The S-block Elements: Overview, Questions, Preparation

A 10.1

Common physical features of alkali metals are:

• Physical appearance: All the alkali metals are silvery white, soft and light metals.
• Density: Because of the large size, these elements have low density which increases down the group except for potassium which is lighter than sodium (most likely due to an unexpected increase in the atomic size.).
• The melting and boiling points of the alkali metals are low indicating weak metallic bonding due to the presence of only a single valence electron in them.
• Atomic volume:The atomic volume, atomic and ionic radii rise as the group number reduces from Li to Cs.
• Melting and boiling points:The weak crystal lattice bonding causes low melting and boiling points of all alkali metals.
• The alkali metals and their salts impart characteristiccolour to an oxidizing flame and thus can be detected bythe respective flame tests and can bedetermined by flame photometry or atomic absorption spectroscopy.
• These elements whenirradiated with light, the light energy absorbedmay be sufficient to make a can atom lose electron.This property makes caesium and potassiumuseful as electrodes in photoelectric cells.
• Ionisation power:Alkali metals have the lowest initial ionisation energy among their periods.The ionisation energy drops as one moves down the group.
• Electropositivity character:Alkali metals are highly electropositive or metallic in nature, and this feature grows in strength as the alkali metals progress from Li to Cs.
• Electrode potentials:Alkali metals are strongly electropositive and oxidise easily in water, releasing hydrogen gas.
• Reducing nature:Due to their low ionisation energies, alkali metals are good reductants. Their lowering character is Na, Cs, Rb, K, Li (in aqueous solution).
• Electronegativity:The electronegativity of alkali metals decreases as one moves down the periodic table.
• Oxidation states:The outermost valence shell of all alkali metals contains only one electron. After the penultimate shell is completed, these atoms lose one electron to the nearest inert gas. As a result, the monovalent elements have an oxidation state +1.
• Hydration of ions:All alkali metal salts (excluding lithium) are ionic and soluble in water due to a strong hydration tendency of the alkali metal cations;
• Colour of flame:All alkali metals and associated salts give a non-luminous flame a distinct colour.Because of the absorption of energy, the electrons of an alkali metal and its salt are stimulated to higher energy levels when placed in a burning flame. This is a volatile state. Thus, a specific light colour is released as the electrons return towards the ground state. As one moves from Li to Cs, the ionisation energy decreases, increasing the frequency of the light emitted.
• The photoelectric effect:Alkali metals, in particular K and Cs, have a photoelectric effect because of their low ionisation activity and work function.
• The electrical conductivity:Since the valence electrons of all alkali metals may freely travel through the crystal structure, they are excellent heat and electric conductors. Electrical conductivity increases as one moves down the group.

Common chemical features of alkali metals are:

• The alkali metals are highly reactive due totheir large size and low ionization enthalpy. Thereactivity of these metals increases down thegroup.
• Reactivity towards air: The alkali metalstarnish in dry air due to the formation oftheir oxides which in turn react withmoisture to form hydroxides. They burnvigorously in oxygen forming oxides.Because oftheir high reactivity towards air and water,alkali metals are normally kept in kerosene oil.
• Reactivity towards water: The alkalimetals react with water to form hydroxideand dihydrogen. Reaction of lithium with water is less vigorous thanthat of sodium due to itssmall size and very high hydration energy.Other metals of the group react explosivelywith water.They also react with proton donors suchas alcohol, gaseous ammonia and alkynes.
• Reactivity towards dihydrogen: Thealkali metals react with dihydrogen atabout 673K to formhydrides. All the alkali metal hydrides areionic solids with high melting points.
• Reactivity towards halogens: The alkalimetals readily react vigorously withhalogens to form ionic halides, butlithium halides are somewhatcovalent due to highpolarisation capability of lithium ion.
• Reducing nature: The alkali metals arestrong reducing agents, lithium being themost and sodium the least powerful. With the small size of its ion, lithium hasthe highest hydration enthalpy whichaccounts for its high negative E0 value andits high reducing power.
• Solutions in liquid ammonia: The alkalimetals dissolve in liquid ammonia givingdeep blue solutions which are conductingin nature.

A 10.2

The general characteristics and gradation in properties of alkaline earth metals are:

• Atomic size goes on increasing down the group.
• Ionisation energy goes on decreasing down the group.
• They are harder than alkali metals.
• They are less electropositive than alkali metals.
• Electropositive character increases on going down the group.

A 10.3 All the alkali metals have one valence electron, ns1 outside the noble gas core. The loosely held s-electron in the outermost valence shell of these elements makes them the mostelectropositive metals, i.e. they readily lose electron to give monovalent M+ ions. Hence, they are never found in free state in nature.

A 10.4

Let x be the oxidation state of Na in Na₂O₂     

Then, 2x + 2 (-1) = 0  

      =>2x – 2 = 0  

      => x =  +1

A 10.5  It is because ionization enthalpy ∆H_i of potassium = 419 kJ mol^-1. Ionization enthalpy of sodium = 496 KJ mol^-1. Since Ionization enthalpy of potassium is less than that of sodium, potassium is more reactive than sodium.

A 10.6

(i) Ionization enthalpy. Because of high nuclear charge the ionization enthalpy
of alkaline earth metals are higher than those of the corresponding alkali metals.
(ii) Basicity of oxides. Basicity of oxides of alkali metals are higher than that of alkaline earth metals.
(iii) Solubility of hydroxides of alkali metals are higher than that of alkaline earth metals. Alkali metals due to lower ionization enthalpy are more electropositive than the corresponding group 2 elements.

A 10.7

1. Both lithium and magnesium are harder and lighter than other elements in the respective groups.
2. Lithium and magnesium react slowly with water. Their oxides and hydroxides are much less soluble and their hydroxides decompose on heating. Both form a nitride, Li3N and Mg3N2, by direct combination with nitrogen.
3. The oxides, Li2O and MgO do not combine with excess oxygen to give any superoxide. (iv) The carbonates of lithium and magnesium decompose easily on heating to form the oxides and CO2. Solid hydrogen carbonates are not formed by lithium and magnesium.
4. Both LiCl and MgCl2are soluble in ethanol.
5. Both LiCl and MgCl2are deliquescent and crystallise from aqueous solution as hydrates, LiCl·2H2O and MgCl26H2O.

A 10.8

Alkali and alkaline earth metals are reducing agents. That is why these metals are not obtained by chemical reduction methods.

A 10.9 Potassium and caesium have much lower ionization enthalpy than that of lithium. As a result, these metals easily emit electrons on exposure to light. Due to this, K and Cs are used in photoelectric cells rather than lithium

A 10.10 Different concentrations of alkali metals in liquid ammonia gives different colours. The dilute solutions impart blue colour due to presence of ammoniated electrons whereas the concentrated solutions have copper bronze colour as  ammoniated metal ions are bound by free electrons.

A 10.11 Due to small size, the ionization enthalpies of Be and Mg are much higher than those of other alkaline earth metals. Therefore, a large amount of energy is needed to excite their valence electron, and that’s why they do not impart colour to the flame.

A 10.12

Solvay process is also known as ammonia soda process, wherein carbon dioxide is passed through a brine solution (containing about 28 % NaCl) which is saturated with ammonia to form sodium carbonate.

2NH₃ + H₂O + CO₂ → (NH₄)₂CO₃

(NH₄)₂CO₃ + H₂O + CO₂ → 2NH₄HCO₃

NH₄HCO₃ + NaCl → NaHCO₃↓ + NH₄Cl

The precipitate of sodium bicarbonate is filtered, dried and ignited to form sodium carbonate.
2NaHCO₃​ →​ Na₂​CO₃​ + CO₂ ​+ H₂​O
The carbon dioxide required for the reaction can be obtained by heating limestone (calcium carbonate) to 1300 K in a lime klin. Lime dissolves in water to form calcium hydroxide which is then transferred to the ammonia recovery tower.
CaCO₃​→​ ​CaO+CO₂​
CaO+H_2​O→Ca(OH)_2​
Ammonia required for the process can be prepared by heating ammonium chloride with calcium hydroxide.
2NH_4​Cl+Ca(OH)₂→ ​2NH₃​+CaCl₂​+H₂​O
Hence, the only by product of the reaction is calcium chloride.

A 10.13 Potassium carbonate cannot be prepared by Solvay process because potassium hydrogencarbonate is too soluble to be precipitated by the addition of ammonium hydrogencarbonate to a saturated solution of potassium chloride.

A 10.14 Li₂CO₃ is a covalent compound whereas Na₂CO₃ is an ionic compound. Therefore, lattice energy of Na₂CO₃ is higher than that of Li₂CO₃. Thus, LiCO₃ is decomposed at a lower temperature as compared to Na₂CO₃

A 10.15

(a) Nitrates of both group 1 and group 2 elements are soluble in water because hydration energy is more than the lattice energy.
Nitrates of both group 1 and group 2 elements are thermally unstable but they decompose differently except LiCO₃ e.g.

2NaNO₃ → 2NaNO₂ + O₂

2KNO₃ → 2KNO₂ + O₂

4LiNO₃ → 2MgO + 4 NO₂ + O₂

2Mg(NO₃)₂ → 2MgO + 4NO₂ + O₂

(b) Carbonates of group 1 elements are soluble in water except Li₂CO₃ They are also thermally stable except Li₂CO₃.

Li₂CO₃ → Li₂O + CO₂
Group 2 carbonates are insoluble in water because their lattice energy is higher than hydration energy.
Thermal stability of carbonates of group 2 increases down the group because lattice energy goes on increasing due to increase in ionic character.
(c) Sulphates of group 1 are soluble in water except Li₂SO₄. They are thermally stable.
Solubility of sulphates of group 2 decreases down the group because Lattice energy dominates over hydration energy.
Sulphates of group 2 elements are thermally stable and increasing down the group due to increases in Lattice energy.

A 10.16

(i) Sodium metal is manufactured by electrolysis of a fused mass of NaCl 40% and CaCl₂ 60% in Down’s cell at 873 K, using iron as cathode and graphite as anode. Na is liberated at the cathode.
At cathode:
Na⁺ + e^– → Na (l)
At anode:
2Cl^– (melt) → Cl₂ (g) + 2e^–.
(ii) Sodium hydroxide is manufactured by electrolysis of an aqueous solution of NaCl (brine) in Castner-Kellner cell.
At cathode:
Na⁺ + e^– → Na
2Na + Hg → Na – Hg + 2H₂O
2Na- Hg + 2H₂O → 2NaOH +H₂ +Hg
At anode:
Cl^– – e^– → Cl
Cl + Cl→ Cl₂
(iii) Sodium peroxide:

Sodium is heated in excess of oxygen to form sodium peroxide. Initially, sodium oxide is obtained which reacts with more oxygen to form sodium peroxide.

4Na+O₂→2Na₂O

2Na₂O+O2→Na₂O₂
(iv)Sodium carbonate is obtained by Solvay ammonia process.
NaCl + NH₃ + H₂O + CO₂ → NaHCO₃↓ + NH₄Cl

2NaHCO₃→ Na₂CO₃ + H₂O + CO₂

A 10.17

(i) Magnesium is burnt in air to form magnesium oxide and magnesium nitride.

2Mg + O₂ → 2MgO

3Mg + N₂ → Mg₃N₂.

(ii) Quick lime is heated with silica above 1273 K to obtain calcium silicate

CaO+SiO₂ → CaSiO₃.

(iii) Chlorine reacts with slaked lime to form calcium hypochlorite- a constituent of bleaching powder.

2Ca(OH)₂ + 2Cl₂ → CaCl₂ + Ca(OCl)₂ + 2H₂O.

(iv) Calcium nitrate is heated to obtain CaO,NO₂ and O₂.

2Ca(NO₃)₂→2CaO+4NO₂+O₂.

A 10.18

(i) Caustic soda
It is used in (a) the manufacture of soap, paper, artificial silk and a number of chemicals, (b) in petroleum refining, (c) in the purification of bauxite, (d) in the textile industries for mercerising cotton fabrics, (e) for the preparation of pure fats and oils, and (f) as a laboratory reagent.
(ii) Sodium carbonate
(a) It is used in water softening, launderingand cleaning.

(b) It is used in the manufacture of glass,soap, borax and caustic soda.

(c) It is used in paper, paints and textileindustries.

(d) It is an important laboratory reagent bothin qualitative and quantitative analysis
(iii) Quick lime

(a) It is an important primary material formanufacturing cement and is the cheapest

form of alkali.

(b) It is used in the manufacture of sodiumcarbonate from caustic soda.

(c) It is employed in the purification of sugarand in the manufacture of dye stuffs.

A 10.19  BeCl₂ (vapour)

A 10.20 Since group 1 hydroxides and carbonates due to large size contain higher hydration energy than the lattice energy so, they are easily soluble in water. Whereas, in magnesium and calcium due to small size their lattice energy dominates over hydration energy they are sparingly soluble in water.

A 10.21

Limestone:

• A raw material for cement.
• It is used as a building material in the form ofmarble and in the manufacture of quick lime.
• used in the manufactureof high quality paper. It is also used as anantacid, mildabrasive in tooth paste, aconstituent of chewing gum, and a filler incosmetics.

Cement:

• It is used in concrete and reinforcedconcrete, in plastering and in the construction
• of bridges, dams and buildings

Plaster of Paris:

• It is used in the building industry as well as plasters.
• It is usedfor immobilising the affected part of organ wherethere is a bone fracture or sprain.
• It is alsoemployed in dentistry, in ornamental work andfor making casts of statues and busts.

A 10.22 Li⁺ can polarize water molecules easily than the other alkali metal ions because of its small size.

A 10.23 It is due to high lattice energy of LiF as compared to LiCl.
LiCl is soluble in water because its hydration energy is higher than its lattice energy.

A 10.24

Sodium:

• Na⁺ions participate in the transmission of nerve signals, in regulating the flow of water across cell membranes.
• In the transport of sugars and amino acids into cell.

Potassium:

• K⁺ions activate many enzymes.
• Participate in the oxidation of glucose to produce ATP.

Magnesium:

• All enzymes that utilise ATP in phosphate transfer require magnesium as a cofactor.
• Mg is the central metal ion present in chlorophyll pigment in plants.

Calcium:

• Ca²⁺ ions are present in bones.
• plays important roles in neuromuscular function.

A 10.25 (i) 2Na + 2H₂O → 2NaOH + H₂
(ii) 2Na + O₂ → Na₂O₂
(iii) Na₂O₂ + 2H20 → 2NaOH + H₂O₂

(a) The mobilities of the alkali metal ions in aqueous solution are Li⁺ < Na⁺  + < Rb ⁺ < Cs ⁺
(b) Lithium is the only alkali metal to form a nitride directly.
(c) E^ϴ for M²⁺ (aq) + 2e^– → M(s) (where M = Ca, Sr, or Ba) is nearly constant.
A 10.26 (a) Smaller the size of the ion, more highly it is hydrated and hence greater is the mass of the hydrated ion and thus the ionic mobility become lesser. The extent of hydration decreases in the order.
Li⁺ < Na⁺  + < Rb ⁺ < Cs ⁺
Thus the mobility of Cs ⁺ will be the highest.
(b) Due to its smaller size lithium can form nitride directly as an exceptional behaviour unlike other alkali metals.
(c) It is because reduction potential depends upon sublimation energy, ionisation energy and hydration energy. Their resultant is almost constant for these ions.

A 10.27

(a) Na₂CO₃ is a salt of a weak acid, carbonic acid (H₂CO₃) and a strong base NaOH. Thus, it undergoes hydrolysis to produce strong base NaOH and its aqueous solution isalkaline in nature.
Na₂CO₃(s) + H₂O(l) → 2NaOH
(b) Because the discharge potential of alkali metals is much higher than that of hydrogen, therefore when the aqueous solution of any alkali metal chloride is subjected to electrolysis, H₂, instead of the alkali metal, is produced at the cathode. Therefore, alkali metals are prepared by electrolysis of their fused chlorides.
(c) Since potassium is move reactive than sodium and it is found in nature to a less extent than Na, sodium is found to be more useful.

(a) Na₂O₂ and water
(b) KO₂ and water
(c) Na₂O and CO₂

A 10.28

(a) Na₂O₂ + 2H₂O → 2NaOH + H₂O₂
(b) 2KO₂ + 2H₂O → 2KOH + O₂+ H₂O₂
(c) Na₂O+ CO₂ → Na₂CO₃

A 10.29

(i) Lattice energy of BeO is comparatively higher than the hydration energy. Therefore, it is almost insoluble in water. Whereas, BeSO₄ is ionic in nature and its hydration energy dominates the lattice energy.
(ii) Both BaO and BaSO₄ are ionic compounds but the hydration energy of BaO is higher than the lattice energy therefore it is soluble in water.
(iii) Since the size of Li⁺ ion is very small in comparison to K⁺ ion, it polarises the electron cloud of I^– ion to a great extent. Thus, Lil dissolves in ethanol more easily than the KI.

A 10.30

(d) Cs

Size of Cs is the biggest thus, its melting point is the lowest, (d) is correct.

A 10.31

(a) Li

Li⁺ is the smallest in size and thus, has the highest charge density and hence attracts the water molecules more strongly to form hydrated salts

(a) MgCO₃ (b) CaCO₃ (c) SrCO₃ (d) BaCO₃(Intermediate)
A 10.32 (d) BaCO₃

Thermal stability is directly proportional to the size of the cation i.e., larger the size of the atom, greater is its thermal stability. Ba being the biggest cation among the given compounds, has a greater thermal stability.