The Periodic Law

The modern periodic law states that the properties of elements are a periodic function of their atomic numbers.
Atomic number is the number of protons present in the nucleus of an atom.
Shells or energy levels are spherical regions around the nucleus around which electrons revolve. The shells are situated at various distances from the nucleus. The shells are named K, L, M, N, O… which are assigned numbers 1, 2, 3, 4, 5, …All electrons in a given shell possess approximately equal energy.
The maximum number of electrons a shell can hold is given by 2n2. Viz:
2(1)2 =2
2(2)2 =8, etc.
The electronic configuration of the first twenty elements show that electrons are added to a shell until a stable duplet (for K shell) or octet state is reached and the electrons move to the shell with the next higher energy level.
Note: Besides the eight groups of elements on the periodic table, we have the transition group which lie between group 2 and 3, and have incompletely filled d-orbitals. Among group 6 and 7 elements are the Lanthanides and Actinides which lie between group 2 and 3 of periods 6 and 7.
Periodic Table and atomic propoerties
Atomic size or Radius: This is half the distance of closest approach between the nuclei of atoms in the elemental substance.
Across the period as we move from left to right, more electrons are added to the same shell at about same same distance from the nucleus. At the same time, more protons are added. At this point, the nuclear charge exerts more attractive force on the electrons, pulling them towards the nucleus, and also force of repulsion exists amongst the electrons. The force of attraction which is higher than the repulsion decreases the size of the atom as atomic number increases. Thus, atomic radius decreases across the period as atomic number increases due to increase in attractive force, and increases down the group.
Ionic size/Radius
Positive ions are smaller than the corresponding atomic radius because they are formed by loss of electrons while ionic radius of negative ions is greater than the corresponding atomic radius because they are formed by gain of electrons.
Variation of some properties
Melting and boiling point: From group 1-3, melting and boiling point increase because they are metals. Group 4 have very high boiling point because they can exist in crystallines form, while from group 5-0 , they have low boiling and melting point because they are covalent compounds which require low energy to break their bonds.
Ionization Energy: This is the energy required to remove a valence electron from an atom of the element to form an ion. It generally increases across/within a period.
The first ionization energy is the amount of energy required to remove one electron from each atom in a mole of gaseous atom producing one mole o gaseous ion with a positive charge.
Factors affecting Ionization energy of an atom
Distance of the outermost electron from the nucleus: Across the period, the attractive force of the nucleus on the outermost electrons increases. Thus, more energy will be needed to remove the electrons. This makes ionization energy increase across the period
Size of the positive nuclear charge: As the nuclear charge increases, nuclear attraction for outermost electrons increases. Hence, to remove electrons from outermost shell, more energy will be required. Thus, as nuclear charge increases, ionization energy increases.
Screening effects of inner electrons: As we move across the period (left to right), the number of electrons increases. However this increase, the electrons still occupy the same shell. Hence, screening effect is almost the same.
Specifically, despite the general increase in ionization energy across the period, He has a higher ionization energy than Li because it has a more stable electronic configuration.
He- 1s2
Li- 1s2, 2s1
Assignment 1
Explain the variation in Ionization energy of
Be and B
N and O
Electron affinity: Ionization energy deals with electron loss while electron affinity deals with electron gain. Electron affinity is the energy released when a gaseous atom gains an electron to form a gaseous negative ion.
A(g) + e- —- A-(g)
Electron affinity generally increases across the period except noble gases which do not gain electrons i.e. they have zero electron affinity. It increases across the period with decrease in atomic radius, and decreases down the group with increase in atomic radius.
Electronegativity: This is the power of an atom in a molecule to attract electron to itself. It increases across the period but decreases dow the group.
Transition Metals (Properties)
They are elements between group 2 and 3 on the periodic table
They are all metals with metallic lustre such as high boiling point, high melting point, etc.
They have variable oxidation state
They form coloured salts
They form complex ions.
Examples are Iron, Copper, Cobalt, etc
Assignment 2
State three properties each of group 1 to group 8 elements stating their family names.