ATAR CHEM PERIODIC TABLE/TRENDS

Ionisation Energy

the amount of energy required to remove an electron from an isolated atom or molecule.

Electronegativity

Electronegativity is an atom’s ability to attract shared electrons in a chemical bond.

Atomic Radius

Atomic radius is the distance from the nucleus of an atom to the outermost electron shell.

Trend for atomic radius on the periodic table down a group and across a period:

Across a period: Atomic Radius decreases.

Down a group: Atomic Radius increases.

Trend for Ionisation Energy on the periodic table down a group and across a period:

Across a period: Increases

Down a group: Decreases

Trend for Electronegativity on the periodic table down a group and across a period:

Across a period: Increases

Down a group: Decreases

Give at least one reason why electronegativity increases across a period:

1. More Protons in the Nucleus → As you move from left to right, atoms have more protons, creating a stronger positive charge to pull electrons closer.

2. Smaller Atomic Radius → The valence electrons are closer to the nucleus, so the nucleus has a stronger pull on bonding electrons.

3. Same Number of Electron Shells → The number of shells stays the same, so there’s no extra shielding to weaken the attraction.

Give at least one reason why Atomic Radius increases down a group:

• More Electron Shells: As you move down a group, atoms have more electron shells, which makes the atom larger.

• Increased Shielding: Inner electron shells block the outer electrons from the full pull of the nucleus, reducing the attraction between the nucleus and outer electrons.

• Weaker Nucleus-Electron Attraction: As the number of electron shells increases, the outer electrons are farther from the nucleus, so the pull from the nucleus is weaker, allowing the atom to expand.

Give atleast one reason why ionisation energy increases across a period:

1. More Protons: As the number of protons increases, the nucleus has a stronger positive charge, which pulls electrons closer and makes them harder to remove.

2. Smaller Atomic Radius: As the atomic radius decreases, electrons are closer to the nucleus, so it takes more energy to remove an electron.

3. Less Shielding: There are no additional electron shells to block the nucleus’s pull, so the outer electrons experience a stronger attraction, making them harder to remove.