Periodic Table & Periodic Trends — Simple Exam Notes
Concise, exam-oriented notes covering: periods & groups, s/p/d/f blocks, atomic/ionic/covalent radii, ionization potential, electronegativity, electron gain enthalpy, effective nuclear charge and screening effect.
1. The Periodic Table — Periods and Groups
Periodic table: elements arranged by increasing atomic number (Z).
- Periods (horizontal): 7 rows. Period number = highest occupied principal quantum number (n).
- Groups (vertical): 18 columns. Elements in a group have similar valence electron configuration and similar chemical properties.
2. Blocks of the Periodic Table
Blocks are defined by the orbital that receives the last electron.
- s‑block (Groups 1–2): last electron in
ns. Metals, reactive, strong reducing agents. - p‑block (Groups 13–18): last electron in
np. Contains metals, nonmetals and metalloids; includes halogens and noble gases. - d‑block (transition elements) (Groups 3–12): last electron in
(n-1)d. Variable oxidation states, colored ions, complex formation. - f‑block (inner transition) (lanthanides & actinides): last electron in
(n-2)f. Characteristic +3 states, lanthanide contraction.
3. Modern Periodic Trends (With Simple Reasoning)
Atomic radius
Definition: approximate distance from nucleus to outermost electron.
- Across a period: decreases — nuclear charge (Z) increases while shielding changes little, so electrons pulled in (↑ Zeff).
- Down a group: increases — extra shells are added (higher n).
Ionic radius
Definition: size of an ion.
- Cations are smaller than the parent atom (loss of electron(s) → decreased electron–electron repulsion and higher Zeff per electron).
- Anions are larger than the parent atom (gain of electron(s) → increased repulsion among electrons).
- Trend across period: generally decreases; down group: increases.
Covalent radius
Half the inter-nuclear distance in a covalent bond between identical atoms. Follows the same qualitative trend as atomic radius.
Ionization potential (Ionization energy, I.E.)
Definition: Energy required to remove an electron from a gaseous atom (usually the first ionization energy).
- Across a period: increases (atoms smaller and more strongly bound).
- Down a group: decreases (outer electrons are farther and more easily removed).
- Important exceptions: Be (Group 2) & B (Group 13) — B has slightly lower I.E. than Be; N (Group 15) & O (Group 16) — O has slightly lower I.E. than N due to electron pair repulsion.
Electronegativity
Definition: tendency of an atom (in a molecule) to attract shared electrons. (Pauling scale commonly used.)
- Across a period: increases.
- Down a group: decreases.
- Fluorine is the most electronegative element on the Pauling scale.
Electron gain enthalpy (Electron affinity)
Definition: enthalpy change when an electron is added to a gaseous atom.
- Across a period: becomes more negative (higher attraction for the added electron).
- Down a group: becomes less negative (added electron is farther from nucleus).
- Notes & exceptions: Noble gases have positive (unfavorable) electron gain enthalpy; fluorine is less negative than chlorine because of strong electron–electron repulsion in the small 2p orbital.
4. Effective Nuclear Charge (Zeff)
Definition: the net positive charge experienced by an electron in a multi-electron atom; approximated by
Zeff = Z − σ (where σ is the screening constant)
Slater's rules give a simple way to estimate σ. Key points:
- Across a period: Zeff increases (atoms hold electrons more tightly).
- Down a group: Zeff does not increase as rapidly; added shells increase size despite some increase in nuclear charge.
5. Screening (Shielding) Effect
Definition: inner electrons repel outer electrons and reduce the full nuclear charge felt by those outer electrons.
- Shielding (screening) ability order: s > p > d > f. That is, s-electrons shield better than p-electrons, and d/f are poor shielders.
- Poor shielding by d and f electrons explains phenomena such as lanthanide contraction — a steady decrease in size of lanthanide ions across the series, which affects subsequent elements (e.g., similar sizes of 4d and 5d elements).
6. Quick Revision (One-line Rules)
7. Useful Short Answers (Ready for 2–4 mark questions)
Q: Define ionization potential.
A: Ionization potential (ionization energy) is the minimum energy required to remove an electron from an isolated gaseous atom in its ground state.
Q: What is effective nuclear charge?
A: Effective nuclear charge is the net positive charge experienced by an electron after accounting for the screening effect of other electrons; estimated by Z−σ.
Q: Why does electronegativity decrease down a group?
A: Electronegativity decreases because added shells place the valence electrons farther from the nucleus and increase shielding, reducing the nucleus's ability to attract shared electrons.
