Periodicity Flashcards

Active recall practice for s, p, d block classification, physical property trends, and electronegativity.

Classification & Concepts: read all 5 questions and answers

What is meant by the term periodicity?
The repeating pattern of physical and chemical properties of elements across different periods.
What defines an element as belonging to the s-block?
Its valence (outermost) electrons occupy an s subshell.
Which groups in the periodic table represent the p-block?
Groups 3 to 8 (or 13 to 18), excluding Helium.
Why is Helium classified as an s-block element even though it is in Group 8/0?
Its electron configuration is 1s2, meaning its outer electrons occupy an s subshell.
Where are the d-block elements located in the periodic table?
Between Groups 2 and 3, representing the transition metals where the d subshell is being filled.

Atomic Radius & Ionisation Energy: read all 6 questions and answers

State the trend in atomic radius across Period 3 and explain why.
Decreases: nuclear charge (proton number) increases while shielding remains constant, pulling outer electrons closer to the nucleus.
State the trend in first ionisation energy across Period 3.
General increase due to increasing nuclear charge and decreasing atomic radius, making it harder to remove an outer electron.
Explain the deviation in first ionisation energy between Magnesium and Aluminium.
Aluminium is lower: its outer electron is in a 3p subshell, which is higher in energy and shielded by the 3s electrons, making it easier to remove.
Explain the deviation in first ionisation energy between Phosphorus and Sulfur.
Sulfur is lower: it has paired electrons in a 3p orbital, and the electron-electron repulsion between them makes it easier to remove an electron.
State the trend in atomic radius down Group 2 and explain why.
Increases: the number of filled shells increases, which increases shielding and overrides the effect of the increasing nuclear charge.
Why does first ionisation energy decrease down Group 2?
Atomic radius and shielding increase, weakening the electrostatic attraction between the positive nucleus and the outer electrons.

Melting & Boiling Points: read all 5 questions and answers

Explain the trend in melting point from Sodium to Aluminium.
Increases: metallic bonding strengthens as the number of delocalised electrons per atom increases and the cation charge density increases.
Why does Silicon have the highest melting point in Period 3?
Silicon has a giant covalent (macromolecular) structure with a network of strong covalent bonds that require a large amount of energy to break.
Explain the trend in melting point for Phosphorus, Sulfur, and Chlorine.
Dependent on molecular size: S8 > P4 > Cl2. S8 has the most electrons, leading to the strongest van der Waals forces between molecules.
Why does Argon have a very low melting point?
It exists as monatomic atoms (Ar) with extremely weak van der Waals forces between atoms.
Explain why the boiling point of Group 7 elements increases down the group.
The number of electrons increases, making the molecules larger and strengthening the van der Waals forces between them.

Electronegativity Trends: read all 4 questions and answers

Define electronegativity.
The power of an atom to attract the pair of electrons in a covalent bond.
State and explain the trend in electronegativity across a period.
Increases: nuclear charge increases and atomic radius decreases with constant shielding, attracting the bonding pair of electrons more strongly.
State and explain the trend in electronegativity down a group.
Decreases: atomic radius and shielding increase, placing the bonding electrons further from the nucleus and weakening the attraction.
Which element is the most electronegative in the periodic table?
Fluorine.