PRE-UNIVERSITYSEMESTER 2CHEMISTRYCHAPTER 4 :

GROUP 2

4.1 Physical Properties of Group 2 Group 2 are also known as alkali earth metal. The elements of

Group 2 and some basic physical properties are described as below

Name , symbol

Z

Atomic

radius/ nm

Melting

point (oC)

1st ionisation energy (kJ/mol)

Electronic configuration

Beryllium Be

4 0.112 1287 900 1s2 2s2

MagnesiumMg

12 0.160 650 738 1s2 2s2 2p6 3s2

CalciumCa

20 0.197 842 590 1s2 2s2 2p6 3s2 3p6 4s2

Strontium Sr

38 0.215 777 5501s2 2s2 2p6 3s2 3p6 3d10 4s2

4p6 5s2

Barium, Ba

56 0.218 727 5031s2 2s2 2p6 3s2 3p6 3d10 4s2

4p6 4d10 5s2 5p6 6s2

4.1.1 Atomic radius• Atomic radius depend on 2 factors – Nuclear charge – Screening effect• When going down to Group 2, both screening effect and

nuclear charge increase. However, the increase in screening effect is more significant, as more shell is used to filling in the electrons. This will cause the effective nuclear charge to decrease, resulting the electron cloud to be further away from the nucleus. Hence atomic radius increase.

4.1.2 Melting point• The melting point of the Group 2 generally decrease when

goes down to group. • All the elements occur as hexagonal closed-packed

structures with the exception of barium and radium, which adopt the more 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 as a result small unit cells.

4.1.3 Ionisation Energy.© The 1st ionisation energy decrease when goes down to

Group 2. © Atomic size when goes down to Group 2 which contribute

the decrease in ionisation energy. Furthermore, with the increase in atomic size, the number of shell also increase thus causing the screening effect to increase. This may also affected the effective nuclear charge as the distance between the electron and the nucleus is getting further.

© The third ionisation energy of Group 2 elements are extremely high, which suggested that the 3rd electron the withdrawn from an inner shell. Thus Group 2 elements only goes through 2nd ionisation energy and form a stable M2+.

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

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

Element Be Mg Ca Sr Ba

1st IE (kJ/mol) 900 740 590 550 500

2nd IE (kJ/mol) 2700 2190 1740 1610 1470

4.2 Chemical Properties of Group 2 Table below shows the E0 value of Group 2 elements, When

going down to Group 2, E0 value become more negative, indicates the reducing ability increase when going down to group, hence stronger reducing agent.

This may also indicates the reactivity increase when going down to Group 2. Henceforth, we shall discuss the reactivity of Group 2 elements with air (oxygen) and water.

Element Be Mg Ca Sr Ba

Eo / V - 1.85 -2.37 -2.87 -2.89 - 2.90

Trend of reducing agent

 Reducing strength increased

4.2.1 Reaction of Group 2 elements with oxygen (air). The Group 2 elements react with O2 to form the oxides. All

the elements except Be also form unstable peroxides (MO2). The oxides of Mg to Ra react with water to form the basic hydroxides while BeO and Be(OH)2 are amphoteric When BeO act as base : BeO + 2 H+ → Be2+ + H2O When BeO act as acid : BeO + 2 OH- + H2O →

Be(OH)42-

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-

On its nature, Beryllium is inert in air as its surface is passivated by the formation of a thin layer of BeO. Magnesium and calcium metals also tarnish in air with the formation of an oxide layer, but will burn completely to their oxides when heated. Strontium and barium, especially in powdered forms, ignite in air and are stored under hydrocarbon oils

• The oxides of the other Group 2 elements can be obtained by direct combination of the elements (except Ba, which forms the peroxide)

• Their melting points decrease down the group as the lattice enthalpies decrease with increasing cation radius. Magnesium oxide is a high-melting-point solid (as is BeO) and is used as a refractory lining in industrial furnaces. Like BeO, MgO has a high thermal conductivity coupled with a low electrical conductivity. This combination of properties leads to its use as an electrically insulating material around the heating elements of domestic appliances and in electrical cables

Element Reaction with oxygen ReactivityMelting point

of oxide

Be 2 Be + O2 2 BeO

 

 

Mg 2 Mg + O2 2 MgO

Ca 2 Ca + O2 2 CaO

Sr Sr + O2 SrO

Ba Ba + O2 BaO2

INC

REA

SE

DEC

RE

ASE

• The peroxides of Mg, Ca, Sr, and Ba are prepared by a variety of routes; only SrO2 and BaO2 can be made by direct reaction of the elements. All the peroxides are strong oxidizing agents and decompose to the oxide:

2 MO2 (s) → 2 MO(s) + O2 (g) Special note :

*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 dependence on the relative radii of the cations and anions. As O2– is smaller than O2

2–, 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

4.2.1.1 Reaction of Group 2 oxide with water : Properties of Group 2 hydroxide Beryllium oxide, BeO, is a white solid, which is insoluble in

water, with coordination number of 4, as expected for the small Be2+ ion. The oxides of the other Group 2 elements all adopt coordination number of 6. This is due to Beryllium does not have empty d-orbital available to coordinate more than 8 electrons at its center, while other Group 2 elements have.

Magnesium oxide is insoluble but reacts slowly with water to form Mg(OH)2; likewise CaO reacts with water to form the partially soluble Ca(OH)2.

The oxides of Sr and Ba, SrO and BaO, dissolve in water to form the strongly basic hydroxide solutions:

BaO(s) + H2O (l) → Ba2+ (aq) + 2OH- (aq)

Element

Reaction of metal oxide

with water

Rate of formation of base

Be no reaction / does not dissolve

Mg MgO + H2O ↔ Mg(OH)2   Ca CaO + H2O ↔ Ca(OH)2

Sr SrO + H2O Sr(OH)2

Ba BaO + H2O Ba(OH)2

INC

RE

AS

E

4.2.1 Reaction of Group 2 elements with water. All Group 2 react with water to from metal (II) hydroxide, M(OH)2,

with hydrogen gas liberated The reactivity of Group 2 with water increase (as suggested by

their E0 value). Beryllium react slowly under hot steam to form a white precipitate of beryllium hydroxide. Magnesium reacts similarly as beryllium does, however, compare to Be, the rate of reaction is higher.

Magnesium hydroxide, Mg(OH)2, is basic but only very sparingly soluble; beryllium hydroxide, Be(OH)2, is amphoteric and in strongly basic solutions it forms the tetrahydroxyberyllate ion, Be(OH)4

Calcium react slowly with water under room condition, to form a cloudy calcium hydroxide (also known as lime water). Limewater is well known to test the presence of carbon dioxide, where CO2 will turn limewater chalky and form white precipitate of calcium carbonate, which then dissolved when on further reaction with CO2 to form the hydrogencarbonate (also known as bicarbonate) ion

Ca(OH)2 (aq) + CO2 (g) → CaCO3 (s) + H2O (l)CaCO3 (s) + H2O (l) + CO2 (g) → Ca(HCO3)2 (aq)

Strontium and barium can react even in cold water to form a water soluble strong base of strontium hydroxide and barium hydroxide respectively. However, rate of reaction of barium is greater than strontium, hence more vigorous

Element

Condition of water

Reaction equation

Rate of

reaction

Ksp (mol3

dm-9)Solubilit

y

Be Hot steamBe + 2 H2O Be(OH)2 +

H2

 

6.92 x 10-

22

 

Mg Hot steamMg + 2H2O Mg(OH)2

+ H2

5.61 x 10-

12

Ca

Water at room

temperature

Ca + 2 H2O Ca(OH)2 +

H2

5.50 x 10-6

Sr Cold waterSr + 2 H2O Sr(OH)2 +

H2

7.24 x 10-6

Ba Cold waterBa + 2H2O Ba(OH)2 +

H2

2.54 x 10-4

INC

REA

SE

INC

REA

SE

Element

Group 2 carbonate Group 2 Nitrate

Formula

Decomposition

temperature

Stability Formula Stability

Be BeCO3 1590C

 

Be(NO3)2

 

Mg MgCO3 3500C Mg(NO3)2

Ca CaCO3 8320C Ca(NO3)2

Sr SrCO3 13400C Sr(NO3)2

Ba BaCO3 14500C Ba(NO3)2

INC

REA

SE

INC

REA

SE

4.3 Thermal Decomposition of Nitrates and Carbonates

• All nitrates of the Group 2 elements are decomposed by heat to form metal oxides, nitrogen dioxide and oxygen gases.

2 M(NO3)2 (s) 2 MO (s) + 4 NO2 (g) + O2 (g)• All carbonates of the alkaline-earth metals also decompose on

heating, producing metal oxides and releasing carbon dioxide gas.

MCO3 (s) MO (s) + CO2 (g)

The thermal stabilities of Group 2 nitrates and carbonates increase down the group from beryllium to barium. This means that the temperature needed to decompose the nitrates and carbonates increases down the group.

The trend of decomposition for Group 2 nitrate and carbonate can be explained below Magnesium nitrate and magnesium carbonate decompose easily at low

temperatures. This shows that the metal oxide is more stable than the nitrate and

carbonate. This can be explained by the fact that the size of the oxide ion, O2-, is smaller than that of the nitrate, NO3

-, and carbonate, CO32-

ions. As such, the oxide ion can approach closer to the Mg2+ cation forming a shorter and stronger bond

Magnesium nitrate, MgNO3 Magnesium nitrate, MgCO3

Mg2+

NO

O

O

2-

Mg2+

CO

O

O

2-

• Besides, magnesium ion has a high charge density ratio giving the ion a high polarisation power To polarise the electron clouds of the nitrate and carbonate ions.

• The electron clouds of the NO3- ion and CO3

2- ion are easily distorted, rendering the nitrogen-oxygen bonds in the NO3

- ions and the carbon-oxygen bonds in the CO3

2- ions are weak and easily broken.

• The smaller O2- ions are left attached to the magnesium ions.

2 Mg(NO3)2 (s) 2 MgO (s) + 4 NO2 (g) + O2 (g)MgCO3 (s) MgO (s) + CO2 (g)

• Barium nitrate and barium carbonate appear to be more stable than that of it's counterparts of the magnesium. A higher temperature is needed to decompose the salts. This is because the large size of the barium ion lowers the charge density ratio of the ion. The polarity of the ion depends directly on this ratio. A lower charge density ratio means the cation is less polarising. When barium ion approaches radius anion like NO3

- ion and CO32- ion, the electron clouds of the

anions will not be as distorted as when bonded with the magnesium ion. The bonds between Ba-NO3 and Ba-CO3 are more ionic and are much stronger than that of Ba-O.

4.4 Solubility of the Group 2 Sulphate

1. The solubility solubility of an ionic compound depends mainly on two factors:

a. lattice energy b. hydration energy

2. Heat of hydration is defined as the energy released when one mole of gaseous ions is hydrated by water molecules to form an infinite dilute solution under standard condition.

For cation : M+ (g) + water M+ (aq) ΔHhyd = – ve kJ/mol

For anion : X- (g) + water X- (aq) ΔHhyd = – ve kJ/mol

*Hydration energy is always exothermic since it involves the attraction of ions in the solute for water molecules. Similar to lattice energy, its magnitude depends on :

*Charge of the ion – higher the charge, the greater the heat of hydration. This is due to more heat energy is released as stronger bond are formed between the ion and molecules

*Size of the ion – the smaller the ion, the greater the heat of hydration ; the more heat released

3 The lattice energy of a crystalline substance refers to the amount of energy released when one mole of the ionic substance is produced from its ions in the gaseous state.

M+ (g) + X- (g) MX (s) ΔH = Lattice energy = – ve kJ/mol

*In a lattice that consists of cations and anions with charges Z+ and Z– between ionic distance (r+ + r–), the lattice energy can be expressed using the relation below.

*So, higher the lattice energy (more negative or more exothermic), the more stable the ionic compound formed, the more the energy required to break the strong electrostatic forces between the 2 opposite charged ions.

rr

QQenergyLattice

.

Ion Be2+ Mg2+ Ca2+ Sr2+ Ba2+

Charge density 64.5 30.8 20.2 17.7 14.8

ΔHhyd (kJ / mol) –2486 –1925 –1577 –1446 –1308

4. All sulphates of Group 2 elements are white crystalline and non-deliquescent solids. The solubility of these sulphates decrease down the group. Magnesium sulphate is soluble in water, calcium sulphate is slightly soluble and barium sulphate is insoluble.

5. This can be explained using the standard enthalpy of solution, ΔHsoln, where it is defined as the heat absorbed or released when one mole of crystal lattice is dissolved in water to form ionic aqueous solution under standard condition.

MX (s) M+ (aq) + X- (aq) ΔHsoln = + / – kJ / mol

*If the ΔHsoln, is positive, the salt will be insoluble in water & if ΔHsoln, is negative, it is soluble in water

6. Using Hess Law, the relationship between ΔHsoln, ΔHhyd, and LE can be explained using the chart below

M X (s) M+ (aq) + X- (aq)

negative

M+ (g) + X- (g)

7. A salt is soluble in water if its ΔHsolution is negative or exothermic. This happens when the salt has a high hydration energy and a low lattice energy. The more negative the enthalpy of solution, the more soluble will be the salt.

solutionH

energy LatticeH hydrationH

8. (a) The solubility of the sulphates decreases down Group 2 because the hydration energies of the ions decrease more rapidly than the lattice energies with increasing ionic size in the order Mg2+<Ca2+< Sr2+< Ba2+

(b) The lattice energies of the Group 2 sulphates decrease relatively slower because the magnitude depends on (r+ + r –). The anion radius (r–) for sulphate ion, SO4

2–, is too big compare to the cation radius (r+).

(c) As a results, the overall (r+ + r –). does not show any signifcant increase down the group. The lattice energy does not decrease very much from magnesium sulphate to barium sulphate.

9. On the other hand, the small size of the magnesium ion plus its high charge result in a lot of heat being released as hydration energy. This hydration energy decreases rapidly down the group from magnesium ion to barium ion because there is a significant increase in the size of the cations down the group. The resulting enthalpy of solution for barium sulphate becomes less negative compared to that of magnesium sulphate. Consequently, the solubility decreases down the group in the order:

BeSO4 > MgSO4 > CaSO4 > SrSO4 > BaSO4

Group 2 sulphate Be SO4 Mg SO4 Ca SO4 Sr SO4 Ba SO4

ΔHsolution (kJ / mol) -95.3 -91.2 + 17.8 + 18.70 +19.4

Solubility (g / 100mL) 41.0 36.4 0.21 0.010 0.00025

BeSO4 Mg SO4 CaSO4 Sr SO4 BaSO4

∆Hhydration

∆Hlattice energy

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, combined with 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

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. .

4.6 The anomalous properties of beryllium• The small size of Be2+ (ionic radius 27 pm) and its consequent high charge

density and polarising 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 significant covalent contribution to the bonding in compounds such as

the beryllium halides BeCl2, BeBr2, and BeI2 and the hydride, BeH2. A greater tendency to form complexes, with the formation of molecular

compounds such as Be4O(O2CCH3)6. 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.

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

• Another important general feature of Be is its strong diagonal relationship with Al: Both Be and Al form covalent hydrides and halides; the analogous

compounds of the other Group 2 elements are predominantly ionic. The oxides of Be and Al are amphoteric whereas the oxides of the rest

of the Group 2 elements are basic. 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.

Both elements form structures based on linked tetrahedra: 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.