Periodicity

Reconstruct Your Chemistry With Prince Sir

Page 1

CHEMISTRY

ri

PERIODIC PROPERTIES OF ELEMENTS DOBEREINER’S TRIADS According to Dobereiner when elements of same properties are kept in the increasing order of atomic weights, the atomic weight of middle element is equal to the mean atomic weight of remaining two elements. Such a group of elements is called Dobereiner’s triad.

c ni

Mean of first and last element

Li

Na

7

23

Be 8 P 31

Cl

y B 35.5

r P Mg 24

As 75

Br 80

S e K

39

Ca 40

Sb

120

I 123

Dobereiner could arrange only a few elements as triads and there are some elements present in a triad, whose atomic weights are approximately equal Fe Co Ni Ru Rh Pd There fore, this hypothesis was not acceptable for all elements. NEWLAND’S RULE OF OCTAVE As in music, the eighth node is same as the first node. If the elements are arranged in the increasing order of atomic weights, on starting with an element, the first element will exhibit similarities with the eighth element e.g. Name of element Li 7 Name of element Na 23

Be 9 Mg 24

B 11 Al 27

C 12 Si 28

N 14 P 31

O 16 S 32

F 19 Cl 35.5

It is clear from the above table that sodium is the eighth element from lithium, whose properties resemble that of lithium. This type of classification was limited up to only 20 elements.

Page 2

Reconstruct Your Chemistry With Prince Sir MENDELEEF’S PERIODIC LAW “According to Mendeleef’s periodic law, the physical and chemical properties of elements are periodic function of their atomic weights.” MENDELEV’S PERIODIC TABLE Mendeleve’s Periodic table is based on atomic weight.  In the periodic table, the horizontal lines are called periods and the vertical lines are called groups.

ri

The periodic table consists of seven periods and nine groups (The earlier periodic table had only 8 groups. The noble gases were added later in the zero group because these were not discovered when Mendeleef put forward his periodic table.

S e

All the groups (except VIII and Zero groups) are divided into subgroups A and B. 2, 8, 18 and 32 are called magic numbers. ORIGINAL MENDELEEV”S PERIODIC TABLE

y B

r P

c ni

Page 3

Reconstruct Your Chemistry With Prince Sir



MERITS OF MENDELEEF’S PERIODIC TABLE Classification of elements then known, was done for the first time and the elements having similar properties were kept in the same group.



It encouraged research and led to discovery of newer elements.



Mendeleef had even predicted the properties of many elements not discovered at that time. This helped in the discovery of these elements. For example. Mendeleef predicted the properties of the following elements. (a) Eka-boron- This was later called scandium (Sc) (b) Eka-aluminium - This was later called gallium (Ga) (c) Eka-silicon - This was later called germanium (Ge)

c ni

S e

DEFECTS OF MENDELEEF’S PERIODIC TABLE

ri

(1)

Position of Hydrogen – Like alkali metals hydrogen has one electron in its outermost level and like halogens, it has one electron less than the next inert gas. Therefore, placing hydrogen with alkali metals is equally appropriate as placing it with halogens.

(2)

Position of Isotopes – The isotopes have different atomic weights and the periodic table is based on atomic weights. Therefore, isotopes should get different places in the periodic table on the basis of atomic weights.

(3)

The periodic table is not fully based on increasing order of atomic weights. For example :

y B

r P

Ar(39.94) was placed before Ca (39.098) Te (127.6) was placed before I(126.9) Co(58.93) was placed before Ni (58.69) (4)

It is not proper to place together the elements having differing properties, such as coinage metals (Cu, Ag and Au) with alkali metals; Zn, Cd and Hg with alkaline earth metals and metal like Mn with halogens. Similarly. Pt and Au having similar properties have been placed in different groups.

(5)

There is no indication whether lanthanides and actinides are associated with group IIIA or group IIIB.

(6)

Position of Isobars – These elements have different groups when mass remains same.

Page 4

Reconstruct Your Chemistry With Prince Sir MODERN PERIODIC LAW AND MODERN PERIODIC TABLE Mosley proved that the square root of frequency () of the rays, which are obtained from a metal on showering high velocity electrons is proportional to the atomic number of the atom. This can be represented by the following expression.

 = a (Z–b) where Z is nuclear charge on the atom and a and b are constants. The nuclear charge on an atom is equal to the atomic number.

ri

According to modern periodic law. “The properties of elements are the periodic functions of their atomic numbers”

S e

MODERN PERIODIC TABLE One the basic of the modern periodic law, Bohr proposed a long form of periodic table that was prepared by Rang and Warner.

c ni

In the periodic table the horizontal lines are periods and the vertical lines are groups.

r P

The periodic table has a total of seven periods and 18 groups. But the number of groups is 16, because the eighth group has been divided into three groups. There are two elements in the first period eight elements in each of the second and third periods, eighteen elements in each of the fourth and fifth period thirty two elements in the sixth period and only nineteen elements till now in the seventh period.

y B

The first period is very short period, second and third are short periods fourth and fifth are long periods sixth is very long period, while the seventh is incomplete period. The lanthanides (Elements from atomic numbers 58 and 71) and actinides (elements from atomic numbers 90 to 103) are included in the sixth and seventh groups through these have been kept outside the periodic table. Period - The details about the seven periods are as follows.

Reconstruct Your Chemistry With Prince Sir

Page 5

Period

Number of elements

First Second Third Fourth Fifth Sixth

Atomic number From to H (1) He (2) Li (3) Ne (10) Na (11) Ar (18) K (19) Kr (36) Rb (37) Xe (54) Cs (55) Rn (86)

Seventh

Fr (87)

19 (including actinides)

Ha (105)

2 8 8 18 18 32 (including lanthanides)

S e

ri

Group - The modern periodic table has 18 vertical columns and according to CAS system there are 16 groups having the following number of elements.

(a) (b) (c) (d) (e) (f) (g) (h) (i)

(j) (k) (l) (m) (n) (o) (p)

Group IA or group1 IIA or group2 IIIA or group13 IV A group14 V A group 15 VI A group16 VII A group17 Zero group18 III B group3

y B

r P

IV B group4 V B group5 VI B group6 VII B group7 VIII (3) group8,9,10 I B group11 II B group12

c ni

Number of Elements 7 (H, Li, Na, K, Rb, Cs Fr) Alkali metals 6 (Be, Mg, Ca , Sr, Ba, Ra) Alkaline earth metals 5 (B, Al , Ga , In, Tl) Boron family 5 (C, Si, Ge , Sn , Pb) Carbon family 5 (N, P , As, Sb, Bi) Nitrogen family 5 (O, S , Se, Te , Po) Oxygen family (chalcogen) 5 (F, Cl, Br, I, At) - Halogen family 6 (He , Ne , Ar, Kr, Xe, Rn) Inert elements 32(Sc, Y, La, Ac amd 14 lanthanide elements and 14 actinide elements.These are elements of IIIB group, which which could not be accommodated in one column and therefore written separately outside the periodic table. 4 (Ti, Zr, Hf, Ku) 4 (V, Nb, Ta, Ha) 3 (Cr, Mo, W) 3 (Mn, Tc, Re) 9 (Fe, Co, Ni, Ru, Rh, Pd, Os, Ir, Pt) 3 (Cu, Ag, Au) 3 (Zn, Cd, Hg)

Page 6

Reconstruct Your Chemistry With Prince Sir MERITS OF LONG FORM OF PERIODIC TABLE OVER MANDELEEF’S PERIODIC TABLE Positions of Isotopes and Isobars - Isotopes have same atomic number and the periodic table is based on atomic numbers. Therefore, various isotopes of the same elements have to be provided the same position in the periodic table. Isobars gave same atomic weights but different atomic numbers and therefore they have to be placed at different positions. The positions of actinides and lanthanides is more clear now because these have been placed in IIIB groups and due to paucity of space, these are written at the bottom of the periodic table.

ri

In the periodic table a diagonal line from B to At separates the metals, nonmetals and metalloids from one another. The elements like B, Si As. Te At, that are situated near this line are metalloids i.e. these have characteristic of both metals and nonmetals.

c ni

S e

The general electronic configurations of the elements remains same in group. DEFECTS OF LONG FORM OF PERIODIC TABLE The position of hydrogen is still disputable as it was there in Mendeleef periodic table in group I A as well as IVA & VIIA.

r P

Helium is an inert gas but its configuration is different from that of the other inert gas elements

y B

Lanthanide and actinide series could not be adjusted in the main periodic table and therefore they had to be provided with a place separately below the table.   

  

S-BLOCK ELEMENTS The elements of the periodic table in which the last electron enters in s– orbital, are called s–block elements. s-Orbital can accommodate a1maximum of two electrons. 2 Their general formulae are ns and ns respectively, where n = (1 to 7). I A group elements are known as alkali metals because they react with water to form alkali. II A group elements are known as alkaline earth metals because their oxides react with water to form alkali and these are found in the soil or earth. The total number of s block elements are 14. 87 88 Fr and Ra are radioactive elements while H and He are gaseous elements Cs and Fr are liquid elements belonging to s-block.

Page 7

 

Reconstruct Your Chemistry With Prince Sir P - BLOCK ELEMENTS The elements of the periodic table in which the last electron gets filled up in the p-orbital, called p-block elements. A p-orbital can accommodate a maximum of six electrons. Therefore, p-block elements are divided into six groups which are III A, IV A, V A, VI, A VII A and zero groups. 2

6

ri



The general formulae of p block elements are ns np



The zero group elements having general formula ns np are inert, because their energy levels are fully filled. The total number of p block elements in the periodic table is 30 (excluding He)



2

(where n = 2 to 6) 6

S e



There are nine gaseous elements (Ne, Ar, Kr, Xe, Rn, F 2, Cl2, O2 and N2) belonging to p-block. Gallium (Ga) and bromine (Br) are liquids.



The step-like thick lines drawn in the periodic table in the p-block divides elements into metals nonmetals and metalloids. B to At, which are present on this line are metalloids. D-BLOCK ELEMENTS The elements of the periodic table in which the last electron gets filled up in the d orbital, called d block elements.



y B

r P

c ni



The d block elements are placed in groups named IIIB, IV B, V B, VI B, VII B, VIII, I B and II B.



In d block elements the electron gets filled up in the d orbital of the penultimate shell. That is why these elements are known as transition elements.



Through the total number of d block elements is 33 in the periodic table but there are only 30 transition elements. Becuase only those elements are transition in which d orbital partilly filled.



2 6 1–10 1–2 The general formula of these elements is (n–1)s , p , d ns where n = 4 to 7.



All of these elements are metals



Out of all the d block elements mercury is the only liquid element.



Reconstruct Your Chemistry With Prince Sir F-BLOCK ELEMENTS The elements of the periodic table in which the last electron gets filled up in the f orbital, called f block elements. The f block elements are from atomic number 58 to 71 and from 90 to 103. The lanthanides occur in nature in low abundance and therefore, these are called rare earth elements. There are 28 f block elements in the periodic table. The elements from atomic number 58 to 71 are called lanthanides because comes after lanthanum (57). The elements from 90 to 103 are called actinides because comes after actinium (89). All the actinide elements are radioactive



All the elements after atomic number 82 (i.e., Pb ) are radioactive.



All the elements after atomic number 92 (i.e.U



The general formula of these elements is

Page 8

    

82

c ni

S e 92

2 6 10 (1–14) 2 6 0–1 2 (n–2)s p d f (n–1)s p d ns where n = 6 & 7. 

r P

ri

) are transuranic elements.

EFFECTIVE NUCLEAR CHARGE In a polyelectronic atom, the internal electrons repel the electrons of the outermost orbit. This results in decrease the nuclear attraction on the 1 R ep 2 electrons of the outermost orbit. 0 u l si e on –1 e0 + –1 Attraction Nucleus

y B



Therefore, only a part of the nuclear charge is effective on the electrons of the outermost orbit. Thus, the inner electrons protect or shield the nucleus and thereby decrease the effect of nuclear charge towards the electrons of the outermost orbit.



Thus the part of the nuclear charge works against outer electrons, is known as effective nuclear charge (Z*) = (Z - ) Where  is screening constant .

Shielding (Screening) effect : The repulsion of valency electrons by the electrons in peunltimate shell to reduce effective nuclear charge (thereby reducing ionization enthalpy value). Zeff. = Z –  (where Z = atomic number and  = shielding constant)

Reconstruct Your Chemistry With Prince Sir

Page 9

Calculation of Shielding Constant () In order to estimate extent of shielding, Slater proposed a set of empirical rules. These rules are simplified generalizations based upon average behaviour of various electrons. These calculations are useful in explaining atomic size and electronegativity. To calculate the shielding constant for an electron in the following rules are used :

ri

A.

Write the electronic configuration of element in the following order and grouping (1s), (2s, 3p), (3s, 3p) (3d), (4s, 4p) (4d) (4f), (5s, 5p) etc.

B.

Electrons in any group to the right of the (ns, np) group contribute nothing to shielding constant. All the electrons in the (ns, np) group shield the valence electron to the extent of 0.35 each except for the 1s (i.e., n= 1) for 1s s is taken as 0.3.

C.

D. E. F.

c ni

S e

All the electrons in the (n – 1) shell shield to an extent of 0.85 each. All the electrons in (n – 2) or lower shield completely and their contribution is 1.00 each.

r P

When electron being shielded is in an nd and nf group, rules (B) and (C) are same but rule (D) and (E) are changed. Here all the electrons in groups lying to the left of the nd or nf group contribute 1.00 each.

y B

Thus,  = [0.35 × Number of electrons in nth shell excluding test (valence) electron] + [0.85 × Number of electrons in (n – 1)th shell] + [1.0 × Number of electrons ininner shell]

e.g., (i)7N : Consider the valence electron in 2p orbital, then excluding it 4 electrons are present in (2s,2p) group. N : (1s)2 (2s, 2p)5 7 s = [0.35 × 4] + [0.85 × 2] = 3.10

(ii)

Zn : (1s)2, (2s, 2p)8, (3s, 3p)8, (3d)10, (4s)2

30

[Valence electron in 4s, then excluding it one electron is present in 4s group] s = [0.35 × 1] + [0.85 × 18] + [1.00 × 10] = 25.65

Reconstruct Your Chemistry With Prince Sir

Page 10

Cationic Radius An atom forms a cation on loss of electron/s. The cationic radius can be defined as the distance between the nucleus and the limit of the electron cloud scattered around the nucleus. The size of a cation is smaller in comparison to the size of its corresponding atom. This is because of the fact that an atom on losing electrons/s form a cation, which has lesser number of electrons/s than the number of proton/s. This results in increase in the effective nuclear charge. Size of cation Examples - (1) Mn > Mn (2) Pb

+2

> Pb

+2

> Mn

+3

> Mn

+4

+4

c ni

+6

> Mn

ri

+7

S e > Mn

Anionic Radius When a neutral atom gains electron/s it becomes a negatively charged ion called an anion. The distance between the nucleus of an anion and the limit of the electron cloud scattered around the nucleus, is called its anionic radius.

r P

The size of an anion is greater than the size of its corresponding atom, because the number of electrons present in the anion gets larger than the number of protons due to gain of electron/s. This results in decrease in the effective nuclear charge.

y B

0 –1 –2 O < O < O . Atomic Radius Problem in calculating actual size of atom and hence distance between nuclei is calculated giving rise to three type of radii for atoms. Size of anion

(a)

Covalent radius :

Cr =

d 2

Cr < actual atom size [Slight difference] [Used for H2, Cl2 and such molecules] Half of the internuclear distance between two atoms of the element held by a single covalent bond.

Page 11

A.

Reconstruct Your Chemistry With Prince Sir Calculation for covalent radius can be made as : In homodiatomic molecule (A—A) : r=

B.

dA –A 2

In heterodiatomic molecule (A—B) : (i) When (A – B) is very small : dA–B = rA + rB

(ii)

ri

dA–B = Bond length rA & rB = Covalent radii of A & B respectively. A & B = Electronegativities of A & B.

where

(b)

(dA–A = Bond length in molecule)

S e

When (A – B) is large : dA–B = rA + rB – 0.09 (A – B)

Metallic Radius :

c ni

Half of the internuclear distance between two nearest atoms in the metallic lattice. d

r P Mr =

[Used for metals]

(c)

y B

d 2

Vanderwaal radius :

Half of the internuclear distance between two nearest atoms belonging to two adjacent molecules of noble gases.

Vr =

d 2

Vr >> actual size [very large difference]

In general for an element Vander waal radius > Metallic radius > Covalent Radius (d) Ionic Radius : A cation is smaller than parent atom . An anion is larger than parent atom. SIZE OF ISOELECTRONIC SERIES The species, which have same number of electrons but different nuclear charges, constitute an isoelectronic series. Size in an isoelectronic series 432-1 + 2+ 3+ 4+ 5+ 6+ C >N >O >F >Ne>Na >Mg >Al >Si >P >S (left to right size decreases)

Reconstruct Your Chemistry With Prince Sir

Page 12

IONISATION ENERGIES

The energy required to remove the most loosely bound electron from the outermost orbit of one mole of isolated gaseous atoms of an element, is called ionisation energy or ionisation potential (IP). This ionisation is an endothermic or energy-absorbing process. 



An electron cannot be removed directly from an atom in solid state. For this purpose, the solid state is converted to gaseous state and the energy required for this is called sublimation energy.

1

+

M (g) 

(i)

2+

® M (g) + e

IE

2

r P

c ni

S e

IE < IE < IE because as the electrons go out of the atom, the ionic size goes 1 2 3 on decreasing and the amount of positive charge goes on increasing. Factors Affecting Ionisation Potential Atomic size : When the size of an atom is very large the electron of the outermost orbit bound to the nucleus by weaker attractive forces. Such an electron will be readily removed from the atom. Therefore. The value of ionisation potential will be low.

y B

Ionisation Energies

(ii)

ri

Sucssesive Ionisation Energies The energy required to remove one electron from a neutral gaseous atom to convert it to monopositive cation, is called first ionisation energy (IE ). The 1 energy required to convert a monopositive cation to a diapositive cation is called second ionisation energy (IE ) + 2 M(g) ® M (g) + e IE

1 Atomic radii

Effective Nuclear Charge : Atomic size decreases with increase in effective nuclear charge because, higher the effective nuclear charge stronger will be the attraction of the nucleus towards the electron of the outermost orbit and higher will be the ionisation potential

Ionisation Energies  Z effective (iii)

Shielding Effect : The electrons of internal orbits repel the electrons of the electron of the outermost orbit due to which the attraction of the nucleus towards the electron of the outermost orbit decreases and thus atomic size increases and the value of ionisation potential decreases.

Page 13

Reconstruct Your Chemistry With Prince Sir Ionisation Energies 

1 Sheilding effect

(iv)

Stability of half filled and fully filled orbitals : The atoms whose orbitals are 3 5, 7 2 6 10 14 half-filled (p ,d f ) or fully-filled (s , p , d , f ) have greater stability than the others. Therefore, they required greater energy to for removing out electron. However stability of fully filled orbitals is greater than that of the half filled orbitals

(v)

Penetration power : In any atom the s orbital is nearer to the nucleus in comparison to p, d and f orbitals. Therefor, greater energy is required to remove out electron from s orbital than from p, d and f orbitals. Thus the decreasing order of ionisation potential of s, p, d and f orbitals is as follows s>p>d>f

(a)



c ni

S e

ri

Periodic Table & Ionisation Potential In a Period :- The value of Ionisation potential normally increase on going from left to right in a period, because effective nuclear charge increases and atomic size decreases. Exceptions In second period ionisation potential of Be is greater than that of B, and in the third period ionisation potential of Mg is greater than that of Al due to high stability of fully filled orbitals.

y B

r P



In second period ionisation potential of N is greater than O and in the third period ionisation potential of P is greater than that of S, due to stability of half filled orbitals.



The increasing order of the values of ionisation potential of the second period elements is Li < B < Be < C < O < N < F < Ne The increasing order of the values of ionisation potential of the third period elements is Na < Al < Mg < Si < S < P < Cl < Ar Ionisation Potential of Transition Elements In transition elements, the value of ionisation potential has very little increase on going from left to right in a period because the outermost orbit remains the same but electrons get filled up in the (n–1)d orbitals resulting in very little increase in the values of ionisation potential.

Page 14

Reconstruct Your Chemistry With Prince Sir The value of ionisation potential of transition elements depends on the following two important factors. (a) The value of ionisation potential increases with increase in effective nuclear charge. (b) The value of ionisation potential decreases with increase in shielding effect when the number of electrons increases in (n-1)d orbitals. In the first transition element series the first ionisation potential normally increases on going from left to right from Sc to Cr because shielding effect is much weaker in comparison to effective nuclear charge. The value of first ionisation potential of Fe, Co and Ni remains constant, because shielding effect and effective nuclear charge balance one another. The value of ionisation potential shows slight increase from Cu to Zn because they have fully filled s and d orbitals. The value of first ionisation potential of Mn is maximum because it has maximum stability due to fully filled s and d orbitals.

c ni

S e

ri

Inner Transition Elements The size of inner transition elements is greater than that of d block elements. Therefore the value of ionisation potential of f block elements is smaller than that of d block elements and due to almost constant atomic size of f block elements in a period the value of their ionisation potential remains more constant than that of d block elements.

r P

In a Group The value of ionisation potential normally decreases on going from top to bottom in a group because both atomic size and shielding effect increase.

y B

Exception : The value of ionisation potential remains almost constant from Al to Ga in the III A group. The actual variation for this group is (B > Tl > Ga> Al > In) In the periodic table the element having highest value of ionisation potential is He. The values of ionisation potential of noble gases are extremely high, because 2 6 the orbitals of outermost orbit are fully-filled (ns np ) and provide great stability.

Page 15

Reconstruct Your Chemistry With Prince Sir In a period, the element having least value of ionisation potential is an alkali metal (group I A ) and that having highest value is inert gas (Group 0) Applications of Ionisation Potential The elements having high values of ionisation potential have low reactivity, e.g. inert gases. The value of ionisation potential decreases more on going from top to bottom in a group in comparison to a period. Therefore, reactivity increases and the atom forms a cation by loss of electron.

S e

ri

The elements having low value of ionisation potential readily loose electron and thus behave as strong reducing agents.

c ni

 The elements having low value of ionization potential readily loose electron and thus exhibit greater metallic property. The elements having low value of ionisation potential readily loose electron and thus have basic property.

r P

ELECTRON AFFINITY

y B

The energy released on adding up one mole of electron to one mole of – neutral atom (A) in its gaseous state to form an anion (A ) is called electron affinity of that atom. Since the electron adds up in the outermost orbit, energy is given out. Therefore, electron affinity is associated with an exothermic process. – – A(g) + e  A (g) , H = –En When one electron adds up to a neutral atom, it gets converted to a unit negative ion and energy is released. On adding one more electron to the mononegative anion, there is a repulsion between the negatively charged electron and anion. In order to counteract the repulsive forces, energy has to be provided to the system. Therefore, the value of the second electron affinity is positive. A

–

(g) + e

–



A

–2

(g) ,

H = + En

Page 16

Reconstruct Your Chemistry With Prince Sir Factors Affecting Electron Affinity 1.Effective Nuclear charge When effective nuclear charge is more then, the atomic size less. E. A.  Effective Nuclear Energy

2. Atomic Size or Atomic Radius When the size or radius of an atom increases, the electron entering the outermost orbit is more weakly attracted by the nucleus and the value of electron affinity is lower. E. A. 

1 Atomic size

ri

3.Shielding Effect Shielding effect is directly proportional to atomic size and atomic size is inversely proportional to electron affinity.

c ni

S e

4.Stability of Fully-Filled and Half-Filled Orbitals 6 10, 14 The stability of the configuration having fully-filled orbitals (p ,d f ) and 3 5 7 half-filled orbital (p , d , f ) is relatively higher than that of other configurations.

r P

Periodic Table and Electron Affinity In a period, atomic size decreases with increase in effective nuclear charge and hence increase in electron affinity.

y B

Exception :

6 7 On going from C to N in the second period, the values of electron affinity decreases in stead of increasing. This is because there are 3 half-filled (2p ) orbitals in the outermost orbit of N, which are more 2 stable. On the other hand, the outermost orbit in C has 2p configuration. In the third period, the value of electron affinity of Si is greater than that of P. This is because electronic configuration of the outermost orbit in 3 P atom is 3p , which being half-filled, is relatively more stable The values of electron affinity of inert gases are zero or +ive, because there outermost orbit has fully-filled p orbitals, So they donot have tendency to accept electrons or energy has to be given to add an extra electron. In a period, the value of electron affinity goes on decreases on going from group IA to group IIA. The value of electron affinity of the elements of group IIA is zero because ns orbitals are fully-filled and such orbitals have no tendency to accept electrons.

Page 17

Reconstruct Your Chemistry With Prince Sir In a Group The values of electron affinity normally decrease on going from top to bottom in a group because the atomic size increases which decreases the actual force of attraction by the nucleus. Exceptions The value of electron affinity of F is lower than that of Cl, because the size of F is very small and compact and the charge density is high on the surface. Therefore, the incoming electron experiences more repulsion in comparison to Cl . That is why the value of electron affinity of Cl is highest in the periodic table. The order among the halogens is Cl > F > Br >I

S e

ri

For the elements of oxygen family the order of electronegativity is S > Se > Te > Po > O

c ni

The values of electron affinity of alkali metals and alkaline earth metals can be regarded as zero, because they do not have tendency to form anions by accepting electron. Imortant Note While comparing electron affinity we will only consider magnitude but while comparing Electron Gain Enthalpy we will consider sign also. For e.g. following E.A. values O = -144 kj/mole, S = -200 kj/mole, Se = -195 kj/mole Te = -190 kj/mole Po = -174 kj/mole Order of E. A. is : S > Se >Te > Po > O (consider magnitude only ) Order of E.G.E. is : S < Se H–Cl > H–Br > H–I

ri

Reactivity Bond energy is proportional to stability and stability is inversely proportional to reactivity. H–F < H –Cl < H –Br < H–I

c ni

S e

(v) Nature of Oxides If difference of the two electronegativities (X –X ) is 2.3 or more then oxide O A will be basic in nature. Similarly if value of X –X is lower than 2.3 then the O A compound will be first amphoteric then acidic in nature. Oxide (X –X ) O A Nature

Na O 2 2.6 Strong basic

r P

MgO 2.3 Basic

y B

Al O 2 3 2.1 Amphoteric

SiO 2 1.8 weak

P O 2 5 1.5 Acidic

SO 3 1.1 Strong acidic

Cl O 2 7 0.5 Strongest acidic

Basic character of oxides decreases with increasing difference of electronegativities of central atom and oxide Therefore basic character decreases in the period and acidic character increases. NATURE OF HYDROXIDES According to Gallis if electronegativity of A in a hydroxide (AOH) is more than 1.7 then it will be acidic in nature whereas it will be basic in nature if electronegativity is less than 1.7 For example NaOH and ClOH Electronegativity (X ) 0.9 3.00 A Nature Basic Acidic If the value is more than X –X , then that hydroxide will be basic otherwise O H it will be acidic.

Page 23

Reconstruct Your Chemistry With Prince Sir Acidic and Basic Nature of Hydroxides of Elements Acidic and basic nature of hydroxide of an element AOH depends on ionisation potential of A. If ionisation potential of A is low then it will give its electron to oxygen easily thus AOH will be basic. NATURE OF OXY ACIDS Those inorganic acids in which X–OH group is present are called oxy acids. (Here X = a non metal) like HOCl is an oxy acid.

ri

In a Period Moving from left to right in a period strength of oxy acids increases like Order of strength of oxy acid of elements of second period H BO < H CO < HNO 3 3 2 3 3

c ni

S e

Order of strength of oxy acids of elements of third period H SiO < H SO < HClO 2 3 2 4 4

r P

In a Group

Moving upwards to downward in a group, strength of oxy acids decreases Like - in VA group – HNO > H PO 3 3 4

y B

When amount of oxygen increases in an oxy acid of an element (i.e. on increasing oxidation state of element) strength of acid increases like (a) H SO < H SO 2 3 2 4

(b) HClO < HClO

2

< HClO

3

< HClO

4

Stability of acids also increases in the same order NATURE OF HYDRIDES Moving from left of right in a period, nature of hydrides of nonmetal elements becomes basic to acidic. Like In second period NH H O HF 3 2 Weak base Neutral Acidic (Amphoteric) In third period

PH 3 Very weak acid

H S 2 Weak acid

HCl Strong acid

Moving upwards to downwards in a group, acidic nature and reducing power

Page 24

Reconstruct Your Chemistry With Prince Sir of hydrides of nonmetal elements both increases but stability decreases. Like Acid strength and reducing strength : HF < HCl < HBr < HI stability : HF > HCl > HBr > HI DENSITY Density in solid state changes on change in atomic number of elements. Moving left to right in the period densities of solid and liquid elements first increases, it is maximum in the middle and then decreases. Moving upwards to downwards in a group, density generally increases but there is irregularity in this order. Density decreases instead of increasing on moving towards K in IA group and moving from Mg to Ca in IIA group. Because penultimate value of K and Ca increases, due to which density gets reduced. Os (22.6 gm/cc) is most dense in solid elements and in liquids Hg is most dense (13.6 gm/cc)

c ni

S e

ri

DIAGONAL RELATIONSHIP Some elements of second period Li, Be, B shows dissimilarities with other elements of this group but shows similarities with elements of third group like Mg,. Al, Si situated diagonally to them. It is called diagonal relationship. (a)

r P

Similarities between properties of Li and Mg are as follows Li and Mg both reacts directly with nitrogen to form lithium nitride (Li 3N) and magnesium nitride (Mg3N2) whereas other alkali metals of IA group does not form nitride.

y B

(b)

Fluoride carbonate and phosphate of Li and Mg are insoluble in water whereas these compounds of other alkali metals are soluble.

(c)

Li and Mg both are hard metals, whereas other metals of IA group are soft.

(d)

LiOH and Mg(OH)2 both are weak base, whereas hydroxides of other elements of IA group are strong base.

(f)

Metallic bond in Li and Mg both are strong compare to other alkali metals .

(g)

Their melting and boiling points are high.

(h)

By thermal disintegration of LiNO3 and Mg (NO3)2 Li2O and MgO is obtained respectively.

(i)

Thermal stability of Li2CO3 and Mg CO3 is very less compare to other alkali metals and they liberates CO2 gas easily.

Page 25

Reconstruct Your Chemistry With Prince Sir

(a)

Similarly Be shows similarity to Al of IIIA group They both elements does not provide colour to Bunsen burner.

(b)

They both are comparatively stable in air.

(c)

Both are insoluble in NH3 therefore does not form blue coloured solution.

(d)

There is not tendency of making peroxide and superoxide in them.

(e)

Reducing power is very less due to low value of standard electrode potential in the form of oxidation potential.

(f)

Be and Al both forms halogen bridge halides.

c ni

S e

Some Important Points (1) Triad rule – Dobereiner (2) Octet rule – Newland (3) Study of atomic volume – Lothar Mayer (4) Inventor of atomic number – Moseley (5) Godfather of periodic table – Mandeleef (6) Maker of modern periodic table – Bohr (7) Mg is bridge element, which joins metals of IIA and II B groups. (8) Elements after atomic number 92 are transuranic elements. 43 (9) Artificial element is Tc . (10) Liquid non-metal – Br (11) Liquid metal – Hg, Ga, Cs, Fr (12) Solid volatile non-metal – Iodine (13) Lightest metal – Li (14) Heaviest metal – Ir (15) Hardest metal – W (16) Nobel metals – Pc, Pt, Au, Ag (17) Element most found on earth – Al (18) Gaseous elements – 11 (He, Ne, Ar, Kr, Xe, Rn, H , N , O , Cl , F ) 2 2 2 2 2 (19) Liquid elements – 5(Br, Hg, Ga, Cs, Fr) (20) Submetals – 5(B, Si, As, Te, At) (21) Inert gases – 6 (22) Lowest electronegativity : Cs (23) Highest electronegativity : F (24) Highest ionisation potential : He (25) Lowest ionisation potential : Cs (26) Highest electron affinity : Chlorine (Cl) (27) Least electropositive element : Fluorine (F)

y B

r P

ri

Page 26

Reconstruct Your Chemistry With Prince Sir (28) Most reactive solid element : Li (29) Most reactive liquid element : Cs (30) Most stable element : Te (31) Largest atomic size : Cs (32) Most electropositive element : Cs ; Fr (in stable element) (In all element) (33) Group containing maximum no. : Zero gp ; next to zero gp is VII gp or of gaseous elements in periodic table halogen gp (F and Cl ) 2 2 (34) Total number of gaseous : 11 (H , He, N , O , F Ne, Cl 2 2 2 2 2 elements in periodic table Ar, Kr, Xe, Rn) (35) Total number of liquid elements : 4 (Ga, Br, Cs, Hg) in periodic table (Fr and Eka are also liquid) (36) Volatile d-block elements : Zn, Cd, Hg (37) Most abundant element on earth : Oxygen followed with Si (38) Most stable carbonate : Cs CO 2 3 (39) Strongest alkali : Cs(OH) (40) Element kept in water : P (41) Elements kept in kerosene oil : Na, K, I, Cs (42) Liquid non metal : Br 2 (43) Bridge metals : Na, Mg (44) Noble metals : Au, Pt (45) Lightest element : H (46) Poorest conductor of current : Pb(metal), S (non metal) (47) Most abundant gas : N 2 (48) Lightest solid metal : Li 3 (49) Heaviest solid metal : Os(highest density 22.6 g/cm ) (50) Natural explosive : NCl 3 (51) Dry ice : CO 2 (52) First Nobel prize of chemistry was given to : vant Hoff * Core charge - Atomic number – Kernel of electron * Penultimate shell – Shell present inside one shell (n – 1) from outermost shell, is called penultimate shell. * Prepenultimate shell – Shell present inside two shells (n – 2) from outermost shell, is called prepenultimate shell About 75% of known elements are metals. The non-metals (20 ) are located at the top right of the periodic table.Elements at the border line of metals and non-metals are called metalloids or semimetals e.g., B,Ge, As. Total number of solid elements are about 90. Total number of radioactive elements are about 42. Liquid radioactive element : Francium. Rarest element in the earth’s crust : Astatine. Most poisonous metal : Plutonium. Element having maximum number of natural isotopes (10) : Tin. First man made element : Teachnetium (At. no. = 43). van der Waals’ radius > Metallic radius > Covalent radius. (for an atom)

y B

r P

c ni

S e

ri

Page 27

Reconstruct Your Chemistry With Prince Sir Size of anion > Size of atom > Size of cation e.g., I¯ > I >I +. Size of largest atom (Cs) is approximately 4.4 times greater than that of H-atom. Helium has the highest value of I.E. and the lowest value of electron gain enthalpy among all the elements. Fluorine is the most electronegative element whereas cesium is the least electronegative element. Size of isoelectronics decreases with increase in atomic number. Hydrogen is the most abundant element in universe. Oxygen is the most abundant metal on earth. Al is the most abundant metal on earth. The abundance of first five element on earth is O > Si > Al > Fe > Ca. Metallic nature decreases along the period but increases down the group. Non-metallic nature increases along the period but decreases down the group. Out of 17 non-metals, 11 are gases, one is liquid (Br) and five are solid C, P, S, Si & 1. The most electronegative element is F(EN = 4.0 on the Pauling scale) while the least electronegative elements are Cs and Fr with EN = 0.7. Metalloids have electronegativity values closer to 2.0. Lanthanoids were also called rare earth elements since their oxides were rarely found on earth in earlier days.

y B

r P

c ni

S e

ri

Page 28

Reconstruct Your Chemistry With Prince Sir

Exerxise –1 Only one choice is correct

Q.1. Q.2. Q.3. Q.4. Q.5. Q.6.

Q.7.

Q.8.

Q.9.

Q.10. Q.11. Q.12. Q.13. Q.14. Q.15. Q.16. Q.17.

Q.18. Q.19.

Among O, C, F, Cl, Br the correct order of increasing atomic radii is : (A) F < O < C < Cl < Br (B) F < C < O < Br < Cl (C) F < Cl < Br < O < C (D) C < O < F < Cl < Br Atomic radii of fluorine and neon in Angstrom units are respectively given by : (A) 0.72, 1.60 (B) 1.60, 1.60 (C) 0.72, 0.72 (D) None of these The first ionisation energy of Na, Mg, Al and Si are in the order : (A) Na < Mg > Al < Si (B) Na > Mg > Al > Si (C) Na < Mg < Al > Si (D) Na > Mg > Al < Si The first ionization energy will be maximum for : (A) Uranium (B) He (C) Lithium (D) Iron Which has maximum value for proton : (A) IE (B) Radius (C) Charge (D) Hydration energy Ionisation potential of Na would be numerically the same as : (A) Electron affinity of Na+ (B) Electronegativity of Na+ (C) Electron affinity of He (D) Ionisation potential of Mg. The IP1, IP2, IP3, IP4 and IP5 of an element are 7.1, 14.3, 34.5, 46.8, 162.2 eV respectively. The element is likely to be : (A) Na (B) Si (C) F (D) Ca Ionic radius are : (A) Inversely proportional to effective nuclear charge (B) Inversely proportional to sqaure of effective nuclear charge (C) Directly proportional to effective nuclear charge (D) Directly proportional to sqaure of effective nuclear charge. Which one is incorrect statement : (A) IE1 of He is maximum among all elements (B) EA1 for noble gases is zero (C) Electronegativity is maximum for fluorine (D) IE1 for N < IE1 for O Which set of elements have nearly the same atomic radii : (A) F, Cl, Br, I (B) Na, K, Rb, Cs (C) Li, Be, B, C (D) Fe, Co, Ni, Cu In which pair, the second atom is larger than first : (A) Br, Cl (B) Na, Mg (C) Sr, Ca (D) N, P An element X occurs in short period having configuration ns2np1. The formula and nature of its oxide is : (A) XO3, basic (B) XO3, acidic (C) X2O3, amphoteric (D) X2O3, basic An example of a non-stoichiometric compound is (A) Al2O3 (B) Fe3O4 (C) NiO (D) PbO Mendeleeff corrected the atomic weight of : (A) Be (B) ln (C) Os (D) All of these. The ionic radii of N3–, O2– and F– are respectively given by : (A) 1.36, 1.40, 1.71, (B) 1.36, 1.71, 1.40 (C) 1.71, 1.40, 1.36 (D) 1.71, 1.36, 1.40 Most of the known elements are : (A) Metals (B) Non-metals (C) Transition metals (D) Inner transitional metals Zn and Cd do not show variable valency like ‘d’ block elements due to : (A) Softness (B) Completely filled ‘d’ orbital (C) Two electrons in outermost orbit (D) Low m.pt. Which of the following is paramagnetic : (A) O2– (B) CN– (C) CO (D) NO+ The element having 18 electrons in its outermost shell is : (A) 28Ni (B) 46Pd (C) 29Cu (D) None of these.

y B

r P

c ni

S e

ri

Page 29

Reconstruct Your Chemistry With Prince Sir

y B

r P

c ni

S e

ri

Reconstruct Your Chemistry With Prince Sir Page 30 Q.20. The order of first electron affinity of O, S and Se is : (A) O > S > Se (B) S > O > Se (C) Se > O > S (D) Se > S > O Q.21. Which of the following sequence regarding the first ionisation potential of coinage metal is correct : (A) Cu > Ag > Au (B) Cu < Ag < Au (C) Au > Cu > Ag (D) Ag > Cu < Au Q.22. Which is not arranged in the correct sequence : (A) d5, d3, d1, d4 — increasing magnetic moment (B) MO, M2O3, MO2, M2O5 — decreasing basic strength (C) Sc, V, Cr, Mn— increasing number of oxidation states (D) Co2+, Fe3+, Cr3+, Sc3+ — increasing stability Q.23. Na+, Mg2+, Al3+, Si4+ are isoelectronics. Their ionic size follows the order : (A) Na+ < Mg2+ < Al3+ < Si4+ (B) Na+ > Mg2+ < Al3+ < Si4+ + 2+ 3+ 4+ (C) Na < Mg > Al > Si (D) Na+ > Mg2+ > Al3+ > Si4+. Q.24. Which transition involves maximum amount of energy : (A) M¯(g)  M(g) + e (B) M(g)  M+(g) + e (C) M+(g)  M2+(g) + e (D) M2+(g)  M3+(g) + e. Q.25. Which set of metals becomes passive in conc. HNO3 : (A) Fe, Co, Ni, Al, Cr (B) Na, Mg, K, Cs (C) Mn, Ba, Sr, Ca, Al (D) Pb, Ag, Hg, Cu st Q.26. Low melting point of manganese in the 1 transition series is due to : (A) Strong metallic bond due to d10 configuration (B) Weak metallic bond due to d5 configuration (C) Weak metallic bond due to d7 configuration (D) None of these. Q.27. The order of basic character of given oxides is : (A) Na2O > MgO > Al2O3 > CuO (B) MgO > Al2O3 > CuO > Na2O (C) Al2O3 > MgO > CuO > Na2O (D) CuO > Na2O > MgO > Al2O3. Q.28. Which pair of elements is chemically most similar : (A) Na, Al (B) Cu, S (C) Ti, Zr (D) Zr, Hf Q.29. Decreasing order of atomic weight of the elements given below is : (A) Fe > Co > Ni (B) Ni > Co > Fe (C) Co > Ni > Fe (D) Co > Fe > Ni. Q.30. The largest number of known elements are placed in period : (A) IV (B) V (C) VI (D) VII Q.31. In the transition elements, the incoming electron occupies (n – 1)d sublevel in preference to : (A) np (B) ns (C) (n – 1) d (D) (n + 1)s Q.32. The correct increasing order of bond energy of the following is : (A) N2 < O2 < Cl2 < F2 (B) Cl2 < F2 < N2 < O2 (C) F2 < Cl2 < O2 < N2 (D) O2 < Cl2 < F2 < N2. Q.33. Which of the following is paramagnetic molecule : (A) ClO2 (B) SO2 (C) SiO2 (D) CO2. Q.34. Which arrangement represents the correct order of electron gain enthalpy (with negative sign) of the given atomic species : (A) S < O < Cl < F (B) O < S < F < Cl (C) Cl < F < S < O (D) F < Cl < O < S

y B

r P

c ni

*****

S e

ri

Page 31

Reconstruct Your Chemistry With Prince Sir

Exerxise –2

Q.1. Q.2. Q.3. Q.4. Q.5. Q.6. Q.7.

Passage–1 Ionisation energies of five elements are given below : Atom I.E. (in kcal/mol) I II III A 300 549 920 B 99 734 1100 C 118 1091 1652 D 176 347 1848 E 497 947 1500 Which element is a noble gas : (A) A (B) E (C) C (D) D Which element form stable unipositive ion : (A) A (B) B (C) C (D) D The element whose msot stable oxidation state is + 2 : (A) B (B) C (C) D (D) E Which is a non-metal : (A) A (B) B (C) C (D) D Most reactive metal is : (A) A (B) B (C) C (D) D If B reacts with fluorine and oxygen, the molecular formula of fluoride and oxide will be respectively : (A) BF3, B2O3 (B) BF, B2O (C) BF2, BO (D) None of these. Which of the following pair represents elements of same group : (A) B, C (B) A, B (C) A, D (D) B, D Passage–2 Read the following paragraph and answer the questions given below : Bohr’s periodic table has seven horizontal rows called periods and eighteen vertical columns called groups. Alternatively, the elements of the Bohr’s table are divided into four blocks namely s-, p-d- and fblocks. Group, period and block of given element can be predicted with the help of its electronic configura tion. (i) The principle quantum number of valence shell refers to period of element. (ii) The nature of subshell (s, p or d) having last electron refers the block to which the element belongs. (iii) The number of electrons present in outermost or penultimate shell predicts the group of element. For s-block, the group number = No. of valence electron For p-block, the group number = 10 + numbers of valence electron For d-block, the group number = 2 + (n – 1)d electrons The element of atomic number 36 belongs to (A) 3rd period and p-block (B) 4th period and p-block th (C) 4 period and s-block (D) 3rd period and d-block The element with atomic number 58 belongs to (A) s-block (B) p-block (C) d-block (D) f-block. An element with atomic number 24 is an example of (A) Alkali metal (B) Alkaline earth metal (C) Inert gas (D) Transition element Number of groups of s-block element, possible in a period is/are (A) 1 (B) 2 (C) 3 (D) 4 Number of groups possible for s + p + d block elements in a period are (A) 6 (B) 2 (C) 10 (D) 18

y B

Q.8.

Q.9. Q.10. Q.11 Q.12.

r P

c ni

S e

ri

Page 32

Reconstruct Your Chemistry With Prince Sir

Passage–3 The amount of energy required to remove an electron from the last orbit of an isolated (free) atom in gaseous state is known as ionization potential or first ionization potential of the element. Similarly, the energy required for the removal of the electron from the unipositive ion (produced above) is reffered to as second ionization potential and thus the third, fourth etc., Ionization potentials may be defined in general as the number of ionization potentials of an element may be as much as its number of electrons. Q.13. The relationship between IE3 and IE1 of an element is : (A) IE3 > IE1 (B) IE3 < IE1 (C) IE3 = IE1 (D) None of the above Q.14. Which of the following statement is correct ? (A) IE of elements increases along the period. (B) IE of the 3rd group elements is less than that of the second group elements. (C) IE of group 16 elements is less than that of group 15 elements. (D) All of the above. Q.15. IE1, IE2 and IE3 of an element are 9.32, 18.21 and 153.83 eV. What information these data convey ? (A) The element has two s electrons in the valence shell. (B) The element has two p-electrons in the valence shell. (C) Both of the above (D) None of the above Q.16. Ionization energy increases on moving down a group due to (A) increase in the number of electrons (B) Decrease in size of elements (C) Screening effect (D) None of the above. Q.17. Ionization energy of nitrogen is more than that of oxygen which can be explained by (A) High-effective nuclear charge (B) Low screening (C) Half-filled stability (D) None of the above Assertion & Reason Type Questions In each sub question below a statement S and an explanation E is given. Choose the correct answers from the codes, A, B, C, D given for each question. (A) S is correct but E is wrong. (B) S is wrong but E is correct. (C) Both S and E are correct and E is correct explanation of S. (D) Both S and E are correct but E is not correct explanation of S. Q.18. S : Iodine possess metallic nature. E : It has violet crystalline nature and shows ionic compounds IPO4. Q.19. S : The third period contains only 8 electrons and not 18 like 4th period. E : In III period filling starts from 3s1 and complete at 3p6 whereas in IV period it starts from 4s1 and complete atfter 3d10 and 4s2. Q.20. S : Transition elements show horizontal as well as vertical relationship. E : This is due to shielding effect and similar electronic configuration. Q.21. S : Cs and F2 combines violently to form CsF. E : Cs is most electropositive and F is most electronegative. Q.22. S : Each transiting series possesses ten molecule. E : In transition elements filling of electrons occurs in d sub-shells. Q.23. S : A jump in 3rd ionisation energy is noticed in case of alkaline earth metals. E : The jump in ionisation energy is due to change in major energy shell during successive removal of electron. Q.24. S : First ionisation energy of Be is more than that of IE1 of B. E : Removal of electron in Be occurs from 2s sub-shell whereas in B from 2p sub-shell.

y B

r P

c ni

S e

ri

Reconstruct Your Chemistry With Prince Sir Page 33 Q.25. S : Formation of Cl¯ is exothermic whereas formation of O2– is endothermic. E : EA2 of oxygen is endothermic and greater than its exothermic value EA1 of oxygen. Q.26. S : BiCl5 does not exist. E : In Bi inert pair effect is more prediominant. Q.27. S : Second EA for halogens is zero. E : Fluorine has maximum value of electron affinity. Q.28. S : Pure NaCl is hygroscopic in nature. E : Pure NaCl is non-hygroscopic in nature. Q.29. S : Sulphate is estimated as BaSO4 and not MgSO4. E : Ionic radii of Mg2+ is smaller than that of Ba2+. Q.30. S : First ionization enthalpy of Be is greater than that of B. E : 2p orbital is lower in energy than 2s orbital. Q.31. S : Nobel gases have highest ionization enthalpies in their respective periods. E : Nobel gases have stable closed shell electronic configuration.

Q.32. Q.33. Q.34. Q.35. Q.36.

Q.37. Q.38.

Q.39. Q.40. Q.41. Q.42.

c ni

S e

ri

More than One Correct The elements which are radioactive and have been named after the names of planet are (A) Hg (B) Np (C) Pu (D) Ra Which of the process do not involve absorption of energy ? (A) S(g) + e–  S–(g) (B) O– + e–  O2– (g) (C) Cl(g) + e–  Cl–(g) (D) O(g) + e–  O– (g) Which of the following pairs of elements have almost similar atomic radii ? (A) Zr, Hf (B) Mo, W (C) Co, Ni (D) Nb, Ta Which of the following elements have similar value of electronegativity ? (A) H (B) S (C) Te (D) P Mark the correct statements out of the following (A) He has highest IE1 in the periodic table. (B) Cl has the highest EA out of all elements in the periodic table. (C) Hg and Br are liquid at room temperature. (D) In any period, the atomic radius of the noble gas is lowest. Which of the following show amphoteric behaviour ? (A) Zn(OH)2 (B) BeO (C) Al2O3 (D) Pb(OH)2 Ionization energy is influenced by (A) Size of atom (B) Charge of nucleus (C) Electrons present in inner shells (D) None of the above Which is correct in increasing order of ionic character ? (A) AlCl3 < MgCl2 < NaCl (B) LiF < LiCl < LiBr < Lil (C) NaCl > MgCl2 < AlCl3 (D) None of the above. Which of the following is correct in order of increasing size ? (A) I+ < I < I– (B) Fe < Fe2+ < Fe3+ (C) Fe3+ < Fe2+ < Fe (D) All of the above. Which is the correct increasing order of ionization energy ? (A) Li < B < Be (B) Be < B < Li (C) Li < Na < K (D) O < N < F The first eight ionization energies for a particular neutral atom is as given below. All values are expressed in MJ mol–1. Which oxidation state(s) is/are not possible for the atom ? 1st 2nd 3rd 4th 5th 6th 7th 8th 1.31 3.39 5.30 7.47 10.99 13.33 71.33 94.01 (A) – 2 (B) – 3 (C) – 6 (D) 6

y B

r P

*****

Page 34

Reconstruct Your Chemistry With Prince Sir

Exerxise –3 Previous Years IIT Based Questions Only one option is correct : Q.1. The hydrogen energy of Mg2+ is larger than the of : (1984) 3+ + 2+ 3+ (A) Al (B) Na (C) Be (D) Mg Q.2. The first ionization potential in electron volts of nitrogen and oxygen atoms are respectively given by : (A) 14.6, 13.6 (B) 13.6, 14.6 (C) 13.6, 14.6 (D) 14.6, 14.6 (1987) Q.3. Atomic radii of fluorine and neon in Angstrom units are respectively given by : (1987) (A) 0.72, 1.60 (B) 1.60, 1.60 (C) 0.72, 0.72 (D) None of these Q.4. The electronegativity of the following elements increases in the order : (1987) (A) C, N, Si, P (B) Na, Si, C, P (C) Si, P, C, N (D) P, Si, N, C Q.5. The first ionization potential of Na, Mg, Al and Si are in the order : (1988) (A) Na < Mg > Al < Si (B) Na > Mg > Al > Si (C) Na < Mg < Al > Si (D) Na > Mg > Al < Si Q.6. Which of the following is the smallest is size ? (1989) 3– 2– + (A) N (B) O (C) F¯ (D) Na Q.7. Amongst the following elements (whose electronic configurations are given below), the one having the highest ionization energy is : (1990) (A) [Ne] 3s2 3p1 (B) [Ne] 3s2 3 p3 (C) [Ne] 3s2 3p2 (D) [Ar] 3d10 4s2 4p3. Q.8. The statement that is not correct for the periodic classification of elements is : (1992) (A) The properties of elements are the periodic functions of their atomic numbers. (B) Non-metallic elements are lesser in number than metallic elements. (C) The first ionization energies of elements along a period do not vary in a regular manner with increase in atomic number. (D) For transition elements the d-sub shells are filled with electrons monotonically with increase in atomic number. Q.9. Which has most stable +2 oxidation state : (1995) (A) Sn (B) Pb (C) Fe (D) Ag Q.10. Which of the following has the maximum number of unpared electrons : (1996) 2+ 3+ 3+ 2+ (A) Mg (B) Ti (C) V (D) Fe . Q.11. The incorrect statement among the following is : (1997) (A) The first ionisation potential of Al is less than the first ionisation potential of Mg (B) The second ionsation potential of Mg is greater than the second ionisation potential of Na (C) The first ionisation potential of Na is less than the first ionisation potential of Mg (D) The third ionisation of Mg is greater than third ionisation potential of Mg Q.12. Which of the following compounds is expected to be coloured ? (1997) (A) Ag2SO4 (B) CuF2 (C) MgF2 (D) CuCl. Q.13. The following question of an Assertion in column 1 and the Reason in column 2 : (1998) Column I (Assertion) Column -2 (Reason) F atom has a less negative electron affinity Additional electrons are repelled more effectively than Cl atom by 3p electrons in Cl atom than by 2p electrons in F atom. Which of the following is correct ? (A) If both assertion and Reason are correct, and reason is the correct explanation of the assertion (B) If both assertion and reason are correct, but reason is not the correct explanation of the assertion (C) If assertion is correct but reason is incorrect (D) If assertion is incorrect but reason is correct Q.14. The correct order of radii is : (2000) 2– 3– 3+ (A) N < Be < B (B) F¯ < O < N (C) Na < Li < K (D) Fe < Fe2+ < Fe4+.

y B

r P

c ni

S e

ri

Reconstruct Your Chemistry With Prince Sir Page 35 Q.15. This question contain an Assertion in column 1 and the Reason in column 2. Use the following key to choose the appropriate answer : (2000) Column I (Assertion) Column -2 (Reason) The first ionisation energy of Be is 2p orbital is lower in energy than 2s. greater than that of B (A) If both assertion and Reason are correct, and reason is the correct explanation of the assertion (B) If both assertion and reason are correct, but reason is not the correct explanation of the assertion (C) If assertion is correct but reason is incorrect (D) If assertion is incorrect but reason is correct Q.16. The set representing the correct order of first ionization potential is : (2001) (A) K > Na > Li (B) Be > Mg > Ca (C) B > C > N (D) Ge > Si > C More than one is correct Q.17. The statements that are true for the long form of the periodic table are : (1988) (A) It reflects the sequence of filling the electrons in the order of sub-energy level s, p, d and f. (B) It helps to predict the stable valency states of the elements (C) It reflects tends in physical and chemical properties of the elements (D) It helps to predict the relative ionicity of the bond between any two elements Q.18. Sodium sulphate is soluble in water whereas barium sulphate is sparingly soluble because :(1989) (A) The hydration energy of sodium sulphate is more than its lattice energy (B) The lattice energy of barium sulphate is more than its hydration energy (C) The lattice energy has no role to play ion solubility (D) The hydration energy of sodium sulphate is less than its lattice energy Q.19. The first ionization potential of nitrogen and oxygen atoms are related as follows : (1989) (A) The ionization potential of oxygen is less than the ionization potential of nitrogen. (B) The ionization-potential of nitrogen is greater than the ionization potential of oxygen (C) The two ionization-potential values are comparable. (D) The difference between the two ionzation-potential too large. Q.20. Ionic radii of : (1999) 4+ 7+ 35 37 + 3+ 5+ (A) Ti < Mn (B) Cl¯ < Cl¯ (C) K > Cl¯ (D) P > P . Subjective Type Questions: Q.21. Arrange the following in order of their (1985) (i) Decreasing ionic size Mg2+, O2–, Na+, F¯ (ii) Increasing first ionization energy Mg, Al, Si, Na (iii) Increasing bond length F2, N2, Cl2, O2. Q.22. Arrange the following in the order of their increasing size : (1986) 2– 2+ Cl¯, S , Ca , Ar Q.23. Explain the following : (1989) “The first ionzation energy of carbon atom is greater than that of boron atom whereas, the reverse is true for the second ionzation energy” . Q.24. Arrange the following as stated : (1991) “Increasing order of ionic size” N3–, Na+, F¯, O2–, Mg2+ Q.25. Compare qualitatively the first and second potentials of copper and zinc. Explain the observation. (1996) Q.26. Arrange the following ions in order of their increasing radii (1993) + 2+ + 3+ Li , Mg , K , Al .

y B

r P

c ni

S e

ri

Page 36

Reconstruct Your Chemistry With Prince Sir

ANSWERS Exercise - 1 1. A 6. A 11. D 16. A 21. C 26. D 31. A

2. A 7. B 12. C 17. B 22. A 27. A 32. C

3. A 8. A 13. B 18. A 23. D 28. D 33. A

4. B 9. D 14. D 19. B 24. D 29. C 34. B

2. B 7. A 12. D 17. C 22. C 27. A 32. B, C 37. A, B, C, D 42. B, C

3. C 8. B 13. A 18. C 23. C 28. A 33. A, C, D 38. A, B, C

4. A 9. D 14. D 19. C 24. C 29. D 34. A, B, C, D 39. A, B

Exercise - 2 1. B 6. B 11. B 16. C 21. C 26. C 31. C 36. A, B, C 41. A, C, D

Exercise - 3 1. B 6. D 11. B 16. B

2. A 7. B 12. B 17. A, C, D

y B

r P

c ni

3. A 8. D 13. C 18. A, B

S e 4. C 9. B 14. B 19. A, B, C

5. D 10. D 15. C 20. B 25. A 30. C

ri

5. B 10. D 15. C 20. C 25. C 30. A 35. A, C, D 40. A, C

5. A 10. D 15. C 20. D

Page 37

Reconstruct Your Chemistry With Prince Sir

y B

r P

c ni

S e

ri