Chapter 14: Classification of Elements

Matter around us exists in pure and impure forms — as elements, compounds, and mixtures. So far, 118 elements have been discovered, of which 92 are natural and 26 are artificial. Elements having similar properties are grouped together, and elements with different properties are kept in different groups — this is the basis of classification of elements.

SEE Focus: This chapter covers the periodic table, its development, characteristics of the modern periodic table, electronic configuration based on sub-shells, and how properties of elements (valency, atomic size, electropositivity/electronegativity, reactivity) vary across periods and groups.

A scientific table made for the study of elements by keeping elements of similar properties in the same group, and elements of different properties in different groups, is called a periodic table.

Dmitri Mendeleev, Russian scientist who created the first periodic table

Russian scientist Dmitri Mendeleev created a periodic table by classifying elements on the basis of their atomic weight. He formulated a law known as Mendeleev's Periodic Law.

Mendeleev's Periodic Law: The physical and chemical properties of elements are the periodic functions of their atomic weight.

This means when elements are arranged in increasing order of atomic weight, their physical and chemical properties recur periodically. Elements with similar properties lie in the same vertical column (group), and elements with different properties lie in a horizontal row (period).

Very Important for SEE: The forms of the same element having different atomic weights are called isotopes. The number of protons in an atom is always constant, but the number of neutrons may differ, thus forming isotopes.

Henry Moseley and the Periodic Table

In 1913 AD, English scientist Henry Moseley discovered that the properties of elements depend on atomic number rather than atomic weight. He formulated the modern periodic law on this basis.

Modern Periodic Law: The physical and chemical properties of elements are the periodic function of their atomic number.

Moseley made a new periodic table, called the modern periodic table or long form of periodic table, with elements arranged in increasing order of atomic number. Elements with similar properties are kept in the same vertical column (group), and elements with the same number of valence shells lie in the same horizontal row (period).

Elements of the Modern Periodic Table

The path in which electrons revolve around the nucleus of an atom is called an orbit or shell. A shell may be divided into one or more sub-shells, and each sub-shell has one or more orbitals in which electrons are found.

Since the electronic configuration based on shells alone could not explain all properties of elements, scientists developed electronic configuration based on sub-shells. The K, L, M, and N shells have 1, 2, 3, and 4 sub-shells respectively.

The s, p, d, and f sub-shells can accommodate a maximum of 2, 6, 10, and 14 electrons respectively.

Aufbau's Principle — order of filling of sub-shells

The electronic configuration of elements based on sub-shells is guided by Aufbau's Principle, which states that electrons are filled in a sub-shell in the increasing order of energy of the sub-shells.

The order of increasing energy (filling order) of sub-shells is: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p ...

SEE Focus: Electrons are not filled in the 2s sub-shell until the 1s sub-shell is completely filled — sub-shells are filled strictly in their increasing-energy order, not shell by shell.

Elements are classified into three groups — metals, non-metals, and metalloids — on the basis of their metallic character.

Metals are placed on the left side of the modern periodic table. All elements of group IA to IIIA (except boron) are metals. Metals are good conductors of heat and electricity.

Elements having one electron in their outermost shell lie in group IA (e.g., Li, Na, K). Their valence shell electronic configuration is ns1 (n = shell number). For example, the electronic configuration of sodium is 1s2, 2s2 2p6, 3s1. These are reactive metals that form a strong base or alkali when dissolved in water, so they are called alkali metals. They are soft and have low density.

Elements having two electrons in their outermost shell lie in group IIA (e.g., Mg, Ca). Their valence shell electronic configuration is ns2. For example, the electronic configuration of magnesium is 1s2, 2s2 2p6, 3s2. They are called alkaline earth metals because their oxides are soluble in water and are found on the earth's surface.

Non-metals are kept on the right side of the periodic table. Groups VA, VIA, VIIA, and VIIIA (0) contain non-metals.

Elements having seven electrons in their valence shell are kept in group VIIA (17) — e.g., F, Cl, Br, I. Their outermost electronic configuration is ns2np5. For example, the electronic configuration of chlorine is 1s2, 2s2 2p6, 3s2 3p5. They easily gain one electron to complete their octet, so they are very reactive — the most reactive non-metals. They are soft and have low density. This group is called halogens.

Elements having eight electrons in their valence shell (two for helium, which has only the first shell) are kept in group 0 — e.g., He, Ne, Ar, Kr, Xe, Rn. Their valence shell electronic configuration is ns2np6. For example, the electronic configuration of argon is 1s2, 2s2 2p6, 3s2 3p6. Since they already have a complete octet, they do not take part in chemical reactions, so they are known as noble gases or inert gases.

Elements that lie between metals and non-metals in the periodic table, showing some properties similar to metals and some similar to non-metals, are called metalloids. They are poor conductors / semiconductors of electricity — their electrical conductivity is less than metals but more than non-metals. Silicon (Si), Germanium (Ge), and Bismuth (Bi) are examples of metalloids.

The properties of elements depend on their group and period in the periodic table. This is why the periodic table is important for studying the properties of elements.

The outermost shell of an atom is called its valence shell, and the electrons present in it are called valence electrons. The number of shells stays the same across a period, but the number of valence electrons increases from left to right, so valency changes across a period.

As we move from group IA to VIIA and group 0 across a period, the valencies of elements are 1, 2, 3, 4, 3, 2, 1, and 0 respectively.

The valency of an element is determined by the number of valence electrons, and all elements of the same group have the same valency. For example, groups IA and VIIA have valency 1; groups IIA and VIA have valency 2; groups IIIA and VA generally have valency 3.

Atomic size is determined by the distance of the valence shell from the nucleus.

Across a period (left to right): The atomic size decreases. As atomic number increases, protons and electrons increase while the number of shells stays the same, so electrons are added to the same shell. The increased nuclear charge attracts the shell electrons more strongly, contracting the atom.

Down a group (top to bottom): The atomic size increases, because the number of shells increases. For example, lithium has K and L shells while sodium has K, L, and M shells — so sodium's atom is larger than lithium's, even though both are in group IA.

Electropositivity is the property of an element to lose its valence electrons and form positive ions (cations). Electronegativity is the property of an element to gain electrons in its valence shell and form negative ions (anions).

Across a period (left to right): Atomic size decreases, so the tendency to lose electrons (electropositivity/metallic character) decreases, while the tendency to gain electrons (electronegativity/non-metallic character) increases.

Down a group (top to bottom): Atomic size increases, making it easier to lose electrons, so electropositivity increases. But the increased distance from the nucleus makes it harder to attract additional electrons, so electronegativity decreases.

Across a period: The chemical reactivity of metals decreases, while the chemical reactivity of non-metals increases. The element at the extreme right of a period (noble gas) is inert.

In the third period, sodium (largest atom) is the most electropositive/reactive metal, and chlorine is the most electronegative/reactive non-metal, while argon (smallest, but with a complete octet) is inert.

Down a group: The reactivity of metals increases (they lose electrons more easily as size increases) — e.g., Be < Mg < Ca in group IIA. The reactivity of non-metals decreases (harder to gain electrons as size increases) — e.g., O > S > Se in group VIA.