Chapter 1- “Periodic Properties and Variations of Properties – Physical & Chemical”
MENDELEEV’S
PERIODIC LAW: “The physical and chemical properties of elements are periodic functions
of their atomic weights.”
MODERN PERIODIC
LAW: “The
physical and chemical properties of elements are periodic functions of their
atomic numbers.”
Cause
of Periodicity: Similar electronic configuration or the same number of
valence electrons.
LONG FORM OF PERIODIC TABLE
(i)
There
are 18 VERTICAL COLUMNS (Groups) in
the periodic table.
Ø All
the elements in the same group should have similar electronic configurations.
Ø Elements in GROUP 1 and GROUP 2
placed to the extreme left of the periodic table are METALS.
Ø Elements in GROUP 13 TO 17 placed to the right side of the periodic table are NON-METALS.
Ø Elements in GROUP 18 are INERT GAS or
NOBLE GAS elements.
Ø Elements in GROUP 3 TO GROUP 12 are transition elements which generally exhibit
intermediary properties.
(ii)
There
are Seven Horizontal Rows (Periods)
in the periodic table.
Ø Each period represents a Different Principal Energy Level (shell). In
each element, valence electrons are placed in the principal energy level
(valence orbit).
PERIODIC TRENDS IN PROPERTIES
1.
Atomic Radius:
Ø
Size
INCREASES as you go down a GROUP as new shells are added.
Ø Size DECREASES as you go across a PERIOD as
the nuclear charge increases.
Isoelectronic Series: The size of the atom decreases as
the number of protons increases.
Ø Size of a cation is less than the parent
Atom.
Ø Size of the anion is more than the parent Atom.
2.
Metallic Character:
Ø Metallic character DECREASES while
non-metallic character INCREASES in a PERIOD.
Ø Metallic character INCREASES while
non-metallic character DECREASES in a GROUP.
3.
Ionization Energy Or Ionization Enthalpy Or
Ionization Potential
“Ionization
potential is
defined as the minimum amount of energy that is required to remove the most a loosely bound electron from an isolated gaseous atom in its ground state so as
to convert it into a gaseous cation.”
Ø The ionization energy INCREASES IN A PERIOD
as the atomic numbers increase.
Ø The ionization energy DECREASES in a GROUP
as the atomic size increases.
Order of successive Ionization Energy (IE): IE1 <IE2 <IE3
Factors influencing
ionization energy
(i) Atomic number
(ii) Atomic size
(iii) Electronic
configuration
4.
Electron Gain Enthalpy Or Electron
Affinity:
“Electron affinity an element may be defined as
the energy released when a neutral isolated gaseous atom accepts an extra
electron to form a gaseous anion, i.e. a negative ion.”
Ø The electron gain enthalpy INCREASES on
moving left to right in a period.
Ø
The
electron gain enthalpy DECREASES on moving down the group.
5.
Electronegativity(EN):
Ø INCREASES across the PERIOD from left to
right.
Ø DECREASES down the GROUP from top to bottom
GROUP 1: THE ALKALI METALS
·
Atomic Radii: Increases down the group from top to bottom.
·
Ionic Radii: Increases on moving down the group from top to bottom.
·
Soft metals: Softness
increases with increase in atomic number.
·
Melting and Boiling Points: Low melting and
boiling points due to weak
cohesive forces.
·
Flame Coloration: Reason of flame coloration is low
ionization energy.
Element
|
Colour
|
Lithium (Li)
|
Crimson Red
|
Sodium (Na)
|
Golden Yellow
|
Potassium (K)
|
Pale Violet
|
Cesium (Cs)
|
Blue
|
·
The
alkali metals are highly reactive elements: It increases from Li to Cs.
·
Hydrides. Alkali metals react with dry hydrogen to form ionic
metallic hydrides.
LiH + H2O → LiOH + H2
NaH + H2O → NaOH + H2
These
hydrides are strong reducing agents and
their reducing nature increases down the
group.
·
Oxides. Alkali metals react with air and form BASIC OXIDES.
Basic Nature of Alkali Metal
Hydroxides: Strong bases and highly soluble in water and stable to heat. Exception:
Lithium hydroxide.
·
Reaction with water: Alkali metals react with water and the other
compounds containing acidic hydrogen atoms such as hydrogen halides (HX)
liberate hydrogen gas.
2Na + 2H2O → 2NaOH + H2
2Na + 2HX → 2NaX + H2 (where X = halogen)
·
Reaction with Halogens. Alkali metals react with halogens to form
metal halides, which are ionic crystal solids having general formula M+X-.
·
Reaction with Non-Metals. Alkali metals, on heating, react with
non-metals like sulfur and phosphorus to form sulfides and phosphides
respectively.
2M + S → M2S
3M + P → M3P
GROUP 17 ELEMENTS – Halogen Family
·
Colour: F2 (Pale yellow), Cl2 (greenish-yellow), Br2 (dark red), I2 (Violet)
·
Nature of Bonds: With metals - ionic
halides. With non-metals - covalent compounds.
·
Ionisation Energy: Very high values of IE1.
These IE1 values decrease in moving down the group.
·
Electronegativity: Most electronegative
elements within their respective periods. Electronegativity decreases in moving down the group as
given below:
F(4.0); Cl(3.0); Br(2.9); I(2.7)
·
Electron Affinity: Highest electron affinities within their periods.
Chlorine shows an exceptionally higher value than fluorine. The order of EA1 is Cl > F >
Br > I
·
Hydrides:
(i) All halogens form hydrides
with formula HX (HF, HCl, HBr, HI).
(ii) Acidic nature of these
hydrides is HI > HBr > HCl > HF
·
Oxidizing power: Strong oxidizing agents
and the oxidizing power decrease down
the group from fluorine to Iodine.
COMPARISON BETWEEN ALKALI METAL AND
HALOGEN
S.No
|
Property
|
Alkali
metal
|
Halogen
|
1
|
Name of element
|
Lithium(Li),
Sodium(Na)
Potassium(K),
Rubidium(Rb)
Cesium(Cs),
Francium(Fr)
|
Fluorine(F),
Chlorine(Cl)
Bromine(Br),
Iodine(I)
Astatine(At)
|
2
|
Physical state
|
Solid
|
F,Cl–gas; Br-liquid;
I-solid
|
3
|
Atomic number
|
Li – 3; Na – 11; K
– 19; Rb -37; Cs – 55; Fr – 87
|
F- 9; Cl- 17; Br-
35; I -53;
At -85
|
4
|
Electronic configuration
|
Li – 2,1
Na – 2,8,1
K – 2,8,8,1
Rb – 2,8,18,8,1
|
F- 2,7
Cl- 2,8,7
Br- 2,8,18,7
I -2,8,18,18,7
|
5
|
Valence electron
|
1
|
7
|
6
|
Ion formation
|
Forms
cation (+ve ion)
|
Forms anion (-ve ion)
|
7
|
Electrovalence
|
+1
|
-1
|
8
|
Type of compounds
|
Electrovalent
|
Electrovalent and covalent
|
9
|
Atomic size
|
Highest in the period
Increases from Li
to Fr
|
Lowest in the period
Increases from F to
At
|
10
|
Thermal conductivity
|
Good conductor
|
Bad conductor
|
11
|
Electrical
conductivity
|
Good conductor
|
Bad conductor
|
12
|
Ionization energy
|
Lowest in the period
|
Highest in the period
|
13
|
Electron affinity
|
Lowest in the period
|
Highest in the period
|
14
|
Electronegativity
|
Lowest in the period
|
Highest in the period
|
15
|
Metallic/non-metallic character
|
Metallic nature
|
Non – metallic
nature
|
16
|
Nature of oxide
|
Basic
|
Acidic
|
17
|
Reducing agent
|
Strong reducing
agent
|
Weak reducing agent
|
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