Radioactivity 
 
In 1896, Henry Becquerel observed that the uranium salt was emitting a penetrating type of radiation spontaneously. A little later, Curie suggested that these radiations were from a uranium compound. He termed this an atomic phenomenon and named this Radioactivity.
The phenomenon of spontaneously emitting active radiations from an unstable atomic nucleus is called radioactivity. The substances emitting such radiations are called Radioactive. Such elements in the periodic table can be distinguished by their atomic weights given in parentheses.
In 1905, Rutherford analyzed the radiations emitted by naturally occurring radioactive elements as three types depending on the effect of a magnetic or electric field on them. These radiations are named alpha ( a?), beta ( b?) and gama ( g?). They are schematically shown in figure 4.

For chemical purposes, the most important types of radiation are alpha and beta particles
Every emission of an a - particle leaves a new element with atomic mass 4 units less and atomic number 2 units less than the parent; but the emission of a b - particle leaves the atomic mass unchanged and increases the atomic number by one unit.

An a - particle is a 2He4 nucleus with 2 protons and 2 neutrons

Transmutation :
The phenomenon of the conversion of a radioactive element like Ra, Pa, Ac etc. into a new element by the emission of radiations is called Transmutation (meaning conversion).
The process of conversion of one element into another by bombarding with fast moving particles such as protons, neutrons, deuterons etc is known as artificial transmutation.
The new element formed in the transmutation may be radioactive or non-radioactive.
The other way to induce transmutation is by negative beta particles which are electrons emitted from within atomic nuclei.

92U²³⁹ emits b-particle and forms an element with atomic number 93. This element is called neptunium.





Ions

Looking back at Table 3 indicates that all atoms have equal number of protons (positively charged) and electrons (negatively charged) as a result of which the net charge in the atom is 0.
But when any atom gains or loses one or more electrons we get atoms with imbalance charge which are termed as ions. Ions can be positively charged (loses electron) or negatively charged (gains electron).
Different charge states for Iron is

Fe0 - neutral Fe+2 - Ferrous ion (loses 2 electrons) Fe+3 - Ferric ion (loses 3 electrons)

The atomic charge of various elements influences its chemical behavior. This is clearly depicted in the periodic table. Elements of the same group have the same charge on their ions.

Class	Position	Charge	Examples
Alkali metals	Group IA	+1	Nað++, K+ð+, Cs+ð+
Alkaline earth metals	Group IIA	+2	Mg2+ð+, Ca2+ð+, Sr2ð++
Halogens	Group VIIA	1	F-, Cl-, Br-, I-
Noble gases	Group O	-	He, Ne, Ar

Orbitals

The region in the space around the nucleus where the probability of finding an electron is maximum is known as an orbital of that electron.
Actually, according to the concepts of wave mechanics (electron behaving as a wave), the negative charge of the electron is spread around the nucleus which is referred to as ’cloud of negative charge.’ The charge density can vary and it is maximum in particular regions. This region of higher probability of electron cloud is called as orbital.
Types of orbitals :
There are four different types of orbitals s, p, d and f which are used to express the electronic configuration (arrangement of electrons in different orbitals) of elements. All these have a particular energy and three dimensional shape.
Quantum numbers :
The state of the electron in the atom is described by its location and energy level. These states are governed by the laws of quantum mechanics. The numbers used to identify these states are termed as quantum numbers. They specify the location and the energy of an electron.

Principle quantum number (symbolized n) :
It gives the main energy level to which the electron belongs. It is related to the distance of the electron cloud from the nucleus and hence indicates the size of the electron cloud.
It is represented by integer and letter designation. It can have any integral from 1 to infinity but only 1 to 7 have been established for known elements.
If n = 1 it corresponds to the lowest energy level and is designated as ’K’ level or shell. The subsequent letters L, M, N, ... etc. denote progressive higher energy levels.

The total number of electrons in the same shell is given by 2n2
Subsidiary or orbital quantum number ( symbolized’l’ ) :
It indicates sublevel or subshell in which the electron is present. It also indicates the shape of the electron cloud.
The values of ’ l ’ range from 0 to n-1

’ l ’	=	0	1	2	3
letter designation	=	s	p	d	f
shape	=	spherical	dumb-bell	four-lobe planar	complicated
No.of orbitals	=	1	3	5	7
Max.no.of electrons	=	2	6	10	14

The maximum number of electrons each orbital can accommodate is two.
A set of electron orbitals with the same values ’n’ and ’l’ is called a subshell which is represented by the following notation:



Figure 5 subshell notation
Electronic configuration of various atoms are given in following table with valence subshell.
Table 4

Atomic No.	Element	Atomic wt.	Eletronic configuration			Valence subshell	Common valences
			n=1	n=2	n=3
1	Hydrogen	1	1s1			1s1	+1, -1
2	Helium	4	1s2			1s2	0
3	Lithium	7	1s2	2s1		2s1
	+1
4	Beryllium	9	1s2	2s2		2s2	+2
5	Boron	11	1s2	2s22p1		2p1	+3
6	Carbon	12	1s2	2s22p2		2p2	+4, +2, -4
7	Nitrogen	14	1s2	2s22p3		2p3	+5, +3, -3
8	Oxygen	16	1s2	2s22p4		2p4	-2
9	Fluorine	19	1s2	2s22p5		2p5	-1
10	Neon	20	1s2	2s22p6		2p6	0
11	Sodium	23	1s2	2s22p6	3s1	3s1	+1
12	Magnesium	24	1s2	2s22p6	3s2	3s2	+2
13	Aluminium	27	1s2	2s23s2	3p1	3p1	+3
14	Silicon	28	1s2	2s22p6	3s23p2	3p2	+4
15	Phosphorous	31	1s2	2s22p6	3s23p3	3p3	+5 , +3 , -3