Group 3

Group 3 consists of boron, aluminium, gallium, indium, and thallium. Boron is a metalloid, but the other elements are metals. All elements in this group have three valence electrons.

Boron

Boron is a metalloid, exhibiting characteristics of both metals and non-metals. It is never found in its pure form in Earth, but it is a component of some minerals. However, pure boron can form as a result of chemical reactions. It can be a brown amorphous solid, or a crystalline solid. In the crystalline form, it is black and very hard.

Boron is a poor conductor of electricity at normal temperatures, but at high temperatures it does conduct electricity. Its chemical properties are mostly, but not entirely, non-metallic. Boron is the least electronegative non-metal, so it readily loses electrons when reacting. Boron is almost entirely in the +3 oxidation state. Boron has an oxide B₂O₃, and it is acidic.

${\hbox{BCl}}_{3}+3{\hbox{H}}_{2}{\hbox{O}}\to {\hbox{B}}({\hbox{OH}})_{3}+3{\hbox{HCl}}$	Boric halides are volatile substances that attack water with the formation of boric acid B(OH)₃ and hydrogen halides.
${\hbox{B}}({\hbox{OH}})_{3}+{\hbox{H}}_{2}{\hbox{O}}\to {\hbox{B}}({\hbox{OH}})_{4}^{-}+{\hbox{H}}^{+}$	Boric acid reacts with water to release a proton. It is a weak acid.
$2{\hbox{B}}_{2}{\hbox{O}}_{3}+7{\hbox{C}}\to {\hbox{B}}_{4}{\hbox{C}}+6{\hbox{CO}}$	Boric oxide, at extremely high temperatures, can react with carbon to form boron carbide. Boron carbide is incredibly strong, finding use as tank armour.

Boron has interesting properties when it bonds covalently. It has three electrons, so it usually forms three single covalent bonds, an exception to the octet rule. Substances like boron trichloride (BCl₃), boric acid, and the borate ion (BO₃^3-) are examples of boron's tendency to form three covalent bonds.

Aluminum

Aluminium (American spelling aluminum) is a metal, although it is too reactive to exist in pure form. Instead, it is found in a wide variety of minerals including bauxite, the primary source for extracting aluminium. Aluminium is the third most abundant element in Earth's crust.

Unlike boron, aluminium is an almost stereotypical metal in its physical properties: it is malleable and ductile, and it conducts heat and electricity very well. Because of its lightness, resistance to corrosion, attractiveness (approaching that of silver), and inexpensiveness to process along with adequate strength for many commercial and domestic uses, it is the third-most used of all metals. Only iron and zinc exceed it in production.

Aluminium compounds typically exist in the +3 oxidation state. Because of its durability and light weight, aluminium is used in many alloys, mixtures of metals. Two metals, each with desirable characteristics, can be mixed together to produce an alloy that has the properties of both metals. Aluminium used to be used for household wiring, but it expands too much when heated. The expansion causes the wires to come loose, and house fires can start.

Aluminium seems unreactive because of a thin, hard surface of aluminium oxide Al₂O₃ which does not react with water or oxygen and preserves the aluminium metal underneath it from chemical attack. Pure aluminium will immediately form this layer when exposed to air. Although aluminium is very conductive, aluminium oxide does not conduct electricity.

${\hbox{Al}}_{2}{\hbox{O}}_{3(s)}+3{\hbox{H}}_{2}{\hbox{O}}_{(l)}+2{\hbox{OH}}_{(aq)}^{-}\to 2{\hbox{Al}}({\hbox{OH}})_{4(aq)}^{-}$ ${\hbox{Al}}_{2}{\hbox{O}}_{3(s)}+6{\hbox{H}}_{(aq)}^{+}\to 2{\hbox{Al}}_{(aq)}^{3+}+3{\hbox{H}}_{2}{\hbox{O}}_{(l)}$	Aluminium oxide is amphoteric—it reacts violently to neutralise strong bases or strong acids.
$2{\hbox{Al}}+3{\hbox{Cl}}_{2}\to 2{\hbox{AlCl}}_{3}$	Reacting aluminium with a halogen results in an aluminium halide. Aluminium chloride can exist as an ionic solid or a covalent solid, but either way it is a weak conductor.
$2{\hbox{AlCl}}_{3(aq)}+3{\hbox{NaOH}}_{(aq)}\to {\hbox{Al}}({\hbox{OH}})_{3(aq)}+3{\hbox{NaCl}}_{(aq)}$	Aluminium chloride is an acid, able to neutralise bases.
${\hbox{Fe}}_{2}{\hbox{O}}_{3}+2{\hbox{Al}}\to 2{\hbox{Fe}}+{\hbox{Al}}_{2}{\hbox{O}}_{3}$	Molten aluminium reacts violently with some metal oxides in the thermite reaction, in which a substance like iron oxide is reduced to iron and aluminium is oxidized to its oxide. This is a highly exothermic reaction, releasing enough heat to weld metal.

Generally, aqueous aluminium is found as a hydroxide Al(OH)₃.

Others

Gallium is a brittle metal that melts slightly above room temperature, so holding it in your hand would liquefy it. One magic trick, the 'melting spoon' involves stirring a hot liquid (usually hot water) with a spoon specially made of gallium. The spoon melts in the water and 'disappears'. (The gallium simply sinks to the bottom of the cup; after the magic trick is done the 'magician removes the water and the molten gallium and pours the gallium into a mold that makes a new spoon for doing the trick again). Gallium is used in many semiconductor devices and LEDs (light emitting diodes). Gallium(III) arsenide is an important semiconductor used in many electronics.

Indium's most common isotope is radioactive and decays slowly into tin. Indium typically has the +3 oxidation state.

Thallium tends to appear in the +1 (thallous) oxidation state as well as the +3 (thallic) oxidation state. Oddly, thallium acts much like an alkali metal in its physical and chemical properties. Thallous oxide (Tl₂O) and the metal react violently with water to form a hydroxide TlOH that completely dissociates into Tl⁺ and OH^- ions as if it were an alkali metal hydroxide. Its halides are quite salt-like, dissociating into ions into solutions that, like solutions of salts of the alkali metals, conduct electricity.

Group 4

Group 4 consists of carbon, silicon, germanium, tin, and lead. Carbon is a non-metal, silicon and germanium are metalloids, and tin and lead are metals.

With 4 valence shell electrons, elements of the carbon family tend to form covalent compounds. With increasing mass and atomic radius these elements become increasingly metallic and have lower melting and boiling points.

Group 4 elements form gaseous hydrogen compounds with difficulty. These are either unstable or combustible. All but lead form oxides, sulfides, and halides in the +4 oxidation state. The +4 oxidation state predominates in carbon, silicon, and germanium; the +2 and +4 oxidation states both appear in tin, and the +2 oxidation state prevails in lead. Halides in the +4 state form for all of these elements, and they are covalent.

Carbon compounds are much more covalent than analogous compounds of silicon, germanium, tin, or lead. Even more significantly, carbon forms double and even triple bonds with itself or other elements, forming compounds that the heavier elements of this group cannot form like acetylene (C₂H₂). Silicon and the heavier elements of this group can form only single bonds.

Thus carbon dioxide CO₂ is a gas at normal temperatures because the double bonds between carbon and oxygen create single molecules, but silicon dioxide SiO₂ forms a hard rock known as quartz because it is a covalent network solid. Each silicon atom bonds to four different oxygen atoms with single bonds, and each oxygen atom bonds with two silicon atoms. Similar properties apply to the oxides of germanium, tin, and lead. Carbon dioxide dissolves in water to form carbonic acid, a weak acid that reacts with bases to form carbonates; oxides of the other elements of this group are practically nonreactive in water.

Carbon

Carbon is a very important element. It is abundant in the earth and atmosphere, and it is found in the substances that make all living things. Carbon has many properties that make it different from other elements, so it deserves thorough study.

Allotropes

Allotropes are different forms of a pure element. Carbon has several allotropes, three of which are common.

Amorphous carbon is coal and soot. The carbon molecules are covalently bonded, but there is no order or arrangement.
Graphite occurs when carbon forms flat covalent networks. These flat "sheets" are not bonded to each other, making them free to slide past each other. Graphite composes the "lead" in pencils.
Diamond occurs when carbon forms a three-dimensional covalent network. Diamonds are much different from graphite and amorphous carbon. They are transparent, brilliant-looking, and incredibly hard. Diamond only forms at heat and intense pressures.

There are also several rare and exotic allotropes of carbon, including:

Buckeyballs, or fullerenes, are spherical shaped balls of carbon. Covalent bonds join the carbon atoms into a soccer ball pattern that looks much like the geodesic domes of Buckminster Fuller. The most common buckeyball has the molecular formula C₆₀. Buckeyballs are large enough for a small atom to get trapped inside.
Lonsdaleite forms upon meteorite impact with the Earth.
Carbon nanotubes are incredibly small but rigid tubes made of carbon. They are created in labs and are a subject of research.

Keep in mind that allotropes are composed of only one element. In this case, these allotropes contain only carbon atoms and no other elements.

Inorganic Compounds

Although carbon is known mainly for its organic compounds, it does form many important inorganic compounds.

Oxides

Oxides of carbon contain only carbon atoms and oxygen atoms. There are two oxides that occur commonly:

Carbon monoxide (CO): a poisonous gas released when carbon-based fuels burn in limited oxygen.
Carbon dioxide (CO₂): found naturally in the air, but too much is considered pollution. Animals exhale carbon dioxide, and plants absorb it. It is slightly acidic. Carbon dioxide, when solid, is "dry ice".

There is other oxides that could form, but they are unstable or unnatural:

Carbon trioxide

Carbon suboxide (C₃O₂): Consists of double bonds, with oxygens at the ends. O = C = C = C = O. It breaks apart into carbon dioxide and dicarbon monoxide.
Dicarbon monoxide (C₂O): Very reactive. Contains only double bonds, but the end carbon has a non-bonding pair.
Carbon trioxide (CO₃): Exists in three different shapes, very unstable

Compounds Derived from Oxides

Carbonic acid forms when carbon dioxide is dissolved in water. It is given by the reaction:

{\hbox{CO}}_{2}+{\hbox{H}}_{2}{\hbox{O}}\leftrightarrow {\hbox{H}}_{2}{\hbox{CO}}_{3}

Carbonate and bicarbonate are two ions that carbon forms.

Their formulas are CO₃^2- and HCO₃^-, respectively.

Ionic Compounds

Along with the oxyanions (carbonate and bicarbonate), carbon can form several other ions.

Cyanide (CN^-)
Cyanate (OCN^-)
Thiocyanate (SCN^-)
Carbides (C₂^2- and C₃^4-)

Alloys

Carbon is used in some alloys, mixtures of metals. If a small amount (between 0.2 % and 2.1 % by weight) of carbon is mixed into iron, the result is steel.

Organic Compounds

Compounds containing carbon (except for the above inorganic compounds) are considered organic. They were once thought to be produced only by living things, but they have since been created in laboratories. Most organic compounds contain hydrogen as well as carbon.

Many substances are organic compounds. Polymers are organic compounds consisting of long chains of repeating patterns containing carbon and other atoms. Plastics, rubbers, and nylon are all organic polymers. Hydrocarbons are compounds containing only hydrogen and carbon, methane being a simple example. Crude oil is a sludge of various hydrocarbons mixed together. Propane, butane, and octane are well-known hydrocarbons used for fuel. Perhaps the most interesting type of organic compound is the biomolecule. Carbohydrates, proteins, lipids (fats), and nucleic acids (like DNA) are the most basic biomolecules. They, too, are polymers (except for lipids), being made of long chains of small, repeating chemicals that have bonded together. Biomolecules make up the chemicals found in the living cells that compose all living things.

Organic chemistry and biochemistry are very broad and thorough topics. They are far outside the scope of General Chemistry. Fortunately, your knowledge of General Chemistry is sufficient to begin the Wikibooks Organic Chemistry and Biochemistry if you are interested.

Silicon

Silicon is found in semiconductors, the basis of all electronic devices. Its electron configuration allows it to donate or accept electrons. When pure silicon is "doped" with elements that have more or fewer electrons than silicon, the slightly-impure silicon becomes a semiconductor. These substances make up complex electronics by acting like a switch that can turn on or off depending on electrical signals.

Silicon does not exist uncombined in nature; it most commonly occurs in silica (including the very common rock quartz and most sand grains) and in silicates. Most silicates are insoluble.

Silicon dioxide, a hard substance that melts only at high temperatures, is very different in its chemical properties from carbon dioxide, a gas until it freezes into dry ice. Molten silica can be cast as glass, a hard and useful material resistant to attack by almost all chemicals except fluorine, hydrofluoric acid, and strong alkalis. Glass is extremely useful in household containers and drinking utensils because it is resistant to chemicals, heat, and the attack of micro-organisms. Glass can be very clear if pure or containing certain chemicals, or it can take on attractive colours, making it a favourite material for art objects. Some small living creatures turn dissolved silica in the sea into their shells to create structure. A great variety of silicon compounds known as silicones have widespread and varied uses.

Others

Germanium is another element used in semiconductors. Before transistors were in common use, germanium diodes were heavily used in radios.

Tin is considered a "poor metal". It has two allotropes at STP: grey tin and white tin. Grey tin has non-metallic characteristics, but white tin is metallic. Tin is used frequently as an alloy. "Tin cans" are actually steel cans with a tin plating to resist corrosion. Pewter is an alloy of copper and tin, containing mostly tin. Bronze is an alloy of copper and tin containing mostly copper. Solder is an alloy of tin and lead used for its low melting point for attaching wires.

Lead is a heavy, grey metal. It had a tremendous number of uses, but it is now known to be a neurotoxin if ingested. Water-carrying pipes were made out of lead, but they are now made from copper or plastic because lead could contaminate the water. Lead compounds especially if soluble in water (like lead acetate used as a preservative and sweetener of wines in Roman times) or stomach acids (like the lead oxide once used in paints) or in gaseous or liquid form (like tetraethyl lead once used in gasoline), and powdered lead are very dangerous. Use of lead in foodstuffs, paint, and vehicle fuels is now illegal almost everywhere.

Lead-crystal glass has lead oxide in it but so tightly bound with silica that it can't escape easily. Pencil 'leads' are not lead at all, but instead harmless graphite (really carbon) bound with clay. Lead remains useful in automobile batteries (that contain necessary sulfuric acid much more dangerous than lead) and in nuclear use as a shield against radiation much more dangerous than lead metal.

$2{\hbox{PbO}}+{\hbox{PbS}}\to 3{\hbox{Pb}}+{\hbox{SO}}_{2}$	Lead oxides are easily reduced.
$3{\hbox{Pb}}_{(s)}+8{\hbox{HNO}}_{3(aq)}\to 3{\hbox{Pb}}_{(aq)}^{2+}+6{\hbox{NO}}_{3(aq)}^{-}+2{\hbox{NO}}_{(g)}+4{\hbox{H}}_{2}{\hbox{O}}_{(l)}$	Lead does not dissolve in hydrochloric acid or sulfuric acid, but it will dissolve in nitric acid because nitric acid is a strong oxidizer.
${\hbox{PbO}}_{(s)}+2{\hbox{OH}}_{(aq)}^{-}+{\hbox{H}}_{2}{\hbox{O}}_{(l)}\to {\hbox{Pb}}({\hbox{OH}})_{4(aq)}^{2-}$	Lead(II) oxide will form plumbite ions when added to basic solutions.
${\hbox{Pb}}({\hbox{OH}})_{4(aq)}^{2-}+{\hbox{Cl}}_{2(aq)}\to {\hbox{PbO}}_{2(s)}+2{\hbox{Cl}}_{(aq)}^{-}+2{\hbox{H}}_{2}{\hbox{O}}_{(l)}$	Plumbites form lead(IV) dioxide when chlorinated.
${\hbox{PbO}}_{2(s)}+2{\hbox{OH}}_{(aq)}^{-}+2{\hbox{H}}_{2}{\hbox{O}}_{(l)}\to {\hbox{Pb}}({\hbox{OH}})_{6(aq)}^{2-}$	Adding lead(IV) dioxide to a basic solution will form plumbate ions.

Group 5

Group 5 contains nitrogen, phosphorus, arsenic, antimony, and bismuth. Elements in Group 15 have five valence electrons. Because the elements can either gain three electrons or lose five to gain a stable configuration, they more often form covalent compounds unless bonded to an active metal. Their electron affinities are not very large. Metallic properties increase markedly from gaseous nitrogen to barely-metallic bismuth with increasing size and mass. Nitrogen and phosphorus are non-metallic, and arsenic and antimony are metalloids.

These elements are much less reactive than the elements of Group 16, and their chemistries are more complicated. Most of the chemistry of these elements is in the +3 or +5 oxidation states, although they form gaseous compounds with hydrogen in the -3 oxidation state: ammonia NH₃, phosphine PH₃, arsine AsH₃, stibine AsH₃, and bismuthine BiH₃; these all burn in oxygen to give oxides or the free element (in the case of nitrogen).

4{\hbox{NH}}_{3}+3{\hbox{O}}_{2}\to 6{\hbox{H}}_{2}{\hbox{O}}+2{\hbox{N}}_{2}

All form oxides—nitrogen with difficulty, the others with ease. Most of the oxides are acidic, exceptions being nitrous oxide N₂O, nitric oxide NO, and bismuth oxide Bi₂O₃. Except for nitrogen, typical oxides are in the +3 or +5 (bismuth excluded) oxidation states. All form halides—nitrogen with difficulty, but those of phosphorus, arsenic, and antimony fully hydrolyze in water.

Nitrogen and phosphorus form important acids in the +5 oxidation state. Nitrogen forms nitric acid HNO₃, a substance used to create medicines and explosives (but this acid is corrosive and dangerous, so don't touch it or even spill it on anything), and salts known as nitrates such as potassium nitrate KNO₃, an important fertilizer. Phosphorus forms phosphoric acid, H₃PO₄; phosphates are salts of phosphoric acid. Some phosphates are essential to respiration and thus life itself. General covalency of nitrogen family is 3.

Nitrogen

Nitrogen occurs naturally as the diatomic gas N₂. It comprises about 78% of the air we breathe. The bond holding the two nitrogen atoms together is triple covalent, so it is very strong. Because of that, nitrogen is very unreactive. It is used in many places when an inert gas is needed. However, nitrogen will react with some substances:

$6{\hbox{Li}}_{(s)}+{\hbox{N}}_{2(g)}\to 2{\hbox{Li}}_{3}{\hbox{N}}_{(s)}$ $3{\hbox{Mg}}_{(s)}+{\hbox{N}}_{2(g)}\to {\hbox{Mg}}_{3}{\hbox{N}}_{2(s)}$	Although nitrogen gas is usually considered inert, it does react with some elements by burning.
${\hbox{Li}}_{3}{\hbox{N}}_{(s)}+3{\hbox{H}}_{2}{\hbox{O}}\to 3{\hbox{LiOH}}_{(aq)}+{\hbox{NH}}_{3(g)}$ ${\hbox{Mg}}_{3}{\hbox{N}}_{2(s)}+6{\hbox{H}}_{2}{\hbox{O}}\to 3{\hbox{Mg}}({\hbox{OH}})_{2(aq)}+2{\hbox{NH}}_{3(g)}$	The nitrides react violently with water to form ammonia gas and a basic solution.

In its pure form, nitrogen is not very useful, and it will suffocate any animal that breathes pure nitrogen. It is much more important when it is a component of ammonia, nitrate, oxides, or biomolecules like protein. Due it is very unreactive nature, it is difficult to get nitrogen to react and form these useful substances. Any process that can convert elemental nitrogen into a nitrogen compound is called nitrogen fixation. Nitrogen fixation is biologically important because amino acids, proteins, and enzymes contain nitrogen. It is commercially important because it is used in explosives, rocket fuels, and fertilizers.

There are many nitrogen fixation reactions:

${\hbox{N}}_{2}+8{\hbox{H}}^{+}+8{\hbox{e}}^{-}+energy\to 2{\hbox{NH}}_{3}+{\hbox{H}}_{2}$	This occurs in bacterial enzymes. The ammonia (NH₃) quickly becomes ammonium (NH₄⁺). The nitrogen in the bacteria enters the soil where plants can absorb it. Humans and animals that eat those plants can get the nitrogen in useful compounds.
${\hbox{N}}_{2}+3{\hbox{H}}_{2}\leftrightarrow 2{\hbox{NH}}_{3}$	The Haber process is used for commercially producing ammonia. This reaction only occurs at very high pressures and temperatures (around 20 MPa and 500 °C) and in the presence of an iron catalyst. Also, the reaction occurs in somewhat complex equipment that must input pure reactants and extract the ammonia.

In the cold, dense, hydrogen-rich atmospheres of Jupiter, Saturn, Uranus, and Neptune, nitrogen ordinarily exists combined with hydrogen as ammonia.

Keep in mind that ammonia is a gas at STP. The household product called "ammonia" is actually an aqueous solution of ammonium hydroxide (NH₄OH) that forms when ammonia gas is dissolved in water. Ammonia, quite unlike hydrogen compounds of Groups 16 and 17, is a base in its reactions, forming salts with weak and strong acids alike. Such a substance as ammonium chloride (NH₄Cl) is a soluble, strongly ionic salt.

Nitrogen compounds are often extremely unstable because nitrogen atoms in nitrogen compounds tend to seek each other to recombine as nitrogen gas. Many nitrogen compounds are literal explosives, including TNT and nitroglycerin. These explosives are in common use in construction projects for the demolition of buildings and other obstacles to new construction, or to get access to minerals in mining operations.

Phosphorus

Phosphorus has two common allotropes: red phosphorus and white phosphorus. White phosphorus (P₄) has a waxy appearance and turns yellow when exposed to light. When exposed to oxygen in the dark, it glows pale green.

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White phosphorus ignites under all but the most delicate conditions. The combustion of white phosphorus produces phosphorus(V) oxide:

{\hbox{P}}_{4}+5{\hbox{O}}_{2}\to {\hbox{P}}_{4}{\hbox{O}}_{10}

One of its most common uses is in military weapons that cause severe burning of the object hit by the weapon.

Red phosphorus is an amorphous solid. It is more stable and explodes at temperatures higher than those of white phosphorus. It is still, however, dangerously reactive. Both forms of phosphorus are insoluble in water and can be interconverted with various applications of heat, pressure, and light. Because white phosphorus is the most easily made allotrope, it is used as the standard state for calculations of energy changes. This gives red phosphorus the unusual exothermic Δ_fH value of -17.6 kJ mol^-1.

There also exist black phosphorus and violet phosphorus. Black phosphorus has a Δ_fH value of -39.3 kJ mol^-1 and violet phosphorus has a Δ_fH value of -37.4 to -43.2 kJ mol^-1. Unlike nitrogen, phosphorus will not readily form a diatomic molecule with a triple bond. Diphosphorus does exist, but only between a temperature range of 1200 °C and 2000 °C.

Phosphorus pentoxide strongly reacts with water to form phosphoric acid, a substance that removes rust from iron, especially on ships; it is often known as "naval jelly". But phosphoric acid is corrosive to flesh and not to be touched.

Phosphorus is essential to life in the form of phosphates in bones and in substances known as ADP and ATP that transform food into useful energy in cells.

Others

Arsenic is similar to phosphorus. It has three allotropes: grey arsenic, yellow arsenic, and black arsenic. Grey arsenic is the most common form. Its structure is similar to graphite.

Antimony does not have physical properties of a metal, but behaves chemically as a non-metal.

Bismuth is a brittle, silvery metal. Bismuth is actually radioactive, decaying into thallium-205. Because its half-life is 19 x 10¹⁸ years, about a million times the age of the universe, bismuth is usually considered stable.

Bismuth is much less radioactive than the nearly-harmless and unavoidable radioactive isotopes of carbon and potassium in living things. Unlike arsenic and antimony, its compounds aren't toxic unless something else in the compound is itself toxic or the substance is very acidic or alkaline. In fact, a bismuth compound is very common in a heavily-used stomach medication that requires no prescription.