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Let the oxidation state of Na be x. The oxidation state of oxygen, in case of peroxides, is –1.

Therefore,

Therefore, the oxidation sate of sodium is +1.

Atomic size increases as we move down the alkali group. As a result, the binding energies of their atoms in the crystal lattice decrease. Also, the strength of metallic bonds decreases on moving down a group in the periodic table. This causes a decrease in the melting point. Among the given metals, Cs is the largest and has the least melting point.

Physical properties of alkali metals are as follows.

(1) They are quite soft and can be cut easily. Sodium metal can be easily cut using a knife.

(2) They are light coloured and are mostly silvery white in appearance.

(3) They have low density because of the large atomic sizes. The density increases down the group from Li to Cs. The only exceptionto this isK, which has lower density than Na.

(4) The metallic bonding present in alkali metals is quite weak. Therefore, they have low melting and boiling points.

(5) Alkali metals and their salts impart a characteristic colour to flames. This is because the heat from the flame excites the electron present in the outermost orbital to a high energy level. When this excited electron reverts back to the ground state, it emits excess energy as radiation that falls in the visible region.

(6) They also display photoelectric effect. When metals such as Cs and K are irradiated with light, they lose electrons.

Chemical properties of alkali metals

Alkali metals are highly reactive due to their low ionization enthalpy. As we move down the group, the reactivity increases.

(1) They react with water to form respective oxides or hydroxides. As we move down the group, the reaction becomes more and more spontaneous.

(2) They react with water to form their respective hydroxides and dihydrogens. The general reaction for the same is given as

(3) They react with dihydrogen to form metal hydrides. These hydrides are ionic solids and have high melting points.

(4) Almost all alkali metals, except Li, react directly with halogens to form ionic halides.

Since Li⁺ ion is very small in size, it can easily distort the electron cloud around the negative halide ion. Therefore, lithium halides are covalent in nature.

(5) They are strong reducing agents. The reducing power of alkali metals increases on moving down the group. However, lithium is an exception. It is the strongest reducing agent among the alkali metals. It is because of its high hydration energy.

(6) They dissolve in liquid ammonia to form deep blue coloured solutions. These solutions are conducting in nature.

The ammoniated electrons cause the blue colour of the solution. These solutions are paramagnetic and if allowed to stand for some time, then they liberate hydrogen. This results in the formation of amides.

In a highly concentrated solution, the blue colour changes to bronze and the solution becomes diamagnetic.

Alkali metals include lithium, sodium, potassium, rubidium, cesium, and francium. These metals have only one electron in their valence shell, which they lose easily, owing to their low ionization energies. Therefore, alkali metals are highly reactive and are not found in nature in their elemental state.

In alkali metals, on moving down the group, the atomic size increases and the effective nuclear charge decreases. Because of these factors, the outermost electron in potassium can be lost easily as compared to sodium. Hence, potassium is more reactive than sodium.

(i) Uses of caustic soda

(a) It is used in soap industry.

(b) It is used as a reagent in laboratory.

(ii) Uses of sodium carbonate

(a) It is generally used in glass and soap industry.

(b) It is used as a water softener.

(iii) Uses of quick lime

(a) It is used as a starting material for obtaining slaked lime.

(b) It is used in the manufacture of glass and cement.

LiF is insoluble in water. On the contrary, LiCl is soluble not only in water, but also in acetone. This is mainly because of the greater ionic character of LiF as compared to LiCl. The solubility of a compound in water depends on the balance between lattice energy and hydration energy. Since fluoride ion is much smaller in size than chloride ion, the lattice energy of LiF is greater than that of LiCl. Also, there is not much difference between the hydration energies of fluoride ion and chloride ion. Thus, the net energy change during the dissolution of LiCl in water is more exothermic than that during the dissolution of LiF in water. Hence, low lattice energy and greater covalent character are the factors making LiCl soluble not only in water, but also in acetone.

(a) Na2O2 and water,
When Na2O2 reacts with water sodium hydroxide and Hydrogen peroxide are formed.
e.g., Na2O2 + 2H2O ---> 2NaOH + H2O2

(b) KO2 and water,
when KO2 reacts with water potassium hydroxide , hydrogen peroxide and oxygen gas are formed.
e.g., 2KO2 + 2H2O --->2KOH + H2O2 + O2

(c) Na2O and CO2,
when Na2O2 reacts with CO2 sodium carbonate is formed.
e.g., Na2O + CO2 ---> Na2CO3

(i) BeO is almost insoluble in water and BeSO₄ is soluble in water. Be²⁺ is a small cation with a high polarising power and O^2– is a small anion. The size compatibility of Be²⁺ and O^2– is high. Therefore, the lattice energy released during their formation is also very high. When BeO is dissolved in water, the hydration energy of its ions is not sufficient to overcome the high lattice energy. Therefore, BeO is insoluble in water. On the other hand, ion is a large anion. Hence, Be²⁺ can easily polarise ions, making BeSO₄ unstable. Thus, the lattice energy of BeSO₄ is not very high and so it is soluble in water.

(ii) BaO is soluble in water, but BaSO₄ is not. Ba²⁺ is a large cation and O^2– is a small anion. The size compatibility of Ba²⁺ and O^2– is not high. As a result, BaO is unstable. The lattice energy released during its formation is also not very large. It can easily be overcome by the hydration energy of the ions. Therefore, BaO is soluble in water. In BaSO₄, Ba²⁺ and are both large-sized. The lattice energy released is high. Hence, it is not soluble in water.

(iii) LiI is more soluble than KI in ethanol. As a result of its small size, the lithium ion has a higher polarising power than the potassium ion. It polarises the electron cloud of the iodide ion to a much greater extent than the potassium ion. This causes a greater covalent character in LiI than in KI. Hence, LiI is more soluble in ethanol.

From the given options the most thermally stable alkaline earth metal carbonate among Magnesium Carbonate, Calcium Carbonate, Strontium Carbonate, Barium Carbonate is Magnesium Carbonate.
This is because it has the “smallest atomic number” due to which the size of the molecules or atom is the smallest and most stable. Other options are not as stable as it is because they have a higher atomic number, bigger atomic size and thus the outer electrons are comparatively loosely bound and less “thermally stable”.

General characteristics of alkaline earth metals are as follows.

(i) The general electronic configuration of alkaline earth metals is [noble gas] ns².

(ii) These metals lose two electrons to acquire the nearest noble gas configuration. Therefore, their oxidation state is +2.

(iii)These metals have atomic and ionic radii smaller than that of alkali metals. Also, when moved down the group, the effective nuclear charge decreases and this causes an increase in their atomic radii and ionic radii.

(iv)Since the alkaline earth metals have large size, their ionization enthalpies are found to be fairly low. However, their first ionization enthalpies are higher than the corresponding group 1 metals.

(v) These metals are lustrous and silvery white in appearance. They are relatively less soft as compared to alkali metals.

(vi)Atoms of alkaline earth metals are smaller than that of alkali metals. Also, they have two valence electrons forming stronger metallic bonds. These two factors cause alkaline earth metals to have high melting and boiling points as compared to alkali metals.

(vii) They are highly electropositive in nature. This is due to their low ionization enthalpies. Also, the electropositive character increases on moving down the group from Be to Ba.

(viii) Ca, Sr, and Ba impart characteristic colours to flames.

Ca – Brick red

Sr – Crimson red

Ba – Apple green

In Be and Mg, the electrons are too strongly bound to be excited. Hence, these do not impart any colour to the flame.

The alkaline earth metals are less reactive than alkali metals and their reactivity increases on moving down the group. Chemical properties of alkaline earth metals are as follows.

(i) Reaction with air and water: Be and Mg are almost inert to air and water because of the formation of oxide layer on their surface.

(a) Powdered Be burns in air to form BeO and Be₃N₂.

(b) Mg, being more electropositive, burns in air with a dazzling sparkle to form MgO and Mg₃N₂.

(c) Ca, Sr, and Ba react readily with air to form respective oxides and nitrides.

(d) Ca, Ba, and Sr react vigorously even with cold water.

(ii) Alkaline earth metals react with halogens at high temperatures to form halides.

(iii) All the alkaline earth metals, except Be, react with hydrogen to form hydrides.

(iv) They react readily with acids to form salts and liberate hydrogen gas.

(v) They are strong reducing agents. However, their reducing power is less than that of alkali metals. As we move down the group, the reducing power increases.

(vi) Similar to alkali metals, the alkaline earth metals also dissolve in liquid ammonia to give deep blue coloured solutions.

Similarities between lithium and magnesium are as follows.

(i) Both Li and Mg react slowly with cold water.

(ii) The oxides of both Li and Mg are much less soluble in water and their hydroxides decompose at high temperature.

(iii) Both Li and Mg react with N₂ to form nitrides.

(iv) Neither Li nor Mg form peroxides or superoxides.

(v) The carbonates of both are covalent in nature. Also, these decompose on heating.

(vi) Li and Mg do not form solid bicarbonates.

(vii) Both LiCl and MgCl₂ are soluble in ethanol owing to their covalent nature.

(viii) Both LiCl and MgCl₂ are deliquescent in nature. They crystallize from aqueous solutions as hydrates, for example, and .

In the process of chemical reduction, oxides of metals are reduced using a stronger reducing agent. Alkali metals and alkaline earth metals are among the strongest reducing agents and the reducing agents that are stronger than them are not available. Therefore, they cannot be obtained by chemical reduction of their oxides.

Concept:- Metals have a high tendency to lose electrons. (lower the ionization energy, higher is the tendency to lose electrons) that's why metals are used in photoelectric cells.
Potassium and Caesium have much lower Ionization enthalpy than that of Lithium. Therefore, these metals on exposure to light, emit electrons efficiently, but Lithium doesn't. That's why K and Cs rather than Li, are used in photoelectric cells.

When an alkali metal dissolves in liquid ammonia, the solution can acquire different colours. Meaning, dilute solutions of alkali metals in liquid ammonia exhibit dark blue colour this is because of ammoniated electrons absorb energy in the visible region of light.
e.g., M + (x + y)--->[M()x]^+ + [e-()y]{ammoniated electrons}
However, if the concentration increases above 3M, the colour changes to copper -bronze, and it becomes diamagnetic.

When an alkaline earth metal is heated, the valence electrons get excited to a higher energy level. When this excited electron comes back to its lower energy level, it radiates energy, which belongs to the visible region. Hence, the colour is observed. In Be and Mg, the electrons are strongly bound. The energy required to excite these electrons is very high. Therefore, when the electron reverts back to its original position, the energy released does not fall in the visible region. Hence, no colour in the flame is seen.

Solvay process is used to prepare sodium carbonate.

When carbon dioxide gas is bubbled through a brine solution saturated with ammonia, sodium hydrogen carbonate is formed. This sodium hydrogen carbonate is then converted to sodium carbonate.

Step 1: Brine solution is saturated with ammonia.

This ammoniated brine is filtered to remove any impurity.

Step 2: Carbon dioxide is reacted with this ammoniated brine to result in the formation of insoluble sodium hydrogen carbonate.

NH3+H2O+CO2→NH4HCO3NaCl+NH4HCO3→NaHCO3+NH4ClNH3+H2O+CO2→NH4HCO3NaCl+NH4HCO3→NaHCO3+NH4Cl

Step 3: The solution containing crystals of NaHCO₃ is filtered to obtain NaHCO₃.

Step 4: NaHCO₃ is heated strongly to convert it into NaHCO₃.

Step 5: To recover ammonia, the filtrate (after removing NaHCO₃) is mixed with Ca(OH)₂ and heated.

Ca(OH)2+2NH4Cl→2NH3+2H2O+CaCl2Ca(OH)2+2NH4Cl→2NH3+2H2O+CaCl2

The overall reaction taking place in Solvay process is

Formation of Na2CO3 in Solvay ammonia process, on passing of  CO through brine solution saturated with ammonia. Sodium bicarbonate is precipitated out.
NaCl + NH3 + H2O --->NaHCO3 ↓
 In the case of Potassium,  preparation cannot be done by the Solvay process. This because, potassium bicarbonate being more soluble than sodium bicarbonate, doesn't get precipitated when CO2 is passed through a concentrated solution of KCl saturated with ammonia.

As we move down the alkali metal group, the electropositive character increases. This causes an increase in the stability of alkali carbonates. However, lithium carbonate is not so stable to heat. This is because lithium carbonate is covalent. Lithium ion, being very small in size, polarizes a large carbonate ion, leading to the formation of more stable lithium oxide.

Therefore, lithium carbonate decomposes at a low temperature while a stable sodium carbonate decomposes at a high temperature.

(i) Nitrates

Thermal stability

Nitrates of alkali metals, except LiNO₃, decompose on strong heating to form nitrites.

LiNO₃, on decomposition, gives oxide.

Similar to lithium nitrate, alkaline earth metal nitrates also decompose to give oxides.

As we move down group 1 and group 2, the thermal stability of nitrate increases.

Solubility

Nitrates of both group 1 and group 2 metals are soluble in water.

(ii) Carbonates

Thermal stability

The carbonates of alkali metals are stable towards heat. However, carbonate of lithium, when heated, decomposes to form lithium oxide. The carbonates of alkaline earth metals also decompose on heating to form oxide and carbon dioxide.

Solubility

Carbonates of alkali metals are soluble in water with the exception of Li₂CO₃. Also, the solubility increases as we move down the group.

Carbonates of alkaline earth metals are insoluble in water.

(iii) Sulphates

Thermal stability

Sulphates of both group 1 and group 2 metals are stable towards heat.

Solubility

Sulphates of alkali metals are soluble in water. However, sulphates of alkaline earth metals show varied trends.

BeSO₄ Fairly soluble

MgSO₄ Soluble

CaSO₄ Sparingly soluble

SrSO₄ Insoluble

BaSO₄ Insoluble

In other words, while moving down the alkaline earth metals, the solubility of their sulphates decreases.

(a) Sodium can be extracted from sodium chloride by Downs process.

This process involves the electrolysis of fused NaCl (40%) and CaCl₂ (60 %) at a temperature of 1123 K in Downs cell.

Steel is the cathode and a block of graphite acts as the anode. Metallic Na and Ca are formed at cathode. Molten sodium is taken out of the cell and collected over kerosene.

(ii) Sodium hydroxide can be prepared by the electrolysis of sodium chloride. This is called Castner–Kellner process. In this process, the brine solution is electrolysed using a carbon anode and a mercury cathode.

The sodium metal, which is discharged at cathode, combines with mercury to form an amalgam.

(iii) Sodium peroxide

First, NaCl is electrolysed to result in the formation of Na metal (Downs process).

This sodium metal is then heated on aluminium trays in air (free of CO₂) to form its peroxide.

(iv) Sodium carbonate is prepared by Solvay process. Sodium hydrogen carbonate is precipitated in a reaction of sodium chloride and ammonium hydrogen carbonate.

These sodium hydrogen carbonate crystals are heated to give sodium carbonate.

(i) Magnesium burns in air with a dazzling light to form MgO and Mg₃N₂.

(ii) Quick lime (CaO) combines with silica (SiO₂) to form slag.

(iii) When chloride is added to slaked lime, it gives bleaching powder.

(iv) Calcium nitrate, on heating, decomposes to give calcium oxide.

(a) Structure of BeCl₂ (solid)

BeCl₂ exists as a polymer in condensed (solid) phase.

In the vapour state, BeCl₂ exists as a monomer with a linear structure.

The atomic size of sodium and potassium is larger than that of magnesium and calcium. Thus, the lattice energies of carbonates and hydroxides formed by calcium and magnesium are much more than those of sodium and potassium. Hence, carbonates and hydroxides of sodium and potassium dissolve readily in water whereas those of calcium and magnesium are only sparingly soluble.

(i) Chemically, limestone is CaCO₃.

Importance of limestone

(a) It is used in the preparation of lime and cement.

(b) It is used as a flux during the smelting of iron ores.

(ii) Chemically, cement is a mixture of calcium silicate and calcium aluminate.

Importance of cement

(a) It is used in plastering and in construction of bridges.

(b) It is used in concrete.

(iii) Chemically, plaster of Paris is 2CaSO₄.H₂O.

Importance of plaster of Paris

(a) It is used in surgical bandages.

(b) It is also used for making casts and moulds.

Lithium is the smallest in size among the alkali metals. Hence, Li⁺ ion can polarize water molecules more easily than other alkali metals. As a result, water molecules get attached to lithium salts as water of crystallization. Hence, lithium salts such as trihydrated lithium chloride (LiCl.3H₂O) are commonly hydrated. As the size of the ions increases, their polarizing power decreases. Hence, other alkali metal ions usually form anhydrous salts.

Importance of sodium, potassium, magnesium, and calcium in biological fluids:

(i) Sodium (Na):

Sodium ions are found primarily in the blood plasma. They are also found in the interstitial fluids surrounding the cells.

(a) Sodium ions help in the transmission of nerve signals.

(b) They help in regulating the flow of water across the cell membranes.

(c) They also help in transporting sugars and amino acids into the cells.

(ii) Potassium (K):

Potassium ions are found in the highest quantity within the cell fluids.

(a) K ions help in activating many enzymes.

(b) They also participate in oxidising glucose to produce ATP.

(c) They also help in transmitting nerve signals.

(iii) Magnesium (Mg) and calcium (Ca):

Magnesium and calcium are referred to as macro-minerals. This term indicates their higher abundance in the human body system.

(a) Mg helps in relaxing nerves and muscles.

(b) Mg helps in building and strengthening bones.

(c) Mg maintains normal blood circulation in the human body system.

(d) Ca helps in the coagulation of blood

(e) Ca also helps in maintaining homeostasis.

(a) On moving down the alkali group, the ionic and atomic sizes of the metals increase. The given alkali metal ions can be arranged in the increasing order of their ionic sizes as:

Li^{+ + + + +}

Smaller the size of an ion, the more highly is it hydrated. Since Li⁺ is the smallest, it gets heavily hydrated in an aqueous solution. On the other hand, Cs⁺ is the largest and so it is the least hydrated. The given alkali metal ions can be arranged in the decreasing order of their hydrations as:

Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺

Greater the mass of a hydrated ion, the lower is its ionic mobility. Therefore, hydrated Li⁺ is the least mobile and hydrated Cs⁺ is the most mobile. Thus, the given alkali metal ions can be arranged in the increasing order of their mobilities as:

Li^{+ + + + +}

(b) Unlike the other elements of group 1, Li reacts directly with nitrogen to form lithium nitride. This is because Li⁺ is very small in size and so its size is the most compatible with the N^3– ion. Hence, the lattice energy released is very high. This energy also overcomes the high amount of energy required for the formation of the N^3–ion.

(c) Electrode potential (E°) of any M²⁺/M electrode depends upon three factors:

(i) Ionisation enthalpy

(ii) Enthalpy of hydration

(iii) Enthalpy of vaporisation

The combined effect of these factors is approximately the same for Ca, Sr, and Ba. Hence, their electrode potentials are nearly constant.