Quantum+No[1].+&+Periodic+Table

CHEMISTRY LECTURE NOTES

COURSE - VIJAY (R)

(LECTURE No. 1 TO 3)

TOPIC : QUANTUM NUMBERS & PERIODIC

PROPERITES

LECTURE # 1

QUANTUM NUMBERS FOR ELEMENTS :

Orbital :

An orbital may be defined as the region of space around the nucleus where the probability of finding an electron is maximum (90% to 95%)

Orbitals do not define a definite path for the electron, rather they define only the probability of the electron being in various regions of space around the nucles.

Shape of the orbitals :

Shape of the orbitals are related to the solutions of Schrodinger wave equation, and gives the space in which the probability of finding an electron is maximum.

s- orbital :

Shape

spherical

s- orbital is non directional and it is closest to the nucleus, having lowest energy. s-orbital can accomodate maximum no. of two electrons.

p-orbital : Shape

dumb bell

Dumb bell shape consists of two loops which are separated by a region of zero probability called node.

d - Orbital :

Shape

double dumb bell

d - orbital can accomodate maximum no. of 10 electrons.

f - orbital :

Shape

leaf like

f - orbital can accomodate maximum no. of 14 electrons.

Quantum Numbers :

The set of four numbers required to define an electron completely in an atom are called quantum numbers. The first three have been derived from Schrodinger wave equation.

(i) Principal quantum number (n) : (Proposed by Bohr)

It describes the size of the electron wave and the total energy of the electron. It has integral values 1, 2, 3, 4 ...., etc., and is denoted by K, L, M, N. ..., etc.

* Number of subshell present in nth _{shell = n}

n subshell

1 s

2 s, p

3 s, p, d

4 s, p, d, f

* Number of orbitals present in nth _{shell = n}2 _.

* The maximum number of electrons which can be present in a principal energy shell is equal to 2n2_.

No energy shell in the atoms of known elements possesses more than 32 electrons. * Angular momentum of any orbit =



2 nh

(ii) Azimuthal quantum number () : (Proposed by Sommerfield)

It describes the shape of electron cloud and the number of subshells in a shell. * It can have values from 0 to (n – 1)

* value of  subshell

0 s

1 p

2 d

3 f

* Number of orbitals in a subshell = 2 + 1

* Maximum number of electrons in particular subshell = 2 × (2 + 1)

* Orbital angular momentum L =

 2 h ) 1 (  =  (1)         2 h 

i.e. Orbital angular momentum of s orbital = 0, Orbital angular momentum of p orbital =



2 h

2 _,

Orbital angular momentum of d orbital =



2 h 3

(iii) Magnetic quantum number (m) : (Proposed by Linde)

It describes the orientations of the subshells. It can have values from _– to +  including zero, i.e.,

total (2 + 1) values. Each value corresponds to an orbital. s-subshell has one orbital, p-subshell three

orbitals (p_x, p_y and p_z), d-subshell five orbitals (d_xy,d_yz,d_zx,d_x2 _y2,d_z2)

 and f-subshell has seven orbitals.

The total number of orbitals present in a main energy level is ‘n2’.

(iv) Spin quantum number (s) : (Proposed by Goldschmidt & Uhlenbeck)

It describes the spin of the electron. It has values +1/2 and –1/2. (+) signifies clockwise spinning and

(–) signifies anticlockwise spinning.

* Spin magnetic moment

eh

 s(s1) or



* It represents the value of spin angular momentum which is equal to

 2 h ) 1 s ( s 

* Maximum spin of atom = 2 1

x No. of unpaired electron. Ex. (NCERT)

What is the total number of orbitals associated with the principal quantum number n = 3 ?

Sol. For n = 3, the possible values of  are 0, 1 and 2. Tthere is one 3s orbital (n = 3,  = 0 and m = 0) ; there are

three 3p orbitals (n = 3,  = 1 and m

= – 1, 0, + 1) ; there are five 3d orbitals

(n = 3,  = 2 and m

= – 2, – 1, 0, + 1+, + 2).

Therefore, the total number of orbitals is 1 + 3 + 5 = 9

The same value can also be obtained byusing the relation; number of orbitals = n2_{, i.e. 3}2 _{= 9.}

Ex. (NCERT)

Using s, p, d, f notations, describe the orbital with the following quantum numbers (a) n = 2,  = 1, (b) n = 4, = 0, (c) n = 5,  = 3, (d) n = 3, = 2 Sol. n  orbital (a) 2 1 2p (b) 4 0 4s (c) 5 3 5f (d) 3 2 3d

Ex. Find orbital angular momentum of an electron in (a) 4s subshell and (b) 3p subshell Ans. (a) 0 (b)  2 h 2

Ex. Orbital angular momentum of an electron in a particular subshell is



h

5 then find the maximum number of electrons which may be present in this subshell.

Sol. Orbital angular momentum = (1)

 2 h  (1)  2 h = 5  h ) 1 (  = 2 5 = 20   = 4.

hence maximum number of electrons in this subshell = 2(2 + 1) = 18. Ans. 18

Ex. Which of the following set of quantum numbers is not valid. (A) n = 3, l = 2, m = 2, s = + 2 1 (B) n = 2, l = 0, m = 0, s = – 2 1 (C) n = 4, l = 2, m = –1, s = + 2 1 (D*) n = 4, l = 3, m = 4, s = – 2 1 ( m >  is not possible)

ELECTRONIC CONFIGURATION :

PAULI’S EXCLUSION PRINCIPLE :

No two electrons in an atom can have the same set of all the four quantum numbers, i.e., an orbital cannot have more than 2 electrons because three quantum numbers (principal, azimuthal and magnetic) at the most may be same but the fourth must be different, i.e., spins must be in opposite directions.

AUFBAU PRINCIPLE :

Aufbau is a German word meaning building up. The electrons are filled in various orbitals in order of their increasing energies. An orbital of lowest energy is filled first. The sequence of orbitals in order of their increasing energy is :

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, .... The energy of the orbitals is governed by (n + ) rule.

‘n + Rule :

The relative order of energies of various sub-shell in a multi electron atom can be predicated with the help of ‘n + ’ rule

 The sub-shell with lower value of (n + _) has lower energy and it should be filled first.

eg. 3d 4s

(n +) = 3 + 2 (n +) = 4 + 0

= 5 = 4

Since, (n + ) value of 3d is more than 4s therefore, 4s will be filled before 3d.

 If two sub-shell has same value of (n + _) then the sub-shell with lower value of n has lower energy and it

should be filled first.

eg. 3d 4p (n +) = 3 + 2 = 4 + 1 = 5 = 5 3d is filled before 4p. MEMORY MAP : 1s 2s 3s 4s 5s 6s 2p 3p 4p 5p 6p 3d 4d 4f 5d 5f 1 – s 2 3 – s,p 4 – s,p – s,d,p 5 – s,d,p 6 – s,f,d,p 6 – s,f,d,p 7 – s,f,d,p

HUND’S RULE :

No electron pairing takes place in the orbitals in a sub - shell until each orbital is occupied by

one electron with parallel spin. Exactly half filled and fully filled orbitals make the atoms more stable, i.e., p3_,

p6_{, d}5_{, d}10_{, f}7 _{and f}14 _{configuration are most stable.}

Ex. Write the electronic configuration and find the no. of unpaired electrons as well as total spin for the following atoms : (1) ₆C (2) ₈O (3) ₁₅P (4) ₂₁Sc (5) ₂₆ Fe (6) ₁₀Ne (i) ₆C  1s2, 2s2,2p2

No. of unpaired electrons  2.

Total spin = 2 2  or 2 2  (ii) ₈O  1s2, 2s2, 2p4 1s 2s 2p

 No. of unpaired electrons = 2

Total spin = 2 2  or 2 2  (iii) ₁₅P 1s2, 2s2, 2p6, 3s2, 3p3 3s 3p

 No. of unpaired electrons = 3

Total spin = 2 3  or 2 3  (iv) ₂₁Sc  1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d1 or [Ar] 4s2 _3d1 3d 4s [Ar] 3d1 _4s2

 No. of unpaired electrons = 1  Total spin = 2 1  or 2 1  (v) ₂₆Fe  1s2, 2s2 2p6, 3s2, 3p6, 4s2 3d6 or [Ar] 4s2_{, 3d}6 3d 4s

No. of unpaired electrons = 4

 Total spin = 2 4  or 2 4 

(vi) ₁₀Ne  1s2, 2s2 2p6

No. of unpaired electrons = 0 Total spin = 0

EXCEPTIONS :

(1) ₂₄Cr = [Ar] 4s2_{, 3d}4 _{(Not correct)}

[Ar] 4s1_{, 3d}5 _{(correct : as d}5 _{structure is more stable than d}4 _structure)

(2) ₂₉Cu = [Ar] 4s1_{, 3d}10 _{(correct : as d}10 _{structure is more stable than d}9 _structure)

Ex. Find the electronic configuration of Fe2+ _{and Cr}3+ _{and their magnetic moments.}

Ans.

(i) Fe  [Ar] 3d6, 4s2

Fe2+

[Ar] 3d6

3d

No. of unpaired electrons = 4

 Magnetic moment = n(n2) B.M. = 4(42)B.M. = 46 B.M. = 24 B.M. = 2 6 B.M. (ii) Cr  [Ar] 3d5, 4s1 Cr3+  [Ar] 3d3  No of unpaired electrons = 3  Magnetic moment = n(n2) B.M. = 3(32)B.M. = 35 B.M. = 15 B.M.

Ex. Write all four quantum numbers for the last electrons of Na 1s2_{, 2s}2_,2p6_{, 3s}1 n = 3  = 0 m = 0 s = ± 2 1 Home - Work NCERT (Reading) 2.6.4, 2.6.5 NCERT (Exercise) 2.23 - 2.31, 2.62 - 2.67

LECTURE # 2

COMPLETE THE THEORY CONTENT GIVEN BELOW UP TO THE EFFECTIVE NUCLEAR CHARGE IN LIMITTED TIME OF 10 TO 15 MINUTE :

Mendeleev’s Periodic Table : Mendeleev’s Periodic’s Law

According to him the physical and chemical properties of the elements are a periodic function of their atomic masses.

He arranged then known elements in order of their increasing atomic masses considering the facts that elements with similar properties should fall in the same vertical columns and leaving out blank spaces where necessary.

The table is divided into nine vertical columns called groups and seven horizontal rows called periods.

Periods No. of Elements Called as

(1)st n = 1 2 Very short period

(2)nd n = 2 8 Short period

(3)rd n = 3 8 Short period

(4)th n = 4 18 Long period

(5)th n = 5 18 Long period

(6)th n = 6 32 Very long period

(7)th n = 7 19 Incomplete period

The groups were numbered as I, II, III,IV,V,VI,VII,VIII and Zero group

Long form of the Periodic Table or Moseley’s Periodic Table

He studied (1909) the frequency of the X-ray produecd by the bombardment of a strong beam of electrons on metal target. He found that the square root of the frequency of X-rays ( _) is directly proportional to number of effective nuclear charge (z) of metal i.e. to atomic number and not to atomic mass of the atom of that metal.(as nuclear charge of metal atom is equal to atomic number)

i.e. ( _) = a (z - b) (Will be taught to you in Physics(Modern Physics) later)

Where ‘a’ is the proportionality constant and ‘b’ is a constant for all the lines in a given series of X-rays.

Therefore, he, concluded that atomic number was a better fundamental property of an element than its atomic weight He suggested that the atomic number (z) instead of atomic weight should be basis of the classification of the elements.

Modern Periodic Law (Moseley’s Periodic Law)

Physical and chemical properties of elements are the periodic functions of their atomic number.If the elements are arranged in order of their increasing atomic number, after a regular interval ,element with similar properties are repeated.

Periodicity

The repetition of the properties of elements after regular intervals when the elements are arranged in the order of increasing atomic number is called periodicity.

Cause of Periodicty:

The periodic repetition of the properties of the elements is due to the recurrence of similar valence shell electronic configuration after certain regular intervals. For example ,alkail metals have same electronic

configuration ns

The long form of periodic table is the contribution of Range , Werner, Bohr and Bury

This table is also referred to as Bohr‘s table since it follows Bohr’s scheme of the arrangements of elements

into four types based on electronic configuration of elements

Periods :

There are seven periods numbered as 1, 2, 3, 4, 5, 6 and 7.

(i) Each period consists of a series of elements having same valence shell.

(ii) Each period corresponds to a particular principal quantum number of the valence shell present in it. (iii) Each period starts with an alkali metal having outermost electronic configuration ns1_.

(iv) Each period ends with a noble gas with outermost electronic configuration ns2_np6 _{except helium having}

outermost electronic configuration 1s2_.

(v) Each period starts with the filling of new energy level.

(vi) The number of elements in each period is twice the number of atomic orbitals available in energy level that is being filled. To illustrate



for which,

n = 1,  = 0 (s-subshell) and m = 0.

Only one orbital (1s) is available and thus it contains only two elements.



For which,

n = 3,  = 0, 1, 2 and m = 0, 3, 5

no. of orbitals 1 3 5 (3s) (3p) (3d)

––––––––––––––––

Total no. of orbitals 9

––––––––––––––––

But the energy of 3d orbitals are higher than 4s orbitals. Therefore, four orbitals (one 3s and three 3p orbitals) corresponding to n = 3 are filled before filling in 4s orbital (next energy elevel). Hence 3rd _{period contains eight}

elements not eighteen elements. Groups :

There are eighteen groups numbered as 1, 2, 3, 4, 5, ... 13, 14, 15, 16, 17, 18. Group consists of a series of elements having similar valence shell electronic configuration.

CLASSIFICATION OF THE ELEMENTS :

It is based on the type of orbitals which receives the differentiating electron (i.e., last electron).

(a) s-block elements

When shells upto (n – 1) are completely filled and the last electron enters the s-orbital of the outermost (nth)

shell, the elements of this class are called s-block elements.

 Group 1 & 2 elements constitute the s-block.  General electronic configuration is [inert gas] ns1-2

 s-block elements lie on the extreme left of the periodic table.  This block includes metals.

(b) p-block elements

When shells upto (n – 1) are completely filled and differentiating electron entres the p-orbital of the nth orbit,

elements of this class are called p-block elements.

 Group 13 to 18 elements constitute the p-block.  General electronic configuration is [inert gas] ns2 np1-6

 p-block elements lie on the extreme right of the periodic table.  This block includes some metals, all non-metals and metalloids.

 s-block and p-block elements are collectively called normal or representative elements.

(c) d-Block elements

When outermost (nth_{) and penultimate shells (n}

– 1)th shells are incompletely filled and differentiating electon

enters the (n – 1) d orbitals (i.e., d-orbital of penultimate shell) then elements of this class are called d-block

elements.

 Group 3 to 12 elements constitute the d-block.

 General electronic configuration is [inert gas] (n–1) d1-10 ns0-2

 d-block elements are classified into four series

Series Elements (n – 1)d being filled

3d ₂₁Sc – 30Zn 3d

4d ₃₉Y – 48Cd 4d

5d ₅₇La, ₇₂Hf – 80Hg 5d

6d ₈₉Ac, ₁₀₄Rf – 112Uub 6d (incomplete series)

 Those elements which have partially filled d-orbitals in neutral state or in any stable oxidation state

are called transition elements

(d) f-Block elements

When n, (n – 1) and (n – 2) shells are incompletely filled and last electron enters into f-orbital of antepenultimate

i.e., (n – 2)th shell, elements of this class are called f-block elements., General electronic configuration is

(n – 2) f1-14(n – 1) d0, 1 ns2

 All f-block elements belong to 3rd group.  They are metals

 Within each series, the properties of the elements are quite similar.  The elements coming after uranium are called transuranium elements.

 They are also called as inner-transition elements as they contain three outer most shell

incomplete and were also referred to as rare earth elements since their oxides were rare in earlier days.

The elements of f-blocks have been classified into two series.

 The actinides and lanthanides have been placed at the bottom of the periodic table to avoid the

undue expansion of the periodic table.

1. st inner transition or 4 f-series, contains 14 elements 58Ce to 71Lu. Filling of electrons takes place

in 4f subshell.

2. IInd inner transition or 5 f-series, contains 14 elements ₉₀Th to ₁₀₃Lr. Filling of electrons takes place in 5f subshell.

PREDICTION OF PERIOD, GROUP AND BLOCK :







Group number = the no. of valence electrons (b) For p-block elements

Group number = 10 + no. of valence electrons (c) For d-block elements

Group number = no. of electrons in (n – 1) d sub shell + no. of electrons in valence shell.



where n = number of unpaired electrons.

Que. A particular atom having atomic number between 22 to 30 has magnetic moment equal to 3.73 BM. Then find the atomic number of the element which is just below it in the periodic table. Also predicts the group number, block and period of the former element.

Sol. We know that magnetic moment = n(n2) = 3.73

 n (number of unpaired electron) = 3

Since atom has three unpaired electrons hence it must be ₂₇Co.

therefore the atomic number of the element which is just below it is = 27 + 18 = 45

TYPICAL ELEMENTS:

 Third period elements are called as typical elements. These include Na, Mg, Al, Si, P, S, Cl.  The properties of all the elements belonging to a particular group resemble the properties of the

corresponding typical element of that group. For example, the general properties of alkali metals (A) can be predicted from the properties of Na, not Li, the first member of the group.

 The properties of the elements of second period differ in the many respect belonging to the same

group due to the smaller atomic size and absence of vacant d-orbitals.

DIAGONAL RELATIONSHIP :

Some elements of certain groups of 2nd _{period resemble much in properties with the elements of third period}

of next group i.e. elements of second and third period are diagonally related in properties. This phenomenon is known as diagonal relationship.

2nd _{period Li Be B C}

3rd _{period Na Mg Al Si}

Diagonal relationship arises because of (i) similar size of atom and ions

(Li = 1.23 Å & Mg = 1.36 Å ; Li+ = 0.60 Å & Mg2+ = 0.65 Å)

(ii) similar polarising powers (charge to radius ratio) (iii) similarity in electronegativity values

(Li = 1.0 & Mg = 1.2 ; Be = 1.5 & Al = 1.5)

THE PERIODICIT Y OF ATOMIC PROPERTIES :

( I )

EFFECTIVE NUCLEAR CHARGE :

Between the outer most valence electrons and the nucleus of an atom, there exists finite number of shells containing electrons. Due to the presence of these intervening electrons, the valence electrons are unable to experience the attractive pull of the actual number of protons in the nucleus. These intervening electrons act as shield between the valence electrons and protons in the nucleus. Thus, the presence of intervening (shielding) electrons reduces the electrostatic attraction between the protons in the nucleus and the valenece electrons because intervening electrons repel the valence electrons. The concept of effective nuclear charge allows us to account for the effects of shielding on periodic properties.

The effective nuclear charge (Z_eff) is the charge felt by the valence electron. Z_eff is given by Z_eff = Z – .

Where Z is the actual nuclear charge (atomic number of the element) and  is the shielding (screening)

constant.

2 .

ATOMI C RADIU S :

(A) Covalent radius :

It is one-half of the distance between the centres of two nuclei (of like atoms) bonded by a single covalent bond. Covalent radius is generally used for non-metals.

Single Bond Covalent Radius, SBCR (bond length) -(a) For homodiatomic molecules d_A

–A = rA + rA or 2rA so, rA =

A A

d 2



 (b) For heterodiatomic molecules in which electronegativity remains approximately same.

 For heteronuclear diatomic molecule, A–B, where difference between the electronegativity values of atom A

and atom B is relatively larger, d_A

– B = rA + rB– 9.0 

Electronegativity values are given in pauling units and radius in picometers.

 = X_A– X_B where X_A and X_B are electronegativity values of high electronegative element A and less

electrone-gative element B. This formula is given by Stevenson & Schomaker. Note : Modified and more accurate

d_A

– B = rA + rB– 7.0 () 2

This formula was proposed by Porterfield.

Multiplicity of the bond causes a shortening of the bond length. (Source : Huheey and Keiter)

(B) Van der Waals radius (Collision radius) :

It is one-half of the internuclear distance between two adjacent atoms in two nearest neighbouring molecules of the substance in solid state.

 Van der Waal’s radius does not apply to metals. Its magnitude depends upon the packing of the

atoms when the element is in the solid state.

Comparision of covalent radius and van der Waal’s radius

(i) The van der Waal’s force of attractions are weak, therefore, their internuclear distances in case of atoms held

by van der Waal’s forces are much larger than those of between covalently bonded atoms. Therefore van der

Waal’s radii are always larger than covalent radii.

(ii) A covalent bond is formed by the overlaping of two half-filled atomic orbitals, a part of the orbital becomes common. Therefore, covalent radii are always smaller than the van der Waals radii. For example,

Elements H O F S Br

Covalent radius (Å) 0.37 0.66 0.64 1.04 1.11

van der Waal's radius (Å) 1.20 1.40 1.35 1.85 1.95

(C) Metallic radius (Crystal radius) :

It is one-half of the distance between the nuclei of two adjacent metal atoms in the metallic crystal lattice.

 Metallic radius of an element is always greater than its covalent radius. It is due to the fact that

metallic bond (electrical attraction between positive charge of an atom and mobile electrons) is weaker than covalent bond and hence the internuclear distance between the two adjacent atoms in a metallic crystal is longer than the internuclear distance between the covalently bonded atom.

For example : Metallic radius Covalent radius

K 231 pm 203 pm

Na 186 pm 154 pm

 rcovalent < rcrystal < rvander Walls

Variation in a period: On moving left to right due to increased nuclear charge the size decreases. Variation in a group: On moving top to bottom due extra addition of a shell the size increases.

---This portion is not properly discussed so do explain all these a bit slow

Some Irregularities :

(a) The atomic radius of inert gas (zero group) is given largest in a period because it is represented as van der Waals’s radius which is generally larger than the covalent radius. The van der Waal’s radius of inert gases

also increases from top to bottom in a group.

(b) In the transition series (e.g. in first transition series), the covalent radii of the elements decrease from left to right across a row until near the end when the size increases slightly. On moving from left to right, extra protons are placed in the nucleus and the extra electron are added. The orbital electrons shield the nuclear charge incompletely. Thus the nuclear charge attracts all the electrons more strongly, hence a contraction in size occurs. The radii of the elements from Cr to Cu, are very close to one another because the extra electron being added increasses the repulsion between the electrons and counter balances the increased nuclear charge on the outer electrons (4s). As a result of this, the size of the atom does not change much in moving from Cr to Cu and for zinc this repulsion even dominates the nuclear charge so size slightly in-creases.

Element Sc Ti V Cr Mn Fe Co Ni Cu Zn

Atomic radius (Å) 1.44 1.32 1.22 1.18 1.17 1.17 1.16 1.15 1.17 1.25

(c) r_{Al } r_Ga

because of d orbital contraction.

(d) 4d 5d (Zr – Hf)

( lanthanide contraction)

The lanthanide contraction counter balances almost exactly the normal size increase on descending a group of transition elements. Thus covalent and ionic radii of Nb (5th peroid) and Ta (6th period) are almost same due to poor shielding of f-orbitals electrons.

---The following portion has already been taught so just revise and be a bit fast

IONI C RADI US :

 The sizes of ions increases as we go down a group (considering the ions of same charge). For

example :

Li+ _{< Na}+ _{< K}+ _{< Rb}+

Be2+ _{< Mg}2+ _{< Ca}2+ _{< Sr}2+

F– _{< Cl}– _{< Br} – _<

–

 Cation is smaller than parent atom but anion is bigger than parent atom.

 The species containing the same number of electrons but differ in the magnitude of their nuclear

charges are called as isoelectronic species. For example, N3– _{, O}2–_{, F}–_{, Ne, Na}+ _{, Mg}2+ _{and Al}3+ _are

all isoelectronic species with same number of electrons (i.e 10) but different nuclear charges of +7, +8, +9, +10, +11, +12 and +13 respectively.

Within a series of isoelectronic species as the nuclear charge increases, the force of attraction by the nucleus on the electrons also increases. As a result, the ionic radii of isoelectronic species decrease with increases in the magnitude of nuclear charges. For example,

Al3+ _Mg2+ _Na+ _F– _O2– _N3–

Ionic radii increase As effective nuclear charge decrease.

 Zero group elements should not be considered while comparing the size or ionic radii as their atomic

radii are expressed as van der Wall’s radii.

Ionisation Energy ( Ionisation potential)

Min amount of energy required to remove an electron from an isolaled gaseous atom / species. (I.E.)₁ M(g)  M+(g) + e–

H = +ve

(I.E.)₂ M+_{(g)  M}2+ _{(g) + e}–

(I.E.)₃ > (I.E)₂ > (I.E.)₁ – Difficult to remove an electron from a positively charged ion than from a neutral

Factors affecting I.E.

(a) Size Size _  I.E. _

(b) Z_eff Z_{eff }  I.E.  but  (i.e. Zeff ) E

(c) Electronic configuration – For a stable electronic configuration i.e. half filled or fully filled configurations

(stable because of symmetry) I.E. will be larger

(d) The relative extent to which the various orbital penetrate the electron clouds of of other orbitals is s > p > d > f. Thus for any given principal quantum number n, an electron will experience the greatest effective nuclear charge when in s-orbital than p-orbital and so on. Hence the order of I.E. is as follows : s > p > d > f ---Teach following portion a bit slowly

Periodicity

L – R I.E. (Zeff )

T – B I.E.  (Size) rregularities

(a) In  period ( also in  period )

Be > B, N > O, Ne >> F (b) .E._Ga> .E._Al

size same extra nuclear charge _ (c) 5d > 3d > 4d

(lanthanide contraction)

Note : I.E. can be correlated with the reactivity e.g., (a) Noble gases are inert towards chemical reactivity. (b) low I.E. of alkalimetals correlate with their high reactivity.

---LECTURE # 3

Electron Affinity (E.A)

The energy released when an electron is added to an isolated gaseous atom to produce a monovalent anion is known as e.a. and enthalpy change of this process is known as electron gain enthalphy

M(g) + e–   M–(g) _{E.A. ve} ve Heg      energy released ve . A . E ve Heg      energy absorbed

(E.A.)₂ is always +ve – because of electrostatic repulsion between anion and electron (having same charge).

M–_{(g) + e}–



 M2–(g) Heg2 = + ve

Factors affecting E.A.

(a) Size Size_  E.A. 

(b) Zeff Zeff _  E.A. _  _  E.A._

(c) E.C. Stable E.C. will have smaller or –ve E.A

Periodicity

L  R Zeff   E.A.

Irregularities (a) nd period

(b) E.A. 3P > 2P

E.A (low) B C N O F

(added electron goes to the samller n = 2 and suffer significant repulsion from the other electrons present in this level.)

E.A.(high) Al Si P S Cl

(added electron goes to the bigger n = 3 i.e. occupies larger region of space and the electron-electron repulsion is much less.)

(c) Noble gases have large positive electron gain enthalpies because the electron has to enter the next higher principal quantum level leading to a very unstable electronic configuration.

e.g.H in kJ / mol of groups 16

th _{and 17}th O S Se Te Po –141 –200 –195 –190 –174 F Cl Br I At –328 –349 –325 –295 –270 Electonegativity.

A qualitative measure of the ability of an atom in a chemical compound to attract shared electrons to itself iscalled electronegativity not a measurable quantity. A number of numerical scales of electronegativity have been proposed.

Differences in E.A. & E.N.

(a) E.A. is defined in isolated gaseous state while E.N. is defined in bonded state

(b) E.A. is a absobule term, it has proper units while E.N. is a relative concept has no units Different Scales of Measurement of Electonegativity.

(a) Pauling’s Scale

 = X_A– X_B = O.208 E.AB EAAEBB

E_A-B = Bond enthalpy/ Bond energy of A – B bond.

E_{A - A} = Bond energy of A – A bond

E_B

–B = Bond energy of B – B bond

All are in kcal / mol)

 = XA– XB = O.1017 E.AB EAAEBB

(b) Mulliken’s scale E.N. = 2 . A . E . E .  

Valid only if E.A. = + ve values are in ev / atom

 pauling mullian N . E N . E 2.8

(c) Allred–Rochow’s Electronegativity X_AR = 2

A r Zeff 359 . O + 0.744 r_A Covalent radii in Aº

Factors affecting Electonegativity.

(a) Size Size _  EN _

(b) Zeff Zeff _  EN _ _  EN_

(c) Charge on cationic species

A3+ _{> A}2+ _{> A}+1 _{(for the same element)}

Greater the charge on cation, greater E.N value (d) Charge on anionic species

A3– _{< A}2– _{< A}–

(e) State of hybridization

Greater the %s character  greater the attraction on the shared pair  Greater will be E.N

sp > sp2 _{> sp}3

Periodicity

L – R  Zeff _ E.N _

T – P  Size_ E.N _

Irregularity

Noble gases have very low E. N. values

---BE SLOW AND EXPLAIN PROPERLY

Applications

(a) Prediction of nature of bond

(a) X_A = X_B  pure covalent bond

(b) X_A X_B  partly ionic + covalent

Acc. to pauling, if

X_A– XB = 1.7 50% ionic

 < 1.7 more covalent less ionic  > 1.7 more ionic less covalent

Henny smith formula % ionic character = 16 + 3.5 2

= | XA– XB|

Bond to be 50% ionic  = 2.1

Q. Calculate % ionic character of bond formed between the most electropositive element Cs (x = .7) and most electronegative element F (x = 4.0). According to Hanny Smith formula % ionic character

= 16(3.3) + 3.5 (3.3)2

= 52.8 + 38 .115 = 91.915 %

(b) To decide nature of oxides

Types of Oxides :

(a) Acidic oxides

(a) Solution in water will be acidic in nature (b) will react with base but not with an acid Ex. SO_{2  }H_2O _H 2SO3 Sulphurous acid SO_{3  }H2_O H 2SO4 sulphsic acid CO_{2  }H_2O _H 2CO3 Carbonic acid

generally non-metallic oxider are acidic oxides

(b) Basic oxides

(a) Solution in water will be basic in nature (b) will react with an acid but not with a base Ex. Na₂O _{ }H2_O 2NaOH

CaO _{ }H_2O _Ca(OH)

2

Generally, metallic oxides are basic oxides

(c) Amphoteric oxides

Can react with an acid as well as with a base.

Ex. Generally, metalloids or elements close to metalloids can form ampholeric oxides. BeO

Be O + 2HCl  BeCl₂ + H₂O Be O + 2NaOHNaAlO₂

Sodiuni meta aluminate ZnO

ZnO + HCl ZnCl₂

ZnO + NaOH Na2ZnO2

sodium zincate 2 SnO  ( stannous oxide) SnO + HCl SuCl2

SnO + NaOH  Na2SnO2

SnO₂ ( stannic oxide)

SnO₂ HCl  SnCl4

(d) Neutral oxides

will donot react acids or bases CO, N₂O, NO

(e) Amphiprotic oxide

which can accept and release H+ _ions

Ex.  H2O

Periodicity in nature of oxides

(a) L  R, metallic character  and non-metallic character . So, basic character of oxides  and acidic

character _ basic strongly O Na₂ basic MgO AmphAl2O3 acidic weakly SiO₂ acidic O P2 5 acidic strongly SO₃ acidic strongly O Cl_{2 7}

(b) T  B, metallic character  so, basic character of oxide will 

LiOH LI₂O

NaOH Na₂O

K O H K₂O

RbOH Rb₂O

CsOH Cs₂O

Basic character increases down the group

(c) If the same element is forming oxides in diff OXn _{stats, then greater the Oxd}n _{no greater will nature}

2 basic more MnO < ₂3 ₃ basic O Mn  < 4₂ neutral MnO < 6 3 acidic MnO  < 7 7 2 acidic more O Mn

oxides are anlydrides of oxyacid or hydroxides

4 3 2SO H     H2O 4 2 SO  ; 6 4 2SO H     H2O +5SO3 H₂CO_{3 }H2_O_ CO 2 ; 5 4 3PO H    +5P2O5 (P4O10) H₃PO₃ P₂O₃ (P₄O₆) ; HNO₃  N2O5 NHO_{2 }N₂O₃ ; HClO_{4 }Cl₂O₇ HClO_{3 }Cl₂O₅ ; HClO_{2 } Cl₂O₃

HClO _ Cl₂O ; NaOH _ Na₂O ; Ca(OH)_{2 } CaO

Mixed anhydrides. 3 2 HNO  + 5 2 ClO     H2O 4 4 2O N  (NO₂) ; 4 2 ClO     H2O 5 3 HClO  + 3 2 HClO 

For oxy acids

(a) On moving L  R non metallic character  acidic strength 

H₃BO₃ < H₂CO₃ < HNO₃ H₂SiO₃ < H₃PO₄ < H₂SO₄ < HClO₄ (b) On moving T B, non – metallic character _ acidic strength _

HNO₃ > H₃PO₄ > H₃AsO₄ HClO₄ > HBrO₄ > HO₄

(c) If same element is forming oxycids in different oxidation states then, Greater the oxidation no. of the element greater will be the acidic strength

HNO₃ > HNO₂ > H₂N₂O₂ HClO₄ > HClO₃ > HClO₂ > HClO

Periodicity of Valence or Oxidation States

The valence of representative elements is usually (though not necessarily) equal to the number of electrons in the outermost orbitals and / or equal to eight minus the number of outermost electrons. Nowadays the term oxidation state is frequently used for valence.

Consider the two oxygen containing compounds : OF₂ and Na₂O. The order of electronegativity of the three elements involved in these compounds is F > O > Na. Each of the atoms of fluorine, with outer electronic

configuration 2s

2 molecule. Being highest electronegative

element, fluorine is given oxidation state –1. Since there are two fluorine atoms in this molecule, oxygen with

outer electronic configuration 2s2_2p4 _{shares two electrons with fluorine atoms and thereby exhibits oxidation}

state +2. In Na₂O, oxygen being more electronegative accepts two electrons, one from each of the two sodium atoms and thus, shows oxidation state –2. On the other hand sodium with electronic configuration

3s1 _{loses one electron to oxygen and is given oxidation state +1. Thus, the oxidation state of an element in}

a particular compound can be defined as the charge acquired by its atom on the basis of electronegative consideration from other atoms in the molecule.

There are many elements which exhibit variable valence. This is particularly characteristic of transition elements and actinoids.

Periodic Trends and Chemical Reactivity

As the periodicity is related to electronic configuration, all chemical and physical properties are a manifesta-tion of the electronic configuramanifesta-tion of elements.

The atomic and ionic radii, as we know, generally decrease in a period from left to right. As a consequence, the ionization enthalpies generally increase (with some exceptions) and electron gain enthal-pies become more negative across a period. In other words, the ionization enthalpy of the extreme left element in a period is the least and the electron gain enthalpy of the element on the extreme reight is the highest negative (note : noble gases having completely filled shells have rather positive electron gain en-thalpy values). This results in high chemical reactivity at the two extremes and the lowest in the centre. Thus, the maximum chemical reactivity at the extreme left (among alkali metals) is exhibited by the loss of an electron leading to the formation of cation and at the extreme right (among halogens) shown by the gain of an electron forming an anion. This property can be related with the reducing and oxidizing behaviour of the elements which you will learn later. However, here it can be directly related to the metallic and non-metallic character of elements. Thus, the metallic character of an element, which is highest at the extremely left decreases and the non-metallic character increases while moving from left to right across the period. The chemical reactivity of an element can be best shown by its reactions with oxygen and halogens. Here, we shall consider the reaction of the elements with oxygen only. Elements on two extremes of a period easily combine with oxygen to form oxides. The normal oxide formed by the element on extreme left is the most basic (e.g. Na₂O), whereas that formed by the element on extreme right is the acidic (e.g. Cl₂O₇). Oxides of elements in the centre are amphoteric (e.g. Al₂O₃, As₂O₃) or neutral (e.g., CO, NO, N₂O). Amphoteric oxides behave as acidic with bases and as basic with acids, whereas neutral oxides have no acidic or basic properties.

Among transition metals (3d-series), the change in atomic radii is much smaller as compared to those of representative elements across the period. The change in atomic radii is still smaller among inner-transition metals (4ƒ series). The ionization enthalpies are intermediate between those of s- and p-blocks.

As a consequence, they are less electropositive than group 1 and 2 metals.

In a group, the increase in atomic and ionic radii with increase in atomic number generaly results in a gradual decrease in ionization enthalpies and a regular decrease (with exception in some third period elements) in electron gain enthalpies in the case of main group elements.

Thus the metallic character increases down the group and non-metallic character decreases. This trend can be related with their reducing and oxidizing property. In the case of transition elements, however, a reverse trend is observed. This can be explained in terms of atomic size and ionization enthalpy.

nonmetalic chara cter metalic charac ter E le c tr o n e g a ti v it y A to m ic R a d iu s Io n iz a ti o n E n th a lp y E le c tr o n G a in E n th a lp y Ionization Enthalpy Atomic Radius Electronegativity Electron Gain Enthalpy