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. 1H1, 1H2, 1H3. |
Isobars |
Atoms having the same mass number but different atomic numbers, e.g. 15P32 and 16S32 |
Isotones |
Atoms having the same number of neutrons but different number of protons or mass number, e.g. 6C14, 8O16, 7N15 |
Isoelectronic |
Atoms, molecules or ions having same number of electrons e.g. N2,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., N2 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 |
3×1014 to 3 ×107 |
Micro waves |
3×109 to 3 ×106 |
Infrared (IR) |
6×106 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, RH = Rydberg constant (108978 cm-1)
n1 and n2 have integral values as follows
n1 |
n2 |
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 vice-versa.
Energy absorbed or released in an electron jump, (dE) is given by dE = E2 – E1 = hν
Energy of stationary state oh hydrogen atom (En) = -RH (1/n2)
For an hydrogen like species i.e. He+, Li2+ with atomic number Z
Radius of nth orbit (rn ) = 52.9 × n2/z pm
Energy of nth orbit (En) = -2.18×10-18(Z2/n2) = –13.6 ×(Z2/n2) eV = 313.6 ×(Z2/n2) kcal /mole
Velocity of electron (v) = (2.18 ×108) 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
1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d9
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
1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 3d10
Fully filled and half filled orbitals are more stable