DAT General Chemistry: Alkali and Alkaline-Earth Metals

DAT High-Yield Concept

Group 1 (Alkali Metals): Valence configuration $ns^1$ , form $\ce{M^+}$ cations, soft, low melting points, and highly reactive.
Group 2 (Alkaline-Earth Metals): Valence configuration $ns^2$ , form $\ce{M^{2+}}$ cations, harder and denser than Group 1 metals.
Water Reactivity: All Group 1 metals react vigorously with water via $\ce{2M(s) + 2H2O(l) -> 2M^+(aq) + 2OH^-(aq) + H2(g)}$ ; Group 2 metals are less reactive ( $\ce{Be}$ is inert, $\ce{Mg}$ requires steam).
Periodic Trends: Reactivity increases down both groups as atomic radius increases and first ionization energy decreases.

1. Group 1: Alkali Metals vs. Group 2: Alkaline-Earth Metals

The s-block of the periodic table consists of the alkali metals (Group 1: $\ce{Li}$ , $\ce{Na}$ , $\ce{K}$ , $\ce{Rb}$ , $\ce{Cs}$ , $\ce{Fr}$ ) and alkaline-earth metals (Group 2: $\ce{Be}$ , $\ce{Mg}$ , $\ce{Ca}$ , $\ce{Sr}$ , $\ce{Ba}$ , $\ce{Ra}$ ). Group 1 elements contain a single valence electron ( $ns^1$ ), granting them low first ionization energies and a stable $\ce{+1}$ oxidation state. Group 2 elements contain two valence electrons ( $ns^2$ ) and form $\ce{+2}$ oxidation states. Highly electronegative nonmetals like halogens (Group 17, covered in depth in its own lesson) readily accept these valence electrons to form ionic salts.

Key Distinction: Successive Ionization Energies
Removing the first valence electron from an alkali metal requires low energy (

\text{IE}_1

). Removing a second electron requires breaking a noble-gas core shell, resulting in an immense jump in

\text{IE}_2

. In contrast, alkaline-earth metals readily lose both valence electrons because

\text{IE}_2

remains energetically accessible, forming stable

\ce{M^{2+}}

cations.

2. Periodic Trends in s-Block Metals

Two primary periodic trends govern s-block metal behavior: atomic radius and ionization energy.

Atomic Radius: Expands down both Group 1 and Group 2 because additional principal energy levels ( $n$ ) are filled. Across a period from Group 1 to Group 2, increasing effective nuclear charge ( $Z_{\text{eff}}$ ) pulls electrons closer, making Group 2 atoms smaller than adjacent Group 1 atoms.
Ionization Energy: Decreases down both groups because increasing atomic radius and shielding by core electrons reduce the electrostatic attraction between the nucleus and outermost valence electrons.

Because less energy is required to remove valence electrons from larger atoms, chemical reactivity increases down Groups 1 and 2.

Mnemonic for Reactivity: "Down the Group, Radius Grows, Valence Electron Goes!"
As atomic radius expands down a group, ionization energy drops, accelerating metal oxidation.

3. Reactions of Alkali Metals with Water

Alkali metals react exothermically with liquid water to form aqueous hydroxides and flammable hydrogen gas:

\ce{2M(s) + 2H2O(l) -> 2M^+(aq) + 2OH^-(aq) + H2(g)}

This redox process pairs metal oxidation ( $\ce{M(s) -> M^+(aq) + e^-}$ ) with water reduction ( $\ce{2H2O(l) + 2e^- -> 2OH^-(aq) + H2(g)}$ ). Reaction intensity escalates down Group 1: $\ce{Li}$ fizzes steadily, $\ce{Na}$ melts into a sphere and ignites, while $\ce{K}$ , $\ce{Rb}$ , and $\ce{Cs}$ react explosively upon contact with water at $\pu{25 degC}$ .

Group 2 metals are significantly less reactive toward water. $\ce{Be}$ does not react with water or steam. $\ce{Mg}$ reacts very slowly with cold water but reacts with steam ( $\ce{H2O(g)}$ ) to yield magnesium oxide ( $\ce{MgO}$ ) and $\ce{H2(g)}$ . Heavier Group 2 metals ( $\ce{Ca}$ , $\ce{Sr}$ , $\ce{Ba}$ ) react with cold liquid water to yield hydroxide solutions:

\ce{Ca(s) + 2H2O(l) -> Ca(OH)2(aq) + H2(g)}

4. Physical Properties and Emission Spectra

Alkali metals are unusually soft, easily sliced with a knife, and possess exceptionally low densities ( $\ce{Li}$ , $\ce{Na}$ , and $\ce{K}$ float on water). Heating s-block metals promotes valence electrons to higher energy levels. Upon relaxing back to ground state, they emit light at distinct wavelengths, producing characteristic flame spectra used for analytical identification.

5. Summary Comparison

Property	Group 1: Alkali Metals	Group 2: Alkaline-Earth Metals
Valence Electron Configuration	$ns^1$	$ns^2$
Common Oxidation State	$\ce{+1}$	$\ce{+2}$
First Ionization Energy ( $\text{IE}_1$ )	Lower	Higher
Atomic Radius (Same Period)	Larger	Smaller
Reactivity with Water	Vigorous/explosive at $\pu{25 degC}$	Inert ( $\ce{Be}$ ), steam only ( $\ce{Mg}$ ), moderate ( $\ce{Ca}$ - $\ce{Ba}$ )
Reactivity Trend Down Group	Increases ( $\ce{Li}$ < $\ce{Na}$ < $\ce{K}$ < $\ce{Rb}$ < $\ce{Cs}$ )	Increases ( $\ce{Be}$ < $\ce{Mg}$ < $\ce{Ca}$ < $\ce{Sr}$ < $\ce{Ba}$ )

6. Worked Quantitative Examples

Example 1: Stoichiometry of Alkali Water Reaction

Problem: Calculate the volume of $\ce{H2(g)}$ produced at $\pu{298 K}$ and $\pu{1.00 atm}$ when $\pu{4.60 g}$ of solid sodium ( $\ce{Na}$ , molar mass = $\pu{22.99 g/mol}$ ) reacts completely with excess water.

Dimensional Analysis Solution:

Calculate moles of $\ce{Na}$ : $\pu{4.60 g Na} \times \frac{\pu{1 mol Na}}{\pu{22.99 g Na}} = \pu{0.200 mol Na}$
Apply mole ratio from balanced equation $\ce{2Na(s) + 2H2O(l) -> 2NaOH(aq) + H2(g)}$ : $\pu{0.200 mol Na} \times \frac{\pu{1 mol H2}}{\pu{2 mol Na}} = \pu{0.100 mol H2}$
Calculate gas volume using $V = \frac{nRT}{P}$ : $V = \frac{(\pu{0.100 mol}) \times (\pu{0.08206 L atm mol^-1 K^-1}) \times (\pu{298 K})}{\pu{1.00 atm}} = \pu{2.45 L}$

DAT Trap: Forgetting the $2:1$ mole ratio between $\ce{Na}$ and $\ce{H2}$ in the balanced chemical equation. Assuming a 1:1 ratio incorrectly yields $\pu{4.89 L}$ .

Example 2: s-Block Reactivity Ranking

Problem: Rank $\ce{Mg(s)}$ , $\ce{Ca(s)}$ , $\ce{Na(s)}$ , and $\ce{K(s)}$ in order of increasing reactivity with water at $\pu{25 degC}$ .

Reasoning:

Group 1 metals are more reactive than Group 2 metals in the same period due to lower ionization energy.
Reactivity increases down a group as principal quantum number increases and atomic radius expands.
$\ce{Mg}$ (Group 2, Period 3) reacts sluggishly with liquid water. $\ce{Ca}$ (Group 2, Period 4) reacts moderately. $\ce{Na}$ (Group 1, Period 3) reacts vigorously. $\ce{K}$ (Group 1, Period 4) ignites spontaneously.

Answer: $\ce{Mg}$ < $\ce{Ca}$ < $\ce{Na}$ < $\ce{K}$ .

DAT Trap: Confusing kinetic reactivity with standard electrode potentials. Although lithium has the most negative reduction potential ( $E^\circ = \pu{-3.04 V}$ ) due to the high hydration energy of $\ce{Li^+}$ , its kinetic reaction with water is calmer than sodium or potassium due to its higher melting point and higher first ionization energy.

Critical Thinking Questions

Why does beryllium ( $\ce{Be}$ ) fail to react with liquid water or steam, whereas calcium ( $\ce{Ca}$ ) reacts readily with cold water?
How does core electron shielding explain why potassium ( $\ce{K}$ ) has a lower first ionization energy than sodium ( $\ce{Na}$ )?
If $\pu{0.10 mol}$ of $\ce{Ca(s)}$ is added to water versus $\pu{0.10 mol}$ of $\ce{Na(s)}$ , which reaction yields more moles of $\ce{H2(g)}$ ? Justify using balanced equations.

References

Sources returned by Google Search grounding during lesson verification.

7.7: Group Trends for the Active Metals - Chemistry LibreTexts — chem.libretexts.org
Reactions of Main Group Elements with Water - Chemistry LibreTexts — chem.libretexts.org

Alkali and Alkaline-Earth Metals

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