Guides · 05
The p-Block: Metalloids, Halogens and Noble Gases
A tour of groups 13–18: from metals to nonmetals, metalloids in semiconductors, reactive halogens, and the noble gases that turned out to form xenon compounds.
The six columns on the right of the periodic table, groups 13 to 18, make up the p-block, where the last electron enters a p orbital. The s-block and d-block contain only metals, but the p-block contains metals, metalloids and nonmetals side by side. Carbon, nitrogen and oxygen, which make up living things, silicon for electronics and the argon in the air all live here, which makes it the most varied part of the table.
The big picture
Properties across the p-block change in two directions.
- Left to right: atoms increasingly tend to gain electrons rather than lose them, so metallic character falls and nonmetallic character rises.
- Top to bottom: atoms grow larger and lose their outer electrons more easily, so metallic character rises.
That is why the lower left of the p-block (aluminium, tin, lead, bismuth) is metallic, the upper right (nitrogen, oxygen, fluorine) is nonmetallic, and a diagonal band of metalloids runs between them from boron down toward astatine.
The number of outer-shell electrons follows from the group number: three in group 13 (s²p¹), four in group 14, five in group 15, six in group 16, seven in group 17 and eight in group 18 (helium has two).
Group 13: the boron group
Boron is a metalloid; aluminium, gallium, indium and thallium are metals. Aluminium is the most abundant metal in Earth's crust. It is light, and a thin, tight oxide film protects it from corrosion. Gallium melts at about 30 °C, low enough to melt in your hand, and gallium nitride is the key material in blue LEDs. Indium is used in indium tin oxide, the transparent conductor in touchscreens.
Moving down group 13, the +1 oxidation state becomes more stable relative to +3, and thallium commonly forms +1 compounds. This happens because in heavy atoms the outermost pair of s electrons is reluctant to take part in bonding, an effect known as the inert pair effect.
Group 14: the carbon group
Carbon is a nonmetal, silicon and germanium are metalloids, and tin and lead are metals. It is one of the clearest examples of a single group spanning the full range from nonmetal to metal.
Carbon can bond to itself in long chains and rings, forming the backbone of millions of organic compounds. Depending on how the atoms are linked, pure carbon can be diamond, graphite, graphene or fullerenes, allotropes with strikingly different properties. Silicon is the second most abundant element in the crust after oxygen, found mostly as silicon dioxide in sand and rock. Ultra-pure silicon crystals are the foundation of computer chips and solar cells. Lead, because of the inert pair effect, is more stable in the +2 state.
Group 15: the nitrogen group
Nitrogen makes up about 78 percent of the air by volume, yet the two atoms in an N₂ molecule are held by a triple bond so strong that the gas is quite unreactive. Phosphorus has several allotropes, including white, red and black phosphorus; white phosphorus is reactive enough to ignite spontaneously in air. Nitrogen and phosphorus are essential to DNA and ATP and are the main nutrients in fertilizers. Arsenic and antimony are metalloids and bismuth is a metal. Arsenic is notorious as a poison but also appears in electronic materials such as gallium arsenide.
Group 16: the oxygen group (chalcogens)
Oxygen is the most abundant element by mass in both Earth's crust and the human body. Besides O₂ it forms ozone, O₃, which absorbs ultraviolet light in the stratosphere. Sulfur is a yellow solid found around volcanoes and is the raw material for sulfuric acid. Selenium and tellurium show semi-metallic behavior, and polonium is a radioactive metal. The group name, chalcogens, comes from Greek words meaning "ore formers".
Group 17: the halogens
Fluorine, chlorine, bromine, iodine, astatine and tennessine form group 17. "Halogen" means "salt former", because these elements combine readily with metals to form salts such as sodium chloride.
A halogen atom has seven outer electrons and needs only one more to reach the stable arrangement of a noble gas, so it attracts electrons strongly. The appearance of the elements changes steadily down the group.
| Element | State at room temperature | Color |
|---|---|---|
| Fluorine (F₂) | gas | pale yellow |
| Chlorine (Cl₂) | gas | yellow-green |
| Bromine (Br₂) | liquid | red-brown |
| Iodine (I₂) | solid | dark violet-black (violet vapor) |
The molecules get larger down the group, so the forces between them grow and melting and boiling points rise.
Reactivity increases toward the top. Fluorine has the highest electronegativity of any element and reacts violently even with glass and water. Displacement reactions show the order clearly. Bubbling chlorine through a solution of potassium bromide releases bromine, and the solution turns orange:
Cl₂ + 2KBr → 2KCl + Br₂
Bromine cannot displace chloride in the reverse direction. The sequence F > Cl > Br > I reflects how strongly each element pulls in an extra electron.
Halogens are all around us: chlorine disinfects tap water and swimming pools, fluoride in toothpaste helps prevent tooth decay, and iodine is used as an antiseptic and is needed for thyroid hormones. The elemental gases and liquids themselves, however, are toxic and corrosive.
Group 18: the noble gases
Helium, neon, argon, krypton, xenon, radon and oganesson make up group 18. Their outer shells are full, and for a long time they were believed to form no compounds at all; they were often called the "inert gases".
That belief ended in 1962. Neil Bartlett, a British chemist working in Canada, had found that platinum hexafluoride (PtF₆) is a strong enough oxidizer to pull an electron from an oxygen molecule. Noticing that xenon's first ionization energy is almost the same as that of O₂, he reacted xenon with PtF₆ and obtained an orange-yellow solid, the first noble gas compound. In the same year another team made xenon tetrafluoride (XeF₄) directly from xenon and fluorine. Many more xenon compounds followed, including XeF₂, XeF₆ and the oxide XeO₃, and krypton was later shown to form compounds such as KrF₂.
The heavier noble gases have larger atoms whose outer electrons are farther from the nucleus, so their ionization energies are lower and compound formation becomes possible. No stable neutral compounds of helium or neon are known. Since these discoveries, the group is described as "noble" (very unreactive) rather than "inert" (completely unreactive).
Their uses all exploit that low reactivity:
- Helium: balloons and airships, and cooling the superconducting magnets in MRI scanners.
- Neon: the red-orange glow of neon signs.
- Argon: shielding gas for welding and filler gas for incandescent bulbs. At about 0.93 percent of the air, it is the most common noble gas.
- Krypton and xenon: high-intensity lamps, some car headlights, and propellant for spacecraft ion engines.
- Radon: a radioactive gas that seeps from soil and rock and can build up in basements, raising the risk of lung cancer.
Metalloids revisited
The p-block metalloids, boron, silicon, germanium, arsenic, antimony and tellurium, mix metallic and nonmetallic traits. Most importantly, their electrical conductivity depends strongly on temperature and on tiny amounts of impurities. Adding a little phosphorus to silicon gives an n-type semiconductor with extra electrons; adding boron gives a p-type semiconductor with a shortage of electrons. This technique, doping a group 14 element with group 13 or group 15 atoms, underlies every modern electronic device. The group numbers on the periodic table map directly onto the design rules of semiconductor engineering.