Periodic Table Trends

The atomic radius is defined as half the distance between the nuclei of two identical atoms joined by a single covalent bond. For metals, it is half the distance between the nuclei of adjacent atoms in the crystal lattice. It is a measure of the size of an atom.

Across a period (left to right): Atomic radius generally decreases. This is because the nuclear charge (number of protons) increases, pulling the outer electrons more strongly towards the nucleus. The shielding effect from inner electrons remains relatively constant as electrons are added to the same main energy level. Down a group (top to bottom): Atomic radius generally increases. This is due to the increasing number of main energy levels (electron shells). Each new shell is further from the nucleus, and the increased shielding effect from the inner electrons reduces the effective nuclear charge experienced by the outer electrons, allowing them to spread out further.

The first ionisation energy (IE₁) is the minimum energy required to remove the most loosely bound electron from a gaseous atom in its ground state to form a gaseous ion with a +1 charge. The equation for this process is: X(g) → X⁺(g) + e⁻. It is a measure of how difficult it is to remove an electron from an atom.

Across a period (left to right): First ionisation energy generally increases. This is because the nuclear charge increases, and the atomic radius decreases, meaning the outer electrons are held more tightly by the nucleus. The shielding effect is relatively constant. More energy is therefore required to remove an electron. Down a group (top to bottom): First ionisation energy generally decreases. This is because the atomic radius increases, and the shielding effect from inner electrons increases. The outer electron is further from the nucleus and experiences a weaker attraction, making it easier to remove.

Electronegativity is the relative attraction an atom has for a shared pair of electrons in a covalent bond. It is a measure of an atom's ability to attract electrons in a chemical bond. The Pauling scale is commonly used to assign electronegativity values.

Across a period (left to right): Electronegativity generally increases. This is because the nuclear charge increases, and the atomic radius decreases, meaning the nucleus has a stronger pull on bonding electrons. The shielding effect is relatively constant. Down a group (top to bottom): Electronegativity generally decreases. This is because the atomic radius increases, and the shielding effect from inner electrons increases. The nucleus's attraction for a shared pair of electrons is weaker as the bonding electrons are further from the nucleus.

Metallic character refers to the chemical properties associated with metals, primarily the tendency of an atom to lose electrons and form positive ions (cations). Elements with high metallic character are good conductors of heat and electricity, are malleable and ductile, and have low ionisation energies.

Across a period (left to right): Metallic character generally decreases. This is because ionisation energy increases, making it more difficult for atoms to lose electrons. Electronegativity increases, meaning atoms are more likely to gain electrons or share them unequally. Down a group (top to bottom): Metallic character generally increases. This is because ionisation energy decreases, making it easier for atoms to lose electrons. The atomic radius increases, and shielding increases, reducing the attraction between the nucleus and the outer electrons, thus promoting electron loss.