Alkaline earth metals Questions in English

(A) The substance known as 'blanc fixe' is chemically pure barium sulfate $(BaSO_4)$.
It is prepared by the reaction of barium chloride with sodium sulfate.
It is widely used as a white pigment in paints and as a filler in paper and rubber industries.

(D) Among the alkaline earth metals,the tendency to form peroxides increases down the group as the size of the metal cation increases and the lattice energy becomes more favorable for the larger peroxide ion $(O_2^{2-})$.
$Be$ and $Mg$ form only oxides $(MO)$.
$Ca$ and $Sr$ form peroxides $(MO_2)$ to a limited extent.
$Ba$ has the largest size among the given options and forms $BaO_2$ (Barium peroxide) most readily.

(C) Mortar is a paste used to bind building blocks such as stones,bricks,and concrete together.
It is prepared by mixing slaked lime $(Ca(OH)_2)$,sand,and water.
Therefore,the correct option is $C$.

(C) Celestite is a mineral composed of strontium sulfate,with the chemical formula $SrSO_4$.
Therefore,it is a mineral of $Sr$ (Strontium).

(B) The basic strength of hydroxides depends on the ease with which they release $OH^-$ ions in an aqueous solution.
$NaOH$ and $KOH$ are strong bases as they are alkali metal hydroxides.
$Ca(OH)_2$ and $Ba(OH)_2$ are alkaline earth metal hydroxides.
Among the given options,$Ba(OH)_2$ is more basic than $Ca(OH)_2$ because the basic strength of alkaline earth metal hydroxides increases down the group due to the increase in ionic character and solubility.
Therefore,$Ca(OH)_2$ is the weakest base among the given options.

(A) During the manufacturing of cement,gypsum $(CaSO_4 \cdot 2H_2O)$ is added to the clinker in small amounts (about $2-3 \%$).
This is done to slow down the process of setting of cement,which allows sufficient time for mixing and laying the concrete.

(C) Beryllium $(Be)$ and Aluminum $(Al)$ exhibit a diagonal relationship due to similar charge-to-size ratios.
$1$. Both $Be_2C$ and $Al_4C_3$ produce methane $(CH_4)$ upon hydrolysis.
$2$. Both metals become passive towards concentrated $HNO_3$ due to the formation of an inert oxide layer.
$3$. Both $Be(OH)_2$ and $Al(OH)_3$ are amphoteric in nature,not basic. Therefore,the statement that $Be(OH)_2$ is basic like $Al(OH)_3$ is incorrect.
$4$. Both form complex anions: $Be$ forms beryllates $([Be(OH)_4]^{2-})$ and $Al$ forms aluminates $([Al(OH)_4]^-)$.

(A) Among the alkaline earth metals,$Be$ (Beryllium) has a very small atomic size and high ionization enthalpy. Due to its high charge density and the availability of vacant $2p$ orbitals,it has a strong tendency to form covalent bonds and complex compounds (e.g.,$[Be(H_2O)_4]^{2+}$,$[BeF_4]^{2-}$). Other alkaline earth metals are larger and tend to form ionic compounds.

(D) The tendency to form ionic bonds increases as we move down the group due to the increase in atomic size and decrease in ionization energy.
$BeCl_2$ is covalent due to the small size and high polarizing power of the $Be^{2+}$ ion.
$MgCl_2$ is predominantly covalent with some ionic character.
$CaCl_2$ and $SrCl_2$ are ionic in nature.
Among the given options,$Sr$ is the most electropositive and forms a purely ionic chloride.

(A) Alkaline earth metals belong to Group $2$ of the periodic table.
They have $2$ valence electrons in their outermost shell.
To achieve a stable noble gas configuration,they lose these $2$ electrons,forming ions with a $+2$ charge.
Therefore,the general formula for these ions is $M^{2+}$.

(A) The flame test is used to identify metal ions based on the characteristic color they impart to the flame.
Among the alkaline earth metals,$Be^{2+}$ and $Mg^{2+}$ do not impart any characteristic color to the Bunsen flame because the energy of the flame is not sufficient to excite their electrons to higher energy levels.
$BeCl_2$ is soluble in water.
Therefore,the cation present is $Be^{2+}$.

(C) When magnesium $(Mg)$ is burnt in air,it forms magnesium oxide $(MgO)$ and magnesium nitride $(Mg_3N_2)$.
$2Mg + O_2 \rightarrow 2MgO$
$3Mg + N_2 \rightarrow Mg_3N_2$
When the resulting ash (containing $Mg_3N_2$) is moistened,it undergoes hydrolysis to produce ammonia $(NH_3)$ gas,which has a characteristic pungent smell.
$Mg_3N_2 + 6H_2O \rightarrow 3Mg(OH)_2 + 2NH_3 \uparrow$

(C) The statement '$Mg$ cannot form complex salts' is incorrect. $Mg^{2+}$ ions,due to their small size and high charge density,are capable of forming various complex compounds,such as chlorophyll (a magnesium complex) and various coordination complexes with ligands like $EDTA$ or $NH_3$.

(B) As we move down the group in alkaline earth metals ($Be$ to $Ba$),the atomic size increases.
$1$. Ionization enthalpy decreases due to the increase in atomic size.
$2$. Electronegativity decreases as the atomic size increases.
$3$. The solubility of alkaline earth metal hydroxides increases down the group because the lattice energy decreases more rapidly than the hydration energy.
$4$. The solubility of alkaline earth metal sulfates decreases down the group because the lattice energy decreases less rapidly than the hydration energy.
Therefore,the solubility of their hydroxides increases with an increase in atomic number.

(B) The reactions are as follows:
$1$. Heating $X$ $(CaCO_3)$: $CaCO_3 \xrightarrow{\Delta} CaO + CO_2 \uparrow$
$2$. Residue $(CaO)$ + water $(H_2O)$: $CaO + H_2O \rightarrow Ca(OH)_2$ $(Y)$
$3$. Passing excess $CO_2$ through $Y$ $(Ca(OH)_2)$: $Ca(OH)_2 + 2CO_2 \rightarrow Ca(HCO_3)_2$ $(Z)$
$4$. Boiling $Z$ $(Ca(HCO_3)_2)$: $Ca(HCO_3)_2 \xrightarrow{\Delta} CaCO_3 + H_2O + CO_2 \uparrow$ (Regenerates $X$)
Therefore,$X$ is $CaCO_3$.

(A) For a compound to be highly soluble in water,its hydration energy must be greater than its lattice energy. \\ The solubility of alkaline earth metal sulfates decreases down the group from $BeSO_4$ to $BaSO_4$. \\ $BeSO_4$ and $MgSO_4$ are highly soluble in water because their hydration energy is significantly higher than their lattice energy. \\ As the size of the cation increases down the group,the lattice energy decreases,but the hydration energy decreases much more rapidly. \\ Therefore,$MgSO_4$ is the only compound among the options where the hydration energy is greater than the lattice energy,making it highly soluble.

(A) The assertion is correct because $Be$ can expand its coordination number up to $4$ due to the presence of only $s$ and $p$ orbitals in its valence shell $(n=2)$.
$Al$ can expand its coordination number to $6$ because it has vacant $3d$ orbitals in its valence shell $(n=3)$.
Since $Be$ lacks vacant $d-$ orbitals,it cannot form $BeF_6^{3-}$,which would require a coordination number of $6$.
Thus,the reason correctly explains the assertion.

(B) The melting point of alkaline earth metal chlorides depends on their ionic character.
As the size of the cation increases from $Be^{2+}$ to $Ba^{2+}$,the polarizing power of the cation decreases,leading to an increase in the ionic character of the $M-Cl$ bond.
Greater ionic character results in stronger electrostatic forces of attraction,which increases the melting point.
Therefore,the order of increasing melting points is $BeCl_2 < MgCl_2 < CaCl_2 < SrCl_2 < BaCl_2$.

(D) $Be$ (Beryllium) is the first member of the alkaline earth metals and exhibits anomalous properties compared to the rest of the group. Due to its small size and high charge density,it shows a diagonal relationship with $Al$ (Aluminium). Both $Be$ and $Al$ have an electronegativity value of approximately $1.5$ on the Pauling scale.

(B) The thermal stability of alkaline earth metal carbonates increases as we move down the group due to the increase in the size of the metal cation,which decreases the polarizing power of the cation.
The order of thermal stability for alkaline earth metal carbonates is: $BeCO_3 < MgCO_3 < CaCO_3$.
$K_2CO_3$ is a carbonate of an alkali metal (Group $1$). Alkali metal carbonates are generally much more thermally stable than alkaline earth metal carbonates (Group $2$).
Therefore,the overall order of increasing thermal stability is: $BeCO_3 (IV) < MgCO_3 (II) < CaCO_3 (III) < K_2CO_3 (I)$.

(C) The chemical $A$ is calcium hydroxide,$Ca(OH)_2$,also known as slaked lime.
$1$. It is used for water softening to remove temporary hardness by reacting with calcium bicarbonate: $Ca(HCO_3)_2 + Ca(OH)_2 \to 2CaCO_3 \downarrow + 2H_2O$.
$2$. It reacts with sodium carbonate $(Na_2CO_3)$ to produce caustic soda $(NaOH)$: $Ca(OH)_2 + Na_2CO_3 \to 2NaOH + CaCO_3$.
$3$. When $CO_2$ is passed through lime water $(Ca(OH)_2)$,it turns milky due to the formation of insoluble calcium carbonate: $Ca(OH)_2 + CO_2 \to CaCO_3 \downarrow + H_2O$.

(B) The assertion is correct because $Ba^{2+}$ ions are toxic and have no known biological role in humans.
The reason is also correct because $Ba$ is an alkaline earth metal and exhibits a constant oxidation state of $+2$,not variable oxidation states.
However,the reason does not explain why $Ba$ is not required for biological functions.
Therefore,both are correct but the reason is not the correct explanation of the assertion.

(B) The assertion is correct because $BeCl_2$ is covalent due to the high polarizing power of the small $Be^{2+}$ ion (Fajans' rule),while $MgCl_2$ and $CaCl_2$ are ionic.
The reason is also correct because $Be$ is indeed the first member of group $2$.
However,the reason is not the correct explanation of the assertion. The covalent nature of $BeCl_2$ is due to the small size and high charge density of the $Be^{2+}$ ion,not simply because it is the first member of the group.

(B) . Plaster of paris $= CaSO_4.\frac{1}{2}H_2O$
$B$. Epsomite $= MgSO_4.7H_2O$
$C$. Kieserite $= MgSO_4.H_2O$
$D$. Gypsum $= CaSO_4.2H_2O$
Therefore,the correct matching is $A-ii, B-iii, C-iv, D-i$.

(B) All enzymes that utilize $ATP$ in phosphate transfer require magnesium $(Mg^{2+})$ as the cofactor.
Magnesium ions play a crucial role in stabilizing the negative charges on the phosphate groups of $ATP$,facilitating the transfer of the phosphate group.

(B) In the alkaline earth metals group,as we move down the group,the metallic character increases,which leads to an increase in the basic nature of the oxides. $BeO$ is amphoteric in nature,while $MgO$,$CaO$,and $BaO$ are basic. Among the given options,$BeO$ is the most acidic (or least basic) because it has the smallest size and highest polarizing power.

(A) The ionic character of metal hydrides depends on the electronegativity difference between the metal and hydrogen,or alternatively,on the polarizing power of the cation.
According to Fajan's rule,smaller cations have higher polarizing power,which increases the covalent character of the bond.
As we move down the group $2$ ($Be$ to $Ba$),the size of the cation increases,and its polarizing power decreases.
Therefore,the covalent character decreases and the ionic character increases down the group.
The order of ionic character is $BeH_2 < CaH_2 < BaH_2$.

(D) Gypsum is $CaSO_4 \cdot 2 H_2 O$.
When heated to $393 \; K$,it loses three-fourths of its water of crystallization to form Plaster of Paris.
The reaction is: $CaSO_4 \cdot 2 H_2 O \xrightarrow{393 \; K} CaSO_4 \cdot 0.5 H_2 O + 1.5 H_2 O$.
Thus,the product formed is $CaSO_4 \cdot 0.5 H_2 O$.

(N/A) In alkaline earth metal hydroxides,the anion $(OH^-)$ is common,so the cationic radius influences the lattice enthalpy. As we move down the group,the cationic size increases. The lattice enthalpy decreases more significantly than the hydration enthalpy with an increase in ionic size. Consequently,the solubility of these hydroxides increases down the group.

(N/A) The size of the anions (like $CO_3^{2-}$ and $SO_4^{2-}$) is much larger compared to the cations. Consequently,the lattice enthalpy remains almost constant as we move down the group. However,the hydration enthalpy of the cations decreases significantly down the group. Since the solubility depends on the balance between lattice enthalpy and hydration enthalpy,the decrease in hydration enthalpy leads to a decrease in solubility down the group.

(N/A) General characteristics of alkaline earth metals are as follows:
$(i)$ The general electronic configuration of alkaline earth metals is $[\text{noble gas}] ns^{2}$.
$(ii)$ 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 $Be_{3}N_{2}$.
$(b)$ $Mg$,being more electropositive,burns in air with a dazzling sparkle to form $MgO$ and $Mg_{3}N_{2}$.
$(c)$ $Ca, Sr,$ 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.
$M + X_{2} \longrightarrow MX_{2} (X = F, Cl, Br, I)$
$(iii)$ 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.
$M + 2HCl \longrightarrow MCl_{2} + H_{2(g)} \uparrow$
$(v)$ 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.
$M + (x + y)NH_{3} \longrightarrow [M(NH_{3})_{x}]^{2+} + 2[e(NH_{3})_{y}]^{-}$

(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 $Mg$,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,and the energy released upon their return to the ground state does not fall within the visible region of the spectrum.
Therefore,$Be$ and $Mg$ do not impart any characteristic colour to the flame.

(N/A) $(i)$ Magnesium burns in air with a dazzling light to form $MgO$ and $Mg_3N_2$.
$2Mg + O_2 \xrightarrow{\Delta} 2MgO$
$3Mg + N_2 \xrightarrow{\Delta} Mg_3N_2$
$(ii)$ Quick lime $(CaO)$ combines with silica $(SiO_2)$ to form calcium silicate (slag).
$CaO + SiO_2 \xrightarrow{\Delta} CaSiO_3$
$(iii)$ When chlorine reacts with slaked lime,it forms bleaching powder.
$Ca(OH)_2 + Cl_2 \xrightarrow{\Delta} CaOCl_2 + H_2O$
$(iv)$ 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) $(i)$ Chemically,limestone is $CaCO_{3}$.
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 the construction of bridges.
$(b)$ It is used in concrete.
$(iii)$ Chemically,plaster of Paris is $CaSO_{4} \cdot \frac{1}{2}H_{2}O$.
Importance of plaster of Paris:
$(a)$ It is used in surgical bandages.
$(b)$ It is also used for making casts and moulds.

(A) $(i)$ $BeO$ is almost insoluble in water because $Be^{2+}$ and $O^{2-}$ are both small ions,leading to high lattice energy that hydration energy cannot overcome. $BeSO_4$ is soluble because the large $SO_4^{2-}$ anion reduces the lattice energy.
$(ii)$ $BaO$ is soluble because $Ba^{2+}$ and $O^{2-}$ have a size mismatch,resulting in lower lattice energy. $BaSO_4$ is insoluble because both $Ba^{2+}$ and $SO_4^{2-}$ are large,leading to high lattice energy due to favorable packing.
$(iii)$ $LiI$ is more soluble than $KI$ in ethanol due to the higher covalent character of $LiI$ caused by the high polarising power of the small $Li^+$ ion,which makes it more compatible with the organic solvent ethanol.

(D) Thermal stability of alkaline earth metal carbonates increases as the size of the cation increases down the group.
This is because the larger cation stabilizes the large carbonate anion more effectively.
The order of cationic size for the given alkaline earth metals is $Mg^{2+} < Ca^{2+} < Sr^{2+} < Ba^{2+}$.
Therefore,the order of thermal stability is $MgCO_3 < CaCO_3 < SrCO_3 < BaCO_3$.
Thus,$BaCO_3$ is the most thermally stable carbonate.

(N/A) The group-$2$ elements comprise beryllium $(Be)$, magnesium $(Mg)$, calcium $(Ca)$, strontium $(Sr)$, barium $(Ba)$, and radium $(Ra)$. They follow alkali metals in the periodic table.
These (except beryllium) are known as alkaline earth metals. The first element, beryllium, differs from the rest of the members and shows a diagonal relationship with aluminium.
Atomic and Physical Properties of the Alkaline Earth Metals:

Property	Beryllium $(Be)$	Magnesium $(Mg)$	Calcium $(Ca)$	Strontium $(Sr)$	Barium $(Ba)$	Radium $(Ra)$
Atomic number	$4$	$12$	$20$	$38$	$56$	$88$
Atomic mass $(g \ mol^{-1})$	$9.01$	$24.31$	$40.08$	$87.62$	$137.33$	$226.03$
Electronic configuration	$[He] \ 2s^{2}$	$[Ne] \ 3s^{2}$	$[Ar] \ 4s^{2}$	$[Kr] \ 5s^{2}$	$[Xe] \ 6s^{2}$	$[Rn] \ 7s^{2}$
Ionization enthalpy $(I)$ $(kJ \ mol^{-1})$	$899$	$737$	$590$	$549$	$503$	$509$
Ionization enthalpy $(II)$ $(kJ \ mol^{-1})$	$1757$	$1450$	$1145$	$1064$	$965$	$979$
Hydration enthalpy $(kJ \ mol^{-1})$	$-2494$	$-1921$	$-1577$	$-1443$	$-1305$	$-$
Metallic radius $(pm)$	$111$	$160$	$197$	$215$	$222$	$-$
Ionic radius $(M^{2+})$ $(pm)$	$31$	$72$	$100$	$118$	$135$	$148$
$m.p.$ $(K)$	$1560$	$924$	$1124$	$1062$	$1002$	$973$
$b.p.$ $(K)$	$2745$	$1363$	$1767$	$1655$	$2078$	$(1973)$
Density $(g \ cm^{-3})$	$1.84$	$1.74$	$1.55$	$2.63$	$3.59$	$(5.5)$
Standard potential $E^{\ominus}$ $(V)$ $(M^{2+} / M)$	$-1.97$	$-2.36$	$-2.84$	$-2.89$	$-2.92$	$-2.92$
Occurrence in lithosphere	$2^{*}$	$2.76^{**}$	$4.6^{**}$	$384^{*}$	$390^{*}$	$10^{-6 *}$

$^{*}$ ppm (part per million), $^{**}$ Percentage by weight

Alkaline earth metals have two electrons in the $s$-orbital of the valence shell. Their general electronic configuration is represented as $[Noble \ gas] ns^{2}$.

Element	Electronic configuration
Beryllium $(Be)$	$1s^{2} 2s^{2}$
Magnesium $(Mg)$	$[Ne] 3s^{2}$
Calcium $(Ca)$	$[Ar] 4s^{2}$
Strontium $(Sr)$	$[Kr] 5s^{2}$
Barium $(Ba)$	$[Xe] 6s^{2}$
Radium $(Ra)$	$[Rn] 7s^{2}$

(N/A) The alkaline earth metals have low ionisation enthalpies due to the fairly large size of their atoms.
As the atomic size increases down the group,the ionisation enthalpy decreases.
The first ionisation enthalpies of the alkaline earth metals are higher than those of the corresponding Group-$1$ metals because of their smaller atomic size compared to the alkali metals.
It is interesting to note that the second ionisation enthalpies of the alkaline earth metals are smaller than those of the corresponding alkali metals.

(N/A) The hydration enthalpies of alkaline earth metal ions decrease with an increase in ionic size down the group. The order is $Be^{2+} > Mg^{2+} > Ca^{2+} > Sr^{2+} > Ba^{2+}$.
The hydration enthalpies of alkaline earth metal ions are larger than those of alkali metal ions due to their smaller size and higher charge.
Consequently,compounds of alkaline earth metals are more extensively hydrated than those of alkali metals.
For example,$MgCl_{2}$ and $CaCl_{2}$ exist as $MgCl_{2} \cdot 6H_{2}O$ and $CaCl_{2} \cdot 6H_{2}O$,whereas $NaCl$ and $KCl$ do not form such hydrates.

(N/A) The alkaline earth metals include $Be, Mg, Ca, Sr, Ba,$ and $Ra$. Their general characteristics and trends are as follows:
$1$. Electronic Configuration: They have a general valence shell electronic configuration of $ns^2$.
$2$. Atomic and Ionic Radii: The atomic and ionic radii increase down the group due to the addition of new shells.
$3$. Ionization Enthalpy: They have low ionization enthalpies,which decrease down the group as the atomic size increases.
$4$. Hydration Enthalpy: Hydration enthalpies decrease down the group as the ionic size increases. $Be^{2+}$ has the highest hydration enthalpy.
$5$. Physical Properties: They are silvery white,lustrous,and relatively soft,though harder than alkali metals. They impart characteristic colors to the flame (except $Be$ and $Mg$).
$6$. Chemical Reactivity: They are highly reactive,though less than alkali metals. Reactivity increases down the group.
$7$. Reducing Nature: They are strong reducing agents,with $E^{\circ}$ values becoming more negative down the group.

(N/A) The alkaline earth metals are generally silvery white,lustrous,and relatively soft,though they are harder than the alkali metals.
$Be$ and $Mg$ appear somewhat greyish. The melting and boiling points of these metals are higher than the corresponding alkali metals due to their smaller atomic sizes.
However,the trend in melting and boiling points is not entirely systematic. Due to their low ionisation enthalpies,they are strongly electropositive in nature.
The electropositive character increases down the group from $Be$ to $Ba$. $Ca$,$Sr$,and $Ba$ impart characteristic brick red,crimson,and apple green colours to the flame,respectively.
In the flame,electrons are excited to higher energy levels,and when they return to the ground state,energy is emitted as visible light.
The electrons in $Be$ and $Mg$ are too strongly bound to be excited by the flame,so they do not impart any colour.
The flame test for $Ca$,$Sr$,and $Ba$ is useful for their detection in qualitative analysis and estimation via flame photometry.
Like alkali metals,alkaline earth metals exhibit high electrical and thermal conductivities,which are typical metallic characteristics.

(N/A) $(i)$ Reactivity towards air and water: Beryllium and magnesium are kinetically inert to oxygen and water because of the formation of an oxide film on their surface.
However,powdered beryllium burns brilliantly on ignition in air to give $BeO$ and $Be_{3}N_{2}$.
Magnesium is more electropositive and burns with dazzling brilliance in air to give $MgO$ and $Mg_{3}N_{2}$.
Calcium,strontium,and barium are readily attacked by air to form the oxide and nitride. They also react with water with increasing vigour even in cold to form hydroxides.
$(ii)$ Reactivity towards the halogens: All the alkaline earth metals combine with halogens at elevated temperatures forming their halides,$MX_{2}$ $(X = F, Cl, Br, I)$.
Thermal decomposition of $(NH_{4})_{2}BeF_{4}$ is the best route for the preparation of $BeF_{2}$,and $BeCl_{2}$ is conveniently made from the oxide: $BeO + C + Cl_{2} \xrightarrow{600-800 \ K} BeCl_{2} + CO$.
$(iii)$ Reactivity towards hydrogen: All the elements except beryllium combine with hydrogen upon heating to form their hydrides,$MH_{2}$. $BeH_{2}$,however,can be prepared by the reaction of $BeCl_{2}$ with $LiAlH_{4}$: $2BeCl_{2} + LiAlH_{4} \rightarrow 2BeH_{2} + LiCl + AlCl_{3}$.
$(iv)$ Reactivity towards acids: The alkaline earth metals readily react with acids liberating dihydrogen: $M + 2HCl \rightarrow MCl_{2} + H_{2}$.
$(v)$ Reducing nature: Like alkali metals,the alkaline earth metals are strong reducing agents. This is indicated by large negative values of their reduction potentials. However,their reducing power is less than those of their corresponding alkali metals.
$(vi)$ Solutions in liquid ammonia: Like alkali metals,the alkaline earth metals dissolve in liquid ammonia to give deep blue-black solutions forming ammoniated ions: $M_{(s)} + (x+y)NH_{3} \rightarrow [M(NH_{3})_{x}]^{2+} + 2[e(NH_{3})_{y}]^{-}$.

(N/A) $(i)$ Reactivity towards air and water: $Be$ and $Mg$ are kinetically inert to oxygen and water due to the formation of an oxide film on their surface. Powdered $Be$ burns in air to give $BeO$ and $Be_{3}N_{2}$. $Mg$ burns with dazzling brilliance to give $MgO$ and $Mg_{3}N_{2}$. $Ca$,$Sr$,and $Ba$ react readily with air to form oxides and nitrides,and react vigorously with water to form hydroxides.
$(ii)$ Reactivity towards dihydrogen: All elements except $Be$ combine with $H_{2}$ upon heating to form hydrides $(MH_{2})$. $BeH_{2}$ is prepared by the reaction: $2BeCl_{2} + LiAlH_{4} \rightarrow 2BeH_{2} + LiCl + AlCl_{3}$.
$(iii)$ Reactivity towards acids: Alkaline earth metals react with acids to liberate $H_{2}$ gas: $M + 2HCl \rightarrow MCl_{2} + H_{2}$.
$(iv)$ Reducing nature: Alkaline earth metals are strong reducing agents due to large negative reduction potentials,though less than alkali metals. $Be$ has a less negative reduction potential,but its reducing nature is attributed to high hydration energy of $Be^{2+}$ and high atomization enthalpy.
$(v)$ Liquid ammonia: They dissolve in liquid ammonia to form deep blue-black solutions: $M + (x+y)NH_{3} \rightarrow [M(NH_{3})_{x}]^{2+} + 2[e(NH_{3})_{y}]^{-}$.

(N/A) $Beryllium$ is used in the manufacture of alloys. $Copper-beryllium$ alloys are used in the preparation of high-strength springs.
Metallic $beryllium$ is used for making windows of $X$-ray tubes. $Magnesium$ forms alloys with $aluminium$,$zinc$,$manganese$,and $tin$.
$Magnesium-aluminium$ alloys,being light in mass,are used in aircraft construction. $Magnesium$ (powder and ribbon) is used in flash powders and bulbs,incendiary bombs,and signals.
$A$ suspension of $magnesium$ $hydroxide$ in water (called $milk$ $of$ $magnesia$) is used as an antacid in medicine. $Magnesium$ $carbonate$ is an ingredient of toothpaste.
$Calcium$ is used in the extraction of metals from oxides which are difficult to reduce with carbon. $Calcium$ and $barium$ metals,owing to their reactivity with oxygen and nitrogen at elevated temperatures,have often been used to remove air from vacuum tubes.
$Radium$ salts are used in radiotherapy,for example,in the treatment of cancer.