The S-Block Elements (Group 1 - Alkali Metals)

Group 1 Elements: Alkali Metals

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The Group 1 elements, also known as the alkali metals, consist of Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Cesium (Cs), and Francium (Fr). They are highly reactive metals with characteristic properties.

Electronic Configuration

General Configuration: Alkali metals have the general outer electronic configuration $ns^1$, where $n$ is the principal quantum number of the outermost shell.

• Li: $[He] 2s^1$
• Na: $[Ne] 3s^1$
• K: $[Ar] 4s^1$
• Rb: $[Kr] 5s^1$
• Cs: $[Xe] 6s^1$
• Fr: $[Rn] 7s^1$

Significance: The presence of a single valence electron in the outermost shell is responsible for many of their characteristic properties, particularly their high reactivity and tendency to form $+1$ ions.

Atomic And Ionic Radii

Atomic Radii:

• Trend: Atomic radii increase gradually on moving down the group from Li to Cs.
• Reason: As we move down the group, the number of electron shells increases. The outermost electron is further from the nucleus, and the inner electrons shield the valence electron from the nuclear charge more effectively. This results in a larger atomic size.

Ionic Radii:

• Trend: Ionic radii of alkali metal ions ($M^+$) also increase down the group (e.g., $Li^+ < Na^+ < K^+ < Rb^+ < Cs^+$).
• Reason: Similar to atomic radii, the increase in the number of electron shells leads to larger ionic radii.
• Comparison: The ionic radii ($M^+$) are smaller than the corresponding atomic radii ($M$) because the removal of the outermost electron reduces the electron-electron repulsion and increases the effective nuclear charge per electron.

Ionization Enthalpy

Definition: Ionization enthalpy is the energy required to remove the outermost electron from a gaseous atom.

Trend: Alkali metals have very low first ionization enthalpies. This is because they have only one valence electron, which is relatively far from the nucleus and well-shielded by inner electrons.

Trend Down the Group: Ionization enthalpies decrease from Li to Cs.

Reason: Down the group, the atomic size increases, and the shielding effect of inner electrons increases. This weakens the attraction between the nucleus and the valence electron, making it easier to remove.

Significance: The low ionization enthalpy explains why alkali metals readily lose their valence electron to form $+1$ cations ($M^+$).

Hydration Enthalpy

Definition: Hydration enthalpy is the energy released when gaseous ions are dissolved in water.

Trend: The hydration enthalpies of alkali metal ions decrease from $Li^+$ to $Cs^+$.

Reason: Hydration enthalpy is primarily influenced by the charge density of the ion (charge/size ratio). Smaller ions have higher charge density, leading to stronger electrostatic attraction with polar water molecules and thus higher hydration enthalpy (more energy released).

Order: $| \Delta H_{hydration}(Li^+) | > | \Delta H_{hydration}(Na^+) | > | \Delta H_{hydration}(K^+) | > | \Delta H_{hydration}(Rb^+) | > | \Delta H_{hydration}(Cs^+) |$

Impact on Properties: The high hydration enthalpy of $Li^+$ contributes to the unusually high ionization enthalpy of lithium and the high solubility of lithium salts.

Physical Properties

Appearance: Alkali metals are silvery-white, soft metals.

Softness: They are very soft and can be easily cut with a knife. Their softness increases down the group (Li is hardest, Cs is softest).

Low Melting and Boiling Points: They have relatively low melting and boiling points compared to other metals, which is due to the weak metallic bonding (only one valence electron per atom contributing to the 'sea' of electrons).

Density: They are generally light metals with low densities. Their density decreases slightly down the group from Li to K, then increases from Rb to Cs (due to the large increase in atomic size outweighing the increase in mass for Rb and Cs).

Flame Coloration: Alkali metals and their salts impart characteristic colors to a flame due to the excitation and emission of electrons. These are used for qualitative analysis:

• Li: Crimson red
• Na: Golden yellow
• K: Lilac
• Rb: Reddish violet
• Cs: Blue

Chemical Properties

Alkali metals are characterized by their high reactivity, primarily due to their low ionization enthalpies and electropositive nature.

1. Reactivity: Their reactivity increases significantly on moving down the group from Li to Cs. This is because the ionization enthalpy decreases, making it easier to lose the valence electron.

2. Formation of Ionic Compounds: They readily lose their valence electron to form $+1$ cations ($M^+$), which form ionic compounds with non-metals.

3. Action with Air (Oxygen):

• Lithium reacts slowly with oxygen in air to form lithium oxide ($Li_2O$).
• Sodium reacts vigorously with oxygen in air to form sodium peroxide ($Na_2O_2$) and a small amount of sodium superoxide ($NaO_2$).
• Potassium, Rubidium, and Cesium react rapidly with oxygen in air to form superoxides ($MO_2$) and oxides.
• They tarnish rapidly in air due to oxide, nitride, and carbonate formation. Hence, they are usually stored under kerosene or inert liquid.

4. Action with Water:

• Alkali metals react vigorously with water to liberate hydrogen gas and form the corresponding metal hydroxide (an alkali).
• Their reactivity with water increases down the group.
• The reaction is highly exothermic and can ignite the hydrogen gas produced.

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

• Lithium reacts moderately.
• Sodium reacts vigorously and melts into a spherical ball, moving on the surface.
• Potassium, Rubidium, and Cesium react so violently that the hydrogen produced ignites.

5. Action with Halogens: They react vigorously with halogens to form ionic halides.

$2M(s) + X_2 \rightarrow 2MX(s)$ (where X = F, Cl, Br, I)

6. Formation of Hydrides: React with hydrogen upon heating to form ionic hydrides ($MH$).

$2M(s) + H_2(g) \rightarrow 2MH(s)$

7. Reducing Nature: They are strong reducing agents. Their reducing power increases down the group ($Li < Na < K < Rb < Cs$).

8. Liquid Ammonia Reactions: They dissolve in liquid ammonia to form a blue solution due to the formation of solvated electrons ($e^-(am)$). These solutions are good conductors of electricity. On standing, or with the addition of a catalyst (like iron), they slowly react to form the amide ($MNH_2$) and $H_2$ gas.

$M(s) + (x+y)NH_3 \rightarrow M^+(am) + e^-(am)$

$2M^+(am) + 2e^-(am) + 2NH_3 \rightarrow 2MNH_2(s) + H_2(g)$

Oxo-acids: Alkali metals do not form oxo-acids. However, their oxides and hydroxides are basic.