PERIODICITY

Development of the Periodic Table

Groups & Periods

In the modern periodic table, the elements are arranged in order of increasing atomic number, Z, with elements having similar chemical and physical properties placed underneath each other in vertical columns called groups.

The groups are numbered from 1 to 18; certain groups have their own names:

The current periodic table consists of 118 elements.

The horizontal rows of elements numbered from 1 to 7 are termed periods. The period number is equal to the principal quantum number, n, and shows the highest occupied energy level in an element.

For example, Ca, Z = 20, is in period 4, so it has 4 shells. We can check this through electronic configuration as well.

Metals and non-metals

The periodic table can also be split broadly into metals and non-metals by a stepped diagonal line on the right side of the periodic table.

The elements to the left of the line are metals (except hydrogen) and non-metals lie to the right.

The elements along the line are metalloids (exhibit both metallic and non-metallic properties).

*Elements exhibit more metallic properties down the group

*Melting and boiling points increase down the group

*Metals get oxidized and non-metals get reduced.

*Oxides of metals are basic, oxides of non-metals are acidic, oxide of Al is amphoteric.

s, p, d, and f blocks

The periodic table can also be further divided into main group elements (groups 1-2 & groups 13-18) and transition elements (groups 3-11). The properties of main group elements can be predicted based on their position in the periodic table; this is less true for the properties of the transition elements.

The periodic table can be split into four blocks based on the s, p, d, f sublevels. The occupancy of electrons for each sublevel is shown below:

What can we find by looking at the periodic table?

Valence electrons (through group number)
No. of shells (through period number)
Metal/non-metal/metalloid (placement along stepped diagonal line)
Electronic configuration (s, p, d, f block placement)

Atomic Radius

The radius of a circle, R, is the distance from the centre of the circle to a point on the circumference.

Atomic radius is the distance between two nuclei, halfed.

Covalent Radius (x2 non metals)
Ionic Radius (metal + non metal)
Van der Waal's Radius (noble gases)

The bonding atomic radius is always smaller than the non-bonding atomic radius.

The valence electrons in the outer shell increase from left to right across a period. This leads to an increase in the effective nuclear charge (less shielding effect across the period), thus a larger pull on valence electrons from the nucleus. This increased pull leads to a decrease in atomic radius.

As you move down a group, the number of shells increase, leading to an increased atomic radius.

Across a period, atomic radius decreases.

Down a group, atomic radius increases.

Cations lose valence electrons and outer electrons and are more strongly attracted to the nucleus. Cations of an atom have a smaller atomic radii.

Anions gain more electrons and there is greater repulsion amongst valence electrons, leading in enlargement of shell. Anions of an atom have a larger atomic radii.

Ionization Energy

Ionization Energy is the minimum energy required to remove an electron from a neutral gaseous atom in its ground-state.

Going across a period, the ionization energy values increase for the following reasons:

As the effective nuclear charge increases from left to right across a period, the valence electrons are pulled closer to the nucleus, so the attraction between the nucleus and electrons increases. This makes it more difficult to remove an electron from the atom.
Atomic radii decreases across the period - because the distance between the valence electrons and the nucleus decreases, it becomes more difficult to remove an electron from the atom.

Going down a group, the ionization energy values decrease for the following reasons:

Atomic radii increases down the group, making it easier to remove an electron from the atom.
The shielding effect of the core electrons increases faster than the nuclear charge, weakening the attractive forces between the nucleus and outer electrons in the atom.

Electron Affinity

Electron affinity is the energy required to detach an electron from the singly charged negative ion in gaseous state. Another definition is, the energy released when 1 mol of electrons is attached to 1 mol of neutral atoms or molecules in the gas phase:

Generally, electron affinity increases across a period (except some elements) . The group 17 elements, the haolgens, have the most negative values for electron affinity. This is expected because halogens need to gain an electron to achieve the stable noble gas configuration.

Elements that do not follow these rules can be justified by looking at their electronic configuration and orbital diagrams (e.g. As - has 3 electrons in p orbital, adding one will lead to adding an electron to an orbital which already contains one, this repulsion).

There are no clear trends for electron affinity down a group.

Electronegativity

Electronegativity is the relative attraction that an atom has for the shared pair of electrons in a covalent bond.

It is measured using the Pauling scale.

Across a period, from left to right, electronegativity values increase because the effective nuclear charge and atomic radii both increase.

Down a group, electronegativity values decrease because atomic radii increases and although the nuclear charge, Z, increases, its effect is shielded by the core electrons.

Most electronegative element: Fluroine (F)

Least electronegative element: Francium (Fr)