JEE Chemistry Classification of Elements and Periodicity in Properties Notes Set 01

Development of the Periodic Table

1. Dobereiner (triads): in a group of three similar elements, the atomic mass of the middle one is about the mean of the other two (Ca, Sr, Ba; Cl, Br, I; S, Se, Te).
2. Newlands (law of octaves): arranged by increasing atomic mass, every eighth element repeats the properties of the first. It failed beyond calcium.
3. Lothar Meyer: plotted atomic volume against atomic weight; alkali metals sit at the peaks.
4. Mendeleev: properties are a periodic function of atomic mass. He left gaps and predicted eka-aluminium (Ga), eka-silicon (Ge) and eka-boron (Sc). His table had anomalies (Ar-K, Co-Ni, Te-I) where a heavier element precedes a lighter one.
5. Moseley and the modern periodic law: properties are a periodic function of atomic number. This removes the anomalies.

Long Form of the Periodic Table

The table has 7 periods and 18 groups. The periods contain 2, 8, 8, 18, 18, 32 and (incomplete) 32 elements. Elements in a group have the same valence-shell configuration; elements in a period have the same number of shells. Blocks are named after the subshell receiving the last electron.

To place an element: the period is the principal quantum number of the outermost shell; for s and p blocks the group number is 1 or 2 for s, and 10 + the number of valence electrons for p (group 13 to 18); for the d-block it is the sum of (n-1)d and ns electrons. Example: the configuration ending 5s2 5p3 is period 5, group 15 (antimony); 4d3 5s2 is period 5, group 5 (niobium).

IUPAC Names of Elements with Z above 100

Write the digits of the atomic number as the roots nil (0), un (1), bi (2), tri (3), quad (4), pent (5), hex (6), sept (7), oct (8), enn (9) and end with -ium. Z = 104 is unnilquadium, Z = 105 is unnilpentium, Z = 106 is unnilhexium and Z = 112 is ununbium. Z = 114 would be in the carbon family and Z = 118 would be a noble gas.

Screening and Effective Nuclear Charge

Inner electrons reduce the attraction of the nucleus on valence electrons. The effective nuclear charge is \( Z_{\text{eff}} = Z - \sigma \), where \( \sigma \) is the screening (shielding) constant. Across a period, Zeff rises, because the added electron goes into the same shell and screens poorly. Down a group, the screening rises, so Zeff changes only slightly. By Slater's rules, Zeff of the last electron is 1.30 for Li and 2.20 for Na and K. Screening by d and f electrons is poor, which explains the d-block contraction (Ga smaller than Al) and the lanthanoid contraction.

Atomic and Ionic Radius

Types: covalent radius (half the bond length of identical atoms), metallic radius, and van der Waals radius (the largest, used for noble gases). Trends: the radius decreases across a period and increases down a group. In the 3d series the radius decreases at first, stays nearly constant in the middle (Fe, Co, Ni), and rises slightly at the end (Cu, Zn). Because of lanthanoid contraction, the 4d and 5d elements of a group (Zr, Hf; Nb, Ta) have almost equal sizes.

Ionic radii: a cation is smaller than its atom; an anion is larger. For isoelectronic species the radius decreases as the nuclear charge increases: \( \ce{N^3- > O^2- > F^- > Na+ > Mg^2+ > Al^3+} \), and \( \ce{S^2- > Cl- > Ar > K+ > Ca^2+} \). For the same element the radius falls as the positive charge rises: \( \ce{Fe^2+ > Fe^3+} \).

Ionisation Energy

Ionisation energy (IE) is the energy needed to remove the outermost electron from a gaseous atom: \( \ce{M(g) -> M+(g) + e^-} \). It is always positive and \( \text{IE}_1 \lt \text{IE}_2 \lt \text{IE}_3 \). It increases across a period and decreases down a group. The exceptions in period 2 and 3: Be is higher than B (filled 2s), N is higher than O (half-filled 2p), and Mg is higher than Al, P is higher than S. Ordering for period 2: Li < B < Be < C < O < N < F < Ne.

A sudden large jump between successive IEs shows the number of valence electrons: a jump between IE1 and IE2 means one valence electron (group 1); between IE2 and IE3 means two (group 2); and so on. The most stable oxidation state of an element follows the same jump. For one-electron species IE = 13.6 \( Z^2 \) eV, so He+ (54.4 eV) is higher than H (13.6 eV). Cr and Cu have a high IE2 because the cation is 3d5 or 3d10.

Electron Gain Enthalpy (Electron Affinity)

It is the enthalpy change when a gaseous atom takes an electron: \( \ce{X(g) + e^- -> X^-(g)} \). It is negative (exothermic) for most atoms' first addition, and positive for the second addition (O to \( \ce{O^2-} \): -141 then +780 kJ/mol). Elements with filled or half-filled subshells (Be, Mg, N, noble gases) have values near zero or positive. In halogens the order of magnitude is Cl > F > Br > I; in group 16, S > Se > O. Second-period elements O and F have lower EA than S and Cl because their small 2p subshells have strong electron repulsion. The order of EA magnitude, with sign ignored, for F, Cl, O, S is O < S < F < Cl.

Electronegativity

Electronegativity is the tendency of an atom in a bond to attract the shared electrons. It is measured on the Pauling scale (from bond energies: F = 4.0, O = 3.5, Cl = 3.0, N = 3.0, C = 2.5, H = 2.1, Cs = 0.7), the Mulliken scale, \( X = \dfrac{IE + EA}{2} \), and the Allred-Rochow scale. It increases across a period and decreases down a group. A bond between atoms of equal EN (for example P and H) is nonpolar. For an A-B bond, the percentage of ionic character from Hannay-Smith is \( 16\,\Delta X + 3.5\,\Delta X^2 \); for \( \Delta X = 2 \) this gives 46 percent. In a binary compound the less electronegative element is named first (iodine chloride, ICl).

Chemical Character and Periodic Trends

1. Metallic character and basicity of oxides rise down a group and fall across a period. The oxide of the most electropositive elements (K2O, Na2O) is strongly basic; MgO is basic, Al2O3, BeO, ZnO, SnO and PbO are amphoteric; SiO2, P4O10, SO3 and Cl2O7 are acidic.
2. The acidic strength of hydrides increases down a group (H2S < H2Se < H2Te; HF < HCl < HBr < HI) and across a period (CH4 < NH3 < H2O < HF).
3. Diagonal relationships: Li-Mg, Be-Al, B-Si.
4. The strongest reducing agent in a group is the element with the lowest IE (Cs among stable alkali metals).
5. The size of Ga, In, Tl in group 13 does not rise smoothly: B < Ga < Al < In < Tl.

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