Classification of Elements and Periodicity — Practice Questions with Answers

The modern periodic law states that the physical and chemical properties of elements are a periodic function of their atomic number.

Mendeleev's original law was based on atomic mass, not atomic number — this earlier version led to a few inconsistencies/anomalies in element placement (e.g., some elements needed to be placed out of strict atomic mass order to correctly match their chemical properties), which were fully resolved once atomic number (established later by Moseley) was adopted as the correct organising basis.

Across a period: Electrons are added to the same outermost shell, while nuclear charge (number of protons) simultaneously increases — this increased nuclear attraction pulls the same-shell electrons in more tightly, decreasing atomic radius, despite the increasing electron count.

Down a group: A new electron shell is added at each successive element, significantly increasing the overall distance between the nucleus and outermost electrons — this shell addition effect dominates over the also-increasing nuclear charge, resulting in an overall increase in atomic radius down a group.

Ionisation enthalpy is the minimum energy required to remove the most loosely bound electron from an isolated, gaseous atom in its ground state.

Across a period: Ionisation enthalpy generally increases, due to increasing nuclear charge and decreasing atomic size, both making electron removal progressively harder.

Down a group: Ionisation enthalpy generally decreases, since the outermost electron is progressively farther from the nucleus (in a higher shell) and more effectively shielded by inner electrons, making it easier to remove.

Electron gain enthalpy is the energy change when an electron is added to a neutral, isolated, gaseous atom to form a negative ion.

Across a period, increasing nuclear charge and decreasing atomic size mean the nucleus exerts a progressively stronger attractive pull on any newly added electron — this stronger attraction makes electron addition an increasingly energetically favourable (more exothermic, more negative) process, which is why electron gain enthalpy generally becomes more negative across a period (with some exceptions/irregularities, particularly related to stable half-filled or fully-filled configurations).

Electronegativity is a measure of an atom's tendency to attract shared electrons towards itself specifically within a covalent bond (a relative, comparative concept, without fixed absolute units in most common scales like Pauling's).

Electron gain enthalpy, by contrast, is a measurable thermodynamic quantity (with defined energy units) describing the energy change when a completely isolated, gaseous atom gains an electron to form an anion — it applies to an isolated atom, not specifically to an atom already engaged within a covalent bond, distinguishing it conceptually from electronegativity.

Removing the first electron creates a positively charged cation. Removing a second electron from this already-positive cation requires overcoming an even stronger effective nuclear attraction, since the remaining electrons now experience a greater net positive charge (having one fewer electron to partially shield/repel them from the nucleus) than in the original neutral atom.

This progressively stronger nuclear attraction on the remaining electrons — as successive electrons are removed from an increasingly positive species — is precisely why each successive ionisation enthalpy (second, third, etc.) is consistently greater than the one before it.

Noble gases possess a completely filled outermost electron shell (a stable octet, except helium with a filled duet) — this exceptionally stable configuration gives them very low chemical reactivity, very high ionisation enthalpy, and generally positive/unfavourable electron gain enthalpy (little tendency to gain or lose electrons).

This behaviour is fundamentally different from halogens (which readily gain one electron to achieve a stable configuration) or alkali metals (which readily lose one electron) — this distinctly unique chemical behaviour, arising directly from their complete valence shell, is why noble gases warrant their own separate periodic group, distinct from neighbouring reactive groups.

A diagonal relationship refers to notable similarities in chemical properties between certain pairs of elements located diagonally adjacent to each other in the periodic table (typically an element in period 2 and the element diagonally below-right in period 3), rather than the expected similarity within the same group.

Lithium and magnesium exhibit such a diagonal relationship — despite being in different groups (Group 1 and Group 2 respectively), they show notably similar properties (e.g., both form relatively covalent compounds compared to their other group members, both form a nitride directly with nitrogen gas), attributed to their similar ionic size/charge density ratios despite differing atomic numbers.

A cation forms when an atom loses one or more electrons — this loss reduces electron-electron repulsion in the remaining electron cloud, and the same (now reduced) number of electrons experiences an effectively stronger net attractive pull from the unchanged nuclear charge, causing the electron cloud to contract, resulting in a smaller radius than the parent atom.

An anion forms when an atom gains one or more electrons — this addition increases electron-electron repulsion within the electron cloud, and the same nuclear charge must now attract a greater number of electrons, resulting in a somewhat weaker effective pull per electron, causing the electron cloud to expand, resulting in a larger radius than the parent atom.

While the general trend of increasing ionisation enthalpy across a period holds overall, minor exceptions occur due to additional stability associated with fully-filled or half-filled subshells.

For example, beryllium (Be, configuration $[He]2s^2$) has a completely filled, extra-stable 2s subshell, making its electron somewhat harder to remove than expected from the simple trend — whereas boron (B, configuration $[He]2s^22p^1$) has its outermost electron in a slightly higher-energy, less-stabilised 2p orbital, making it comparatively easier to remove despite boron's higher atomic number/nuclear charge — this specific extra-stability effect (from the fully-filled 2s subshell in Be) is precisely why Be's first ionisation enthalpy is anomalously higher than B's, contrary to the simple expected trend.

All four elements are in period 3, where atomic radius decreases with increasing atomic number (increasing nuclear charge, same outermost shell), so radius decreases in the order they appear in the periodic table (Na→Mg→Al→Si).

Order of increasing atomic radius: Si < Al < Mg < Na

Sodium (leftmost, lowest nuclear charge among these four) has the largest radius; silicon (rightmost, highest nuclear charge among these four) has the smallest.

All three are Group 1 (alkali metal) elements. Down a group, ionisation enthalpy decreases, since the outermost electron is progressively farther from the nucleus (in successively higher shells) and more effectively shielded.

Li is in period 2, Na in period 3, K in period 4 — so ionisation enthalpy decreases in the order Li → Na → K

Order of increasing ionisation enthalpy: K < Na < Li

Fluoride ion ($F^-$) forms when a neutral fluorine atom gains one additional electron, increasing its total electron count from 9 to 10, while the nuclear charge (9 protons) remains unchanged.

This same nuclear charge (9+) must now attract and hold one additional electron (10 electrons total, rather than 9) — since the ratio of nuclear charge to electron count has decreased, the effective attraction per electron is weaker, and increased electron-electron repulsion (from the additional electron) causes the electron cloud to expand.

This expansion is precisely why the ionic radius of $F^-$ (approximately 133 pm) is measurably larger than the atomic radius of neutral fluorine (approximately 72 pm) — gaining an electron without a corresponding increase in nuclear charge always results in radius expansion.

Nitrogen's electronic configuration is $[He]2s^22p^3$ — the 2p subshell is exactly half-filled (one electron in each of the three 2p orbitals, with parallel spins), which provides extra stability due to symmetric electron distribution and favourable exchange energy.

Oxygen's electronic configuration is $[He]2s^22p^4$ — one 2p orbital must now contain a pair of electrons (opposite spins), disrupting the extra-stable half-filled arrangement; the resulting increased electron-electron repulsion within this now-paired orbital makes one of these paired electrons comparatively easier to remove than expected from a simple, straightforward trend.

This combination of nitrogen's extra half-filled-subshell stability and oxygen's electron-pairing repulsion is precisely why nitrogen's first ionisation enthalpy (1402 kJ/mol) is anomalously higher than oxygen's (1314 kJ/mol), despite oxygen's higher atomic number and nuclear charge.

Fluorine, despite being higher in electronegativity, has an unusually small atomic size, with its 2p subshell being quite compact. This small size means an incoming additional electron experiences significant electron-electron repulsion from the already densely packed existing electrons within this small volume — this repulsion partially offsets/reduces the favourable energy release upon electron addition.

Chlorine, being larger (in period 3, with a more diffuse, spacious 3p subshell), experiences comparatively less electron-electron repulsion upon gaining an additional electron, since the added electron has relatively more space to occupy — this reduced repulsion allows chlorine's electron gain process to release comparatively more energy (be more exothermic/negative) than fluorine's, despite fluorine's overall greater electronegativity (which reflects a different, though related, property — the tendency to attract shared electrons within an existing bond, rather than gain a completely new, isolated electron).

The lanthanide contraction refers to the gradual, cumulative decrease in atomic/ionic radii observed across the lanthanide series (elements 57-71), which — although small at each individual step — accumulates to a significant overall size reduction by the end of the series.

This occurs because, across the lanthanide series, additional electrons are being added specifically to the deeply buried, inner 4f subshell — these 4f electrons are relatively poor at shielding the outer valence electrons from the correspondingly increasing nuclear charge (since 4f orbitals have unusual, diffuse shapes that don't shield outer electrons very effectively compared to typical s or p orbitals).

This weak shielding by the 4f electrons means each successive element in the series experiences an increased effective nuclear charge on its outer electrons (despite the added 4f electrons technically increasing total electron count) — this stronger effective pull causes gradual, cumulative radius contraction across the series, which is why this specific effect is termed the "lanthanide contraction."

Across period 3, from Na to Cl, nuclear charge steadily increases while atomic radius steadily decreases — this combination progressively increases ionisation enthalpy and electronegativity across the period.

Metallic character (associated with a tendency to lose electrons, correlating with low ionisation enthalpy) is highest for sodium (leftmost, lowest ionisation enthalpy) and progressively decreases across the period, becoming essentially absent by chlorine.

Non-metallic character (associated with a tendency to gain electrons, correlating with high electronegativity/favourable electron gain enthalpy) is essentially absent for sodium and progressively increases across the period, becoming most pronounced for chlorine (rightmost, excluding the noble gas argon) — this inverse relationship between metallic and non-metallic character across a period directly reflects the underlying, steadily changing nuclear charge and atomic size.

Magnesium's electronic configuration is $[Ne]3s^2$ — a completely filled 3s subshell, providing extra stability due to the fully-paired, symmetric arrangement.

Aluminium's electronic configuration is $[Ne]3s^23p^1$ — its outermost (differentiating) electron occupies the higher-energy, less-penetrating 3p subshell, rather than the extra-stable, lower-energy 3s subshell.

Since removing an electron from the higher-energy, less-stabilised 3p orbital of aluminium requires less energy than removing an electron from magnesium's extra-stable, fully-filled 3s subshell, aluminium's first ionisation enthalpy is anomalously lower than magnesium's, despite aluminium having one additional proton (higher nuclear charge) — this is directly analogous to the similar Be/B anomaly discussed for period 2.

Potassium (K, Z=19) and calcium (Ca, Z=20) are adjacent elements in period 4, with calcium having one additional proton (higher nuclear charge) compared to potassium.

Atomic radius: Since both elements are in the same period, increasing nuclear charge (moving from K to Ca) pulls the outer electron cloud in more tightly, without adding a new shell — therefore, potassium has the larger atomic radius of the two.

Ionisation enthalpy: Following the general periodic trend of increasing ionisation enthalpy across a period (due to increasing nuclear charge and decreasing radius), potassium has the lower first ionisation enthalpy of the two — consistent with potassium's well-known high reactivity as an alkali metal, readily losing its single valence electron.

Sodium's electronic configuration is $[Ne]3s^1$ — removing the single 3s valence electron (first ionisation) leaves behind the extremely stable, completely filled neon-like core configuration ($[Ne]$). Removing a second electron (second ionisation) would require breaking into this exceptionally stable, filled inner shell, requiring vastly more energy — this is why sodium shows an enormous jump between its first and second ionisation enthalpies.

Magnesium's electronic configuration is $[Ne]3s^2$ — removing its first and second valence electrons (both from the 3s subshell) is comparatively straightforward, leaving behind the same stable $[Ne]$ core after the second ionisation. Only upon attempting a third ionisation (breaking into this now-exposed, stable $[Ne]$ core) does magnesium show its own dramatic jump — analogous to sodium's jump, but occurring one ionisation step later, specifically because magnesium has two valence electrons (rather than sodium's one) available for relatively easier removal before reaching its own stable core configuration.