Chapter 3: Classification of Elements and Periodicity in Properties (Periodicity)

To arrange the elements so that elements with similar properties fall in the same group. In the modern table they are placed in order of increasing atomic number, and similar valence configurations (hence similar properties) recur at regular intervals. This makes the study of over a hundred elements systematic: we study the trends of a group or a period instead of each element separately.

Similar elements are grouped together (by increasing atomic number), so their study becomes systematic.

He used atomic weight (mass). He did not stick to it strictly: where the order by weight would put an element among unlike elements, he gave priority to similar properties. For example, he placed iodine (126.9) after tellurium (127.6) so that I fell with the halogens. He also left gaps for elements not yet discovered.

Atomic weight; no, he gave priority to similarity of properties (e.g. Te before I).

Mendeleev: the properties of elements are a periodic function of their atomic weights. Modern law: the properties are a periodic function of their atomic numbers. (Moseley showed that Z, not the mass, is the fundamental property; this also removes anomalies such as Te–I.)

Atomic weight (Mendeleev) versus atomic number (modern law).

In the 6th period (n = 6) the orbitals filled, in order of energy, are 6s, 4f, 5d and 6p: 1 + 7 + 5 + 3 = 16 orbitals. Each holds 2 electrons, so 32 electrons can be accommodated, i.e. 32 elements.

6s + 4f + 5d + 6p = 16 orbitals = 32 electrons = 32 elements.

Configuration: [Rn]5f¹⁴ 6d¹⁰ 7s² 7p². Valence shell n = 7; p-block with 4 valence electrons: group 10 + 4 = 14.

7th period, group 14 (below lead: flerovium, Fl)

Third period, group 17: 1s² 2s² 2p⁶ 3s² 3p⁵, which is chlorine.

Z = 17

(i) Lawrencium, Lr (Z = 103) (also berkelium, Bk, 97) (ii) Seaborgium, Sg (Z = 106)

Properties depend mainly on the outer (valence) electrons. Elements of a group have the same valence shell configuration and number of valence electrons (e.g. group 1: ns¹), so they behave alike.

Same valence-shell electronic configuration.

Atomic radius: the distance from the centre of the nucleus to the outermost electrons, found as half the internuclear distance: the covalent radius (half the bond length between two like atoms, e.g. Cl₂: 99 pm) or the metallic radius (half the distance between neighbouring metal atoms, e.g. Cu: 128 pm).

Ionic radius: the radius of a cation or anion, estimated from the distances between cations and anions in ionic crystals.

Size of an atom (from bond lengths) and size of an ion (from ionic crystals).

Across a period it decreases: electrons enter the same shell while the nuclear charge rises by one each step, so the effective nuclear charge increases and the electron cloud is pulled in.

Down a group it increases: each step adds a new shell (n increases), and the extra inner electrons shield the outer ones; this outweighs the rise in nuclear charge.

Decreases across a period (greater effective nuclear charge); increases down a group (new shells).

Isoelectronic species are atoms or ions with the same number of electrons.

(i) F⁻ (10 e⁻): Ne, Na⁺, O²⁻, Mg²⁺ (ii) Ar (18 e⁻): Cl⁻, K⁺, S²⁻, Ca²⁺ (iii) Mg²⁺ (10 e⁻): Na⁺, Ne, F⁻, Al³⁺ (iv) Rb⁺ (36 e⁻): Kr, Br⁻, Sr²⁺

Same number of electrons; e.g. (i) Na⁺ (ii) K⁺ (iii) Ne (iv) Kr

(a) All are isoelectronic: 10 electrons each (the neon configuration 1s² 2s² 2p⁶).

(b) The nuclear charge rises from N (7) to Al (13), pulling the same 10 electrons closer, so size falls as Z rises.

(a) 10 electrons each (b) Al³⁺ < Mg²⁺ < Na⁺ < F⁻ < O²⁻ < N³⁻

Cation: it has fewer electrons but the same nuclear charge, so the effective nuclear charge per electron is larger and the cloud contracts (often the whole outer shell is lost: Na 186 pm → Na⁺ 95 pm).

Anion: the added electron(s) increase electron–electron repulsion and lower the effective nuclear charge per electron, so the cloud expands (F 64 pm → F⁻ 136 pm).

Cations: more nuclear pull per electron; anions: more repulsion, less pull per electron.

Isolated gaseous atom: in a solid or liquid, neighbouring atoms affect the energy. In a dilute gas the atoms are far apart and do not interact, so the measured value belongs to the atom alone.

Ground state: an excited atom needs less energy to lose an electron. Taking the most stable (lowest-energy) state gives one well-defined value for every element, so the elements can be compared.

Both conditions make the values belong to the atom alone and comparable between elements.

Ionisation takes the electron to E = 0, so $2.18\times 10^{-18}$ J is needed per atom.

Per mole: $2.18\times 10^{-18}\times 6.022\times 10^{23}=1.312\times 10^6$ J mol⁻¹.

1.312 × 10⁶ J mol⁻¹ (1312 kJ mol⁻¹)

(i) Be (2s²) loses a 2s electron while B (2s² 2p¹) loses a 2p electron. The 2s electron penetrates closer to the nucleus and is held more tightly, whereas the 2p electron of B is shielded by the 2s electrons. Also 2s² is a completely filled subshell. So ΔᵢH(Be) > ΔᵢH(B).

(ii) N (2p³) has three singly occupied 2p orbitals: a stable half-filled subshell. In O (2p⁴) two electrons share one orbital and repel each other, so one of them is removed more easily: ΔᵢH(O) < ΔᵢH(N). F has a higher nuclear charge and smaller size than O, so its ΔᵢH is higher.

(i) 2s penetrates more than 2p (and 2s² is filled) (ii) half-filled 2p³ of N is stable; paired electron in O is easier to remove; F has higher Z.

First: Na (3s¹) and Mg (3s²) both lose a 3s electron, but Mg has a higher nuclear charge and smaller size (and a filled 3s), so ΔᵢH₁(Na) < ΔᵢH₁(Mg).

Second: Na⁺ has the stable neon configuration; its next electron comes from the inner 2p shell, close to the nucleus, which needs very high energy. Mg⁺ (3s¹) still has a valence electron to lose. So ΔᵢH₂(Na) > ΔᵢH₂(Mg).

Second electron of Na comes from a noble-gas core; of Mg from the valence shell.

(i) Atomic size increases (a new shell each step), so the outer electron is farther from the nucleus. (ii) Shielding by inner electrons increases. These outweigh (iii) the increase in nuclear charge, so the effective attraction on the outer electron falls and ΔᵢH decreases.

Larger size and greater shielding outweigh the higher nuclear charge.

B → Al: falls as expected (larger size).

Ga ≥ Al: Ga has ten 3d electrons, which shield the nucleus poorly, so the effective nuclear charge on its 4p electron is higher than expected and ΔᵢH does not fall.

In: drops slightly (larger size).

Tl > In: Tl has fourteen 4f and ten 5d electrons, which shield even more poorly, so the effective nuclear charge rises.

Poor shielding by d (Ga) and f + d (Tl) electrons raises the effective nuclear charge.

(i) F: it is one electron short of a noble-gas configuration and has a higher nuclear charge and smaller size than O, so adding an electron releases more energy.

(ii) Cl: in F the incoming electron enters the small, compact 2p subshell and is strongly repelled by the electrons already there; in the larger 3p subshell of Cl the repulsion is less.

(i) F (ii) Cl

Positive. The first step (O + e⁻ → O⁻) releases energy. The second electron must be added to the negative ion O⁻ against strong electrostatic repulsion, so energy has to be supplied (about +780 kJ mol⁻¹).

Positive: the electron is added to a negative ion.

Electron gain enthalpy: the enthalpy change when an isolated gaseous atom gains an electron; an experimental, measurable energy (kJ mol⁻¹).

Electronegativity: the tendency of an atom in a molecule to attract the shared electron pair of a bond; a relative number (Pauling scale), not directly measurable, and it varies with the bonding situation.

Electron gain enthalpy: energy for an isolated atom; electronegativity: relative pull on a shared pair in a molecule.

The statement is not correct. Electronegativity is not a fixed property of an atom: it depends on the atoms it is bonded to, its oxidation state and its hybridisation (an sp nitrogen is more electronegative than an sp³ one; N in NO₂ differs from N in NH₃). 3.0 is only an average value.

Not correct; it varies with bonding, oxidation state and hybridisation.

(a) Gains: the anion has more electrons than protons; electron–electron repulsion increases and the effective nuclear charge per electron falls, so the radius increases.

(b) Loses: the cation has fewer electrons but the same nuclear charge, so each electron is pulled harder (and a whole shell may be lost): the radius decreases.

(a) Radius increases (b) radius decreases

The same. Isotopes have the same nuclear charge and the same electronic configuration, and ionisation enthalpy depends on these, not on the number of neutrons (the tiny effect of different nuclear mass is negligible).

The same

Metals: low ionisation enthalpy; lose electrons to form cations (electropositive, reducing); lustrous, malleable, ductile, good conductors of heat and electricity; mostly solids with high melting points; form basic oxides.

Non-metals: high ionisation enthalpy and very negative electron gain enthalpy; gain electrons to form anions (electronegative, oxidising); dull, brittle, poor conductors (except graphite); solids or gases with low melting points (bromine is a liquid); form acidic oxides.

Metals lose electrons, are lustrous, malleable conductors with basic oxides; non-metals gain electrons, are brittle non-conductors with acidic oxides.

(a) A p⁵ element (ns² np⁵): F or Cl. (b) Group 2: Mg or Ca. (c) Group 16: O or S. (d) Group 17: F₂, Cl₂ are gases, Br₂ a liquid, I₂ a solid non-metal, and astatine is metallic in character.

(a) Cl (b) Mg (c) O (d) group 17

Alkali metals react by losing their ns¹ electron. Down the group the size increases and ΔᵢH falls, so the electron is lost more easily: reactivity increases from Li to Cs.

Halogens react by gaining an electron. Down the group the size increases and the attraction for an extra electron (electronegativity, oxidising power) falls, so reactivity decreases from F to I.

Alkali metals: ease of losing an electron increases down; halogens: ease of gaining one decreases down.

s: ns¹⁻²; p: ns² np¹⁻⁶ (He 1s² is an exception); d: (n−1)d¹⁻¹⁰ ns⁰⁻²; f: (n−2)f¹⁻¹⁴ (n−1)d⁰⁻¹ ns²

(i) 3s² 3p⁴: period 3, group 10 + 6 = 16, p-block (sulphur).

(ii) 3d² 4s²: period 4, group 2 + 2 = 4, d-block (titanium).

(iii) 4f⁷ 5d¹ 6s²: period 6, group 3, f-block lanthanoid (gadolinium, Z = 64).

(i) Period 3, group 16 (S) (ii) period 4, group 4 (Ti) (iii) period 6, group 3, f-block (Gd)

(a) V: highest ΔᵢH and positive ΔegH, a noble gas (He).

(b) II: lowest ΔᵢH₁ (like K).

(c) III: high ΔᵢH and the most negative ΔegH (like F).

(d) IV: high ΔᵢH and negative ΔegH, but less so than III (like I).

(e) VI: low ΔᵢH₁ and a ΔᵢH₂ that is not too high: an alkaline earth metal (like Mg), forming MgX₂.

(f) I: ΔᵢH₂ very high (alkali metal) but ΔᵢH₁ the highest among the metals (like Li); the small Li⁺ polarises the anion, so LiX is largely covalent.

(a) V (b) II (c) III (d) IV (e) VI (f) I

Valencies: Li 1, O 2, Mg 2, N 3, Al 3, I 1, Si 4, P 3 or 5, F 1, Lu (Z = 71, a lanthanoid) 3.

(a) Li₂O (b) Mg₃N₂ (c) AlI₃ (d) SiO₂ (e) PF₃ or PF₅ (f) LuF₃

The period number equals the principal quantum number n of the valence shell.

(c)

A d-subshell has 5 orbitals and holds 10 electrons, so the d-block has 10 columns (groups 3–12).

(b)

Neutrons add mass but no charge, so nuclear mass does not change the attraction on the valence electrons.

(c)

All have 10 electrons in the same shells; the size falls as Z rises (F 9, Ne 10, Na 11): F⁻ > Ne > Na⁺.

(a)

An electron with lower n is closer to the nucleus and held more tightly, so it is harder to remove.

(d)

Metallic character increases down a group (B → Al) and decreases across a period (K > Mg > Al).

(d)

Non-metallic character increases across period 2 (B < C < N < F); Si (period 3) is the least non-metallic.

(c)

Oxidising power follows electronegativity: F (4.0) > O (3.5) > Cl (3.0) > N (3.0); chlorine is a stronger oxidant than nitrogen.

(b)