CHAPTER 3 PERIODIC CLASSIFICATION

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CHAPTER 3PERIODIC CLASSIFICATION

www.chemzblog.wordpress.com.....Zaid Mansuri 9824662116

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EARLIER ATTEMPTS OF CLASSIFICATION OF ELEMENTS

1. Doberenier’s Triads2. Newland’s Law of Octaves3. Mendeleev’s Periodic Table4. Long form of Periodic Table


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Doberenier’s Triads Certain similar elements exist in group of

three elements which he named as triads. The At. Wt. of middle member was the

arithmetic mean of the other two members of the triad.

Properties of the middle element was intermediate of the other two.


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Newland’s Law of Octaves Elements were arranged in increasing order

of atomic weights. Eight element, starting from a given one is

a kind of repetition of the first. – like musical notes


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Mendeleev’s Periodic Table “The properties of elements are a periodic

function of their atomic weights”. Main criterion of the judgment of similarities

in the properties was valency of the elements.


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Modern Periodic Law The physical & chemical properties of the

elements are the periodic function of their atomic numbers.

Cause of Periodicity: The periodic repetition is due to the

recurrence of similar valence shell configurations after certain regular intervals.


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Long form of Periodic Table


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About Periodic tableNo. of periods: 7 No. of Groups: 18No. of periods represents the highest principal quantum number (n) of the

elements present in it.First Period: n=1…..(1s) two elements (h & He)Second Period n=2….(2s & 2p) eight elements (Li to Ne)Third Period n=3 …(3s & 3p) eight elements (Na to Ar)Fourth Period n=4 …(4s, 3d & 4p) eighteen elements (K to Kr)Fifth Period n=5…(5s, 4d & 5p) eighteen elements (Rb to Xe)Sixth Period n=6…(6s, 4f, 5d & 6p) thirty two elements (Cs to Rn)(Lanthanoids …Ce to Lu)….14 elements (4f)Seventh Period n=7…(7s, 5f, 6d, 7p) (Actinoids…Th to Lr)


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About Periodic tablePeriod no. 2 & 3 are called ………..Short PeriodsPeriod no. 4 & 5 are called ………..Long PeriodsPeriod no. 6 is called ………………Longest Period


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s-block elements (ns1-2)Also known as representative elements or main

group elements“When last electron enters s-subshell, it is an s-block

element.”

Gr. 1: Alkali metals, Gr. 2: Alkaline earth metalsProperties of s-block elements:1. Low IE, High e-+ve character.2. Very reactive & hence do not occur in native state.3. Good reducing agents4. Compounds are predominantly ionic


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p-Block elements (ns2np1-6)Also known as representative elements or

main group elements“The elements in which the last electron enters

the p-subshell of their outermost energy level are called p-block elements”

1. Exhibit variable oxidation states2. They form ionic as well as covalent

compounds3. They have relatively high values of IE4. Most of them are non-metals, highly

electronegative, form acidic oxides.


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d-Block elements [(n-1)1-10ns2]“The elements in which the last electron enters the d-

subshell of their outermost energy level are called d-block elements”

They1. Are Hard, high Melting metals,2. Have Variable oxidation states3. Form coloured complexes4. Form ionic as well as covalent compounds5. Most of them exhibit paramagnetism, possess

catalytic properties6. Form alloys,7. Are good conductors of heat & electricity


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f-Block elements [(n-2)f1-14(n-1)d0-10ns2]

Also known as inner transition elements or f-transition elements (Lanthanoids & Actinoids)

“The elements in which the last electron enters the f-subshell of their outermost energy level are called f-block elements”

They1. Show variable oxidation states2. Have high MP, high densities3. Form complexes, most of which are coloured4. Most of the actinoid series are radioactive.


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Atomic Radius“The distance from the centre of nucleus of

the atom to the outermost shell of electrons.”

1. Covalent Radius: One-half of the distance between the centres of the nuclei of two similar atoms bonded by a single covalent bond.

2. Metallic Radius: One-half of the internuclear distance between two adjacent atoms in the metallic lattice.

***The metallic radius is always larger than its covalent radius.


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Variation of atomic radius Increases down the group Decreases across the period.


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Ionic radius“Ionic radius is defined as the effective

distance from the nucleus fo the ion to the point up to which it has an influence in the ionic bond”

(a) The size of cation is smaller than parent atom becoz of increase in the effective nuclear charge per electron.

(b) The size of anion is greater than parent atom becoz of decrease in effective nuclear charge per electron.


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Ionization Enthalpies“The amount of energy required to remove

the most loosely bound electron from its isolated gaseous atom in the ground state”


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Ionization Enthalpy depends on…1. Size of atom Decreases with the increasing size of atom as the electrons are less tightly held with

increasing distance2. Magnitude of nuclear charge Higher the nuclear charge, higher the IE.3. Screening effect of the inner electrons (Outermost electrons are shielded or screened by the inner electrons. This is screening

effect.) IE decreases with increase in screening effect. More the no. of inner electrons, greater is the screening and lower is the IE.4. Penetration effect of the electrons: Penetration effect for a given ‘n’ s>p>d>f. Greater the penetration, lower the shielding by other electrons, higher the IE.5. Electronic configuration: Atom having more stable (half filled, fully filled subshells) config. has less tendency to

lose e-, hence higher the IE. noble gases (ns2np6) elements like N: [He] 2s22px12py12pz1 & P: [Ne] 3s23px13py13pz1 have half-filled stable

configuration. elements like Be: 1s22s2, Mg:[Ne]3s2 have electrons paired


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Variation of IE across the periodIE across the period increases due

to..1. Increase in Nuclear charge2. Addition of e-s in the same energy level3. Decrease in atomic size.Exceptions:1. Decrease from Be to B:

Reason: (a) penetration of 2s>2p (b) more shielding faced of 2p by inner electrons (c) more stable config of Be.

2. Decrease from N to O:Relatively stable half-filled configuration of

N: [He] 2s22px12py12pz1.3. Large increase from F to Ne:Fully-filled energy level of Ne.


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Variation of IE down the groupIE decreases in general down the group

due to..1. addition of new energy levels2. increase in screening effect.


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Electrongain enthalpies“The enthalpy change taking place when

an isolated gaseous atom of the element accepts an electron to form a monovalent gaseous anion”

X(g) + e- → X-(g) Larger the negative EGE, greater the

tendency to accept electron.


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Factors affecting EGE1. Nuclear charge:

Greater the NC, more attraction, large –ve is the EGE

2. Atomic size:Smaller the size, higher the attraction, large –ve is the EGE

3. Electron configuration:More stable e- configuration, less tendency to accept the e-, less –ve EGE.For example: Noble gases have high +ve EGE.


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Variation of EGE across the period Across the period, atomic size decreases &

nuclear increases, therefore electron gain enthalpies tend to be more –ve.

Some irregularities…1. in group 2 → filled ns subshells2. in group 15 → half-filled np subshells3. in group 18 → fully filled subshells

These elec config are relatively stable & hence these have +ve or very low –ve EGEs.


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Variation of EGE down the group Down the group, the atomic size & nuclear

charge both increase. But increase in atomic size is more pronounced. Therefore, EGE becomes less –ve down the group.

Some irregularities…1. F(-328) < Cl(-349). Reverse expected. Becoz, when an electron is added to F, it goes to relatively compact n=2 energy level. As a result it experiences significant e-e repulsion. 2. Same is the case with O(-141) < S(-200).


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Successive EGE X(g) + e- → X-(g) ∆egH1 X-(g) + e- → X-2(g) ∆egH2 Always, ∆egH2 is +ve . This is becoz, when

e- is added to uninegative ion, it experiences significant repulsion. Hence energy has to be supplied to overcome the repulsive force to add electron.

Hence values of successive EGEs are positive.

For example:O(g) + e- → O-(g) ∆egH1 =-141kJO-(g) + e- → O-2(g) ∆egH2 =+780kJ


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Electronegativity“The tendency of an atom in a molecule

to attract the shared pair of electron towards itself”

Factors affecting electronegativity are…1. Effective nuclear charge

Greater the Nuclear charge, greater is the EN

2. Atomic radiussmaller the Atomic radius, greater the EN

***EN for any given element is not constant but varies depending on the element to which it is bound.


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Variation of ENIncreases across the period… becoz of increasing nuclear charge and

decreasing atomic size.Decreases down the group…

becoz of increasing atomic size.-------------------------------------------------*** non-metallic character is directly related

to EN.


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electronegativity


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Electropositivity or metallic character“ Tendency of atoms of an element to

lose electrons and form +ve ions is known as Electropositivity”

A more electro+ve element has more metallic character.

Variation of Electropositive character…Decreases across the period

due to increase in IEIncreases down the group

due to decrease in IE


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Valence“ The valence of an element may be

defined as the combining capacity of element ”

Valence= no. of H or Cl or double the no. of O atoms that combine with an atom of an element.

Electrons present in the outermost shell are called valence e-s.

Down the group, the valency remains the same.

Across the period, increases from 1to 4 & then decreases from 4 to 0.


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Nature of oxides1. Oxides of elements at the extreme left of

periodic table are BASIC in nature. (metallic oxides)2. Extreme right, ACIDIC. (non-metallic oxides)3. A Basic oxide is one which when dissolves in

water gives a base. Na2O + H2O → 2NaOH

Basic oxide Base4. An Acidic oxide is one which when dissolved in

water gives an acidCl2O7 + H2O → 2HClO4.Acidic oxide Acid


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Contd…5. An Amphoteric oxide exhibits acidic

behaviour in presence of base and basic behaviour in presence of acid.Al2O3 + 6HCl → 2AlCl3 + 3H2OAl2O3 + 2NaOH → 2Na[Al(OH)4]

6. A Neutral oxide exhibits neither acidic nor basic properties.

* Since metallic character increases down the group, the basic character of oxides also increases.


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Anomalous properties of second period“ First member of each group (from Li to F) is different from rest members of the

same group”For example: Li → covalent compounds while other members of group 1 → ionic compounds Reason: 1. Small size of the first element2. Large charge/radius ratio3. High EN4. Absence of d-orbitals in the valence shell of the first element:

1st element → n=2, no d-orbitals in this energy level∴ no d- orbitals available

∴ maximum covalency =4On the other hand, n=3 onwards d-orbitals are available. ∴ covalency can be expanded beyond 4.

5. Ability to form pπ-pπ multiple bonds:1st member due to its small size, forms pπ-pπ multiple bonds with itself & other members of 2nd period.eg: C=C, C≡C, C=O, C ≡N.Other members do no form pπ-pπ bonds due to their larger sizes.


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Diagonal Relationship“ An element of the 2nd period exhibits

certain similarities with the 2nd element of the following group ”

For example: 1 2 13 14Li Be B CNa Mg Al Si

“ Diagonal relationship is the similarity between a pair of elements in different groups and different periods and located diagonally in the periodic table”


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End of chapter

CHAPTER 3 PERIODIC CLASSIFICATION

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