| Attempt | Basis | Limitation |
|---|---|---|
| Döbereiner's Triads | Groups of 3 elements where the middle element's atomic mass ≈ average of the other two | Worked for only a few elements; most didn't fit any triad |
| Newlands' Law of Octaves | Every 8th element (by atomic mass) showed similar properties, like musical notes | Broke down after calcium; worked only for lighter elements |
| Mendeleev's Periodic Table | Elements arranged by increasing atomic mass; properties are a periodic function of atomic mass | Some elements needed to be placed out of atomic mass order; no place for isotopes; couldn't explain why properties repeated |
Mendeleev left gaps for undiscovered elements and successfully predicted the properties of "eka-aluminium" (later found to be gallium) and "eka-silicon" (later germanium) — a major reason his table gained wide acceptance despite its limitations.
The physical and chemical properties of elements are a periodic function of their atomic number (number of protons), not atomic mass. This resolved the anomalies in Mendeleev's table and gave periodicity a solid physical basis.
The table has 7 periods (horizontal rows) and 18 groups (vertical columns). A period corresponds to the filling of a particular principal shell (n); elements in the same period have the same number of shells. A group contains elements with similar valence shell electronic configurations, giving them similar chemical properties.
For newly synthesised elements with atomic number greater than 100, IUPAC assigns a temporary systematic name derived directly from the atomic number, using numerical roots (0 = nil, 1 = un, 2 = bi, 3 = tri, 4 = quad, 5 = pent, 6 = hex, 7 = sept, 8 = oct, 9 = enn), combined and ending in "-ium".
Derive the IUPAC systematic name for the element with atomic number 104.
104 → 1-0-4 → un-nil-quad → Unnilquadium (symbol Unq) — this element's accepted permanent name is now Rutherfordium (Rf).
The length of each period matches the number of orbitals available for filling in that energy range: period 1 has 2 elements (1s), periods 2 and 3 have 8 each (ns, np), periods 4 and 5 have 18 each (ns, (n-1)d, np), and periods 6 and 7 have 32 each (ns, (n-2)f, (n-1)d, np).
| Block | Last Electron Enters | Groups |
|---|---|---|
| s-block | s sub-shell | 1, 2 (+ H, He) |
| p-block | p sub-shell | 13–18 |
| d-block | d sub-shell | 3–12 (transition elements) |
| f-block | f sub-shell | Lanthanoids & actinoids |
Since electron clouds don't have a sharply defined boundary, atomic radius is measured indirectly — as half the distance between the nuclei of two bonded (covalent radius) or adjacent, non-bonded (van der Waals radius) atoms.
| Trend | Direction | Reason |
|---|---|---|
| Across a period (left to right) | Decreases | Nuclear charge increases while electrons add to the same shell — increased effective nuclear charge pulls electrons in closer |
| Down a group | Increases | A new shell is added at each step, and increased shielding by inner electrons outweighs the increase in nuclear charge |
A cation is always smaller than its parent atom (fewer electrons, same nuclear charge pulls remaining electrons in closer, and often a shell is lost entirely). An anion is always larger than its parent atom (extra electrons increase electron-electron repulsion, and effective nuclear charge per electron decreases).
Arrange Na, Na⁺, and Mg²⁺ in order of increasing size, given all are isoelectronic (same number of electrons, 10).
For isoelectronic species, the one with higher nuclear charge (more protons) pulls the same number of electrons in more tightly, giving a smaller radius. Nuclear charges: Na (11), Na⁺ (11, but fewer electrons than atom), Mg²⁺ (12).
Order of increasing size: $\text{Mg}^{2+} < \text{Na}^+ < \text{Na}$
Ionization enthalpy (IE) is the minimum energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state, forming a cation. Always endothermic (positive value).
| Trend | Direction | Reason |
|---|---|---|
| Across a period | Generally increases | Increasing effective nuclear charge holds valence electrons more tightly |
| Down a group | Decreases | Increasing atomic size and shielding make the outermost electron easier to remove |
Group 13 elements (e.g., B) have lower IE than group 2 (e.g., Be) because the removed electron in group 13 comes from a p-orbital (slightly higher energy, less penetrating, more shielded) rather than a filled, more stable s-orbital. Similarly, group 16 (e.g., O) has lower IE than group 15 (e.g., N), because N has an extra-stable, exactly half-filled 2p³ configuration that resists electron removal, while O's 2p⁴ configuration has one electron paired, which experiences extra repulsion and is easier to remove.
Explain why the first ionization enthalpy of nitrogen (1402 kJ/mol) is higher than that of oxygen (1314 kJ/mol), even though oxygen comes after nitrogen in the same period.
Nitrogen's configuration is $2p^3$ — exactly half-filled, which is extra stable due to symmetric distribution and maximum exchange energy. Removing an electron from this stable arrangement requires more energy. Oxygen's configuration is $2p^4$, where one orbital has a paired electron — the resulting inter-electron repulsion in that paired orbital makes it comparatively easier to remove, lowering oxygen's IE below nitrogen's despite oxygen's higher nuclear charge.
Electron gain enthalpy is the enthalpy change when an electron is added to an isolated gaseous atom to form an anion. Usually negative (energy released) since the added electron is attracted by the nucleus, but can be positive for some species.
| Trend | Direction (more negative = more favourable) | Reason |
|---|---|---|
| Across a period | Becomes more negative | Increasing nuclear charge attracts the incoming electron more strongly |
| Down a group | Becomes less negative | Increasing atomic size means the incoming electron is farther from the nucleus and less strongly attracted |
Noble gases have positive electron gain enthalpy (very unfavourable) since the incoming electron would have to enter a new, higher-energy shell against a stable, filled configuration. Elements with stable half-filled or fully-filled configurations (like N, with 2p³) also have less negative (less favourable) electron gain enthalpies than expected. Fluorine's electron gain enthalpy is actually less negative than chlorine's, despite F's smaller size and expected stronger attraction — because F's very small atomic size causes strong electron-electron repulsion in its already compact 2p orbital when a new electron is added.
Electronegativity is the tendency of an atom in a molecule (already bonded) to attract the shared pair of electrons towards itself. Unlike ionization enthalpy or electron gain enthalpy, it's not a measurable property of an isolated atom — it applies specifically to a bonded atom, and there's no single universally agreed absolute scale (Pauling's scale is most common).
| Trend | Direction | Reason |
|---|---|---|
| Across a period | Increases | Increasing nuclear charge and decreasing atomic size increase the pull on shared/bonding electrons |
| Down a group | Decreases | Increasing atomic size and shielding weaken the atom's pull on shared bonding electrons |
Fluorine is the most electronegative element overall (value 4.0 on the Pauling scale); caesium/francium are among the least electronegative.
For main-group elements, valence is generally related to the number of valence electrons — often equal to the number of valence electrons (for groups 1–14) or (8 − number of valence electrons) for groups 15–18. Elements in the same group typically show the same valence, since they share the same valence-shell configuration pattern.
Moving across a period, oxides change character from strongly basic (left side, e.g., Na₂O) through amphoteric (middle, e.g., Al₂O₃) to strongly acidic (right side, e.g., Cl₂O₇) — reflecting the increasing non-metallic character across a period.
The first element of each group (period 2: Li, Be, B, C, N, O, F) often behaves differently from other group members, due to its unusually small size, high electronegativity, and the absence of d-orbitals in its valence shell (limiting its maximum covalency, e.g., to 4).
Certain elements in period 2 show notable similarities to elements diagonally placed in period 3 (one group to the right), due to comparable size and charge density: Li and Mg, Be and Al, B and Si. E.g., like Al, Be forms an amphoteric oxide (BeO) and covalent halides, unlike the rest of its own group (alkaline earth metals), which are more ionic/basic.
Explain why lithium (unlike other alkali metals) forms a normal, stable oxide (Li₂O) rather than a peroxide or superoxide when burnt in excess oxygen.
Lithium's very small ionic size gives it a high charge density, favouring the formation of a simple, small oxide ion (O²⁻) lattice with high lattice energy. Larger alkali metal ions (Na, K, etc.) cannot stabilise the small O²⁻ ion as effectively and instead form larger anions like peroxide (O₂²⁻) or superoxide (O₂⁻), which pack better with their larger cation size.