The S -block Elements

The alkali metals are physically soft, silvery white and light metals. Due to the large atomic size, these elements have low density which increases down the group from Li to Cs. Caesium has the largest atomic size out of the given examples, and had weak metallic bonding ability due to the presence of a single valence electron. Hence it has the lowest melting point of 302°K or 28°C, hence it will melt if the room temperature rises to 30°C. The correct answer is (iv).

Alkali metals react with water and form hydroxides and dihydrogen.

2M + 2H₂O → 2M⁺ + 2OH^- + H₂ where M is an alkali metal. Lithium has the most negative E^Ө value, however it has the least vigorous reaction with water. It is even less vigorous than sodium, which has the least negative E^Ө value. This is because of the small size of Li and its high hydration energy. It is also because Li has the highest melting point. The correct answer is (i).

Alkali metals are strong reducing agents, with lithium being the strongest and sodium being the weakest reducing agent. Li is has the smallest atomic size and the highest hydration enthalpy. Hydration enthalpy is the amount of energy released when one mole of ions undergo hydration. Due to this, Li is the strongest reducing agent.

Carbonates of all alkaline earth metals decompose to give metal oxide and carbon dioxide. The thermal stability of a metal carbonate depends on the cationic size, it increases with increasing cationic size. The carbonates form the positively charged metal cation and negatively charged CO₃^2-. The larger the cationic size, the more electropositive is the ion. The electropositivity of the alkaline earth metals increases down the group from Be to Ba. Due to these reason, BaCO₃ is the most stable metal carbonate.

The correct option is (iv).

Beryllium carbonate is unstable and is kept in CO₂ atmosphere to avoid decomposition. Beryllium is also the least electropositive element out of the group due to the atomic size. BeCO₃ decomposes to form BeO and CO₂ as carbonates of all alkaline earth metals decompose to give metal oxide and carbon dioxide. This reaction is reversible, hence to shift the equilibrium to the right, BeCO₃ is placed in CO₂ atmosphere.

Alkaline earth metals burn in oxygen to form monoxides which are basic in nature (except BeO) and react with water to form sparingly soluble hydroxides. The solubility, thermal stability and the basic character of these hydroxides increase with increasing atomic number from Mg(OH)₂ to Ba(OH)₂. Ionization enthalpy increases from Mg to Ba, hence the M—O bond becomes weaker from Mg to Ba, and hence Mg will form the weakest base. Hence, Mg(OH)₂ is the least basic in nature. The correct option is (i).

Group 2, or alkaline earth metals form halides and except Beryllium, the metal halides are ionic in nature. Beryllium chloride however is covalent in nature due to the small size, high ionization enthalpy and high electronegativity. Due to this, BeCl₂ is soluble in organic solvents such as ethanol. The correct answer is (i).

Ionization enthalpy is the energy required to displace an electron from the valence shell of an atom in ground state. The smaller the size of the atomic radii, more energy is required to separate the valence electron. In the given metals, the size of the atomic radii in decreasing order is Li > N > K > Rb. The correct option is (iii).

Since lattice enthalpy decreases much more than the hydration enthalpy with increasing ionic size for alkali metals, the solubility of LiF is the lowest in water due to highest lattice enthalpy.

Amphoteric hydroxides are compounds which can behave both as an acid and base. Beryllium hydroxide is amphoteric in nature as it reacts with acid and alkali both. Beryllium hydroxide reacts with sodium hydroxide to form beryllate ion which is soluble in the same. The reaction is given as

Be(OH)₂ + 2OH^–→ [Be(OH)₄]^2–

Hence, the correct answer is (i).

Sodium carbonate is prepared from sodium by Solvay process. The equations in the process involved are as follows:

2NH₃ + H₂O + CO₂→ (NH₄)₂CO₃

(NH₄)₂CO₃ + H₂O + CO₂→ 2NH₄HCO₃

NH₄CO₃ + NaCl → NH₄Cl + NaHCO₃

2NaHCO₃→ Na₂CO₃ + CO₂ + H₂O

2NH₄Cl + Ca(OH)₂→ 2NH₃ + CaCl₂ + H₂O

Carbon dioxide is passed through a concentrated solution of sodium chloride saturated with ammonia, which forms ammonium carbonate followed by ammonium hydrogen carbonate. Ammonium hydrogen carbonate crystals separate and they are heated to form sodium carbonate. NH₃ is recovered from the solution which contains NH₄Cl is heated and treated with Ca(OH)₂. Calcium chloride is obtained as a by-product. The correct answer is (i).

Alkali metals like sodium dissolve in liquid ammonia to give a deep blue colour. The reaction can be given as

M + (x+y)NH₃ → [M(NH₃)_x]⁺ + [e(NH₃)_y]^-

The deep blue colour is due to the ammoniated electron. It absorbs energy in the visible region of light and thus imparts blue colour to the solution. The correct answer is (i).

Gypsum is CaSO₄.2H₂O. Gypsum is added in 2-3% of weight to cement along with CaO (lime) and clay. Gypsum is added in order to increase the setting time of cement. Slowing the setting time of cement makes it harden sufficiently. The correct answer is (ii).

Dead burnt plaster is formed when gypsum CaSO₄.2H₂O is heated to 393K. This causes loss of water of crystallisation and forms anhydrous calcium sulphate.

2(CaSO₄.2H₂O) → 2(CaSO₄).H₂O + 3H₂O

Suspension of slaked lime in water is known as milk of lime. It is a suspension of calcium hydroxide in water. The correct answer is (iii).

Alkali metals react with dihyrogen at high temperatures to form hydrides. Beryllium has a small atomic size and has a high ionisation enthalpy. Due to this, it does not react on direct heating with dihydrogen. The correct answer is (i).

Soda ash is formed when sodium carbonate decahydrate, washing soda, is heated above 373K. This causes Na₂CO₃.10H₂O to lose all water of crystallisation to form anhydrous sodium carbonate or soda ash. The correct answer is (iv).

Compounds containing calcium burn with a brick red flame. Hence Calcium nitrate will break down on heating to give a brick red flame and a brown gas which is NO₂.

2Ca(NO₃)₂ → 2CaO + 4NO₂ + O₂

Calcium hydroxide or Ca(OH)₂ reacts with chlorine to form hypochlorite, which is a component of bleaching powder. The reaction is as follows.

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

The correct answer is (i).

Chemical A used in the preparation of washing soda Na₂CO₃ to recover ammonia is Ca(OH)₂. It is a reaction involved in the Solvay process for preparation of sodium carbonate.

2NH₄Cl + Ca(OH)₂→ 2NH₃ + CaCl₂ + H₂O

When CO₂ is bubbled through a solution of Ca(OH)₂, it forms milky precipitate of CaCO₃.

Ca(OH)₂ + CO₂→ CaCO₃ + H₂O

Ca(OH)₂ is used as slaked lime as white wash due its disinfectant nature.

The substances get wet on contact with air as they are hygroscopic in nature and can absorb moisture around them. The tendency to form halide hydrates gradually decreases down the group. For example, MgCl₂ .8H₂O, CaCl₂.6H₂O, SrCl₂.6H₂O and BaCl₂.2H₂O. Hence, all the properties are correct and the correct answer is (iv).

Out of the given properties, alkali metals are characterised by their high negative standard electrode potential and large atomic size. Alkali metals are characterised by being soft and light. They have low boiling point and also low density due to the large atomic size.

NaOH is used in textile industries as a mercerising agent for cotton fabrics, Na₂CO₃ is also used in textiles and as water softening, laundering and cleaning agents.

Out of the given compounds, BeSO₄ and MgSO₄ are readily soluble in water. The solubility of the sulphates decreases from CaSO₄ to BaSO₄. The compounds are readily soluble because the greater hydration enthalpies of Be²⁺ and Mg²⁺ ions overcome the lattice enthalpy factor.

Zeolite is hydrated sodium aluminium silicate, or Al₂Si₂O₈. Hard water consists mainly of calcium and magnesium salts, hence sodium ions are exchanged with mainly Mg²⁺ and Ca²⁺ ions. The reaction is given as

Na₂Z + M²⁺→ MZ + 2Na⁺

Where M is Mg²⁺ or Ca²⁺ and Z is Zeolite.

The tendency to make halide hydrates decreases down the group of alkaline earth metals. Examples are MgCl₂ .8H₂O, CaCl₂.6H₂O, SrCl₂.6H₂O and BaCl₂.2H₂O. Hence the correct answers are (i) and (iii).

Beryllium is chemically inert to air and water and is not readily attacked by acids because of the presence of a BeO film on the surface of the metal. BeSO₄ is readily soluble in water. This is because Be²⁺ ions overcome the lattice enthalpy factor. Beryllium does not have a coordination number of more than 4. Beryllium oxide is not purely acidic but amphoteric in nature.

Lithium shows anomalous behaviour where it shows similar properties to alkaline earth metals, known as diagonal relationship. This is due to the exceptionally small size of the atom. Anomalous behaviour is also due to the high polarising power of the lithium ion. Polarisation is the distortion of electron cloud of the anion by the cation. Options (iii) and (iv) are also properties of lithium but they are also properties of alkali metals. The correct answers are (i) and (ii).

Alkali metals are strong reducing agents, with lithium being the strongest of the group. The reducing power of an element can be determined by its standard electrode potential, which is a measure of the tendency of the element to lose electrons in aqueous solution. It is determined by three factors: (i) Sublimation enthalpy (ii) Ionization enthalpy and (iii) Enthalpy of hydration. Sublimation enthalpy of alkali metals is similar. Lithium has highest negative E^Ө value, which is –3.04V. Lithium has a small atomic size, the highest ionization enthalpy but it is compensated by its high hydration enthalpy. Due to this, the reducing power of lithium is highest in an aqueous solution.

Alkali metals are highly reactive and their chemical reactivity increases down the group. Alkali metals burn vigorously in oxygen forming oxides. Lithium forms monoxide, sodium forms peroxide, and potassium form superoxide. The reactions are as follows:

4LiO +O₂→ 2Li₂O (oxide)

2Na + O₂→ Na₂O₂ (peroxide) with Na₂O as a minor product

K + O₂→ KO₂ (superoxide)

The formation of oxide to superoxide is

O₂^- + 1/2O₂→ O₂^2- + O₂→ 2O₂^-

Oxide peroxide superoxide

The increasing stability of the peroxide or superoxide, as the size of the metal ion increases, is due to the stabilisation of large anions by larger cations through lattice energy effects.

(i) Peroxide ions react with H₂O to form hydrogen peroxide.

O₂^2- + 2H₂O → H₂O₂ + 2OH^-

(ii) Superoxide ions react with H₂O to form hydrogen peroxide and O₂.

O₂^- + 2H₂O → H₂O₂ + O₂ + 2OH^-

Lithium has some properties which resemble the properties of the second element of Group 2, magnesium. This is known as a diagonal relationship in the periodic table.

(i) Lithium and magnesium are both lighter and harder than the other metals in their respective groups.

(ii) Halides of both elements, LiCl and MgCl₂ are soluble in ethanol.

(iii) Both lithium and magnesium form a nitride Li₃N and Mg₃N₂ respectively.

These two elements have similar properties because of their similar atomic and ionic radii.

The element from Group 2 is Beryllium. Beryllium reacts with oxygen to give BeO which is covalent in nature, unlike other elements of Group 2 which form ionic compounds. Beryllium oxide is amphoteric in nature unlike the other compounds which are basic in nature. Sulphates of Group 2 are soluble in water and BeSO₄ is readily soluble in water. This is because of the hydration enthalpy of Be²⁺ ions being much higher than the lattice enthalpy of BeSO₄.

(i) All Group 2 elements form carbonates. All carbonates of alkaline earth metals are insoluble in water. These carbonates decompose on heating to give carbon dioxide and the metal oxide. The thermal stability of these carbonates increases down the group, as the solubility reduces. The order is BeCO₃ < MgCO₃ < CaCO₃ < SrCO₃ < BaCO₃. Beryllium carbonate is unstable and requires to be in an atmosphere of CO₂. The thermal stability of the carbonates increases with increases with increasing cationic size. The more stable the oxide of an alkaline earth metal, the less stable is the carbonate of the same. Hence BeCO₃ is highly unstable as BeO is stable.

(ii) All alkaline earth metals form oxides with oxygen in the formula MO where M is the metal. The alkaline earth metals burn in oxygen to give metal oxides. The compounds are ionic in nature except BeO which is covalent. The oxides are basic in nature except BeO which is amphoteric in nature and behaves as both acid and base. The basic nature increases down the group due to the increase in ionic nature. The enthalpy of formation of these oxides are high and hence they are thermally stable. They also react with water to form sparingly soluble hydroxides. As the size of the cations increase, the lattice energy decreases. BeO and MgO have the highest lattice energy and they are insoluble in water.

The solubility of the alkaline earth metal sulphates decrease down the group. While BeSO₄ and MgSO₄ are readily soluble in water, CaSO₄, SrSO₄, and BaSO₄ are insoluble. This is because the greater hydration enthalpies of Be²⁺ and Mg²⁺ ions overcome the lattice enthalpy factor and therefore their sulphates are soluble in water. The hydration enthalpies of the rest compounds do not match the lattice enthalpy factor, thus making these sulphates insoluble in water.

All compounds of alkali metals are easily soluble in water but lithium compounds are more soluble in organic solvents because the alkali metal compounds form ionic compounds due to their large ionic size and low ionization enthalpy, while lithium forms compounds of covalent nature due to their small ionic size, high ionization enthalpy, and high electronegativity. Ionic compounds dissolve in water, and covalent compounds dissolve in organic solvents.

We cannot obtain sodium carbonate directly by treating the (NH₄)₂CO₃ solution with NaCl, sodium chloride. In the Solvay process, carbon dioxide is passed through a concentrated solution of sodium chloride saturated with ammonia, which forms ammonium carbonate followed by ammonium hydrogen carbonate. Ammonium hydrogen carbonate crystals separate and they are heated to form sodium carbonate. NH₃ is recovered from the solution which contains NH₄Cl is heated and treated with Ca(OH)₂. The reaction of (NH₄)₂CO₃ with NaCl gives two products, Na₂CO₃ and NH₄Cl which are both soluble in water which does not shift the equilibrium to the right.

(NH₄)₂CO₃⇌ NH₄Cl + Na₂CO₃

The Lewis structure of O₂^- is . Oxygen atom with zero charges has 6 electrons, therefore the oxidation state is 0. When the oxygen atom has a negative charge, it has 7 electrons. Hence, the oxidation state is -1. The average oxidation state is 0 + (-1)/2 = -1/2.

Alkaline earth metals except Be and Mg impart a characteristic colour to the Bunsen burner flame in the flame test. The different colours arise due to the different energies (provided by heat in this case) for the states of electron excitation and de-excitation. Be and Mg electrons are tightly bound to the atom due to the small atomic and ionic size. The electrons do not gain excitation from the energy provided from the flame, hence they do not impart any colour to the flame.

Beryllium chloride is covalent in nature.

The structure of BeCl₂ in solid state is a polymeric chain structure and each Be atom is surrounded by four Cl atoms. Two Cl atoms are covalently bonded and two are bonded by coordinate bonds. The gaseous/vapour state is different than the solid state. BeCl₂ tends to form a chloro-bridged dimer at temperatures below 1200K and dissociates into a linear monomer at high temperatures of the order of 1200 K.

(i) → (c)

As Li has the most negative E^⊖ value due to the high hydration energy.

(ii) → (b)

As Alkali metals are more acidic than the given alkaline earth metals, while Li base is covalent in nature.

(iii) → (d)

Calcium oxalate is insoluble in water

(iv) → (a)

Due to the high hydration energy.

(iv) → (e)

Electronic configuration of Ba is 1s²2s²2p⁶3s²3p⁶3d¹⁰4s²4p⁶4d¹⁰5s²5p⁶6s².

(i) → (c)

Calcium carbonate used in manufacturing high quality paper because the paper and cardboard industries use lime-based coating pigments and filters such as GCC(Ground Calcium Carbonate). GCC is made from concentrated and fine-ground calcium carbonate and used to make fine paper, cardboard packaging and pulp-based paper.

(ii) → (d)

Calcium Hydroxide is used in white washing because for white washing walls as it slowly reacts with CO₂in air to form a thin layer of Calcium Carbonate on the walls. The layer of Calcium Carbonate gives a shiny finish to the walls.

Ca(OH)₂ + CO₂→ CaCO₃ + H₂O

(iii) → (b)

CaO is used in manufacturing of sodium carbonate from caustic soda.

Above process is also known as Solvay Process. This reaction occurs in two steps.

First step,

NaCl + CO₂ + NH₃ + H₂O → NaHCO₃ + NH₄Cl(l)

First Ammonia converts into Sodium Hydrogen Carbonate and Ammonium Chloride.

Second Step,

2NaCl + CaCO₃→Na₂CO₃ + CaCl₂

So, in the next or second step, there is the formation of Sodium Carbonate.

(iv) → (a)

CaSO₄ is used in Dentistry, Ornamental work because CaSO₄with adequate quantity of water, it forms a plastic mass that sets into a hard solid.

The colour of the flame imparted depends on the energy required for electron excitation and deexcitation. This happens because the heat from the flame excites the outermost orbital electron to a higher energy level. When the excited electron comes back to the ground state, there is emission of radiation in the visible region of the spectrum.

(i) → (f)

The colour of the flame imparted depends on the energy required for electron excitation and deexcitation. This happens because the heat from the flame excites the outermost orbital electron to a higher energy level. When the excited electron comes back to the ground state, there is emission of radiation in the visible region of the spectrum.

So, Cs imparts blue colour.

(ii) → (d)

The colour of the flame imparted depends on the energy required for electron excitation and deexcitation. This happens because the heat from the flame excites the outermost orbital electron to a higher energy level. When the excited electron comes back to the ground state, there is emission of radiation in the visible region of the spectrum.

Thus, Na imparts yellow colour.

(iii) → (b)

The colour of the flame imparted depends on the energy required for electron excitation and deexcitation. This happens because the heat from the flame excites the outermost orbital electron to a higher energy level. When the excited electron comes back to the ground state, there is emission of radiation in the visible region of the spectrum.

Thus, Potassium imparts violet colour.

(iv) → (c)

The colour of the flame imparted depends on the energy required for electron excitation and deexcitation. This happens because the heat from the flame excites the outermost orbital electron to a higher energy level. When the excited electron comes back to the ground state, there is emission of radiation in the visible region of the spectrum.

Thus, Ca imparts brick red colour.

(v) → (e)

The colour of the flame imparted depends on the energy required for electron excitation and deexcitation. This happens because the heat from the flame excites the outermost orbital electron to a higher energy level. When the excited electron comes back to the ground state, there is emission of radiation in the visible region of the spectrum.

Thus, Sr imparts crimson red colour.

(vi) → (a)

The colour of the flame imparted depends on the energy required for electron excitation and deexcitation. This happens because the heat from the flame excites the outermost orbital electron to a higher energy level. When the excited electron comes back to the ground state, there is emission of radiation in the visible region of the spectrum.

Thus, Barium imparts apple green colour.

(i) Both A and R are correct and R is the correct explanation of A.

Explanation: Unlike other alkali metal carbonates which are stable in heat Lithium being very small in size polarises a large CO₃^2– ion leading to the formation of more stable Li₂O and CO₂.

Li₂CO₃→ Li₂O + CO₂

Polarisation is the distortion of electron cloud of the anion CO₃^2- by the cation Li²⁺.

(i) Both A and R are correct and R is the correct explanation of A.

Explanation: All alkaline earth metal carbonates are stable at room temperature except beryllium carbonate. It decomposes to give BeO and CO₂, hence it is kept in a CO₂ atmosphere so that the equilibrium shifts to the right of the equation.

BeCO₃⇌ BeO + CO₂

The s-block elements contain two groups on the periodic table, the alkali metals (Group 1) and the alkaline earth metals (Group 2). The nature of their oxides, halides and oxosalts is given as follows:

1) Alkali metals:

(i) Oxides: Alkali metals form oxides on combustion in excess air. Lithium forms oxide Li₂O and peroxide Li₂O₂ in minor quantities. Sodium forms the peroxide, Na₂O₂ (and superoxide NaO₂ in minor quantities) whilst potassium, rubidium and caesium form the superoxides, MO₂. Going down the group, the size of the atoms increases and the ability to form peroxides and superoxides also increases. This happens due to the stabilisation of large anions by larger cations through lattice energy effects.

(ii) Halides: Alkali metal halides can be produced by reaction with the appropriate carbonate, oxide or hydroxide and hydrohalic acid (HX). The melting and boiling points follow the trend flouride > chloride > bromide > iodide. All alkali metal fluorides have high negative enthalpy of formation. The values for fluorides become less negative down the group and for chlorides, bromides and iodides the values become more negative down the group. All of these halides are highly soluble in water except LiF which has high lattice energy.

(iii) Oxosalts: All alkali metals react with oxo-acids like carbonic acid H₂CO₃ and sulphuric acid H₂SO₄ to form oxo-salts like carbonate and sulphate respectively. They are soluble in water and thermally stable. The carbonates and hydrogencarbonates, M₂CO₃ and MHCO₃ respectively are thermally stable. The stability of these compounds increases with increasing electropositive character down the group. The exception is Li₂CO₃ which is not thermally stable due to the small size of the Li atom, which polarizes the large CO₂^3- ion. This leads to the decomposition and formation of LiO and CO₂. The hydrogencarbonate does not exist as a solid.

2) Alkaline Earth Metals

(i) Oxides: Alkaline earth metals burn in oxygen to form metal oxides in the form of MO. The oxides are ionic in nature except BeO which is covalent in nature. The enthalpies of formation of oxides is high. The oxides are basic in nature except BeO which is amphoteric.

(ii) Halides: Halides of alkaline earth metals are ionic in nature with the exception of BeCl₂ which is covalent and is soluble in organic solvents. The fluorides are relatively less soluble than the chlorides owing to their high lattice energies.

(iii) Oxosalts: Alkaline earth metals form oxosalts with oxo-acids. They form carbonates with carbonic acid, sulphates with sulphuric acid and nitrates with nitric acid. Carbonates are insoluble in water and they decompose on heating to form the oxides and carbon dioxide. Beryllium carbonate is unstable and has to be placed in a CO₂ atmosphere to stop it from decomposing. The sulphates are stable in heat. BeSO₄ and MgSO₄ are readily soluble in water because of hydration enthalpies of Be²⁺ and Mg²⁺ overcoming lattice enthalpy, and the solubility reduces from CaSO₄ to BaSO₄. Nitrates show decreasing tendency to form hydrates with increasing size and decreasing hydration enthalpy in the group. All of them decompose to give metal oxide MO and NO₂ and O₂.

Alkali metals dissolve in liquid ammonia to give a deep blue colour. These solutions are also conducting in nature. The blue colour is obtained because of ammoniated electron which absorbs energy in the visible region of light and thus imparts blue colour to the solution. The reaction is given as

M + (x+y)NH₃ → [M(NH₃)_x]⁺ + [e(NH₃)_y]^-

As the concentration of the metal increases, the colour of the solution become bronze in colour, which is due to the formation of metal ion clusters.

If the product formed is kept standing for some time, it slowly liberates hydrogen resulting in the formation of amide.

M⁺_(am) + e^- +NH₃(l) → MNH_2(am) +1/2H_2(g)

The equation of formation of oxide, peroxide, and superoxide is

O₂^- + 1/2O₂→ O₂^2- + O₂→ 2O₂^-

Oxide peroxide superoxide

The increasing stability of the peroxide or superoxide, as the size of the metal ion increases, is due to the stabilisation of large anions by larger cations through lattice energy effects.

The size of the cations down group 1 increases from Li to Cs. Li has the smallest atomic and ionic size and has a strong positive field around it. Li⁺ combines with small anion called oxide ion (O^2-) with strong negative field and form Lithium oxide, Li₂O and Lithium peroxide in very small amounts.

Na⁺ ion is larger than Li⁺ ion and hence with less positive field around it. It combines with peroxide ion (O₂^2-) with less negative field and form sodium peroxide and sodium superoxide in small quantity.

K⁺, Rb⁺ and Cs⁺ ions are bigger in size with even lesser positive field around them. Hence they can stabilize the bigger superoxide ion (O^2-) with less negative field and form super oxides.

Compound (A) is Calcium oxide, CaO. Addition of water to CaO forms Compound (B), which is Ca(OH)₂.

CaO + H₂O → Ca(OH)₂

When CO₂ is passed through the Ca(OH)₂ solution, it turns milky and forms a white precipitate which is Compound (C) which is calcium carbonate, CaCO₃.

Ca(OH)₂ + CO₂→ CaCO₃ + H₂O

On passing excess of CO₂, the precipitate dissolves to form calcium hydrogencarbonate, which is compound (D).

CaCO₃ + CO₂ + H₂O → Ca(HCO₃)₂

The milkiness dissolves due to formation of a soluble compound calcium hydrogencarbonate.

Beryllium cannot react with dihydrogen like other alkaline earth metals to form beryllium hydride. Lithium hydride on the other hand is formed by heating with dihydrogen at 1073K. Steps to form beryllium hydride from lithium hydride are as follows.

Lithium hydride reacts with anhydrous aluminium chloride in presence of ether to form LiAlH₄, lithium aluminium hydride.

LiH + Al₂Cl₆→ 2LiAlH₄ + 6LiCl

LiAlH₄ reacts with BeCl₂ to form BeH₂.

2BeCl₂ + LiAlH₄→ 2BeH₂ + LiCl + AlCl₃

The only element from Group 2 which forms a covalent oxide and is amphoteric in nature is Beryllium, Be. The covalent oxide is BeO which is amphoteric. When dissolved in water, it forms Be(OH)₂. The reactions of Be(OH)₂ with an acid and a base are as follows.

Be(OH)₂ reacts with a base to forms beryllate ion.

Be(OH)₂ + 2OH^–→ [Be(OH)₄]^2–

Be(OH)₂ reacts with an acid to form beryllate chloride.

Be(OH)₂ + 2HCl + 2H₂O → [Be(OH)₄]Cl₂

Sodium ions participate in transfer of nerve signals, regulating the flow of water across cell membranes and in the transport of sugars and amino acids into cells. Sodium imparts a golden yellow colour to the flame in a flame test. This happens because the heat from the flame excites the outermost orbital electron to a higher energy level. When the excited electron comes back to the ground state, there is emission of radiation in the visible region of the spectrum, and for sodium is yellow at 589.2 λ/nm.

The reaction of sodium with oxygen to form sodium peroxide is given as

2Na + O₂→ Na₂O₂ (peroxide)

The stability of the sodium peroxide is due to the stabilisation of large anion O₂^2- by larger cation Na⁺.