Representative Metals, Metalloids and Nonmetals

The main groups are the two tall columns on the left of the periodic table and the six on the right. They hold the elements you meet every day. Table salt, lime, sand, glass, rust proofing, fertilizer, bleach and battery acid all come from them. This chapter walks group by group through those elements. You will see why sodium explodes in water while magnesium barely fizzes. You will see why carbon can be the hardest solid or the slipperiest one. And you will see why the noble gases sit out almost every reaction. The rules behind all of it come from two trends you already know.

Section 1Periodicity of the main groups

One outer shell, one personality

The main group elements fill groups 1, 2 and 13 through 18. They are also called the representative elements. What ties a group together is its outer shell. Every group 1 atom has one electron in its outer s orbital. Every group 17 atom has seven outer electrons. That count is the valence electron count. It sets the charges an element forms. It also sets how many bonds the atom makes.

Two trends explain most of this chapter. The first is size. The second is how tightly the nucleus grips the outer electrons. Chemists measure that grip with effective nuclear charge, written Z_eff. It is the charge an outer electron really feels. Inner electrons screen part of the true nuclear charge.

Across a period the grip tightens

Walk from sodium to chlorine. Each step adds one proton and one outer electron. The new electron joins a shell that is already there, so it screens badly. Z_eff climbs at every step. Atoms shrink and hold their electrons harder. Ionization energy rises from 496 kJ/mol for sodium to 1251 kJ/mol for chlorine. Electronegativity rises as well, from 0.93 to 3.16.

So the left side hands electrons out and the right side takes them in. That shift is a loss of metallic character. Sodium is a soft shiny metal. Silicon is a metalloid, shiny but a weak conductor. Chlorine is a gas that grabs electrons from almost anything.

Down a group the grip loosens

Now move from lithium to cesium. Each step adds a whole new shell. The outer electron sits farther out, and more core electrons screen it. The atomic radius grows from 152 pm to 265 pm. Ionization energy drops from 520 kJ/mol to 376 kJ/mol. A looser grip means the atom gives an electron up more easily. So metallic character grows down every main group.

Group 14 shows this best. Carbon at the top is a nonmetal. Silicon and germanium are metalloids. Tin and lead at the bottom are metals soft enough to hammer flat. One group, three kinds of element, and the only change is shell count.

Oxides give the game away

You can place an element by how its oxide treats water. A basic oxide comes from a metal and makes a base:

An acidic oxide comes from a nonmetal and makes an acid:

Between them sits the amphoteric oxide, which reacts with acid and with base. Aluminum oxide is the textbook case. It dissolves in strong acid and in strong base, and it shrugs off water:

Section 2Alkali and alkaline earth metals

Group 1: one electron to lose

The alkali metals are lithium, sodium, potassium, rubidium, cesium and francium. Each one has a single outer s electron. Each one loses it to make a +1 ion. None of them is ever found free in nature. All are soft enough to cut with a knife, and all are stored under oil. Air and water attack them on contact.

In water they all do the same thing. They make hydrogen gas and a hydroxide:

The equation is the same for every member. The speed is not. Lithium fizzes steadily. Sodium melts into a ball and darts across the surface. Potassium lights its own hydrogen with a lilac flame. Cesium reacts so fast it can crack the glass. Ionization energy explains the ladder. It takes 520 kJ/mol to strip lithium and only 376 kJ/mol to strip cesium.

Oxygen gives three different products, and which one you get depends on the metal. Lithium makes a plain oxide. Sodium makes a peroxide, with an O–O bond. Potassium and the heavier metals make a superoxide:

Lithium has one more trick. It is the only group 1 metal that attacks nitrogen gas at room temperature:

Sodium metal itself is made by electrolysis. Molten sodium chloride runs in a Downs cell at about 600 °C. Calcium chloride is mixed in to drop the melting point. Sodium forms at the cathode and chlorine at the anode:

Group 2: two electrons, a tighter hold

The alkaline earth metals are beryllium, magnesium, calcium, strontium, barium and radium. Each has two outer s electrons and forms a +2 ion. Pulling two electrons off costs much more than pulling one off. So group 2 metals are harder, denser and far less reactive than their group 1 neighbours.

Water shows the gap clearly. Beryllium does not react at all, even in steam. Magnesium needs steam to get going. Calcium works away in cold water at a steady pace:

Magnesium burns in air with a white glare hot enough to melt steel. That is why it once filled flash bulbs and still fills emergency flares:

The lime cycle

Limestone is calcium carbonate. Heat it past about 900 °C in a kiln and it gives up carbon dioxide. The product is quicklime, calcium oxide:

Add water to quicklime and it hisses, swells and gets hot. That gives slaked lime, calcium hydroxide:

Lime mortar then closes the loop. Over months it pulls carbon dioxide back out of the air and turns to stone again. Farmers spread lime on sour soil, and power plants use it to trap sulfur dioxide from flue gas. Hard water is the same chemistry working against you. Dissolved Ca²⁺ and Mg²⁺ ions leave scale in kettles and pipes.

Section 3Boron, aluminum, carbon and silicon

Boron is short of electrons

Boron is the only member of group 13 that is not a metal. It is a hard, dark metalloid that melts near 2075 °C. It has three valence electrons, which is one short of a filled set of four bonds. That shortage runs all of its chemistry. Boron compounds hunt for electron pairs, so they act as Lewis acids. Boron trifluoride, BF₃, grabs a pair from almost any donor.

Diborane, B₂H₆, shows the shortage best. It has 12 valence electrons but needs 16 for eight normal bonds. Two hydrogens solve this by bridging. Each bridge is a three-center two-electron bond, with one pair shared across B–H–B. Boron is electron deficient, not broken.

Boric acid, H₃BO₃, is a mild antiseptic and a very weak acid. It does not hand over its own proton. It takes a hydroxide from water instead, which leaves an H⁺ behind:

Borax, Na₂B₄O₇·10H₂O, is the mined ore. Boron oxide, B₂O₃, goes into borosilicate glass, which barely expands when heated.

Aluminum is common, armoured and two-faced

Aluminum is the most plentiful metal in the crust, near 8% by mass. It is never free, because it holds oxygen too tightly. Pulling it out needs electricity. In the Hall–Héroult process, alumina is dissolved in molten cryolite, Na₃AlF₆, near 960 °C. Carbon anodes burn away as the oxide gives up its oxygen:

Fresh aluminum looks like it should corrode away. It does not, because the surface grows a thin tight film of Al₂O₃ that seals the metal below. Anodizing thickens that film on purpose. Aluminum is also an amphoteric metal. It dissolves in strong acid and in strong base, giving hydrogen both times:

Aluminum wants oxygen so badly that it will steal it from other metal oxides. The thermite reaction uses that hunger to weld rail track in the field:

Carbon: one atom, many solids

Carbon forms more allotropes than any other element. An allotrope is a different solid form of the same element. In diamond every carbon bonds to four others at 109.5°, with bonds 154 pm long. The whole crystal is one giant molecule. Nothing can slide and no electron is free, so diamond is the hardest natural solid and a poor conductor.

In graphite each carbon bonds to only three others in a flat sheet. Those bonds are 142 pm long and very strong. The fourth electron sits in a p orbital above and below the sheet, and it roams. Sheets lie 335 pm apart and are held by weak forces only. So graphite conducts along its sheets and slides apart in your pencil.

Two newer allotropes matter too. A single graphite sheet on its own is graphene. Roll it up and you get a nanotube. Buckminsterfullerene, C₆₀, is a hollow ball of 20 hexagons and 12 pentagons.

Silicon builds the ground you stand on

Silicon is second only to oxygen in the crust, near 28% by mass. You never find it free. Every silicon atom in rock is tied to oxygen in a silicate. The unit is a tetrahedron, one silicon in the middle with four oxygens at the corners. Those tetrahedra share corners, and how many they share decides the mineral.

Carbon and silicon sit in the same group, so why is CO₂ a gas and SiO₂ a rock? Carbon is small, so its p orbitals overlap sideways well and give strong double bonds. One carbon with two double bonds is a complete molecule. Silicon is bigger and that sideways overlap is poor. So silicon makes four single bonds to oxygen instead, and the solid never ends.

Pure silicon starts as sand reduced with coke in an arc furnace. Chip grade silicon needs a second pass through a liquid:

The trichlorosilane is distilled, then broken back down to deposit pure silicon. Chips need fewer than one stray atom per billion.

Section 4Nitrogen and phosphorus

Nitrogen is locked in a triple bond

Nitrogen gas is 78% of the air you breathe, yet almost nothing reacts with it. The reason is one number. The N≡N triple bond is worth 945 kJ/mol and is only 110 pm long. It is one of the strongest bonds known. Breaking it is the whole problem of nitrogen fixation, which means turning N₂ into a form life can use.

Industry solves it with the Haber process. Nitrogen from air meets hydrogen from steam reforming over an iron catalyst:

The forward reaction gives out heat and shrinks the gas volume. So cold and squeezed should give the best yield. But cold is also slow. Real plants run near 450 °C and 200 atm, which is a deal struck between yield and speed. Ammonia is pulled out as a liquid as it forms, which drags the balance forward.

Ammonia itself is a trigonal pyramid with H–N–H angles near 107°. It boils at −33 °C, so it is easy to store under mild pressure. In water it is a weak base:

The nitrogen oxides

Nitrogen forms an oxide at every oxidation state from +1 to +5. Dinitrogen oxide, N₂O, is the dental gas. Nitrogen monoxide, NO, is a colourless gas with an odd number of electrons, which makes it a radical. Your own blood vessels use it as a signal. Nitrogen dioxide, NO₂, is the brown haze over a busy city. It is also a radical, and it pairs up when cooled:

Brown NO₂ fades to colourless N₂O₄ in an ice bath. Warm the tube and the colour floods back. It is one of the clearest equilibrium demonstrations there is.

Nitric acid from ammonia

The Ostwald process turns ammonia into nitric acid in three steps. A platinum and rhodium gauze does the first one near 850 °C:

The NO from the last step is piped back and used again. Nitric acid is a strong acid, but it is also a strong oxidizer, and that second job matters more. Copper will not touch hydrochloric acid. It dissolves in nitric acid at once, because the nitrate ion does the work. The product depends on how concentrated the acid is:

Phosphorus comes in three solids

White phosphorus is built from P₄ molecules shaped like a tiny tetrahedron. The bond angles are forced to 60°, far from the 109.5° the atom wants. That strain makes it wildly reactive. White phosphorus catches fire in air near 34 °C, so it is kept under water. Heat it gently and it turns into red phosphorus, a chain polymer safe enough for match strikers. Heat it under pressure and you get black phosphorus, a layered solid much like graphite.

Phosphorus burns in air to give one of two oxides. Starve it of oxygen and you get P₄O₆. Give it plenty and you get P₄O₁₀, a white powder so thirsty that chemists use it to dry other gases:

The phosphorus acids

Phosphoric acid, H₃PO₄, gives up three protons, one at a time, each much harder than the last. Its three acid constants are 7.5 × 10⁻³, 6.2 × 10⁻⁸ and 4.2 × 10⁻¹³. The middle one sits close to blood pH, which is why phosphate is a body buffer. Most of it is made by pouring sulfuric acid on phosphate rock:

Now the trap. Phosphorous acid is written H₃PO₃, so it looks triprotic. It is not. One of its three hydrogens is bonded straight to phosphorus, and a P–H hydrogen never leaves in water. Only the two O–H hydrogens are acidic, so H₃PO₃ is diprotic. Hypophosphorous acid, H₃PO₂, has two P–H bonds and gives up just one proton. The rule to carry away is simple. Count the O–H groups.

Section 5Oxygen and sulfur

Oxygen has a second form

Oxygen gas is 21% of the air by volume. Its double bond is worth 498 kJ/mol, just over half of the nitrogen triple bond, which is why oxygen reacts so readily. O₂ also has two unpaired electrons, so a pool of liquid oxygen clings to a magnet. Industry gets it by chilling air until it liquefies, then boiling the nitrogen off at −196 °C and the oxygen at −183 °C.

Oxygen also comes as ozone, O₃, an allotrope with a bent shape and a 117° angle. Both O–O bonds are 128 pm, which sits between a single bond and a double bond. Neither bond is whole, because the real structure is an average of two drawings. Making ozone costs energy:

High in the air, ultraviolet light pays that bill and ozone shields the ground. At ground level the same gas is a lung irritant. Position, not chemistry, decides whether ozone helps you.

Oxide, peroxide, superoxide

Oxygen is almost always −2, but not always. In a peroxide the ion is O₂²⁻ and each oxygen is −1. In a superoxide the ion is O₂⁻ and each oxygen is −½. Adding electrons fills antibonding levels, so the bond order drops and the O–O distance grows.

These ions are useful. Sodium peroxide gives hydrogen peroxide with water. Potassium superoxide gives oxygen, which is why it fills emergency breathing canisters in mines and submarines:

Hydrogen peroxide is caught in the middle of its own range, so it can be oxidized or reduced. Left alone it does both at once and tears itself apart. A speck of manganese dioxide, or the enzyme in a cut finger, makes the split instant:

Sulfur: rings that open

Sulfur sits under oxygen but behaves very differently. It is too big for strong double bonds, so instead of S₂ it builds eight atom rings shaped like a crown. Solid S₈ is rhombic below 95.3 °C and monoclinic above it, and it melts near 115 °C. Keep heating and the rings snap open into long chains. The liquid gets darker and thicker, which is odd for a hot liquid. Pour it into cold water and the tangle freezes as rubbery plastic sulfur.

Most sulfur now comes out of oil and gas as hydrogen sulfide. The Claus process turns that waste into solid sulfur:

Hydrogen sulfide is a weak diprotic acid with K_a1 near 1.0 × 10⁻⁷. It smells of rotten eggs at tiny amounts. At higher amounts it numbs the nerve that detects it, so the warning disappears while the danger grows.

The contact process

Sulfuric acid is made in larger amounts than any other chemical. The route is the contact process, and it has four steps. Burn sulfur, oxidize the dioxide over a catalyst, absorb the trioxide into strong acid, then add water:

Only the second step is reversible, and it is the one that needs care. Cool gas favours SO₃ but reacts too slowly, so the plant settles near 430 °C with a vanadium oxide catalyst. The third step looks like a detour. It is there because SO₃ met with water throws a fine acid mist that no tower can trap. The oleum route keeps the product in the liquid.

Concentrated sulfuric acid is about 98% by mass, which works out near 18 M. It is a strong acid on the first proton and a weak one on the second, with K_a2 = 1.2 × 10⁻². It is also a powerful drying agent. Pour it on table sugar and it strips out the water, leaving a black column of carbon. Sulfur dioxide gives the weaker sulfurous acid, which bleaches paper, preserves dried fruit and falls as acid rain:

Section 6Halogens and noble gases

The halogens are one electron short

The halogens are fluorine, chlorine, bromine, iodine and astatine. Each has seven outer electrons, written ns²np⁵. Each needs one more to fill the shell, so each is a hungry oxidizer. All five exist as X₂ molecules, and the way they look changes right down the group. Fluorine is a pale yellow gas, chlorine a yellow green gas, bromine a red brown liquid and iodine a grey violet solid.

Their appetite for electrons is measured by the standard reduction potential. Fluorine leads at +2.87 V, then chlorine at +1.36 V, bromine at +1.07 V and iodine at +0.54 V. That order lets you predict every displacement reaction. Any halogen pushes a halide ion from lower in the list out of its compound:

Fluorine does not stop at halides. It takes electrons from water itself:

Chlorine is gentler. In water it splits its own oxidation states, a move called disproportionation. In base the same reaction runs further and gives household bleach:

That bond energy chart holds a surprise. The F–F bond is only 159 kJ/mol, weaker than the Cl–Cl bond at 242 kJ/mol. Fluorine atoms are so small that their lone pairs crowd each other and push the bond apart. A bond that cheap to break, plus a huge gain from forming F⁻, is why fluorine attacks almost everything.

Halogen acids and oxyacids

Three of the four hydrogen halides are strong acids. Hydrofluoric acid is the odd one out. Its K_a is 6.8 × 10⁻⁴, because the H–F bond is worth 569 kJ/mol and does not want to break. Weak does not mean safe. HF soaks through skin and attacks bone, and it is the only common acid that eats glass:

Chlorine also makes a set of oxyacids, and each extra oxygen makes the acid stronger. Hypochlorous acid is very weak, with K_a = 2.9 × 10⁻⁸. Chlorous acid is far stronger at 1.1 × 10⁻². Chloric and perchloric acids are both strong, and perchloric is one of the strongest acids known. The reason is the leftover ion. Extra oxygens pull electron density away and spread the negative charge, so the ion holds together without the proton.

Interhalogens

Halogens also bond to each other. An interhalogen has the formula AB, AB₃, AB₅ or AB₇. The larger, less electronegative halogen always sits in the middle, because only a big atom has room for several neighbours. Chlorine monofluoride, ClF, is linear. Bromine trifluoride, BrF₃, is T shaped. Iodine pentafluoride, IF₅, is a square pyramid. Only iodine is wide enough to hold seven fluorines, so IF₇ is the one AB₇ that exists.

These compounds are savage fluorinating agents. Chlorine trifluoride sets fire to glass wool, sand and concrete. It was tested as a rocket oxidizer and then quietly dropped.

The noble gases hold back

The noble gases are helium, neon, argon, krypton, xenon and radon. Each has a filled outer shell, so each has nothing to gain by bonding. That shows up in their ionization energies, which are the highest in every period. Helium needs 2372 kJ/mol. Argon needs 1521 kJ/mol. All six are gases that boil far below room temperature, and all are single atoms rather than pairs.

Still, the heavier ones do react. Xenon needs only 1170 kJ/mol to lose an electron, which is about the same as an oxygen molecule at 1175 kJ/mol. Since oxygen can be forced to give up an electron, xenon should be able to as well. That single comparison led to the first noble gas compound. Fluorine is the partner, and the amount you use sets the product:

XeF₂ is linear and XeF₄ is square planar, both shaped by extra lone pairs on xenon. Water turns the fluorides into xenon trioxide, which is a violent explosive:

Their everyday jobs come from not reacting. Helium lifts balloons and cools magnets. Argon fills the space around a weld so the hot metal cannot meet oxygen. Neon glows red orange in a tube, and krypton and xenon run bright lamps and camera flashes.

ReferenceWords worth knowing

These are the words this chapter leans on. Come back to them before a quiz.

alkali metal

A group 1 element. One outer s electron, a +1 ion, and a fast reaction with water.

alkaline earth metal

A group 2 element. Two outer s electrons and a +2 ion. Harder and slower to react than group 1.

allotrope

A different solid form of the same element. Diamond and graphite are allotropes of carbon.

amphoteric

Able to react as an acid and as a base. Al₂O₃ and BeO both qualify.

contact process

The four step industrial route to sulfuric acid: burn sulfur, oxidize SO₂ over V₂O₅, absorb SO₃ into acid, then add water.

cryolite

Na₃AlF₆. Molten cryolite dissolves alumina so aluminum can be won at about 960 °C instead of 2072 °C.

disproportionation

One substance is oxidized and reduced in the same reaction. Chlorine in water gives HOCl and HCl.

effective nuclear charge

Written Z_eff. The share of the nuclear charge an outer electron actually feels after inner electrons screen it.

Haber process

Making ammonia from N₂ and H₂ over iron, near 450 °C and 200 atm.

halogen

A group 17 element. Seven outer electrons, an X₂ molecule, and a strong pull for one more electron.

interhalogen

A compound of two different halogens, such as BrF₃ or IF₇. The larger halogen sits in the middle.

metallic character

How readily an element gives electrons up and behaves as a metal. It rises down a group and falls across a period.

metalloid

An element on the staircase border with some metal and some nonmetal behaviour. Boron, silicon, germanium, arsenic, antimony and tellurium.

nitrogen fixation

Turning unreactive N₂ into usable compounds such as ammonia or nitrate.

noble gas

A group 18 element. A filled outer shell, a very high ionization energy, and single atoms rather than molecules.

oleum

H₂S₂O₇, made by feeding SO₃ into concentrated sulfuric acid. It avoids the acid mist that water would cause.

Ostwald process

Making nitric acid from ammonia in three steps, through NO and then NO₂.

peroxide

A compound holding the O₂²⁻ ion, with oxygen at −1 and a single O–O bond.

silicate

A mineral built from SiO₄ tetrahedra sharing corners. Sharing none gives separate ions, sharing all four gives quartz.

superoxide

A compound holding the O₂⁻ ion, with oxygen at −½. KO₂ is the common example.