basic description of Alkali metals

1. S Block Elements
Alkali metals
- Mohamed Minas
Ist M.Sc Chemistry 1CHD6108

2. TABLE OF
CONTENTS
01
Introductiom
02
Characteristics properties of Alkali metals
03
applications
04
References
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3. The elements in the long form of the
periodic table has been divided into
four blocks, namely s, p, d & f blocks.
The elements of group I & II receive
their last electron in s-orbital. So they
are called as s – block elements
The general electronic configuration of
s-block elements is [noble gas]ns1 for
alkali metals and [noble gas}ns2 for
alkaline earth metals..
INTRODUCTION
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4. S block
elements
Na
K
Rb
Cs
Fr
Be
Mg
Ca
Sr
Ba
Ra
IA IIA
IA - Alkali metals
IIA - Alkaline earth metals
Li
3 4
11 12
19 20
37 38
55 56
87 88
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5. Characteristics properties of Alkali Metals
Electronic Configuration: All the alkali metals have one valence
electron, ns1 outside the noble gas core. The loosely held s-electron in the
outermost valence shell of these elements makes them the most
electropositive metals. They readily lose electron to give monovalent M+
ions. Hence they are never found in free state in nature.
Atomic and Ionic Radii: The alkali metal atoms have the largest sizes in
a particular period of the periodic table. With increase in atomic number,
the atom becomes larger. The monovalent ions (M+ ) are smaller than the
parent atom. They increase on moving down the group i.e., they increase
in size while going from Li to Cs.
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6. Characteristics properties of Alkali Metals
Ionization Enthalpy: As we go down the group the size of the atoms increases
due to which the attraction between the nucleus and the electrons in the outermost
shell decreases. As a result, the ionization enthalpy decreases. The ionization
enthalpy of the alkali metals is comparatively lesser than other elements.
Flame Colouration: The alkali metals and their salts impart a characteristic
colour to flame On heating an alkali metal or its salt (especially chlorides due to its
more volatile nature in a flame), the electrons are excited easily to higher energy
levels because of absorption of energy. When these electrons return to their ground
states, they emit extra energy in form of radiations which fall in the visible region
thereby imparting a characteristic colour to the flame.
Li Na K
Colour Crimson Red Golden yellow Pale violet
nm 670.8 589.2 766.5
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7. Atomic and Physical Properties of the Alkali Metals
7

8. Chemical Properties
Reactivity towards air:: The alkali metals tarnish in air due to the formation of an oxide or
hydroxide on the surface. Alkali metals when burnt in air form different kinds of oxides . The alkali
metals on reaction with limited quantity of oxygen form normal oxides of formula, M2O
4M + O2 2M2O (Where M = Li, Na, K, Rb, Cs)
When heated with excess of air, lithium forms normal oxide,Li2O ; sodium forms peroxide, Na2O2,
whereas potassium rubidium and caesium form superoxides having general formula MO2
4LiO2 2Li2O ( Lithium oxide)
2Na + O2 Na2O2 ( Sodium peroxide)
K + O2 KO2 ( Potassium Superoxide)
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10. Chemical Properties
Reactivity towards Water: All alkali metals react violently with water with the formation of
the metal hydroxide and hydrogen. lithium is the least reactive alkali metal and reacts ‘only’
quickly with water, whereas potassium, rubidium and caesium are more reactive and react
violently with water
2Li(s) + 2H2O 2LiOH + H2O(g)
Reactivity towards Hydrogen: Alkali metals react with dry hydrogen at about 673K to form
colourless crystalline hydrides. All the alkali metal hydrides are ionic solids with high melting
points.
2M + H2 2M+ H-
These hydrides react with water to form corresponding hydroxides and hydrogen gas.
LiH + H2O LiOH + H2
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11. Chemical Properties
Reactivity towards Halogen: Alkali metals react vigorously with halogens to form metal
halides of general formula MX, which are ionic crystalline solids.
2M + X2 2MX M = Li, Na, K, Rb or Cs and X = F, Cl, Br or I
Solubility in liquid Ammonia: The alkali metals dissolve in liquid ammonia giving deep blue
solutions which are conducting in nature
M (x+y)NH3 [M(NH3)x]+ + [e(NH3)y]-
The blue colour of the solution is due to the ammoniated electron which absorbs energy in the
visible region of light and thus imparts blue colour to the solution.The solutions are
paramagnetic and on standing slowly liberate hydrogen resulting in the formation of amide.
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Applications
Sodium is by far the most important alkali metal in terms of industrial use. Hundreds of thousands
of tons of commercial compounds that contain sodium are used annually, including common salt
(NaCl), baking soda (NaHCO3), sodium carbonate (Na2CO3), and caustic soda (NaOH).
Potassium has considerably less use than sodium as a free metal. Potassium salts, however, are
consumed in considerable tonnages in the manufacture of fertilizers.
Lithium metal is used in certain light-metal alloys and as a reactant in organic syntheses. An
important use of lithium is in the construction of lightweight batteries.
Rubidium and cesium and their compounds have limited use, but cesium metal vapour is used in
atomic clocks, which are so accurate that they are used as time standards.
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13. References
13
Atkins P.W., Overton T., Rourke, J., Weller, M. and Armstrong, F. Shriver and Atkins
inorganic chemistry, 4th edition, Oxford University Press, 2006.
F. Albert Cotton, Geoffrey Wilkinson, Paul L. Gaus basic inorganic chemistry, 3rd=
edition
Catherine E. Housecroft and Alan G. Sharpe inorganic chemistry, 2nd edition, Pearson
Education Limited.
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