JEE Chemistry Classification Of Elements and Periodicity In Properties Notes Set 04

Development of the Periodic Table

1. Dobereiner grouped elements in triads, where the atomic weight of the middle member is about the mean of the other two (Ca, Sr, Ba and Cl, Br, I). Newlands noted that every eighth element repeats properties (law of octaves). De Chancourtois arranged elements on a cylinder as the telluric helix.
2. Mendeleev (1869) arranged elements by increasing atomic weight into 7 periods and 8 groups (plus a zero group later), leaving gaps for undiscovered elements. Some pairs (such as Ar and K, Te and I) were in reverse order of atomic weight.
3. Moseley's X-ray work showed that atomic number, not atomic mass, is the fundamental property. Modern periodic law: the properties of elements are periodic functions of their atomic numbers.
4. Using atomic number removes the anomalies: isotopes share one position, and the order of Ar/K and Te/I becomes correct.

Long Form of the Periodic Table

1. The long form has 7 periods and 18 groups. Period lengths are 2, 8, 8, 18, 18, 32 and an incomplete 7th period. The lanthanides (Ce to Lu, 4f filling) and actinides (Th to Lr, 5f filling) are placed in a separate block below the table.
2. s-block: groups 1 and 2 (ns^1-2). p-block: groups 13 to 18 (ns²np^1-6). d-block: groups 3 to 12 [(n-1)d^1-10ns^0-2]. f-block: inner transition elements [(n-2)f^1-14(n-1)d^0-1ns²].
3. Group 1 are alkali metals, group 2 alkaline earth metals, group 16 chalcogens, group 17 halogens, group 18 noble gases. The 6 noble gases are He, Ne, Ar, Kr, Xe and Rn.
4. To place an element: the period equals the highest principal quantum number n; for s-block the group is the number of valence electrons, for p-block it is 10 + number of valence electrons, for d-block it is (n-1)d electrons + ns electrons.
5. Example: 1s²2s²2p⁶3s²3p⁴ has n = 3, 6 valence electrons: period 3, group 16, p-block (sulphur).

Diagonal Relationship and Special Positions

1. Diagonal pairs of the 2nd and 3rd periods resemble each other because their charge/size ratio (polarising power) is similar: Li-Mg, Be-Al, B-Si.
2. Hydrogen is placed in group 1 because it has one valence electron and forms H⁺, but it also forms H^- and is diatomic like halogens, so its position is special.
3. Lanthanides all have the same number of valence electrons and similar chemistry, so they share one place in group 3. The same holds for actinides.
4. Elements beyond 100 are named by IUPAC roots: 0 nil, 1 un, 2 bi, 3 tri, 4 quad, 5 pent, 6 hex, 7 sept, 8 oct, 9 enn, with the ending ium (for example, 104 is unnilquadium).

Atomic and Ionic Radii

1. Covalent radius is half the internuclear distance of a single bond between like atoms. Van der Waals radius is half the distance between nearest non-bonded atoms of neighbouring molecules and is always larger than the covalent radius. Noble gases show only van der Waals radii.
2. Across a period the radius decreases because the nuclear charge rises while the electrons enter the same shell (effective nuclear charge Z* = Z - screening constant increases). Down a group it increases because a new shell is added.
3. A cation is smaller than its atom (fewer electrons, higher Z/e ratio, sometimes loss of a whole shell). An anion is larger than its atom. Greater positive charge means a smaller cation: Fe > Fe²⁺ > Fe³⁺.
4. For isoelectronic species the size decreases as the nuclear charge increases: N^3- > O^2- > F^- > Ne > Na⁺ > Mg²⁺ > Al³⁺.
5. Lanthanide contraction: the steady decrease in size from La to Lu, caused by poor shielding of 4f electrons. It makes Zr and Hf almost equal in size.

Ionisation Energy

1. Ionisation energy (IE) is the energy needed to remove the outermost electron from an isolated gaseous atom: IE₃ > IE₂ > IE₁ because the same nuclear charge holds fewer electrons.
2. IE falls down a group and generally rises across a period. Exceptions: Be > B and Mg > Al (filled s subshell is stable); N > O and P > S (half-filled p subshell is stable).
3. Factors: larger atomic size lowers IE; larger nuclear charge raises IE; more screening lowers IE; penetration order s > p > d > f means s electrons are harder to remove; filled and half-filled subshells raise IE.
4. A large jump between successive IEs reveals the number of valence electrons. A big jump after IE₁ means one valence electron (group 1), after IE₂ means two (group 2).
5. Noble gases have the highest IE in each period. Order for period 3 (kJ/mol): Na 496 < Al 577 < Mg 738 < Si 787 < S 1000 < P 1012 < Cl 1251 < Ar 1521.

Electron Affinity

1. Electron affinity is the energy released when a gaseous atom gains an electron. It is largest for halogens. Second electron affinity is always endothermic because of repulsion, as in O^- + e^- going to O^2-.
2. Chlorine has a higher electron affinity than fluorine because the compact 2p subshell of fluorine has strong inter-electronic repulsion. Order: Cl > F > Br > I. Similarly S > O and Se > Te for the other families.
3. Noble gases, Be, Mg and N have zero or nearly zero (positive) electron affinity because of full or half-filled subshells.
4. Across a period electron affinity generally increases; down a group it generally decreases, apart from the second-period anomaly.

Electronegativity and Other Trends

1. Electronegativity is the ability of a bonded atom to attract the shared electron pair. On the Pauling scale F = 4.0 is the highest and Cs/Fr (0.7) the lowest. It increases across a period and decreases down a group.
2. Mulliken: EN = (IE + EA)/2 (in suitable units). A large electronegativity difference gives an ionic bond, a small one gives a covalent bond. Fajans: small cation, large anion and high charge increase covalent character (LiCl < BeCl₂).
3. Metallic character decreases across a period and increases down a group; non-metallic character shows the reverse. Oxides change from basic (Na₂O, MgO) through amphoteric (Al₂O₃) to acidic (SO₂, Cl₂O₇) across a period.
4. Valency with respect to hydrogen goes 1, 2, 3, 4, 3, 2, 1 across the period (groups 1 to 17) and valency with respect to oxygen increases from 1 to 7. Oxidation state of an atom depends on its valence electrons.
5. Typical properties: the most electropositive stable element is Cs; the smallest atom is He; the highest density element is Os/Ir; Hg is the only metal liquid at room temperature (apart from Ga, Cs near room temperature).