The s-Block Elements
The reactive metals of Groups 1 and 2 — soft, light and eager to lose electrons, colouring flames and powering everything from washing soda to the cells of your body
- The general character of the s-block — Group 1 alkali metals and Group 2 alkaline earth metals.
- The periodic trends in size, ionization enthalpy, reactivity and reducing power.
- Reactions with oxygen (oxide, peroxide, superoxide), water, hydrogen and halogens, and the flame colours.
- The anomalies of lithium and beryllium, and their diagonal relationships.
- Key sodium compounds and the Solvay process; key calcium compounds and plaster of Paris.
- The biological roles of \(\ce{Na+},\ \ce{K+},\ \ce{Mg^2+}\) and \(\ce{Ca^2+}\).
The s-Block at a Glance
The s-block holds the two leftmost groups, where the outermost electron occupies an \(s\) orbital. Group 1 (\(\ce{Li, Na, K, Rb, Cs, Fr}\)) — the alkali metals — have \(ns^1\); Group 2 (\(\ce{Be, Mg, Ca, Sr, Ba, Ra}\)) — the alkaline earth metals — have \(ns^2\). All are reactive metals that readily lose their outer electrons to reach a noble-gas core, giving fixed oxidation states of \(+1\) and \(+2\).
| Feature | Group 1 (alkali) | Group 2 (alkaline earth) |
|---|---|---|
| Valence config | \(ns^1\) | \(ns^2\) |
| Oxidation state | \(+1\) | \(+2\) |
| Hardness / density | very soft, low density | harder, denser |
| Melting point | low | higher than Group 1 |
| Reactivity | extremely high | high, but less than Group 1 |
Group 1: The Alkali Metals
The alkali metals are the most electropositive elements: large atoms with a single, loosely held electron and the lowest ionization enthalpies in their periods. Going down the group, atoms grow, ionization enthalpy falls, and reactivity rises. They are soft enough to cut with a knife and so light that lithium, sodium and potassium float on water.
| Property (down the group) | Trend | Reason |
|---|---|---|
| Atomic / ionic radius | increases | new shell each period |
| Ionization enthalpy | decreases | outer e⁻ farther, more shielded |
| Metallic / reducing character | increases | easier to lose the electron |
| Melting & boiling point | decreases | weaker metallic bonding |
| Hydration enthalpy | decreases | larger ions hydrate less |
The loosely held electron is easily excited by a flame and emits a characteristic colour as it falls back. This is the basis of the flame test and of sodium-vapour street lamps.
Reactions of the Alkali Metals
Their lone valence electron makes the alkali metals react vigorously with almost everything. The reaction with oxygen is especially revealing — the product changes down the group as the larger cations stabilise larger oxide anions.
Small \(\ce{Li+}\) stabilises only the small oxide ion \(\ce{O^2-}\); larger cations stabilise the bigger peroxide \(\ce{O2^2-}\) and superoxide \(\ce{O2-}\) ions.
| Reacts with | Equation | Note |
|---|---|---|
| Water | \(\ce{2Na + 2H2O -> 2NaOH + H2 ^}\) | vigour rises down group |
| Hydrogen | \(\ce{2Na + H2 -> 2NaH}\) | ionic hydride |
| Halogen | \(\ce{2Na + Cl2 -> 2NaCl}\) | ionic halide |
| Liquid ammonia | \(\ce{Na + (x+y)NH3 -> [Na(NH3)_x]+ + [e(NH3)_y]-}\) | deep-blue solution |
The Lithium Anomaly & Diagonal Relationship
Lithium, the first member, is the odd one out. Its very small size and high polarising power (charge density) give it more covalent character and a chemistry closer to magnesium — the classic diagonal relationship.
| Lithium (unlike its group) | ...resembles magnesium |
|---|---|
| forms only the oxide \(\ce{Li2O}\) (no peroxide) | \(\ce{Mg}\) also forms only \(\ce{MgO}\) |
| \(\ce{LiF},\ \ce{Li2CO3},\ \ce{Li3PO4}\) sparingly soluble | corresponding \(\ce{Mg}\) salts also low-solubility |
| combines with \(\ce{N2}\) to give \(\ce{Li3N}\) | \(\ce{Mg}\) forms \(\ce{Mg3N2}\) |
| \(\ce{LiNO3 ->[\Delta] Li2O + NO2 + O2}\) | \(\ce{Mg(NO3)2}\) decomposes similarly |
Compounds of Sodium & the Solvay Process
Sodium gives some of the most important industrial chemicals. Sodium carbonate (washing soda, \(\ce{Na2CO3.10H2O}\)) is made by the Solvay (ammonia-soda) process, in which cheap brine and limestone are converted using ammonia as a recyclable go-between.
Sparingly soluble \(\ce{NaHCO3}\) precipitates out and is filtered, then heated to \(\ce{Na2CO3}\). Ammonia is recovered from \(\ce{NH4Cl}\) with lime and recycled. The process fails for \(\ce{K2CO3}\) because \(\ce{KHCO3}\) is too soluble to precipitate.
| Compound | Common name | Key use |
|---|---|---|
| \(\ce{NaOH}\) | caustic soda | soap, paper, rayon (Castner–Kellner electrolysis of brine) |
| \(\ce{Na2CO3.10H2O}\) | washing soda | glass, detergents, water softening |
| \(\ce{NaHCO3}\) | baking soda | baking, antacid, fire extinguishers |
| \(\ce{NaCl}\) | common salt | feedstock for most sodium chemicals |
Group 2: The Alkaline Earth Metals
The alkaline earth metals have two valence electrons and form \(+2\) ions. Compared with their Group 1 neighbours they are smaller, harder, denser and higher-melting, because the extra electron and higher nuclear charge tighten the metallic bonding. They are still reactive, but less so than the alkali metals.
Beryllium and magnesium hold their electrons too tightly to be excited by a Bunsen flame, so they give no flame colour — a useful diagnostic.
Trends & Reactions of Group 2
Down the group, reactivity rises just as in Group 1. The most exam-relevant trends concern the solubility and thermal stability of their compounds, which move in opposite directions.
| Down the group | Trend |
|---|---|
| Solubility of hydroxides | increases (\(\ce{Mg(OH)2}\) low, \(\ce{Ba(OH)2}\) high) |
| Solubility of sulphates / carbonates | decreases (\(\ce{BaSO4}\) insoluble) |
| Thermal stability of carbonates | increases |
| Reactivity with water | increases (\(\ce{Be}\) almost none, \(\ce{Ba}\) vigorous) |
The Beryllium Anomaly & Diagonal Relationship
Beryllium, like lithium, breaks ranks. Its tiny size, high electronegativity and lack of \(d\)-orbitals give it strong covalent character and a chemistry that mirrors aluminium diagonally below it.
| Beryllium (unlike its group) | ...resembles aluminium |
|---|---|
| \(\ce{Be(OH)2}\) is amphoteric | \(\ce{Al(OH)3}\) is amphoteric |
| \(\ce{BeCl2}\) is covalent & polymeric | \(\ce{AlCl3}\) is covalent (dimer \(\ce{Al2Cl6}\)) |
| dissolves in alkali giving \(\ce{[Be(OH)4]^2-}\) | \(\ce{Al}\) gives \(\ce{[Al(OH)4]-}\) |
| forms covalent, hydrolysing compounds | \(\ce{Al}\) salts hydrolyse too |
Compounds of Calcium
Calcium compounds run the construction industry. They begin with limestone \(\ce{CaCO3}\), which on heating gives quicklime \(\ce{CaO}\), and from there slaked lime, gypsum and cement.
Heating gypsum to about \(393\,\text{K}\) gives plaster of Paris, \(\ce{CaSO4.\tfrac12 H2O}\); adding water reverses the step, setting it back to a hard mass of gypsum — the basis of casts and moulds.
| Compound | Name | Use |
|---|---|---|
| \(\ce{CaO}\) | quicklime | cement, steel, drying agent |
| \(\ce{Ca(OH)2}\) | slaked lime | whitewash, \(\ce{CO2}\) test (limewater) |
| \(\ce{CaCO3}\) | limestone / marble | building, antacid, toothpaste |
| \(\ce{CaSO4.\tfrac12 H2O}\) | plaster of Paris | casts, moulds, statues |
Biological Importance
Four s-block ions are essential to life. Their roles flow directly from their charge, size and abundance.
| Ion | Role in the body |
|---|---|
| \(\ce{Na+}\) | nerve signal transmission, fluid balance, the \(\ce{Na}\)–\(\ce{K}\) pump |
| \(\ce{K+}\) | intracellular cation, muscle and nerve function |
| \(\ce{Mg^2+}\) | centre of chlorophyll, cofactor for many enzymes |
| \(\ce{Ca^2+}\) | bones and teeth, blood clotting, muscle contraction |
Putting It to Work
Problem. Predict the main product when lithium, sodium and potassium each burn in excess oxygen.
Solution. Cation size decides which oxide-anion is stabilised:
Problem. An unknown salt gives an apple-green flame. Which metal ion is present?
Solution. Apple-green is the signature of barium:
Problem. Arrange \(\ce{MgSO4},\ \ce{CaSO4},\ \ce{BaSO4}\) in order of decreasing solubility.
Solution. Group-2 sulphate solubility falls down the group:
Problem. In the Solvay process, what precipitates when \(\ce{NH4HCO3}\) reacts with brine, and what is it heated to give?
Solution. Sparingly soluble bicarbonate drops out, then decomposes:
Problem. Lithium forms a nitride directly with \(\ce{N2}\). Which Group 2 element behaves similarly, and why?
Solution. The Li–Mg diagonal relationship predicts the match:
Problem. Write the reaction by which gypsum is converted to plaster of Paris, and state how it sets.
Solution. Controlled heating removes most of the water of crystallisation:
Chapter Summary
Group 1 (\(ns^1\), +1) and Group 2 (\(ns^2\), +2) — reactive, electropositive metals.
Down the group: size ↑, IE ↓, reactivity ↑; Li is the strongest aqueous reductant.
Li → oxide, Na → peroxide, K/Rb/Cs → superoxide; flame colours identify the metal.
Li ↔ Mg and Be ↔ Al diagonal relationships; Be is amphoteric and covalent.
NaOH, washing soda, baking soda; the Solvay process makes \(\ce{Na2CO3}\).
Lime cycle, plaster of Paris, cement; \(\ce{Na, K, Mg, Ca}\) are vital to life.
Problems
For each item, first decide which idea it tests — a periodic trend, a reaction, an anomaly, or a named compound — then apply the relevant rule. Difficulty rises down the list.
- State the general electronic configuration of Group 1 and Group 2 elements.
- Why are the alkali metals the most electropositive elements in their periods?
- Explain why lithium is the strongest reducing agent in aqueous solution despite its high ionization enthalpy.
- Write the products formed when Li, Na and K each react with excess oxygen.
- Give the flame colours of \(\ce{Na},\ \ce{K},\ \ce{Ca}\) and \(\ce{Ba}\).
- List three ways lithium resembles magnesium (the diagonal relationship).
- Write the three main steps of the Solvay process and explain why it fails for \(\ce{K2CO3}\).
- Arrange \(\ce{Mg(OH)2},\ \ce{Ca(OH)2},\ \ce{Ba(OH)2}\) in order of increasing solubility.
- Why do beryllium and magnesium not impart colour to a flame?
- Explain why \(\ce{Be(OH)2}\) and \(\ce{BeCl2}\) are described as anomalous, relating them to aluminium.
- Describe the lime cycle and write the reaction for the limewater test for \(\ce{CO2}\).
- How is plaster of Paris prepared from gypsum, and why must the temperature be controlled?