Download ch 3 classification of elements and periodic properties

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CHAPTER 3
PERIODIC CLASSIFICATION
EARLIER ATTEMPTS OF CLASSIFICATION OF ELEMENTS
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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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Doberenier’s Triads
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


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
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

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
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

“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
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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 table
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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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About Periodic table
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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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s-block elements
1-2
(ns )
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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:
1.
Low IE, High e-+ve character.
2.
Very reactive & hence do not occur in native state.
3.
Good reducing agents
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4.
Compounds
are predominantly
ionic

p-Block elements
2
1-6
(ns np )
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Also known as representative elements or main group
elements
“The elements in which the last electron enters the psubshell of their outermost energy level are called pblock elements”
1.
Exhibit variable oxidation states
2.
They form ionic as well as covalent compounds
3.
They have relatively high values of IE
4.
Most of them are non-metals, highly electronegative,
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acidic oxides.
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d-Block elements
1-10
2
[(n-1) ns ]
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“The elements in which the last electron enters the d-subshell of
their outermost energy level are called d-block elements”
They
1.
Are Hard, high Melting metals,
2.
Have Variable oxidation states
3.
Form coloured complexes
4.
Form ionic as well as covalent compounds
5.
Most of them exhibit paramagnetism, possess catalytic
properties
6.
Form alloys,
7.
Are good conductors of heat & electricity
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f-Block elements [(n-2)f1-14(n-1)d0-10ns2]
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Also known as inner transition elements or f-transition
elements (Lanthanoids & Actinoids)
“The elements in which the last electron enters the fsubshell of their outermost energy level are called fblock elements”
They
1.
Show variable oxidation states
2.
Have high MP, high densities
3.
Form complexes, most of which are coloured
4.
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of the actinoid series
are
radioactive.
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Atomic Radius
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“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
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

Increases down the group
Decreases across the period.
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Ionic radius
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“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
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“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…
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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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Variation of IE across the period
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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:
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Fully-filled energy
level of Ne.
Variation of IE down the group
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IE decreases in general down the group due to..
1. addition of new energy levels
2. increase in screening effect.
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Electrongain enthalpies
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“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 EGE
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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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Variation of EGE across the period
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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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Variation of EGE down the group
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

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
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




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
-(g) + e- → O-2(g)
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egH2 =+780kJ
Electronegativity
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“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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Variation of EN
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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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electronegativity
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Electropositivity or metallic character
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“ 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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Valence
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“ 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 oxides
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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
Acid9824662116
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Mansuri
Contd…
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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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Anomalous properties of second period
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“ 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 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.
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Other members
do no form pπ-pπ bonds due toMansuri
their larger
sizes.
Diagonal Relationship
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“ An element of the 2nd period exhibits certain similarities
with the 2nd element of the following group ”
For example:
1
2
13
14
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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End of chapter
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