PERIODIC TRENDS IN PROPERTIES OF ELEMENTS

 Trends in Physical Properties
Atomic Radii: It is defined as the distance from the centre of the nucleus to the outermost shell containing the electrons. Depending upon whether an element is a non-metal or a metal, three different types of atomic radii are used. These are:
(a) Covalent radius (b) Ionic Radius (c) van der Waals radius (d) Metallic radius.
(a) Covalent Radius: It is equal to half of the distance between the centres of the nuclei of two atoms held together by a purely covalent single bond.
(b) Ionic Radius: It may be defined as the effictive distance from the nucleus of an ion upto which it has an influence in the ionic bond.
(c) van der Waals Radius: Atoms of Noble gases are held together by weak van der Waals forces of attraction. The van der Waals radius is half of the distance between the centre of nuclei of atoms of noble gases.
(d) Metallic Radius: It is defined as half of the intemuclear distance between the two adjacent metal ions in the metallic lattice.

• Variation of Atomic Radius in the Periodic Table
Variation in a Period: Along a period, the atomic radii of the elements generally decreases from left to right.

Variation in a group: The atomic radii of the elements in every group of the periodic table increases as we move downwards.

• Ionic Radius
The ionic radii can be estimated by measuring the distances between cations and anions in ionic crystals.
In general, the ionic radii of elements exhibit the same trend as the atomic radii.

Cation: The removal of an electron from an atom results in the formation of a cation. The radius of cation is always smaller than that of the atom.

Anion: Gain of an electron leads to an anion. The radius of the anion is always larger than that ‘ of the atom.

Isoelectronic Species: Some atoms and ions which contain the same number of electrons, we call them isoelectronic species. For example, O2-, F, Na+ and Mg2+ have the same number of electrons (10). Their radii would be different because of their different nuclear charges.

• Ionization Enthalpy
It is the energy required to remove an electron from an isolated gaseous atom in its ground state.
M (g) + I.E ——->M+ (g) + e
The unit of ionization enthalpy is kJ mol-1 and the unit of ionization potential is electron volt per atom.

Successive Ionization Enthalpies
If a gaseous atom is to lose more than one electron, they can be removed one after the other i.e., in succession and not simultaneously. This is known as successive ionization enthalpy (or potential).

• Variation of Ionization Enthalpies in the Periodic Table:

Variation of Ionization Enthalpy Along a Period
Along a period ionization enthalpies are expected to increase in moving across from left to the right, because the nuclear charge increases and the atomic size decreases.

Variation of Ionization Ethalpy in a Group
The ionization enthalpies of the elements decrease on moving from top to the bottom in any group.
The decrease in ionization enthalpies down any group is because of the following factors.
(i) There is an increase in the number of the main energy shells (n) in moving from one element to the other.
(ii) There is also an increase in the magnitude of the screening effect due to the gradual increase in the number of inner electrons. 
M(g)    IE1    M+ + e-
M+(g)    IE2    M2+ + e-
M2+(g)    IE3    M3+ + e-

• Electron Gain Enthalpy
Electron Gain Enthalpy is the energy released when an electron is added to an isolated gaseous atom so as to convert it into a negative ion. The process is represented as:

For majority of the elements the electron gain enthalpy is negative. For example, the electron gain enthalpy for halogens is highly negative because they can acquire the nearest noble gas configuration by accepting an extra electron.
In contrast, noble gases have large positive electron gain enthalpies because the extra electron has to be placed in the next higher principal quantum energy level thereby producing highly unstable electronic configuration.

Successive Electron Gain Enthalpies
We have studied that electrons from a gaseous atoms are lost in succession (i.e., one after the other). Similarly, these are also accepted one after the other, i.e., in succession. After the addition of one electron, the atom becomes negatively charged and the second electron is to be added to a negatively charged ion. But the addition of second electron is opposed by electrostatic repulsion and hence the energy has to be supplied for the addition of second electron. Thus the second electron gain enthalpy of an element is positive.
For example, when an electron is added to oxygen atom to form O– ion, energy is released. But when another electron is added to 0- ion to form O2- ion, energy is absorbed to overcome the strong electrostatic repulsion between the negatively charged 0– ion and the second electron being added. Thus, first electron gain enthalpy:
X(g) + e-             X- (g)

Factors on which Electron Gain Enthalpy Depends
(i) Atomic size: As the size of an atom increases, the distance between its nucleus and the incoming electron also increases and electron gain enthalpy becomes less negative,
(ii)Nuclear charge: With the increase in nuclear charge, force of attraction between the nucleus and the incoming electron increases and thus electron gain enthalpy becomes more negative.
(iii) Symmetry of the Electronic Configuration: The atoms with symmetrical configuration (having fully filled or half filled orbitals in the same sub-shell) do not have any urge to take up extra electrons because their configuration will become unstable.
In that case the energy will be needed and electron gain enthalpy (Δ eg H) will be positive. For example, noble gas elements have positive electron gain enthalpies.

Variation of Electron Gain Enthalpy Across a Period
Electron gain ethalpy becomes more negative with increase in the atomic number across a period.

Variation of Electron Gain Enthalpy in a Group
Electron gain enthalpy becomes less negative as we go down a group.

• Electronegativity 
A qualitative measure of the ability of an atom in a chemical compound to attract shared electrons to itself is called electronegativity. Unlike ionization enthalpy and electron gain enthalpy, it is not a measurable quantity.
However, a number of numerical scales of electronegativity of elements viz, Pauling scale, Milliken- Jaffe scale, Allred Kochow scale have been developed. The electronegativity of any given element is not constant; it varies depending on the element to which it is bound.

Across a Period
Electronegativity generally increases across a period from left to right.

In a Group
It decreases down a group.

• Periodic Trends in Chemical Properties along a Period
(i) Metallic character: Decrease across a period maximum on the extreme left (alkali metals).
(ii) Non-metallic character: Increasess along a period. (From left to right).
(iii) Basic nature of oxides: Decreases from left to right in a period.
(iv) Acidic nature of oxides: Increases from left to right in a period.

• Variation from Top to Bottom on Moving Down a Group
(i) Metallic character. Generally increases because increase in atomic size and hence decrease in the ionizatiort energy of the elements in a group from top to bottom.
(ii) Non-metallic character. Generally decreases down a group. As electronegativity of elements decreases from top to bottom in a group.
(iii) Basic nature of oxides. Since metallic character or electropositivity of elements increases in going from top to bottom in a group basic nature of oxidise naturally increases.
(iv) Acidic character of oxides. Generally decreases as non-metallic character of elements decreases in going from top to bottom in a group.
(v) Reactivity of metals. Generally increases down a group. Since tendency to lose electron increases.
(vi) Reactivity of non-metals. Generally decreases down the group, Higher the electro-negativity of non-metals, greater is their reactivity. Since electronegativity of non-metals in a group decreases from top to bottom, their reactivity also decreases.

• Anomalous Properties of Second Period Elements
The first element of each of the group 1 (lithium) and 2 (beryllium) and group 13-17 (boron to fluorine) differs in many respect from the other members of their respective groups. For example, lithium unlike other alkali metals, and beryllium unlike other alkaline earth metals
form compounds which have significant covalent character; the other members of these groups, pre-dominatly form ionic compounds.
It has been observed that some elements of the second period show similarities with the elements of the third period placed diagonally to each other, though belonging to different groups.
For example,

This similarity in properties of elements placed diagonally to each other is called diagonal relationship.

• Mendeleev’s Periodic Law. Physical and chemical properties of elements are periodic function of their atomic masses.

• Modem Periodic Law. Physical and chemical properties of the elements are periodic function of their atomic numbers.

• Groups. There are 18 groups. These are vertical rows.

• Periods. There are 7 periods. These are horizontal rows.

• Representative Elements. The S and P block of elements are known as representative 
elements.

• Transition Elements. They are also called d-block elements. They have general electronic 
configuration (n – 1) d1-10 ns0-2.

• Inner Transition Elements. Lanthanoids (the fourteen elements after Lanthanum) and actinides (the fourteen elements after actinium) are called inner transition elements. General electronic configuration is (n – 2) f1-14(n – 1) d0-1 ns2.
They are also called f-block elements.

• Metals. Present on the left side of the periodic table. Comprise more than 78% of the known elements.

• Non-metals. Mostly located on the right hand side of the periodic table.

• Metalloids. Elements which line as the border line between metals and non-metals (e.g., Si, Ge, As) are called metalloids or semimetals.

• Atomic Radii and Ionic Radii, increase down the group decrease along the period.

• Ionization Enthalpy. Increases along the period and decreases down the group.

• Noble Gas Elements. Elements with symmetrical configuration are chemically inert in nature.

• Electric Nuclear Charge. Z = Nuclear charge – Screening constant.

• Electronegativity. Increases along a period decreases down the group,

• Chemical Reactivity. Chemical reactivity is highest at the two extremess of a period and lowest 
in the centre.

• Oxides of Elements. Oxides formed of the Elements on the left are basic and of elements
on the right are acidic in nature.
Oxides of elements in the centre are amphoteric or neutral.