Since the discovery of gold an estimated 5,000 years ago, metals have played an integral part in the lives of human beings. We all know, almost instinctively, what consitutes a metal and what doesn't, and those who dimly remember their school science lessons might vaguely remember being taught the following.

Identifying Metals

The important things to remember are that this list does not constitute the definition of a metal and that none of these definitions are exclusive to metals. The list is merely a checklist of properties to which nearly all metals conform. The actual definition is complicated, involves electron configuration and crystallography, and is beyond the scope of this Entry.

Shiny Appearance

Perhaps the most obvious property of many metals is that they reflect a great deal of light, and thus appear shiny. Although not a particularly scientific method of identification, it holds true in all cases — although some metals are more reflective than others. Additionally, many other non-metallic materials are shiny, including silicon and iron pyrites¹.

The reason for metals' high reflective index is linked to their structure. Atoms are closely packed and their electrons are relatively free to move within this closely-packed network. The electrons effectively form a 'shield' that absorbs and reflects light straight back.

High Thermal Conductivity

Again, this is a property frequently and instinctively associated with metals. Metals feel cold to the touch. The reason? They are good heat conductors and quickly move heat away from your body.

High Electrical Conductivity

Allied with the conduction of heat is the conduction of electricity. At a simple level, both operate through the same mechanism — energy is transferred quickly through the electrons, which are free to move. It is important to note that not all metals conduct equally well. Copper is a very efficient conductor of both heat and electricity, hence it is used in heat-exchange tubes and electrical wiring. A metal such as titanium or zinc, on the other hand, would be of little use, as it conducts far less efficiently.

High Melting Point

This is a rather tenuous point. It is true that most common metals² are solid at room temperature, and melt at temperatures above 500°C — a considerably greater temperature than that at the centre of a gas flame. The classic 'exception' to the rule is mercury, whose unusually low melting point (-39°C) makes it a liquid at room temperature. Several other metals have melting points below 100°C, including sodium, potassium, caesium³ and gallium.

The high melting point of most metals is due to the strong attractive forces between the tightly-packed atoms. You need to put in a considerable amount of heat energy in order to break these forces and turn the solid metal into a liquid.

Sonorosity

This one tends to be frowned down upon by serious scientists, but it still holds an element of truth. Metals, when manufactured in a shape that allows them to vibrate, make a pleasant ringing sound when tapped. This is due to the very regular atomic structure which allows vibrations to pass through without interference.

Malleability, Ductility and Elasticity

If metals weren't malleable and ductile, they would be useless. Malleability measures the ease of a material to be bent, stretched, tooled or formed. Ductility is the extent to which a material can be drawn out or deformed without breaking. As a result of these properties, humans can shape metals into pretty much any shape they want or require.

Elasticity is the ability of a material to regain its original shape once it has been deformed. The more it can be deformed and still 'spring back', the more elastic it is. As metals bend rather than break, they have a degree of elasticity which is not found in many other traditional materials, such as concrete, glass or stone, of similar strength. Today, carbon fibre and similar reinforced polymers are used in hi-tech construction because they have an even greater elasticity.

All of these properties derive from metals' very regular structure, which allows layers of atoms to slip over each other as the metal is deformed. A shoebox of marbles is a good analogy — if you tip the shoebox, the marbles all slide over each other, but end up in the same regular pattern.

Strength

Ah, yes. Now you're talking. Metals are indeed strong. In their solid state, they can support loads some 250,000 times their own weight. The ready availability of iron, coupled with its comparatively easy method of extraction and purification (as when making steel) and very high strength, is the direct reason for its dominance over other metals in the construction industry. Other metals combine light weight with very high strength — this is why titanium and aluminium are widely used in the aerospace industry.

Toughness

Toughness? Is that not the same as strength? Well, yes and no. The scientific meaning of toughness is the substance's ability to withstand impact — it is measured in terms of the amount of energy it absorbs before fracturing. Frequently, this measurement is closely allied to the strength of a material, but certain materials (notably ceramics) shatter very easily — hence a low toughness — yet have high strength.

Hardness

This is another tenuous one. Hardness is defined as a material's resistance to being scratched. Most common metals do rate very highly on Moh's scale, which arbitrarily compares hardnesses, but the alkali metals on the far left of the Periodic Table are very soft, and can be cut clean through with a sharp blade.

All the above rules hold true for both elemental metals (those made of a single type of atom) and alloys — mixtures of two or more different metals, usually designed for increased strength, machineability or casting properties. Some of our best-known metals are alloys: brass is copper (60-70%) and zinc (30-40%); bronze is copper (75-95%) and tin (5-25%); steel is very pure iron with a number of other elements added in small quantities. The second half of this Entry, however, will focus mainly on those metals which are pure elements.

Reactivity

Over the years, enterprising scientists have thought of many ways to categorise and arrange the known metallic elements - the Periodic Table of the Elements being the best known - but perhaps the most accessible and useful is the Reactivity Series. The following series includes the best-known metals:

A number of mnemonics exist to help people remember the Reactivity Series; one big problem, however, is that different versions of the series exist, containing different numbers of metals. 'Please Stop Calling My Zebra In Class' is commonly used to remember a seven-metal Reactivity Series of potassium, sodium, calcium, magnesium, zinc, iron and copper.

The reactivity of an element derives from the ease with which a metal atom will lose electrons, hence forming an ion and therefore undergoing a chemical reaction. The 'grip' an atom has on its outermost electrons is dependent on the atomic radius — the distance from the electron to the nucleus — so the more reactive metals are those with higher atomic radii.

So What Use Is It?

Knowing the Reactivity Series is one thing — being able to use it is another. Nearly every use of metals will involve reference to the series at some point; we focus here on some of the most common applications.

Corrosion

Metals used outdoors are corroded by rainwater or acid rain. Metals in seawater corrode much faster because of the presence of dissolved salts. In the case of each, it is useful to know how fast a bare metal will corrode.

The rate of corrosion is directly linked to a metal's reactivity. Potassium and sodium are sufficiently reactive to spontaneously burst into flame on contact with cold water. Magnesium and calcium will not remain intact for a day in a wet environment. Iron, as we all know, will rust (a chemical reaction with water and oxygen in the air) in a month. Even unprotected copper will form verdigris (green copper oxide) in a few decades.

Surprisingly, given its high reactivity, aluminium is not particularly easy to corrode. The reason for this is that exposed aluminium very quickly reacts with oxygen in the air (as would be expected), and the oxide layer that forms as a result effectively protects the metal underneath from any further reaction. It is easy to see this at home if you have an aluminium greenhouse — carefully put a scratch in the aluminium and you will see shiny metal. Within a few seconds this will fade to a whiter shade as aluminium and oxygen react.

Voltage

Electricity is the movement of electrons, and all that is required to make electrons move is two different metals, separated by a solution containing ions. Many people are familiar with the 'potato clock'⁴ which uses zinc and copper strips and the salt solution in a potato to create enough voltage to power a small digital clock.

The voltage of a circuit is the amount of electrical energy available to it, and depends on a number of factors, such as the concentration of the acid and the number of metal strips. Among the factors is the choice of the two metals. It so happens that a metal's position in the Reactivity Series is linked to its electronegativity: the electrical potential energy. So, gold is very electropositive while potassium is very electronegative.

The potential voltage of a chemical cell such as the potato clock is dependant upon the difference between the two metals' electronegativities. Therefore, gold and magnesium strips would make a high-voltage cell; iron and zinc less so. Metals higher than magnesium would quickly react with the weak acids in fruit or vegetables, so would not be suitable.

Displacement

Most metals readily form compounds with oxygen (oxides), chlorine (chlorides), sulphur (sulphides and sulphates) and carbon (carbides and carbonates). A metal can be displaced from its compounds by a more reactive metal — that is, one higher in the Reactivity Series — the products are a compound of the more reactive metal and the pure, unreacted metal that was originaly part of the compound. This is of use in certain specialised reactions such as:

The Thermite Reaction. The reaction between aluminium and iron (III) oxide is a classic example of displacement. The aluminium, being more reactive, displaces or 'steals' the oxide, and aluminium oxide and iron remain. The most significant aspect of this reaction is that it is highly exothermic — that is, it gives out a lot of heat. The heat released is sufficient to melt the iron that is produced. When this reaction is performed in a localised area, triggered by a fuse (usually of magnesium)⁵, the molten iron will effectively create a weld, and so the thermite reaction is used as a low-cost method of welding railway lines in situ.

Plating. It is possible to silver-plate or copper-plate an item made of a more reactive metal very quickly, simply by placing it in a solution of silver nitrate or copper sulphate, respectively. The more reactive metal displaces the silver (or copper) from the compound in solution, allowing the pure silver (or copper) to form on the surface of the item.

Extraction

A metal is useless unless you can extract it from the ground and purify it. Most metals are tied up in ores — chemical compounds of the metal, usually sulphides or oxides. The more reactive a metal, the more strongly bonded it is to its compounds. Hence, those that are more reactive require a great deal of energy. Less reactive metals, such as silver, can be extracted from certain ores simply by heating them so that the ore decomposes. Gold is so unreactive that it doesn't even form ores, and is found in solid lumps underground or in rivers.

The other metals need more specialised methods to extract them from their ores. It would be possible to use a displacement reaction, as described above, but this would mean sacrificing one metal at the cost of another, and is usually unprofitable. What we can do is add certain low-cost non-metals into the Reactivity Series in order to plan displacement reactions. Carbon has a place between aluminium and zinc, and hydrogen between lead and copper.

It is interesting to note in passing that hydrogen is a constituent part of all acids. Indeed, the definition of an acid is based on how many hydrogen ions it contains. When a metal reacts with an acid, what effectively happens is a displacement reaction, releasing hydrogen gas rather than a metal. It is easy to tell which metals will react with acids and which won't: those that are less reactive than hydrogen (copper, silver, gold) will not displace the gas and so no reaction will take place.

Back to extraction: it will be noted that copper is below hydrogen in the Reactivity Series, and so can be extracted by displacement. This is done by passing a hydrogen-rich source (usually methane) over the ore⁶ and heating gently. This must be done under carefully controlled conditions, as methane and hydrogen are very explosive!

Similarly, iron is extracted from its ore⁷ in the blast furnace by being displaced by a carbon source. In this case the source is carbon monoxide, which is produced in the blast furnace by the addition of coke — coal which has been partially oxidised to remove sulphur⁸.

Metals higher up the reactivity series require a more intensive input of energy, and are generally separated by electrolysis — the passing of a strong current through the molten ore to separate the ions. This is a very high-energy and thus expensive process. It costs money not only to melt the ore (often in excess of 2000°C), but to keep it molten and generate large amounts of electricity. This is not a problem for sodium or potassium, where the demand is fairly low, but extremely uneconomical for aluminium, which has a massive global demand⁹. For this reason, aluminium recycling is universally encouraged — it saves some 70% of the energy required to extract aluminium from bauxite, its ore.

Discovery

Since very reactive metals are harder to extract, it is unsurprising that they remained unisolated, and hence undiscovered, for a greater length of time. This Reactivity Series shows quite a good correlation between reactivity and date of discovery: