Main Group Chemistry

The representative (main-group) elements are those in which the s and p orbitals are filling — groups 1, 2, and 13–18. Their chemistry follows directly from their position in the periodic table, so once you know where an element sits you can predict how it behaves.

Two trends do most of the work. First, metallic character increases down a group and decreases from left to right across a period: the alkali metals at the lower left are the most metallic and reactive, while the nonmetals at the upper right are the least metallic. Second, the group number gives the characteristic ion charge — group 1 forms 1+ ions, group 2 forms 2+, group 13 forms 3+, group 16 forms 2−, and group 17 forms 1−. Combining a cation and an anion in the ratio that makes the compound neutral gives the formula (for example Na⁺ and O²⁻ give Na₂O; Al³⁺ and Se²⁻ give Al₂Se₃).

The alkali metals are named for the fact that they and their oxides react with water to form strongly basic (alkaline) solutions; the alkaline-earth metals are named the same way for their heavier oxides.

To compare elements quantitatively, use the periodic trends in size and in how tightly an atom holds its electrons. Atomic radius increases down a group and decreases across a period. Ionization energy and electronegativity decrease down a group and increase across a period.

These trends let you rank a set of elements or pick the “more” member of a pair. Within the halogens, for instance, fluorine (top of group 17) has the highest electronegativity and the highest first ionization energy. Reactivity of the metals runs the opposite way: a larger atom lower in a group holds its valence electron(s) more loosely, so it is oxidized more readily — rubidium reacts with water far more vigorously than sodium, and the alkali metals (group 1) are more reactive than the alkaline-earth metals (group 2) in the same period.

The same reasoning predicts metallic versus covalent behavior: elements low and to the left are the most metallic, and those high and to the right form the most covalent, molecular substances.

Covalent character within a family of ionic compounds is decided by the cation, not the element trend. A smaller, more highly charged cation (higher charge density) pulls on — polarizes — the electron cloud of its anion, sharing it partially and giving the bond covalent character. Among NaCl, CaCl₂, and BeCl₂, the tiny, doubly charged Be²⁺ is by far the most polarizing, so BeCl₂ is the most covalent of the three chlorides even though all three are metal–chloride compounds. Reach for this cation-polarizing-power reasoning whenever a question compares the covalency of salts of the same anion; the ordinary electronegativity trend compares elements, not cations.

The active metals of groups 1 and 2 share three characteristic reactions.

With water they give the metal hydroxide and hydrogen gas: 2M + 2H₂O → 2MOH + H₂ for group 1, and M + 2H₂O → M(OH)₂ + H₂ for group 2. The reaction is more vigorous for the larger, more reactive metals lower in each group.

With oxygen the group-1 product follows a trend set by cation size: lithium gives the simple oxide (Li₂O), sodium gives the peroxide (Na₂O₂), and potassium, rubidium, and cesium give the superoxide (for example KO₂). The alkaline-earth metals give the simple oxide MO.

With the halogens both groups give the ionic metal halide (for example 2Na + Cl₂ → 2NaCl; Sr + Br₂ → SrBr₂). In every case the metal is oxidized and the nonmetal is reduced; balancing the equation is just a matter of matching the ion charges and then conserving atoms.

Allotropes are different structural forms of the same element, and the structural difference is what drives the difference in properties. The recurring theme is the strength of sideways (π) overlap between p orbitals, which is strong for small period-2 atoms and weak for larger atoms below them.

Carbon: graphite is built from stacked sheets with a delocalized π system (soft, electrically conducting), while diamond is a three-dimensional σ network (very hard, insulating). Silicon, being larger, cannot sustain that π system and forms only the diamond-type network.

Nitrogen vs. phosphorus: small nitrogen makes a strong N≡N triple bond and exists as diatomic N₂; larger phosphorus has weak π overlap, so it forms single P–P bonds and builds the tetrahedral P₄ molecule (reactive white phosphorus) or polymeric red phosphorus.

Oxygen vs. sulfur: oxygen forms the double-bonded diatomic O₂ and the higher-energy allotrope ozone, O₃; larger sulfur cannot form strong π bonds and instead links into crown-shaped S₈ rings.

Hydrogen sits by itself: its oxidation state depends on its partner. It is +1 when bonded to a more electronegative nonmetal (as in H₂O or HCl) and −1 in ionic hydrides formed with active metals (as in NaH or CaH₂).

A defining part of hydrogen's chemistry is that it acts as a reducing agent. Hydrogen gas reduces the oxides of less-active metals to the metal or a lower oxide while being oxidized to water, for example H₂ + MnO₂ → MnO + H₂O.

An ionic hydride reacts with water to give the metal hydroxide plus hydrogen gas: CaH₂ + 2H₂O → Ca(OH)₂ + 2H₂. Here the hydride ion (H⁻) is both the Lewis base (it donates the electron pair) and the reducing species (it is oxidized from −1 to 0), while water acts as the Lewis acid and the oxidizing agent — the two H atoms of each product H₂ come from two different sources.

Three named processes turn main-group elements into some of the most heavily produced industrial chemicals, and each illustrates how reaction chemistry, kinetics, and economics jointly determine a process.

The Haber-Bosch process makes ammonia: N₂ + 3H₂ ⇌ 2NH₃ over an iron catalyst. The Ostwald process turns that ammonia into nitric acid in stages: 4NH₃ + 5O₂ → 4NO + 6H₂O, then 2NO + O₂ → 2NO₂, then 3NO₂ + H₂O → 2HNO₃ + NO. The contact process makes sulfuric acid by oxidizing sulfur dioxide over a vanadium(V) oxide catalyst, 2SO₂ + O₂ ⇌ 2SO₃, then absorbing the SO₃ into sulfuric acid to form oleum, which is diluted to H₂SO₄.

The reversible product-forming steps in Haber and the contact process are both exothermic and mole-reducing, so higher pressure and lower temperature favor the products at equilibrium. Their operating choices differ: Haber uses high pressure because the yield gain justifies compression, whereas the contact process operates near atmospheric pressure because SO₃ conversion is already high and further compression is not economical. Both use a moderate temperature and a catalyst to obtain a workable rate.

How a metal is won from its ore depends on how reactive it is. Very active metals are obtained by electrolytic reduction of a molten salt, because no ordinary chemical reductant can reduce them: the Downs cell electrolyzes molten NaCl to give sodium (2NaCl → 2Na + Cl₂), and the Hall-Héroult cell electrolyzes alumina dissolved in molten cryolite to give aluminum (cathode: Al³⁺ + 3e⁻ → Al). Less-active metals are obtained by chemical reduction with carbon or carbon monoxide (for example ZnO + C → Zn + CO; SnO₂ + 2C → Sn + 2CO).

Phosphorus is produced industrially by the same carbon-reduction idea applied to a nonmetal: phosphate rock is heated with sand and coke, 2Ca₃(PO₄)₂ + 6SiO₂ + 10C → 6CaSiO₃ + 10CO + P₄, and the volatile P₄ distills off. Burning that phosphorus gives the acidic oxide P₄O₁₀ (P₄ + 5O₂ → P₄O₁₀), which reacts with water to give phosphoric acid: P₄O₁₀ + 6H₂O → 4H₃PO₄.

One more lever: removing a product can drive an otherwise unfavorable reduction. Calcium normally cannot reduce Cs⁺ (cesium is the more reactive metal), yet Ca + 2CsCl → CaCl₂ + 2Cs works industrially because at the operating temperature the cesium is volatile and distills out of the melt — continuously removing a product pulls the equilibrium forward (Le Chatelier), just as the escape of volatile P₄ drives the phosphorus furnace.

A reliable way to work main-group problems:

1. Read the position. Use the group number for the ion charge and the up/down, left/right trends for metallic character, size, ionization energy, and electronegativity.
2. Match the reaction to the family. Active metal + water → hydroxide + H₂; metal + oxygen → the size-appropriate oxide/peroxide/superoxide; metal + halogen → ionic halide.
3. Let structure explain properties. When comparing allotropes, argue from the strength of π overlap (small atoms make strong multiple bonds; larger atoms make single-bonded networks or rings).
4. For industrial steps, name the process from its inputs and outputs, write the balanced equation, predict the equilibrium shift, and then distinguish that thermodynamic effect from the economic operating choice. Haber uses high pressure; the contact process operates near atmospheric pressure even though higher pressure would favor SO₃.

Common mistakes: forgetting that the group-1 oxygen product changes from oxide to peroxide to superoxide down the group; assuming every element forms the same allotrope as the one above it (silicon has no graphite analog; phosphorus has no stable P₂); confusing the direction of an equilibrium shift with the most economical plant conditions; and trying to reduce sodium or aluminum chemically with carbon when they require electrolysis.