CLASS- 11 (The s-Block Elements)

                          The s-Block Elements

 

 

The s-block elements

The s-block elements are the metals in Group IA and IIA of the Periodic Table.

Group IA - alkali metals

Group IIA - alkaline earth metals

They are called s-block elements because electron(s) in the outermost shell occupies the s-orbitals.

 

Electronic configuration

 

Group IA         

Lithium (Li) 1s2 2s1

Sodium (Na) 3s1

Potassium (K) 4s1

Rubidium (Ru) 5s1

Caesium (Cs) 6s1

Francium (Fr) 7s1

 

Group IIA

Beryllium (Be) 1s22s2

Magnesium (Mg) 3s2

Calcium (Ca) 4s2

Strontium (St) 5s2

Barium (Ba) 6s2

Radium (Ra) 7s2

 

Group IA elements

 

1. Group IA consists of lithium, sodium, potassium, rubidium, cesium, and

francium. Francium is radioactive.

2. They are silvery in appearance.

3. They have an outer electronic configuration of ns1. As the outermost shell

electron is effectively screened from nuclear attraction, the electrostatic

attraction between the delocalized electron cloud and the metal ions is weak.

9 The s-block Elements – 2 That results in their weak metallic bond.

4. Because of the weak metallic bond, they are soft metals.

5. Their metallic crystals have body-centered cubic structures and their densities are comparatively low. Lithium, sodium, and potassium are all less dense than water.

 

Group IIA elements

 

1. Group IIA elements include beryllium, magnesium, calcium, strontium,

barium and radium.

2. They are grey metals.

3. They have an outer electronic configuration of ns2. Compared with Group IA

metals, Group IIA metals have smaller atomic sizes as their effective nuclear

charges are greater. Thus, they have stronger metallic bonding and are harder

than the corresponding Group IA elements.

4. Because their atoms are significantly smaller than those of Group IA elements, Group IIA metals have higher densities.

5. Lattice structures:

Beryllium and magnesium – hexagonal close-packed structure

Calcium and strontium – face-centered cubic structure

Barium – body-centered cubic structure

 

Characteristic properties of the s-block elements

 

Electronegativity

Both Group IA and Group IIA elements are all highly electropositive metals, i.e.

they possess small electronegativity values. This is because the outermost shell

electrons are effectively screened from the nuclear attraction. The electronegativity decreases in both groups.

 

Electronegativity values of Groups IA and IIA elements

 

Group IA Electronegativity

Li      1.0

Na     0.9

K       0.8

Rb     0.8

Cs      0.7

Fr      0.7

 

Group IIA Electronegativity

 

Be      1.5

Mg    1.2

Ca      1.0

Sr      1.0

Ba     0.9

Ra     0.9

 

Metallic character

 

The metallic character refers to the tendency of a metal to lose electrons to form cations. As the removal of an electron from a larger atom is easier, the metallic character increases down both Groups IA and IIA.

 

Oxides

 

The s-block elements, when freshly cut, show silvery-white luster. Except

beryllium and magnesium which have a protective oxide layer, Group IA and IIA metals tarnish quickly and form an oxide layer when exposed to air, owing to The s-block Elements –atmospheric oxidation. To prevent oxidation, all Group IA metals, strontium, and barium of Group IIA metals are stored under paraffin oil to prevent contact with the air. Ca is stored in air-tight containers Group IA metals form one or more of the three types of oxides, namely, normal oxides, peroxides, and superoxides.

 

4Li(s) + O2(g) → 2Li2O(s) lithium oxide, a normal oxide (major product)

4Na(s) + O2(g) → 2Na2O(s)

2Na(s) + O2(g) → Na2O2(s) sodium peroxide (major product)

4K(s) + O2(g) → 2K2O(s)

2K(s) + O2(g) → K2O2(s)

K(s) + O2(g) → KO2(s) potassium superoxide (major product)

2Mg(s) + O2(g) → 2MgO(s)

2Ca(s) + O2(g) → 2CaO(s)

 

*Mg also forms nitride when heated in air: 3Mg(s) + N2(g) → Mg3N2(s)

Except for BeO which is amphoteric, all Group IA and IIA oxides are basic.

 

 

Hydroxides

 

1. Most s-block metals react with cold water or sometimes with steam to form

oxides or hydroxides. The reactivity increases down the group.

 

2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)

 

Ca(s) + 2H2O(l) → Ca(OH)2(aq) + H2(g)

 

Mg(s) + 2H2O(l) → Mg(OH)2(aq) + H2(g) …… slow reaction

Mg(s) + H2O(l) → MgO(s) + H2(g)steam

 

 

2. Hydroxides are also formed when Group IA and IIA metal oxides react with

water. BeO and MgO are insoluble in water.

 

Normal oxides:

 

Na2O(s) + H2O(l) → 2NaOH(aq)

 

K2O(s) + H2O(l) → 2KOH(aq)

 

CaO(s) + H2O(l) → Ca(OH)2(aq)

 

BaO2(s) + H2O(l) → Ba(OH)2(aq)

 

Peroxides:

 

Na2O2(S) + 2H2O(1) → 2NaOH(aq) + H2O2(aq)

 

Superoxides:

 

2KO2(s) + 2H2O(1) → 2KOH(aq) + H2O(aq) + O2(g)

 

Both oxides and hydroxides of s-block elements are basic in nature, that is, they react with acid to form the salt.

 

CaO(s) + 2H+(aq) → Ca2+(aq) + H2O(l)

 

Mg(OH)2(aq) + 2H+(aq) → Mg2+(aq) + 2H2O(l)

 

 

Bonding and oxidation state in compounds

 

Group IA elements

 

Group IA elements are electropositive metals. Because of their relatively large

atomic size and effective screening of the single outermost shell electron, they

have low first ionization enthalpies. However, their second ionization

enthalpies are extremely high, owing to the stable octet electronic

configurations of the M+ ions. So, they form predominantly ionic compounds

such as oxides (M2O), hydrides (MH), and chlorides (XCl) with non-metals by

losing their single outermost shell electrons and having a fixed oxidation state

of +1.

 

 

Group IIA elements

 

Group IIA elements are also electropositive metals. They have relatively low

first and second ionization enthalpies. However, their third ionization

enthalpies are extremely high, owing to the stable octet electronic

configurations of the M2+ ions. They form predominantly ionic compounds

such as oxides (MO), hydrides (MH2), and chlorides (XCl2) with non-metals by

losing their two outermost shell electrons. They have a fixed oxidation state of

+2. However, beryllium and magnesium, have smaller atomic sizes and

greater electronegativities, and their compounds possess some degree of covalent

character.

 

Tendency to form complexes (coordination compound)

 

A complex is an ion or a compound in which a central metal atom or ion is

attached to a group of surrounding molecules or ions by dative (coordinate

covalent) bonds.

 

        Li

        Na

        K

        Rb

        Cs

all soluble in water

 

        MOH

base strength increasing

 

        Be

        Mg

        Ca

        Sr

        Ba

insoluble  solubility increase

 

M(OH)2 amphoteric  base strength increase

 

 

Reasons for the low tendency of s-block elements to form complexes :

 

1. Relatively small charges and large in size of their ions.

2. They have completely filled inner orbitals. However, owing to its higher charge/radius ratio and higher tendency to form covalent bonds, beryllium forms the most stable complexes. The metal or its hydroxide can dissolve in strong bases to give the beryllate ion, [Be(OH)4]2

 

Flame color

On heating, the outermost shell electrons are excited to higher energy levels.

When these electrons return to their ground states, energy is released in the form of light. Since the amount of energy is quantized, radiation of certain wavelengths is emitted and the flame color is a characteristic property of each element. This forms the basis of flame tests to identify the metals.

The ease of exciting their outermost shell electrons and emitting radiation in the visible light region enables the carrying out of flame tests on Group IA and IIA elements.

 

Characteristic flame colors of Groups IA and IIA elements

 

Group I A Elements   Flame Color      Group IIA Element    Flame Color

 

          Li                          crimson                                   Be               brick red

          Na                        golden yellow                          Mg              crimson

          K                          lilac                                         Ca               apple green

          Rb                        violet                                       Sr

          Cs                         Blue                                          Ba

 

 

Variation in physical properties of the s-block elements

 

Atomic radius

 

Variation of atomic radii of;

 

Group I and      II elements

          Li                          Be

          Na                        Mg

          K                          Ca

          Rb                        Sr

          Cs                         Ba

          Fr                         Ra

1. Across a period, the atomic radii decrease as the effective nuclear charges

increase. Therefore, Group I elements are larger than Group II elements in the

same period.

2. Down both groups, the atomic radii increase, owing to the increase in the

a number of completely filled inner shells.

Ionization enthalpy

 

1. The first ionization enthalpies of the s-block elements are low. It is because the

atomic sizes of the s-block elements are large, their valence electrons are

therefore at a greater distance from the nucleus compared with atoms of other

groups. As a result, these valence electrons experience less nuclear attraction.

 

For the same period, the first ionization enthalpies of group IIA elements is

greater than that of Group I elements.

 

Reasons: For Group IIA elements, the effective nuclear charge is greater and

the atomic size is smaller.

 

Going down the group, the first ionization enthalpy decreases.

 

Reasons: As we down the group, there are more filled inner electron shells,

and thus the screening effect will be greater. Also, the distance between the valence electron and the nucleus increases.

 

Group IA: The second ionization enthalpies of Group IA elements are

much higher than the first ionization enthalpies, i.e. 1st I.E. <<2nd I.E.

 

Reasons: It is because the second electron has to be removed from the

positively charged ion, so the amount of energy required is

much larger. Also, the second electron is removed from the

inner shell which is closer to the nucleus.

 

Group IIA: 1st I.E. < 2nd I.E << 3rd I.E

 

Reasons: The two outermost shell electrons are screened by inner shell

electrons, so they can be removed easily. The third ionization enthalpies of Group II elements are much higher than the first and second ionization enthalpies. It is because a much larger amount of energy is required to remove

an electron from an ion that carries two positive charges already. Also, the third electron is removed from the inner shell which is closer to the nucleus.

 

 

 

The 1st I.E. of Group IIA elements is higher than that of Group IA elements.

 

Reasons: The size of Group IIA elements is smaller than those of Group I A

elements, so the distance between the outermost shell electrons and the nuclei is shorter for Group IIA elements. As a result, it is more difficult to remove the 1st electron from Group IIA elements than from Group IA elements.

 

Melting point

 

Generally speaking, as the atomic sizes of Groups IA and IIA metals are large, their outermost shell electrons are weakly held by their nuclei. Therefore, the strength of the metallic bond in these metals is weak and requires a relatively small amount of energy to break. As a result, the s-block elements have low melting points.

 

Down the group: Melting point decreases

 

Reasons: For elements in the same group, the number of delocalized (valence)

electron per atom remains the same. However, as going down the group, the atomic sizes increase. Hence, the strengths of the metallic bond (attractions between the delocalized electron and the metal ion) decrease.

 

Melting points of Group IIA elements are greater than those of Group IA

Elements

 

Reasons: • The atoms of Group IIA metals have two delocalized electrons and

their sizes are smaller than those of Group IA metals. Group IIA metals have a closed-packed structure while Group IA metals have open structures (body-centered cubic).

The pattern for the melting points of Group IIA metals is irregular. This is

because metals have different metallic crystal structures. Calcium has an

exceptionally high melting point because of its ability to use its 3d vacant orbitals to strengthen the metallic bonding.

 

Hydration enthalpy

 

Hydration enthalpy (ΔHhyd) is the amount of energy released when one mole of

aqueous ions are formed from their gaseous ions.

 

i.e. energy released for M (g) + aq → M+(aq)

 

Hydration enthalpy is results from the attraction between ions and water

molecules. The amount of energy released is therefore dependent on the

charge/radius ratio of the ion. A larger amount of energy is given out if the charge on the ion is greater and if the size is smaller which causes an increase in charge density.

 

Group IIA > Group IA

Group IIA metal ions carry two positive charges and have smaller ionic sizes

compared with the corresponding Group IA metal ions, hence their ions exert

a greater attractive force on water molecules. Therefore, Group IIA metal ions

have greater hydration enthalpies than Group IA metal ions.

 

Down the group, enthalpy of hydration decreases. .

Although the charge on the ions remains unchanged, the ionic sizes increase

down the group. Therefore, the attractive force between the ions and the water

molecules becomes smaller. Hence, hydration enthalpy decreases down the

group.

 

Variation in chemical properties of the s-block elements

Metals react by losing electrons and therefore their reactivities depend on the

magnitude of their ionization enthalpies. Group IA metals have lower ionization enthalpies than Group IIA metals, they are therefore more electropositive and more reactive than Group IIA metals.

 

Reaction with hydrogen

1. All Group IA metals react with hydrogen at about 400oC to form hydrides

which are white crystalline solids. This is called the direct synthesis of metal

hydrides and is their usual method of preparation.

For example,

2Na(s) + H2(g) → 2NaH(s)

 

The most important alkali metal hydride is lithium hydride, which on

treatment with aluminum chloride in a dry ethereal solution yields one of the

 

Be2+

Mg2+

Ca2+

Sr2+

Ba2+

Li+

Na+

K+

Rb+

 

Variation of hydration enthalpies of Cs+ Group IA and IIA metal ions.

 

The most useful reducing agents in organic chemistry, lithium aluminum hydride,

2. In the electrolysis of molten hydrides, hydride ions are discharged at the

The anode forms hydrogen gas.

 

H−(l) → H2(g) + 2e−

Group IIA metals also react with hydrogen to form hydrides, but at a higher

temperature (about 600oC to 700oC).

 

For example,

 

Ca(s) + H2(g) → CaH2(s)

 

Group IIA hydrides are ionic except beryllium and magnesium hydrides which

are predominantly covalent with bridged polymer-type structures, representing

a transition between covalent and ionic hydrides. Beryllium hydride has the

following structure :

 

Reaction with oxygen

 

1.     With the exception of Li, all alkali metals react with O2 to produce more than one oxide, peroxide, and superoxide. All three types of oxides are ionic:

 

4Li(s) + O2(g) → 2Li2O(s) lithium oxide

4Na(s) + O2(g) → 2Na2O(s)

2Na(s) + O2(g) → Na2O2(s) sodium peroxide (major product)

4K(s) + O2(g) → 2K2O(s)

2K(s) + O2(g) → K2O2(s)

K(s) + O2(g) → KO2(s) potassium superoxide (major product)

 

2. Group IIA metals react with oxygen to form metal oxide.

Beryllium and magnesium are relatively unreactive toward oxygen at room

temperature. However, both of them burn with a brilliant white fume when

heated with a Bunsen flame. Strontium and barium are so reactive that they are stored under paraffin oil to prevent contact with air. Strontium and barium also form peroxides.

For example,

2Be(s) + O2(g) → 2BeO(s)

2Mg(s) + O2(g) → 2MgO(s)

2Ca(s) + O2(g) → 2CaO(s)

2Ba(s) + O2(g) → 2BaO(s)

Ba(s) + O2(g) → BaO2(s) barium peroxide

 

 

3. Stability of peroxides increases down the group because peroxides are large

anions and can be stabilized by less polarizing and large cations. If the charge

density on the cation is large such as in the case of Mg2+, their peroxide tends to be unstable and tends to decompose to give oxide and oxygen.

Reaction with chlorine

All s-block metals react with chlorine directly to form chlorides. The reactivity of the metals with chlorine increases down the groups. For example, beryllium reacts with chlorine only when heated but barium burns immediately when put in contact with chlorine.

 

2. Group IA metals react rapidly with dry chlorine to form colorless ionic

chlorides with high melting and boiling points.

For example,

2Na(s) + C12(g) → 2NaCl(s)

2K(s) + C12(g) → 2KCl(s)

 

Group IIA metals react with chlorine to give ionic chloride salts.

For example,

Mg(s) + C12(g) → MgC12(s)

Ca(s) + C12(g) → CaC12(s)

 

Beryllium is the most unreactive metal in Group IIA. A high temperature is

required for the synthesis of beryllium chloride, which is covalent in nature

owing to the high charge and small size of the beryllium ion.

 

Be(s) + Cl2(g) → BeCl2(s)

 

Reaction with water

 

1. With the exception of Be, all Group IA and IIA metals react readily with cold

water or steam to form hydroxides and hydrogen.

 

2Li(s) + 2H2O(l) → 2LiOH(aq) + H2(g) ………….. gentle reaction

2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g) …………. violent

2K(s) + 2H2O(l) → 2KOH(aq) + H2(g) …………. explosive

Mg(s) + H2O(g) → MgO(s) + H2(g) …………. vigorous reaction

Mg(s) + 2H2O(l) → Mg(OH)2(aq) + H2(g) ……….. very slow

Ca(s) + 2H2O(l) → Ca(OH)2(aq) + H2(g) ……….. moderate

Sr(s) + 2H2O(l) → Sr(OH)2(aq) + H2(g) ………. moderate

 

2. The reactivity of metals with water increases down the group as the relative

ease of donating the outermost shell electron increases with increasing atomic

size. For example, lithium reacts with water with vigorous bubbling as

hydrogen is released, sodium reacts violently with water and dashes about on

the water's surface and potassium reacts almost explosively with water, the

hydrogen produced bursting instantly into a lilac flame. Rubidium and cesium react explosively with water.

 

Variation in properties of the compounds of the s-block elements

 

(i) The oxides

(a) Reaction with water

 

1. All Group IA oxides are basic in nature and react vigorously with water

to form hydroxides. The basicity of the oxides increases down the group.

 

Normal oxides:

Na2O(s) + H2O(l) → 2NaOH(aq) (strongly alkaline)

K2O(s) + H2O(l) → 2KOH(aq) (strongly alkaline)

 

Peroxides:

Na2O2(S) + 2H2O(1) → 2NaOH(aq) + H2O2(aq)

 

Superoxides:

2KO2(s) + 2H2O(1) → 2KOH(aq) + H2O2(aq) + O2(g)

 

2. Group IIA metals, except beryllium oxide and magnesium oxide, all

their oxides react with water to form weakly alkaline solutions. For

example,

CaO(s) + H2O(l) → Ca(OH)2(aq) (weakly alkaline)

SrO(s) + H2O(l) → Sr(OH)2(aq) (weakly alkaline)

 

3. Group IIA metal oxides are generally less basic than Group IA metal

oxides and forms a weakly alkaline solutions with water.

 

 Reaction with acids

 

All three types of metal oxides react with acids.

Normal oxides neutralize the acids to form salts and water.

 

Na2O(s) + 2HCl(aq) → 2NaCl(aq) + H2O(l)

CaO(s) + 2HCl(aq) → CaCl2(aq) + H2O(l)

 

Peroxides react with acids to form salts and hydrogen peroxide.

Na2O2(s) + 2HCl(aq) → 2NaCl(aq) + H2O2(l)

BaO2(s) + 2HCl(aq) → BaCl2(aq) + H2O2(l)

 

Superoxides react with acids to form salts, hydrogen peroxide, and oxygen.

2KO2(s) + 2HCl(aq) → 2KCl(aq) + H2O2(l) + O2(g)

 

Reaction with alkalis

 

In general, there is no reaction between the oxides of the s-block elements

with alkalis. However, beryllium oxide, which is amphoteric, reacts with

sodium hydroxide to give sodium beryllate.

 

BeO(s) + 2NaOH(aq) + H2O(1) → Na2Be(OH)4(aq)

 

 

The hydrides

 

All s-block metal hydrides are hydrolyzed by water to form hydroxides and

hydrogen. The stability of the hydrides decreases down both groups, suggesting

that the reactivity of the metal hydrides increases down both groups.

 

NaH(s) + H2O(1) → NaOH(aq) + H2(g) vigorous

KH(s) + H2O(1) →KOH(aq) + H2(g) very vigorous

CaH2(s)+ 2H2O(1) → Ca(OH)2(aq) + H2(g) fairly vigorous

 

 

The chlorides

 

1. Group IA chlorides are basically ionic in nature. They dissolve in water to

form metal cations and chloride ions. The ions are not hydrolyzed by water

and the resulting solutions are neutral. They are insoluble in organic solvents.

 

2. Group IIA chlorides show some degree of covalent character as the

electronegativity of Group IIA metals is higher than that of Group IA metals.

The degree of ionic character of Group IIA chlorides increases down the group.

Beryllium chloride is basically covalent and is hydrolyzed by water to form an

acidic solution.

BeCl2(aq) + 2H2O → Be(OH)2(s) + 2HCl(aq)

 

Magnesium chloride is intermediate between ionic and covalent in character. It dissolves readily, with very slight hydrolysis. Other Group IIA chlorides just

dissolve in water without hydrolysis.

 

Relative thermal stability of the carbonates and hydroxides

 

The thermal stability of an ionic compound M+X−(s) depends on:

1. Charge of the ions - The higher the charge of the ions, the stronger is the

the attraction between the ions, the more stable the ionic compound.

 

2. Size of the ions - The smaller the ions, the closer the distance between

the ions, the more stable the compound.

Hence, an ionic compound becomes more stable as ionic charge increases or

ionic radius decreases.

 

For compounds with large anions, thermal stability is affected by the

polarizing power of the cations. If the cation has a greater polarizing power

(smaller size and higher charge, e.g. Li+ and Mg2+), it will distort (attract)

the electron cloud of the neighboring large anion to the cation. This weakens

the bond in the anion (i.e. cause the compound to be less stable), and leads to

the decomposition to a smaller oxide ion.

M2+

 

O2− ions are more stable than CO3 2− and OH− because O2− is smaller and hence the charge density on the ion is higher. Oxide ions are therefore more strongly

attracted to the cations in the compound. Also, the internuclear distance

between oxide ions and cations is smaller than those between CO3 2− ions or

OH− ions and the cations. Therefore, most carbonates and hydroxides decompose readily on heating to oxides.

 

Smaller anions formed by thermal decomposition are less easily polarized.

 

Down each group, as the sizes of cations increase, the polarizing power of

cations decreases and compounds with large cations become more stable.

Hence, the thermal stability of carbonates and hydroxides of both Groups IA

and IIA metals increase down the group.

Effect of the sizes of cations on the thermal stability of compounds

 

The carbonates

 

Group IA carbonates

 

All Group IA carbonates (except lithium carbonate) are stable at a temperature of around 800oC. Lithium carbonate decomposes around 700oC forming lithium oxide and carbon dioxide. The relative instability of lithium carbonate is due to the small size Li+, and hence a very large charge/radius ratio of Li+ ion.

 

LiCO3(s) → LiO(s) + CO2(g)

 

Group IIA carbonates

 

Group IIA metal ions are much smaller than Group IA metal ions and carry a

higher charge. Therefore, their polarizing power is greater. Electron clouds of

carbonate anions are more easily distorted by Group IIA metal ions than by

Group IA metal ions. Carbonates of Group IIA metals are therefore less stable

to heat and decompose to metal oxide and carbon dioxide.

Equation Decomposition temp /oC

 

BeCO3(s) → BeO(s) + CO2 (g) 100

MgCO3(s) → MgO(s) + CO2 (g) 540

CaCO3(s) → CaO(s) + CO2 (g) 900

SrCO3(s) → SrO(s) + CO2 (g) 1290

BaCO3(s) → BaO(s) + CO2 (g) 1360

 

 

The hydroxides

 

Group IA hydroxides

 

All Group IA hydroxides, except lithium hydroxide, are stable when heated

with a Bunsen burner. Lithium hydroxide is the least stable in the group

because the extremely small Li+ ion has very high polarizing power and

distorts the electron cloud of the hydroxide ion. This weakens the O–H

bond and finally decomposes the hydroxide.

 

2LiOH(s) → Li2O(s) + H2O(g)

 

 

Group IIA hydroxides

 

Group IIA metal ions are much smaller than Group IA metal ions and carry

a higher charge. Consequently, their polarizing power is greater. The electron

cloud of the hydroxide anions is more easily distorted by Group IIA metal

ions than by Group IA metal ions. This weakens the O–H bond and leads to

the decomposition of the hydroxides. Hydroxides of Group IIA metals are

therefore less stable to heat than those of Group IA metals. The thermal stability of Group IIA hydroxides increases down the group because the polarizing power of cations decreases with increasing size, as can be seen from the enthalpy change of their decomposition.

 

Be(OH)2(s) → BeO(s) + H2O ΔH = +54 kJ mol−1

Mg(OH)2(s) → MgO(s) + H2O ΔH = +81 kJ mol−1

Ca(OH)2(s) → CaO(s) + H2O ΔH = +109 kJ mol−1

Sr(OH)2(s) → SrO(s) + H2O ΔH = +127 kJ mol−1

Ca(OH)2(s) → CaO(s) + H2O ΔH = +146 kJ mol−1

 

Relative solubility of the sulphates(VI) and hydroxides

 

When an ionic crystal dissolves in water,

1. the ions become separated and energy is required to break up the lattice

2. cations and anions are solvated and energy is released in the process.

 

The energy change involved when one mole of a substance is dissolved

completely in a solution is called the enthalpy change of solution (ΔHsoln).

 

For dissolving a salt MX in water, different enthalpy changes are related as

shown by the following enthalpy cycle.

 

ΔHsoln = ΔHhyd – ΔHlattice

 

If the magnitude of ΔHhyd > ΔHlattice, then ΔHsoln is negative (exothermic) and the compound will be soluble. The compound will also be soluble in water if

ΔHsoln has a small positive value.

 

The solubility depends mainly on the size of the ions.

 

1. The larger the cation, the smaller the hydration enthalpy.

2. Lattice enthalpy depends on the sum of the ionic radii, i.e. (r+ + r−). The

greater the value of (r+ + r−) , the smaller will the lattice enthalpy be.

 

Group IA metal ions are comparatively large in size and carry small charges,

so their effective nuclear charges are not very high. Their lattice enthalpies and hydration enthalpies have similar values. Hence, their enthalpy changes of

the solution, ΔHsoln, are usually negative or slightly positive in values.

 

Group IIA metal ions have higher charges and smaller sizes, so, ΔHlattice, of

their compounds are larger in magnitude than those of Group IA compounds.

Consequently, their ΔHsoln, are less negative or more positive.

 

 

The sulphates(VI)

 

All Group IA sulfates (V1) are soluble in water.

Group IIA sulfates (VI) are less soluble.

 

The solubility of sulfates (VI) decreases down the group.

Reasons: • SO4 2− is a large anion (much larger than the cations of Group

IIA metals), therefore on going down the group, ΔHlattice of sulfate does not vary a lot because the change in the size of the cations does not cause a significant change in the sum of ionic radii.

• ΔHhyd decreases rapidly as the size of the cation increases. As a

the result, ΔHsoln becomes less negative and the solubility of Group

ΔHlattice α 1

r+ + r−

 

IIA sulfates (VI) decrease down the group.

 

For example;

 beryllium sulfate (VI) is very soluble whereas barium sulfate(VI) is insoluble in water.

 

The hydroxides

 

For hydroxides, as the size of a cation increases, the change in (r+ + r−) is very

significant because the size of OH− is not large. Hence, the decrease in ΔHlattice

is quite significant.

 

All Group IA hydroxides are soluble in water. The solubility generally

increases down the group because of the decrease in lattice enthalpy.

 

For Group IIA hydroxides

 

• As the size of a cation increases, the change in (r+ + r−) is very significant

because the size of OH− is not large. Hence, the decrease in ΔHlattice is quite

significant. However, the change in ΔHhyd is relatively insignificant. Therefore, the decrease in ΔHhyd is smaller than the decrease in ΔHlattice on descending the group.

 

• As a result, the ΔHsoln becomes more negative, and the solubility of Group

IIA hydroxides increase down the group.

For example,

beryllium hydroxide is insoluble but barium hydroxide (solubility = 1.5 × 10−3 mol per 100 g of water) is soluble in water.

 

 

 

 

 

Uses of the compounds of the s-block elements

 

Sodium Carbonate

 

1. Sodium carbonate is used in the manufacture of glass. Soda glass is made by

fusing the carbonates with silica (from sand) at 1500oC and is thus a mixture

of sodium silicate and calcium silicate.

CaCO3(s) + SiO2(s) → CaSiO3(s) + CO2 (g)

Na2CO3(s) + SiO2(s) → Na2SiO3(s) + CO2 (g)

 

2. Used in sewage treatment and water softening as carbonate ions can precipitate

Mg2+ (aq) and Ca2+(aq) in hard water.

Mg2+(aq) + CO3 2−(aq) → MgCO3(s)

Ca2+(aq) + CO3 2−(aq) → CaCO3(s)

 

Sodium hydrogen carbonate

 

Sodium hydrogen carbonate is commonly found in baking powder which also

contains a solid acid. On adding water, the acid reacts with sodium

hydrogen carbonate to produce carbon dioxide. Besides, sodium

hydrogen carbonate decomposes at high temperatures giving off carbon dioxide.

The carbon dioxide gas produced makes the cake rise and become spongy.

 

HCO3− (aq) + H+(aq) → H2O(1) + CO2 (g)

 

from solid acid when dissolved in water

 

2NaHCO3(s) → Na2CO3(s) + CO2 (g) + H2O(1)

 

Sodium hydrogen carbonate is also used extensively in making soft drinks.

 

Sodium hydroxide

 

Sodium hydroxide is used to manufacture soaps, detergents, dyes, paper, and

drugs. Soaps are made by saponification in which fats and oils are hydrolyzed by sodium hydroxide.

 

Magnesium hydroxide

Magnesium hydroxide is a weak base and is used as an active ingredient in many antacids to neutralize the excess hydrochloric acid produced in the stomachs of people suffering from gastric pain.

Mg(OH)2(s) + 2HCl(aq) → MgC12(aq) + 2H2O(1)

Because magnesium hydroxide is relatively insoluble and would remain in the

patient’s stomach to act on any acids produced at a later time.

 

Calcium oxide and hydroxide

 

Calcium carbonate decomposes to quicklime (calcium oxide) when heated. The

addition of water to quicklime, CaO, produces slaked lime, Ca(OH)2.

CaCO3(s) → CaO(s)+ CO2(g)

CaO(s) + H2O(l) → Ca(OH)2(s)

 

Slaked lime is used to neutralize the acids in industrial effluents.

 

Ca(OH)2(s) + 2H+(aq) → Ca2+(aq) + 2H2O(l)

 

Strontium compounds

Used in fireworks because they give a persistent and intense red flame while

burning.