CBSE Class 11 Chemistry Classification Of Elements And Periodicity In Properties Of Elements Notes Set 07

These Class 11 CBSE Chemistry notes for Chapter 3 take up the story after the history of the periodic table and show how the long form follows from electronic configuration. They explain how to read the period, group and block of an element, describe the s-, p-, d- and f-blocks, metals, non-metals and metalloids, and give the IUPAC rules for naming elements above atomic number 100. The second half deals with covalent, metallic and van der Waals radii, and with the trends in atomic and ionic size. Eight solved examples show the working step by step.

Important Notes and Formulas

Moseley and the Basis of the Modern Periodic Law

1. In 1913 Henry Moseley studied the characteristic X-rays given out by different elements when they are bombarded with high-speed electrons.
2. He found that the square root of the frequency of a given X-ray line increases regularly with the atomic number of the element: \( \sqrt{\nu} = a(Z - b) \), where a and b are constants. A graph of \( \sqrt{\nu} \) against Z is a straight line, but a graph against atomic mass is not.
3. Atomic number (the number of protons, i.e. the nuclear charge) is therefore the more fundamental property, and Mendeleev's law was modified into the modern periodic law.
4. This removed the cases in which an element of higher atomic mass had to be placed before one of lower mass (Ar and K, Te and I) and gave isotopes a single position.

Electronic Configuration and the Structure of the Long Form

The long form is a direct picture of the aufbau principle. Electrons fill the orbitals in the order 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p. Every period begins with the filling of an ns orbital and ends with a completed np sub-shell (ns²np⁶), so the length of a period is set by the orbitals that fill within it.

How to Find the Period, Block and Group of an Element

1. Period = the highest principal quantum number n of the occupied shells.
2. Block = the type of sub-shell (s, p, d or f) that receives the last electron.
3. Group (s-block): the number of electrons in the ns orbital (1 or 2).
4. Group (p-block): 10 + the number of electrons in the outer ns and np orbitals (this gives groups 13 to 18).
5. Group (d-block): the number of ns electrons + the number of (n−1)d electrons (this gives groups 3 to 12).
6. f-block: these elements are accommodated in group 3 and shown separately at the bottom.

The s-, p-, d- and f-Blocks

1. Group 1 elements (except hydrogen) are the alkali metals (ns¹) and group 2 elements are the alkaline earth metals (ns²). Group 16 elements are the chalcogens, group 17 the halogens and group 18 the noble gases.
2. The p-block ends with the noble gases, ns²np⁶. Helium (1s²) is an s-block element but is placed in group 18 because its valence shell is completely filled and it behaves like the other noble gases.
3. Hydrogen (1s¹) is placed in group 1, although it is a non-metallic gas and not an alkali metal.
4. The 3d, 4d, 5d and 6d series have 10 elements each, and the f-block has two rows of 14 elements, the 4f and 5f series.

Types of Elements Based on Electronic Configuration

1. Main group (representative) elements: the s- and p-block elements of groups 1, 2 and 13 to 17. Their outermost shell is incomplete and the inner shells are complete.
2. Noble gases: group 18, with the outer configuration ns²np⁶ (He is 1s²). A completely filled valence shell makes them very unreactive and gives them zero valency.
3. Transition elements: the d-block elements of groups 3 to 12. They are all metals and show variable valence and oxidation states.
4. Inner transition elements: the lanthanoids (Z = 58 to 71, 4f series) and actinoids (Z = 90 to 103, 5f series). Their common oxidation state is +3 and all actinoids are radioactive. The elements beyond uranium (Z > 92) are called transuranium elements and are man-made.

Metals, Non-metals and Metalloids

1. More than 78% of the known elements are metals. They occupy the left and centre of the table: all s-, d- and f-block elements and the heavier members of the p-block.
2. Metals are usually solids with a metallic lustre, are malleable and ductile, and are good conductors of heat and electricity. Mercury is the only metal that is liquid at room temperature; gallium and caesium melt only slightly above it.
3. Non-metals are fewer than 20 in number and lie in the upper right-hand corner of the table. At room temperature they are gases (such as H₂, N₂, O₂ and the noble gases), brittle solids (such as C, S and P) or, in the case of bromine, a liquid. Most are poor conductors of heat and electricity.
4. Metalloids (semi-metals) lie along the zig-zag line that separates metals from non-metals. B, Si, Ge, As, Sb and Te are the usual examples; they show properties in between the two classes. Silicon and germanium are semiconductors.
5. The metallic character of an element decreases across a period and increases down a group, so the staircase line slants down to the right in the p-block.

IUPAC Nomenclature of Elements with Atomic Number Above 100

Newly made elements were once named by their discoverers and the name was approved by IUPAC. To avoid disputes, IUPAC adopted a systematic temporary name that is built directly from the atomic number.

1. Write the roots for the digits of the atomic number in order and end the name with -ium.
2. The symbol is made from the first letters of the roots, written as one capital letter followed by lower-case letters.
3. Example: Z = 120 is un + bi + nil + ium = unbinilium, symbol Ubn. Z = 118 is un + un + oct + ium = ununoctium (Uuo), which now carries the official name oganesson (Og).

Elements 111 to 118 have also received official names: roentgenium (Rg), copernicium (Cn), nihonium (Nh), flerovium (Fl), moscovium (Mc), livermorium (Lv), tennessine (Ts) and oganesson (Og).

Atomic Radius

1. The size of an isolated atom cannot be measured exactly, because the electron cloud has no sharp boundary and the outer electrons have a small chance of being found far from the nucleus. The radius is therefore found from the distance between the nuclei of bonded atoms. It is expressed in picometres (1 pm = 10⁻¹² m).
2. Covalent radius: half the distance between the nuclei of two identical atoms joined by a single covalent bond. For Cl₂ the Cl–Cl distance is 198 pm, so the covalent radius of Cl is 99 pm. It is mainly used for non-metals.
3. Metallic radius: half the distance between the nuclei of two adjacent metal atoms in the close-packed metallic lattice. For sodium it is 186 pm.
4. van der Waals radius: half the distance between the nuclei of two similar atoms of neighbouring molecules that touch each other in the solid state. For chlorine it is 180 pm.
5. For the same element: van der Waals radius > metallic radius > covalent radius. The van der Waals radius is the largest because the atoms are not bonded and are held apart by weak forces only.
6. Noble gases form no covalent bonds, so only van der Waals radii can be given for them. This is why their radii (for example Ne 160 pm) look much larger than the covalent radius of the preceding halogen (F 64 pm).

(All values are covalent radii except those of the alkali metals Li to Cs, which are metallic radii.)

1. In a period the radius falls from left to right. The outer electrons enter the same shell, while the nuclear charge increases by one at each step. The effective nuclear charge rises and pulls the electrons closer.
2. In a group the radius rises from top to bottom. The principal quantum number n increases and a new shell is added each time. The extra inner electrons shield the valence electrons, and this effect outweighs the increase in nuclear charge.

Remember: the three factors are the number of shells, the nuclear charge and the shielding effect of the inner electrons.

Ionic Radius

1. The radius of an ion is the effective distance from its nucleus up to which it influences the electron cloud. The trends follow those of atomic radii: ionic radii rise down a group and fall across a period for ions of the same charge.
2. A cation is smaller than its parent atom because it has fewer electrons for the same nuclear charge (often a whole shell is lost), so the effective nuclear charge per electron rises. Na = 186 pm and Na⁺ = 95 pm.
3. An anion is larger than its parent atom because the added electrons increase repulsion while the nuclear charge is unchanged. F = 64 pm and F⁻ = 136 pm.
4. For one element, the radius falls as the positive charge rises: M > M⁺ > M²⁺ (for example Fe > Fe²⁺ > Fe³⁺).
5. Isoelectronic species have the same number of electrons but different nuclear charges. The ion with the higher nuclear charge has the smaller radius, because its electrons are pulled in more strongly. The cation with the greater positive charge is the smallest, and the anion with the greater negative charge is the largest.

Ionic radius decreases from N³⁻ to Al³⁺ as Z ÷ e increases: N³⁻ > O²⁻ > F⁻ > Ne > Na⁺ > Mg²⁺ > Al³⁺. The 18-electron series P³⁻, S²⁻, Cl⁻, Ar, K⁺, Ca²⁺, Sc³⁺ behaves in the same way.

Solved Examples

Example 1. On the basis of quantum numbers, justify that the sixth period of the periodic table should have 32 elements.
Answer:
The sixth period corresponds to n = 6. The orbitals filled in it are 6s (one orbital), 4f (seven orbitals), 5d (five orbitals) and 6p (three orbitals).
Total number of orbitals = 1 + 7 + 5 + 3 = 16. Each orbital holds two electrons, so 16 × 2 = 32 electrons can be accommodated.
Each added electron gives one new element, so the sixth period has 32 elements.

Example 2. What are isoelectronic species? Name the species that are isoelectronic with (i) F⁻, (ii) Ar, (iii) Mg²⁺ and (iv) Rb⁺.
Answer:
Isoelectronic species are atoms or ions with the same number of electrons but different nuclear charges.
(i) F⁻ has 9 + 1 = 10 electrons: N³⁻, O²⁻, Ne, Na⁺, Mg²⁺ and Al³⁺ are isoelectronic with it.
(ii) Ar has 18 electrons: P³⁻, S²⁻, Cl⁻, K⁺ and Ca²⁺.
(iii) Mg²⁺ has 12 − 2 = 10 electrons: the same set as in (i), N³⁻, O²⁻, F⁻, Ne, Na⁺ and Al³⁺.
(iv) Rb⁺ has 37 − 1 = 36 electrons: Br⁻, Kr and Sr²⁺.

Example 3. To which block does the element of atomic number 50 belong? Find the period, group and block of the element of atomic number 43.
Answer:
Z = 50: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ 4d¹⁰ 5s² 5p². The last electron enters a 5p orbital, so it is a p-block element (tin) in period 5, group 10 + 2 + 2 = 14.
Z = 43: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ 4d⁵ 5s². The last electron enters a 4d orbital, so it is a d-block element (technetium). The highest n is 5, so it is in period 5, and the group number is 5 + 2 = 7.

Example 4. Arrange the following in increasing order of size: Ca²⁺, S²⁻, P³⁻, K⁺.
Answer:
All four ions have 18 electrons (isoelectronic). Their nuclear charges are P +15, S +16, K +19 and Ca +20.
The greater the nuclear charge, the smaller the ion. So the order of increasing size is Ca²⁺ < K⁺ < S²⁻ < P³⁻.

Example 5. Cl⁻ and K⁺ are isoelectronic. Why is the ionic size of Cl⁻ greater than that of K⁺?
Answer:
Both ions have 18 electrons. The nuclear charge of K⁺ is +19, whereas that of Cl⁻ is +17.
The stronger pull of the nucleus of K⁺ draws the same 18 electrons closer, so K⁺ is smaller. In Cl⁻ the nucleus has two fewer protons, so the electron cloud is more spread out and Cl⁻ is larger.

Example 6. Which species will have the largest and the smallest size among Mg, Mg²⁺, Al and Al³⁺?
Answer:
Mg (Z = 12) and Al (Z = 13) are in the same period, and the atomic size falls from left to right. So Mg is larger than Al.
A cation is smaller than its parent atom, so Mg²⁺ < Mg and Al³⁺ < Al. Al³⁺ has a larger nuclear charge than Mg²⁺ for the same 10 electrons, so Al³⁺ is smaller than Mg²⁺.
Order: Mg > Al > Mg²⁺ > Al³⁺. Mg is the largest and Al³⁺ is the smallest.

Example 7. The outer electronic configurations of some elements are (a) 3s² 3p⁴, (b) 3d¹⁰ 4s², (c) 3s² 3p⁶ 4s² and (d) 6s² 4f³. To which block does each belong?
Answer:
(a) The last electron enters 3p, so the element (S) is in the p-block.
(b) The last electron enters 3d (Zn, group 12), so it is in the d-block.
(c) The last electrons enter 4s (Ca), so it is in the s-block.
(d) The last electron enters 4f (Pr), so it is in the f-block.

Example 8. The Cl–Cl distance in a chlorine molecule is 198 pm and the van der Waals radius of chlorine is 180 pm. Find (a) the covalent radius of chlorine and (b) the shortest distance between two chlorine atoms of neighbouring molecules in solid chlorine.
Answer:
(a) Covalent radius = half of the bond length = 198 ÷ 2 = 99 pm.
(b) The atoms of neighbouring molecules touch each other at the van der Waals radii, so the distance = 2 × 180 pm = 360 pm.
This is much longer than the bonded distance of 198 pm, which shows that the van der Waals radius is larger than the covalent radius.