Periodic Properties

Periodic Laws

Periodic table is a table of elements in which the elements with similar properties are placed together. In order to arrange the elements, various classifications were made by chemists. The classifications were based on different properties of elements.

Dobereiner’s Law of Triads 

The elements with similar properties were arranged in groups of three where the atomic mass of the central element was the arithmetic mean of the atomic mass of the other two elements.

e.g. Li Na K and Cl Br I

7 23 39 35.5 80 127

Newland’s Law of Octaves

According to this law, when the elements were arranged in order of their increasing atomic mass, every eighth element had properties similar to those of the first just like the eighth note of an octave of music. This law worked only for lighter elements.

Mendeleev’s Periodic Table

It is based on Mendeleev’s periodic law according to which physical and chemical properties of elements are periodic function of their atomic masses

Structural features of the Mendeleev’s periodic table:

1. Originally, it consisted of eight vertical columns called groups (IVIII) and seven horizontal rows called periods (1 – 7). With the discovery of Noble gases, the table was modified and another vertical column in the form of zero group was added.
2. All groups except VIII (and zero) have been further subdivided into two subgroups called A and B.
3. Few elements like gallium and germanium were not discovered at the time when Mendeleev proposed his periodic table. Mendeleev named these elements as eka-aluminum and eka-silicon because he believed that they would be similar to aluminium and silicon respectively.

MODERN PERIODIC TABLE 

According to modern periodic law, physical and chemical properties of the elements are periodic function of their atomic number.

In modern periodic table, similar properties reoccur after the interval of 2, 8, 8, 18, 18 and 32 elements, if arranged in increasing atomic number. These numbers are called magic numbers.

1. Characteristics of long form of the periodic table.

(i) The table is divided into four portions called blocks.

These blocks are named on the basis of the valence orbital, i.e. orbital in which last electron enters.

(ii) It has 7 horizontal rows (periods) and 18 vertical columns (groups) consisting 112 elements.

(iii) s and p block elements are called representative elements, d-block elements are called transition elements and f-block elements are called inner-transition elements.

2. Diagonal relationship: Certain elements of IInd period show similarity with their diagonal elements in the 3rd period as shown below.

Thus, Li resembles Mg, Be resembles Al and B resembles Si. This happens because of almost same charge/size ratio of these elements.

Ex.: Alkali metals belong to

(A) s-block

(B) b-block

(C) d-block

(D) f-block

Solution: (A).

Differentiating electron of alkali metals goes to s-subshell so are called as s-block elements.

Ex.: The 3^rd period of the periodic table contains

(A) 8 elements

(B) 32 elements

(C) 3 elements

(D) 18 elements

Solution: (A). 

3d is filled after 4s so 3^rd period is having eight elements only.

3. Periodic properties: Properties which show a regular gradation when we move from left to right in a period or from top to bottom in a group. These are the properties, which directly or indirectly, related to the electronic configuration of the elements.

(i) Atomic radius: It refers to the distance between atomic nucleus and outermost shell of electrons of an atom. Since the absolute value of the atomic size cannot be determined, it is usually expressed in terms of the following operational definitions.

(a) Covalent radius: One half of distance between the nuclei of two covalently bonded atoms of same element in a molecule is called covalent radius.

(b) Metallic radius: It is defined as one half of the distance between the centres of nuclei of the two adjacent atoms in the metallic crystal. Metallic radius of an element is always greater than its covalent radius.

(c) van der Waal’s radius: It is defined as one half of the distance between the nuclei of two non-bonded isolated atoms or two adjacent atoms belonging to two neighbouring molecule of an element in the solid state. van der Waal’s radius is larger than covalent radius.

(ii) Ionic radius: An atom can be changed to a cation by loss of electrons and to an anion by gain of electrons. A cation is always smaller than the parent atom because during its formation nuclear charge increases and sometimes a shell may also decrease. On the other hand, the size of an anion is always larger than the parent atom because during its formation effective nuclear charge decreases.

(iii) Isoelectronic species: The atoms or ions which consist of same number of electrons but differ in nuclear charge are called isoelectronic species. The size of the isoelectronic species depends upon their nuclear charge. Greater the nuclear charge smaller the size.

Periodicity in Size

(i) Moving across the period, the electrons are filled in the same shell. So, atomic size decreases as nuclear charge increases.

(ii) In a group, number of shells in an atom increases so atomic size increases.

(iii) The atomic radius of inert gases are largest in a period.

(iv) In case of transition elements, moving left to right initially atomic size decreases (as expected) and then starts increasing (due to shielding effect).

Ex.: Which among the following is largest in size?

(A) Cl–

(B) S2–

(C) Na+

(D) F–

Solution: (B). 

Among isoelectronic species, higher is the anionic charge larger is the size. 

Ionization Energy or Ionization potenitla

The amount of energy required to remove the most loosely bound electron from one mole of an isolated gaseous atom is known as ionization energy. The amount of energies required to remove the first, second, third etc. electrons, from the isolated gaseous atom are called successive ionization energies and are designated as, IE1, IE2, IE3, etc. 

It is measured in electron volts or k. cal / mole. The factors affecting the ionization energies are:

(i) Atomic radius

(ii) Effective nuclear charge

(iii) Penetration effect of the electrons,

(iv) Shielding effect

(v) Effect of exactly half-filled and completely filled orbitals.

Periodicity in Ionization Potential

(i) For normal elements, on moving from left to right in a period, generally ionization potential increases because the outermost shell will be the same and effective nuclear charge increase while on moving down in a group, it decreases, the outermost shell will become far away from nucleus.

(ii) In a period, ionization energies of 2nd group elements are more than the 3rd group elements due to higher penetration effect or fully filled s-orbital. Similarly, for group 5 elements, I.P. value is more than group 6 elements because of half-filled configuration (in group 5 elements) which is more stable.

(iii) In a group, I.P. of III and VI period elements are higher than II and V period elements. Although, there is increase in number of shells but shielding of nucleus is less as electrons are filled in 3d and 4f shells. 

Ex.: Which of the following statement concerning ionization energy is not correct?

(A) The I. E2. is always more than the first.

(B) Within a group, there is a gradual increase in ionization energy because nuclear charge increases.

(C Ionization energies of Be is more than B.

(D)Ionization energies of noble gases are high

Solution: (B).

In a group size is increase and effective nuclear charge is not much as size thus IE decrease down the group.

Ex.: In which of the following pairs, the ionization energy of the first species is less than the second? 

(A) N, P

(B) Be2+, Be

(C) N, N–

(D) S, P

Solution: (D).

p has half filled orbital.

ELECTRON AFFINITY

The amount of energy released when an electron is added to an isolated gaseous atom to produce monovalent anion is called electron affinity.

Similarly, second and third electrons can be added to form gaseous binegative and tri negative ions. The electron affinity of an element depends upon, (i) atomic size, (ii) effective nuclear charge and (iii) electronic configuration.

First electron affinity is exothermic in nature since a neutral atom has a natural tendency to accept an electron whereas second electron affinity is endothermic in nature because electron has to be added to a negatively charged ion. It is measured in electron volts or k/ cal per mole. Factors responsible to influence magnitude of electron affinity are

(i) Atomic size 

(ii) Effective nuclear charge 

(iii) Screening effect 

Periodicity in Electron Affinity

(i) Across the period, E.A. increases from left to right as the effective nuclear charge increases but E.A. of inert gases are taken to be zero as they do not have tendency to accept electrons due to the increase of atomic size.

(ii) Down the group, E.A. generally decreases. But E.A. of IIIrd period elements is higher than IInd period elements because of very small size of IInd period elements due to which electron-electron repulsions are appreciable. While IIIrd period onwards, size is large enough to easily accommodate incoming electrons.

Illustration 7: Which of the elements have strongest tendency to form anions?

(A) No, Cl, Al

(B) Cu, Ag, Au

(C) Be, F, N

(D) F, Cl, Br

Solution: (D).

Tendency to form anion depends upon electron affinity value which is highest for halogens in a period. 

Electronegativity

The tendency of an atom in a molecule to attract shared pair of electrons towards itself is called electronegativity. It is a relative phenomenon. Its value can be calculated by theoretical equations. The standard value of electronegativity for fluorine is 4. (which is highest). It depends upon 

(i) atomic size

(ii) nuclear charge

(iii) oxidation state

(iv) hybridization state and 

(v) electronic configuration.

Periodicity of Electronegativity

(i) Across the period, it increases on moving from left to right as nuclear charge increases.

(ii) Down the group, it decreases as atomic size increases.

Applications of Electronegativity

(i) It helps to predict the polarity of bonds and dipole moment of molecules.

(ii) It gives an idea about the atomic size. Higher the electronegativity, smaller is the atomic size.

(iii) It helps to have an idea about the bond length. Higher the electronegativity difference, smaller is the bond length.

Scales of Electronegativity: 

(i) Poling scales 

X_A – X_B = 0.208

Where 

X_A = E.N. of atom A 

X_B = E.N. of atom B

E_A-B = bond dissociation energy of molecule A – B 

E_A-A = bond dissociation energy of molecule A – A

E_B-B = bond dissociation energy of molecule B – B

(ii) Mulliken Scale 

when IP and EA values are expressed in ev

when IP and EA values are expressed in K. cal

Example 8 : Which of the following does not have any unit?

(A) Electronegativity

(B) Electron affinity

(C) Ionization potential

(D) Atomic radii

Solution: (A).

Electronegativity is the relative term with out any units. 

Other Porperties

Atomic Volume

Atomic volume may be defined as the volume occupied by one mole atom of the element at its melting point in solid state.

Periodicity of Atomic Volume

(i) In a period, from left to right atomic volume decreases initially, then becomes minimum in middle and finally increases. This is due to different packing arrangements of atoms in different elements in the solid state. For example, P, S etc. 

(ii) On moving down the group, atomic volume increases gradually as atomic size increases.

Density 

(i) In a period, from left to right density increases first and then decreases. 

(ii) In a group, density increases as we move down.

Although, atomic size increases down the group but rise in atomic weight is more than the increase in atomic volume.

Melting and Boiling Points

These properties vary differently in different blocks.

(i) In s-block, M.Pt. and B.Pt. decreases down the group.

(ii) In d-block, M.Pt. and B.Pt. increases down the group.

(iii) In p-block,M.Pt. and B.Pt. decreases down the group for group 13 and 14 but it increases for group 15 to 18.

[Note: Tungsten (W) and mercury (Hg) has the maximum and minimum boiling points among metals while in non-metals carbon has maximum melting point and helium has minimum melting point].

Nature of Oxides and Hydroxides

The oxide or hydroxide of an element may act either as a base or as an acid depending upon its ionization energy. If the ionization energy is low, it acts as a base and if the ionization energy is high, it acts as an acid. The ionization energies of alkali metals are lowest, therefore, their oxides and hydroxides are the strongest bases. Further, the ionization energies of alkali metals decreases down the group, the basic character of their hydroxides increases in the same order: CsOH > RbOH > KOH > NaOH > LiOH. 

The ionization energies of halogens are quite high, therefore, their oxides are the strongest acids. Further, the ionization energies decrease down the group, therefore, the acidic character of their oxides and hydroxides decreases in the same order: HClO₄ > HBrO₄ > HIO₄.

As we move from left to right within a period, the ionization energies of the elements usually increase and hence their oxides and hydroxides show a gradual variation from strongly basic through amphoteric to strongly acidic character.

For example:

Ex.: Which of the following element represents highly electropositive as well as highly electronegative element in its period?

(A) Nitrogen

(B) Fluorine

(C)Hydrogen

(D) None of the above

Solution: (C).

In 1st period only H and He are present and He is noble gas. 

Ex.: Which of the following halides is the most acidic? 

(A) PCl5

(B) SbCl3 

(C) BiCl3

(D) CCl4

Solution: (C).

Acidic nature is explained by the deficiency of electron and BiCl3 is more electron deficient among the following. 

Points to Remeber

1. The elements showing horizontal relationship (due to shielding effect) and vertical relationship (due to same general configuration) are transition elements. 

2. Elements with atomic number 110 to 112 are still to be studied and named properly. 

3. ₄₃Tc and ₈₆Rn are man-made elements along with transuranic elements. 

4. Fluorine has highest electronegativity in periodic table. 

5. Cl, S and P have high electron affinity than F, O and N respectively. 

6. Ionization energy of 2^nd and 15^th group elements are exceptionally high due to fully filled and half filled configuration. 

7. In case of transition elements, the atomic and ionic radii first decrease, till d-subshell is exactly half-filled and thereafter these nearly remain the same since the increase in size due to increasing shielding effect of inner d-electrons almost cancels the contractive effect of the increased nuclear charge. 

8. For isoelectronic ions, the size depends upon the nuclear charge. Higher the nuclear charge, similar is the size. 

9. Metalloids have electronegativity values close to 2.0 and the differentiation value of electronegativity between metals and non-metals is 2.1. 

10. For the same element, the electronegativity depends upon the state of hybridization. For example, electronegativity of C in the three states of hybridization varies as: sp > sp² > sp³. As the s-character of hybrid orbitals decreases, the electronegativity also decreases. 

11. The first attempt to classify all the known elements was made by Newland in 1864 and the law given by him known as law of octave.

12. All s-block elements (except H₂ and He) are soft metals. 

13. Effective nuclear charge (Z_eff) is given by Z_eff = Z – S, where Z is nuclear charge and S is screening constant (determined with the help of Slater rules). 

14. The correct decreasing order of electron affinities is Cl > F > Br > I > S > Si. 

15. Ionization energy for completely filled orbitals > half filled orbitals > partly filled orbitals. 

SOLVED EXAMPLES

1. In which of the following binary compound the ratio, rcation/ranion is least?

(A) LiI

(B) CsI

(C) LiF

(D) CsF

Sol. (A).

I^- is the largest anion and Li+ is the smallest cation out of given. 

2. The I. E. of Al is smaller than that of Mg because

(A) atomic size of Al > Mg.

(B) atomic size of Al < Mg.

(C) I. E. in Al pertains to the removal of p-electron which is relatively easy. 

(D) unpredictable.

Sol. (C).

Penetration effect of s-subshell is more than p – subshell.

3. From the ground state electronic configurations of the elements given below, pick up the one with highest value of second ionization energies

(A) 1s², 2s², 2p⁶, 3s²

(B) 1s², 2s², 2p6, 3s¹

(C) 1s², 2s², 2p⁶

(D) 1s^₂, 2s², 2p⁵

Sol. (B).

After removal of one electron (B) attain noble gas configuration. 

4. Which of the following iso electronic ions lowest ionization enthalpy? 

(A) K+

(B) Ca²⁺ 

(C) Cl-

(D) S^2- 

Sol. (D).

S^2- has the largest size and hence the lowest I.E. 

5. Which of the following is incorrect? 

(A) An element which has high electronegativity always has high electron gain enthalpy

(B) Electron gain enthalpy is the property of an isolated atom 

(C) Electron negativity is the property of bonded atom 

(D) Both electronegativity and electron gain are usually directly related to nuclear charge and inversely related to atomic size. 

Sol. (A).

Elements with electronegativity usually but not always have high electron gain enthalpies. For example both N and Cl have an electronegativity of 3 but electrons gain enthalpy of N is zero while that f chlorine is the highest in the periodic table. 

6. The order of which the following oxides are arranged according to decreasing basic nature 

(A) Na^₂O, MgO, Al²O³, CuO

(B) CuO, Al²O³, MgO, Na²O 

(C) Al²O³, CuO, MgO, Na²O

(D) CuO, MgO, Na2O, Al²O³ 

Sol. (A).

As we move from left to right in a period, the basic character of the oxides of s- and p- block elements decreases while their acidic character increase. The basic character of oxides of transition elements is however lower than alkali and alkaline earth metal. Thus Na²O is most basic followed by MgO and Al_²O³ while CuO is least basic. 

7. The screening effect of d-electron is 

(A) equal to p-electron

(B) much more than p-electron 

(C) same as f-electrons

(D) less than p-electrons 

Sol. (D).

The order of screening effect is s > p > d > f 

8. Which of the following represents the electronic configuration of the most electropositive element? 

(A) [He]2s^₁

(B) [Xe] 6s¹ 

(C) [He]2s² 

(D) [Xe]6s²

Sol. (D).

Element (a) and (b) are alkali metals whereas (c) and (d) are alkaline earth metals. Amongst alkali and alkaline earth metals, alkali metals with the biggest size is the most electropositive element. 

9. The ionization of hydrogen atom would give rise to 

(A) hydride ion

(B) hydronium ion 

(C) proton

(D) hydroxyl ion 

Sol. (C). whereas (a) is H-, (b) H₃O+ and (d) is OH-. 

10. Which is true about the electronegativity order of the following elements? 

(A) P > Si

(B) C > N 

(C) Br > Cl

(D) Sr > Ca 

Sol. (A). Electronegativity of P(2.1) is greater than Si (1.8).