The Periodic Table

Structure of the Periodic Table

The periodic table is arranged in order of increasing atomic number, which is the number of protons in an atom's nucleus. This arrangement shows repeating patterns in element properties.

Elements are organised into periods (horizontal rows) and groups (vertical columns):

• Periods indicate the number of electron shells an element has. For example, elements in period 2 have two electron shells.
• Groups contain elements with the same number of electrons in their outer shell, giving them similar chemical properties.

The periodic table also separates elements into:

• Metals: Found on the left and centre; they are typically shiny, good conductors, malleable, and ductile.
• Non-metals: Found on the right side; they are usually dull, poor conductors, and brittle if solid.
• Metalloids: Elements with properties between metals and non-metals, found along the zig-zag line dividing metals and non-metals.

For example, silicon is a metalloid with some metallic and some non-metallic properties.

Diagram suggestion: Include a labelled periodic table highlighting periods, groups, metals, non-metals, and metalloids.

History and Development

The first widely recognised periodic table was created by Dmitri Mendeleev in 1869. He arranged elements by increasing atomic mass but also grouped elements with similar properties together.

Mendeleev left gaps in his table where he predicted undiscovered elements would fit. He accurately predicted the properties of these elements, which were discovered later, confirming his table's usefulness.

For example, Mendeleev predicted an element he called "eka-silicon" which was later discovered as germanium.

The modern periodic table is arranged by atomic number, not atomic mass, following the discovery of protons. This fixed inconsistencies in Mendeleev's table and improved the organisation of elements.

Trends in the Periodic Table

Several important trends can be observed across periods and down groups:

Metallic and Non-metallic Character

Across a period (left to right), elements become less metallic and more non-metallic. For example, in period 3, sodium (Na) is a metal, while chlorine (Cl) is a non-metal.

Down a group, elements become more metallic. For example, in Group 1, lithium is less metallic than cesium.

Atomic Radius

Atomic radius is the distance from the nucleus to the outermost electron shell:

• Across a period, atomic radius decreases because the number of protons increases, pulling electrons closer.
• Down a group, atomic radius increases because additional electron shells are added.

Ionisation Energy

Ionisation energy is the energy needed to remove one electron from an atom in the gaseous state:

• Across a period, ionisation energy increases due to stronger attraction between nucleus and electrons.
• Down a group, ionisation energy decreases because outer electrons are further from the nucleus and shielded by inner shells.

For instance, sodium has a lower ionisation energy than lithium because its outer electron is further from the nucleus.

Example: Sodium has 3 electron shells, chlorine has 3 as well, but chlorine has more protons (17 vs 11). The increased nuclear charge pulls electrons closer, so chlorine has a smaller atomic radius than sodium.

For instance, as you move from lithium to sodium down Group 1, the atomic radius increases because sodium has more electron shells, making its outer electron further from the nucleus.

Metals and Non-Metals

Metals and non-metals differ in both physical and chemical properties and are positioned differently on the periodic table.

Position in the Periodic Table

Metals are found on the left and centre of the periodic table, including Groups 1, 2, and the transition metals. Non-metals are mainly on the right side, including Groups 5, 6, and 7. Metalloids lie along the dividing line.

Physical Properties

• Metals are generally shiny, good conductors of heat and electricity, malleable (can be hammered into sheets), and ductile (can be drawn into wires). They have high melting and boiling points.
• Non-metals are dull, poor conductors, brittle when solid, and have lower melting and boiling points compared to metals.

Chemical Properties

• Metals tend to lose electrons to form positive ions (cations) during chemical reactions. For example, sodium (Na) loses one electron to form Na⁺.
• Non-metals tend to gain electrons to form negative ions (anions) or share electrons in covalent bonds. For example, chlorine (Cl) gains one electron to form Cl⁻.

Example: Sodium (a metal) reacts with chlorine (a non-metal) to form sodium chloride (NaCl). Sodium loses one electron to become Na⁺, and chlorine gains one electron to become Cl⁻. The oppositely charged ions attract, forming an ionic bond.