4.1 Properties and Trends of Group 1 and 2 Elements

Periodic Trends

Atomic Properties and Electronic Structure

Group 1 (alkali metals) and Group 2 (alkaline earth metals) make up the s-block of the periodic table. Their shared feature is that their outermost electrons occupy s orbitals: Group 1 elements have an $ns^1$ configuration, and Group 2 elements have $ns^2$. This simple electronic structure drives most of their chemistry.

• Atomic radius increases down each group as new electron shells are added. Across a period (Group 1 → Group 2), radius decreases slightly because the extra proton in Group 2 pulls the same shell inward more tightly.
• Ionization energy decreases down each group because the valence electron sits farther from the nucleus and is better shielded by inner electrons. Group 2 elements have higher first ionization energies than their Group 1 neighbors because of the greater nuclear charge at similar size.
• Electronegativity follows the same trend as ionization energy: it increases across a period and decreases down a group. That said, all s-block metals have low electronegativities compared to the rest of the periodic table, which is why they form predominantly ionic compounds.
• Electron configurations follow the Aufbau principle, Hund's rule, and the Pauli exclusion principle. For example, potassium is $[Ar]\,4s^1$ and calcium is $[Ar]\,4s^2$.

Chemical Behavior and Relationships

Group 1 metals always adopt a +1 oxidation state, and Group 2 metals adopt +2. Removing additional electrons beyond these would mean breaking into a noble-gas core, which costs far too much energy.

Diagonal relationships are an important pattern in the s-block. Elements that sit diagonally adjacent in the periodic table can behave similarly because moving down a group increases size while moving right across a period increases charge. These effects roughly cancel, giving the diagonal pair a comparable charge density (charge-to-size ratio):

• Li and Mg: Both form nitrides directly with $N_2$, both have carbonates that decompose on heating, and both form organometallic compounds with significant covalent character.
• Be and Al: Both form amphoteric oxides and have a strong tendency toward covalent bonding due to high charge density.

Anomalous behavior of the first member is a recurring theme. Lithium and beryllium often deviate from the rest of their groups. Their small ionic radii and high charge densities give them unusually high polarizing power, which introduces covalent character into bonds that are fully ionic for the heavier members.

Group 1: Alkali Metals

Physical and Chemical Properties

The alkali metals (Li, Na, K, Rb, Cs, Fr) each have a single valence electron that is easily lost. They are soft, low-melting-point metals with low densities. Li, Na, and K are actually less dense than water.

• Reactivity increases down the group. As atomic radius grows, the valence electron is held less tightly, so ionization energy drops. Lithium reacts steadily with water; sodium reacts vigorously; potassium ignites on contact.
• Metallic bonding is weak because only one electron per atom contributes to the "electron sea." This explains the softness, low melting points, and low densities.
• Flame tests produce characteristic colors because thermal energy excites the single outer electron to a higher energy level. When it relaxes back, it emits a photon of a specific wavelength:

Aqueous Behavior and Applications

When alkali metals react with water, the products are a metal hydroxide and hydrogen gas:

$2M + 2H_2O \rightarrow 2MOH + H_2$

These hydroxides are strong bases, and their solubility increases down the group (LiOH is the least soluble; CsOH is the most).

Hydration enthalpy (the energy released when a gaseous ion is surrounded by water molecules) decreases down the group because larger ions have a lower charge density and attract water molecules less strongly. Lithium's small $Li^+$ ion actually has the highest hydration enthalpy in the group, which partly explains why many lithium salts are more soluble than you might expect from lattice energy arguments alone.

Key applications:

• Lithium: rechargeable batteries, psychiatric medication ($Li_2CO_3$)
• Sodium: sodium vapor lamps, $NaOH$ in industrial chemistry, $NaCl$ as a ubiquitous reagent
• Potassium: essential biological electrolyte for nerve impulse transmission, fertilizers ($KNO_3$)

Group 2: Alkaline Earth Metals

Chemical Characteristics and Reactivity

The alkaline earth metals (Be, Mg, Ca, Sr, Ba, Ra) have two valence electrons and form $M^{2+}$ cations. They are harder, denser, and have higher melting points than their Group 1 neighbors because two electrons per atom strengthen the metallic bond.

• Reactivity increases down the group, but is generally lower than the corresponding alkali metal in the same period. The higher nuclear charge and smaller radius of Group 2 elements mean more energy is needed to remove both valence electrons.
• Beryllium is the outlier. It does not react with water or steam under normal conditions, and its compounds are predominantly covalent. $BeCl_2$, for instance, is a linear molecule in the gas phase, not an ionic solid.
• Magnesium reacts very slowly with cold water but burns vigorously in steam. Calcium and the heavier members react readily with cold water.

Flame tests for Group 2:

Coordination Chemistry and Applications

Alkaline earth metal cations typically have coordination numbers of 4, 6, or 8, depending on the size of the cation and the ligands involved. Beryllium, being very small, strongly prefers a coordination number of 4 (tetrahedral geometry). The larger ions like $Ca^{2+}$, $Sr^{2+}$, and $Ba^{2+}$ can accommodate 6–8 or even more donor atoms.

$Ca^{2+}$ and $Mg^{2+}$ ions in natural water are responsible for water hardness. They react with soap (fatty acid salts) to form insoluble scum, and they precipitate as scale ($CaCO_3$) in pipes and kettles.

Key applications:

• Magnesium: lightweight structural alloys in aerospace and automotive industries
• Calcium: cement and concrete production (from $CaCO_3$), biological role in bones and teeth
• Strontium: red color in pyrotechnics ($SrCO_3$), formerly used in CRT television screens
• Barium: $BaSO_4$ as a radiocontrast agent in medical imaging (safe because of its extreme insolubility)