# CHAPTER # 3 ATOMIC STRUCTURE

**CHAPTER # 3**

ATOMIC STRUCTURE

ATOMIC STRUCTURE

Introduction

The theory explains that matter was made from small indivisible particles called atoms. According to ancient era, this theory was put on a sound scientific basis by John Dalton in 1808.

Dalton’s Concept

In early 1800s, John Dalton proposed his theory that matter was composed of minute discrete particles called atoms (Greek = indivisible).

Recent Concept

Today, it is well established that atoms are complex organization of matter and energy. Many particles have been discovered within the atom. These sub atomic particles include electron, proton, neutron, several types of mesons, hyprons, etc.

Evidence for the presence of electrons, protons and neutrons in the atom is derived through the experiment.

__FARADAY’S EXPERIMENT__**Passage of Electricity through Solutions**

**Clue about Electrons**

**Conclusion**

There is some elementary unit of electric charge associated with these ions which can be calculated. The ions were observed to carry some integral multiple of this charge and the basic charge was later named by as an electron.

**Example**

Ionization of Electrolytic Solution

NaCl Current Na + Cl^{+}^{-}

__A__^{+}__ ____anode (oxidation) __(lose of particles)

2Cl^{-} Cl_{2} + 2e^{-}

At cathode (Reduction) (gain of particle)

2 Na^{+} + 2e^{-} 2 Na

**CROOK’S TUBE OR DISCHARGE TUBE EXPERIMENT**

**(Passage of electricity through gases under low pressure)**

**Introduction**

The conduction of electricity through gas was first studied by GEISSLER followed by W-CROOKES in 1879. Electron was the first sub-atomic particles of an atom which was discovered in an electrical discharge tube called Crook’s Tube by J.J.Thomson in 1896.

**Construction of the discharge tube**

The discharge tube consists of long glass tube with circular electrodes sealed into each end. A small opening is made on its one side which is connected to a vacuum pump. The two electrodes are connected to the positive and negative terminal of a very high voltage battery.

**Working of Discharge tube**

At normal pressure (760mm Hg), a very high voltage about 4000 volt is connected to the tube. A sharp noise is heard by reducing pressure with the help of vacuum pump. The noise disappears, the gas attains a violet colour, then pink and finally whole tube begins to glow.

When pressure is lowered to about 1m.m. Hg, a dark space appears which is separated from cathode by a bluish glow at about 0.1mm. The rest of the glow in the tube is broken into several alternate bright and dark patches called “striations”. At further low pressure, the blue glow separates from cathode and a dark space appears in between called Crook’s Dark Space. At further low pressure, the Crook’s and Faraday’s dark space appear growing in size & the positive column shrinks, till at 0.001 mm Hg. The whole tube is filled with crook’s dark space and radiation called “cathode rays” emit from cathode electrode.

**Conclusion**

**Discovery of Protons**

**Goldstein’s experiment**

**PROPERTIES OF CANAL RAYS**

**Conclusion**

**Natural Radioactivity**

(Confirmation of electron and proton)

**Definition:**The spontaneous and continuous emission of rays at constant rate by same elements is called Radioactivity. And, the elements are called Radioactive elements.

Examples of radioactivity elements are:

Uranium (U) Radium (Ra) Polonium (Po) Radon (Rn) Tritium (T)

History The phenomena was first observed by Henry Becqueral while working on a mineral pitch blinde. Most of rays emitted by this mineral consisted of electron.

Later, Pierrie Curie and Marie curie isolated Radium, which was highly radioactive.

**“Experiment for Detection of Radiation”**

Rutherford used simple method for the detection and separation of this radiation. He placed a piece of radioactive substance (e.g. Ra) in a lead pot. The rays emitted by the radioactive substance were allowed to pass through two oppositely charged plates. The rays bending towards the negative plate carry +ve charge and are called Alpha (α). Those deflected towards the positive plate carry a - ve charge and are named as Beta (β) rays. Those rays which remained undeflected carry no charge and are called Gamma ray ( γ )

Result of Radioactivity: An element after giving out radiation breaks down to a more stable element, e.g. U on emission of α - particles would be converted to Th.

^{238}U ^{234} Th + He^{4}

The emission of radiation would continue until the formation of lead is produced which is more stable.

**Conclusion**

The evidence of radioactivity shows that the atom is not an indivisible particle. If it can emit electrons and Helium nuclei, it must have a substructure of its own.

**PROPERTIES OF RADIATION**

**PROPERTRIES OF ALPHA ( α) RAYS**

i) Charge: They carry positive charge.

**PROPERTIES OF BETA PARTICLES**

**PROPERTIES OF GAMMA RAYS**

**CHAD WICK EXPERIMENT (Discovery of Neutrons)**

_{4}Be^{9}+_{2}He^{4 }_{6}C^{12}+_{0}n^{1 }

With the help of artificial radioactivity, the
particle “Neutron” was discovered.

**PLANK’S QUANTUM THEORY**

**Postulate**

**Conclusion**

**Spectroscopy**

**Types of spectra**

**Emission Spectrum**

**Continuous Spectrum**

**Line Spectrum**

**Rutherford’s atomic model**

**Ruther ford’s Experiment**

**Assumption Drawn from this model**

**Conclusion**

**Defects of Rutherford Model (Weakness of Rutherford’s Theory)**

**1) Classical Mechanics**

**2) Spectroscopic Theory**

**BOHR’S ATOMIC THEORY**

**Assumptions of Bohr’s theory**

If E_{2} = Energy of higher state.

E_{1}= Energy of lower
state.

Then ∆E = E_{2}- E_{1}

∆E = H υ

Where “h” is called plank’s constant. Each
transition of electron leaves a spectrum.
4. Stationary states are only those orbits for which the product of
momentum ( mν) and circumefrence ( 2π*r *) is equal to plank’s
constant (h) or integral multiple of (h) this product is also called “ THE ACTION” Mathematically.

m v x 2 π r = nh

m v r = nh / 2π

(mvr is called Angular Momentum )

According to equation,

Only those stationary states or orbits are possible for which the angular momentum ( mvr ) is an integral multiple of h/2π .

“ n “ is called principal Quantum Number . Its values are n =
1,2,3,4,----

It is equal to number of orbit .

**Expression for the Radius of BOHR’S ATOM **

( Radius of nth orbit of an atom)

If we know the atomic number of an atom and orbit number (n), we can find the radius of nth orbit as follows:

Derivation:

Let Z = atomic No

N = orbit No

R = radius

E- charge of electron in nth orbit.

E+ = charge of proton in nucleus

Ze+ = total charge on nucleus

According to columb’s Law.

The electrostatic force between e- and nucleus will be

Fe = __k ze+ . e__^{-}__ __(Centripetal force which is denoted by Fe) r^{2}

The values of e+ and e’
are same = 1.6 x 10^{-19} columb.

Fe = k.
Ze^{2}

-------------

r^{2}

There is another force
in the atom which keeps the electron in their orbits, this is called
centrifugal force Which is denoted by
Fc

Fc =__ mν__^{2}

r

For equation: Fe = Fc

__KZe __²
= __mν__²

r² r

Kze² = __mv²r__^{2}

r

Kze^{2} = mv^{2}r

r = __k
Ze__^{2}__ __

mv^{2 }

r = __K. Ze__^{2 }. __1 __1

m v^{2 }

Now, we will find the velocity v. According to
second postulate of Bohr:

mvr = __n
h__

2π

Squaring both the sides

v^{2 }= __n__^{2}__ ____h__^{2}

4π^{2} m^{2} r^{2}

__1 __= __4____π__^{2 }__m__^{2 }__v__^{2}

V^{2 }n^{2 }h^{2 }

Putting the value of __1 _ __in equation (1),
we get

v^{2 }

^{ }r =__k Ze__^{2 }_{.}^{ }__4 ____π__^{2}__ ____m r__^{2}

m n2 h^{2}

1/r = __KZe__^{2 }. __4 π__^{2}__mr__^{2}

m n^{2}h^{2 }

^{ }1/r = __KZe__^{2}__. 4 π__^{2}__m__

n^{2}h^{2 }

r = __n__^{2}__h__^{2}__ __(PROVED)

4 π^{2}mKZe^{2 }

__Radius of Bohr’s atom for nth orbit __** **

__Radius of 1__^{st}__ ____Orbit of Hydrogen atom__** **

(r_{1 or }x =1 )

For Hydrogen Atomic Number = x_{o} = 1

For 1^{st} Orbit n = 1

Putting in equation 2, we get

r_{1} = a_{o} = h^{2} / 4 π m e^{2}

Putting the value of constant h, π, m, e, we
get

a_{o} = 0.529 A^{0}

This a_{0} is the value of radius of 1^{st} orbit of hydrogen atom,
and is called Bohr’s radius. Relation
between r and a_{0 }

r = n^{2} __h__^{2 }.

4 π^{2} m Ze^{2}

But for hydrogen Z =1 and h^{2} / 4π^{2} m Ze^{2} = a_{0}

r = n^{2} a_{0}

This expression can be
used to find radius of any orbit of hydrogen atom. For 1^{st} Orbit = n =1

2^{nd} Orbit = n = 2

3^{rd} Orbit = n = 3

“n” is number of orbit which is called Principal
Quantum Number.

**Calculate the radius of second orbit of hydrogen atom in Angstrom unit. Q) Bohr’s radius is**** 0.529 A**^{0}**. ****Calculate radius of 2nd and 3rd Orbit of hydrogen atom. **

** DATA **

Bohr’s radius = a_{0} = 0.529 A^{0}

For 2^{nd} Orbit = n =2

3^{rd} Orbit = h = 3

**Solution**

Now

R = n^{2} a_{o}

For 2^{nd} Orbit

r = (2)^{2} x 0.529

r = 4 x
0.529

r = 2.11
A^{0}

For 3^{rd} Orbit

r = (3)^{2} x 0.529

r = 9 x
0.529

r = 4.76
A^{0}

**Derivation of energy of electron in an orbit**

KE = ½ mv^{2}

PE = - Ze^{2} / r

E_{n} = KE + PE

= ½ mv^{2} + ( -Ze^{2} / r)

= ½ mv^{2} - _{r}*Ze*^{2} --------------- eq
1

we know that

2 2

^{mv }~~=~~

*Ze *

*r *

*r *

2

mv^{2} = _{r}*Ze*^{2}

now subs the value of mv^{2} in eq (1)

E = _{2}^{1} mv^{2} = _{r}*Ze*^{2}

E = _{2}^{1}_{r}*Ze*^{2} = _{r}*Ze*^{2}

Substituting the value
of r = [ n^{2} h^{2} / 4 π^{2}m Ze^{2}] In equation ( 1 ) we get,

E = - __Ze__^{2}__ __= __-2 ____π__^{2}__m Z__^{2}__e__^{4}

2[n^{2}h^{2}/4π^{2}mZe^{2}] n^{2} h^{2}

**Expression for frequency and Wave number**

According to Bohr’s quantum theory when electron jumps from higher energy. Sate (E_{2}) to lower energy state (E_{1}) the energy difference is equal to the energy of photon emitted i.e.,

h ν = E_{2} – E_{1}

h ν = __-2 ____π__^{2}__m Z__^{2}__e__^{4}__ __– (__-2 ____π__^{2}__m Z__^{2}__e__^{4}__)__

n_{2}^{2} h^{2} n_{1}^{2}h^{2}

h ν = __-2 ____π__^{2}__m Z__^{2}__e__^{4}__ __+ __2 ____π__^{2}__m Z__^{2}__e__^{4}

n_{2}^{2} h^{2} n_{1}^{2}h^{2}

h ν = __2 ____π__^{2}__m Z__^{2}__e__^{4}__ __+ __2 ____π__^{2}__m Z__^{2}__e__^{4}__)__

n_{1}^{2} h^{2} n_{2}^{2}h^{2}

h ν = __2 ____π__^{2}__m Z__^{2}__e__^{4}__ __- __2 ____π__^{2}__m Z__^{2}__e__^{4}

n_{1}^{2} h^{2} n_{2}^{2}h^{2}

h ν = __2 ____π__^{2}__m Z__^{2}__e__^{4 }( 1/n_{1}^{2} –1/n_{2}^{2})

h^{2}

h ν = __2 ____π__^{2}__m e__^{4 }Z^{4} ( 1/n_{1}^{2} –1/n_{2}^{2})

h^{2}

ν = __2 ____π__^{2}__m e__^{4 }Z^{2} ( 1/n_{1}^{2} –1/n_{2}^{2})

h^{3}

This is the expression of frequency. Now ν = *v *. c

Where ν is wave number. It is no
of wave/cm.

*v *. c = __2 ____π__^{2}__m e__^{4 }Z^{2} ( 1/n_{1}^{2} –1/n_{2}^{2})

h^{3}

*v *= __2 ____π__^{2}__m e__^{4 }Z^{2}( 1/n_{1}^{2} –1/n_{2}^{2})

ch^{3}

~~No~~w for hydrogen z =1. on
substituting the vale of constant is

__2 ____π__^{2}__m e__^{4 }, it was found to be equal to 109678cm^{-1}. This value is denoted by R_{H} is called Rydberg
constant ch^{3}

now.

*v *= R_{H} ( 1/n_{1}^{2} –1/n_{2}^{2}) (PROVED)

**Hydrogen Atom Specturm (proved)**

**Balmer Series**

The simplest element is hydrogen which contains only one electron in its valence shell. Balmer studied the spectrum of hydrogen for this purpose in 1885. He used hydrogen gas in the discharge tube. Balmer observed that hydrogen atom spectrum consist of a series of lines called Balmer Series. Balmer determined the wave number of each of the line in the series and found that the series could be derived from a simple formula.

ν = R_{H} ( 1/n_{1}^{2} –1/n_{2}^{2})

where n_{1} = 2 and n_{2} = 3^{rd}, 4^{th}, 5^{th}, etc.

and R_{n} is a constant called Rydberg Constant for hydrogen

RH = 109678 cm^{-1}

**Lyman Series **

Lyman series is obtained
when the electron returns to the ground state, i.e. from higher energy level to
lower energy level.

Where n_{2} = 2,3,4,5,6, etc.

This series of lines
belongs to ultraviolet region of spectrum. It is represented by the following
equation: ν = R_{H} ( 1/n_{1}^{2} –1/n_{2}^{2})

**Pasahen Series **

Paschen series is obtained when the electron
returns to the 3^{rd} shell.

n = 3 from the higher energy levels.

n_{2} = 4, 5, 6…… etc.

This series belongs to
the infrared region. The equation for Pasahen series may be written as: ν = R_{H} ( 1/3^{2} –1/n_{2}^{2})

n_{2} = 4,5,6, etc.

**Bracket Series **

This series is obtained
when an electron jumps from a higher energy level to the 4^{th} energy level. The equation
for this series:

V = R_{H} ( 1/4^{2} –1/n_{2}^{2})

n_{2} = 5,6,7… etc.

**P Fund Series **

P-Fund series is
obtained when an electron jumps from an energy level to 5^{th} energy level. V = R_{H} ( 1/5^{2} –1/n_{2}^{2})

n = 6,7,8… etc.

Heisenberg’s Uncertainty Principle

Heisenberg’s Uncertainty Principle

According to Bohr’s theory, an electron was considered to be particle. However, electron also behaves like a wave according to de-Borglie. Due to this dual nature of electron, Heisenberg gave a principle known as Heisenberg uncertainty principle. It is stated as: It is impossible to calculate the position and momentum of a moving electron simultaneously. If one was known exactly, it would be impossible to know the other exactly. Therefore, if the uncertainty in the determination of momentum is px and the uncertainty is x, then according to this principle, the product of these two uncertainties may be written as:

So, if one of these quantities 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 Energy Levels **

According to Bohr’s atomic theory, electrons revolve 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. Hence, these orbits are referred to as “Energy Levels” and are designed as 1,2,3,4… etc., or L, M, N, etc.

The spectral line which corresponds to the transition of an electron from one energy level to another consists of several separate close lines as double triplets and so on. It indicates that some of the electrons of the given energy level have different energies or the electron belonging to some energy level may differ in their energy. So energy levels are accordingly divided into sub energy levels and are denoted by letters s, p, d, f (sharp, principal, diffuse and fundamental).

The numbers of sublevels in a given energy level or shell is equal to the value of n, e.g. in third shell where n = 3, three sublevels (s, p and) are possible.

The maximum number of electrons in each shell can be calculated by using formula: Total number of electrons in shell n = 2n^{2}

Total numbers of electron in shell K = (n = 1)
are 2 (1)^{2} = 2

Total numbers of electron in shell L = (n = 2)
are 2 (2)^{2} = 8

Total numbers of electron in shell M = (n = 3)
are 2 (3)^{2} = 18

Total numbers of electron in shell N = (n = 4)
are 2 (4)^{2} = 32

The number of sub energy
levels present in an energy level (orbit) are equal to its values of n. K Shell ( n = 1) has only one orbit “s”.

L Shell ( n = 2) has only one orbit “s, p”.

M Shell ( n = 3) has only one orbit “s, p,
d”.

N Shell ( n = 4) has only one orbit “s, p, d,
f”.

**Orbitals and Quantum Numbers**

By using mathematical methods known as wave mechanics, Schrodingre in 1926, calculated the probability of locating the electrons in a region of space about the nucleus. Thus, on the basis of wave mechanics, the path of electrons cannot be probably determined. However, we can say that there are the regions in which probability of finding an electron is maximum. Such regions are called orbitals. Each orbital in an atom is completely described by four quantums.

**Quantum numbers**

• Principal quantum number

• Azimuthal quantum number or subsidiary quantum number

• Spin quantum number

• Magnetic quantum number

**Principal Quantum Number
(n) **

Principal quantum number
specifies the size of the orbital. It is denoted by symbol n. Its value may be
1, 2, 3, 4… etc., for respective
orbitals. Therefore, the size (radius) of the orbital = 0.529 A.n^{2} (n = 1, 2, 3 …) and

energy of orbital n = En = __2 ____π__^{2}__m e__^{2 }__Z__^{2}

n^{2} h^{2}

The energy level K, L, M, N, O, etc.

Corresponds to n = 1, 2, 3, 4, 5, … etc.

If n = 1 the electron is in K = shell

n = 2 the electrons is in L = shell

n = 3 the electrons is in M = shell

**Azimuthal Quantum Number **

It cannot measure the angular momentum of electron, as value of “l” increases, the angular momentum also increases. It depends upon the energy of electron. It also indicates the shape of the orbital. If n is known, then the value “ι”of can be calculated using formula (1 = n – 1)

If n = 1 l = 0 ⮴ n - 1

l = 0 ⮴ 1 - 1

l = 0 ⮴ 0 = (s) circular (2, electrons)

If n = 2 l = 0 ⮴ n - 1

l = 0 ⮴ 2 – 1

l = 0 ⮴ 1 – (dumbbell)

l = 0 ⮴ 0 = (s, p) circular (8 electron)

If n = 3 l = 0 ⮴ n - 1

ι = 0 ⮴ 3 – 1

ι = 0 ⮴ 2

ι = 0,1,2 (s, p, d) circular (18 electron) style="color: black; font-family: "Times New Roman","serif"; mso-fareast-font-family: "Times New Roman";">The energy of sub shell are in the order of s>p>d>f.

**Magnetic Quantum Number (m)**

Due to angular momentum of an electron, magnetic field is developed. This magnetism can be measured by magnetic quantum number. It depends upon the value of ‘l’ and gives different orientations of an orbital in space in applied magnetic field. It can be calculated by the formula **For –p –orbital we can calculate the Orientation for p – orbital** = 1, where l = 1 then m = -1 ⮴ 0 ⮴ + 1

m = -1, 0, + 1

(px),
(py), (pz) (Orientation)

d-orbital = l=2 then

m = -2 ⮴ -1 ⮴ 0 ⮴ + 1⮴ +2

m = -2, -1, 0, 1, 2

pz^{2} , px^{2}y^{2} , pxy, pyz, pxz
(Orientation)

**Spin Quantum Number (s)**

It specifies the spin of electron in an orbital. Clockwise spin is denoted by (+1/2) and anticlockwise is denoted by (-1/2).

Example

Spin of two electrons of He atom in s – orbital
can be shown as above:

S = 1/2

__Pauli’s Exclusion Principle__**
**

__Introduction__** **

When rules of Quantum
mechanics are applied to an atom having maximum than one electron. This
rule puts a limit to the values of
quantum numbers of an electron.

**Definition **

This rule was given by
Wolf Gang Pauli in 1925. According to this rule, “In an atom, no two electrons
can have the same set of all four
quantum number.”

**Explanation **

According to this rule,
in an element, two atoms are not exactly alike. They must differ in some
respect. Thus, all maximum two electrons
can have only three quantum numbers are same. The fourth one is always different.

__Example__

Consider He (atomic number = 2)

Its configuration is 1s^{2}

Because both electrons are in the first orbit,
so for both n = 1

Because both electrons are in the s orbit, so
for both l= 0

Because both electrons are in l = 0, so for both
m = 0

Now, according to
Plank’s principle, maximum three quantum can be same; and hence, the fourth
quantum must be different.

So for one electron = S = +1/2

For other electron = S = -1/2

They should have opposite spin.

**Shapes of Orbitals**

The shape of s orbital is spherical so the probability of finding the electrons is maximum because the electrons are uniformly distributed around the nucleus. The p – orbitals are dum-bel shaped

These are oriented in space along the three mutually perpendicular axis (x, y, z) and are called px, py, pz. These are degenerated orbitals of equal energy. Each p-orbital has two lobes, one of which is labeled (+) and the other (-).

**Electronic Configuration**

The distribution of electron in the available orbitals is proceeded according to the following rules: • Pauli’s Exclusion Principle

• Aufbau Principle

• (n + 1) rule

The details of these rules and principles are
given below:

**Aufbau Principle**

It states that, “the orbitals are filled up with electrons in the increasing order of their energy” starting with the 1st orbital. OR

Hypothetically, the electronic configuration of the atoms can be constructed by placing the electrons in the lowest available orbitals until the total number of electrons added is equal to the atomic number Z.

1s^{2}

2s^{2} 2p^{6}

3s^{2} 3p^{6} 3d^{10}

4s^{2} 4p^{6} 4d^{10} 4f^{14}

5s^{2} 5p^{6} 5d^{10}

6s^{2} 6p^{6 }

7s^{2}

In order is :

**1s**^{2}** ****, 2s**^{2}**, 2p**^{6}**, 3s**^{2}**, 3p**^{6}**, 4s**^{2}**, 3d**^{10}**, 4p**^{6 }** **

**5s**^{2}**, 4d**^{10}**, 5p**^{6}**, 6s**^{2}**, 4f**^{14}**, 5d**^{10}**, 6s**^{2}**, 7s**^{2}** **

__n + 1 rule__** **

According to this rule,

The orbital with the lowest value of (n + 1) is filled first. However, when the two orbitals have the same value of (n + 1), the orbital with the lower value of n is filled first.

The electronic configuration according to (n + 1) rule is stated as:

**1s**^{2}** ****, 2s**^{2}**, 2p**^{6}**, 3s**^{2}**, 3p**^{6}**, 4s**^{2}**, 3d**^{10}**, 4p**^{6 }** **

**5s**^{2}**, 4d**^{10}**, 5p**^{6}**, 6s**^{2}**, 4f**^{14}**, 5d**^{10}**, 6s**^{2}**, 7s**^{2}** **

**Example: **

** **1s 2s 2p 3s 3p

n =1 n =
2 n = 2 n = 3 n = 3

l = 0 l =
0 l = 1 l = 0 l = 1

n+1=1 (n+1)=2 (n+1)=3 (n+1) = 3 (n+1) =4 Hund’s rule

When there are available
orbitals of equal energy of electrons, they would go in separate particles
having same spin; rather then go in the
same orbitals and have paired spin. In other words, the arrangement (1^{2}) is more stable than (1^{1}) provided that these two orbitals are of equal energy.

Px Py Pz Px Py Pz

**A B **

According to Hund’s rule structure (A) is possible.

or

In other words, we can say that electrons are distributed in the orbitals of a sub shell in such a way as to give the same direction of spin.

Example:

_{16}O^{8}

Z = 8 = 1s^{2}, 2s^{2}, 2px, 2py, 2pz

_{14 }N^{7} _{19} F^{9}

Z = 7 = 1s^{2}, 2s^{2}, 2px, 2py, 2pz Z = 9 = 1s^{2}, 2s^{2}, 2px, 2py, 2pz

**Atomic radius**

Definition

It is defined as half of the distance between two homo nuclear atoms in diatomic molecules; such as, H_{2}, O_{2}, N_{2}, etc.

__Range and units__** **

Atomic radii are expressed in:

• angstrom (A)

• Pi meter (pm)

• Nanometer (n.m.)

The range of atomic radii is 0.7A^{0} to 2.9A^{0}.

1A^{o} = 10^{-10}m = 100m = 0.1n.m

**Explanation**

In case of hetronuclear molecules; like AB, the bond length is calculated which is (rA + rB) and if radii of any one is known, the radii of the other can be calculated For the elements present in the periodic table, the atomic radius decreases from left to right due to more attraction on the valence shell electron, but it increases down the group with increase of number of shell.

**Ionic Radius**

Definition

Ionic radius is defined as the distance between the nucleus of an ion and the point up to which the nucleus has influence on its electron cloud.

**Explanation**

When an electron is removed from a neutral atom, the atom is left with an excess of positive charge called a cation.

Example

Na ⮴ Na^{+} + e^{-}

Mg ⮴ Mg^{+2} + 2e^{ }

But when an electron is
added in a neutral atom, a negative ion or anion is formed. Cl’ + e′ ⮴ Cl^{ }

O + 2e′ ⮴ O^{-2 }

**Cations have smaller
ionic radii than parent neutral atoms or K**^{+1}**
****is
smaller than K**^{0}**. **

This is because due to loss of electrons.

• The effective nuclear charge increases, pulling the remaining electrons more strongly.

**Example **

Li -e Li^{+ }

Li > Li^{+}

**3A**^{0}** ****1.2A**^{0}** **

Anions have larger ionic radii than neutral atom

This is due to gain of electrons. The repulsion between the valence electrons increases which results in the increase in the size of –ve ion as compared to parent atom. Thus,

F -e^{+} F^{- }

**1.3A**^{0}** ****2.7A**^{0}** ****F < F **^{-}** **

**Ionization Potential**

The energy required to completely remove the most loosely bounded electron from an atom in gaseous state is the energy of that atom.

or

The energy required to remove the most loosely held electrons from an isolated gaseous atom in its ground state is called ionization potential.

or

__Unit__** **

Its units are:

• Kj/mole

• Electron volt (ev)

Example:

To remove an electron of
H atom, we require 1312K.J mole^{-1} energy, so it is the I.P
of H-atom. H 1. P H^{+} + e^{-}

1312Kj/mole

**Factors affecting
I.P: **

It is effected by,

**Atomic Radius : **Greater the atomic radius, the less is the I.P values

I. P ∝ 1/_{Atomic radius }

**Nuclear Charge: **Greater the nuclear charge, greater the I.P values

I. P ∝ _{Nuclear Charge }

**Shielding effect: **Greater the shielding effect, less the value of I.P

Types of Ionization potential

Types of Ionization potential

I.P is of following three types:

1st Ionization potential

1st Ionization potential

The energy requires to remove the first electron from the neutral gaseous atom is called 1st ionization potential.

M _{(g) }⮴ M^{+ }_{(g)} + e^{/ }

**Electron Affinity**

Definition

The amount of energy liberated by an atom when an electron is added in it is called electron affinity. or

The energy released when an electron is added to a gaseous atom to produce gaseous negative ionis called electron affinity.

It shows that this process is an exothermic change which is represented as:

Cl^{-} + e′ ⮴ Cl^{-} ΔH = -348kj/mole

__Unit__** **

The unit of electron affinity is kilo-joule per
mole.

**Explanation**

when a neutral gaseous atom gains an electron, it looses some energy. Therefore, the electron is generally exothermic.

O^{-1} + e′ ⮴ O^{-2} Δ E = -141kj/mole

O + e′ ⮴ O Δ E = +844kj/mole

Hence, the first electron affinity is exothermic, but the successive electron affinities are endothermic.

**Factors Affecting Electron Affinity**

The magnitude of electron affinity depends on the following factors:

**Atomic Size:**With the increase in atomic size electron affinity decreases.

Nuclear Charge: Electron affinity value increases with the increase in nuclear charge. Shielding effect: With the increase in shielding effect, the electron affinity value increase.

**Electron affinity and Periodic table**

In the periodic table the electron affinity value increases from left to right in a period and decreases from top to bottom in a group.

Electron Affinity period Increasing

Group

Decreasing

**Electronegativity**

The electronegativity of an element is a measure of the attraction that an atom has for the shared electron in combined state.

The ability of an atom to attract shared electrons to itself is called electronegativity. or

The tendency of the bonded atom in a molecule to attract shared pair of electron is called electronegativity.

**Explanation**

When an atom forms a covalent bond with the other atom, then a pair of electron is shared between the atoms. Now each atom tends to acquire this shared pair of electrons. Therefore, both the atoms exert a force to attract the electron pair. This is the electronegativity of the atoms.

**Factors Affecting Electronegativity**

The magnitude of electronegativity depends on the following factors:

**Atomic Size:**With the increase in atomic size Electronegativity decrease.

**Nuclear Charge:**Increases the nuclear charge increases the Electronegativity.

Shielding Effect: Increases the shielding effect, the Electronegativity value increases from left to right in a period and decreases from top to bottom in a group.

**Electronegativity and Periodic Table**

In the periodic table, the electronegativity value increases from left to right in a period and decreases from top to bottom in a group.