s-Block Elements (Group 1 and Group 2 elements)

The Alkali Metals comprise the six elements of Group 1 of the periodic table: Lithium ($\text{Li}, Z=3$), Sodium ($\text{Na}, Z=11$), Potassium ($\text{K}, Z=19$), Rubidium ($\text{Rb}, Z=37$), Caesium ($\text{Cs}, Z=55$), and Francium ($\text{Fr}, Z=87$, radioactive). They are designated "alkali" metals because their oxides and hydroxides dissolve in water to produce strongly alkaline (basic) caustic solutions.

• Atomic and Ionic Radii: Alkali metals possess the largest atomic and ionic radii in their respective periods. Down the group from $\text{Li}$ to $\text{Cs}$, radii increase progressively due to the addition of successive principal electron shells: $$\text{Atomic Radii: } \text{Li} < \text{Na} < \text{K} < \text{Rb} < \text{Cs}$$ Upon losing the single valence $ns^1$ electron, the unipositive cation ($M^+$) has a substantially smaller radius than the parent neutral atom because the nuclear charge pulls the remaining electron core inward.
• Ionization Enthalpies ($\Delta_i H$): The first ionization enthalpies of alkali metals are the lowest among all elements in the periodic table, decreasing smoothly down the group ($\text{Li} > \text{Na} > \text{K} > \text{Rb} > \text{Cs}$) due to increasing atomic radius and superior screening (shielding) of the valence electron by inner closed shells.
  Second Ionization Enthalpy ($\Delta_i H_2$): The second ionization enthalpies are astronomically high because removing a second electron requires disrupting a stable, closed noble-gas octet configuration ($[\text{Ne}], [\text{Ar}]$, etc.). Consequently, alkali metals strictly exhibit a $+1$ oxidation state in all their chemical compounds.
• Density & The Potassium Anomaly: Because of their large atomic volumes, alkali metals have unusually low densities compared to other metals; $\text{Li, Na},$ and $\text{K}$ are all lighter than water!
  Density generally increases down the group from $\text{Li}$ to $\text{Cs}$, but there is a famous anomaly: Potassium ($\text{K}$) is lighter than Sodium ($\text{Na}$)! $$\text{Density Order: } \text{Li} (0.53\text{ g/cm}^3) < \mathbf{\text{K}} (0.86\text{ g/cm}^3) < \mathbf{\text{Na}} (0.97\text{ g/cm}^3) < \text{Rb} (1.53\text{ g/cm}^3) < \text{Cs} (1.90\text{ g/cm}^3)$$ Explanation: Going from $\text{Na}$ to $\text{K}$, there is an unusually large increase in atomic volume (from $23.7\text{ cm}^3/\text{mol}$ to $45.9\text{ cm}^3/\text{mol}$) caused by the presence of empty $3d$-orbitals in potassium core structure. The percentage increase in volume outweighs the increase in atomic mass, making potassium less dense than sodium.

When gaseous alkali metal cations enter an aqueous medium, they become surrounded by polar water molecules through ion-dipole electrostatic attractions, releasing hydration enthalpy ($\Delta_{\text{hyd}} H^\circ$):

When an alkali metal or its volatile salt (typically chloride, moistened with concentrated $\text{HCl}$) is introduced into the non-luminous oxidizing zone of a Bunsen flame, the thermal energy of the flame excites the loosely bound outermost $ns^1$ valence electron to higher electronic energy levels.

When this excited electron falls back to the ground state, the absorbed excess energy is emitted as electromagnetic radiation in the visible region of the spectrum ($\Delta E = h\nu = hc/\lambda$). Because the energy level spacing is unique for each element, each alkali metal produces a characteristic flame emission color:

Alkali metals tarnish rapidly in moist air due to the spontaneous formation of surface oxide, hydroxide, and carbonate coatings. When burned in excess air or pure dioxygen, different alkali metals yield distinctly different classes of oxygen compounds:

Why does the stability of peroxides and superoxides increase as we descend Group 1?

• Reactivity with Water: Alkali metals react violently and exothermically with water to form basic metal hydroxides and liberate flammable hydrogen gas: $$2M(s) + 2\text{H}_2\text{O}(l) \to 2M^+(aq) + 2\text{OH}^-(aq) + \text{H}_2(g) \uparrow$$ The reaction becomes increasingly violent down the group. Sodium melts into a silvery globule darting across the water surface, while potassium catches fire spontaneously with a lilac flame.
• Reducing Nature: Alkali metals are among the strongest reducing agents known. In aqueous solution, reducing power is governed by standard reduction potentials ($E^\circ$).
  Surprisingly, Lithium is the strongest reducing agent in aqueous solution ($E^\circ = -3.05\text{ V}$), despite having the highest first ionization enthalpy!
  Thermodynamic Reason: Standard electrode potential is the sum of three thermodynamic steps: $$\text{Sublimation: } M(s) \xrightarrow{\Delta_{\text{sub}} H} M(g)$$ $$\text{Ionization: } M(g) \xrightarrow{\Delta_i H} M^+(g) + e^-$$ $$\text{Hydration: } M^+(g) + \text{aq} \xrightarrow{\Delta_{\text{hyd}} H} M^+(aq)$$ Although lithium has high sublimation and ionization enthalpies, its exceptionally large, negative hydration enthalpy ($\Delta_{\text{hyd}} H^\circ \approx -506\text{ kJ/mol}$) due to its tiny ionic size easily compensates, giving $\text{Li}$ the most negative $E^\circ$ in the entire periodic table.

Alkali metals dissolve smoothly in anhydrous liquid ammonia without chemical evolution of hydrogen, producing remarkable solutions with unique physical properties:

As the first member of Group 1, lithium differs substantially from its congeners ($\text{Na, K, Rb, Cs}$). These anomalies stem directly from three structural attributes:

1. Exceptionally Small Atomic and Ionic Size: Bare atomic radius is only $152\text{ pm}$ and $\text{Li}^+$ ionic radius is merely $76\text{ pm}$.
2. Enormous Polarizing Power: According to Fajans' Rules, polarizing power is proportional to charge density ($\text{Charge} / \text{Radius}^2$). The high polarizing power of $\text{Li}^+$ distorts the electron clouds of large anions, imparting pronounced covalent character to lithium compounds.
3. High Electronegativity & Ionization Enthalpy: Electronegativity is $1.0$ (highest among alkali metals) and $\Delta_i H$ is $520\text{ kJ/mol}$.

In the periodic table, elements of Period 2 often exhibit striking chemical resemblances to the element situated diagonally to their right in Period 3. Lithium ($\text{Li}$, Group 1, Period 2) displays a famous diagonal relationship with Magnesium ($\text{Mg}$, Group 2, Period 3):

1. Direct Nitride Synthesis: Both $\text{Li}$ and $\text{Mg}$ react directly with nitrogen gas at elevated temperatures to form ionic nitrides: $$6\text{Li} + \text{N}_2 \to 2\text{Li}_3\text{N}; \quad 3\text{Mg} + \text{N}_2 \to \text{Mg}_3\text{N}_2$$ Both nitrides hydrolyze in water to liberate ammonia gas: $\text{Li}_3\text{N} + 3\text{H}_2\text{O} \to 3\text{LiOH} + \text{NH}_3 \uparrow$.
2. Thermal Decomposition of Carbonates: Both carbonates decompose upon heating, releasing carbon dioxide: $$\text{Li}_2\text{CO}_3 \xrightarrow{\Delta} \text{Li}_2\text{O} + \text{CO}_2 \uparrow; \quad \text{MgCO}_3 \xrightarrow{\Delta} \text{MgO} + \text{CO}_2 \uparrow$$
3. Nitrate Decomposition Behavior: Both nitrates decompose on heating to yield metal oxides, reddish-brown $\text{NO}_2$ gas, and $\text{O}_2$: $$4\text{LiNO}_3 \xrightarrow{\Delta} 2\text{Li}_2\text{O} + 4\text{NO}_2 + \text{O}_2; \quad 2\text{Mg}(\text{NO}_3)_2 \xrightarrow{\Delta} 2\text{MgO} + 4\text{NO}_2 + \text{O}_2$$
4. Deliquescence & Hydration of Chlorides: Both $\text{LiCl}$ and $\text{MgCl}_2$ are deliquescent, separating from aqueous solution as hydrated crystals ($\text{LiCl}\cdot 2\text{H}_2\text{O}$ and $\text{MgCl}_2\cdot 6\text{H}_2\text{O}$), and both are appreciably soluble in organic solvents (alcohol, pyridine).
5. Low Solubility of Fluorides, Carbonates & Phosphates: Unlike the salts of other alkali metals, the fluorides ($\text{LiF}, \text{MgF}_2$), carbonates ($\text{Li}_2\text{CO}_3, \text{MgCO}_3$), and phosphates ($\text{Li}_3\text{PO}_4, \text{Mg}_3(\text{PO}_4)_2$) of lithium and magnesium are sparingly soluble in water.

The Alkaline Earth Metals occupy Group 2 of the periodic table, consisting of Beryllium ($\text{Be}, Z=4$), Magnesium ($\text{Mg}, Z=12$), Calcium ($\text{Ca}, Z=20$), Strontium ($\text{Sr}, Z=38$), Barium ($\text{Ba}, Z=56$), and Radium ($\text{Ra}, Z=88$, radioactive). They were named alkaline earths because their oxides ("earths" in historical alchemical terminology) are basic and found in mineral earth deposits.

• Atomic and Ionic Radii: In any given period, the atomic and ionic radii of alkaline earth metals are smaller than those of the corresponding alkali metals because of increased nuclear charge ($Z$) pulling the electron cloud inward: $$\text{Radii: } \text{Be} < \text{Mg} < \text{Ca} < \text{Sr} < \text{Ba}$$
• Ionization Enthalpies:
  • First Ionization Enthalpy ($\Delta_i H_1$): Higher than that of Group 1 metals due to smaller atomic size and the extra stability of a completely filled $ns^2$ subshell.
  • Second Ionization Enthalpy ($\Delta_i H_2$): Much smaller than the second ionization enthalpy of alkali metals because removing the second electron yields a stable noble-gas configuration. Hence, alkaline earth metals form divalent $M^{2+}$ cations rather than monovalent $M^+$ cations.
• Hydration Enthalpies: Because of their doubly positive charge ($+2$) and smaller ionic radii, the hydration enthalpies of alkaline earth metal ions are roughly four to five times larger than those of alkali metal ions! $$\Delta_{\text{hyd}} H^\circ: \mathbf{\text{Be}^{2+} > \text{Mg}^{2+} > \text{Ca}^{2+} > \text{Sr}^{2+} > \text{Ba}^{2+}}$$ Consequently, Group 2 salts are extensively hydrated compared to Group 1 salts: e.g., $\text{MgCl}_2\cdot 6\text{H}_2\text{O}$ and $\text{CaCl}_2\cdot 6\text{H}_2\text{O}$, whereas $\text{NaCl}$ and $\text{KCl}$ crystallize as anhydrous salts.
• Hardness, Melting & Boiling Points: Due to two valence electrons participating in metallic bonding and smaller interatomic distances, alkaline earth metals possess much stronger metallic bonds, resulting in significantly higher densities, hardness, melting points, and boiling points than alkali metals.

When salts of $\text{Ca, Sr,}$ and $\text{Ba}$ are vaporized in a Bunsen flame, electrons are thermally promoted and emit characteristic visible light upon relaxation:

• Reactivity Towards Oxygen & Air:
  • $\text{Be}$ and $\text{Mg}$ are kinetically inert at room temperature because a thin, impervious, protective film of oxide forms on their surface. When ignited, $\text{Mg}$ burns with a blinding white light forming $\text{MgO}$ and $\text{Mg}_3\text{N}_2$.
  • $\text{Ca, Sr,}$ and $\text{Ba}$ burn readily in air to form normal monoxides ($MO$). Heavy barium also forms peroxide ($\text{BaO}_2$).
  • Basicity of Oxides: $\text{BeO}$ is amphoteric; $\text{MgO}$ is weakly basic; $\text{CaO, SrO,}$ and $\text{BaO}$ are strongly basic, with basicity increasing down the group: $$\text{Basicity: } \mathbf{\text{BeO (Amphoteric)} < \text{MgO} < \text{CaO} < \text{SrO} < \text{BaO (Strongest Base)}}$$
• Reactivity Towards Water:
  • $\text{Be}$ does not react with water or steam even at red heat.
  • $\text{Mg}$ reacts very slowly with cold water, but decomposes steam rapidly: $\text{Mg} + \text{H}_2\text{O}(g) \to \text{MgO} + \text{H}_2 \uparrow$.
  • $\text{Ca, Sr,}$ and $\text{Ba}$ react vigorously with cold water: $$M(s) + 2\text{H}_2\text{O}(l) \to M(\text{OH})_2(aq) + \text{H}_2(g) \uparrow$$
  • Solubility of Hydroxides: Increases down the group ($\text{Be(OH)}_2 < \text{Mg(OH)}_2 < \text{Ca(OH)}_2 < \text{Sr(OH)}_2 < \text{Ba(OH)}_2$) because lattice energy decreases much faster than hydration energy.
  • Solubility of Sulfates: Decreases down the group ($\text{BeSO}_4 > \text{MgSO}_4 \gg \text{CaSO}_4 > \text{SrSO}_4 > \text{BaSO}_4$) because the large sulfate anion maintains almost constant lattice energy, while cation hydration enthalpy drops dramatically from $\text{Be}^{2+}$ to $\text{Ba}^{2+}$.

Beryllium ($\text{Be}$) differs sharply from the rest of Group 2 due to its exceedingly small atomic/ionic size ($r_{\text{Be}^{2+}} = 31\text{ pm}$), high electronegativity ($1.5$), and lack of available $d$-orbitals, limiting its maximum covalency to 4. Instead, Beryllium displays a striking diagonal relationship with Aluminum ($\text{Al}$):

1. Action of Nitric Acid (Passivation): Both $\text{Be}$ and $\text{Al}$ are rendered completely passive by concentrated nitric acid ($\text{HNO}_3$) due to the immediate formation of an impervious, protective oxide film on their surfaces.
2. Amphoteric Oxides & Hydroxides: Unlike other Group 2 basic oxides, $\text{BeO}$ and $\text{Be(OH)}_2$ are amphoteric, dissolving in both acids and strong alkalis, exactly like $\text{Al}_2\text{O}_3$ and $\text{Al(OH)}_3$: $$\text{Be(OH)}_2 + 2\text{OH}^- \to [\text{Be(OH)}_4]^{2-} \quad (\text{Beryllate ion})$$ $$\text{Al(OH)}_3 + \text{OH}^- \to [\text{Al(OH)}_4]^- \quad (\text{Aluminate ion})$$
3. Chloro-Bridged Halide Structures:
   • In the vapor phase, $\text{BeCl}_2$ exists as a chloro-bridged planar dimer ($\text{Cl}-\text{Be}-(\mu\text{-Cl})_2-\text{Be}-\text{Cl}$), and in the solid state as a polymeric chain of coordinate chloro-bridges.
   • Similarly, aluminum chloride exists as a chloro-bridged dimer ($\text{Al}_2\text{Cl}_6$) in the vapor phase and non-polar organic solvents. Both act as strong Lewis acids and catalyze Friedel-Crafts alkylations!
4. Carbide Hydrolysis: The carbides of both beryllium ($\text{Be}_2\text{C}$) and aluminum ($\text{Al}_4\text{C}_3$) are methanides; they react with water to yield pure methane gas ($\text{CH}_4$): $$\text{Be}_2\text{C} + 4\text{H}_2\text{O} \to 2\text{Be(OH)}_2 + \text{CH}_4 \uparrow; \quad \text{Al}_4\text{C}_3 + 12\text{H}_2\text{O} \to 4\text{Al(OH)}_3 + 3\text{CH}_4 \uparrow$$

The industrial synthesis of Sodium Carbonate (Washing Soda, $\text{Na}_2\text{CO}_3 \cdot 10\text{H}_2\text{O}$) relies on the Solvay Process, pioneered by Ernest Solvay in 1861. The process exploits the low aqueous solubility of sodium hydrogen carbonate ($\text{NaHCO}_3$) in the presence of concentrated sodium chloride brine.