22.Group II Lecture Notes

Hwa Chong Institution 22 – Group II

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22 Group II


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3 INTRODUCTION The Group II metals are also known as the alkaline earth metals for two reasons: (a) Their oxides (white solids) form alkaline solutions when dissolved in water.

CaO(s) + H2O(l) Ca(OH)2(s)

lime slaked lime

Ca(OH)2(s) + H2O(l) Ca(OH)2(aq) slaked lime limewater

(alkali) (b) An artefact of history from the predecessors of chemistry – alchemists

In ancient times, alchemists attempted to obtain pure metals from their ores by

heating. They often encountered the Group II oxides in their furnaces. These Group

II oxides melted at such high temperatures that they remained as solids in the

alchemists’ fires.

These solids were referred to as “earth” which to the alchemists were materials

which did not melt just as the sand and soil found in the ores also did not melt on

account of their extremely high melting point.

3.1 ELECTRONIC CONFIGURATION

The outermost shell electronic configuration is ns2 Group II metals have a fixed oxidation state +2

Table 3.1 – Electronic configuration and ionisation energies of Group II elements

Element Electronic configuration 1st I.E.

/ kJ mol1

2nd I.E.

/ kJ mol1

3rd I.E.

/ kJ mol1

Be 1s22s2 900 1760 14800

Mg 1s22s22p63s2 736 1450 7740

Ca 1s22s22p63s23p64s2 590 1150 4940

Sr 1s22s22p63s23p63d104s24p65s2 548 1060 4120

Ba 1s22s22p63s23p63d104s24p64d105s25p66s2 502 966 3390


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Why do Group II metals have a fixed oxidation state of +2? The +2 oxidation state

The outermost 2 electrons in the ns shell are easily removed as they are shielded from the nucleus

by the inner core electrons. Hence +2 oxidation state is easily and mainly formed.

The +1 oxidation state

Although the +1 oxidation state seems possible given that the 2nd IE is considerably greater than that

of the 1st IE (about twice as high), the much higher lattice energy of MX2 compounds will cause MX

compounds to disproportionate readily to give MX2 and M.

2MX MX2 + M

The enthalpy change of formation of MX2 compounds is about twice as negative as that for MX

compounds and so MX2 compounds are relatively more stable.

The +3 oxidation state

The +3 oxidation state is not possible as the 3rd IE is significantly greater than either the 1st or 2nd. As

the 3rd electron would be removed from an inner quantum shell, this process requires too much

energy which cannot be compensated by the higher lattice energy of MX3 compounds.

Lecture Exercise 3.1

1 Which statement explains why calcium and chlorine react to form CaCl2 instead of CaCl? A Less energy is required to remove one electron from the calcium atom than to

remove two electrons. B More energy is released in forming chloride ions from chlorine molecules in the

formation of CaCl2 rather than in the formation of CaCl. C The lattice energy of CaCl is less exothermic than that of CaCl2. D When CaCl is formed from its elements, more energy is released than when CaCl2

is formed from its elements.


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3.2 PHYSICAL PROPERTIES Table 3.2 – Physical properties of Group II elements

Element Boiling point

/ oC Melting point

/ oC Atomic radius

/ nm Ionic radius

/ nm Density

/ g cm3

Be 2471 1287 0.112 0.031 1.85

Mg 1090 650 0.160 0.065 1.74

Ca 1484 842 0.197 0.099 1.54

Sr 1382 777 0.215 0.113 2.64

Ba 1897 727 0.217 0.135 3.62

3.2.1 Atomic (Metallic) & Ionic Radii

The metallic and ionic radii increase down Group II.

Figure 3.3 – Trends in metallic and ionic radii of Group I and II

Down the group, an increasing number of quantum shells increases the size of the atom.


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3.2.2 1st & 2nd Ionisation Energies

Both the 1st and 2nd ionisation energies for the Group II metals show a decreasing trend on going down the group.

Figure 3.4 – Trends in 1st and 2nd ionization energy for Group II metals

Down the group, the atomic radii increases (number of quantum shells increases), and the valence electrons are further from the nucleus. Hence down the group, less energy is required to remove the valence electrons, leading to the decrease in 1st and 2nd ionisation energies.

3.2.3 Melting & Boiling Points

The melting points of the Group II metals generally decrease down the group. However, there are no simple explanations to the irregularities in the trends; it may be due to the different crystal structures of the metals.

Figure 3.5 – Trends in melting and boiling point for Group II metals

As the size of the cations increases down the group, the attraction between the sea of delocalised electrons and the metal cations decreases. So metallic bonding weakens down the group.

2nd I.E.

Ist I.E.


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Comparing melting and boiling points of Group I and Group II metals

The melting and boiling points of the Group II metals are significantly higher than that for

the Group I metals (alkali metals). For example, the melting point of Ca (850 oC) would be

sufficient to ensure the vaporisation of K (766 oC).

The reasons for the relatively higher melting and boiling points of the Group II metals

(compared to Group I metals) are:

1) Group II metals have two valence electrons but their Group I counterparts have only

one

2) Group II metals have smaller cationic size due to greater number of protons

So, Group II cations have higher charge density than their Group I counterparts and thus,

more energy is required to break the stronger metallic bonding in Group II metals which

results in their higher melting and boiling points.

3.2.4 Hardness & Densities

The Group II metals are much harder metals when compared to the Group I metals but

much softer when compared to the transition metals like Fe.

Hardness is dependent on the metallic bond strength and the atomic packing in the metal

lattice.

1) The Group II metals have two valence electrons compared to one valence electron for

Group I metals.

2) The size of the Group II cations is smaller than that for Group II cations.

Metallic bonding is stronger in Group II metals compared to Group I metals and Group II

metal cations are closer-packed in the metallic lattice than Group I metals. As such, Group II

metals are harder as compared to Group I metals.

The density of the Group II metals generally increases down the group. The reason is that

down the group, the atomic mass increases faster than the atomic volume.


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3.2.5 Conductivities of Heat & Electricity

The Group II metals are all good conductors of heat and electricity. The mobile valence electrons are able to transmit thermal energy via rapid translations

within the metal lattice, exciting the lattice by cation-electron collisions so that heat is

transmitted via thermal vibrations of the metal lattice.

The presence of mobile valence electrons enables them to be charge carriers when a

voltage is applied. Hence the metals conduct electricity.

Table 3.6 – Thermal conductivity and electrical resistivity values for Group II metals (and sulfur)

Element Thermal conductivity / W m1 K1 Electrical resistivity / 108 W m

Mg 160 4.4

Ca 200 3.4

Sr 35 13

Ba 18 35

S 0.205 > 1023

3.2.6 Solubility – Sulfates and Hydroxides

It is known that there are trends in the solubility of Group II sulfates and hydroxides.

Solubility of Group II sulfates decreases down the group whereas solubility of Group II

hydroxides increases down the group.

An easy way to remember the trend is to know the solubilities of the following:

1) MgSO4 and Ba(OH)2 are soluble.

2) BaSO4 and Mg(OH)2 are insoluble or sparingly soluble.

To understand these trends, we have to make use of what we have learnt under Chemical

Energetics. To simplify, we can compare lattice energies and hydration enthalpies to predict

the trend of solubilities [not in syllabus]. Refer to Appendix B for the explanations.


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1 The Group II metals have higher melting points than the Group I metals. Which of the following factors could contribute to the higher melting points?

1 There are smaller inter-atomic distances in the metallic lattices of the Group II metals. 2 Two valence electrons are available from each Group II metal atom for bonding in the

metallic lattice. 3 Group II metals have higher first ionisation energies.

2 Which of the following elements is likely to be in Group II of the Periodic Table? [ = ohm]

Element Melting point / oC Density / g cm3 Electrical conductivity / 1 m1 A 98 0.97 2.4 × 107 B 113 2.07 5.0 × 1016 C 649 1.74 2.2 × 107 D 1744 11.3 6.0 × 107

3.3 CHEMICAL PROPERTIES

3.3.1 Relative Strength of Group II Metals as Reducing Agents

Reactivity of the Group II metals increases down the group. As atomic radii increases, the metal atoms lose their electrons more readily (1st and 2nd ionisation energies decrease) going down the group. So, they form M2+ cations more easily.

M2+(aq) + 2e M(s)

Therefore, the metals become better reducing agents on going down the group i.e. reducing power increases (tendency to be oxidised increases) down the group. This is illustrated in the increasingly negative Eo values down the group.

Table 3.7 – Standard electrode potential values of Group II metals

Element Standard electrode potential, Eo / V

Mg2+(aq) + 2e− Mg(s) −2.37

Ca2+(aq) + 2e− Ca(s) −2.87

Sr2+(aq) + 2e− Sr(s) −2.89

Ba2+(aq) + 2e− Ba(s) −2.91


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3.3.2 Reaction of Group II Metals with Oxygen

All Group II metals except Be tarnish in air, forming a thin oxide layer. Be reacts only when ignited in air.

2M(s) + O2(g) 2MO(s) The reactivity of Group II metals with oxygen increases down the group, a consequence of the increase in reducing strength of the metals. Table 3.8 – Reaction of Group II metals with oxygen

Element Reaction with O2 Colour of flame Reactivity

Mg 2Mg(s) + O2(g) 2MgO(s) Brilliant white Very slow

Ca 2Ca(s) + O2(g) 2CaO(s) Brick-red Slow

Sr 2Sr(s) + O2(g) 2SrO(s) Crimson red Fast

Ba 2Ba(s) + O2(g) 2BaO(s) — Explosive, kept under oil

3.3.3 Reaction of Group II Metals with Water

All Group II metals except Be reduce water to hydrogen gas (see Appendix A for Anomalous Behaviour of Beryllium).

M(s) + 2H2O(l) M(OH)2 + H2(g)

Reactivity of the metals with water increases on going down the group. Table 3.9 – Reaction of Group II metals with water

Element Reaction with H2O Reactivity

Mg Mg(s) + H2O(g) MgO(s) + H2(g) Reacts with steam

Mg(s) + 2H2O(l) Mg(OH)2(s) + H2(g) Very slow reaction with cold water

Ca Ca(s) + 2H2O(l) Ca(OH)2(aq) + H2(g) Slow reaction with cold water

Sr Sr(s) + 2H2O(l) Sr(OH)2(aq) + H2(g) Vigorous reaction with cold water

Ba Ba(s) + 2H2O(l) Ba(OH)2(aq) + H2(g) Very vigorous reaction with cold water

Effervescence of H2(g) is seen. The hydroxides formed are not very soluble, but they get more soluble as you go down the group (see Appendix B for Solubility of Group II Compounds). Note: Calcium reacts with small amounts of water to form the white powdery solid slaked

lime Ca(OH)2(s). In large amounts of water, limewater Ca(OH)2(aq) is formed. All the hydroxides are basic except for Be(OH)2 which is amphoteric.


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3.3.4 Reaction of Group II Oxides with Water (a) All the metal oxides of Group II:

are white solids

are ionic compounds

have extremely high melting point (highly refractory)

are basic except for the amphoteric BeO

Table 3.10 – Melting points of Group II oxides

Compound MgO CaO SrO BaO

Melting point / oC 2830 2900 2530 1973

Group II metal oxides can be prepared by the thermal decomposition of the metal carbonates, except for BeO.

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

(b) All the metal oxides react with water to give the corresponding metal hydroxide except for

BeO. Reaction with water increases with vigour down the group as the magnitude of the lattice energies of the metal oxides decreases.

MO(s) + H2O(l) M(OH)2(s)

(c) The resistance of MgO to hydration by water passivates the metal such that it has no further

reaction with oxygen or moisture in the atmosphere once the oxide film is formed over the surface.

Table 3.11 – Reaction of Group II oxides with water

Compound Reaction with H2O Reactivity

MgO

MgO(s) + H2O(l) Mg(OH)2(s)

Mg(OH)2 is sparingly soluble in water to give an alkaline solution.

Reacts very slowly with cold water

CaO

CaO(s) + H2O(l) Ca(OH)2(aq)

Although Ca(OH)2 has limited solubility, it still dissolves in water to give an alkaline solution.

Reacts rapidly with heat given off

SrO

SrO(s) + H2O(l) Sr(OH)2(aq)

Sr(OH)2 dissolves in water to give a strongly alkaline solution.

Reacts vigorously with heat given off

BaO

BaO(s) + H2O(l) Ba(OH)2(aq)

Ba(OH)2 dissolves in water to give a strongly alkaline solution.

Very vigorous reaction with a lot of heat given off


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O

C

O

M2+ O2

− +

HEAT

metal carbonate

-O

C

O-

O

M2+

oxide carbon dioxide

Lecture Exercise 3.3 Which of the following statements concerning Group II elements, calcium, strontium and barium are correct?

1 Their oxides are amphoteric. 2 Aqueous solutions of their hydroxides have a pH greater than 7. 3 The elements react with cold water liberating hydrogen.

3.4 THERMAL STABILITY

Group II nitrates, carbonates and hydroxides are all unstable towards heat. Thermal decomposition of these compounds gives stable oxides.

Nitrate : M(NO3)2(s) MO(s) + 2NO2(g) + ½O2(g)

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

Hydroxide : M(OH)2(s) MO(s) + H2O(g)

Table 3.12 – Decomposition temperatures of Group II carbonates, nitrates and hydroxides

Element Decomposition Temperatures / oC

Carbonates Nitrates Hydroxides

Mg 400 450 300

Ca 900 575 390

Sr 1280 635 466

Ba 1360 675 700

Group II nitrates, carbonates and hydroxides become more thermally stable as you go down the group. The ones lower down in the group have to be heated more strongly before they will decompose, i.e. the ease of decomposition decreases down the group, hence they will have a higher decomposition temperature. The thermal stability of the Group II compounds is affected by the polarising power of the cations and the polarisability of the large anions. (a) Polarising Power of the Cations

Polarising power refers to the ability of a cation to distort the electron cloud of another ion, atom or molecule. In general, the higher the charge density, the greater the polarising power of the cation. The small and highly charged M2+ cation in these Group II compounds polarises the anion’s larger electron cloud, weakening the covalent bonds within the anion. Hence these compounds decompose on heating as the weakened covalent bonds within the anion are easily broken.


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For Group II,

Cationic radius increases down the group while charge remains the same

Therefore charge density of cation decreases down the group

Polarising power of the cation decreases down the group and is less able to distort the

electron cloud of the anion, weakening the covalent bonds within the anion

Covalent bonds within the anion are less likely to be broken down the group

The ease of decomposition decreases (or thermal stability increases) down the group so

higher temperature is required to decompose the compound

(b) Polarisability of the Anions CO32, NO3

& OH

Polarisability refers to the ease with which an anion, atom or molecule’s electron cloud can be

distorted by another ion.

Large anions are susceptible to polarisation of their electron cloud.

Only polyatomic anions are susceptible to decomposition as monoatomic anions (e.g. O2,

Cl ) cannot be broken down further.

There are other factors that may contribute to ease of decomposition. One factor is the way

the ions are packed in the lattice structure.

Polarisation of the Carbonate ion

Now imagine what happens when this ion is placed next to a positive ion. The positive ion attracts the electron cloud of the carbonate ion towards itself. The carbonate ion becomes polarised.

-O

C

O-

O

O

N

O-

O

O H

carbonates nitrates hydroxides

2−


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Products of Decomposition

Comparison with Group I metals

The Group I metal nitrates also undergo thermal decomposition but not to the same extent as the

corresponding Group II compounds. The reason for the difference is the lower charge density of the

Group I metal cations which are unable to polarise the nitrate’s electron cloud to the same extent

as the Group II metal cations.

The Group I nitrates decompose on heating to give the corresponding nitrite and oxygen.

MNO3(s) MNO2(s) + ½O2(g)

The Group I carbonates are resistant to decomposition except for Li2CO3 (why?). The charge density

of Li+ is high enough to distort the carbonate’s electron cloud so that the compound decomposes on

heating.

Li2CO3(s) Li2O(s) + CO2(g)

The other Group I metal cations have low charge density due to their larger radii and so are unable

to distort the carbonate’s electron cloud. Therefore the carbonates after lithium do not decompose

on heating.


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1 M is a Group II metal which can undergo reaction via two routes:

Which set below contains 3 different compounds?

A P Q U B P R T C Q S U D Q S T

2 Which of the following is the correct trend for Group II elements from Mg to Ba? A The oxide becomes less basic. B The elements become more electronegative. C The calcium cation has the highest charge density. D The decomposition temperature of the nitrates increases.

3 Which of the following statements concerning Group II elements, magnesium, calcium and barium are correct?

1 Their reactivity increases with increasing atomic mass. 2 The only oxidation state exhibited in their stable compounds is +2. 3 On strong heating, the nitrates liberate only oxygen gas.


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3.5 APPENDIX A - ANOMALOUS BEHAVIOUR OF BERYLLIUM

Beryllium shows quite different chemical behaviour when compared to the rest of the group.

In fact, it is closer to aluminium in its chemical properties (diagonal relationship).

Beryllium compounds tend to be covalent

Beryllium does not react with water in any way

Beryllium oxide does not react with water

Beryllium oxide and hydroxide are amphoteric

Beryllium forms complexes which the rest of group does not

Beryllium salts are appreciably acidic but the rest of the group is not

Note: Beryllium compounds are exceedingly poisonous

A) Covalency of Be compounds

Beryllium compounds tend to be covalent due to extremely high charge density of the cation.

The high polarising power of Be2+ causes significant distortion of the anion’s electron cloud leading to overlap of the cation and anion orbitals – resulting in covalent compounds or with high degree of covalency.

Compounds which carry large anions like BeCl2 are completely covalent.

Even compounds with small highly charged anions like BeO have significant covalent character.

B) Non-reaction of Be & BeO with water

Beryllium forms an impervious oxide layer on the surface of the metal on exposure to air.

This BeO layer is resistant to hydration by water molecules and hence the metal is protected

from reacting with water. Beryllium does not react with water even when heated red-hot

which is testament to the extremely high stability of the BeO layer.

Hence neither Be nor BeO reacts with water.

Note that MgO also forms an impervious oxide layer which also accounts for the lack of

reactivity of MgO with water.

C) Amphoteric nature of BeO and Be(OH)2

Ionic oxides are basic while covalent oxides are acidic. The amphoteric oxides possess both

ionic and covalent character.

The amphoteric nature of BeO and Be(OH)2 can be traced back to the partial covalent

character of the compounds.

Basic BeO(s) + H2SO4(aq) BeSO4(aq) + H2O(l)

Acidic BeO(s) + 2NaOH(aq) + H2O(l) Na2[Be(OH)4](aq)


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D) Formation of Be complexes

Beryllium forms complexes which the rest of the Group II metals do not:

Be(OH)2(aq) + 2OH (aq) [Be(OH)4]2 (aq)

BeF2(s) + 2F (aq) [BeF4]2 (aq)

The reason for beryllium forming complexes lies in its electronic configuration. In binary compounds (e.g. BeCl2), beryllium has only 4 valence electrons and is only 2-coordinate. There is a tendency to achieve maximum coordination (and hence valence octet) by acting as a Lewis acid and form a complex.

Beryllium has a maximum coordination number of 4 as it has only 4 available valence orbitals – one 2s & three 2p orbitals. There are no 2d orbitals and hence six-coordination is impossible.

This explains why in gaseous phase, BeCl2 is a linear discrete molecule but in the solid state, BeCl2 consists of linear polymers [compare it with AlCl3 dimer].

E) Acidity of Be salts in water

The beryllium cation forms a tetrahedral hydrated complex in water:

H2O

H2O

Be

OH2

OH2

Hydrolysis of the hydrated complex occurs readily in water due to the very high charge density of the beryllium cation:

[Be(H2O)4]2+(aq) + H2O(l) [Be(H2O)3(OH)]+(aq) + H3O

+(aq)

Cl Be Cl

180o

Gas phase

Solid phase

2+


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3.6 APPENDIX B - SOLUBILITY OF GROUP II COMPOUNDS

The solubilities of the Group II compounds are varied and depend on solution enthalpies which in turn depend on both the lattice energies of the compounds and hydration energies of the ions (refer to Chemical Energetics notes).

Hsolution = − LE + Hhydration

Generally soluble salts tend to have an exothermic Hsolution but this is not always the case (for example the process of dissolving sodium nitrate is endothermic). In order to know

whether a salt is soluble, we also have to look at Ssolution but we will not discuss it here.

Table 3.13 – Solubility of Group II sulfates, carbonates and hydroxides

Element Solubility (moles per 100 g of water)

Sulfates Carbonates Hydroxides

Be 2.4 × 101 Insoluble 8.0 × 107

Mg 2.2 × 101 1.3 × 104 1.6 × 105

Ca 1.5 × 103 1.3 × 105 2.5 × 103

Sr 7.1 × 104 7.0 × 106 3.4 × 103

Ba 1.1 × 106 9.0 × 106 4.1 × 102

TREND Solubility decreases

down the group

Solubility generally decreases down the

group

Solubility increases down the group

A) Lattice energy

The lattice energy measures how strongly the ions are attracted to one another. In general, ions with high charge density form salts with very exothermic lattice energies.

1LE

r r

The magnitude of the lattice energy is dependent on the sum of the ionic radii. Hence changes in radii affect the lattice energy less significantly than the hydration energy.

B) Hydration energy

The hydration energy measures how strongly an ion forms ion-dipole interactions with water. In general, ions with high charge density are very soluble in water as the hydration energies are very exothermic

hydration

1 1H

r r

The magnitude of the total hydration energy is dependent on individual ionic radii. Hence changes in radii affect the hydration energy more significantly than the lattice energy.


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Solubility of Group II Sulfates and Carbonates

The decreasing trend in solubility of the sulfates and carbonates down the group is due to the

increasing radii of the cations down the group.

Both the lattice energy and hydration energy decreases down the group as cationic radii increases.

But the hydration energy decreases faster than the lattice energy. This is because the change in

cationic radii affects the hydration energy more significantly than the lattice energy. Hence Hsolution

is increasingly positive down the group.

Solubility of Group II Hydroxides

The increasing trend in solubility of the hydroxides down the group is opposite to that for the

sulfates and carbonates.

Both the lattice energy and hydration energy decreases down the group as cation radii increases. But

the hydration energy decreases more slowly than the lattice energy now. This is because of the small

OH− size which dominates the value of the hydration energy. Therefore changes in the cationic radii

do not affect the hydration energy as much as the lattice energy. Hence Hsolution is increasingly

negative down the group.

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22.Group II Lecture Notes