Group 2

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STPM CHEMISTRY
SEMESTER 2
CHAPTER 4 :
GROUP 2

CHAPTER 4 : GROUP 2
4.1 Selected Group 2 elements and their compounds
4.2 Anomalous behaviour of beryllium
4.3 Uses of Group 2 compounds

Topic 2009 2010 2011 2012 2013 2013 U 2014 2014 U
P1 P2 P1 P2 P1 P2 P1 P2 A B,C A B,C A B,C A B,C

17, 19b 20
4. Group 2 1 -- 1 -- -- 3a -- 8 1 18d 2 -- 1 20b 1 b, c

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4.1 Selected Group 2 elements and their compounds

Alkaline earth metals are a group of chemical
elements in the periodic table with very similar
properties. They are all shiny, silvery-white, somewhat
reactive metals at standard temperature and
pressure and readily lose their two outermost
electrons to form cations with charge 2+ and an
oxidation state (number) of +2. In the modern IUPAC
nomenclature, the alkaline earth metals comprise the
Group 2 elements. The elements in Group 2 discussed
are beryllium (Be), magnesium (Mg), calcium (Ca),
strontium (Sr) and barium (Ba).

Name , Atomic Electronic configuration
symbol Melting 1st IE

Z radius/ point (oC) (kJ/mol
nm

Beryllium, 1280 900 1s2 2s2
4 0.112 738 1s2 2s2 2p6 3s2
590 1s2 2s2 2p6 3s2 3p6 4s2
Be 550 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 5s2
503 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 5s2 5p6 6s2
Magnesium 651
12 0.160

, Mg

Calcium, 851
20 0.197

Ca

Strontium, 800
38 0.215

Sr

Barium, 56 0.218 725
Ba

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a) Atomic radius : - when going down to Group 2, atomic radius
increased. This is due to as the number of proton increased, nuclear
charge also increased. However, the number of electrons filling into the
shells increased, the screening effect increased. As a result, the
effective nuclear charge decreased, which caused the electrostatic
attraction forces between nucleus and outermost electron less stronger
hence increased the atomic radius.

b) Melting point :- The melting point of Group 2 generally decreased.
This is due to beryllium, magnesium and calcium have hexagonal
closed-packed structures while barium and radium have open body-
centred cubic structure. The density decreases from Be to Mg to Ca as
a result of very strong metallic bonding in the Group 2 elements, which
leads to short metal–metal distances in the lighter elements (225 pm in
beryllium, for instance) and hence a small unit cells.

c) Ionisation energy : The ionisation energy of Group 2 elements generally

decreased down Group 2. As atomic radius increased as a result of increasing

nuclear charge and screening effect, the effective nuclear charge decreased

down Group 2. Therefore, lesser energy required to remove one mole of

electrons from Group 2 atoms. Group 2 have greater tendency to form ion with

charge +2, due to very high third ionisation energy (as third electron is removed

from an inner shell with greater effective nuclear charge).

First Ionisation energy : M (g)  M+ (g) + e-

Second Ionisation energy : M+ (g)  M2+ (g) + e-

Element Be Mg Ca Sr Ba
1st ionisation energy (kJ/mol) 900 740 590 550 500
2nd ionisation energy (kJ/mol) 2700 2190 1740 1610 1470

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2. The chemical properties of Group 2 elements can be related with
the E0 value of the elements in Group 2. Table below shows the E0 value
of Group 2 elements

Element Be Mg Ca Sr Ba
- 1.85 -2.37 -2.87 -2.89 - 2.90
Eo / V
Trend of reducing Stronger reducing agent

agent

a) The ionisation energies of the elements decrease down the group
as the radius increases, and the elements become more reactive and more
electropositive as it becomes easier to form +2 ions. This decrease in
ionisation energy is reflected in the trend in standard potentials for the
M2+/M couples, which become more negative down the group.

b) The most significant of Group 2 metals reaction is the reaction
with water. Most of the Group 2 metals react with water and, therefore,
with any aqueous solution giving effectively an M2+, according to the
following equation :

M (s) + 2 H2O (l) → M2+ (aq) + 2 OH- (aq) + H2 (g)

Element Be Mg Ca Sr Ba

Reaction Does not React slowly React with React with cold water
condition react with steam water

Solubility of No product MgO Ca(OH)2 Sr(OH)2 Ba(OH)2
products in (weak base Slightly soluble Soluble in Soluble in
and insoluble
water in water water water
in water

Basic [Be(OH)2 is Base strength increased
properties amphoteric]
of product

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i. Beryllium is insoluble in water as it is naturally coated with an
impervious oxide layer of BeO (similar to aluminium). Therefore no
reaction take place when beryllium is dissolved in water.
ii. Compare to beryllium, magnesium is slightly more reactive as it
can react with water. However, magnesium will only slowly react when its
passed though with flow of steam over magnesium, to form magnesium
oxide according to the equation :

Mg (s) + H2O (g) → MgO (s) + H2 (g)
iii. Calcium, strontium and barium react with water under room
temperature. The reactivity increased down Group 2, which implies the
greater strength as reducing agent as the E0 value become increasingly
negative. The general equation for the reaction can be written as

M (s) + 2 H2O (l) → M(OH)2 (aq) + H2 (g)
Solubility of hydroxide in water increased down the group to form
strong base of metal hydroxide solution.

c) Group 2 elements also react with oxygen to form Group 2 oxide
and in some cases peroxide.

Element Be Mg Ca Sr Ba

Product BeO MgO CaO SrO2 BaO2
formed
Thermal Thermal stability of peroxide (MO2) increased down Group 2
stability
of oxide Thermal stability of oxide (MO) decreased down Group 2

and
peroxide

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i. When beryllium, magnesium and calcium heated with oxygen (air),
the metals ignite brightly to form metal oxide, MO. The equation can
be described as
2 M (Be, Mg, Ca) + O2 (g) → 2 MO (s)
Both beryllium oxide (BeO) and magnesium oxide (MgO) are white
solid with high melting point and low reactivity. Both possessed
excellent thermal conductivity but low electrical conductivity, which
lead to its use as a refractory material. However, since BeO is more
toxic, therefore, most of the time, MgO is preferred over BeO.

ii. Calcium oxide (as lime or quicklime) is used in large quantities in the
steel industry to remove P, Si, and S. When heated, CaO is
thermoluminescent and emits a bright white light (hence ‘limelight’).
Calcium oxide is also used as a water softener to remove hardness
by reacting with soluble carbonates and hydrogen carbonates to
form the insoluble CaCO3. It reacts with water to form Ca(OH)2,
which is sometimes known as slaked lime and is used to neutralize
acidic soils

iii. When strontium and barium are heated with oxygen, a metal
peroxide (MO2) is formed instead of metal oxide. The equation for the
reaction can be written as :

M (Sr, Ba) + O2 (g) → MO2 (s)
All Group 2 peroxides, MO2, are strong oxidising agent and are
reduced according to the equation :

2 MO2 (s) → 2 MO (s) + O2 (g)
iv. The thermal stability of the peroxides increases down the group
as the radius of the cation increases. This trend is explained by
considering the lattice enthalpies of the peroxide and the oxide and their
depend on the relative radii of the cations and anions. As O2–is smaller
than O22–, the lattice enthalpy of the oxide is greater than that of the
corresponding peroxide. The difference between the two lattice
enthalpies decreases down the group as both values become smaller with
increasing cation radius, therefore the tendency to decompose decreases.
Magnesium peroxide, MgO2, is consequently the least stable peroxide

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4.1.2 Compounds of Group 2 elements

1. Oxide of Group 2 are formed by direct heating of Group 2
metal with oxygen, except for strontium oxide and barium oxide.
However, generally most of these oxide are prepared via
decomposition of Group 2 carbonate, where

MCO3 (s) → MO (s) + CO2 (g)

Element Be Mg Ca Sr Ba

Formula of BeO MgO CaO SrO BaO
oxide

Solubility in Insoluble Slightly soluble in water Soluble in water
water Soluble in steam

Acid-base Amphoteric Basic
properties

pH in -- 8 - 10 12 - 14
aqueous

a) Beryllium oxide, BeO, is insoluble in water due to its high covalent

properties (as it has a high charge and small ionic radius), therefore is not

easily dissolve in water. However, BeO reacts with both acid and alkali

solution according to the equations :

When BeO act as base : BeO + 2 H+ → Be2+ + H2O
When BeO act as acid : BeO + 2 OH- + H2O → Be(OH)42-
b) Magnesium oxide, MgO, and calcium, CaO, are slightly dissolve in

water but easily dissolve in steam to form basic solution of magnesium

hydroxide and calcium hydroxide according to the equation :

MgO or CaO in steam : MgO (s) + H2O (g) → Mg(OH)2 (aq)
Though the basicity formed is not as strong as in SrO and BaO in water.

c) Strontium oxide, SrO, and barium oxide, BaO, are soluble in water to

form strong alkaline solution of strontium hydroxide, Sr(OH)2, and barium
hydroxide, Ba(OH)2.

SrO or BaO in steam : BaO (s) + H2O (l) → Ba(OH)2 (aq)

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2. Hydroxide of Group 2 have the general formula of M(OH)2. They
were mostly formed from the reaction between metal or metal oxide with
water (as describe above)

Reaction of metal with water : M (s) + 2 H2O (l) → M(OH)2 + H2 (g)
Reaction of metal oxide with water : MO (s) + H2O (l) → M(OH)2 (aq)

Element Be Mg Ca Sr Ba

Formula of Be(OH)2 Mg(OH)2 Ca(OH)2 Sr(OH)2 Ba(OH)2
hydroxide

Ksp of 6.92 x 10-22 5.61 x 10-12 5.50 x 10-6 7.24 x 10-6 2.54 x 10-4
hydroxide

Solubility Solubilities increased down Group 2 hydroxide
in water

Acid-base Amphoteric Basic
properties

a) Beryllium hydroxide is a white gelatinous solid insoluble in water. Due to
its high covalency yet an ionic compound, it gives the amphoteric
properties of Be(OH)2.
When Be(OH)2 act as base : Be(OH)2 + 2 H+ → Be2+ + 2 H2O
When Be(OH)2 act as acid : Be(OH)2 + 2 OH- → [Be(OH)4]2-

b) The hydroxides become apparently more basic down the group because
their solubility increases from Mg(OH)2 to Ba(OH)2. Magnesium
hydroxide, Mg(OH)2, is sparingly soluble and forms a mildly basic
solution because a saturated solution contains a low concentration of the
OH– ions

Dissociation of Mg(OH)2 : Mg(OH)2 (s) ↔ Mg2+ + 2 OH- (aq)
Since barium hydroxide has high Ksp value, therefore, the basicity is the
strongest, with the dissociation is considered as one way reaction.

Dissociation of Ba(OH)2 : Ba(OH)2 (s) → Ba2+ + 2 OH- (aq)

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c) Calcium hydroxide (commercially known as lime water) when react
with carbon dioxide, a white precipitate of calcium carbonate is first
observed

Reaction of Ca(OH)2 with CO2 :
Ca(OH)2 (aq) + CO2 (g) → CaCO3 (s) + H2O (l)

However, under excess CO2, white precipitate eventually dissolved to
form a clear solution of calcium hydrogencarbonate.

Further reaction : CaCO3 (s) + CO2 (g) + H2O (l) → Ca(HCO3)2 (aq)

1. Group 2 carbonate and nitrates when heated will decomposed to
form different products. The decomposition equations for these compound are

Dissociation of Group 2 nitrate :

2 M(NO3)2 (s) → 2 MO (s) + O2 (g) + 4 NO2 (g)
Dissociation of Group 2 carbonate :

MCO3 (s) → MO (s) + CO2 (g)

Formula of BeCO3 MgCO3 CaCO3 SrCO3 BaCO3
carbonate 1340 1450
159 350 832
Decomposition
temperature /0C Stability of Group 2 carbonate increase

Stability of
Group 2
carbonate

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a) When going down to Group 2, thermal stability of Group 2
carbonate increased. The equation for the decomposition of Group
2 carbonate is shown below

Dissociation of Group 2 carbonate : MCO3 (s) → MO (s) + CO2 (g)
b) This can be explained based on the cationic radius of Group 2. As

cationic radius of Group 2 increased from Be2+ < Mg2+ < Ca2+ <
Sr2+ < Ba2+, the charge density of the cation decreased.
c) Small size and high charge of Be2+ has high charge density. As a
result, Be2+ has high polarising power. This allow Be2+ to polarise
the large carbonate ion, CO32- and caused the O in carbonate to
attract closer to Be2+, rendering (weakening) the C-O bond in
carbonate, hence decomposed. Furthermore, beryllium oxide, BeO,
is more stable than beryllium carbonate, BeCO3, which makes the
reaction more easily to occur at low temperature.

DECOMPOSITION OF GROUP 2 NITRATES

Formula of Be(NO3)2 Mg(NO3)2 Ca(NO3)2 Sr(NO3)2 Ba(NO3)2
carbonate Stability of Group 2 nitrate increase
Stability of

Group 2

a) When going down to Group 2, thermal stability of Group 2 nitrate
increased.

Dissociation equation : 2 M(NO3)2 (s) → 2 MO (s) + O2 (g) + 4 NO2 (g)
b) This can be explained based on the cationic radius of Group 2. As
cationic radius of Group 2 increased from Be2+ < Mg2+ < Ca2+ < Sr2+ <
Ba2+, the charge density of the cation decreased.
c) Small size and high charge of Be2+ has high charge density. As a
result, Be2+ has high polarising power. This allow Be2+ to polarise the large
nitrate ion, NO3- and caused the O in nitrate to attract closer to Be2+,
rendering (weakening) the N-O bond in nitrate, hence decomposed.
Furthermore, beryllium oxide, BeO, is more stable than beryllium nitrate,
Be(NO3)2, which makes the reaction more easily to occur at low temperature.

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4.2 Anomalous behaviour of beryllium

1. The small size of Be2+(ionic radius 27 pm) and its consequent high
charge density andpolarising power results in the compounds of Be being
largely covalent; the ion is a strong Lewis acid. The coordination number
most commonly observed for this small atom is 4 and the local geometry
tetrahedral. Some consequences of these properties are:

a) A significant covalent contribution to the bonding in compounds such as
the beryllium halides (BeCl2, BeBr2, and BeI2) and the hydride, BeH2.

b) A greater tendency to form complexes, with the formation of molecular
compounds such as Be4O(OCOCH3)6.

c) Hydrolysis (deprotonation) of beryllium salts in aqueous solution, forming
species such as [Be(H2O)3OH]+and acidic solutions. Hydrated beryllium
salts tend to decompose by hydrolysis reactions, where beryllium oxo or
hydroxo salts are formed, rather than by the simple loss of water.

d) Beryllium forms many stable organometallic compounds, including
methylberyllium (Be(CH3)2), ethylberyllium (Be(CH3CH2)2), t-
butylberyllium (Be(C(CH3)2).

e) Another important general feature of Be is its strong diagonal
relationship with Al:

i. Both Be and Al form covalent hydrides and halides; the analogous
compounds of the other Group 2elements are predominantly ionic. Both
BeCl2 and AlCl3 exist as dimer

BeCl2 AlCl3

ii. The oxides of Be and Al are amphoteric whereas the oxides of the rest
of the Group 2 elements are basic.

iii. In the presence of excess OH– ions, Be and Al form [Be(OH)4]2– and
[Al(OH)4]–respectively, however, no equivalent chemistry is observed for
Mg.

iv. Both elements form structures based on linked tetrahedral : Be forms
structures built from [BeO4]n- and [BeX4]n- tetrahedra (X = halide) and Al
forms numerous aluminates and aluminosilicates containing the[AlO4]n-
unit.

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Beryllium and Aluminium Magnesium, Calcium,
Strontium, Barium

Doubtful existance of simple Be2+ or Al3+ ion All exist as M2+ ions and
; as compound are mainly covalent due to main compounds are
high charge density (high charge and ionic.
small ionic radius)

Metal react with concentrated alkali No reaction with alkali
solution (regardless
Be + 2 OH- + 2 H2O → Be(OH)42- + H2 concentrated or diluted)
2 Al + 2 OH- + 6 H2O → 2 [Al(OH)4]- + 3 H2

Both BeO and Al2O3 are amphoteric
BeO + 2 OH- + H2O → Be(OH)42-
BeO + 2 H+ → Be2+ + H2O
All MO are basic

Al2O3 + 2 OH- + 3 H2O → 2 [Al(OH)4]-
Al2O3 + 6 H+ → 2 Al3+ + 3 H2O

Both Be(OH)2 and Al(OH)3 are amphoteric
Be(OH)2 + 2 OH- → Be(OH)42-
Be(OH)2 + 2 H+ → Be2+ + 2 H2O
Al(OH)3 + OH- → [Al(OH)4]- All M(OH)2 are basic.

Al(OH)3 + 3 H+ → 2 Al3+ + 3 H2O

Both BeO and Al2O3 are insoluble in Solubility in water increase down
water and resistant to the oxidation Group 2 oxide, and the reactivity
with air also increase.
by air due to the protective oxide

layer.

BeCl2 and AlCl3 are covalently
bonded molecules. They exist as Group 2 chloride are ionic

dimer, hence cannot conduct compounds, which can conduct

electricity in molten state. Their electricity in molten or aqueous

chlorides are readily to hydrolysed in state. Though MgCl2 can slightly

water. hydrolysed in water (hence slightly

acidic) where :

[Al(H2O)6]3+ + H2O → [Mg(H2O)6]2+ + H2O→
[Al(OH)(H2O)5]- + H3O+ [Mg(OH)(H2O)5]+ + H3O+

[Be(H2O)4]2+ + H2O →
[Be(OH)(H2O)3]- + H3O+

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4.5 Application of Group 2 elements and compounds
4.5.1 Beryllium and its compound
• Beryllium is unreactive in air on account of a passivating layer of an

inert oxide film on its surface, which makes it very resistant to
corrosion. This inertness, coBmebryinlleiudmwith the fact that it is one of the
lightest metals, results in its use in alloys to make precision
instruments, aircraft, and missiles.
• It is highly transparent to X-rays due to its low atomic number (and
thus electron count) and is used for X-ray tube windows.
• Beryllium is also used as a moderator for nuclear reactions (where it
slows down fast-moving neutrons through inelastic collisions)
because the beryllium nucleus is a very weak absorber of neutrons
and the metal has a high melting point.
• As beryllium oxide is extremely toxic and carcinogenic by inhalation
and soluble beryllium salts are mildly poisonous, the industrial
applications of beryllium compounds are limited; BeO is used as an
insulator in high-power electrical devices where high thermal
conductivity is also necessary

4.5.2 Magnesium and its compound

• Most of the applications of elemental magnesium are based on the formation of light
alloys, especially with aluminium, that are widely used in construction in applications
where weight is an issue, such as aircraft. A magnesium–aluminium alloy was
previously used in warships but was discovered to be highly flammable when
subjected to missile attack.

• Some of the uses of magnesium are based on the fact that the metal burns in air with
an intense white flame, and so it is used in fireworks and flares.

• Various applications of magnesium compounds include ‘Milk of Magnesia’, Mg(OH)2,
which is a common remedy for indigestion, and ‘Epsom Salts’, MgSO4.7H2O, which
is used for a variety of health treatments, including as a treatment for constipation, a
purgative, and a soak for sprains and bruises.

• Magnesium and calcium are of great biological importance. Magnesium is a
component of chlorophyll but also it is coordinated by many other biologically
important ligands, including ATP (adenosine triphosphate). It is essential for human
health, being responsible for the activity of many enzymes. The recommended adult
human dose is approximately 0.3 g per day and the average adult contains about 25
g of magnesium

• Magnesium oxide, MgO, is used as a refractory lining for furnaces. Organo-
magnesium compounds are widely used in organic synthesis as Grignard reagents

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4.5.3 Calcium and its compound
• The compounds of calcium are much more useful than the element itself.

Calcium oxide (as lime or quicklime) is a major component of mortar and
cement. It is also used in steelmaking and papermaking.
• Calcium sulfate dihydrate, CaSO4.2 H2O is widely used in building
materials, such as plasterboard, and anhydrous CaSO4 is a common drying
agent.
• Calcium carbonate is used in the Solvay process for the production of
sodium carbonate and as the raw material for production of CaO.
• Calcium fluoride is insoluble and transparent over a wide range of
wavelengths. It is used to make cells and windows for infrared and
ultraviolet spectrometers.
4.5.4 Strontium and its compound
• Strontium is used in pyrotechnics phosphors, and in glasses for the now
rapidly declining market for colour television tubes.

4.5.6 Barium and its compound
• Barium compounds, taking advantage of the large number of electrons of

each Ba2+ ion, are very effective at absorbing X-rays: they are used as
‘barium meals’ and ‘barium enemas’ to investigate the intestinal tract.
Barium is highly toxic, so the insoluble sulfate is used in this application.
• Barium carbonate is used in glassmaking and as a flux to aid the flow of
glazes and enamels. It is also used as rat poison.
• Barium sulphide (BaS) has been used as a depilatory, to remove
unwanted body hair.
• Barium sulphate (BaSO4) is pure white, with no absorption in the visible
region of the electromagnetic spectrum, and it is used as a reference
standard in UV-visible spectroscopy. Soon after its discovery, radium was
used to treat malignant tumours; its compounds are still used as
precursors for radon used in similar applications. .

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