Transcript 12.S-elements of the ІІ А group.Alkaline earth metals.р

Lecture 12. S-elements of the ІІ А group.
Alkaline earth metals.
р-Elements of the ІІІА group.
Boron and Aluminium
Be
Ra
Mg
Ba
Sr
Ca
PhD Halina Falfushynska
Members of the s-Block Elements
IA IIA
Li
Be
Na
Mg
K
Ca
Rb
Sr
IA Alkali metals
Cs
Fr
Ba
Ra
IIA Alkaline Earth
metals
• High tendency to lose e- to form
positive ions
• Metallic character increases down
both groups
Group II
Be 1.5
Mg 1.2
Ca 1.0
• Low nuclear attraction for outer
electrons
• Highly electropositive
• Small electronegativity
Sr
1.0
Ba 0.9
Ra 0.9
Occurrence and Extraction
These elements are widely distributed in rock structures. The
main minerals in which magnesium is found are carnellite,
magnesite and dolomite.
Calcium is found in chalk, limestone, gypsum and anhydrite.
Magnesium is the eighth most abundant element in the
Earth’s crust, and calcium is the fifth.
It is extracted from sea-water by the addition of calcium
hydroxide, which precipitates out the less soluble magnesium
hydroxide. This hydroxide is then converted to the chloride,
which is electrolysed in a Downs cell to extract magnesium
metal. MgCl2 → Ca + Cl2;
cathode: Mg2+ + 2e → Mg0 ; anode: 2Cl- – 2e → Cl02
Beryllium: BeF2 + Mg  MgF2 + Be.
Biological occurrence of elements of group IIA
Beryllium's low aqueous solubility means it is rarely available to
biological systems, is usually highly toxic.
Magnesium and calcium are ubiquitous and essential to all
known living organisms. They are involved in Mg/Ca ion pumps,
magnesium functioning as the active center in some enzymes,
and calcium salts taking a structural role (e.g. bones).
Strontium and barium have a lower availability in the biosphere.
Strontium plays an important role in marine aquatic life,
especially hard corals. They use strontium to build their
exoskeleton. These elements have some uses in medicine, for
example "barium meals" in radiographic imaging, whilst
strontium compounds are employed in some toothpastes.
Flame test
Mg brilliant white
Ca brick red
Sr blood red
Ba apple green
HCl(aq)
sample
Oxide
Hydroxides
BeO
Be(OH)2
MgO
Mg(OH)2
CaO
Ca(OH)2
SrO
Sr(OH)2
BaO, Ba2O2 Ba(OH)2
Amphoteric to basic, base strength increase
Basic oxides, hydroxides
Reaction with water:
Oxide: O2- + H2O  2OH-
Predominantly ionic with fixed
oxidation state
Group II: Electropositive metals.
Low first and second I.E. but very high third
I.E.. Have a fixed oxidation state of +2.
Be and Mg compounds possess some degree
of covalent character.
Compare to
Group I: Most electropositive metals.
Low first I.E. and extremely high second I.E.
Atomic radii (nm)
Li
0.152 Be
0.112
Na
0.186 Mg
0.160
K
0.231 Ca
0.197
Fr
Li
Rb
0.244 Sr
0.215
Cs
0.262 Ba
0.217
Fr
0.270 Ra
0.220
Be
Ra
Ionization Enthalpy
1st I.E.
2000
600
Li
500
400
300
Na
Be+
1500
K
Rb
Cs
2nd IE
Ca+
1000
Ba+
Be
500
Ca
1st IE
Ba
Variation in Melting Points
1250
Be
1000
Ca
Sr
750
Ba
Mg
500
250
Li
Na
10
K
20
Rb
30
40
Cs
50
60
Reactions with oxygen
S-block elements reacts readily
with oxygen.
they have to be stored under
liquid paraffin
to prevent contact with the
atmosphere.
Normal
Oxide
Formed by
S-block elements are strong
reducing agents.
Their reducing power increases
down both groups.
(As the atomic size increases, it
becomes easier to
remove the outermost electron)
Peroxide
Be, Mg, Ca, Ba
Sr
Superoxide
None
Reaction with water
M(s)  M+(aq) + eH2O(l) + e-  OH-(aq) + ½ H2(g)
Li -3.05 volt
Na -2.71
K -2.93
Rb -2.99
Cs -3.20
Be -1.85 volt
Mg -2.38
Ca -2.87
Sr -2.89
Ba -2.90
Reaction with
hydrogen
All the II A elements
except Be react directly
with
hydrogen.
Ca(s) + H2(g)  CaH2(s)
The reactivity increases
down the group.
Only BeH2 and MgH2 are
covalent, others are ionic.
Reaction with chlorine
All the s-block metals react
directly with chlorine to
produce chloride.
Ba (s) + Cl2(g)  BaCl2 (s)
BeCl2 is essentially covalent
The lower members in group
II form essentially ionic
chlorides, with Mg having
intermediate properties.
Reactions of
chlorides
Group II chlorides show
some degree of covalent
character.
BeCl2 hydrolysis to
form Be(OH)2(s) and
HCl(aq).
MgCl2 is intermediate, it
dissolves and
hydrolysis slightly.
Other group II chlorides
just dissolve without
hydrolysis.
Reaction with nitrogen
These reactions can not
occur without extreme
circumstances.
A compound may be
created via really high
temperatures.
3Mg(s) + N2(g) -> Mg3N2(s)
Reactions
with acid
All the IIA metals react
directly with acid to produce
salt and hydrogen.
Mg (s) + H2SO4(l)→ MgSO4 (s)
+ H2 (g)
Mg + 2HCl -----> MgCl2 + H2
Reactions with alkali. Only for Beryllium!!!
Be (s) + 2NaOH(l) + 2H2O = Na2[Be(OH)4] (s) + H2
Reactions of oxides and hydroxides
1. All group I oxides reacts with water to form
hydroxides
Oxide: O2- + H2O  2OHPeroxide: O22- + 2H2O  H2O2 + 2OHSuperoxide: 2O2- + 2H2O  2OH- + H2O2 + O2
2. All group I oxides/hydroxides are basic and the
basicity increases down the group.
Thermal Stability of carbonates
BeCO3  BeO + CO2 ( at 100oC)
MgCO3  MgO + CO2 ( at 540oC)
CaCO3  CaO + CO2 ( at 900oC)
SrCO3  SrO + CO2 ( at 1290oC)
BaCO3  BaO + CO2 ( at 1360oC)
Li2CO3  Li2O + CO2 ( at 700oC)
All other group I carbonates are stable at ~800oC
Thermal Stability of hydroxides
Be(OH)2(s)  BeO(s) + H2O(g) H = +54 kJ/mol
Mg(OH)2(s)  MgO(s) + H2O(g) H = +81 kJ/mol
Ca(OH)2(s)  CaO(s) + H2O(g) H = +109 kJ/mol
Sr(OH)2(s)  SrO(s) + H2O(g) H = +127 kJ/mol
Ba(OH)2(s)  BaO(s) + H2O(g) H = +146 kJ/mol
All group I hydroxides are stable except LiOH
at Bunsen temperature.
Explanation of Thermal Stability
+
Decreasing
polarizing
power
+
+
-
Increasing
stability
-
-
Relative solubility of Group II
hydroxides
Compou Solubility / mol per
nd
100g water
Mg(OH)2
0.020 x 10-3
Ca(OH)2
1.5 x 10-3
Sr(OH)2
3.4 x 10-3
Ba(OH)2
15 x 10-3
Solubility of hydroxides
increases down the group.
Solubility of Group II sulphates
Compound Solubility / mol per 100g
water
MgSO4
3600 x 10-4
epsom salt
CaSO4
11 x 10-4
SrSO4
0.62 x 10-4
BaSO4
0.009 x 10-4
Solubility of sulphates
increases up the group.
Explanation of solubility
1. Group I compounds are more soluble than Group II
because the metal ions have smaller charges and
larger sizes. H lattice is smaller, and H solution is
more exothermic.
H solution =
-H lattice + H hydration
2. For Group II sulphates, the cations are much smaller
than the anions. The changing in size of cations does
not cause a significant change in H lattice (proportional
to 1/(r+ + r-).
However, the changing in size of cations does cause
H hydration (proportional to 1/r+ and 1/r-) to become less
exothermic, and the solubility decreases when
descending the Group.
SO42-
MgSO4
SO42-
SrSO4
3. For the smaller size anions, OH-.
Down the Group, less enthalpy is required to break the lattice
as the size of cation increases. However the change in H
solution is comparatively smaller due to the large value of 1/r- .
As a result, H solution becomes more exothermic
and the solubility increases down the Group.
Mg(OH)2
Sr(OH)2
Comparison of the physical properties and chemical
properties between alkaline earth and alkali metals
---The main difference is the electron configuration, which is ns2 for
alkaline earth metals and ns1 for alkali metals. For the alkaline earth
metals, there are two electrons that are available to form a metallic
bond, and the nucleus contains an additional positive charge. Also, the
elements of group 2A (alkaline earth) have much higher melting
points and boiling points compared to those of group 1A
---The alkaline earth metals are much harder and denser compare to
alkali metals.
---The elements of the group 2A contain a smaller atomic radius and
much higher ionization energy than the group 1A. Even though the
group 2A contains much higher ionization energy, they still form an
ionic compound with 2+ cations.
Uses of IIA group compounds
• Magnesium
– Magnesium hydroxide - Milk of magnesia, an antacid
– Magnesium can be obtained by eating food which is rich in
Magnesium, such as nuts and certain vegetables or by eating
supplementary diet pills.
• Calcium hydroxide
– To neutralize acids in waste water treatment
• Strontium compound
– Strontium ranelate is used in the treatment of osteoporosis. It is a
prescription drug in the EU, but not in the USA.
– Strontium chloride is sometimes used in toothpastes for sensitive
teeth. One popular brand includes 10% total strontium chloride
hexahydrate by weight.
Barium
Uses of barium sulfate include being
a radiocontrast agent for X-ray
imaging of the digestive system
Properties and Trends
in Group IIIA
Boron is a non-metal, whereas the other Group IIIA
(13) elements are metals.
• The most common oxidation state for Group 13 is +3.
Gallium, indium, and thallium also often form 1+ ions
by retaining their ns2 electrons; this is called the inert
pair effect.
• Aluminium and gallium are amphoteric, but indium
and thallium show more metallic character and do not
dissolve in alkalis.
Production of Aluminum
• The Hall–Heroult process is used to isolate
aluminum.
• Bauxite (aluminum ore) is first treated with NaOH
to form aluminate ion, Al(OH)4–.
• This solution is then diluted with water and
acidified slightly, precipitating Al(OH)3. The
Al(OH)3 is then heated to about 1200 oC and
decomposes to pure Al2O3.
• The Al2O3 is then electrolyzed using cryolite,
Na3AlF6, as the electrolyte and graphite (solid
carbon) as the electrodes.
Electrolysis Cell for
Aluminum Production
Carbon
aluminium
Boron
• Most of the chemistry of boron compounds is based on
the lack of an octet of electrons about the central boron
atom.
• These compounds are electron-deficient; the deficiency
causes them to exhibit some unusual bonding features.
• Boron hydride (BH3) forms a coordinate covalent bond
with another atom that has a lone pair of electrons to
complete its octet; this is called an adduct.
• Borax, Na2B4O7·10 H2O, a
hydrated borate, is the primary
source of boron.
• Sodium perborate is used as a
bleach; perborate’s structure is:
Structure of
Diborane, B2H6
The “regular” B—H bond
distance is shorter than the
B—H distance in the threecenter bond; why?
Properties and Uses
of Aluminum
• The reduction of Al3+(aq) to Al(s) occurs with difficulty.
• Thus, aluminum metal is a good reducing agent.
• As an active metal, aluminum readily reacts with acids to
produce hydrogen gas.
• Aluminum also dissolves in basic solutions.
• Because its combustion is a highly exothermic reaction,
powdered aluminum is used as a component in rocket
propellants and fireworks.
• Perhaps the most familiar use of aluminum is in beverage
cans, cookware, and as a foil for wrapping foods.
Reactivity towards air
• Boron is unreactive in crystalline form. Aluminium
forms a very thin oxide layer on the surface which
protects the metal from further attack. Amorphous
boron and aluminium metal on heating in air form
B2O3 and Al2O3 respectively. With dinitrogen at high
temperature they form nitrides.
Oxides, hydroxides, acids
Boron forms the oxide B2O3 and a large number of borate
anions containing trigonal planar BO3 units and/or tetrahedral
BO4 units.
• Boric acid (B(OH)3) is a monobasic acid. It acts as a Lewis acid,
interacting with water to form B(OH)3(OH2), which loses a
proton to form [B(OH)4]–
• A12O3 and Ga2O3 are amphoteric, but In2O3 and T12O3 are
basic.
• [M(H2O)6]3+ salts are acidic in aqueous solution due to
hydrolysis.
Anodized Aluminum
• The oxide layer on aluminum
protects it from further corrosion.
• The oxide layer may be enhanced
by making an aluminum object the
anode in an electrolysis apparatus.
• The thick, hard oxide layer that
results is porous and may be dyed.
• When the pores are sealed, the
color is a permanent part of the
object.
Reactivity towards acids and alkalies
• Boron does not react with acids and alkalies even
at moderate temperature; but aluminium
dissolves in mineral acids and aqueous alkalies
and thus shows amphoteric character.
Aluminium dissolves in dilute HCl and liberates
dihydrogen.
2Al(s) + 6HCl (aq) → 2Al3+ (aq) + 6Cl-(aq) + 3H2 (g)
However, concentrated nitric acid renders
aluminium passive by forming a protective oxide
layer on the surface.
Aluminium also reacts with aqueous alkali and
liberates dihydrogen
Reaction with halogens
The IIIA elements react with
halogens to form
trihalides (except TlI3).
2E(s) + 3 X2 (g) → 2EX3 (s)
(X = F, Cl, Br, I)
Halides
properties
The Lewis acidity of the
boron halides increases in
the order BF3 < BCl3 <
BBr3 < BI3. This is because it takes less
energy to distort the trigonal planar
geometry with larger halides. This trend can
be explained by weaker pπ–pπ bonding.
• Aluminium fluoride is a
solid with high ionic
character, but the other
aluminium halides have
structures containing
covalent bonds.