Periodicity and Trends
Specification Reference Topic 1, points 24–26 (Edexcel A-Level Chemistry)
Quick Notes
- Periodicity refers to repeating trends in the properties of elements across different periods of the periodic table.
- Trends across Periods 2 and 3 include:
- Electronic configuration: elements fill the same outer shell across a period.
- Atomic radius: decreases across a period (increased nuclear charge).
- Melting and boiling points: increases to silicon (metallic to giant covalent), then decreases (simple molecules).
- Ionisation energy: generally increases across a period.
- Melting/boiling points depend on structure and bonding:
- Giant metallic/covalent = high melting point
- Simple molecular = lower melting point
- Noble gases = very low melting point
- Ionisation energy increases across a period due to stronger nuclear attraction.
Full Notes
Periodicity
Periodicity is the term used to describe patterns that repeat at regular intervals across the periods (horizontal rows) of the periodic table.
As you move across a period:
- The number of protons increases, which increases nuclear charge.
- Electrons are added to the same outer shell, so shielding stays similar.
- This results in clear, predictable trends in atomic properties.
Trends Across Periods 2 and 3
Atomic Radius
Atomic radius decreases across a period
For example Across period 3, atoms of each element get smaller.
Reason:
- Nuclear charge increases (more protons in the nucleus).
- Electron shielding remains the same or similar (all outer electrons are in the same energy level).
- Stronger attraction pulls outer electrons closer to nucleus.
Melting and Boiling Temperatures
Melting and boiling points are based on the structure and bonding of the element:
Giant metallic structures (e.g. Na, Mg, Al):
- High melting points due to strong electrostatic forces between positive ions and delocalised electrons
Giant covalent structures (e.g. Si in Period 3, C in Period 2):
- Very high melting points due to extensive covalent bonding
Simple molecular substances (e.g. P4, S8, Cl2, O2, N2, F2):
- Lower melting/boiling points due to weak London forces between molecules
Noble gases (Ne, Ar):
- Very low melting/boiling points due to weak dispersion forces
For exampleFor Period 3:
Explanation:
- Metals (Na, Mg, Al) have high melting points
Metallic bonding increases in strength from Na → Al due to more delocalised electrons and greater positive charge of metal ions. - Silicon (Si) has the highest melting point
Giant covalent structure with strong covalent bonds requires a lot of energy to break. - Non-metals (P4, S8, Cl2, Ar) have low melting points
Weak London forces hold molecules together.
S8 has a higher melting point than P4 and Cl2 because it is a larger molecule, meaning stronger London forces.
First Ionisation Energy
First Ionisation Energy Increases as a trend across a period
Explanation:
- Nuclear charge increases (more protons in nucleus).
- Atomic radius decreases, meaning outer electrons are closer to the nucleus.
- Electrons are more strongly attracted to nucleus, requiring more energy to remove.
Exceptions to the trend (see first ionisation energy and atomic structure):
- Aluminium (Al) has a lower first ionisation energy than Mg
Reason: The outer electron in Al is in the 3p subshell, which is higher in energy than the outer electron in Mg (3s) and easier to remove. - Sulfur (S) has a lower first ionisation energy than phosphorus (P)
Reason: In sulfur, the 3p orbital starts to pair electrons, causing electron pair repulsion, making it easier to remove an electron.
Summary
- Periodicity is the repetition of trends across periods of the periodic table.
- Across a period: nuclear charge increases, radius decreases, ionisation energy increases.
- Melting/boiling points vary with bonding: metals high, giant covalent highest, molecular substances low, noble gases very low.
- Exceptions to ionisation energy trends support sub-shell structure (Al, S).