The Periodic Table S3.1.3 Periodicity of Properties IB DP Chemistry Study Notes - New Syllabus 2025
The periodic table: Classification of elements- IB DP Chemistry- Study Notes
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Structure 3.1.3 – Periodicity
Structure 3.1.3 – Periodicity
Periodicity refers to the regular and repeating patterns in the physical and chemical properties of elements as you move across a period (left to right) or down a group (top to bottom) in the periodic table.
These patterns arise due to predictable changes in atomic structure, such as:
- Increasing nuclear charge (number of protons in the nucleus)
- Changes in electron configuration and sublevel filling (s, p, d, f orbitals)
- Electron shielding by inner shells, which reduces the effect of nuclear charge on outer electrons
Understanding periodicity is essential to explaining and predicting the behavior of elements in terms of:
- Atomic radius
- Ionic radius
- Ionization energy
- Electron affinity
- Electronegativity
These trends help classify elements into groups (such as alkali metals, halogens, noble gases, etc.) and predict their chemical behavior in reactions.
Example
Why is sodium more reactive than lithium, even though both are alkali metals in Group 1?
▶️Answer/Explanation
Sodium is lower in the group, so its outer electron is further from the nucleus and more shielded. This makes it easier to lose, increasing reactivity. This trend—higher reactivity down the group—is an example of periodicity.
Periodicity of Atomic Properties
As you move across a period or down a group in the periodic table, elements display clear trends in physical and chemical properties. These trends are due to changes in nuclear charge, shielding, and the number of occupied energy levels. Below are the five key periodic properties:
1. Atomic Radius
Definition: The distance from the nucleus to the outermost electron shell.
Across a period: Atomic radius decreases.
Why? The number of protons increases across a period, increasing the nuclear charge. Electrons are pulled closer to the nucleus. No new energy levels are added, so shielding remains constant.
Down a group: Atomic radius increases.
Why? New electron shells are added, increasing the distance of outer electrons from the nucleus, despite increased nuclear charge.
Example
Atomic radius of fluorine is smaller than that of lithium. Both are in Period 2, but fluorine has more protons and thus a stronger nuclear pull on electrons.
2. Ionic Radius
Definition: The radius of an ion. Cations (positive ions) are smaller than their atoms; anions (negative ions) are larger than their atoms.
Across a period: Ionic radius decreases (within isoelectronic series).
Why? Increasing nuclear charge pulls the remaining electrons closer after ionization.
Down a group: Ionic radius increases.
Why? Additional energy levels are added, increasing the size of the ion.
Example
The radius of Na+ is smaller than that of Na because it has lost an entire electron shell. Cl− is larger than Cl because added electron repulsion causes the shell to expand.
3. First Ionization Energy
Definition: The energy required to remove one mole of electrons from one mole of gaseous atoms.
Across a period: Ionization energy increases.
Why? Higher nuclear charge and smaller atomic radius mean electrons are more tightly held and harder to remove.
Down a group: Ionization energy decreases.
Why? Outer electrons are farther from the nucleus and more shielded, making them easier to remove.
Example
The first ionization energy of potassium is lower than that of sodium because the outermost electron in potassium is in the fourth shell, further from the nucleus and more shielded.
4. Electron Affinity
Definition: The energy change when one mole of electrons is added to one mole of gaseous atoms.
Across a period: Electron affinity becomes more negative (more energy is released).
Why? Increasing nuclear charge makes it more favorable for an atom to gain an electron.
Down a group: Electron affinity becomes less negative.
Why? Added electrons are farther from the nucleus and experience less attraction.
Note: Electron affinity becomes more negative across a period, but Group 2 and Group 15 elements are exceptions because their stable, filled, or half-filled electron configurations (like Group 2’s filled ns subshell or Group 15’s half-filled np subshell) make it energetically unfavorable to add an electron.
Example
Chlorine has a more negative electron affinity than bromine. Chlorine is smaller and the nucleus exerts a stronger attraction on the added electron.
5. Electronegativity
Definition: The tendency of an atom to attract a bonding pair of electrons in a covalent bond.
Across a period: Electronegativity increases.
Why? Higher nuclear charge and smaller atomic radius result in stronger pull on bonding electrons.
Down a group: Electronegativity decreases.
Why? Atoms are larger, and the bonding electrons are farther from the nucleus and more shielded.
Example
Fluorine is the most electronegative element. As you go down Group 17, electronegativity decreases from fluorine to iodine.
Summary Table
| Property | Across a Period | Down a Group |
|---|---|---|
| Atomic Radius | Decreases | Increases |
| Ionic Radius | Decreases (within charge types) | Increases |
| Ionization Energy | Increases | Decreases |
| Electron Affinity | Becomes more negative | Becomes less negative |
| Electronegativity | Increases | Decreases |
Example
Compare lithium (Li) and fluorine (F) in terms of atomic radius, first ionization energy, and electronegativity.
▶️Answer/Explanation
- Atomic Radius: Li has a larger atomic radius than F. Both are in Period 2, but F has more protons, pulling electrons closer.
- Ionization Energy: F has a higher ionization energy than Li. Its electrons are held more tightly due to higher nuclear charge and smaller radius.
- Electronegativity: F is much more electronegative (3.98) than Li (0.98), meaning it attracts bonding electrons much more strongly.
Conclusion: As you move across a period, atomic radius decreases, ionization energy increases, and electronegativity increases.
Example
Compare chlorine (Cl) and bromine (Br) in terms of electron affinity and electronegativity.
▶️Answer/Explanation
- Electron Affinity: Cl has a more negative electron affinity than Br. Although both are in Group 17, Cl is smaller and its nucleus attracts the added electron more strongly.
- Electronegativity: Cl (3.16) is more electronegative than Br (2.96). As we go down the group, electronegativity decreases due to increased atomic size and electron shielding.
Conclusion: Down a group, electron affinity becomes less negative, and electronegativity decreases due to increased shielding and atomic radius.
Example
Compare magnesium (Mg) and calcium (Ca) in terms of atomic radius, ionization energy, and electronegativity.
▶️Answer/Explanation
- Atomic Radius: Ca has a larger atomic radius than Mg. Although both are in Group 2 (alkaline earth metals), Ca is in Period 4 and has an additional electron shell.
- Ionization Energy: Ca has a lower first ionization energy than Mg. The outermost electron in Ca is farther from the nucleus and more shielded, so it is easier to remove.
- Electronegativity: Ca is less electronegative (1.00) than Mg (1.31). Down a group, electronegativity decreases due to increased atomic size and electron shielding.
Conclusion: Down a group, atomic radius increases, ionization energy decreases, and electronegativity decreases due to additional energy levels and shielding.