As we have seen, the electron configurations of the elements show a periodic variation with increasing atomic number. Consequently, there are also periodic variations in physical and chemical behavior.

A. Atomic size (Atomic radius)

The radius of an atom is an important factor in determining the ease with which the atom gain or lose electrons to form molecules or compounds. Atomic radius may be defined as a half distance between two chemically bonded atoms in a molecule

What is Atomic radius of an element?

In going down a group, the number of shells increases and atomic radius also increases.

While moving across a period in the Periodic Table, there is a decrease in atomic size of the elements as electrons are added to the same energy level.

Generally, in the Periodic Table, elements on the left side have larger atomic size, while elements on the right side have smaller atomic size.

B. Ionization Energy

The ionization energy (IE) is the amount of energy required to remove the outermost electronic in an isolated gaseous atom or ion.

M(g) + Energy → M⁺ (g) + e^-

^{The process of changing a neutral to an electrically charged species is called ionization. The minimum energy required to remove the first valence electron from an isolated neutral gaseous atom is called the first ionization energy.}

^{Factors affecting ionization energy}

• ^{Atomic size – with increasing atomic size, the attraction between the nucleus and the electrons is reduced. Thus, less energy is required to ionize the atom.}
• ^{Effective nuclear charge – the force of attraction of the nucleus for the electron is directly proportional to the nuclear charge.}
• ^{Type of electron – in a given energy level, s-electrons are closer to the nucleus than p-electrons.}

C. Electron Affinity

Electron affinity is the energy liberated when an extra electron is added in isolated neutral gaseous atom to form a free ion. It is a measure of attraction between the nucleus and the extra electrons added to the atom. The smaller the atomic size, the larger the effective nuclear charge and the higher the electron affinity.

D. Electronegativity

The term electronegativity was proposed by the American Chemist Linus Pauling in 1931. Elecronegativity is the power of an atom to attract electrons towards itself when bonded to other atoms.

Atoms of the elements in the upper right of the Periodic Table (small, non-metal atoms) attract bonding electrons most strongly.

Atoms of the elements to the left side of the table (large, metal atoms) have a weaker hold on electrons. Typical active metals have electronegativities of about 1.0 or less.

Within a period of the Periodic Table, elements become more electronegative from left to right.

E. Metallic Character

Metallic character refers to the chemical properties associated with elements classified as metals. These properties arise from the elements ability to lose electrons.

Down a group, the metallic character increases, due to the lesser attraction from the nucleus to the valence electrons.

Non-metallic Character

Non-metals tend to gain electrons in chemical reactions and have a high attraction for electrons within a compound. Since the noble gases are a special group because of their lack of reactivity, the element fluorine is the most reactive non-metal. It is not found in nature as a free element.

Non-metallic character increases from left to right across a period and deceases down a group.

G. Ionic and covalent bonding in compounds

A molecule or compound is made when two or more atoms form a chemical bond that links them together. There are two types of bonds:

• ionic bonds and
• covalent bonds

In an ionic bond, the atoms are bound together by the electrostatic forces of attraction between ions of opposite charge. Ionic bonds usually occur between metal and non-metal ions. For example, sodium (Na), a metal, and chloride (Cl), a non-metal, form an ionic bond to make NaCl.

Covalent bonds usually occur between non-metals. For example, in water (H2O) each hydrogen (H) and oxygen (O) share a pair of electrons to make a molecule of two hydrogen atoms single bonded to a single oxygen atom.

H. Lattice energy

Ionic compounds are more stable because of their elctrostatic force between the two opposite ions. After the formation of ions, they combine together to form ionic compound. The energy released in this process is known as lattice energy or lattice enthalpy.

Thus, we can write

A+ + B– → A+B– + Lattice energy

The strength of ionic bond increases with the increase of lattice energy. Lattice energy depends on two factors: size or radius of ions and charge of ions.

As the radius of ions increases, the lattice energy decreases. This is because with the increase of size of ions, the distance between their nuclei increases. Thus the attraction between them decreases and finally the less lattice energy released during the process.

Lattice energy increases with the increase of charge on the ions because of their more attractive force between them.

Here we can see that the lattice energy of MgO is much greater than the lattice energy of NaCl.