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TOPIC: PERIODIC TABLE
CONTENT:
- Periodic properties
- Periodic gradation of the element in the third period
- Diagonal relationships
PERIOD 1 AND 2: PERIODIC PROPERTIES. Some properties of the atom change along a group or across a period on the periodic table. Atomic radius which is measured of 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.
- Atomic radius: This has been 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. 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:
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- The atom radius increases down a group
- 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 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 number increase the positive charge of the nucleus increases giving rise to greater attraction between the positive nucleus and negative electrons. This is turn result in contraction of the electrons cloud resulting in a smaller atom. Atomic radii therefore decrease across a given period on the periodic table.
- 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 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 phosphorous to chlorine.
- Ionization energy: Ionization occurs when gaseous atom loses electrons from its outer most 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 kilo- joules per mole of atoms ionized.
First ionization energy increase across the period with noble gases having the highest. As we go down, he the group, the value of first ionization energy decreases.
FIRST IONIZATION ENERGIES OF ALKALI METALS
| Element | LI | Na | K | Rb | Cs |
| First ionization energy KJMOL-1 | 520 | 500 | 420 | 400 | 380 |
FIRST IONIZATION ENERGIES OF THE ELEMENTS IN THE THIRD PERIOD OF THE PERIODIC TABLE
| Element | Na | Mg | Al | Si | P | S | Cl | Ar |
| First ionization energy KJMOL-1 | 496 | 737 | 577 | 786 | 1012 | 999 | 1255 | 1521 |
Three factors that affect the ionization potential of an atom
Ionization potential of an atom is affected by.
- Distance of the outer most electrons from the nucleus.
- Size of the positive or effective nuclear change.
- 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 due to the fact that the nuclear charge is increasing across the period. This in turn causes a decrease in atomic radius that 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 electron are properly shielded from the effect of nuclear charge
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