Periodic Table Ipe

Periodic Table

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PERIODIC TABLE

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MENDELEEF’S PERIODIC TABLE Mendeleef proposed following periodic law and arranged elements in the increasing order of atomic masses. Mendeleef’s Periodic law "Physical and chemical properties of elements are periodic functions of their atomic weights" * The elements are arranged in horizontal rows and vertical columns. * The horizontal rows are called periods. The number of elements in a period may vary. The first three periods of the Mendeleef table are called as short periods. The other periods are known as long periods. * The vertical columns in the table are known as groups or families. The groups are sub divided into two subgroups ‘A’ and ‘B’. The elements arranged in a group possess similar properties. * The long period of the Mendeleef periodic table consists of two rows of elements called as series. * He considered the similarities in the formulae and the properties of the compounds formed by the elements. * Mendeleef observed that elements with similar properties had (i) Either almost the same atomic weights. e.g. Fe(56), Co(59), Ni(59) Os(191), Ir(193), Pt(195) (ii) Or atomic weights which showed a constant increase e.g. K(39), Rb(85), Cs(133) Ca(40), Sr(88), Ba(137) * The elements with low atomic weights are called typical elements. These are arranged in three short periods of the periodic table. * Group VIII of the Mendeleef table contains three triads, namely, (Fe, Co, Ni and Ru, Rh, Pd and Os, Ir, Pt). These triads are called transition elements which include Sc(21) to Zn (30), lanthanides and actinides. * From a study of adjacent elements and their compounds, Mendeleef was able to predict the characteristics of certain elements which were found to be very accurate. e.g.1) Eka Al - Gallium 2) Eka Si - Germanium 3) Eka B - Scandium Comparison of properties predicted by Mendeleef with those observed. S.No.

Property

1

Name of the element

As predicted by Mendeleeff Eka Aluminium

2

Atomic weight

68

70

-3

As observed experimentally Gallium

3

Density (g.cm )

5.90

5.94

4

Formula of oxide

(Eka Al)2 O3

Ga2O3

5

Formula of Chloride

(Eka Al)Cl3

GaCl3

* He corrected the atomic weights of some elements like Beryllium, Indium, Uranium.

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Limitations : * Zero group elements were not known at the time of Mendeleef. * In every group lanthanides were placed though they have dissimilar properties. * At some places it violates the increasing order of atomic weight rule. Ar40 & K39 Co59 & Ni58 Te128 & I127 Th 232 & Pa 231 These four pairs are called anomalous pairs or inverted pairs. * Elements with dissimilar properties were grouped together. E.g. ' Th' is placed in III group and Ag is placed in I group. * Atomic mass is taken as fundamental property.

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  a  z  b

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MODERN LONG FORM OF PERIODIC TABLE Modern periodic table was constructed by Neils Bohr based on modern periodic law. Modern periodic law The chemical and physical properties of elements are the periodic functions of their atomic numbers and electronic configurations. Above law was proposed by Moseley. He found the relation between atomic numbers (z) and the frequencies ( ) of x- rays produced by them.

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Where a & b are constants characteristic of elements

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SALIENT FEATURES OF MODERN LONG FORM OF PERIODIC TABLE Modern long form of periodic table consists of seven rows called periods and eighteen columns called groups. Groups: * The 18 groups in the periodic table are numbered as IA , IIA, IIIB...... VIII (3 groups) IB ,IIB, IIIA - VIIA and zero group. * I A to VII A group elements are called representative elements. * Zero group elements are called inert gases. * IB to VII B and VIII B group element are called transition elements. * The elements present in a group show similar chemical properties as they have similar outer electronic configuration. Periods: * Each period starts with alkali metals (IA) and ends with inert gas elements. * The first period is a very short period with only two elements i.e., Hydrogen (H) & Helium (He). * The second period starts with Lithium (Li) and ends with Neon (Ne) and contains 8 elements. It is called first short period. * The third period also contain 8 elements i.e., from Sodium (Na) to Argon (Ar). It is called second short period. * The fourth period is the first long period with 18 elements , it starts Potassium (K) & ends with Krypton (Kr). It also includes 10 elements belonging to 3d series i.e., from Scandium (Sc) to Zinc (Zn). * The fifth period is the second long period with 18 elements, it starts with Rubidium (Rb) and ends with Xenon (Xe). It also includes 10 elements belonging to 4d series i.e. from Yttrium (Y) to

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Cadmium (Cd). * The sixth period is the longest period with 32 elements. It not only includes 10 elements belonging to 5d series i.e., Lanthanum (La), Hafnium (Hf) to Mercury (Hg) but also contain 14 elements belonging the 4f series called lanthanides (Cerium(Ce) to Lutetium (Lu)). * The seventh is an incomplete period which starts with Fr. It includes the 14 elements belonging to 5f series called actinides (Thorium (Th) to Lawrencium (Lr)). * Lanthanides and actinides are placed below the period separately. Advantages: 1. The chemical and physical properties of the elements can be studied easily. 2. Position of metals and nonmetals can be know. 3. The chemistry of transition elements can be studied along with d-block elements.

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Total no of electrons filled in all the sub-levels 2 8 8 18 18 32 --

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Main and sub-energy levels {as (nl)} 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d

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No. of the period 1. 2. 3. 4. 5. 6. 7.

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THE RELATION BETWEEN THE NUMBER OF ELECTRONS FILLED INTO THE SUB SHELLS AND THE NUMBER OF ELEMENTS IN A PERIOD In the long form of periodic table, there are seven periods according to the number of main shells being filled up by electrons. Each period starts with filling up of a new quantum level. In each period, the number of elements is equal to the number of electrons filled into the subshells as shown below. Total number of elements in the period 2 8 8 18 18 32 incomplete

CLASSIFICATION INTO S, P, D & F BLOCKS The elements in the periodic table are divided into s, p, d & f -blocks based on the type of sub shell into which the differentiating electron enters. 1) s - block * The differentiating electron enters into s - orbitals of outer shell in s - block elements. * The general outer electronic configuration is ns12 * These elements are present at the left hand side of the periodic table. * IA group [Alkali metals] and IIA group [alkaline earth metals] elements belong to this block. * These elements are strong metals and good reducing agents. * The common oxidation states are +1 & +2 respectively for IA and IIA group elements. 2) p - block * In p-block elements, the differentiating electrons enter into p-orbitals of outer shell. * The general electronic configuration is ns 2 np1 6 * The IIIA to 'zero' group elements belong to this block. These are present at the right hand side of the periodic table.

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Group III A Boron family ns2np1 Group IV A Carbon family ns2np2 Group V A Nitrogen family ns2np3 Group VI A Oxygen family ns2np4 Group VII A Halogen family ns2np5 Group O Inert gases ns2np6 (except He which has 1s2 configuration) * These elements constitute nonmetals, metals and metalloids and inert gases. * These elements can exhibit both negative and positive states. 3) d - block * In d - block elements, the differentiating electron enters into d - orbitals of (n-1) shell. * The general outer electronic configuration is (n-1)d1-10 ns1,2

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* These elements are present in between s and p-block elements. * Depending on the principal quantum number of d - orbital into which the last electron enters, these elements are again divided into following series. 3d series Sc to Zn 4th Period 4d series Y to Cd 5th Period 5d series La, Hf to Hg 6th Period 6d series Ac - (incomplete) 7th Period * All the d-block elements are metals and are usually called as transition elements. * These elements show variable oxidation states, paramagnetism and color. They can effectively form complex compounds, alloys and non-stoichiometric compounds. * These elements and their compounds are good catalysts. E.g., 1) Ni in hydrogenation of oils 2) Fe / Mo in Haber's process of synthesis of NH3. 3) Platinized asbestos in the manufacture of H2SO4. * These elements can form alloys effectively. E.g., Brass, Bronze, German silver, etc., 4) f -block * In f-block elements, the differentiating electrons enter into f-orbitals of (n-2) shell. * The general outer electronic configuration is (n-2)f 1-14 (n-1)d 0 or 1ns 2 * These elements belong to IIIB group and are placed below the table separately these are also called as inner transition elements. * The 14 elements from Cerium(Ce) to Lutetium (Lu) which follow the elements Lanthanum are called as Lanthanoids or rare earth elements. In these elements the differentiating electrons enter into 4f-orbitals and hence are also called as 4f -series. * The 14 elements from Thorium (Th) to Lawrencium (Lr), which follow Actinium are called Actinoids. In these element the differentiating electrons enter into 5f-orbitals and hence these are also known as 5f-series. * All of these elements are radioactive. The elements after Uranium are called trans uranic elements. These are artificially prepared elements. * These elements show +3 common oxidation state.

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p-block

s-block d-block

f-block

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Classification based on chemical properties The elements in the periodic table are classified into four groups based on their chemical properties and the electronic configuration.

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1) Noble gas elements : * The zero group elements are called noble or inert gases (He, Ne, Ar, Kr, Xe, Rn). * The general outer electronic configuration is ns 2 np6 except for He (1s2).

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* In these elements, all the shells are completely filled. * Due to stable electronic configuration, these elements do not participate in bond formation. * These elements exist as mono atomic gases.

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2) Representative elements : * The s & p-block elements except zero group elements are called representative elements. * The outer electronic configuration is ns1-2 np 0 5

* In these elements, the outer shell is incompletely filled. * This group includes metals, nonmetals and metalloids. * These elements can share or gain or lose electrons to get octet configuration. 3) Transition elements : * The d-block elements except II B group (Zn, Cd and Hg), are called as transition elements. * In these elements the last two shells (n and n-1) are incompletely filled * The general outer electronic configuration is (n-1)d1-10 ns1,2 * Due to small size, high nuclear charge and presence of incompletely filled d-orbitals, these elements show characteristic properties as given below. i) These elements are very hard and heavy metals with high melting and boiling points. ii) These are good conductors of heat and electricity. iii) They exhibit variable oxidation states. E.g., Fe exhibits +2 and +3 oxidation states. iv) The transition metal ions show colors due to d-d transitions. v) They show paramagnetism due to presence of unpaired electrons. vi) Transition metals and their compounds are good catalysts. E.g., 1) Ni in hydrogenation of oils. 2) Fe / Mo in Haber's process of synthesis of NH3. 3) Platinized asbestos in the manufacture of H2SO4.

Periodic Table

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vii) These elements can form alloys effectively. E.g., Brass, Bronze, German silver, etc., 4) Inner Transition elements : * The f-block elements are called inner transition elements. * In these elements the last three shells (n, n-1 and n-2) are incompletely filled. * The general outer electronic configuration is (n-2)f 1-14 (n-1)d 0 or 1ns 2 .

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* These elements belong to IIIB group and are placed below the table separately. * The 14 elements from Cerium(Ce) to Lutetium (Lu) which follow the elements Lanthanum are called as Lanthanoids or rare earth elements. In these elements the differentiating electrons enter into 4 f-orbitals and hence are also called as 4f -series. * The 14 elements from Thorium (Th) to Lawrencium (Lr), which follow Actinium are called Actinoids. In these elements the differentiating electrons enter into 5f-orbitals and hence these are also known as 5f-series. * All of these elements are radioactive. The elements after Uranium are called trans uranic elements. These are artificially prepared elements. * These elements show +3 common oxidation state.

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PERIODICITY The recurrence of similar chemical and physical properties at regular interval of atomic numbers in a periodic table is called Periodicity. Periodic Properties : The properties which show periodicity. E.g., (i) Density (ii) Melting and boiling points (iii) Hardness (iv) Conductivity (v) Magnetic property (vi) Atomic radius (vii) Ionic radius (viii) Ionization potential (ix) Electronegativity (x) Electron affinity (xi) Electropositivity (xii) Valency (xiii) Oxidation number (xiv) Metallic and non metallic nature (xv) Nature of oxides ATOMIC RADIUS : The average distance between the centre of the nucleus and the electron cloud of outermost orbit is called atomic radius. There are three types of atomic radii based on the nature of bonding. (i) Crystal radius ( atomic radius ) : The half of inter nuclear distance between two adjacent atoms in a metallic crystal. It is applicable to metals. e.g. 1) The distance between two Na atoms is 3.72 A0. Hence its crystal radius is

3.72 =1.86 A o . 2

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2) The distance between two Cs atoms is 5.24 A0. Hence its crystal radius is

5.24 =2.62 A o . 2

(ii) Covalent radius : The half of inter nuclear distance between two atoms held together by a covalent bond. 0.99A 0

Bond length of Cl 2 2

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E.g., covalent radius of 'Cl' atom =

1.98 = 0.99A 0 2 (iii) van der Waal's radius : The half of inter nuclear distance between two closest atoms of different molecules attracted by van der Waal's forces. Usually van der Wall's radii are 40% greater than crystal or covalent radii.

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=

3.6 = 1.86A 0 2

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E.g., van der Wall's radius of Cl =

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1.

86

A

0

Periodicity in atomic radius In groups - Atomic radius increases down the group as the differentiating electron enters into the new quantum shell. This outweighs the increase in nuclear charge. E.g., The order of atomic radius in IA group is Li < Na < K< Rb < Cs < Fr In periods : Atomic radius decreases across the period from left to right as the nuclear charge and atomic number increase. The differentiating electron enters into same shell. In a given period IA group element is bigger in size and VII A group element is smaller in size. The abnormal increase in case of zero group element is due to its van der Waal's radius. E.g., Order of atomic size in 2nd period Li > Be > B > C > N > O > F < Ne In transition elements : The atomic size decreases slightly across the period in d-block elements due to shielding effect of inner d- electrons. In Lanthanoids : The atomic radius decreases with increase in atomic number in Lanthanoids due to poor shielding effect of inner f - electrons. This is called Lanthanoid contraction.

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Consequences of Lanthanoid contraction i) Lanthanoids possesses similar crystalline structures and hence their separation is difficult. ii) Elements of 4d and 5d series show more similarity in their properties. E.g., Zirconium (Zr), ( 4d series) and Hafnium (Hf), ( 5d series) have almost same atomic radii.

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IONIC RADIUS * The ionic radius decreases with increase in effective nuclear charge. * In case of cations, the ionic radius decreases with increase in the positive charge. A+ > A2+ > A3+ .............. * In case of anions, the size increases with increase on the negative charge. A- < A2- < A3- .............. * In a group, for same type of ions the ionic radius increases from top to bottom. E.g., The ionic radii of M+ ions in I A group elements increase in the following order. Li+ < Na+ < K+ < Rb+ < Cs+ < Fr+ * In a given period, the ionic radius decreases with increase in effective nuclear charge for isoelectronic ions. C4- > N3- > O2- > FNuclear charge 6 7 8 9 No. of electrons 10 10 10 10

A

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IONIZATION ENERGY (I.E) The minimum amount of energy required to remove an outer most electron from a gaseous neutral atom in the ground state is called first ionization energy. It is represented by I.E1. +

I.E1

 

A+

+

1e-

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neutral atom The minimum amount of energy required to remove an outer most electron from a gaseous unipositive ion is called second ionization energy (I.E2).

A+

+

I.E 2

 

A 2+

+

1e -

unipositive ion

Second ionization energy is always greater than the first as the effective nuclear charge increases from A to A+. In general I.E1 < I.E2 < I.E3 < I.E4 _ _ _ _ _ Factors affecting ionization energy i) Atomic radius Ionization energy decreases with increase in atomic radius as the nuclear attraction over outer electron decreases. Ionization energy 

1 Atomic size

E.g., In IA group, ionization energies decreases with increase in atomic radius down the group. Order ionization energy Li > Na > K > Rb > Cs > Fr 2) Nuclear charge : Ionization energy increases with increase in nuclear charge. I.E  nuclear charge

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3) Screening or shielding effect : Reduction of nuclear attraction over the outer electrons due to presence of inner electrons is called screening or shielding effect. The order of shielding ability for different types of orbitals is given below. s>p>d>f Ionization energy decreases with increase in screening effect I.E 

1 Screening effect

4) Penetration power : Ability to come closer towards nucleus is called penetration power. Greater the penetration power, greater is the ionization energy. I.E  Penetraction power

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Penetration power of different orbitals is given below. s>p>d>f 5) Electronic Configuration Atoms with ns2 np6 configuration or with half filled or completely filled sub shells are extra stable. Hence their ionization energies are very high. E.g. IIA group elements with ns2 (completely filled) and VA group elements with ns2 np3 (half filled) configurations possess higher ionization potentials.

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Periodicity in ionization energy In groups Ionization energy decreases down the group as the atomic radius increases. E.g. The order of ionization energy in IA group elements is Li > Na >K > Rb >Cs > Fr

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In periods : In general, the ionization energy increases from left to right in a period with decrease in atomic radius. But this increment is not regular. In a given period, IIA (ns2) and VA (ns2np3) group elements have higher ionization energies than elements in their next groups i.e., III A and VI A respectively. It is due to stable electronic configurations. E.g., The order of ionization energies in 2nd period is given below Li < Be > B < C < N > O < F << Ne 2s2 2s22p3 2s22p6 2 2 3 'Be' with 2s (full filled) and 'N' with 2s 2p (half filled) configurations are more stable and hence possess higher ionization potentials than 'B' and 'O' respectively.

Ionization Energy

Ne F

2s22p3 N C

2s2 Be

O

B Li Atomic number(Z)

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In a given period alkali metals possess low ionization energies and inert gases possess very high ionization energies. Cs ---- Element with lowest ionization energy He ---- Element with highest ionization energy. In d-block elements : Due to shielding effect of inner d-electrons, the ionization energy is increased slowly across the period in transition elements. ELECTRON GAIN ENERGY (ELECTRON AFFINITY) The amount of energy released when an electron is added to a neutral gaseous atom in the ground state is called electron gain energy.

A

1e-



 

A

 -EA1 

neutralatom

A

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Addition of another electron to uni negative ion is difficult due to repulsions. Hence energy is absorbed during the addition of second electron. Hence second electron affinity is positive.



 

A 2-

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uninegative ion

1e -

 +EA2 

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Electron affinity values are not measured directly but can be indirectly determined from BornHaber cycle.

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Factors effecting electron affinity : * Electron affinity values are higher for smaller atoms with high effective nuclear charge. * The elements with stable electronic configurations possess very low electron affinities. * In the periodic table the element with highest electron affinity is Chlorine (Cl). The inert gases have zero electron affinities. Periodicity in electron affinity In groups In general, the electron affinity decreases with increase in size of the atom from top to bottom in a group. But due to small size, the first element possesses lower electron gain energy than the second element in a given group. E.g., In halogens, fluorine atom is smaller in size and hence the newly added electron experiences repulsion from the electrons in the atom. Hence its electron affinity is less than that of chlorine (Cl) In Periods Generally electron affinities increase from left to right in a period with decrease in atomic size. But this increment is not regular. In a given period, the electron affinities of elements with stable electronic configurations are very low. E.g. The order of electron gain energies in 2nd period is given below Li > Be> B< C>N>Ne Be (2s2) , N (2s2 2p3) and Ne (2s22p6) have stable configurations and hence possess low electron affinity values.

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F

Electron affinity

O C

B Li

N 2s22p3

Be 2s2

Ne

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Atomic number(Z)

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ELECTRONEGATIVITY The tendency of the atom of an element to attract the shared electron pair towards itself in a hetero nuclear diatomic molecule is called Electronegativity. Electronegativity values can be expressed by using following two scales.

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1) Pauling scale : In this scale, the electronegativity values of elements are calculated from the bond energies. The difference in electronegativity of atom A & B in a molecule AB can be given as follows. X A - X B  0.208  AB K .cal / mol

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(or)

X A - X B  0.1017

where

 AB K . J / mol

X A  Electronegativity of A X B  Electronegativity of B  AB  polarity of bond A - B  AB  E A-B  E A-B

E A-B = Experimental bond energy of AB E A-B = The bond energy of AB E A-B 

E AA .E BB

On Pauling scale fluorine has highest electronegativity value of 4.0 Element Electronegativity F 4.0 O 3.5 N 3.0 Cl 3.0 H 2.1 C 2.5 Mulliken Scale : According to this scale, electronegativity is the average of ionization energy and electron affinity.

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Ionization energy + Electon affinity 2 In this scale, the electronegativity values are 2.8 times larger than those of pauling values. The electronegativity values equal to those in pauling scale are obtained by using following formulae. Electronegativity =

Electronegativity =

 I.E in K J / mole 

+  E.A in K.J / mol  544

Periodicity in electronegativity Electronegativity decreases down the group as atomic size increases. Whereas it increases from left to right in a period since the atomic size decreases.

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Applications : * The nature of bond formed between two atoms can be predicted from their electronegativity difference. If the electronegativity difference is 1.7, the bond would be more than 50% ionic. If the electronegativity difference is less than 1.7, the bond formed will be more than 50% covalent. * From the electronegativity values, proper chemical formulae of compounds can be written. * It is possible to calculate oxidation states by comparing electronegativity values.

Valency Hydride

IA 1 LiH

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VALENCY Valency indicates the combing power of an atom. It is equal to the number of hydrogen atoms or double the number of oxygen atoms with which the atom of an element combines. The valency with respect to hydrogen is equal to group number in IA to IVA groups. But in VA to Zero group, valency with respect to hydrogen is equal to (8 - group number). valency with respect to hydrogen. IIA IIIA IVA VA VIA 2 3 4 3 2 BeH2 B2H6 CH4 NH3 H2O

VIIA 1 HF

O O -

The maximum valency with respect to oxygen or fluorine is equal to the group number

Valency Oxide Fluoride

Valency with respect to oxygen or fluorine IA IIA IIIA IVA VA VIA VIIA 1 2 3 4 5 6 7 Na 2 O MgO Al2O3 CO 2 P2 O 5 SO 3 Cl2O 7 NaF MgF2 AlF3 CF4 PF5 SF6 IF7

O O -

* An element can show more than one valency (multiple valency) OXIDATION STATES (OR) OXIDATION NUMBERS The formal charge acquired by an atom in a given species is called oxidation state or oxidation number Periodicity The elements in a given group exhibit mostly one oxidation number.

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* The maximum positive oxidation state shown by IA to VIIA group elements is equal to the group number. E.g.,

Maximum positive oxidation number

IA

IIA

IIIA

IVA

VA

VIA

VIIA

+1

+2

+3

+4

+5

+6

+7

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* The elements of groups from IVA to VIIA also exhibit negative oxidation states equal to group number -8. E.g., IVA VA VIA VIIA Oxidation number -4 -3 -2 -1 * The p-block elements can show variable oxidation states-both negative and positive. E.g., Phosphorus can show +3 (PCl3) and +5 (PCl5) oxidation states. Inert pair effect : The reluctance of the electron pair in the outer 'ns' orbital to get unpaired and involved in the bond formation is called inert pair effect. Inert pair effect is caused due to poor shielding effect of inner f-electrons. Due to inert pair effect, the heavier elements of IIIA, IVA and VA group elements are more stable in oxidation numbers less by two units of the group oxidation number. E.g., (i) Thallium (Tl) is more stable in +1 oxidation state then in +3 state. (ii) Tin (Sn) and Lead (Pb) are more stable in +2 oxidation state than in +4 oxidation state. (iii) Bismuth (Bi) is more stable in +3 state than in +5 state. * d-block elements show variable oxidation numbers. Their common oxidation state is +2 due to presence of two electrons in the ns orbital. * Ruthenium (Ru) and Osmium (Os) can show a maximum oxidation state of +8. * Lanthanoids show +3 as common oxidation state. ELECTROPOSITIVITY The tendency of an element to lose electrons is called electropositivity. Electropositivity increases with increase in the atomic size. It increases down the group and decreases across the period from left to right. Highly electropositive elements are called metals. They possess low ionization energies. IA and IIA group elements are highly electropositive elements. Metallic and Non metallic nature: * The elements with low electronegativity (high electropositivity) are called metals. * The elements with high electronegativity are called non metals. Periodicity * In a given group metallic nature increases from top to bottom with decrease in electronegativity. * In a given period non metallic nature increases from left to right with increase in electronegativity.

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* In a given period IA and IIA group elements are stronger metals whereas halogens are stronger non metals. NATURE OF OXIDES * Oxides are the binary compounds of elements with oxygen. These are of three types as follows. i) Basic oxides : These oxides dissolve in water by giving alkaline solutions. E.g., Na2O, CaO etc....... Na 2 O + H 2 O   2 NaOH

CaO + H 2O   Ca  OH 2 Metal oxides are usually basic in nature ii) Acidic oxides : These oxides dissolve in water by giving acidic solutions. E.g., CO2 , SO3 , Cl2O7 etc.,

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CO 2 + H 2 O   H 2 CO 3 Cl 2 O 7 + H 2 O   2HClO 4 SO 3 + H 2 O   H 2SO 4

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Non metal oxides are usually acidic in nature. iii) Amphoteric oxides: These oxides show both acidic and basic properties. E.g., BeO, Al2O3 etc....

Group Oxides Nature

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Periodicity in nature of oxides In a group, the basic nature of oxides increases down the group. E.g., The increasing order of basic nature in oxides of IA group elements is Li2O < Na2O < K2O < Rb2O < Cs2O In a period, the acidic nature of oxides increases from left to right. E.g., In third period, the nature of oxides vary as follows. IA Na2O

IIA MgO

Strong Base

Weak Base

IIIA Al2O3

IVA SiO2 Very weak acid

Amphoteric

VA P4O10

VIA SO3

Weak acid

Strong acidic

VIIA Cl2O7 Very strong acid

DIAGONAL RELATIONSHIP In the periodic table, an element of second period in a group is similar in properties with the third period element of next group. This is known as diagonal relationship. E.g., Following diagonally placed elements exhibit similar chemical properties. Group

IA

IIA

IIIA

IVA

2nd period

Li

Be

B

C

3rd period

Na

Mg

Al

Si

Diagonally related elements possess similar ionic sizes, similar electronegativities and same polarizing power.

Periodic Table

Prepared by V. Aditya vardhan adichemadi @ gmail.com

15

Polarizing power =

ionic charge ( ionic radius )2

E.g., Be and Al exhibit similar properties as given below. * BeO and Al2O3 amphoteric oxides. * Carbides of both the elements produce methane (CH4) gas on hydrolysis.   4Al(OH)3 Al4C3 + 12H2O + 3CH4 +

 

4H2O

2 Be(OH)2

+

V. AD VA I AG T

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Pr ep ar ed

DE by VI W JU AR N AN IOR G C AL O LL EG E

VA RD HA N

Be2C

CH4