Periodic Table – SS2 Chemistry Lesson Note (First Term)

A periodic table shows the arrangement or grouping of elements in order of increasing atomic number. The modern periodic law states that the properties of elements are periodic functions of their atomic number. Periodic properties such as atomic radius, ionisation energy and electronegativity change down a group and across a period.

Behavioural objectives

By the end of the lesson, students should be able to:

• state the modern periodic law and explain how it differs from Mendeleev’s law;
• state the number of sub-orbitals and the maximum number of electrons in the s, p, d and f orbitals;
• state the Aufbau principle and the order in which orbitals are filled;
• describe how atomic radius, ionic radius, ionisation energy, electron affinity and electronegativity vary across a period and down a group;
• describe how the hydrolysis of chlorides and the metallic character of elements change across the third period.

What is the periodic law?

The periodic law is the basic assumption behind the modern periodic table. It states that the properties of the elements are periodic functions of their atomic number.

How was the periodic table developed?

In chemistry, the classification of elements began with Antoine Lavoisier, a French chemist, in 1789. He grouped the elements into gases, metals, non-metals and earth.

Scientists like Döbereiner, Newlands and Lothar Meyer also attempted to classify elements based on their properties.

In 1869, the Russian scientist Dmitri Mendeleev prepared his first periodic table. He arranged the elements in increasing order of their atomic masses (i.e. atomic weight). He stated the periodic law that the properties of the elements varied periodically with their relative atomic masses. Elements with similar properties kept recurring at regular intervals or periods, forming families of related elements.

With the discovery of the electronic structure of atoms, it became clear that the properties of elements vary regularly with their atomic number rather than their atomic mass. The modern periodic table is therefore based on atomic numbers.

The modern periodic law states that the properties of elements are periodic functions of their atomic number.

What is electronic configuration?

The electronic configuration of an atom is the representation of the arrangement of the electrons distributed among the orbital shells and subshells. It is commonly used to describe the orbitals of an atom in its ground state.

What is an atomic orbital?

An orbital is the region of space around the nucleus where there is a high probability of finding electrons. The four types of orbitals are s, p, d and f. These orbitals have different shapes, and one orbital can hold a maximum of two electrons. The p orbital has three degenerate orbitals, with a maximum of six electrons. The d orbital has five sub-orbitals, with a maximum of ten electrons. The f orbital has seven sub-orbitals, with a maximum of fourteen electrons.

What rules govern the filling of electrons?

Aufbau principle: electrons are filled into orbitals in order of increasing energy level.

The order is as follows: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p, and so on.

What are periodic properties?

Some properties of the atom change down a group or across a period on the periodic table. Atomic radius, which is a measure of the size of an atom, is one of them. The orbiting electrons in an atom are best represented by an electron cloud, which has no distinct limit, so the size of an atom cannot be defined easily.

Atomic radius

Atomic radius is defined as the distance of closest approach to another identical atom in a given bonding situation. There are two types of atomic radii: covalent radius and Van der Waals radius. A covalent radius is half the distance between two identical atoms that are chemically bonded.

For both types of atomic radius, two variations are noticeable:

1. The atomic radius increases down a group.
2. The atomic radius decreases across a period.

Down any group on the periodic table, the number of valence electrons remains constant while the shells increase in size (radius), despite the increase in nuclear charge. The atomic radius of potassium is greater than that of sodium. The atomic radius of caesium is greater than that of rubidium.

Across a period, electrons are added to orbitals in the same shell, so all the valence electrons are at the same energy level. As the atomic number increases, the positive charge of the nucleus increases, giving greater attraction between the positive nucleus and the negative electrons. This contracts the electron cloud and results in a smaller atom. Atomic radii therefore decrease across a given period.

Ionic radius

Ions are formed by the loss or gain of electrons by an atom. A positive ion (cation) is smaller than the original metal atom because the electrons are pulled in by an increase in effective nuclear charge. A negative ion (anion) is bigger than the corresponding non-metal atom because the effective nuclear charge is reduced.

Across the second short period, the cationic radii decrease from sodium to aluminium, while the anionic radii decrease from phosphorus to chlorine.

Ionisation energy

Ionisation occurs when a gaseous atom loses electrons from its outermost shell to become positively charged:

K(g) → K⁺(g) + e^–

The energy required to do this is called ionisation energy or ionisation potential. The first ionisation energy of an element is the energy needed to remove one mole of electrons from one mole of atoms in the gaseous state. It is expressed in kilojoules per mole of atoms ionised.

The first ionisation energy increases across the period, with the noble gases having the highest. As we go down a group, the value of the first ionisation energy decreases.

First ionisation energies of alkali metals

First ionisation energies of the elements in the third period

The ionisation potential of an atom is affected by three factors:

1. The distance of the outermost electrons from the nucleus.
2. The size of the positive or effective nuclear charge.
3. The screening effect of the inner electrons.

Moving from left to right across a period, there is a general rise in the first ionisation energy. This is because the nuclear charge increases across the period. This in turn causes a decrease in atomic radius, which means a decrease in the distance of the outermost electrons from the nucleus. The screening effect is almost the same across the period.

Down a group, ionisation energy decreases because the nuclear charge on the outermost electron is reduced. The outermost electrons are properly shielded from the effect of nuclear charge.

Electron affinity

Electron affinity is the energy released when an electron is added to a gaseous atom in its lowest energy state. Its unit is kJ mol^-1 or electron volts (eV). Electron affinity increases across a period from left to right and decreases down the group.

Group 1 elements (the alkali metals) have the least tendency to add electrons to their neutral atoms. Elements in groups VI and VII have the greatest tendency to accept electrons. Noble gases (group 8 or 0) have a stable electronic configuration.

Electronegativity

Electronegativity is the ability or power of an atom in a molecule to attract a shared pair of electrons. It is more pronounced in heteronuclear molecules, where two dissimilar atoms share one or more pairs of electrons.

Electronegativity increases across the period, going from left to right of the periodic table, and decreases down the group. The steady increase across the period is due to a steady increase in nuclear charge and a decrease in atomic size.

Consequently, the halogen fluorine has the highest electronegativity in its period, due to its strong affinity for electrons. Down the group, the increase in atomic size and the screening effect of the inner shells of electrons decrease the nuclear attraction for shared electrons. The noble gases of group 0 are not assigned electronegativity values because they have completed shells of electrons.

How do properties change across the third period (Na to Ar)?

Properties change gradually with increasing atomic number, both in the elements and in their compounds.

The extent of hydrolysis of the chlorides changes across the third period:

• Sodium chloride is not hydrolysed in aqueous solution.
• Magnesium chloride is also not hydrolysed in solution, although its hydrated crystals undergo hydrolysis when heated, giving off HCl and leaving a basic salt.
• An aqueous solution of aluminium chloride shows appreciable hydrolysis and turns blue litmus red.
• The chlorides of silicon, phosphorus and sulphur are hydrolysed completely in water.

The general conclusion is that, from left to right across any period of representative elements, metallic character decreases and the tendency to lose electrons decreases. Non-metallic character, i.e. the tendency to gain electrons, increases. Also, across the period, ionic property decreases while covalent property increases.

Evaluation

1. State the modern periodic law.
2. Define the first ionisation energy of an element.
3. Explain why atomic radius decreases across a period.
4. Why does the first ionisation energy decrease down a group?
5. State the maximum number of electrons that the s, p, d and f orbitals can hold.

Assignment

1. List three periodic properties of elements that generally increase across the period of the Periodic Table.
2. Explain the term electron affinity and discuss how it varies across the period and down the group of the Periodic Table.