Alkaline earth metals Questions in English

(N/A) The alkaline earth metals burn in oxygen to form the monoxide,$MO$,which,except for $BeO$,have a rock-salt structure.
The $BeO$ is essentially covalent in nature. The enthalpies of formation of these oxides are quite high,and consequently,they are very stable to heat.
$BeO$ is amphoteric,while the oxides of other elements are ionic in nature.

(N/A) The hydroxides of alkaline earth metals are obtained by the reaction of their oxides with water:
$MO + H_2O \rightarrow M(OH)_2$
Properties of these hydroxides:
$1$. The solubility,thermal stability,and basic character of these hydroxides increase down the group from $Mg(OH)_2$ to $Ba(OH)_2$.
$2$. Alkaline earth metal hydroxides are less basic and less stable than alkali metal hydroxides.
$3$. $Be(OH)_2$ is amphoteric in nature,reacting with both acids and bases:
- With bases: $Be(OH)_2 + 2OH^{-} \rightarrow [Be(OH)_4]^{2-}$ (Beryllate ion)
- With acids: $Be(OH)_2 + 2HCl + 2H_2O \rightarrow [Be(H_2O)_4]Cl_2$

(N/A) Except for beryllium halides,all other halides of alkaline earth metals are ionic in nature. Beryllium halides are essentially covalent and soluble in organic solvents. Beryllium chloride has a chain structure in the solid state as shown below:
In the vapour phase $BeCl_{2}$ tends to form a chloro-bridged dimer which dissociates into the linear monomer at high temperatures of the order of $1200 \ K$.
The tendency to form halide hydrates gradually decreases (for example,$MgCl_{2} \cdot 8 H_{2} O$,$CaCl_{2} \cdot 6 H_{2} O$,$SrCl_{2} \cdot 6 H_{2} O$,$BaCl_{2} \cdot 2 H_{2} O$) down the group.
The dehydration of hydrated chlorides,bromides and iodides of $Ca$,$Sr$ and $Ba$ can be achieved on heating; however,the corresponding hydrated halides of $Be$ and $Mg$ on heating suffer hydrolysis. The fluorides are relatively less soluble than the chlorides owing to their high lattice energies.

(A) Salts of oxoacids: The alkaline earth metals form various salts of oxoacids,including: $(a)$ Carbonates,$(b)$ Sulphates,and $(c)$ Nitrates.
$(a)$ Carbonates: Carbonates of alkaline earth metals are insoluble in water and can be precipitated by adding sodium or ammonium carbonate to a solution of their soluble salts. The solubility of these carbonates in water decreases as the atomic number of the metal ion increases. All carbonates decompose on heating to yield carbon dioxide and the metal oxide. $BeCO_{3}$ is unstable and must be stored in an atmosphere of $CO_{2}$. Thermal stability increases with increasing cationic size.
$(b)$ Sulphates: The sulphates of alkaline earth metals are white solids and are stable to heat. $BeSO_{4}$ and $MgSO_{4}$ are readily soluble in water,while solubility decreases from $CaSO_{4}$ to $BaSO_{4}$. The high hydration enthalpies of $Be^{2+}$ and $Mg^{2+}$ ions overcome the lattice enthalpy,making their sulphates soluble.
$(c)$ Nitrates: Nitrates are prepared by dissolving carbonates in dilute nitric acid. $Mg(NO_{3})_{2}$ crystallizes with six molecules of water,whereas $Ba(NO_{3})_{2}$ crystallizes as an anhydrous salt. This reflects a decreasing tendency to form hydrates as the cationic size increases and hydration enthalpy decreases. All nitrates decompose on heating to form the corresponding oxide: $2 M(NO_{3})_{2} \rightarrow 2 MO + 4 NO_{2} + O_{2}$ (where $M = Be, Mg, Ca, Sr, Ba$).

(N/A) Salts of Oxoacids: The alkaline earth metals form salts of oxoacids,including: $(a)$ Carbonates,$(b)$ Sulphates,and $(c)$ Nitrates.
$(a)$ Carbonates: Carbonates of alkaline earth metals are insoluble in water and are precipitated by adding sodium or ammonium carbonate to a solution of a soluble salt of these metals. The solubility of carbonates in water decreases as the atomic number of the metal ion increases. All carbonates decompose on heating to give carbon dioxide and the metal oxide. $BeCO_{3}$ is unstable and can only be stored in an atmosphere of $CO_{2}$. Thermal stability increases with increasing cationic size.
$(b)$ Sulphates: The sulphates of alkaline earth metals are white solids and are stable to heat. $BeSO_{4}$ and $MgSO_{4}$ are readily soluble in water; the solubility decreases from $CaSO_{4}$ to $BaSO_{4}$. The high hydration enthalpies of $Be^{2+}$ and $Mg^{2+}$ ions overcome the lattice enthalpy,making their sulphates soluble.
$(c)$ Nitrates: Nitrates are prepared by dissolving carbonates in dilute nitric acid. $Mg(NO_{3})_{2}$ crystallizes with six molecules of water,while $Ba(NO_{3})_{2}$ crystallizes as an anhydrous salt. This reflects a decreasing tendency to form hydrates with increasing size and decreasing hydration enthalpy. All nitrates decompose on heating to form the oxide: $2M(NO_{3})_{2} \rightarrow 2MO + 4NO_{2} + O_{2}$ (where $M = Be, Mg, Ca, Sr, Ba$).

(N/A) Beryllium,the first member of the Group-$2$ metals,shows anomalous behaviour compared to magnesium and the rest of the members. Further,it shows a diagonal relationship with aluminium.
$(i)$ Beryllium has exceptionally small atomic and ionic sizes and thus does not compare well with other members of the group. Because of its high ionisation enthalpy and small size,it forms compounds which are largely covalent and get easily hydrolysed.
$(ii)$ Beryllium does not exhibit a coordination number greater than four because its valence shell has only four orbitals. The remaining members of the group can have a coordination number of six by making use of $d$-orbitals.
$(iii)$ The oxide and hydroxide of beryllium,unlike the hydroxides of other elements in the group,are amphoteric in nature.

(N/A) Beryllium,the first member of the Group-$2$ metals,shows anomalous behavior compared to magnesium and the rest of the members. It also shows a diagonal relationship with aluminium.
$(i)$ Beryllium has an exceptionally small atomic and ionic size,which makes it distinct from other members of the group. Due to its high ionisation enthalpy and small size,it forms compounds that are largely covalent and easily hydrolysed.
$(ii)$ Beryllium does not exhibit a coordination number greater than four because its valence shell contains only four orbitals. Other members of the group can achieve a coordination number of six by utilizing $d$-orbitals.
$(iii)$ The oxide and hydroxide of beryllium are amphoteric in nature,unlike the hydroxides of other elements in the group,which are basic.

(N/A) Beryllium,the first member of the Group-$2$ metals,shows anomalous behavior compared to magnesium and the rest of the members. Further,it shows a diagonal relationship with aluminium.
$(i)$ Beryllium has exceptionally small atomic and ionic sizes and thus does not compare well with other members of the group. Because of high ionisation enthalpy and small size,it forms compounds which are largely covalent and get easily hydrolysed.
$(ii)$ Beryllium does not exhibit a coordination number greater than four as its valence shell has only four orbitals. The remaining members of the group can have a coordination number of six by making use of $d$-orbitals.
$(iii)$ The oxide and hydroxide of beryllium,unlike the hydroxides of other elements in the group,are amphoteric in nature.

(N/A) Beryllium,the first member of the Group-$2$ metals,shows anomalous behaviour compared to magnesium and the rest of the members. Further,it shows a diagonal relationship with aluminium.
$(i)$ Beryllium has exceptionally small atomic and ionic sizes and thus does not compare well with other members of the group. Because of high ionisation enthalpy and small size,it forms compounds which are largely covalent and get easily hydrolysed.
$(ii)$ Beryllium does not exhibit a coordination number greater than four as its valence shell has only four orbitals. The remaining members of the group can have a coordination number of six by making use of $d$-orbitals.
$(iii)$ The oxide and hydroxide of beryllium,unlike the hydroxides of other elements in the group,are amphoteric in nature.

(N/A) The ionic radius of $Be^{2+}$ is estimated to be $31 \text{ pm}$; the charge/radius ratio is nearly the same as that of the $Al^{3+}$ ion. Hence, beryllium resembles aluminum in some ways. Some of the similarities are:
$(i)$ Like aluminum, beryllium is not readily attacked by acids because of the presence of an oxide film on the surface of the metal.
$(ii)$ Beryllium hydroxide dissolves in excess of alkali to give a beryllate ion, $[Be(OH)_{4}]^{2-}$, just as aluminum hydroxide gives an aluminate ion, $[Al(OH)_{4}]^{-}$.
$(iii)$ The chlorides of both beryllium and aluminum have $Cl^{-}$ bridged chloride structures in the vapor phase. Both chlorides are soluble in organic solvents and are strong Lewis acids. They are used as Friedel-Crafts catalysts.
$(iv)$ Beryllium and aluminum ions have a strong tendency to form complexes like $[BeF_{4}]^{2-}$ and $[AlF_{6}]^{3-}$.

(N/A) The ionic radius of $Be^{2+}$ is estimated to be $31 \ pm$; the charge/radius ratio is nearly the same as that of the $Al^{3+}$ ion. Hence,beryllium resembles aluminium in some ways. Some of the similarities are:
$(i)$ Like aluminium,beryllium is not readily attacked by acids because of the presence of an oxide film on the surface of the metal.
$(ii)$ Beryllium hydroxide dissolves in excess of alkali to give a beryllate ion,$[Be(OH)_{4}]^{2-}$,just as aluminium hydroxide gives an aluminate ion,$[Al(OH)_{4}]^{-}$.
$(iii)$ The chlorides of both beryllium and aluminium have $Cl^{-}$ bridged chloride structures in the vapour phase. Both the chlorides are soluble in organic solvents and are strong Lewis acids. They are used as Friedel-Crafts catalysts.
$(iv)$ Beryllium and aluminium ions have a strong tendency to form complexes like $[BeF_{4}]^{2-}$ and $[AlF_{6}]^{3-}$.

(N/A) Preparation: It is prepared on a commercial scale by heating limestone $(CaCO_{3})$ in a rotary kiln at $1070-1270 \ K$.
$CaCO_{3} \xrightarrow{\text{heat}} CaO + CO_{2}$
The carbon dioxide is removed as soon as it is produced to enable the reaction to proceed to completion.
Properties: Calcium oxide is a white amorphous solid. It has a melting point of $2870 \ K$. On exposure to the atmosphere,it absorbs moisture and carbon dioxide.
$CaO + H_{2}O \rightarrow Ca(OH)_{2}$
$CaO + CO_{2} \rightarrow CaCO_{3}$
The addition of a limited amount of water breaks the lump of lime. This process is called slaking of lime. Quick lime slaked with soda gives solid soda lime. Being a basic oxide,it combines with acidic oxides at high temperatures.
$CaO + SiO_{2} \rightarrow CaSiO_{3}$
$6CaO + P_{4}O_{10} \rightarrow 2Ca_{3}(PO_{4})_{2}$
Uses: $(i)$ It is an important primary material for manufacturing cement and is the cheapest form of alkali. $(ii)$ It is used in the manufacture of sodium carbonate from caustic soda. $(iii)$ It is employed in the purification of sugar and in the manufacture of dye stuffs.

(N/A) Preparation: Calcium hydroxide is prepared by adding water to quick lime,$CaO$.
$CaO + H_2O \rightarrow Ca(OH)_2$
Properties: It is a white amorphous powder. It is sparingly soluble in water. The aqueous solution is known as lime water,and a suspension of slaked lime in water is known as milk of lime. When carbon dioxide is passed through lime water,it turns milky due to the formation of calcium carbonate.
$Ca(OH)_2 + CO_2 \rightarrow CaCO_3 + H_2O$
On passing an excess of carbon dioxide,the precipitate dissolves to form calcium hydrogen carbonate.
$CaCO_3 + CO_2 + H_2O \rightarrow Ca(HCO_3)_2$
Milk of lime reacts with chlorine to form hypochlorite,a constituent of bleaching powder:
$2Ca(OH)_2 + 2Cl_2 \rightarrow CaCl_2 + Ca(OCl)_2 + 2H_2O$
Uses: $(i)$ It is used in the preparation of mortar,a building material. $(ii)$ It is used in white washing due to its disinfectant nature. $(iii)$ It is used in glass making,in the tanning industry,for the preparation of bleaching powder,and for the purification of sugar.

(N/A) Preparation: Calcium carbonate occurs in nature in several forms like limestone,chalk,and marble. It can be prepared by passing carbon dioxide through slaked lime or by the addition of sodium carbonate to calcium chloride.
$Ca(OH)_2 + CO_2 \rightarrow CaCO_3 + H_2O$
$CaCl_2 + Na_2CO_3 \rightarrow CaCO_3 + 2NaCl$
Excess of $CO_2$ leads to the formation of water-soluble calcium hydrogen carbonate.
Properties: It is a white fluffy powder and is almost insoluble in water. When heated to $1200 \ K$,it decomposes to evolve $CO_2$.
$CaCO_3 \xrightarrow{1200 \ K} CaO + CO_2$
It reacts with dilute acids to liberate carbon dioxide.
$CaCO_3 + 2HCl \rightarrow CaCl_2 + H_2O + CO_2$
$CaCO_3 + H_2SO_4 \rightarrow CaSO_4 + H_2O + CO_2$
Uses: $(i)$ It is used as a building material in the form of marble and in the manufacture of quick lime. (ii) $A$ mixture of $CaCO_3$ and $MgCO_3$ is used as a flux in the extraction of metals such as iron. (iii) Specially precipitated $CaCO_3$ is extensively used in the manufacture of high-quality paper. (iv) It is also used as an antacid,mild abrasive in tooth-paste,a constituent of chewing gum,and a filler in cosmetics.

(N/A) Preparation: It is a hemihydrate of calcium sulphate. It is obtained when gypsum $(CaSO_{4} \cdot 2 H_{2} O)$ is heated to $393 \ K$.
$2(CaSO_{4} \cdot 2 H_{2} O) \rightarrow 2(CaSO_{4}) \cdot H_{2} O + 3 H_{2} O$
Above $393 \ K$,no water of crystallisation is left and anhydrous calcium sulphate $(CaSO_{4})$ is formed. This is known as 'dead burnt plaster'.
Properties: It has a remarkable property of setting with water. On mixing with an adequate quantity of water,it forms a plastic mass that sets into a hard solid in $5$ to $15$ minutes.
Uses: $(i)$ The largest use of Plaster of Paris is in the building industry as well as in plasters. $(ii)$ It is used for immobilising the affected part of an organ where there is a bone fracture or sprain. $(iii)$ It is also employed in dentistry,in ornamental work,and for making casts of statues and busts.

$(i)$ Magnesium burns in air with a dazzling light to form $MgO$ and $Mg_{3}N_{2}$.
$2Mg + O_{2} \xrightarrow{\Delta} 2MgO$
$3Mg + N_{2} \xrightarrow{\Delta} Mg_{3}N_{2}$
$(ii)$ Quick lime $(CaO)$ combines with silica $(SiO_{2})$ to form calcium silicate (slag).
$CaO + SiO_{2} \xrightarrow{\Delta} CaSiO_{3}$
$(iii)$ When chlorine is added to slaked lime,it gives bleaching powder.
$Ca(OH)_{2} + Cl_{2} \longrightarrow CaOCl_{2} + H_{2}O$
$(iv)$ Calcium nitrate,on heating,decomposes to give calcium oxide,nitrogen dioxide,and oxygen.
$2Ca(NO_{3})_{2(s)} \xrightarrow{\Delta} 2CaO_{(s)} + 4NO_{2(g)} + O_{2(g)}$

(N/A) Cement is an important building material. It was first introduced in England in $1824$ by Joseph Aspdin.
It is commonly called Portland cement because it resembles the natural limestone quarried in the Isle of Portland,England.
Cement is a product obtained by combining a material rich in lime $(CaO)$ with other materials such as clay,which contains silica $(SiO_{2})$ along with the oxides of aluminium,iron,and magnesium.

(N/A) An adult body contains about $25 \ g$ of $Mg$ and $1200 \ g$ of $Ca$,compared with only $5 \ g$ of iron and $0.06 \ g$ of copper. The daily requirement in the human body is estimated to be $200-300 \ mg$.
All enzymes that utilize $ATP$ in phosphate transfer require magnesium as a cofactor. The main pigment for the absorption of light in plants is chlorophyll,which contains magnesium. About $99 \%$ of body calcium is present in bones and teeth. It also plays important roles in neuromuscular function,interneuronal transmission,cell membrane integrity,and blood coagulation.
The calcium concentration in plasma is regulated at about $100 \ mg \ L^{-1}$. It is maintained by two hormones: calcitonin and parathyroid hormone.
Bone is not an inert and unchanging substance but is continuously being solubilized and redeposited to the extent of $400 \ mg$ per day in humans. All this calcium passes through the plasma.

(D) Thermal stability of alkaline earth metal carbonates increases as the size of the cation increases.
The ionic radii of the alkaline earth metal cations follow the order: $Mg^{2+} < Ca^{2+} < Sr^{2+} < Ba^{2+}$.
Therefore,the thermal stability increases in the order: $MgCO_{3} < CaCO_{3} < SrCO_{3} < BaCO_{3}$.
Thus,$BaCO_{3}$ is the most thermally stable carbonate.

(N/A) When an alkaline earth metal is heated,the valence electrons get excited to a higher energy level.
When this excited electron returns to its ground state,it releases energy in the form of light.
If this energy corresponds to the visible region of the electromagnetic spectrum,a characteristic colour is observed.
In $Be$ and $Mg$ atoms,the valence electrons are very strongly bound to the nucleus due to their small size and high ionization enthalpy.
The energy required to excite these electrons to a higher energy level is very high,which is not available in a standard Bunsen flame.
Consequently,the energy released upon their return to the ground state does not fall within the visible region,and thus,no colour is observed in the flame.

(B) The elements of group-$2$ are known as alkaline earth metals because their oxides and hydroxides are alkaline in nature.
Furthermore,these metal oxides are widely found in the earth's crust.

(A) The flame test colors for alkaline earth metals are as follows:
$(a)$ Calcium $(Ca)$ imparts a brick red color to the flame.
$(b)$ Barium $(Ba)$ imparts an apple green color to the flame.
$(c)$ Strontium $(Sr)$ imparts a crimson red color to the flame.
Therefore,the correct matching is: $a-q, b-r, c-p$.

(A) $(a)-(p, s), (b)-(q), (c)-(r), (d)-(p)$
$1$. $BeO$ is amphoteric and basic in nature.
$2$. $MgCO_3 \cdot CaCO_3$ is known as Dolomite.
$3$. $BeCl_2$ is a covalent compound due to the small size and high polarizing power of $Be^{2+}$ ion.
$4$. $Al(OH)_3$ is amphoteric in nature.

(A) The correct matches are as follows:
$1. \text{Be (Beryllium)}$: Used to prepare windows of $X$-ray tubes because of its low atomic number and high transparency to $X$-rays $(1-c)$.
$2. \text{Ca (Calcium)}$: Used to remove air from vacuumized tubes as a getter $(2-d)$.
$3. \text{Mg (Magnesium)}$: Used to prepare Grignard reagents $(R-Mg-X)$ $(3-a)$.
Therefore,the correct sequence is $1-c, 2-d, 3-a$.

(A) The given compound is dolomite,$CaCO_3 \cdot MgCO_3$.
Upon heating (calcination),both carbonates decompose to form their respective oxides and release carbon dioxide gas.
The balanced chemical equation is:
$CaCO_3 \cdot MgCO_{3(s)} \xrightarrow{\Delta } CaO_{(s)} + MgO_{(s)} + 2CO_{2(g)}$

(N/A) The chemical formulas are as follows:
$1$. Gypsum: $CaSO_4 \cdot 2H_2O$
$2$. Epsom salt: $MgSO_4 \cdot 7H_2O$
$3$. Barite: $BaSO_4$

(B) The group-$2$ elements are known as alkaline earth metals. These elements include beryllium $(Be)$,magnesium $(Mg)$,calcium $(Ca)$,strontium $(Sr)$,barium $(Ba)$,and radium $(Ra)$.

(N/A) Beryllium is used in the manufacture of alloys,particularly copper-beryllium alloys,which are used in the preparation of high-strength springs. It is also used in making windows for $X$-ray tubes.
Magnesium is used in flash powders and bulbs,incendiary bombs,and signals. $A$ suspension of magnesium hydroxide,$Mg(OH)_2$,in water (known as milk of magnesia) is used as an antacid in medicine.

(N/A) Beryllium exhibits amphoteric character because its oxide $(BeO)$ and hydroxide $(Be(OH)_2)$ can react with both acids and bases to form salts and water.
For example,it reacts with $HCl$ (acid) to form $BeCl_2$ and with $NaOH$ (base) to form sodium beryllate $(Na_2[Be(OH)_4])$.

(N/A) The important compounds of calcium are:
$1$. Calcium oxide $(CaO)$
$2$. Calcium hydroxide $(Ca(OH)_2)$
$3$. Calcium carbonate $(CaCO_3)$
$4$. Calcium sulphate $(CaSO_4 \cdot 2H_2O)$
$5$. Cement

(N/A) The addition of a limited amount of water to quicklime $(CaO)$ causes it to break into a powder,forming calcium hydroxide $(Ca(OH)_2)$. This process is known as slaking of lime. The chemical reaction is: $CaO(s) + H_2O(l) \rightarrow Ca(OH)_2(s)$.

(N/A) $i$. It is an important primary material for manufacturing cement.
$ii$. It is employed in the purification of sugar and in the manufacture of dye stuffs.

(B) Calcium hydroxide,also known as slaked lime,is prepared by the addition of a limited amount of water to quick lime $(CaO)$.
This process is known as slaking of lime.
The chemical reaction is: $CaO(s) + H_2O(l) \rightarrow Ca(OH)_2(aq)$.

(N/A) $(i)$ It is used in the preparation of mortar,a building material.
$(ii)$ It is used in the manufacture of sodium carbonate from caustic soda.

(N/A) $Calcium \ carbonate$ $(CaCO_3)$ is used in the manufacture of high-quality paper. It is also used as an antacid in medicine to neutralize excess stomach acid.

(B) Calcium sulphate is obtained by heating gypsum,$CaSO_{4} \cdot 2H_{2}O$,to $393 \ K$.
The reaction is: $CaSO_{4} \cdot 2H_{2}O \xrightarrow{393 \ K} CaSO_{4} \cdot \frac{1}{2}H_{2}O + \frac{3}{2}H_{2}O$.

(N/A) Above $393 \ K$,no water of crystallization is left and anhydrous calcium sulphate,$CaSO_{4}$,is formed. This is known as 'dead burnt plaster'.

(N/A) $(i)$ It is used in the building industry for plastering walls and ceilings.
(ii) It is used in dentistry and for making casts of statues,toys,and busts.

(N/A) Cement is known as Portland cement because its hardened form resembles the natural limestone quarried on the Isle of Portland in England.

(N/A) The average composition of Portland cement is as follows:
$CaO: 50-60 \%$,$SiO_{2}: 20-25 \%$,$Al_{2}O_{3}: 5-10 \%$
$MgO: 2-3 \%$,$Fe_{2}O_{3}: 1-2 \%$,$SO_{3}: 1-2 \%$

(A) The important ingredients present in Portland cement are dicalcium silicate $(Ca_2SiO_4)$ $(26 \%)$,tricalcium silicate $(Ca_3SiO_5)$ $(51 \%)$,and tricalcium aluminate $(Ca_3Al_2O_6)$ $(11 \%)$.

(B) The purpose of adding gypsum $(CaSO_4 \cdot 2H_2O)$ to cement is to slow down the process of setting of the cement.
This allows the cement to remain in a plastic state for a longer duration,ensuring it gets sufficiently hardened and gains proper strength.

(D) Cement is a vital material for national infrastructure,second only to iron and steel.
It is primarily used in the following ways:
$1$. In the preparation of concrete and reinforced concrete.
$2$. In plastering walls and surfaces.
$3$. In the construction of major infrastructure such as bridges,dams,and buildings.

(N/A) $(1)$ Brick red,light green
$(2)$ Flame photometry
$(3)$ $BeO$ and $Be_{3}N_{2}$
$(4)$ Beryllate ion $[Be(OH)_{4}]^{2-}$

(N/A) $(1)$ Quick lime
$(2)$ Milk of lime
$(3)$ Slaked lime
$(4)$ Plaster of Paris

(N/A) $(1)$ Portland cement
$(2)$ Joseph Aspdin

(A) $(1)$ True statement.
$(2)$ False statement: In a group,when moving top to bottom,atomic size increases and hence,ionization enthalpy decreases.
$(3)$ False statement: $BeCl_2$ possesses a polymer-like arrangement in the solid state,whereas it exists as a dimer or monomer in the vapour phase.
$(4)$ False statement: $Ca(OH)_2$ (slaked lime) is used to prepare mortar.

(A) $(1)$ True: Plaster of Paris $(CaSO_4 \cdot \frac{1}{2}H_2O)$ sets into a hard mass when mixed with water,making it useful for orthopedic casts.
$(2)$ True: Cement is manufactured by heating a mixture of limestone $(CaCO_3)$ and clay (which provides soil components like silica,alumina,and iron oxide) in a rotary kiln.

(BERYLLIUM) Beryllium $(Be)$ is the element from Group $2$ that forms an amphoteric oxide,$BeO$,and a water-soluble sulphate,$BeSO_{4}$.
$BeO$ reacts with both acids and bases:
$BeO + 2HCl \rightarrow BeCl_{2} + H_{2}O$
$BeO + 2NaOH + H_{2}O \rightarrow Na_{2}[Be(OH)_{4}]$
$BeSO_{4}$ is soluble in water because its high hydration energy compensates for its lattice energy.

(N/A) $(i)$ Thermal stability increases down the group. On heating,these carbonates decompose to give $CO_{2}$ and the corresponding oxide of the Group $-2$ element.
$BeCO_{3} < MgCO_{3} < CaCO_{3} < SrCO_{3} < Ba_{CO_{3}}$
$BeCO_{3}$ is unstable and is kept stable in an atmosphere of $CO_{2}$.
$BeCO_{3} \rightleftharpoons BeO + CO_{2}$
$(ii)$ All oxides are basic and ionic in nature except $BeO$,which is amphoteric and covalent. Lattice energy of oxides decreases as the size of the cation increases. Basic nature increases down the group.
Except $BeO$ and $MgO$,all oxides are soluble in water and produce a large amount of heat upon dissolution. $BeO$ and $MgO$,due to their high lattice energy,are insoluble in water.