Alkali metals Questions in English

(N/A) The alkali metal halides,$MX$ (where $X = F, Cl, Br, I$),are all high-melting,colourless crystalline solids.
They can be prepared by the reaction of the appropriate oxide,hydroxide,or carbonate with aqueous hydrohalic acid $(HX)$.
All of these halides have high negative enthalpies of formation $(\Delta_{f} H^{\ominus})$.
The $\Delta_{f} H^{\ominus}$ values for fluorides become less negative as we go down the group,while the reverse is true for $\Delta_{f} H^{\ominus}$ for chlorides,bromides,and iodides.
For a given metal,$\Delta_{f} H^{\ominus}$ always becomes less negative from fluoride to iodide.
The melting and boiling points follow the trend: $Fluoride > Chloride > Bromide > Iodide$.
All these halides are soluble in water.
The low solubility of $LiF$ in water is due to its high lattice enthalpy,whereas the low solubility of $CsI$ is due to the smaller hydration enthalpy of its two ions.
Other halides of lithium are soluble in ethanol,acetone,and ethyl acetate; $LiCl$ is also soluble in pyridine.

(N/A) Oxo-acids are compounds in which the acidic proton is on a hydroxyl group with an oxo group attached to the same atom.
Examples include carbonic acid,$H_{2}CO_{3}$ $(OC(OH)_{2})$,and sulphuric acid,$H_{2}SO_{4}$ $(O_{2}S(OH)_{2})$.
Alkali metals form salts with all oxo-acids. These salts are generally soluble in water and thermally stable.
Their carbonates $(M_{2}CO_{3})$ and,in most cases,their hydrogen carbonates $(MHCO_{3})$ are highly stable to heat.
As the electropositive character increases down the group,the stability of the carbonates and hydrogen carbonates increases.
Lithium carbonate is not thermally stable because the small $Li^{+}$ ion polarizes the large $CO_{3}^{2-}$ ion,leading to the formation of $Li_{2}O$ and $CO_{2}$. Additionally,lithium hydrogen carbonate does not exist as a solid.

(N/A) $(i)$ Lithium is much harder. Its melting point $(m.p.)$ and boiling point $(b.p.)$ are higher than the other alkali metals.
$(ii)$ Lithium is the least reactive but the strongest reducing agent among all the alkali metals. On combustion in air,it forms mainly monoxide,$Li_{2}O$,and the nitride,$Li_{3}N$,unlike other alkali metals.
$(iii)$ $LiCl$ is deliquescent and crystallises as a hydrate,$(LiCl \cdot 2H_{2}O)$,whereas other alkali metal chlorides do not form hydrates.
$(iv)$ Lithium hydrogen carbonate is not obtained in the solid form,while all other elements form solid hydrogen carbonates.
$(v)$ Lithium,unlike other alkali metals,forms no ethynide on reaction with ethyne.
$(vi)$ Lithium nitrate when heated gives lithium oxide $(Li_{2}O)$,whereas other alkali metal nitrates decompose to give the corresponding nitrite.
$4LiNO_{3} \rightarrow 2Li_{2}O + 4NO_{2} + O_{2}$
$2NaNO_{3} \rightarrow 2NaNO_{2} + O_{2}$
$(vii)$ $LiF$ and $Li_{2}O$ are comparatively much less soluble in water than the corresponding compounds of other alkali metals.

(N/A) $(i)$ Lithium is much harder. Its melting point $(m.p.)$ and boiling point $(b.p.)$ are higher than the other alkali metals.
$(ii)$ Lithium is the least reactive but the strongest reducing agent among all the alkali metals. On combustion in air,it forms mainly monoxide,$Li_{2}O$,and the nitride,$Li_{3}N$,unlike other alkali metals.
$(iii)$ $LiCl$ is deliquescent and crystallises as a hydrate,$(LiCl \cdot 2H_{2}O)$,whereas other alkali metal chlorides do not form hydrates.
$(iv)$ Lithium hydrogen carbonate is not obtained in the solid form,while all other elements form solid hydrogen carbonates.
$(v)$ Lithium,unlike other alkali metals,forms no ethynide on reaction with ethyne.
$(vi)$ Lithium nitrate when heated gives lithium oxide $(Li_{2}O)$,whereas other alkali metal nitrates decompose to give the corresponding nitrite.
$4LiNO_{3} \rightarrow 2Li_{2}O + 4NO_{2} + O_{2}$
$2NaNO_{3} \rightarrow 2NaNO_{2} + O_{2}$
$(vii)$ $LiF$ and $Li_{2}O$ are comparatively much less soluble in water than the corresponding compounds of other alkali metals.

(N/A) $(i)$ Lithium is much harder. Its melting point $(m.p.)$ and boiling point $(b.p.)$ are higher than the other alkali metals.
$(ii)$ Lithium is the least reactive but the strongest reducing agent among all the alkali metals. On combustion in air,it forms mainly monoxide,$Li_{2}O$,and the nitride,$Li_{3}N$,unlike other alkali metals.
$(iii)$ $LiCl$ is deliquescent and crystallises as a hydrate,$(LiCl \cdot 2H_{2}O)$,whereas other alkali metal chlorides do not form hydrates.
$(iv)$ Lithium hydrogen carbonate is not obtained in the solid form,while all other elements form solid hydrogen carbonates.
$(v)$ Lithium,unlike other alkali metals,forms no ethynide on reaction with ethyne.
$(vi)$ Lithium nitrate when heated gives lithium oxide $(Li_{2}O)$,whereas other alkali metal nitrates decompose to give the corresponding nitrite.
$4LiNO_{3} \rightarrow 2Li_{2}O + 4NO_{2} + O_{2}$
$2NaNO_{3} \rightarrow 2NaNO_{2} + O_{2}$
$(vii)$ $LiF$ and $Li_{2}O$ are comparatively much less soluble in water than the corresponding compounds of other alkali metals.

(N/A) $i$. Lithium is much harder. Its $m.p.$ and $b.p.$ are higher than the other alkali metals.
$ii$. Lithium is the least reactive but the strongest reducing agent among all the alkali metals. On combustion in air,it forms mainly monoxide,$Li_2O$,and the nitride,$Li_3N$,unlike other alkali metals.
$iii$. $LiCl$ is deliquescent and crystallises as a hydrate,$(LiCl \cdot 2H_2O)$,whereas other alkali metal chlorides do not form hydrates.
$iv$. Lithium hydrogen carbonate is not obtained in the solid form,while all other elements form solid hydrogen carbonates.
$v$. Lithium,unlike other alkali metals,forms no ethynide on reaction with ethyne.
$vi$. Lithium nitrate when heated gives lithium oxide $(Li_2O)$,whereas other alkali metal nitrates decompose to give the corresponding nitrite.
$4LiNO_3 \rightarrow 2Li_2O + 4NO_2 + O_2$
$2NaNO_3 \rightarrow 2NaNO_2 + O_2$
$vii$. $LiF$ and $Li_2O$ are comparatively much less soluble in water than the corresponding compounds of other alkali metals.

(N/A) $(i)$ Lithium is much harder. Its melting point $(m.p.)$ and boiling point $(b.p.)$ are higher than the other alkali metals.
$(ii)$ Lithium is the least reactive but the strongest reducing agent among all the alkali metals. On combustion in air,it forms mainly monoxide,$Li_{2}O$,and the nitride,$Li_{3}N$,unlike other alkali metals.
$(iii)$ $LiCl$ is deliquescent and crystallises as a hydrate,$(LiCl \cdot 2H_{2}O)$,whereas other alkali metal chlorides do not form hydrates.
$(iv)$ Lithium hydrogen carbonate is not obtained in the solid form,while all other elements form solid hydrogen carbonates.
$(v)$ Lithium,unlike other alkali metals,forms no ethynide on reaction with ethyne.
$(vi)$ Lithium nitrate,when heated,gives lithium oxide $(Li_{2}O)$,whereas other alkali metal nitrates decompose to give the corresponding nitrite.
$4LiNO_{3} \rightarrow 2Li_{2}O + 4NO_{2} + O_{2}$
$2NaNO_{3} \rightarrow 2NaNO_{2} + O_{2}$
$(vii)$ $LiF$ and $Li_{2}O$ are comparatively much less soluble in water than the corresponding compounds of other alkali metals.

(N/A) The similarity between lithium and magnesium is particularly striking and arises because of their similar sizes.
Atomic radii: $Li \ (152 \ pm); Mg \ (160 \ pm)$
Ionic radii: $Li^{+} \ (76 \ pm); Mg^{2+} \ (72 \ pm)$
The main points of similarity are:
$(i)$ Both lithium and magnesium are harder and lighter than other elements in their respective groups.
$(ii)$ 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,$Li_{3}N$ and $Mg_{3}N_{2}$,by direct combination with nitrogen.
$(iii)$ The oxides,$Li_{2}O$ 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 $CO_{2}$. Solid hydrogen carbonates are not formed by lithium and magnesium.
$(v)$ Both $LiCl$ and $MgCl_{2}$ are soluble in ethanol.
$(vi)$ Both $LiCl$ and $MgCl_{2}$ are deliquescent and crystallise from aqueous solution as hydrates,$LiCl \cdot 2H_{2}O$ and $MgCl_{2} \cdot 8H_{2}O$.

(N/A) The similarity between lithium and magnesium is particularly striking and arises because of their similar sizes.
Atomic radii: $Li$ $(152 \ pm)$; $Mg$ $(160 \ pm)$
Ionic radii: $Li^{+}$ $(76 \ pm)$; $Mg^{2+}$ $(72 \ pm)$
The main points of similarity are:
$(i)$ Both lithium and magnesium are harder and lighter than other elements in their respective groups.
$(ii)$ 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, $Li_{3}N$ and $Mg_{3}N_{2}$, by direct combination with nitrogen.
$(iii)$ The oxides, $Li_{2}O$ 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 $CO_{2}$. Solid hydrogen carbonates are not formed by lithium and magnesium.
$(v)$ Both $LiCl$ and $MgCl_{2}$ are soluble in ethanol.
$(vi)$ Both $LiCl$ and $MgCl_{2}$ are deliquescent and crystallise from aqueous solution as hydrates, $LiCl \cdot 2H_{2}O$ and $MgCl_{2} \cdot 8H_{2}O$.

(N/A) The similarity between lithium and magnesium is particularly striking and arises because of their similar ionic sizes and charge-to-size ratios (ionic potential).
Atomic radii: $Li$ $(152 \ pm)$; $Mg$ $(160 \ pm)$
Ionic radii: $Li^{+}$ $(76 \ pm)$; $Mg^{2+}$ $(72 \ pm)$
The main points of similarity are:
$(i)$ Both lithium and magnesium are harder and lighter than other elements in their respective groups.
$(ii)$ 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, $Li_{3}N$ and $Mg_{3}N_{2}$, by direct combination with nitrogen.
$(iii)$ The oxides, $Li_{2}O$ 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 $CO_{2}$. Solid hydrogen carbonates are not formed by lithium and magnesium.
$(v)$ Both $LiCl$ and $MgCl_{2}$ are soluble in ethanol.
$(vi)$ Both $LiCl$ and $MgCl_{2}$ are deliquescent and crystallise from aqueous solution as hydrates, $LiCl \cdot 2H_{2}O$ and $MgCl_{2} \cdot 8H_{2}O$.

(N/A) Among alkali metal ions,$Li^{+}$ has the smallest size,due to which it has a very high charge density and can easily polarize water molecules. This leads to strong ion-dipole interactions,causing $Li$ salts to be commonly hydrated (e.g.,$LiCl \cdot 2H_2O$). The other alkali metal ions $(Na^{+}, K^{+}, Rb^{+}, Cs^{+})$ are larger in size,resulting in lower charge density and a weaker ability to polarize water molecules. Consequently,their salts are usually anhydrous.

(N/A) $LiF$ is almost insoluble in water due to its very high lattice energy. The fluoride ion $(F^-)$ is very small,leading to a strong electrostatic attraction with the $Li^+$ ion,resulting in a high lattice energy that is not compensated by the hydration energy.
In contrast,$LiCl$ has a lower lattice energy because the chloride ion $(Cl^-)$ is larger. Additionally,$LiCl$ possesses significant covalent character due to the polarizing power of the small $Li^+$ ion (Fajans' rule),which makes it soluble in organic solvents like acetone.
Therefore,the solubility of $LiCl$ in water is driven by favorable hydration energy,while its solubility in acetone is attributed to its covalent nature.

(N/A) The Solvay process is used for the industrial preparation of sodium carbonate $(Na_{2}CO_{3})$. The process relies on the low solubility of sodium hydrogen carbonate $(NaHCO_{3})$,which precipitates when sodium chloride $(NaCl)$ reacts with ammonium hydrogen carbonate $(NH_{4}HCO_{3})$.
$1$. Ammonia is dissolved in a concentrated solution of sodium chloride,and $CO_{2}$ is passed through it to form ammonium carbonate,which then converts to ammonium hydrogen carbonate:
$2NH_{3} + H_{2}O + CO_{2} \rightarrow (NH_{4})_{2}CO_{3}$
$(NH_{4})_{2}CO_{3} + H_{2}O + CO_{2} \rightarrow 2NH_{4}HCO_{3}$
$2$. The $NH_{4}HCO_{3}$ reacts with $NaCl$ to precipitate $NaHCO_{3}$:
$NH_{4}HCO_{3} + NaCl \rightarrow NH_{4}Cl + NaHCO_{3}$
$3$. The $NaHCO_{3}$ crystals are filtered and heated to produce sodium carbonate:
$2NaHCO_{3} \rightarrow Na_{2}CO_{3} + CO_{2} + H_{2}O$
$4$. Ammonia is recovered by treating the $NH_{4}Cl$ solution with calcium hydroxide $(Ca(OH)_{2})$,producing calcium chloride $(CaCl_{2})$ as a by-product:
$2NH_{4}Cl + Ca(OH)_{2} \rightarrow 2NH_{3} + CaCl_{2} + 2H_{2}O$

(N/A) Sodium carbonate $(Na_{2}CO_{3} \cdot 10H_{2}O)$ is generally prepared by the Solvay process. In this process,the low solubility of sodium hydrogen carbonate is utilized,causing it to precipitate during the reaction of sodium chloride with ammonium hydrogen carbonate.
The ammonium hydrogen carbonate is prepared by passing $CO_{2}$ through a concentrated solution of sodium chloride saturated with ammonia,where ammonium carbonate is formed first,followed by ammonium hydrogen carbonate. The equations for the complete process are as follows:
$2NH_{3} + H_{2}O + CO_{2} \rightarrow (NH_{4})_{2}CO_{3}$
$(NH_{4})_{2}CO_{3} + H_{2}O + CO_{2} \rightarrow 2NH_{4}HCO_{3}$
$NH_{4}HCO_{3} + NaCl \rightarrow NH_{4}Cl + NaHCO_{3}$
The sodium hydrogen carbonate crystals separate out. These are then heated to produce sodium carbonate:
$2NaHCO_{3} \rightarrow Na_{2}CO_{3} + CO_{2} + H_{2}O$
In this process,$NH_{3}$ is recovered when the solution containing $NH_{4}Cl$ is treated with $Ca(OH)_{2}$. Calcium chloride is obtained as a by-product:
$2NH_{4}Cl + Ca(OH)_{2} \rightarrow 2NH_{3} + CaCl_{2} + H_{2}O$

(N/A) The Solvay process involves the reaction of $NaCl$ with $NH_3$ and $CO_2$ to form $NaHCO_3$,which precipitates out because it is sparingly soluble in water. In the case of potassium,$KHCO_3$ is formed,but it is fairly soluble in water and does not precipitate out. Therefore,the process cannot be used to prepare $K_2CO_3$.

Sodium carbonate is a white crystalline solid which exists as a decahydrate,$Na_{2}CO_{3} \cdot 10H_{2}O$. This is also called washing soda.
It is readily soluble in water.
On heating,the decahydrate loses its water of crystallization to form monohydrate.
Above $373 \ K$,the monohydrate becomes completely anhydrous and changes to a white powder called soda ash.
$Na_{2}CO_{3} \cdot 10H_{2}O \xrightarrow{375 \ K} Na_{2}CO_{3} \cdot H_{2}O + 9H_{2}O$
$Na_{2}CO_{3} \cdot H_{2}O \xrightarrow{>373 \ K} Na_{2}CO_{3} + H_{2}O$
The carbonate part of sodium carbonate gets hydrolysed by water to form an alkaline solution: $CO_{3}^{2-} + H_{2}O \rightarrow HCO_{3}^{-} + OH^{-}$

(N/A) $(i)$ It is used in water softening,laundering,and cleaning.
$(ii)$ It is used in the manufacture of glass,soap,borax,and caustic soda.
$(iii)$ It is used in paper,paints,and textile industries.
$(iv)$ It is an important laboratory reagent in both qualitative and quantitative analysis.

(N/A) As we move down the alkali metal group,the electropositive character increases,which leads to an increase in the thermal stability of alkali metal carbonates.
However,$Li_{2}CO_{3}$ is relatively unstable to heat because the $Li^+$ ion is very small in size and has a high polarizing power. It polarizes the large $CO_{3}^{2-}$ ion,leading to the formation of stable $Li_{2}O$ and $CO_{2}$ gas.
$Li_{2}CO_{3} \xrightarrow{\Delta} Li_{2}O + CO_{2}$
In contrast,$Na_{2}CO_{3}$ is more stable due to the larger size of the $Na^+$ ion,which has a lower polarizing power,thus requiring a much higher temperature for decomposition.

(B) The most abundant source of sodium chloride is sea water,which contains $2.7$ to $2.9 \%$ by mass of the salt. Common salt is generally obtained by the evaporation of sea water.
Crude sodium chloride,obtained by the crystallization of brine solution,contains impurities like sodium sulphate,calcium sulphate,calcium chloride,and magnesium chloride. Calcium chloride $(CaCl_2)$ and magnesium chloride $(MgCl_2)$ are impurities because they are deliquescent (they absorb moisture easily from the atmosphere).
To obtain pure sodium chloride,the crude salt is dissolved in a minimum amount of water and filtered to remove insoluble impurities. The solution is then saturated with hydrogen chloride $(HCl)$ gas. Crystals of pure sodium chloride separate out,while calcium and magnesium chlorides,being more soluble than sodium chloride,remain in the solution.

(N/A) Preparation: The most abundant source of sodium chloride is sea water,which contains $2.7$ to $2.9 \%$ by mass of the salt. Common salt is generally obtained by the evaporation of sea water.
Crude sodium chloride,obtained by the crystallization of brine solution,contains impurities like sodium sulphate,calcium sulphate,calcium chloride,and magnesium chloride. Calcium chloride $(CaCl_{2})$ and magnesium chloride $(MgCl_{2})$ are impurities because they are deliquescent (absorb moisture easily from the atmosphere).
To obtain pure sodium chloride,the crude salt is dissolved in a minimum amount of water and filtered to remove insoluble impurities. The solution is then saturated with hydrogen chloride gas. Crystals of pure sodium chloride separate out,while calcium and magnesium chlorides,being more soluble than sodium chloride,remain in the solution.
Properties: Sodium chloride melts at $1081 \ K$. It has a solubility of $36.0 \ g$ in $100 \ g$ of water at $273 \ K$. The solubility does not increase appreciably with an increase in temperature.
Uses: $(i)$ It is used as a common salt or table salt for domestic purposes. $(ii)$ It is used for the preparation of $Na_{2}O_{2}$,$NaOH$,and $Na_{2}CO_{3}$.

(N/A) Sodium hydroxide is generally prepared commercially by the electrolysis of sodium chloride in a $Castner-Kellner$ cell. $A$ brine solution is electrolysed using a mercury cathode and a carbon anode. Sodium metal discharged at the cathode combines with mercury to form sodium amalgam. Chlorine gas is evolved at the anode.
Cathode: $Na^{+} + e^{-} \xrightarrow{Hg} Na-\text{amalgam}$
Anode: $Cl^{-} \rightarrow \frac{1}{2} Cl_{2} + e^{-}$
The amalgam is treated with water to give sodium hydroxide and hydrogen gas.
$2Na-\text{amalgam} + 2H_{2}O \rightarrow 2NaOH + 2Hg + H_{2}$
Properties: Sodium hydroxide is a white,translucent solid. It melts at $591 \ K$. It is readily soluble in water to give a strong alkaline solution. Crystals of sodium hydroxide are deliquescent.
The sodium hydroxide solution at the surface reacts with the $CO_{2}$ in the atmosphere to form $Na_{2}CO_{3}$.
Uses: $(i)$ The manufacture of soap,paper,artificial silk and a number of chemicals,$(ii)$ In petroleum refining,$(iii)$ In the purification of bauxite,$(iv)$ In the textile industries for mercerising cotton fabrics,$(v)$ For the preparation of pure fats and oils,and $(vi)$ As a laboratory reagent.

(N/A) Sodium hydroxide is prepared commercially by the electrolysis of brine ($NaCl$ solution) in the Castner-Kellner cell. $A$ mercury cathode and a carbon anode are used.
At the cathode,sodium ions are reduced to sodium metal,which dissolves in mercury to form sodium amalgam $(Na-Hg)$:
$Na^{+} + e^{-} \xrightarrow{Hg} Na-Hg$
At the anode,chloride ions are oxidized to chlorine gas:
$Cl^{-} \rightarrow \frac{1}{2} Cl_{2} + e^{-}$
The sodium amalgam is then treated with water to produce sodium hydroxide and hydrogen gas:
$2Na-Hg + 2H_{2}O \rightarrow 2NaOH + 2Hg + H_{2}$
Properties: Sodium hydroxide is a white,translucent solid with a melting point of $591 \ K$. It is highly soluble in water,forming a strong alkaline solution. Its crystals are deliquescent and react with atmospheric $CO_{2}$ to form $Na_{2}CO_{3}$.
Uses: $(i)$ Manufacture of soap,paper,and artificial silk,$(ii)$ Petroleum refining,$(iii)$ Purification of bauxite,$(iv)$ Mercerising cotton fabrics,$(v)$ Preparation of pure fats and oils,and $(vi)$ As a laboratory reagent.

(N/A) Preparation: Sodium hydrogen carbonate is prepared by saturating a solution of sodium carbonate with carbon dioxide. The white crystalline powder of sodium hydrogen carbonate,being less soluble,precipitates out.
$Na_2CO_3 + H_2O + CO_2 \rightarrow 2NaHCO_3$
Uses:
$(i)$ It is used as a mild antiseptic for skin infections.
$(ii)$ It is used in fire extinguishers.
$(iii)$ It is used as an ingredient in baking powder to make cakes and pastries light and fluffy due to the release of $CO_2$ gas upon heating.

(N/A) $(i)$ Sodium metal: Sodium is extracted from fused $NaCl$ by the Downs process. The electrolyte is a mixture of $NaCl$ $(40 \%)$ and $CaCl_2$ $(60 \%)$ at $1123 \ K$. Steel is the cathode and graphite is the anode. $Na^{+} + e^- \rightarrow Na$ (at cathode); $2Cl^- \rightarrow Cl_2 + 2e^-$ (at anode).
$(ii)$ Sodium hydroxide: It is prepared by the electrolysis of an aqueous solution of $NaCl$ (brine) in a Castner-Kellner cell. $2NaCl + 2H_2O \rightarrow 2NaOH + Cl_2 + H_2$.
$(iii)$ Sodium peroxide: Sodium metal obtained from the Downs process is heated on aluminium trays in an atmosphere of $CO_2$-free air to form sodium peroxide. $2Na + O_2 \rightarrow Na_2O_2$.
$(iv)$ Sodium carbonate: It is prepared by the Solvay process. $2NH_3 + H_2O + CO_2 \rightarrow (NH_4)_2CO_3$; $(NH_4)_2CO_3 + H_2O + CO_2 \rightarrow 2NH_4HCO_3$; $NH_4HCO_3 + NaCl \rightarrow NH_4Cl + NaHCO_3$; $2NaHCO_3 \rightarrow Na_2CO_3 + H_2O + CO_2$.

(N/A) typical $70 \ kg$ man contains about $90 \ g$ of $Na$ and $170 \ g$ of $K$ compared with only $5 \ g$ of iron and $0.06 \ g$ of copper.
Sodium ions $(Na^+)$ are found primarily on the outside of cells,being located in blood plasma and in the interstitial fluid which surrounds the cells.
These ions participate in the transmission of nerve signals,in regulating the flow of water across cell membranes,and in the transport of sugars and amino acids into cells.
Potassium ions $(K^+)$ are the most abundant cations within cell fluids,where they activate many enzymes,participate in the oxidation of glucose to produce $ATP$,and,with sodium,are responsible for the transmission of nerve signals.
There is a considerable variation in the concentration of $Na^+$ and $K^+$ ions across cell membranes. In blood plasma,$Na^+$ is present at $143 \ mmol \ L^{-1}$,whereas $K^+$ is only $5 \ mmol \ L^{-1}$. Inside red blood cells,these concentrations change to $10 \ mmol \ L^{-1}$ for $Na^+$ and $105 \ mmol \ L^{-1}$ for $K^+$.
These ionic gradients are maintained by a discriminatory mechanism called the sodium-potassium pump,which consumes more than one-third of the $ATP$ used by a resting animal.

(N/A) typical $70 \ kg$ human contains about $90 \ g$ of $Na$ and $170 \ g$ of $K$,compared with only $5 \ g$ of iron and $0.06 \ g$ of copper. Sodium ions are found primarily on the outside of cells,located in blood plasma and in the interstitial fluid surrounding the cells.
These ions participate in the transmission of nerve signals,in regulating the flow of water across cell membranes,and in the transport of sugars and amino acids into cells.
Sodium and potassium,although chemically similar,differ quantitatively in their ability to penetrate cell membranes,in their transport mechanisms,and in their efficiency to activate enzymes. Potassium ions are the most abundant cations within cell fluids,where they activate many enzymes,participate in the oxidation of glucose to produce $ATP$,and,along with sodium,are responsible for the transmission of nerve signals.
There is a considerable variation in the concentration of sodium and potassium ions on opposite sides of cell membranes. For example,in blood plasma,sodium is present at $143 \ mmol \ L^{-1}$,whereas the potassium level is only $5 \ mmol \ L^{-1}$. Within red blood cells,these concentrations change to $10 \ mmol \ L^{-1}$ for $Na^{+}$ and $105 \ mmol \ L^{-1}$ for $K^{+}$.
These ionic gradients demonstrate that a discriminatory mechanism,called the sodium-potassium pump,operates across cell membranes. This pump consumes more than one-third of the $ATP$ used by a resting animal and about $15 \ kg$ of $ATP$ per $24 \ h$ in a resting human.

(N/A) The reactivity of alkali metals increases down the group as the ionization enthalpy decreases.
$Sodium$ $(Na)$ is in the $3rd$ period,while $potassium$ $(K)$ is in the $4th$ period.
Due to the larger atomic size of $potassium$,the valence electron is further away from the nucleus compared to $sodium$.
Consequently,the force of attraction between the nucleus and the valence electron is weaker in $potassium$,making it easier to lose the electron.
Therefore,$potassium$ has a lower ionization enthalpy than $sodium$,which makes $potassium$ more reactive than $sodium$.

(N/A) $(i)$ When $Na$ metal is dropped in water,it reacts violently to form sodium hydroxide and hydrogen gas. The chemical equation involved in the reaction is: $2Na_{(s)} + 2H_{2}O_{(l)} \rightarrow 2NaOH_{(aq)} + H_{2(g)}$
$(ii)$ On being heated in air,sodium reacts vigorously with oxygen to form sodium peroxide. The chemical equation involved in the reaction is: $2Na_{(s)} + O_{2(g)} \rightarrow Na_{2}O_{2(s)}$
$(iii)$ When sodium peroxide is dissolved in water,it is readily hydrolysed to form sodium hydroxide and hydrogen peroxide. The chemical equation involved in the reaction is: $Na_{2}O_{2(s)} + 2H_{2}O_{(l)} \rightarrow 2NaOH_{(aq)} + H_{2}O_{2(aq)}$

(A) When $Na_2CO_3$ is added to water,it undergoes hydrolysis to produce $NaHCO_3$ and $NaOH$. Since $NaOH$ is a strong base,the resulting solution becomes alkaline.
$Na_2CO_3 + H_2O \rightarrow NaHCO_3 + NaOH$
$(b)$ Alkali metals are strong reducing agents and highly electropositive,making their extraction by chemical reduction of oxides or displacement reactions impossible. Electrolysis of aqueous solutions is also not feasible because the liberated metals react with water. Therefore,they are prepared by the electrolysis of their fused chlorides.
$(c)$ $Na^+$ ions are primarily found in blood plasma and interstitial fluids,where they regulate nerve signal transmission,water flow across cell membranes,and the transport of sugars and amino acids. While $K^+$ ions are essential within cells,$Na^+$ ions perform these critical extracellular functions,making them highly significant.

$(a)$ Reaction of $Na_{2}O_{2}$ with water:
$2Na_{2}O_{2(s)} + 2H_{2}O_{(l)} \rightarrow 4NaOH_{(aq)} + O_{2(g)}$
$(b)$ Reaction of $KO_{2}$ with water:
$2KO_{2(s)} + 2H_{2}O_{(l)} \rightarrow 2KOH_{(aq)} + H_{2}O_{2(aq)} + O_{2(g)}$
$(c)$ Reaction of $Na_{2}O$ with $CO_{2}$:
$Na_{2}O_{(s)} + CO_{2(g)} \rightarrow Na_{2}CO_{3(s)}$

(N/A) The stability of superoxides increases with an increase in the size of the alkali metal cation due to the increase in lattice enthalpy.
$Lithium$ $(Li^+)$ has a very small ionic radius,which results in a lower lattice enthalpy for $LiO_2$ compared to the superoxides of larger alkali metals.
Therefore,$LiO_2$ is thermodynamically unstable and does not exist under normal conditions.

(N/A) The solubility of ionic compounds in water is determined by the balance between lattice energy and hydration energy.
For $Na^+$ and $K^+$ salts,the lattice energy is relatively low due to the larger size of the cations,which is easily overcome by the hydration energy released upon dissolution.
In contrast,$Mg^{2+}$ and $Ca^{2+}$ are smaller and carry a higher charge $(+2)$,resulting in significantly higher lattice energies for their carbonates and hydroxides.
Since the lattice energy of $Mg$ and $Ca$ salts is much higher than the hydration energy,they are only sparingly soluble in water.

(N/A) Sodium $(Na)$: Sodium ions are found primarily in the blood plasma and in the interstitial fluids surrounding the cells.
Uses: $(i)$ Sodium ions help in the transmission of nerve signals. $(ii)$ They help in regulating the flow of water across the cell membranes. $(iii)$ They also help in transporting sugars and amino acids into the cells.
$(b)$ Potassium $(K)$: Potassium ions are found in the highest quantity within the cell fluids.
Uses: $(i)$ $K^+$ ions help in activating many enzymes. $(ii)$ They participate in oxidizing glucose to produce $ATP$. $(iii)$ They also help in transmitting nerve signals.
$(c)$ Magnesium $(Mg)$: Magnesium is a macro-mineral,indicating its high abundance in the human body.
Uses: $(i)$ $Mg^{2+}$ helps in relaxing nerves and muscles. $(ii)$ $Mg^{2+}$ helps in building and strengthening bones. $(iii)$ $Mg^{2+}$ maintains normal blood circulation.
$(d)$ Calcium $(Ca)$: Calcium is a macro-mineral.
Uses: $(i)$ $Ca^{2+}$ helps in the coagulation of blood. $(ii)$ $Ca^{2+}$ helps in maintaining homeostasis and muscle contraction.

(N/A) Physical properties of alkali metals are as follows:
$(i)$ They are quite soft and can be cut easily. Sodium metal can be easily cut using a knife.
$(ii)$ They are light coloured and are mostly silvery white in appearance.
$(iii)$ They have low density because of the large atomic sizes. The density increases down the group from $Li$ to $Cs$. The only exception to this is $K$,which has lower density than $Na$.
$(iv)$ The metallic bonding present in alkali metals is quite weak. Therefore,they have low melting and boiling points.
$(v)$ 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.
$(vi)$ They also display photoelectric effect. When metals such as $Cs$ and $K$ are irradiated with light,they lose electrons.
$(b)$ 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 hydroxides and dihydrogen gas. The general reaction is: $2M + 2H_2O \rightarrow 2MOH + H_2$
$(2)$ They react with dihydrogen to form metal hydrides. The general reaction is: $2M + H_2 \rightarrow 2M^+H^-$
$(3)$ Almost all alkali metals,except $Li$,react directly with halogens to form ionic halides. $2M + Cl_2 \rightarrow 2MCl$ (where $M = Na, K, Rb, Cs$). Since $Li^+$ ion is very small in size,it can easily distort the electron cloud around the negative halide ion,making lithium halides covalent in nature.
$(4)$ They are strong reducing agents. The reducing power of alkali metals generally increases on moving down the group. However,$Li$ is an exception as it is the strongest reducing agent due to its high hydration energy.

(N/A) In the process of chemical reduction,metal oxides are reduced using a stronger reducing agent.
Alkali metals and alkaline earth metals are themselves among the strongest reducing agents.
Since there are no available reducing agents stronger than these metals,their oxides cannot be reduced to metals by chemical reduction.
Therefore,these metals are typically obtained by the electrolytic reduction of their fused salts.

(N/A) All three elements,lithium $(Li)$,potassium $(K)$,and cesium $(Cs)$,are alkali metals.
However,$K$ and $Cs$ are used in photoelectric cells,whereas $Li$ is not.
This is because $Li$ has a smaller atomic size compared to $K$ and $Cs$,which results in a higher ionization enthalpy,making it difficult to remove an electron.
On the other hand,$K$ and $Cs$ have lower ionization enthalpies,allowing them to lose electrons easily when exposed to light.
This property of $K$ and $Cs$ is utilized in photoelectric cells.

(A) Lithium $(Li)$ and beryllium $(Be)$,the first elements of Group-$1$ and Group-$2$ respectively,exhibit properties that differ from those of the other members of their respective groups. This is due to their small size,high electronegativity,and high polarizing power. Lithium shows diagonal relationship similarities to magnesium $(Mg)$,and beryllium shows similarities to aluminium $(Al)$.

(A) Among the group-$1$ elements,$Li^+$ ion has the highest hydration enthalpy.
This is because the hydration enthalpy depends on the size of the ion.
Smaller ions have a higher charge density,which allows them to attract water molecules more strongly.
Since $Li^+$ is the smallest ion in group-$1$,it has the highest charge density and thus the highest hydration enthalpy.

(N/A) The alkali metals and their salts impart a characteristic color to an oxidizing flame. This is because the heat from the flame excites the outermost orbital electron to a higher energy level. When the excited electron returns to the ground state,it emits radiation in the visible region of the electromagnetic spectrum.

(C) The concentration of alkali metals can be determined by using instrumental techniques such as flame photometry or atomic absorption spectroscopy.
These methods rely on the characteristic emission or absorption of light by the metal ions in a flame.

(B) Caesium $(Cs)$ and potassium $(K)$ are useful as electrodes in photoelectric cells because they have low ionization enthalpies,allowing them to emit electrons when exposed to light.

(N/A) The reaction of alkali metals with dioxygen depends on the nature of the metal:
$1$. Lithium forms mainly the oxide: $4 Li + O_{2} \rightarrow 2 Li_{2}O$
$2$. Sodium forms mainly the peroxide: $2 Na + O_{2} \rightarrow Na_{2}O_{2}$
$3$. Potassium,Rubidium,and Cesium form superoxides: $M + O_{2} \rightarrow MO_{2}$ (where $M = K, Rb, Cs$)

(N/A) The general chemical reaction for the dissolution of an alkali metal $(M)$ in liquid ammonia is given by:
$M + (x+y) NH_{3} \longrightarrow [M(NH_{3})_{x}]^{+} + [e(NH_{3})_{y}]^{-}$
Here,the alkali metal dissolves in liquid ammonia to form a deep blue solution,which is conducting in nature due to the presence of ammoniated cations and ammoniated electrons.

(A) The correct matches are as follows:
$(a)$. Super oxide: $CsO_2$ (contains $O_2^-$ ion).
$(b)$. Peroxide: $Na_2O_2$ (contains $O_2^{2-}$ ion).
$(c)$. Dioxide: $CO_2$ (carbon dioxide).
$(d)$. Sub oxide: $C_3O_2$ (carbon suboxide).
Therefore,the correct sequence is $(a-s), (b-p), (c-q), (d-r)$.

(B) The correct matches are:
$(a)$ Borax is $Na_2B_4O_7.10H_2O$ $(s)$.
$(b)$ Carnallite is $KCl.MgCl_2.6H_2O$ $(r)$.
$(c)$ Rock salt is $NaCl$ $(p)$.
$(d)$ Silvine is $KCl$ $(q)$.
Therefore,the correct sequence is $(a-s), (b-r), (c-p), (d-q)$.

(A) The chemical compositions are as follows:
$(a).$ Spodumene is a lithium aluminium silicate,$LiAl(SiO_3)_2$,containing $Li$.
$(b).$ Borax is $Na_2B_4O_7 \cdot 10H_2O$,containing $Na$.
$(c).$ Sylvine is potassium chloride,$KCl$,containing $K$.
$(d).$ Chile saltpetre is sodium nitrate,$NaNO_3$,containing $Na$.
Therefore,the correct matching is $(a-r), (b-q), (c-p), (d-q)$.

(A) $1-c, 2-a, 3-b$.
$1$. Soda ash $(Na_2CO_3)$ is used in paper and textile industries.
$2$. Calcium oxide $(CaO)$ is used in the purification of sugar.
$3$. Baking soda $(NaHCO_3)$ is used as an antacid.

(A) The correct matches are:
$1$. Lime stone is $CaCO_3$ $(c)$.
$2$. Quick lime is $CaO$ $(a)$.
$3$. Washing soda is $Na_2CO_3 \cdot 10H_2O$ $(d)$.
$4$. Baking soda is $NaHCO_3$ $(b)$.
Therefore,the correct sequence is $1-c, 2-a, 3-d, 4-b$.

(A) The flame colors for alkali metals are as follows:
$1$. $Li$ (Lithium) gives a crimson red flame.
$2$. $K$ (Potassium) gives a violet flame.
$3$. $Cs$ (Cesium) gives a blue flame.
Therefore,the correct matching is $1-a, 2-b, 3-c$.

(C) The correct matches are as follows:
$1$. Lithium is found in Lepidolite $(c)$.
$2$. Barium is found in Witherite $(a)$.
$3$. Magnesium is found in Carnallite $(d)$.
$4$. Strontium is found in Celestine $(b)$.
Therefore,the correct sequence is $1-c, 2-a, 3-d, 4-b$.