# Periodic Trends in Properties of Elements

1. Periodicity in valency or oxidation states:

• Thevalence (or valency) of an element is a measure of its combining power with other atoms when it forms chemical compounds or molecules.
• The valence of elements is usually (though not necessarily) equal to the number of electrons in the outermost orbitals and/ or equal to eight minus the number of outermost electrons.
• The valency of all the elements in a group is same as the electrons in outer shell for elements of a group is same as shown below.
• Going across a period from left to right the number of valence electrons increases.
• Many elements, particularly transition elements and actinoids, exhibit variable valency as shown in the table below.
 “ous: “ic” Iron Fe²⁺ – Ferrous Fe³⁺ – Ferric Copper Cu²⁺ – Cuprous Cu³⁺ – Cupric Mercury Hg²⁺ – Mercurous Hg³⁺ – Mercuric Gold Au²⁺ – Aurous Au³⁺ – Auric Tin Sn²⁺ – Stannous Sn³⁺ – Stannic Lead Pb²⁺ – Plumbous Pb³⁺ – Plumbic

Oxidation state:

1. Nowadays the term oxidation state is frequently used for valence.
2. It is defined as the charge acquired by its atom on the basis of electronegative consideration from other atoms.
3. For example: in the compounds OF₂ and Na2O the electronegativity order for the elements is F > O > Na. the oxidation state of F is -1 and O is +2 in OF₂. The oxidation state of O is -2 and that of Na is +1 in Na₂F.
4. Its value can be zero, positive, negative or fractional.
5. In case of alkali and alkaline earth metals it is fixed i.e., + 1 and + 2 respectively.
6. Higher p-block elements exhibit more stability for lower oxidation state due to inert pair effect.
7. In case of d-block elements,
Oxidation state = (n − 1) d e(unpaired) +ns e.
In general, the stability of the higher oxidation states is in the order 3d << 4d << 5d

Inert pair effect

1. The inert pair effect is the tendency of the electrons in the outermost atomic s-orbitalto remain unionized or unshared in compounds of post-transition metals.
2. The term inert pair effect is often used in relation to the increasing stability of oxidation states that are 2 less than the group valency for the heavier elements of groups 13, 14, 15 and 16
3. As an example in group 13 the +1 oxidation state of Tl is the most stable and Tl3+ compounds are comparatively rare. The stability of the +1 oxidation state increases in the following sequence:

Al1+ < GaI+ < In1+ < Tl1+

2. Anomalous properties of second period elements – diagonal relationship

Certain elements in the 2nd period show similarity with their diagonal elements in 3rd period. This is called diagonal relationship. Thus,

• Li resembles Mg,
• Be resembles Al and
• B resembles Si.

These pairs of element have almost identical ionic radii and polarizing power (i.e. charge/size ratio). Element of second period are known as bridge elements.

This anomalous behaviour is due to

1. small size
2. high electronegativity
3. large charge to radius ratio, and
4. Non availability of d-orbitals for bonding.

Periodic trends and chemical reactivity:

• Chemical reactivity is highest at the two extremes of a period and is lowest at the center. This is because the elements present at left extreme of a period have the lowest ionization enthalpy and can lose an electron to form a cation and that present at the right extreme have the highest electron affinity and can gain an electron to form an anion.
• Among transition metals (3d series), the change in atomic radii is much smaller as compared to those of representative elements across the period.
• The change in atomic radii is still smaller among inner-transition metals (4f series). The ionization enthalpies are intermediate between those of s- and p-blocks. As a consequence, they are less electropositive and thus less reactive than group 1 and 2 metals.
• It can be directly related to the metallic and non-metallic character of elements.
• The metallic character of an element, which is highest at the extremely left decreases from left to right across the period as the electrons are more tightly bound due to increased nuclear charge.
• As you go down the periodic table, the nuclear charge increases but so does the number of shielding electrons. Consequently the dominant factor is that we have more and more energy levels and the electrons are further and further away from the nucleus. Thus it is easier for those electrons to come off. The metallic character increases down the group. Hence, the reactivity increases down the group.
• The non-metallic character increases while moving from left to right across the period.
• Nonmetals usually react by gaining electrons. With nonmetals the greater the tendency to gain electrons, the more reactive it is.
• As you go across a period, there is a greater nuclear charge and thus the electrons should be attracted more readily by elements that are further to the right and the tendency to gain electrons will increase.
• Thus the reactivity of the nonmetals increases as you go from left to right across the periodic table.
• As we go down the group the non-metallic character decrease due to decreased attraction between the nucleus and the electrons because of shielding effect. Thus the reactivity of non-metals decreases down the group.