We all recognise metals when we see them, and know intuitively what sort of properties they have. But what is it that makes them what they are, and why do they account for four-fifths of known elements?
METALS have been important to us for millennia, since our ancestors first discovered 鈥渘ative metals鈥, learned to extract others from their ores, and found ways to fashion them into crude tools, coins, ornaments and weapons. Today, metals still underpin the fabric of an industrial society, piping electricity into homes, factories and offices and providing the raw materials for many items, ranging from robots and automobiles to toys and jewellery.
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Some metals such as gold and silver have always held an allure because of their colour, their ability to resist corrosion and their rarity. Indeed, gold rushes continue to the present day in South American countries such as Brazil. The German chemist, Fritz Haber, even developed a way to extract gold from sea water to help his country pay war reparations after the First World War, but he overestimated the amount of gold present in the ocean. Today, metals such as platinum and rhodium are invaluable as catalysts. Well-known metals used in the nuclear industry include uranium, which exists in ores and plutonium, produced in nuclear reactors.
Mercury is the only metal that exists as a liquid at room temperature, and this explains why it earned the sobriquet 鈥渜uicksilver鈥. But what is it that sets metals apart from nonmetals? And how come four-fifths of the elements are metals?
We recognise metals by their appearance and physical attributes. They are cold to the touch, opaque, and have a shiny lustre. They are strong yet almost without exception ductile, capable of being stretched into wires, hammered into sheets and bent into awkward shapes without breaking. And, compared with nonmetals, they are exceptionally good conductors of heat and electricity. Metals can also mix well together to form alloys, such as brass, without losing their metallic character.
But the boundary between metals and nonmetals is fuzzy, and some elements can behave as either. Often, temperature, pressure and impurities can make the difference between whether an element is a metal or not. Germanium, selenium and tellurium, for example, are nonmetals in their crystalline state under normal conditions, but melt them and they become metallic. Even hydrogen 鈥 a gas under ordinary conditions 鈥 becomes a metal under intense pressure. Almost any substance can become a metal if its atoms are crammed closely enough together. This may be the case in the centre of Jupiter, whose interior is under enough pressure to turn all the elements present into metals.
Some metals are compounds rather than elements. When ammonia dissolves alkali metals such as potassium or sodium, for example, it becomes a metal. Magnetite 鈥 an oxide of iron 鈥 is a metal as well. Some organic polymers such as polyacetylene acquire metallic properties, such as the ability to conduct electricity, when they are 鈥渄oped鈥 with iodine.
Atomic assembly
Filling shells
BUT, as will become clear, all metals owe their distinctive physical appearance and properties to the way that elements are assembled at the atomic level (see Box 1). This also dictates their chemical properties 鈥 the way that they react with nonmetals and other chemical compounds.
More specifically, metallic behaviour arises from the way that electrons fill up orbitals, or shells, around the nucleus of an atom. The first shell accommodates a maximum of 2 electrons, the second contains no more than 8, the third no more than 18 and the fourth a maximum of 32. The electronic configuration of an element shows how many electrons reside in each shell. For example, the configuration of the inert gas neon is (2.8), because it has 2 electrons in its first shell and 8 in its outer shell. Models of the atom devised first by Ernest Rutherford, the New Zealand physicist, and than modified by Niels Bohr, the Danish physicist, help to explain why the shells fill up in this way (see Box 2).
Nearly all the 81 known metals have either one or two electrons in the outermost shell. For instance, lithium (2.1) has one in the outermost shell and beryllium (2.2) has two. The electrons in the outermost shell are usually the most loosely bound ones, and are known as valency electrons. When the number of these is three or more, as in aluminium and germanium respectively, the element becomes less metallic in nature. These valency electrons hold the key to metallic properties. Because they are so loosely bound, they are 鈥渇ree鈥 to perform chemical and physical tasks and functions. They are the active agents in a metal鈥檚 chemical reactions.
Valency electrons help to explain how metals bond to other elements. Two of the most important bonding processes, known as ionic and covalent bonding, occur because individual elements which lack the stability of inert gases such as helium, neon and argon can acquire it by forming octets of electrons with other elements. They can do this by losing, gaining or sharing valency electrons, which explains why metals bond in so many different combinations with other elements.
If metals can find a way to give away their valency electrons to another element, they are left with the electronic configuration of an inert gas. They would also end up having more protons than electrons. Such electrically charged species are called ions. For example, if sodium (2.8.1) loses its valency electron, it has the same electronic configuration as neon (2.8), and becomes the positively charged sodium ion (Na+). Magnesium must lose two valency electrons to become like neon. It therefore becomes an ion with a double positive charge (Mg2+). Metals are known as electropositive elements because they form electrically positive ions.
Nonmetals, by contrast, achieve their stable octets by 鈥渟tealing鈥 electrons to top up outer shells that already almost form an octet. Nonmetals form electronegative ions because in stealing electrons, they acquire a net negative charge. Fluorine (2.7), for example, adopts the electronic configuration of neon (2.8) by gaining a valency electron, and becomes the ion F鈥 with a single negative charge.
When a pair of elements release and accept electrons in this way, they become positive and negative ions respectively. These then attract each other electrically, forming a stable ionic compound. A good example is common salt, or sodium chloride, which crystallises into a regular cubic structure. Salt is stable because sodium donates its single valency electron to chlorine which needs an electron to complete an octet, so that both ions acquire the electronic configurations of an inert gas.
The metallic bond
A sunken grid?
IONIC bonding is a very straightforward way for two elements to acquire mutual stability, provided one has an outer shell that is almost full and the other has one or two electrons in its outer shell. But what happens if the element鈥檚 outermost shell contains four valency electrons, or three or five? Such elements, which can be either metallic or nonmetallic, achieve their octets by sharing rather than trading their outermost electrons. The links they form are called covalent bonds. For example, carbon, which has a valency of four, crystallises as diamond by forming covalent bonds with four neighbouring carbon atoms, each of which contributes one electron to make up a stable, eight-electron ring around each individual carbon atom. Silicon, grey tin and germanium atoms bind to one another by a similar process.
As all the valency electrons are involved in individual bonds, none is available for conducting electricity, so covalent substances are insulators and non-electrolytes. However, substances such as silicon can be made to conduct electricity by the addition of traces of 鈥渋mpurities鈥. Such substances are called semiconductors. Silicon has four valency electrons, whereas phosphorus, for example, has five. This means that if traces of phosphorus are added to a block of silicon, four of the impurity鈥檚 five valency electrons will form covalent bonds with a corresponding number of silicon electrons in the lattice. But there will be one spare phosphorus electron, which is loosely attached and easily ionised. This means it is free to wander through the silicon lattice, creating an electric current as it moves. By contrast, an atom with fewer valency electrons will generate an 鈥渆lectron hole鈥 which is also capable of breaking free and moving through the lattice. Such impurity semiconductors, with either electrons or holes to generate current, form the basis of the transistor.
Broadly speaking, however, nonmetals are insulators. Compounds of nonmetals do form important 鈥渃omplexes鈥 with metals, such as haemoglobin, the molecule in blood that transports oxygen from the lungs to the tissues, but these are beyond the scope of this Inside Science. Almost all the compounds formed by nonmetals link together through covalent bonding. One of the simplest examples is methane (CH4), or natural gas. Since carbon has four valency electrons, it needs another four to make up its stable octet. In methane, it does this by sharing each valency electron with a single electron from each hydrogen atom. The four hydrogen atoms each acquire full outer shells (of two electrons) into the bargain, achieving the stability of the inert gas helium. Now that we know the underlying atomic structure of metals, how they bond and how they react, we can see how these factors combine to give metals their well-known properties.
The typical metal atom has far too few electrons in its outer shell to form covalent bonds with neighbouring metal atoms, and ionic bonding is only possible between metals and nonmetals. Yet solid metals are very strong and the atoms evidently bond together very well, so how do they do it? Again, the valency electrons predominate in the bonding process. Individual metal atoms give up their valency electrons to form a sort of gas, or plasma, of freely moving electrons surrounding a 鈥済rid鈥 of positively charged metal ions. Crudely, we can think of it as a lattice of ions in a 鈥渟ea鈥 of electrons. Because the electrons are relatively free to manoeuvre through the lattice structure, metals make good electrical conductors. The lattice model of metals also explains their other physical hallmarks.
The elasticity found in most metals, for example, results from their need to retain a delicate balance of repulsive and attractive forces within the structure. Neighbouring metal ions repel one another, yet the ions are all attracted to the surrounding sea of electrons. Physical stress, such as the blow of a hammer, can disturb the balance between the ions and the electrons, forcing the ions to squeeze together. But as soon as the stress recedes, the ions reform the original lattice structure. This is what gives metals their elasticity.
As there are no directional bonds 鈥 only an isotropic attraction drawing charged ions equally in all directions to a negatively charged fluid 鈥 metals generally crystallise in simple, dense structures composed of planes of ions. When one plane slides over a lower one by one or two lattice spacings, it rapidly recovers its original symmetry and structural stability. This explains why metals are ductile, and can be drawn out into wires or hammered into sheets.
Moreover, because the lattice structures are so regular, it is easy to substitute atoms of one metal with atoms of another without disturbing the structure of the lattice or the overall properties of the material. This explains why many metals readily form alloys with one another.
When a beam of light strikes a metal, free electrons in the lattice absorb energy from the light to move to higher energy 鈥渆xcited鈥 states. These excited electrons absorb the light at all wavelengths, which makes the metal opaque. However, as the excited electrons return to their lower energy states, they emit nearly all the radiation they originally received. This reflected light produces the familiar lustre of metals.
Metal crystals
Snooker frame
THE best-known crystals, such as diamond, sapphire and quartz, are usually transparent or translucent, and their outer shape reflects their regular atomic structure. Metals are crystalline too, but this is less obvious because they are opaque, and can come in any shape. Moreover, because metals are usually aggregates of small grains called crystallites, their external shape does not reflect their underlying atomic structure.
In a metallic lattice, the ions are attracted equally in all directions by the free electrons. Therefore, metals minimise the energy held in the lattice 鈥 which confers stability 鈥 by packing the ions in regular arrays as closely as possible. The structures of metals can be likened to assemblies of snooker balls. When all the red balls are packed into the triangle before a game of snooker, they arrange themselves into a close-packed array. If the array extended beyond the triangle, covering the entire table or beyond, each ball would be seen to be surrounded by six others in a hexagonal pattern. Imagine adding a second layer (B) of balls on top of the first (A). These would fit in the hollows between the balls in the first layer. Thus, the pattern of the first, closely-packed plane is reproduced exactly, but displaced to one side in relation to the lower layer.
If we add a third layer, a choice arises because the second layer has two sets of hollows. If one set is chosen the balls in the third layer will end up directly above the balls in the first layer. This is referred to as ABABAB stacking. If the other set of hollows is chosen, an identical pattern is obtained, but it is displaced to one side. This stacking is called ABCABCABC. In this arrangement, it is the fourth layer which ends up directly above the first.
In both structures, the ions 鈥 represented by the snooker balls 鈥 achieve maximum packing density. A random assembly of balls would leave 36 per cent of free space. Maximum packing density reduces this to just 26 per cent. Each ion is in direct contact with 12 neighbours, and is said to have a coordination number of 12. ABABAB-stacked structures are known as hexagonal close-packed (hcp) arrays, whereas those with ABCABC stacking are called face-centred cubic (fcc) arrays (see Diagram). Hcp metals include beryllium and cadmium.
A third simple crystal type common among metals is the body-centred cubic (bcc) structure, in which ions at the corners of a basic cube imprison another ion in the centre. The ions are less closely packed than in hcp or fcc, and leave as much as 32 per cent free space. A bcc structure has a coordination number of only eight. Bcc metals include the alkali metals lithium, potassium and sodium. Some metals have different structures at different temperatures. Iron, for example, is bcc below 910掳C and above 1400掳C, and fcc between.
We have already mentioned how metal crystals deform and confer ductility when close-packed atomic planes slide over one another. But if the sliding plane moves en bloc 鈥 rigidly, that is, with all its atoms moving simultaneously 鈥 the force required to make it happen would be about a thousand times as great as that observed. Line defects in the atomic plane provide the explanation. The movement of line defects allows atoms to be displaced sequentially, one by one. The process resembles the trick of creating a small fold, or 鈥渞uck鈥, in a carpet. The movement of the ruck displaces the entire carpet by a small distance. These line defects, known as edge dislocations, are visible under electron microscopes.
Another type of line defect called a screw dislocation turns an otherwise parallel series of crystallographic planes into a continuous spiral. Screw dislocations are important when metal crystals grow from a melt, or from the vapour phase.
Another type of fault, the point defect, arises where an atom is missing from its lattice site, or where an atom is wedged into an interstitial position between lattice-bound atoms. These are known respectively as vacancies and interstitials.
Vacancies are especially important, particularly at high temperatures, where they play a dominant role in processes of diffusion and deformation. More vacancies occur as the temperature rises. In copper, for example, vacancies are 100 000 lattice spacings apart at room temperature, but at 1007掳C they are just 10 spacings apart. Vacancies are like a row of patients in a doctor鈥檚 surgery. As one person sees the doctor, their seat is taken by the next, so that the 鈥渧acant鈥 seat diffuses along.
Interstitials are fewer in number and generally less important than vacancies. But where the interstitial is an atom of a different element, either as an impurity or as a deliberate alloying addition, they can affect properties profoundly. The properties of steel, for example, arise from interstitial atoms of carbon in a lattice of iron.
We now know the underlying atomic architecture that makes metals what they are. But there鈥檚 much more to metals than that. They form a multitude of alloys, for example, and suffer from weaknesses such as corrosion. To do these subjects justice, they may each need an edition of Inside Science to themselves in the future.
1: Meet the atomic family
ATOMS contain three fundamental particles 鈥 protons, neutrons and electrons. An atom of the simplest element, hydrogen, has a single proton in its nucleus, which is orbited by a single electron. Like all atoms of elements, the hydrogen atom is electrically neutral. This is because the positive charge on the proton and the negative charge on the orbiting electron are equal but opposite, cancelling one another out.
As the number of protons in the nucleus increases from element to element (a quantity known as the atomic number) so too does the number of circulating electrons, to keep the whole atom electrically neutral. The atoms of all elements heavier than hydrogen contain neutrons in their nuclei as well as protons. Neutrons add mass to the nucleus, but they do not disturb the electrical equilibrium because they are electrically neutral. A proton鈥檚 mass is almost the same as that of a neutron, and both are roughly 1836 times as heavy as an electron.
Lithium, with the atomic number 3, is the lightest natural metal. It has 3 protons and 4 neutrons in the nucleus, orbited by three electrons.
2. The Rutherford-Bohr model of the atom
METALS usually possess one or two electrons in their outermost electron shells, which give them their unique character. But why should as many as four-fifths of the elements qualify as metals by having this electronic configuration? To understand this, we should look at the way that electron shells fill up. Ernest Rutherford鈥檚 model of the atom consisted of nuclei orbited by electrons. In a brilliant modification of this rather crude model, Niels Bohr assumed that the motions of the electrons obeyed 鈥渜uantum theory鈥 rules, orbiting the nucleus only within fixed orbitals, or shells.
Once a shell was full, additional electrons would have to occupy shells more distant from the nucleus. The theory requires that each shell would only be able to hold a limited number of electrons, dictated by the formula 2n2, where n is 1 for the innermost shell, 2 for the next and so on. Thus the innermost shell would accommodate just 2 electrons (2脳12=2), the second would accommodate 8 electrons (2脳22=8), the third shell 18, the fourth 32, and so on. In practice 鈥 and for reasons beyond the scope of this Inside Science 鈥 no shells contain more than 32.
The table of the first 26 elements shows that as the atomic number rises, the consecutive shells (designated K, L, M, N and so on) fill with electrons sequentially. In atoms of the element helium, for example, the innermost K shell gets filled to capacity with two electrons. Then, as expected for atoms that have no room in their innermost shell, the L shell begins to fill. Likewise, when the L shell has its full complement of eight electrons, the M shell begins to fill.
According to this simple model, the M shell should steadily fill up with 18 electrons. But something unexpected happens after argon, which has an electronic configuration of (2.8.8), corresponding to two electrons in the innermost K shell, eight in the L shell and eight in the M shell. We would expect the next element, potassium, to have the electronic configuration 2.8.9. Instead, the new electron finds it energetically favourable to start occupancy of the next shell in the sequence, leaving the filling of the M shell till later. So potassium has the unexpected configuration 2.8.8.1, and the lone electron in the outermost shell is clearly a valency electron.
Calcium, the next element, adds a further valency electron to the N shell, giving the configuration 2.8.8.2. After that, the elements start filling up the M shell again, so scandium has the configuration 2.8.9.2. As the next eight elements gradually fill up the M shell, they each retain outer shells of two valency electrons, which is why they are all metals.
The group of eight metals from titanium (2.8.10.2) through to copper (2.8.18.1) constitute a very important group called the transition metals. Some of them 鈥 iron, nickel and cobalt 鈥 are ferromagnetic and swivel in magnetic fields. This is because the tiny magnetic moments of the electrons filling up the M shell do not mutually cancel. Such elements can be magnetised and in some cases made into permanent magnets.
The lanthanides (rare-earth metals) and actinides form additional transition metal groups, from cerium to lutetium and from thorium to lawrencium respectively. So by 鈥渂ackfilling鈥 of inner shells, while keeping one or two electrons in the outermost shell, nature ensures that most elements take on the electronic configuration of metals.
Further Reading
To find out more about metals, try A.H. Cottrell鈥檚 books entitled Theoretical Structural Metallurgy (Edward Arnold, 1948) and Portrait of Nature (Charles Scribner鈥檚 Sons, 1975). Other volumes include Metals in the Service of Man, by W. Alexander and A. Street (Penguin, 1994) and Tommorrow鈥檚 Materials, by K. Easterling (The Institute of Metals, 1990). J.E. Gordon has written two books, both published by Penguin, entitled The New Science of Strong Materials (1968) and Structures (1978).