The Periodic Table (Grade 10 NSC Matric Physical Sciences): Revision Notes

The Arrangement of the Elements

Introduction to the periodic table

The periodic table of the elements is a systematic way of organising all chemical elements according to their atomic number. This arrangement reveals important patterns in the properties of elements that help us understand and predict their behaviour.

The modern periodic table is based on the work of Russian chemist Dmitri Mendeleev, who created his version in 1869. Mendeleev was brilliant because he arranged elements by increasing atomic mass and noticed repeating patterns in their properties. He even left gaps for elements he predicted were "missing" and described what properties these unknown elements should have. When these elements were later discovered, Mendeleev's predictions proved to be remarkably accurate.

Basic structure and organisation

Groups and periods

The periodic table is organised into rows and columns that have specific names and meanings:

Groups are the vertical columns in the periodic table. These are numbered from 1 to 18 (left to right) and represent the most important way of classifying elements. Elements in the same group have similar chemical properties because they have the same number of electrons in their outer shell.

Periods are the horizontal rows in the periodic table. These are numbered from 1 to 7 (top to bottom). Elements in the same period have the same number of electron shells.

The diagram above shows how the periodic table is organised, with metals shown in grey, metalloids in light blue, and non-metals highlighted in turquoise.

Block structure

The periodic table can be divided into different blocks based on which type of orbital is being filled with electrons:

• s-block: Groups 1 and 2 (left side of table)
• p-block: Groups 13 to 18 (right side of table)
• d-block: The middle section containing transition elements

This block structure helps us understand electron configurations and predict properties of elements.

Key definitions and concepts

Before examining periodic trends, you need to understand these important terms:

Atomic radius is a measure of the size of an atom. It tells us how large an atom is from its centre to its outermost electrons.

Ionisation energy is the energy needed to remove one electron from an atom in the gas phase. We can measure first, second, third ionisation energies depending on how many electrons we remove.

Electron affinity describes how much an element "wants" to gain electrons. It measures the energy change when an electron is added to a neutral atom.

Electronegativity is the tendency of atoms to attract electrons when they form chemical bonds. The electronegativity scale runs from about 0.7 (Francium) to 4.0 (Fluorine), with fluorine being the most electronegative element.

Using the periodic table to find electron configurations

The position of an element in the periodic table tells us about its electron configuration. The period number shows which energy level is being filled, and the group number helps determine how many electrons are in the outer shell.

Periodic trends across periods

As we move across a period from left to right, several important trends occur:

The diagram above shows how three key properties change across periods:

• Atomic radius decreases (atoms get smaller)
• Ionisation energy increases (harder to remove electrons)
• Electronegativity increases (atoms attract electrons more strongly)

These trends happen because as we move across a period, the number of protons increases while electrons are added to the same energy level. More protons create a stronger positive charge that pulls electrons closer, making atoms smaller and holding electrons more tightly.

Period 3 trends - a worked example

Let's examine how properties change across period 3, from sodium to chlorine:

Element	Na	Mg	Al	Si	P	S	Cl
Chlorides	NaCl	MgCl₂	AlCl₃	SiCl₄	PCl₅ or PCl₃	S₂Cl₂	no chlorides
Oxides	Na₂O	MgO	Al₂O₃	SiO₂	P₄O₆ or P₄O₁₀	SO₂ or SO₃	Cl₂O₇ or Cl₂O
Valence electrons	3s¹	3s²	3s²3p¹	3s²3p²	3s²3p³	3s²3p⁴	3s²3p⁵
Atomic radius	Decreases across the period
First ionisation energy	Generally increases across the period
Electronegativity	Increases across the period
Melting/boiling point	Increases to silicon, then decreases to argon
Electrical conductivity	Increases from sodium to aluminium. Silicon is a semiconductor. The rest are insulators

Key observations from period 3:

• Atomic radius decreases from sodium (largest) to chlorine (smallest) because increasing nuclear charge pulls electrons closer
• Ionisation energy generally increases because electrons are held more tightly by the increasing positive charge
• Electronegativity increases as atoms become more effective at attracting electrons
• Electrical conductivity changes from metals (good conductors) to silicon (semiconductor) to non-metals (insulators)
• Melting points increase up to silicon (which has a giant covalent structure) then decrease for the molecular substances

Note that argon is not included in most comparisons because it's a noble gas with a complete electron configuration [Ne]3s²3p⁶ and doesn't readily form compounds.