Understanding Periodic Trends
The periodic table reveals clear patterns in how element properties change across periods and down groups. These periodic trends are crucial for predicting element behavior. Let’s explore the major trends.
1. Atomic Radius (Atomic Size)
Definition: The atomic radius is half the distance between the nuclei of two identical atoms joined by a single covalent bond.
Atomic Radius – Across a Period (Left to Right)
Trend: DECREASES across a period
Reason:
- Number of protons in the nucleus increases
- Electrons are added to the same shell (same principal quantum number)
- Increased nuclear charge pulls electrons closer
- Screening (shielding) effect by inner electrons remains nearly constant
- Result: Effective nuclear charge increases, pulling valence electrons closer
Example – Period 2: Li → Be → B → C → N → O → F → Ne
Atomic radius (nm): 0.123, 0.089, 0.082, 0.077, 0.070, 0.066, 0.064, -
Lithium is the largest, fluorine is the smallest
Atomic Radius – Down a Group (Top to Bottom)
Trend: INCREASES down a group
Reason:
- New electron shells are added as you go down
- Each new shell is further from the nucleus
- Although nuclear charge increases, the shielding effect of inner electrons is more significant
- Effective nuclear attraction on valence electrons decreases
- Result: Atoms get larger
Example – Group I (Alkali Metals): Li → Na → K → Rb → Cs
Atomic radius (nm): 0.15, 0.19, 0.23, 0.25, 0.26
Lithium is smallest, cesium is largest
2. Ionisation Energy (IE)
Definition: The first ionisation energy is the minimum energy required to remove one mole of electrons from one mole of gaseous atoms to form gaseous ions.
X(g) → X⁺(g) + e⁻ ΔH = First IE
Ionisation Energy – Across a Period
Trend: INCREASES across a period
Reason:
- Atomic size decreases (electrons closer to nucleus)
- Nuclear charge increases
- Effective nuclear charge increases
- Electrons are held more tightly
- More energy needed to remove an electron
General Pattern: Li < Be < B < C < N < O < F < Ne
However, there are exceptions:
- Be > B (half-filled 2s orbital is more stable than 2s²2p¹)
- N > O (half-filled 2p³ is more stable than 2p⁴)
Ionisation Energy – Down a Group
Trend: DECREASES down a group
Reason:
- Atomic size increases
- Outermost electrons are further from nucleus
- Shielding effect by inner electrons increases dramatically
- Effective nuclear attraction decreases
- Electrons are easier to remove
Example – Group I: Li > Na > K > Rb > Cs
First IE (kJ/mol): 520, 500, 420, 400, 380
Why is the ionisation energy of Magnesium higher than Aluminium?
Mg (3s²) has a complete filled 3s orbital which is more stable than Al (3s²3p¹) with an incomplete 3p orbital.
Second Ionisation Energy
The second ionisation energy is always much higher than the first because you’re removing an electron from a positive ion that holds its electrons more tightly.
Example: Na → Na⁺ (easier, removes one electron from valence shell)
Na⁺ → Na²⁺ (extremely difficult, must remove electron from a filled, inner shell)
3. Electron Affinity (EA)
Definition: Electron affinity is the energy change when one mole of electrons is gained by one mole of gaseous atoms.
X(g) + e⁻ → X⁻(g) ΔH = Electron Affinity
Important: Negative EA means energy is released (exothermic, favorable)
Electron Affinity – Across a Period
Trend: INCREASES (becomes more negative) across a period
Reason:
- Atomic size decreases
- Nuclear charge increases
- Added electron is held more tightly
- More energy released when electron is added
Exception: Group II elements (Be, Mg) and Group VIII (noble gases) have high positive values because they already have stable electron configurations.
Electron Affinity – Down a Group
Trend: DECREASES (becomes less negative) down a group
Reason:
- Atomic size increases
- Added electron is further from nucleus
- Less strongly attracted to nucleus
- Less energy released
4. Electronegativity
Definition: Electronegativity is the ability of an atom to pull bonded pair of electrons to itself in a covalent compound.
Note: Electronegativity is NOT the same as electron affinity or ionisation energy
Electronegativity – Across a Period
Trend: INCREASES across a period
Reason:
- Atomic size decreases
- Nuclear charge increases
- Atoms pull electron pairs more strongly
Example: Na < Mg < Al < Si < P < S < Cl
Fluorine is the most electronegative element
Electronegativity – Down a Group
Trend: DECREASES down a group
Reason:
- Atomic size increases
- Shielding effect increases
- Nuclear charge has less effect on bonding electrons
- Atoms pull electron pairs less strongly
Example: F > Cl > Br > I
5. Metallic Character
Definition: Metallic character is the tendency of an element to form positive ions and act as a reducing agent.
Metallic Character – Across a Period
Trend: DECREASES across a period
Reason:
- Ionisation energy increases
- Atoms hold electrons more tightly
- Less likely to form positive ions
- Non-metallic character increases
Metallic Character – Down a Group
Trend: INCREASES down a group
Reason:
- Ionisation energy decreases
- Atoms lose electrons more easily
- More likely to form positive ions
- Greater metallic character
Summary Table of Periodic Trends
| Property | Across Period (L→R) | Down Group (T→B) | Reason |
|---|---|---|---|
| Atomic Radius | Decreases | Increases | Nuclear charge vs. shielding effect |
| Ionisation Energy | Increases | Decreases | Distance from nucleus and shielding |
| Electron Affinity | Increases | Decreases | Attraction to nucleus |
| Electronegativity | Increases | Decreases | Ability to attract bonded electrons |
| Metallic Character | Decreases | Increases | Tendency to lose electrons |
Why These Trends Matter
- Predicting reactivity: More reactive metals have lower IE and higher metallic character
- Bond formation: Electronegativity determines bond polarity
- Oxidation states: Trends help predict common oxidation states
- Compound formation: Understanding trends helps predict compound formulas
- Element properties: Allows prediction of unknown element properties
