron, since non-metals can attract electrons with a greater force due to high electronegativity. So a non-metal is also involved in the formation of ionic bond due to high electron affinity.

Example of Ionic Bond
In order to understand ionic bond consider the example of NaCl. During the formation of Ionic bond between Na and Cl2, Sodium loses one electron to form Na+ ion while chlorine atom gains this electron to form Cl- ion. When Na+ ion and Cl- ion attract to each other NaCl is formed. The stability of NaCl is due to the decrease in the energy. These energy change which are involved in the formation of ionic bond between Na and Cl are as follows.
i. Sodium has one valence electron. In order to complete its octet Na loses its valence electron. The loss of the valence electron required 495 kJ/mole.
Na ----> Na+ + e- ....................... ?H = 495 kJ/mole

ii. Chlorine atom has seven electrons in its valence shell. It require only one electron to complete its octet, so chlorine gains this electron of sodium and release 348 kJ/mole energy.
Cl + e- ----> Cl- ...................... ?H = -348 kJ/mole
Here the energy difference is 147 kJ/mole (495 - 348 = 147). This loss of energy is balanced when oppositely charged ions are associated to form a crystal lattice.

iii. In third step, positively charged Na+ ion and negatively charged Cl- ion attract to each other and a crystal lattice is formed with a definite pattern.
Na+(g) + Cl-(g) ----> Na+Cl- ........... ?H = - 788 kJ/mole
This energy which is released when one mole of gaseous ions arrange themselves in definite pattern to form lattice is called lattice energy.
From this example, we can conclude that it is essential for the formation of ionic bond that the sum of energies released in the second and third steps must be greater than the energy required for the first step.

Characteristics of Ionic Compounds
1. An ionic compounds, the oppositely charged ions are tightly packed with each other, so these compounds exist in solid state.
2. Due to strong attractive forces between ions a larger amount of energy is required to melt or to boil the compound and hence the melting and boiling point of the ionic compound are generally high.
3. Ionic compounds are soluble in water but insoluble in organic solvents like benzene, CCl4. etc.
4. In the aqueous solution, the ionic compounds are good electrolytes, because in water the interionic forces are so weakened that the ions are separated and free to move under the influence of electric current. Due to this free movement of ions, the ionic compounds conduct electricity in their solutions.

Covalent Bond
Definition
A link which is formed by the mutual sharing of electrons between two atoms is called covalent bond.

Explanation
In the formation of covalent bond, mutual sharing of electron takes place. This mutual sharing is possible in non-metals, therefore covalent bond is generally formed between the atoms of non-metals. For example
In Cl2 molecule, two atoms of chlorine are combined with each other to form Cl2 molecule. Each atom of chlorine having seven electrons in its valencies shell. These atoms are united with each other by sharing one of its valence electron as shown.
Cl Cl: ----> :Cl :Cl OR Cl - Cl
In this molecule, one shared pair of electrons forms a single covalent bond between two chlorine the atoms. With the formation of a covalent bond the energy of the system is also decreased.
Cl + Cl ----> Cl - Cl .............. ?H = - 242 kJ / mole
This released energy lowered the energy of the molecule and the stability of the compound is also increased.

Types of Covalent Bond
There are three main types of covalent bond.

1. Single Covalent Bond
When a covalent bond is formed by sharing of one electron from each atom, that it is called single covalent bond and denoted by (-) single line between the two bonded atoms e.g.
Cl - Cl, H - H, H - Br etc.

2. Double Covalent Bond
In a covalent bond, if two electrons are shared from each of the bonded atom then this covalent bond is called double covalent bond and denoted by (=) two lines e.g.
O = O, O : : O

3. Triple Covalent Bond
When a covalent bond is formed by sharing of three electrons from each atom then this type of covalent bond is called triple covalent bond, and denoted by (=) three lines between the two bonded atoms e.g.
N : : N :, N = N
The bond distance of multiple bonds are shorter and the bond energies are higher.

Characteristics of Covalent Compounds
The main characteristics properties of covalent compounds are as follows
1. The covalent compounds exist as separate covalent molecules, because the particles are electrically neutral so they passes solid, liquid or gaseous state. This intermolecular force of attraction among the molecules.
2. Since the covalent compound exist in all the three states of matter so their melting points and boiling point may be high or low.
3. Covalent compounds are non-electrolytes so they do not conduct electricity from their aqueous solution.
4. Covalent compounds are generally insoluble in water and similar polar solvent but soluble in the organic solvents.

Co-Ordinate OR Dative Covalent Bond
Definition
It is a type of covalent bond in which both the shared electrons are donated only be one atom, this type is called co-ordinate covalent bond.
The 8 ordinate covalent bond between two atoms is denoted by an arrow (?). The atom which donates an electron pair is called as a donor of electron and the other atom involved in this bond is called acceptor. E.g.
A + B ----> A : B OR A ? B

Dipole Moment
Definition
The product of the charge and the distance present in a polar molecules is called dipole moment and represented by µ.
OR
The extent of tendency of a molecule to be oriented under the influence of an electric field is called dipole moment.

Mathematical Representation of Dipole Moment
Suppose the charge present on a polar molecule is denoted by e and the separation between the two oppositely charged poles of the molecules is d, then the product of these two may be written as
e x d = µ
Where µ is dipole moment.

Dipole Moment in Diatomic Molecules
The diatomic molecules which are made up of similar atoms will be non-polar and their dipole moment is zero but the diatomic molecules made up of two different atoms e.g. HCl or Hl are polar and have some dipole moment. The value of the dipole moment depends upon the difference of electronegativities of the two bonded atom. If the difference of electronegativity between the atoms is greater, the polarity and also the dipole moment of the molecule is greater e.g.
The dipole moment of HCl = 1.03 debye
Whereas dipole moment of HF = 1.90 debye

Dipole Moment of Poly Atomic Molecules
In poly atomic molecules, the dipole moment of molecules depends upon the polarity of the bond as well as the geometry of the molecule.

Ionic Character of Covalent Bond
In homonuclear diatomic molecules like Cl2, O2, l2, H2 both the atoms are identical so the shared electrons are equally attracted due to identical electronegativities and hence the molecules are non-polar.
When two dissimilar atoms are linked by a covalent bond the shared electrons are not attracted equally by the two bonded atoms. Due to unsymmetrical distribution of electrons one end of the molecules acquire partial positive charge and the other end acquire a partial negative charge. This character of a covalent bond is called Ionic character of a covalent bond.
The ionic character of a covalent bond depends upon the difference of electronegativity of the two dissimilar atoms joined with each other in a covalent bond. E.g., the H-F bond is 43% ionic whereas the H-Cl bond is 17% ionic. The ionic character greatly affects the properties of a molecules e.g., melting point, boiling point of polar molecules are high and they are soluble in polar solvent like H2O. Similarly the presence of partial polar character shortens the covalent bond and increases the bond energies.

Bond Energy
Definition
The amount of energy required to break a bond between two atoms in a diatomic molecule is known as Bond Energy.
OR
The energy released in forming a bond from the free atoms is also known as Bond Energy.
It is expressed in kilo Joules per mole or kCal/mole.
Examples
i. The bond energy for hydrogen molecule is
H - H(g) ----> 2 H(g) .......................... ?H = 435 kJ/mole
OR
H(g) + H(g) ----> H - H ....................... ?H = 435 kJ/mole
It can be observed from this example that the breaking of bond is endothermic whereas the formation of the bond is exothermic.

ii. The bond energy for oxygen molecule is
O = O(g) ----> 2 O(g) ........................ ?H = 498 kJ/mole
OR
O(g) + O(g) ----> O = O .................... ?H = -498 kJ/mole
Bond energy of a molecule also measure the strength of the bond. Generally bond energies of polar bond are greater than pure covalent bond.
E.g.
Cl - Cl ----> 2 Cl ........................ ?H = 244 kJ/mole
H - Cl ----> H+ + Cl- ................... ?H = 431 kJ/mole
The value of bond energy e.g., triple bonds are usually shorter than the double bond therefore the bond energy for triple bond is greater than double bond.

Sigma & PI Bond
Sigma Bond Definition
When the two orbitals which are involved in a covalent bond are symmetric about an axis, then the bond formed between these orbitals is called Sigma Bond.
OR
A bond which is formed by head to head overlap of atomic orbitals is called Sigma Bond.

Explanation
In the formation of a sigma bond the atomic orbital lies on the same axis and the overlapping of these orbital is maximum therefore, all such bonds, in which regions of highest density around the bond axis are termed as sigma bond.

Types of Overlapping in Sigma Bond
There are three types of overlapping in the formation of sigma bond.
1. s-s orbitals overlapping
2. s-p orbitals overlapping
3. p-p orbitals overlapping
In all the three types, when the two atomic orbitals are overlapped with each other two molecular orbitals are formed. In these two molecular orbitals the energy of one orbital is greater than the the atomic orbitals which is known as sigma antibonding orbital while the energy of the other orbital is less than the atomic orbital this orbital of lower energy is called sigma bonding orbital and the shared electron are always present in the sigma bonding orbitals.

1. s-s Orbitals Overlapping
In order to explain s-s overlapping consider the example of H2 molecule. In this molecule is orbital of one hydrogen overlaps with is orbital of other hydrogen to form sigma bonding orbitals. Due to this bonding a single covalent bond is formed between the two hydrogen atoms.
Diagram Coming Soon

2. s-p Orbitals Overlapping
This type of overlapping takes place in H-Cl molecule. 1s orbital of hydrogen overlaps with 1p orbital of chlorine to form a single covalent bond. In this overlapping two molecular orbitals are formed, one of the lower energy while the other orbital is of higher energy. The shapes of these orbitals are as follows.
Diagram Coming Soon

3. p-p Orbitals Overlapping
This type of overlapping takes place in fluorine molecule. In this mole 1p orbital of a fluorine atom is overlapped with 1p orbital of the other fluorine atom. The molecular orbitals formed in this overlapping are given in figure
Diagram Coming Soon

PI Bond
When the two atomic orbital involved in a covalent bond are parallel to each other then the bond formed between them is called pi bond.
In this overlapping, two molecular orbitals are also formed. The lower energy molecular orbitals is called p bonding orbital while the higher energy molecular orbital is called p antibonding orbital. The shape of these molecular orbitals are as follows.
Diagram Coming Soon

Hybridization
Definition
The process in which atomic orbitals of different energy and shape are mixed together to form new set of equivalent orbitals of the same energy and same shape.
There are many different types of orbital hybridization but we will discuss here only three main types.

1. sp3 Hybridization
The mixing of one s and three p orbitals to form four equivalent sp3 hybrid orbitals is called sp3 hybridization. These sp3 orbitals are directed from the center of a regular tetrahedron to its four corners. The angles between tetrahedrally arranged orbitals are 109.5º.
It has two partially filled 2p orbitals which indicate that it is divalent, but carbon behaves as tetravalent in most of its compounds. It is only possible if one electron from 2s orbital is promoted to an empty 2pz orbital to get four equivalent sp3 hybridized orbitals.
Diagram Coming Soon
The four sp3 hybrid orbitals of the carbon atom overlap with 1s orbitals of four hydrogen atoms to form a methane CH4 molecule.
The methane molecule contains four sigma bonds and each H-C-H bond angle is 109.5º.

2. sp2 Hybridization
The mixing of one s and two p orbitals to form three orbitals of equal energy is called sp2 or 3sp2 hybridization. Each sp2 orbital consists of s and p in the ratio of 1:2. These three orbitals are co-planar and at 120º angle as shown
Diagram Coming Soon
A typical example of this type of hybridization is of ethane molecule. In ethylene, two sp2 hybrid orbitals of each carbon atom share and overlap with 1s orbitals of two hydrogen atoms to form two s bonds. While the remaining sp2 orbital on each carbon atom overlaps to form a s bond. The remaining two unhybridized p orbitals (one of each) are parallel and perpendicular to the axis joining the two carbon nuclei. These generates a parallel overlap and results in the formation of 2 p orbitals. Thus a molecule of ethylene contain five s bonds and one p bond.
Diagram Coming Soon

3. sp Hybridization
When one s and one p orbitals combine to give two hybrid orbitals the process is called sp hybridization. The sp hybrid orbitals has two lobes, one with greater extension in shape than the other and the lobes are at an angle of 180º from each other. It means that the axis of the two orbitals form a single straight line as shown.
Now consider the formation of acetylene molecule HC = CH. The two C-H s bonds are formed due to sp-s overlap and a triple bond between two carbon atoms consist of a s bond and two p bond. The sigma bond is due to sp-sp overlap whereas p bonds are formed as a result of parallel overlap between the unhybridized four 2p orbitals of the two carbon.
Diagram Coming Soon

Valence Shell Electron Pair Repulsion Theory
The covalent bonds are directed in space to give definite shapes to the molecules. The electrons pairs forming the bonds are distributed in space around the central atom along definite directions. The shared electron pairs as well as the lone pair of electrons are responsible for the shape of molecules.
Sidwick and Powell in 1940 pointed out that the shapes of the molecules could be explained on the basis of electron pairs present in the outermost shell of the central atom. Pairs of electrons around the central atom are arranged in space in such a way so that the distances between them are maximum and coulombie repulsion of electronic cloud are minimized.
The known geometries of many molecules based upon measurement of bond angles shows that lone pairs of electrons occupy more space than bonding pairs. The repulsion between electronic pairs in valence shell, decreases in the following order.
Lone Pair - Lone Pair > Lone Pair - Bond Pair > Bond Pair - Bond Pair
When we apply this theory we can see the variation of angle in the molecular structures.
Consider the molecular structures of NH3, OH & H2O.
Diagram Coming Soon
Variation from ideal bond angles are caused by multiple covalent bonds and lone electron pairs both of which require more space than single covalent bonds and therefore cause compression of surrounding bond angles.
Thus the number of pairs of electrons in the valency shell determine the overall molecular shape.

Structure of BeCl2
The two bond pairs of electrons in BeCl2 arrange themselves as far apart as possible in order to minimize the repulsion between them.

Structure of BF3 OR BCl3
In this molecule three bond pair are present around boron to arrange themselves as far apart as possible a trigonal structure is formed.

Hydrogen Bond
When hydrogen is bonded with a highly electronegative element such as nitrogen oxygen, fluorine, the molecule will be polarized and a dipole is produced. The slightly positive hydrogen atom is attracted by the slightly negatively charged electronegative atom. An electrostatic attraction between the neighbouring molecules is set up when the positive pole of one molecule attracts the negative pole of the neighbouring molecule. This type of attractive force which involves hydrogen is known as hydrogen bonding.

1styear CHEMISTRY Notes Chapter-3

Chapter-3
ATOMIC STRUCTURE.

Atomic Structure
Introduction
About the structure of atom a theory was put on by John Dalton in 1808. According to this theory matter was made from small indivisible particles called atoms.
But after several experiments many particles have been discovered with in the atom which are electrons, protons, neutrons, positrons etc. For the discovery of these fundamental particles the experiments are as follows.
1. Faraday's experiment indicates the existence of electron.
2. Crook's tube experiment explains the discovery of electron and proton.
3. Radioactivity also confirms the presence of electrons and protons.
4. Chadwick's experiment shows the presence of neutrons.
The details of these experiments are given below.

Faraday's Experiment


Passage of Electricity Through Solution
In this experiment Faraday passed the electricity through an electrolytic solution. He observed that when two metal plates called electrodes are placed in an electrolytic solution and electricity is passed through his solution the ions present in the solution are moves towards their respective electrodes. In other words these ions are moves towards the oppositely charge electrodes to give up their charge and liberated as a neutral particles.
Faraday also determined the charges of different ions and the amount of elements liberated from the electrolytic solution. Due to this experiment presence of charge particles in the structure of atoms is discovered. The basic unit of electric charge was later named as electron by Stoney in 1891.

Crook's Tube Or Discharge Tube Experiment
Passage of Electricity Through Gases Under Low Pressure
Introduction
The first of the subatomic particles to be discovered was electron. The knowledge about the electron was derived as a result of the study of the electric discharge in the discharge tube by J.J. Thomson in 1896. This work was later extended by W. Crooke

Working of Discharge Tube
When a very high voltage about 10,000 volts is applied between the two electrodes, no electric discharge occurs until the part of the air has been pumped out of the tube. When the pressure of the gas inside the tube is less than 1 mm, a dark space appears near the cathode and thread like lines are observed in the rest of 0.01 mm Hg it fills the whole tube. The electric discharge passes between the electrodes and the residual gas in the tube begins to glow. These rays which proceed from the cathode and move away from it at right angle in straight lines are called cathode rays.

Properties of Cathode Rays
1. They travel in straight lines away from the cathode and produce shadow of the object placed in their path.
2. The rays carry a negative charge.
3. These rays can also be easily deflected by an electrostatic field.
4. The rays can exert mechanical pressure showing that these consist of material particle which are moving with kinetic energy.
5. The produce fluorescence when they strike the glass wall of the discharge tube.
6. Cathode rays produce x-rays when they strike a metallic plate.
7. These rays consists of material particle whose e/m resembles with electron.
8. These rays emerge normally from the cathode and can be focused by using a concave cathode.

Positive Rays
In 1890 Goldstein used a discharge tube with a hole in the cathode. He observed that while cathode rays were emitting away from the cathode, there were coloured rays produced simultaneously which passed through the perforated cathode and caused a glow on the wall opposite to the anode. Thomson studied these rays and showed that they consisted of particles carrying a positive charge. He called them positive rays.

Properties of Positive Rays
1. These rays travel in a straight line in a direction opposite to the cathode.
2. These are deflected by electric as well as magnetic field in the way indicating that they are positively charged.
3. The charge to mass ratio (e/m) of positive particles varies with the nature of the gas placed in the discharge tube.
4. Positive rays are produced from the ionization of gas and not from anode electrode.
5. Positive rays are deflected in electric field. This deflection shows that these are positively charged so these are named as protons.

The Information Obtained From Discharge Tube Experiment
The negatively charge particles electrons and the positively charge particles protons are the fundamental particle of every atom.

Radioactivity
In 1895, Henry Becqueral observed that uranium and its compounds spontaneously emitted certain type of radiation which affected a photographic plate in the dark and were able to penetrate solid matter. He called these rays as radioactivity rays and a substance which possessed the property of emitting these radioactivity rays was said to be radioactivity element and the phenomenon was called radioactivity.
On further investigation by Maric Curic, it was found that the radiation emitted from the element uranium as well as its salts is independent of temperature and the source of the mineral but depend upon the mineral but depend upon the quantity of uranium present e.g. Pitchblende U3O8 was found to be about four times more radioactive than uranium.

Radioactive Rays
Soon after the discovery of radium it was suspected that the rays given out by radium and other radioactive substance were not of one kind. Rutherford in 1902 devised an ingenious method for separating these rays from each other by passing them between two oppositely charged plate. It was observed that the radioactive rays were of three kinds, the one bending towards the negative plate obviously carrying positive charge were called a-rays and those deflected to the positive plate and carrying -ve charge were named as ß-rays. The third type gamma rays, pass unaffected and carry no charge.

Properties of a - RAYS
1. These rays consists of positively charged particles.
2. These particles are fast moving helium nuclei.
3. The velocity of a-particles is approximately equal to 1/10th of the velocity of light.
4. Being relatively large in size, the penetrating power of a-rays is very low.
5. They ionize air and their ionization power is high.

Properties of ß - RAYS
1. These rays consists of negatively charged particles.
2. These particles are fast moving electron.
3. The velocity of ß-particles is approximately equal to the velocity of light.
4. The penetrating power of ß-rays is much greater than a-rays.
5. These rays ionizes gases to lesser extent.

Properties of Y - RAYS
1. Gamma rays do not consist of particles. These are electromagnetic radiations.
2. They carry no charge so they are not deflected by electric or magnetic field.
3. Their speed is equal to that of light.
4. These are weak ionizer of gases.
5. Due to high speed and non-material nature they have great power of penetration.

Chadwick Experiment (Discovery of Neutron)
When a light element is bombarded by a-particles, these a-particles leaves the nucleus in an unstable disturbed state which on settling down to stable condition sends out radioactivity rays. The phenomenon is known as "Artificial Radioactivity".
In 1933, Chadwick identified a new particle obtained from the bombardment of beryllium by a-particles. It had a unit mass and carried no charge. It was named "Neutron".

Spectroscopic Experiment
After the discovery of fundamental particles which are electrons, protons & neutron, the next question concerned with electronic structure of atom.
The electronic structure of the atom was explained by the spectroscopic studies. In this connection Plank's Quantum theory has great impact on the development of the theory of structure of atom.

Planck's Quantum Theory
In 1900, Max Planck studied the spectral lines obtained from hot body radiations at different temperatures. According to him,
When atoms or molecules absorb or emit radiant energy, they do so in separate units of waves called Quanta or Photons.
Thus light radiations obtained from excited atoms consists of a stream of photons and not continuous waves.
The energy E of a quantum or photon is given by the relation
E = h v
Where v is the frequency of the emitted radiation and h the Planck's constant. The value of h = 6.62 x 10(-27) erg. sec.
The main point of this theory is that the amount of energy gained or lost is quantized which means that energy change occurs in small packets or multiple of those packets, hv, 2 hv, 3 hv and so on.

Spectra
A spectrum is an energy of waves or particles spread out according to the increasing or decreasing of some property. E.g. when a beam of light is allowed to pass through a prism it splits into seven colours. This phenomenon is called dispersion and the band of colours is called spectrum. This spectrum is also known as emission spectrum. Emission spectra are of two types.
1. Continuous Spectrum
2. Line Spectrum

1. Continuous Spectrum
When a beam of white light is passed through a prism, different wave lengths are refracted through different angles. When received on a screen these form a continuous series of colours bands: violet, indigo, blue, green, yellow and red (VIBGYOR). The colours of this spectrum are so mixed up that there is no line of demarcation between different colours. This series of bands that form a continuous rainbow of colours is called continuous spectrum.

2. Line Spectrum
When light emitted from a gas source passes through a prism a different kind of spectrum may be obtained.
If the emitted from the discharge tube is allowed to pass through a prism some discrete sharp lines on a completely dark back ground are obtained. Such spectrum is known as line spectrum. In this spectrum each line corresponds to a definite wave length.

Identification of Element By Spectrum
Each element produces a characteristics set of lines, so line spectra came to serve as "finger prints" for the identification of element. It is possible because same element always emit the same wave length of radiation. Under normal condition only certain wave lengths are emitted by an element.

Rutherford's Atomic Model
Evidence for Nucleus and Arrangement of Particles
Having known that atom contain electrons and a positive ion, Rutherford and Marsden performed their historic "Alpha particle scattering experiment" in 1909 to know how and where these fundamental particles were located in the structure of atom.
Rutherford took a thin of gold with thickness 0.0004 cm and bombarded in with a-particles. He observed that most of the a-particles passed straight through the gold foil and thus produced a flash on the screen behind it. This indicated that old atoms had a structure with plenty of empty space but some flashes were also seen on portion of the screen. This showed that gold atoms deflected or scattered a-particles through large angles so much so that some of these bounced back to the source.
Based on these observations Rutherford proposed a model of the atom which is known as Rutherford's atomic model.

Assumption Drawn From the Model
1. Atom has a tiny dense central core or the nucleus which contains practically the entire mass of the atom leaving the rest of the atom almost empty.
2. The entire positive charge of the atom is located on the nucleus. While electrons were distributed in vacant space around it.
3. The electrons were moving in orbits or closed circular paths around the nucleus like planets around the sun.
4. The greater part of the atomic volume comprises of empty space in which electrons revolve and spin.

Weakness of Rutherford Atomic Model
According to the classical electromagnetic theory if a charged particle accelerate around an oppositely charge particle it will radiate energy. If an electron radiates energy, its speed will decrease and it will go into spiral motion finally falling into the nucleus. Similarly if an electron moving through orbitals of ever decreasing radii would give rise to radiations of all possible frequencies. In other words it would given rise to a continuous spectrum. In actual practise, atom gives discontinuous spectrum.

X-Rays and Atomic Number
In 1895, W.Roentgen discovered that when high energy electrons from cathode collide with the anode in the Crook's tube, very penetrating rays are produced. These rays were named as X-rays.

Explanation
When an electron coming from the cathode strike with the anode in the crook's tube, it can remove an electron from the inner shell of the atom. Due to removal of t his electron the electronic configuration of this ion is unstable and an electron from an orbital of higher energy drops into the inner orbital by emitting energy in form of a photon. This photon corresponds to electromagnetic radiations in the x-rays region.

Relationship Between Wave Length and Nuclear Charge
In 1911, Mosley stablished a relationship between the wave length and nuclear charge. He found that when cathode rays struck elements used as anode targets in the discharge tube, characteristic x-rays were emitted. The wave length of the x-rays emitted decreases regularly with the increase of atomic mass. On careful examination of his data Mosely found that the number of positive charges on the nucleus increases from atom to atom by single electronic unit. He called the number of positive charges as the atomic number.
Diagram Coming Soon

Bohr's Theory
Rutherford's model of atom fails to explain the stability of atom and appearance of the line spectra. Bohr in 1913 was the first to present a simple model of the atom which explained the appearance of line spectra.
Some of the postulates of Bohr's theory are given below.
1. An atom has a number of stable orbits or stationary states in which an electron can reside without emission or absorption of energy.
2. An electron may pass from one of these non-radiating states to another of lower energy with the emission of radiations whose energy equals the energy difference between the initial and final states.
3. In any of these states the electrons move in a circular path about the nucleus.
4. The motion of the electron in these states is governed by the ordinary laws of mechanics and electrostatic provided its angular momentum is an integral multiple of h/2p
It can be written as
mvr = nh / 2p
Here mvr becomes the angular momentum of the electron. Thus Bohr's first condition defining the stationary states could be stated as
"Only those orbits were possible in which the angular momentum of the electrons would be an integral multiple of h/2p". These stationary states correspond to energy levels in the atom.

Calculation of Radius of Orbits
Consider an electrons of charge e revolving.
Atomic number and e the charge on a proton.
Let m be the mass of the electro, r the radius of the orbit and v the tangential velocity of the revolving electron.
The electrostatic force of attraction between the nucleus and the electron according to Coulomb's law
= Z e x e / r2
Diagram Coming Soon
The centrifugal force acting on the electron.
= mv2 / r
Bohr assumed that these two opposing forces must be balanced each other exactly to keep the electron in an orbit.
Therefore
Ze2 / r2 = m v2 / r
Multiply both sides by r
r x Ze2 / r2 = r x m v2 / r
Ze2 / r = m v2
OR
r = Ze2 / m v2 .................. (1)
The Bohr's postulate states that only those orbits are possible in which
mvr = nh / 2p
Therefore,
V = nh / 2pmr
Substituting the value of V in eq (1)
r = Ze2 / m(nh/2pmr)2
or
r = Ze2 x 4p2 mr2/n2h2
or
1/r = 4p2mZe2/n2h2
cr
r = n2h2 / 4p2mZe2 ............... (2)
This equation gives the radii of all the possible stationary states. The values of constants present in this equation are as follows.
H = 6.625 x 10(-27) ergs sec OR 6.625 x 10(-37) J.s
Me = 9.11 x 10(-28) gm OR 9.11 x 10(-31) kg
E = 4.802 x 10(-10) e.s.u OR 1.601 x 10(-19) C
By substituting these values we get for first shell of H atom
r = 0.529 x 10(-8) m OR 0.529
The above equation may also be written as
r = n2 (h2 / 4p2mZe2) x n2 a0 .................... (3)
For the first orbit n = 1 and r = 0.529. This is the value of the terms in the brackets sometimes written as a0 called Bohr's Radius. For the second shell n = 2 and for 3rd orbit n = 3 and so on.

Hydrogen Atom Spectrum
Balmer Series
The simplest element is hydrogen which contain only one electron in its valence shell.
Balmer in 1885 studied the spectrum of hydrogen. For this purpose he used hydrogen gas in the discharge tube. Balmer observed that hydrogen atom spectrum consisted of a series of lines called Balmer Series. Balmer determined the wave number of each of the lines in the series and found that the series could be derived by a simple formula.

Lyman Series
Lyman series is obtained when the electron returns to the ground state i.e. n = 1 from higher energy level n(2) = 2, 3, 4, 5, etc. This series of lines belongs to the ultraviolet region of spectrum.

Paschen Series
Paschen series is obtained when the electron returns to the 3rd shell i.e. n = 3 from the higher energy levels n2 = 4, 5, 6 etc. This series belongs to infrared region.

Bracket Series
This series is obtained when an electron jumps from higher energy levels to 4th energy level.

Heisenberg Uncertainty Principle
According to Bohr's theory an electron was considered to be a particle but electron also behaves as a wave according to be Broglie.
Due to this dual nature of electron in 1925 Heisenberg gave a principle known as Heisenberg Uncertainty Principle which is stated as,
It is impossible to calculate the position and momentum of a moving electron simultaneously.
It means that if one was known exactly it would be impossible to known the other exactly. Therefore if the uncertainty in the determination of momentum is ?px and the uncertainty in position is ?x then according to this principle the product of these two uncertainties may written as
?px . ?x ˜ h
So if one of these uncertainties is known exactly then the uncertainty in its determination is zero and the other uncertainty will become infinite which is according to the principle.

Energy Levels and Sub-Levels
According to Bohr's atomic theory, electrons are revolving around the nucleus in circular orbits which are present at definite distance from the nucleus. These orbits are associated with definite energy of the electron increasing outwards from the nucleus, so these orbits are referred as Energy Levels or Shells.
These shells or energy levels are designated as 1, 2, 3, 4 etc K, L, M, N etc.
The spectral lines which correspond to the transition of an electron from one energy level to another consists of several separate close lying lines as doublets, triplets and so on. It indicates that some of the electrons of the given energy level have different energies or the electrons belonging to same energy level may differ in their energy. So the energy levels are accordingly divided into sub energy levels which are denoted by letters s, p, f (sharp, principle, diffuse & fundamental).
The number of sub levels in a given energy level or shell is equal to its value of n.
e.g. in third shell where n = 3 three sub levels s, p, d are possible.

Quantum Numbers
There are four quantum numbers which describe the electron in an atom.

1. Principle Quantum Number
It is represented by "n" which describe the size of orbital or energy level.
The energy level K, L, M, N, O etc correspond to n = 1, 2, 3, 4, 5 etc.
If
n = 1 the electron is in K shell
n = 2 the electron is in L shell
n = 3 the electron is in M shell

2. Azimuthal Quantum Number

This quantum number is represented by "l" which describes the shape of the orbit. The value of Azimuthal Quantum number may be calculated by a relation.
l = 0 ----> n - 1
So for different shell the value of l are as
n = 1 K Shell l = 0
n = 2 L Shell l = 0, 1
n = 3 M Shell l = 0, 1, 2
n = 4 N Shell l = 0, 1, 2, 3
when l = 0 the orbit is s
when l = 1 the orbit is p
when l = 2 the orbit is d
when l = 3 the orbit is f

3. Magnetic Quantum Number
It is represented by "m" and explains the magnetic properties of an electron. The value of m depends upon the value of l. It is given by
m = + l ----> 0 ----> l
when l = 1, m has three values (+1, 0, -1) which corresponds to p orbital. Similarly when l = 2, m has five values which corresponds to d orbital.

4. Spin Quantum Number
It is represented by "s" which represents spin of a moving electron. This spin may be either clockwise or anticlockwise so the values for s may be +1/2 or -1/2.

Pauli's Exclusion Principle
According to this principle
No two electrons in the same atom can have the same four quantum number.
Consider an electron is present in 1s orbital. For this electron n = 1, l = 0, m = 0. Suppose the spin of this electron is s = +1/2 which will be indicated by an upward arrow ?. Now if another electron is put in the same orbital (1s) for that electron n = 1, l = 0, m = 0. It can occupy this orbital only if the direction of its spin is opposite to that of the first electron so s = -1/2 which is symbolized by downward arrow ?. From this example, we can observe the application of Pauli's exclusion principle on the electronic structure of atom.

Electronic Configuration
The distribution of electrons in the available orbitals is proceeded according to these rules.
1. Pauli Exclusion Principle
2. Aufbau Principle
3. (n + l) Rule
4. Hund's Rule
The detail of these rules and principles is given below.

1. Aufbau Principle
It is states as
The orbitals are filled up with electrons in the increasing order of their energy.
It means that the orbitals are fulled with the electrons according to their energy level. The orbitals of minimum energy are filled up first and after it the orbitals of higher energy are filled.

2. Hund's Rule
If orbitals of equal energy are provided to electron then electron will go to different orbitals and having their parallel spin.
In other words we can say that electrons are distributed among the orbitals of a sub shell in such a way as to give the maximum number of unpaired electrons and have the same direction of spin.

3. (n + l) Rule
According to this rule
The orbital 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.
For the electronic configuration the order of the orbital is as follows.
1s, 2s, 2p, 3s, 4s, 3d, 4p, 5s, 4d, 5p, 6s etc.

Atomic Radius
For homonuclear diatomic molecules the atomic radius may be defined as
The half of the distance between the two nuclei present in a homonuclear diatomic molecules is called atomic radius.
It may be shown as
In case of hetronuclear molecular like AB, the bond length is calculated which is (rA + rB) and if radii of any one is known the other can be calculated.
For the elements present in periodic table the atomic radius decreases from left to right due to the more attraction on the valence shell but it increases down the group with the increase of number of shells.

Ionic Radius
Ionic radius is defined as
The distance between nucleus of an ion and the point up to which nucleus has influence of its electron cloud.
When an electron is removed from a neutral atom the atom is left with an excess of positive charge called a cation e.g
Na ----> Na+ + c-
But when an electron is added in a neutral atom a negative ion or anion is formed.
Cl + e- ----> Cl-
As the atomic radius, the ionic radii are known from x-ray analysis. The value of ionic radius depends upon the ions that surround it.
Ionic radii of cations have smaller radii than the neutral atom because when an electron is removed. The effective charge on the nucleus increases and pulls the remaining electrons with a greater force.
Ionic radii of anions have a large radii than the neutral atom because an excess of negative charge results in greater electron repulsion.
Radius of Na atom = 1.57
Radius of Na+ atom = 0.95 (smaller than neutral atom)
Radius of Cl atom = 0.99
Radius of Cl- atom = 1.81 (larger than neutral atom)

Ionization Potential
Definition
The amount of energy required to remove most loosely bounded electron from the outermost shell of an atom in its gaseous state is called is called ionization potential energy.
It is represented as
M(gas) ----> M+(gas) + e- ................... ?E = I.P
The energy required to remove first electron is called first I.P. The energy required to remove 2nd or 3rd electron is called 2nd I.P or 3rd I.P
M(gas) ----> M+(gas) + e- ................... ?E = 1st I.P
M+(gas) ----> M++(gas) + e- ................?E = 2nd I.P
M++(gas) ----> M+++(gas) + e- ............ ?E = 3rd I.P
The units of I.P is kilo-Joule per mole.

Factors on which I.P Depends
1. Size of the Atom
If the size of an atom is bigger the I.P of the atom is low, but if the size of the atom is small then the I.P will be high, due to fact if we move down the group in the periodic table. The I.P value decreases down the group.

2. Magnitude of Nuclear Charge
If the nuclear charge of atom is greater than the force of attraction on the valence electron is also greater so the I.P value for the atom is high therefore as we move from left to right in the periodic table the I.P is increased.

3. Screening Effect

The shell present between the nucleus and valence electrons also decreases the force of attraction due to which I.P will be low for such elements.

Electron Affinity
Definition
The amount of energy liberated by an atom when an electron is added in it is called electron affinity.
It shows that this process is an exothermic change which is represented as
Cl + e- ----> Cl- ............ ?H = -348 kJ / mole
Factors on which Electron Affinity Depends

1. Size of the Atom
If the size of atom is small, the force of attraction from the nucleus on the valence electron will be high and hence the E.A for the element will also be high but if the size of the atoms is larger the E.A for these atoms will be low.

2. Magnitude of the Nuclear Charge
Due to greater nuclear charge the force of attraction on the added electron is greater so the E.A of the atom is also high.

3. Electronic Configuration
The atoms with the stable configuration has no tendency to gain an electron so the E.A of such elements is zero. The stable configuration may exist in the following cases.
1. Inert gas configuration
2. Fully filled orbital
3. Half filled orbital

Electronegativity
Definition
The force of attraction by which an atom attract a shared pair of electrons is called electronegativity.

Application of Electronegativity
1. Nature of Chemical Bond
If the difference of electronegativity between the two combining atoms is more than 1.7 eV, the nature of the bond between these atoms is ionic but if the difference of electronegativity is less than 1.7 eV then the bond will be covalent.

2. Metallic Character
If an element possesses high electronegativity value then this element is a non-metal but if an element exist with less electronegativity, it will be a metal.

Factors for Electronegativity
1. Size of the Atom
If the size of the atom is greater the electronegativity of the atom is low due to the large distance between the nucleus and valence electron.

2. Number of Valence Electrons
If the electrons present in the valence shell are greater in number, the electronegativity of the element is high.

1styear CHEMISTRY Notes Chapter-2

Chapter-2
THREE STATES OF MATTER

Three States Of Matter
Matter
It is defined as any thing which has mass and occupies space is called matter.
Matter is composed of small and tiny particles called Atoms or molecules. It exist in three different states which are gaseous, liquid & solid.

Properties of Gas

1. It has no definite shape.
2. It has no definite volume, so it can be compressed or expanded.
3. A gas may diffuse with the other gas.
4. The molecules of a gas are in continuous motion.

Properties of Liquids

1. A liquid has no definite shape.
2. It has a fixed volume.
3. The diffusion of a liquid into the other liquid is possible if both of the liquids are polar or non-polar.
4. It can be compressed to a negligible.

Properties of Solids

1. A solid has a definite shape.
2. It has a fixed volume.
3. The rate of diffusion of solid with each other is very slow.
4. It cannot be compressed easily.