The S-Block Elements (Group 2 - Alkaline Earth Metals)

Group 2 Elements : Alkaline Earth Metals

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The Group 2 elements, known as the alkaline earth metals, include Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba), and Radium (Ra). They are reactive metals, though generally less reactive than alkali metals.

Electronic Configuration

General Configuration: Alkaline earth metals have the general outer electronic configuration $ns^2$, where $n$ is the principal quantum number of the outermost shell.

• Be: $[He] 2s^2$
• Mg: $[Ne] 3s^2$
• Ca: $[Ar] 4s^2$
• Sr: $[Kr] 5s^2$
• Ba: $[Xe] 6s^2$
• Ra: $[Rn] 7s^2$

Significance: The presence of two valence electrons in the outermost s-orbital dictates their tendency to lose these two electrons to form $+2$ ions ($M^{2+}$), leading to their characteristic chemical behavior.

Atomic And Ionic Radii

Atomic Radii:

• Trend: Atomic radii increase down the group from Be to Ra.
• Reason: Similar to alkali metals, the increase in the number of electron shells leads to a larger atomic size.
• Comparison to Group 1: The atomic radii of alkaline earth metals are generally smaller than those of the corresponding alkali metals in the same period. This is because they have a greater nuclear charge (more protons) pulling on the same number of valence electrons.

Ionic Radii:

• Trend: Ionic radii of alkaline earth metal ions ($M^{2+}$) also increase down the group (e.g., $Be^{2+} < Mg^{2+} < Ca^{2+} < Sr^{2+} < Ba^{2+}$).
• Reason: Increase in electron shells.
• Comparison to Group 1: The ionic radii of $M^{2+}$ ions are generally smaller than the ionic radii of $M^+$ ions in the same period (e.g., $Mg^{2+}$ is smaller than $Na^+$). This is due to the higher nuclear charge in $M^{2+}$ pulling the same number of electrons more strongly.

Ionization Enthalpies

Trend: Ionization enthalpies (both first and second) of alkaline earth metals are higher than those of alkali metals in the same period.

Reason:

• Greater Nuclear Charge: Alkaline earth metals have two valence electrons and a higher nuclear charge than alkali metals in the same period.
• Smaller Atomic Size: They are smaller, bringing the valence electrons closer to the nucleus.
• Stable $ns^2$ Configuration: The $ns^2$ configuration is relatively stable, requiring more energy to remove the first or second electron.

Trend Down the Group: Ionization enthalpies decrease from Be to Ra.

Reason: Increase in atomic size and increased shielding effect of inner electrons, weakening the nuclear attraction for the valence electrons.

Second Ionization Enthalpy ($IE_2$): The second ionization enthalpy ($IE_2$) is always greater than the first ($IE_1$). However, the sum of the first and second ionization enthalpies ($IE_1 + IE_2$) is less than the sum of the first two ionization enthalpies of the corresponding alkali metal. This explains why alkali metals preferentially form $+1$ ions and alkaline earth metals $+2$ ions.

Hydration Enthalpies

Trend: Hydration enthalpies of alkaline earth metal ions decrease from $Be^{2+}$ to $Ba^{2+}$.

Reason: Similar to alkali metals, hydration enthalpy is governed by charge density. $Be^{2+}$ has the highest charge density (due to its small size and $+2$ charge), leading to the strongest interaction with water molecules and the highest hydration enthalpy. $Ba^{2+}$ has the lowest charge density and thus the lowest hydration enthalpy.

Comparison to Group 1: The hydration enthalpies of $M^{2+}$ ions are significantly higher than those of $M^+$ ions of similar size (e.g., $| \Delta H_{hydration}(Mg^{2+}) | > | \Delta H_{hydration}(Na^+) |$). This higher hydration energy contributes to the greater stability of alkaline earth metal ions in solution.

Physical Properties

Appearance: They are silvery-white and relatively hard metals.

Hardness: Beryllium and magnesium are harder than other alkaline earth metals. Their hardness decreases down the group.

Melting and Boiling Points: They have higher melting and boiling points than alkali metals. This is due to the stronger metallic bonding, as each atom contributes two electrons to the electron sea.

Density: Their densities are higher than those of alkali metals and generally increase down the group.

Flame Coloration: Unlike alkali metals, most alkaline earth metals impart characteristic colors to the flame, but these colors are less distinct and less used for flame tests.

• Be, Mg: Do not impart color to the flame (magnesium burns with a dazzling white light).
• Ca: Brick red
• Sr: Crimson red
• Ba: Apple green

Chemical Properties

Alkaline earth metals are reactive metals, but they are less reactive than alkali metals.

1. Reactivity: Reactivity increases down the group from Be to Ra. Beryllium and magnesium are less reactive than calcium, strontium, and barium.

2. Formation of $+2$ Ions: They readily lose their two valence electrons to form stable $+2$ cations ($M^{2+}$). Their electropositive character increases down the group.

3. Action with Air (Oxygen):

• Beryllium and Magnesium form a protective layer of oxide on their surface when heated in air, which prevents further reaction.
• Calcium, Strontium, and Barium are readily attacked by air, forming oxides and some nitride.

$2M(s) + O_2(g) \rightarrow 2MO(s)$

$3M(s) + N_2(g) \rightarrow M_3N_2(s)$

4. Action with Water:

• Beryllium is almost inert to water due to the formation of a protective oxide layer. It reacts with steam at high temperatures.
• Magnesium reacts slowly with cold water but vigorously with hot water or steam.

$Mg(s) + 2H_2O(l) \rightarrow Mg(OH)_2(s) + H_2(g)$

• Calcium, Strontium, and Barium react readily with cold water to form hydroxides and hydrogen gas. Reactivity increases down the group.

$M(s) + 2H_2O(l) \rightarrow M(OH)_2(aq) + H_2(g)$

5. Action with Halogens: They react directly with halogens on heating to form ionic halides ($MX_2$).

$M(s) + X_2 \rightarrow MX_2(s)$

6. Action with Acids: They react with acids to liberate hydrogen gas, except beryllium, which reacts slowly.

$M(s) + 2HCl(aq) \rightarrow MCl_2(aq) + H_2(g)$

7. Decomposition of Nitrates and Carbonates:

• Nitrates: Thermal decomposition of nitrates.
  $2M(NO_3)_2(s) \xrightarrow{heat} 2MO(s) + 4NO_2(g) + O_2(g)$
  ($Be$ and $Mg$ nitrates give oxide, while $Ca, Sr, Ba$ nitrates give oxide and $NO_2$, $O_2$).
• Carbonates: Thermal decomposition of carbonates.
  $MCO_3(s) \xrightarrow{heat} MO(s) + CO_2(g)$
  The stability of carbonates increases down the group. $BeCO_3$ and $MgCO_3$ decompose easily, while $BaCO_3$ decomposes only at very high temperatures.

8. Solubility of Hydroxides, Carbonates, and Sulfates:

• Hydroxides: Solubility increases down the group ($Be(OH)_2$ and $Mg(OH)_2$ are sparingly soluble, while $Ca(OH)_2$ is slightly soluble, and $Sr(OH)_2, Ba(OH)_2$ are moderately soluble). Basicity also increases down the group.
• Carbonates: Solubility decreases down the group ($BeCO_3$ and $MgCO_3$ are soluble, $CaCO_3$ is insoluble, $SrCO_3$ and $BaCO_3$ are insoluble).
• Sulfates: Solubility decreases down the group ($BeSO_4$ and $MgSO_4$ are soluble, $CaSO_4$ is sparingly soluble, $SrSO_4$ and $BaSO_4$ are insoluble).

Reducing Nature: They are strong reducing agents, but less strong than alkali metals. Reducing power increases from Be to Ra.