Subatomic Particles:
Name 
Electron 
Proton 
Neutron 
Symbol 
e 
p 
n 
Approximate relative mass 
1/1836 
1 
1 
Mass in kg 
9.109×10^{–31} 
1.673×10^{–27} 
1.675×10^{–27} 
Mass in amu 
5.485×10^{–4} 
1.007 
1.008 
Charge (coulomb) 
1.602×10^{–19} 
1.602×10^{–19} 
0 
Actual Charge (e.s.u) 
4.8 × 10^{–10} 
4.8 × 10^{–10} 
0 
Atomic Models:
Thomson’s Atomic Model (Plum – pudding model):
Postulate: Atom is a sphere of positive charge in which number of electrons are embedded.
Limitations:  Could not satisfactorily explain the results of scattering experiment carried out by Rutherford.
Rutherford’s Model:
Postulates:
Limitations:  Could not explain stability and electronic structure of atom.
Atomic Terms
Terms 
Definition / Explanation 
Atomic Number (Z) 
Number of protons or electrons of neutral atom. 
Mass Number (A) 
Total number of protons and neutrons in an atom 
Nucleons 
Protons and neutrons are present in a nucleus. So, these fundamental particles are collectively known as nucleons 
Isotopes 
Atoms of the element with same atomic number but different mass number e.g. _{1}H^{1}, _{1}H^{2}, _{1}H^{3}. 
Isobars 
Atoms having the same mass number but different atomic numbers, e.g. _{15}P^{32} and _{16}S^{32} 
Isotones 
Atoms having the same number of neutrons but different number of protons or mass number, e.g. _{6}C^{14}, _{8}O^{16}, _{7}N^{15} 
Isoelectronic 
Atoms, molecules or ions having same number of electrons e.g. N_{2},CO, CN^{–} 
Nuclear isomers 
atoms with the same atomic number and same mass number but with different radioactive properties. Example of nuclear isomers is Uranium –X (half life 1.4 min) and Uranium –Z (half life 6.7 hours 
Isosters 
Molecules having same number of atoms and also same number of electrons are called isosters. E.g., N_{2} and CO 
Wave
Terms 
Explanation 
Wave length (λ) 
Distance between two neighbouring troughs or crests. 
Frequency (ν) 
Number of times a wave passes through a given point in a medium in one second. ν = c/λ 
Velocity (c) 
The distance travelled by the wave in one second. c = νλ 
Wave number 
Number of wavelengths per cm.

Amplitude (a) 
Height of the crest or depth of the trough. Determines the intensity of the beam of light. 
Electromagnetic Waves
Radiations 
Wave length (Å) 
Radio waves 
3×10^{14} to 3 ×10^{7} 
Micro waves 
3×10^{9} to 3 ×10^{6} 
Infrared (IR) 
6×10^{6} to 7600 
Visible 
7600 to 3800 
Ultra violet (UV) 
3800 to 150 
X–rays 
150 to 0.1 
Gamma rays 
0.1 to 0.01 
Atomic spectrum of hydrogen atom:
Where, R_{H} = Rydberg constant (108978 cm^{1})
n_{1} and n_{2} have integral values as follows
n_{1} 
n_{2} 
Spectral Series 
Spectral region 
1 
2,3,4… 
Lyman 
UV 
2 
3,4,5… 
Balmer 
Visible 
3 
4,5,6… 
Pascher 
IR 
4 
5,6,7… 
Brackett 
IR 
5 
6,7,8… 
Pfund 
IR 
Substances radiate or absorb energy discontinuously in the form of energy packets
The smallest packet of energy is called quantum. In case of light the quantum is known as photon.
The energy of a quantum is directly proportional to the frequency of the radiation.
E = hv were v is the frequency of radiation and h is Planck’s constant having the value 6.626 × 10^{–27} erg sec or 6.626 × 10^{–34} J sec.
A body can radiate or absorb energy in whole number multiples of quantum hn, 2hν,3hν………..nhν, where n is the positive integer.
Electrons revolve around the nucleus in circular orbits of fixed energy.
Electron revolve only in those orbits whose angular momentum (mvr) is an integral multiple of h/2Π.
Electron absorbs energy in the form of EMR, when it jumps from lower energy level (ground state) to higher energy level (excited state) and viceversa.
Energy absorbed or released in an electron jump, (dE) is given by dE = E_{2} – E_{1} = hν
Energy of stationary state oh hydrogen atom (E_{n}) = R_{H} (1/n^{2})
For an hydrogen like species i.e. He^{+}, Li2^{+} with atomic number Z
Radius of n^{th} orbit (r_{n} ) = 52.9 × n^{2}/z pm
Energy of n^{th} orbit (E_{n}) = 2.18×10^{18}(Z^{2}/n^{2}) = –13.6 ×(Z^{2}/n^{2}) eV = 313.6 ×(Z^{2}/n^{2}) kcal /mole
Velocity of electron (v) = (2.18 ×10^{8}) z/n cms^{1}
Where n = 1,2,3,4…
λ = h/mv = h/p
It is impossible to determine simultaneously, the exact position and exact momentum of an electron.
Principal quantum number (n):
Value of l 
0 
1 
2 
3 
4 
Notation of sub shell 
s 
p 
d 
f 
g 
Value of l 
0 
1 
2 
3 
Notation of sub shell 
s 
p 
d 
f 
Values of m 
0 
1,0,1 
2,2,0,1,2 
3,2,1,0,1,2,3 
Pauli’s Exclusion principle :
An orbital can contain a maximum number of two electrons and these two electrons must be of opposite spin.
Hund’s rule of maximum multiplicity :
Electron pairing in p, d and f orbital cannot occur until each orbital of a given subshell contains one electron each or is singly occupied”.
Stability of half filled and completely filled orbitals
Cu has 29 electrons. Its expected electronic configuration is
1s^{2}, 2s^{2}, 2p^{6}, 3s^{2}, 3p^{6}, 4s^{2}, 3d^{9}
But a shift of one electron from lower energy 4s orbital to higher energy 3d orbital will make the distribution of electron symmetrical and hence will impart more stability.
Thus the electronic configuration of Cu is
1s^{2}, 2s^{2}, 2p^{6}, 3s^{2}, 3p^{6}, 4s^{1}, 3d^{10}
Fully filled and half filled orbitals are more stable