Download ALKALI EARTH METALS Introduction Properties Beryllium

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ALKALI EARTH METALS
Introduction
Properties
Earth`s crust
abundance
Beryllium
1.9x10
Electron configuration
o
Melting point ( C)
o
Boiling point ( C)
Density
Common oxidation
num.
Physical State
Color
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-4
2
Magnesium
Calcium
2.9
5
2
2
[He]2s
[Ne]3s
[Ar]4s
1287
650
842
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2469
1.85
1090
1.74
1484
1.55
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2
solid
grey
2
solid
white
2
solid
White
The word “earth” was used for the oxides of group 2A by the
ancient scientists.
Alkaline earth metals, after the alkali metals, are secondary
metals with strong metallic properties.
The group 2A elements are less active than those of 1A, but
more active than those of group 3A.
Except Be, all form ionic compounds.
Be forms mostly covalent compounds.
The atomic radius increases from top to bottom.
The melting point, boiling point and the hardness of the
elements decrease top to bottom.
Beryllium is the hardest alkaline earth metal
Barium is the softest alkaline earth metal.
-2
-2
Compounds of CO3 and O are insoluble in water.
Malleable and ductile.
+2
+2
Excess Mg and Ca in water means that it is hard.
1. Occurrence
Minerals of Some Alkaline Earth Metals
Beryllium
Calcium
Beryl,
Limestone (marble),
Be3Al2(SiO3)6
CaCO3
Gypsum,
Magnesium
CaSO4.2H2O
Dolomite,
Fluorite,
CaCO3. MgCO3
CaF2
Carnallite,
Strontium
KCI.MgCI2.6H2O
Epsom salt,
Celestite,
MgSO4.7H2O
SrSO4
Strontianite,
SrCO3
Barium
Barytes,
BaSO4
Witherite,
BaCO3
Radium
All isotopes are
radioactive
2. Chemical Properties
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They are Covered by their oxides when they are exposed
to air.
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Burn in air with bright flame.
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Alkaline earth metals form MSO4, MCO3, MO, M(OH)2,
compounds where M alkaline earth metal.
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Activity of the metals increases from Be to Ba.
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Reactions of Ca, Sr and Ba are easy with water.
Reactions of Be and Mg with water are difficult.
+2
Ca(s) + 2H2O(l) → Ca (aq) + 2OH (aq) + H2(g)
Mg(s) + 2H2O(l) → Mg(OH)2(s) + H2(g)
2Mg(s) + O2(g) → 2MgO(s)
Mg (s) + H2(g) → MgH2(s)
Ca (s) + 2HCI(aq) → CaCl2 + H2(g)
Be (s) + 2HCI(aq) → BeCI2 + H2(g)
All oxides and hydroxides of alkaline earth metals show
basic properties.
Oxides and its hydroxides of beryllium show amphoteric
properties.
3. Compounds
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Ionic
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Have high melting and boiling points.
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Thermally stable.
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Less soluble in water than those of 1A metals.
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Solubility of salts increase from Be to Ba.
Calcium Oxide, CaO
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Obtained by heating calcium carbonate (limestone) at
900°C.
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CaCO3(s) → CaO(s) + CO2(g)
Calcium Carbonate, CaCO3
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It forms by evolving CO2 gas from Ca(OH)2 solution
Calcium Carbide, CaC2
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Small gray solid pieces
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Used widely in industry
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Prepared by the reaction of calcium oxide (lime) and coke.
CaO(s) + 3C(s) ↔ CaC2(s) + CO (g)
CaC2(s) + 2H2O(l) → C2H2(g) + Ca(OH)2(s)
Calcium Sulfate, CaSO4
CaSO4 : Gypsum
CaSO4.2H2O : Alabaster
CaSO4.1/2H2O : Plaster of Paris
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Gypsum is a white substance and used for fabrication into
wallboard.
Uses
Beryllium
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Very light element
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To prepare some special alloys.
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2% Be in Cu alloy in springs.
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Making glass for X-ray tubes.
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In the aerospace industry.
Magnesium
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Light metal
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Production of planes, missiles
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Some light household items.
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Flash in photography
Magnesium
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To increase the hardness, durability and resistance to
corrosion of alloys.
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Additive in rocket fuels and signal rockets.
Mg(OH)2 is used as a stomach antacid.
Calcium Compounds
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Construction materials, bleaching agents.
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CaO in the production of some chemicals, steel, glass,
paper and sugar.
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CaC2is used to produce acetylene gas for welding.
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In fertilizers and toothpaste
Barium and Its Compounds
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Ba is used as a gas absorbent in vacuum tubes.
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The alloys of Ba with Ni donate electrons so used in
vacuum tubes and in ignitors and spark plugs.
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Barium nitrate and barium chlorate are used to obtain
green light in fireworks.
Barium and Its Compounds
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BaSO4 is used to take the photos of stomach and
intestines.
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Ba does not allow X-rays to pass like all elements with
large atomic number.
+2
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The Ba ion is toxic, but since the solubility of BaSO4 in
water is very low, its poisonous effect disappears
Strontium
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The salts of strontium in signal rockets and fireworks to
produce red light.
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Nuclear batteries.
Radium
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In self luminous paints and in radiotherapy.
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As a neutron source.
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On luminous watch dials.