1. CHAPTER 3 PERIODIC CLASSIFICATION
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CHAPTER 3 PERIODIC CLASSIFICATION

2. EARLIER ATTEMPTS OF CLASSIFICATION OF ELEMENTS
1. Doberenier’s Triads 2. Newland’s Law of Octaves 3. Mendeleev’s Periodic Table 4. Long form of Periodic Table
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3. 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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4. 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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5. 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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7. 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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8. Long form of Periodic Table
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12. About Periodic table
No. of periods: 7 No. of Groups: 18 No. 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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13. About Periodic table
Period no. 2 & 3 are called ………..Short Periods Period no. 4 & 5 are called ………..Long Periods Period no. 6 is called ………………Longest Period
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14. 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 metals
Properties of s-block elements:
Low IE, High e-+ve character.
Very reactive & hence do not occur in native state.
Good reducing agents
Compounds are predominantly ionic
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15. 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”
Exhibit variable oxidation states
They form ionic as well as covalent compounds
They have relatively high values of IE
Most of them are non-metals, highly electronegative, form acidic oxides.
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16. 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”
They
Are Hard, high Melting metals,
Have Variable oxidation states
Form coloured complexes
Form ionic as well as covalent compounds
Most of them exhibit paramagnetism, possess catalytic properties
Form alloys,
Are good conductors of heat & electricity
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17. 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”
They
Show variable oxidation states
Have high MP, high densities
Form complexes, most of which are coloured
Most of the actinoid series are radioactive.
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18. 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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19. Variation of atomic radius
Increases down the group
Decreases across the period.
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21. 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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22. 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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23. Ionization Enthalpy depends on…
1. Size of atom
Decreases with the increasing size of atom as the electrons are less tightly held with increasing distance
2. 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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24. Variation of IE across the period
IE across the period increases due to.. 1. Increase in Nuclear charge 2. Addition of e-s in the same energy level 3. 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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25. Variation of IE down the group
IE decreases in general down the group due to..
1. addition of new energy levels
2. increase in screening effect.
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26. 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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27. Factors affecting EGE
1. 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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28. 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 subshells
2. in group 15 → half-filled np subshells
3. 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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29. 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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30. 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 =-141kJ
O-(g) + e- → O-2(g) ∆egH2 =+780kJ
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31. 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 radius smaller 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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32. Variation of EN
Increases 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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33. electronegativity
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34. 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 IE
Increases down the group
due to decrease in IE
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36. 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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38. Nature of oxides
1. 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 Base 4. An Acidic oxide is one which when dissolved in water gives an acid Cl2O7 + H2O → 2HClO4. Acidic oxide Acid
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39. Contd…
5. An Amphoteric oxide exhibits acidic behaviour in presence of base and basic behaviour in presence of acid. Al2O3 + 6HCl → 2AlCl3 + 3H2O Al2O3 + 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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40. 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 element 2. Large charge/radius ratio 3. High EN 4. 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 =4 On the other hand, n=3 onwards d-orbitals are available. ∴ covalency can be expanded beyond 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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41. Diagonal Relationship
“ An element of the 2nd period exhibits certain similarities with the 2nd element of the following group ” For example: Li Be B C Na 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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42. End of chapter
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