Periodic Table SS2 Chemistry Lesson Note

A periodic table shows the arrangement or grouping of elements in order of increasing atomic number. 

Periodic Law: This is the basic assumption behind the modern periodic table; it states that the properties of the elements are the periodic function of their atomic number. 

History And Development Of The Periodic Table

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 Doberainer, Newland and Lother Meyer also attempted to classify elements based on their properties.

In 1869, Russian Scientist Dmitri Mendeleev prepared his first periodic table where he arranged 

elements based on the atomic masses (ie ATOMIC WEIGHT) of the elements. He stated the periodic 

law that the properties of the elements varied periodically with their relative atomic masses. 

He arranged the elements in the increasing order of their atomic mass. With this element having similar properties kept recurring at regular intervals or periods forming families of related elements

With the discovery of the electronic structure of the atoms, it becomes clear that elements do not vary regularly not with their atomic masses but with their atomic number, hence in the modern periodic table classification is based on atomic numbers, not atomic masses. 

Therefore, the modern periodic law states that the properties of elements are periodic functions of their atomic number 

ELECTRONIC CONFIGURATION OF THE FIRST THIRTY ELEMENTS

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

MEANING OF ATOMIC ORBITAL 

An orbital is the region of space around the nucleus where there is a high probability of finding electrons. The four different 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, d has five sub-orbitals with a maximum of ten electrons and the f-orbital has seven sub-orbitals with a maximum of fourteen electrons. 

RULES AND PRINCIPLES FOR FILLING IN ELECTRONS 

1. Aufbau Principle: states that electrons are filled in their orbital in order of increasing energy level. 

The order is as follows: 

1s < 2s < 2p< 3s < 3p <4s < 3d <4p etc. 

PERIODIC PROPERTIES. 

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

1. Atomic Radius: This has been defined as the distance of the 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 which are not chemically bonded. 

For the two types of atomic radius two variations are noticeable: 

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

This is because going 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, all the valence electrons are therefore at the same energy level. As atomic numbers increase the positive charge of the nucleus increases giving rise to greater attraction between the positive nucleus and negative electrons. This in turn results in contraction of the electron cloud resulting in a smaller atom. Atomic radii therefore decrease across a given period on the periodic table. 

1. Ionic Radius: Ions are formed by a loss or gain of electrons by an atom. A positive ion (cation) is smaller than the original metal atom because electrons are pulled in due to 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. 

As we move across the second short period, the cationic radii decrease from sodium to aluminium while the anionic radii increase from phosphorus to chlorine. 

3. Ionization Energy: Ionization occurs when a gaseous atom loses electrons from its outermost shell to become positively charged 

K———–𝑒             𝐾+———–

The energy required to do this is called ionization energy or ionization potential. First ionization energy of an element is the energy needed to remove one mole of electron(s) from one mole of atoms in the gaseous state. It is expressed in kilojoules per mole of atoms ionized. 

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

FIRST IONIZATION ENERGIES OF ALKALI METALS 

FIRST IONIZATION ENERGIES OF THE ELEMENTS IN THE THIRD PERIOD OF THE PERIODIC TABLE

Three factors affect the ionization potential of an atom is affected by:

1. Distance of the outermost electrons from the nucleus. 
2. Size of the positive or effective nuclear change.
3. Screening effect of the inner electrons.

Moving from left to right across a period, there is a general rise in the first ionization energy. This is because the nuclear charge is increasing across the period. This in turn causes a decrease in atomic radius which is 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 of the periodic table, ionization energy decreases because the nuclear charge on the outermost electron is reduced. The outermost electrons are properly shielded from the effect of nuclear charge.

4. Electron Affinity: the energy released when an electron is added to a gaseous atom in its lowest energy state. Its unit is kJmol-1 or electron volts (ev). Electron affinity increases across a period from left to right and decreases down the group on the periodic table. 

Group 1 elements, 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 stable electronic configuration 

5. Electronegativity: Electronegativity is the ability or power of 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, ie, going from left to right of the periodic table but decreases down the group i.e. going down the periodic table. The steady increase as one goes across the period is due to a steady increase in nuclear charge and a decrease in atomic size. 

Consequently, the halogen atom, Roorine, has the highest electronegativity in the period, due to the strong affinity for electrons. But down the group, the increase in atomic size due to the screening effect of the inner shells of electrons decreases the nuclear attraction for shared electrons. The noble gases of group O are not assigned electronegativity values since they have completed shells of electrons.

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.

PERIODIC GRADATION OF THE ELEMENTS IN THE THIRD PERIOD (Na-Ar)

Gradation in properties is not confined only to the elements, but it is also found in their compounds with increasing atomic number.

The extent of hydrolysis of the chlorides changes across the third period. Sodium chloride is not hydrolyzed in aqueous solution. The same applies to magnesium chloride although 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 hydrolyze completely in water.

The general conclusion from the above is therefore as follows: From the left-hand side to the right-hand side acrоss any period of representative elements, the metallic characters lie, the tendency to lose electron(s) decreases, and the non-metallic character, i.e, tendency to gain electron(s) increases. Also, as one goes across the period, ionic property decreases while covalent property increases.