Transcript File - cpprashanths Chemistry

S-Block Elements
ALKALI METALS
Chapter summary
• Introduction
• Characteristic properties of the s-block
elements .
• Variation in properties of the s-block
elementsof the First Group(Alkali Metals).
• Physical Properties
Members of the s-Block Elements
IA IIA
Li
Be
Na
Mg
K
Ca
Rb
Sr
IA Alkali metals
Cs
Ba
Fr
Ra
IIA Alkaline Earth
metals
Characteristic properties of s-block
elements
•
•
•
•
•
•
Metallic character
Low electronegativity
Basic oxides, hydroxides
Ionic bond with fixed oxidation states
Characteristic flame colours
Weak tendency to from complex
Metallic character
• High tendency to
lose e- to form
positive ions
• Metallic character
increases down both
groups
Electronegativity
• Low nuclear
attraction for
outer electrons
• Highly
electropositive
• Small
electronegativity
Group I
Group II
Li
1.0
Be
Na
0.9
Mg 1.2
K
0.8
Ca 1.0
Rb
0.8
Sr
Cs
0.7
Ba 0.9
Fr
0.7
Ra 0.9
1.5
1.0
Characteristic flame colours
Na+ Cl- (g)  Na (g) + Cl (g)
Na(g)  Na* (g)
[Ne]3s1 [Ne]3p1
Na*(g)  Na(g) + h (589nm, yellow)
Flame test
Li
Na
K
Rb
Cs
HCl(aq)
deep red Ca brick red
yellow
Sr blood red
lilac
Ba apple green
bluish red
blue
sample
Variation in properties of elements
•
•
•
•
•
Atomic radii
Ionization enthalpies
Hydration enthalpies
Melting points
Reactions with oxygen, water, hydrogen and
chlorine
Atomic radii (nm)
Li
0.152 Be
0.112
Na
0.186 Mg
0.160
K
0.231 Ca
0.197
Rb
0.244 Sr
0.215
Cs
0.262 Ba
0.217
Fr
0.270 Ra
0.220
Fr
Li
Be
Ra
Ionization Enthapy
Group I 1st I.E.
Li
Na
K
Rb
Cs
519
494
418
402
376
2nd I.E.
Group I 1st I.E.
7300
Be
900
4560
Mg
736
3070
Ca
590
2370
Sr
548
2420
Ba
502
2nd I.E. 3rd I.E.
1760
14800
1450
7740
1150
4940
1060
4120
966
3390
Ionization Enthalpy
1st I.E.
2000
600
Li
500
400
300
Be+
Na
K
1500
Rb
Cs
2nd IE
Ca+
1000
Ba+
Be
500
Ca
1st IE
Ba
Ionization Enthalpy
Group I
1. Have generally low 1st I.E. as it is well shielded
from the nucleus by inner shells.
2. Removal of a 2nd electron is much more difficult
because it involves the removal of inner shell
electron.
3. I.E. decreases as the group is descended.
As atomic radius increases, the outer e is further
away from the well-shielded nucleus.
Hydration Enthalpy
M+(g) + aqueous  M+(aq) + heat
-600
M+
-300
Li+ Na+ K+
Rb+ Cs+
Hydration Enthalpy
-2250
-600
-2000
-1750
-300
-1500
Li+ Na+ K+
Rb+ Cs+
Be2+ Mg2+ Ca2+ Sr2+ Ba2+
Hydration Enthalpy
General trends:
1. On going down both groups, hydration enthalpy
decreases.
(As the ions get larger, the charge density of the
ions decreases, the electrostatic attraction between
ions and water molecules gets smaller.)
2. Group 2 ions have hydration enthalpies higher
than group 1.
( Group 2 cations are doubly charged and have
smaller sizes)
Ionization Energy
• Amount of energy required to remove an
electron from the ground state of a gaseous
atom or ion.
– First ionization energy is that energy required to
remove first electron.
– Second ionization energy is that energy required
to remove second electron, etc.
Trends in First Ionization Energies
• As one goes down a
column, less energy is
required to remove the
first electron.
– For atoms in the same
group, Zeff is essentially
the same, but the
valence electrons are
farther from the nucleus.
Ionization Enthalpy of Alkali Metals
• ionization enthalpy decreases down the group
from Li to cs because as we move down a
group the number of valence electrons goes
increasing separating the electrons away from
the nucleus ,there is an increasing shielding of
the nuclear charge by the inner shell electrons
and thus the removal of electrons requires
less energy as we move down.
Physical Properties
•
•
•
•
Silvery White
Soft & Light Metals
Due to large size , elements have low density
Low Melting & Boling Points indicates weak
bonding due to presence of only 1 valence
electron .
• In order of increasing atomic number the alkali metals are:
• Lithium
Increasing
atomic
number
•Sodium
•Potassium
•Rubidium
•Caesium
•Francium
All alkali (group 1) metals react violently with water,
forming Hydrogen gas and Hydroxides (pH above 7):
Alkali metals are:
•Metals found in group 1 of the periodic table.
•Soft when cut (compared to other metals).
•Metals with low melting points and densities.
•Powerful reducing agent and form univalent
compounds.
•Metals which tarnish in air.
On cumbustion in excess of air,
alkali metals form
E.g.
Li2O
1. Oxides- LiO2
Na2O, Na2O2
2. Peroxides- Li2O2, NaO2
3. Superoxides- K2O2 ,Cs2O2
RbO2
K2O2, KO2
Rb2O2, RbO2
Cs2O2, CsO2
E.g.
The oxides of the alkali metal are easily LiOH
hydrolysed by addition of water.
NaOH
E.g.:- LiO2 + H2O → LiOH
KOH
RbOH
CsOH
The alkali metals combine directly with halogens
under appropriate conditions forming halides of
general formula MX.
These halides can also be prepared by the action
of aqueous halogen acids (HX) on metals oxides,
hydroxides or carbonate.
All these halides are colourless, high melting
crystalline solids having high negative enthalpies
of formation.
M2O + 2HX → 2MX + H2O
MOH + HX → MX + H2O
M2CO3 + 2HX → 2MX + CO2 + H2O
(M = Li, Na, K, Rb or Cs)
(X = F, Cl, Br or I)
Since the alkali metals are highly electropositive,
therefore their hydroxides are very strong bases
and hence they form salts with all oxoacids.
( H2CO3, H3PO4, H2SO4, HNO3, HNO2 etc) .
They are generally soluble in water and stable
towards heat.
The carbonates (M2CO3) of alkali metals are
remarkably stable upto 1273 K, above which they first melt
and then eventually decompose to form oxides.
Li2CO3 , however is considerably less stable and decomposes
readily.
Δ
Li2CO3 → Li2O + CO2
This is presumably due to large size difference between
Li+ and CO2-3 which makes the crystal lattice unstable.
SODIUM CHLORIDE
The most abundant source of sodium chloride
is sea water. Crude sodium chloride, generally
obtained by crystallisation of brine solution,
contains sodium sulphate, calcium sulphate,
calcium chloride and magnesium chloride as
impurities.
Calcium
chloride,CaCl2,
and
magnesium chloride MgCl2
are impurities
because they are deliquescent(absorb moisture
from the atmosphere).
To obtain pure sodium chloride–
Crude salt is dissolved in minimum
amount of water and filtered to remove
insoluble impurities.
The solution is then saturated with
hydrogen chloride gas.
Crystals of pure sodium separate out.
Calcium and magnesium chloride, being
more soluble than sodium chloride, remains
in solution.
USES OF NaCl −
1) It is used as a common salt or
table salt for domestic purpose.
2) It is used for the preparation of
Na2O2 , NaOH and Na2CO3.
SODIUM HYDROXIDE(CAUSTIC
SODA), NaOH
Sodium hydroxide is generally prepared
commercially by the electrolysis of sodium
chloride in Castner −Kellner cell.
A brine soln. is electrolysed using a mercury
cathode and a carbon anode. Sodium metal
discharged at the cathode combines with
mercury to form sodium amalgam. Chlorine gas
is evolved at the anode.
NaCl → Na+ + Cl¯
AT ANODE: Cl ¯ ─ e¯ → Cl
Cl + Cl → Cl2
AT CATHODE: Na+ + e ¯ → Na+
Na + Hg → NaHg
Amalgam
The amalgam is treated with water to give sodium
hydroxide, mercury and hydrogen gas.
2NaHg + 2H2O→ 2NaOH + 2Hg + H2
PROPERTIES1. NaOH is a white, translucent solid and it melts
at 591K.
2. it is readily soluble in water to give alkaline
solution.
3. Crystals of NaOH are deliquescent. It reacts with
the CO2 in the atmosphere to form Na2CO3.
USES The manufacture of soap, paper and
no. of chemicals.
In petroleum refining
In textile industry for mercerising cotton
fabrics
As laboratory reagent
For preparation of pure oils and fats
SODIUM
HYDROGENCARBONATE(BAKING
SODA), NaHCO3
Sodium hydrogencarbonate is known as baking soda
because it decomposes on heating to generate bubbles
of carbon dioxide.
It is made by saturating a solution of sodium
carbonate with carbon dioxide. The white crystalline
powder of sodium hydrogencarbonate , being less
soluble, gets separated out.
Na2CO3 + H2O + CO2 → 2NaHCO3
NaHCO3 is a mild antiseptic for skin infections. It is
used in fire extinguishers.
BIOLOGICAL IMPORTANCE
OF SODIUM AND POTASSIUM
A typical 70 kg man contains about 90g of Na and 170g of K
compared with only 5g of iron and 0.06g of copper.
Potassium ions are present in higher concentration inside the
cells than sodium ions and they are present outside the cell in
blood plasma.
Because of large concentration gradient inside and outside the
cells, the transport of sodium ion into the cells is favoured. To
pump out these ions again from the cell to maintain
concentration gradient large driving force is carried out. The
energy for this process is provided by ATP molecules.
Thus both sodium and potassium ions are essential for living
organisms.
Anomalous properties of Lithium
Lithium
Symbol – Li
Atomic no. - 3
Atomic Weight – 6.94u
Electronic Configuration - 1s22s1
Group no. – 1
Period no. – 2
Group name – Alkali Metals
Block name – ‘s’
Standard State(298 K)- Solid
Color – Silvery-white/grey
Classification - Metallic
Anomalous Properties
• High melting & boiling point.
• Much harder than other alkali metals.
• Reacts with oxygen least readily to form normal
oxide(E.g. Li2O), whereas other alkali metals form
peroxides and superoxides(E.g. MO2,M2O2).
4Li + O2 → 2Li2O
• Unlike other alkali metals lithium reacts directly
with carbon to form an ionic carbide.
• The carbonates, hydroxides and nitrates of lithium
decompose on heating unlike those of other alkali
metals which are somewhat stable towards heat.
4LiNO3 → 2Li2O + 4NO2 + O2
2LiCO3 → 2Li2O + CO2
2LiOH → Li2O + H2O
• LiOH is a weaker base than hydroxides of other alkali
metals.
• Unlike elements of group 1, Lithium forms nitride
with nitrogen.
3Li + N → Li3N
• Lithium halides are have more covalent nature than
halides of other members of group 1.
Due to appreciable covalent nature, the halides and alkyls
of lithium are soluble in organic solvents.
Li + has very high hydration energy and charge/radius ratio,
therefore it acts as an excellent reducing agent in solution.
Li +
Li +
Li +
Li +
• Small size of atom results in relatively
high cohesive properties associated with
relatively strong inter-metallic bonding;
large atoms usually form weak bonds.
B
Diagonal Relationship
The properties of lithium are quite different from
the properties of other alkali metals. On the
other hand, it shows greater resemblance with
magnesium, which is diagonally opposite element
of it.
Similarly properties of Beryllium & Boron represent
that of Aluminium & Silicon respectively.
The main reasons for the anomalous behavior of
lithium are -:
C
The Reasons -:
(i) The extremely small size of Lithium
& its ion.
(ii) Greater polarizing power of
lithium ion ( Li+), due to its small size
which result in the covalent character in its
compounds.
(iii) Least electropositive character
and highest ionization energy as
compared to other alkali metals.
(iv) Non availability of vacant
d-orbitals in the valence shell.
Some More Examples
Examples For Diagonal Relationship
• Li and Mg form only normal oxides whereas Na forms
peroxide and metals below Na, in addition, forms
superoxide.
• Li is the only Group 1 element which forms nitride, (Li3N).
Mg, as well as other Group 2 elements, also form nitride.
• Lithium carbonate, phosphate and fluoride are sparingly
soluble in water. The corresponding Group 2 salts are
insoluble. (Think lattice and solvation energies).
• Both Li and Mg form covalent organometallic compounds.
LiMe and MgMe2 (of Grignard reagents) are both valuable
synthetic reagents. The other Group 1 and Group 2
analogues are ionic and extremely reactive (and hence
difficult to manipulate).
THE ALKALI METALS ARE
HIGHLY REACTIVE. CAUSING
CONTRIBUTING FACTORS ARE LARGE SIZE AND LOW
IONIZATION ENTHALPY.
THE REACTIVITY OF THESE METALS INCREASES DOWN
THE GROUP.
1. ALKALI METALS TARNISH IN DRY
AIR DUE TO FORMATION OF
THEIR OXIDES.
2. THEY BURN IN OXYGEN
VIGOURSLY.
3. THEY REACT WITH MOISTURE
FORMING HYDROXIDES.
EX4LI+O2
2LI2O
2NA+O2
NA2 O2
M+O2
MO2
OXIDATION STATE- +1
LITHUIM IS AN EXEPTION
REACTING DIRECTLY WITH
NITROGEN OF AIR TO FORM
THE NITRIDE.
DUE TO THEIR HIGH REACTIVITY
TO AIR AND WATER THEY ARE
KEPT IN
KEROSENE OIL.
2. REACTIVITY TOWARDS WATER2M+2H2O
2M+2OH+H2
THE ALKALI METALS REACT WITH
H2O TO FORM HYDROXIDE AND
H2.
REACTION WITH H2O IS
EXPLOSIVE FOR ALKALI METALS.
3. REACTIVITY TOWARDS
HALOGENSTHE ALKALI METALS READILY
REACT VIGOURSLY WITH
HALOGENS TO FORM IONIC
HALIDES,M+X-.
LITHUIM HALIDES ARE
CONVALENT DUE TO
POLARISATION.
5. REDUCING NATURE
LITHUIM-STRONGEST
SODIUM-LEAST POWERFUL
REDUCING AGENT.
A) M - M
(SUBLIMATION
ENTHALPY)
B) M=+M+E- (IONISATION
ENTHALPY)
C)+M+H2O=M+ (HYDRATION
ENTHALPY)
THE ALKALI METALS DISSOLVE IN
LIQUID AMMONIA GIVING DEEP
BLUE SOLUTIONS - CONDUCTING
IN NATURE.
THE PRODUCTS OF THIS
REACTION ARE -
1.HYDROGEN.
2. AMIDE ION.
THIS REACTION TAKES PLACE AT
SLIGHTLY ELEVATED
TEMPERATURES OR IN THE
PRESENCE OF CATALYST.
WHEN LIQUID AMMONIA IS
EXPOSED TO LIGHT OF
SPECTRA REGION OF UV LIGHT
THERE IS NO OBSERVABLE
CHANGE.
WHEN METTALIC SOLUTION ARE
SO EXPOSED,REACTION
OCCURS.THIS REACTION IN
PRESENCE OF UV IS COMPLETE
PHOTOCHEMICAL REACTION.
REACTION-
+M + E- + NH3
MNH2 +1/2H2
THE BLUE COLUR IS DUE TO THE
AMMONIATED ELECTRON
WHICH ABSORBS ENERGY IN
VISIBLE REGION OF LIGHT
IMPARTING BLUE COLOUR TO
SOLUTION. THE SOLUTION IS
PARAMAGNETIC.
IN CONCENTRATED SOLUTION THE
BLUE COLOUR CHANGES TO
BRONZE COLOUR AND BECOME
DIMAGNETIC.
1.LITHUIM IS USED TO MAKE
ALLOYS.EX-WITH LEAD TO MAKE
WHITE METAL,BEARINGS FOR
MOTORS ENGINES,WITH ALUMINIUM
TO MAKE AIRCRAFTS PARTS AND
WITH MAGNESIUM TO MAKE ARMOUR
PLATES.
LITHUIM IS ALSO USED TO MAKE
ELECTROCHEMICAL CELLS.
2. SODIUM IS USED TO MAKE A Na
or Pb ALLOY NEEDED TO MAKE
PbMe4. used as anti-knock additives
to petrol.
Liquid sodium metal is used as a
coolant in fast breeder nuclear
REACTORS.
3. POTASSIUM HAS A VITAL ROLE
IN BIOLOGICAL SYSTEMS.
POTASSIUM CHLORIDE IS USED AS
A FERTILISER.
POTASSIUM HYDROXIDE IS USED AS
AN EXCELLENT ABSORBENT OF
CARBON DIOXIDE.
4.CAESIUM IS USED IN DEVISING
PHOTOELECTRIC CELLS.
.