Metals and Non-metals

Class 7 sorted the world cleanly: metals shine, bend, stretch and ring; non-metals do not. Try applying that rule without exception, though, and it starts to crack almost immediately. Gallium and caesium are unmistakably metals, yet both have such low melting points that either one would melt sitting in your palm. Iodine is unmistakably a non-metal, yet it is lustrous, shiny in a way non-metals are supposed never to be. Carbon, a non-metal, exists in more than one form, called allotropes, and one of them, graphite, conducts electricity well enough to be used in electrodes, a property this whole chapter's rulebook assigns only to metals. Even among metals themselves, the alkali metals, lithium, sodium, potassium, are so soft they can be sliced with a knife, hardly the confident, solid image "metal" usually conjures. None of this means Class 7's properties were wrong, they remain useful first checks, but it does mean physical properties alone cannot be trusted to sort every element correctly. A more reliable dividing line is needed, and it will not come from touch or shine at all, but from chemistry: how an element actually behaves when it reacts.

The physical properties from Class 7 still matter, they are simply not the whole story. Metals in their pure state have metallic lustre. Most are hard, though the alkali metals are a clear exception. Many can be beaten into sheets (malleable) or drawn into wires (ductile); gold and silver top both lists, and one gram of gold can still be drawn into a wire roughly two kilometres long. Metals are generally good conductors of heat, silver and copper best of all, while lead and mercury conduct comparatively poorly. Almost all metals are solid at room temperature, mercury the one liquid exception, and most have high melting points, though gallium and caesium melt at temperatures barely above room warmth. Non-metals, by contrast, are mostly solids or gases, with bromine the one liquid exception, and there are simply far fewer non-metals than metals to begin with. Iodine's lustre and graphite's conductivity are genuine exceptions on the non-metal side too. Compiled honestly into one table, metal versus non-metal, these physical properties form a strong pattern with real, documented exceptions on both sides, useful as a first pass, but not sturdy enough to rely on alone.

Burn a strip of magnesium and it flares into a dazzling white flame; heat a strip of copper and nothing so dramatic happens, but its surface slowly darkens under a coat of black copper(II) oxide. Almost every metal, given the chance, combines with oxygen to form a metal oxide: 2Cu + O2 -> 2CuO, 4Al + 3O2 -> 2Al2O3. As a rule, metal oxides are basic, but a few, aluminium oxide and zinc oxide among them, are amphoteric, reacting with both acids and bases to form a salt and water, Al2O3 + 6HCl -> 2AlCl3 + 3H2O, and Al2O3 + 2NaOH -> 2NaAlO2 + H2O. Most metal oxides are insoluble in water, but sodium oxide and potassium oxide are exceptions, dissolving to form strongly alkaline solutions, Na2O + H2O -> 2NaOH. What this activity really reveals, though, is that metals do not all react with oxygen at the same rate. Potassium and sodium react so vigorously they must be stored fully submerged in kerosene to prevent them catching fire on contact with air. Magnesium, aluminium and zinc develop a thin, tightly-bound protective oxide layer that shields the metal underneath from further reaction, aluminium's is put to deliberate use in anodising, a process that thickens this same natural layer electrolytically to make the metal more corrosion-resistant. Iron does not burn as a solid block, but its filings burn vigorously when sprinkled into a flame. Copper darkens but does not burn. Silver and gold barely react with oxygen at all, even at high temperature, a stubbornness that turns out to matter enormously later in this chapter.

Drop a sliver of sodium into cold water and the reaction is instant and violent, so much heat is released that the hydrogen gas produced can catch fire on the spot: 2Na + 2H2O -> 2NaOH + H2 + heat. Calcium reacts too, but far more gently, Ca + 2H2O -> Ca(OH)2 + H2, and calcium hydroxide is only sparingly soluble, so the metal is often seen floating, held up by hydrogen bubbles clinging to its surface rather than by any change in the metal itself. Magnesium barely reacts with cold water at all, but switch to hot water and it reacts to form magnesium hydroxide and hydrogen, floating for the same bubble-clinging reason as calcium. Aluminium, iron and zinc go further still: they do not react with cold water or hot water, only with steam, forming a metal oxide and hydrogen, 2Al + 3H2O(g) -> Al2O3 + 3H2, 3Fe + 4H2O(g) -> Fe3O4 + 4H2. Lead, copper, silver and gold do not react with water in any form at all, not cold, not hot, not steam. Lay these five outcomes end to end, violent with cold water, gentle with cold water, only with hot water, only with steam, no reaction at all, and a reactivity gradient starts to emerge that oxygen alone could not fully reveal.

Metals react with dilute acids to give a salt and hydrogen gas, Metal + Dilute acid -> Salt + Hydrogen, but not all metals do this at the same speed. Drop magnesium into dilute hydrochloric acid and bubbles form almost immediately, the reaction visibly exothermic, a thermometer dipped into the mixture climbing fast. Aluminium and zinc react too, though less energetically, and iron slower still. Copper, dropped into the very same dilute acid, produces no bubbles and no temperature change at all, it simply does not react with dilute HCl. Lined up by how vigorously they react, magnesium, aluminium, zinc and iron fall into a decreasing order: Mg > Al > Zn > Fe. Nitric acid complicates this picture in one specific way: it is such a strong oxidising agent that it oxidises the hydrogen a metal reaction would normally produce straight into water, so no hydrogen gas is seen at all with most metals, magnesium and manganese being the rare exceptions that still release some. One especially striking exception belongs to gold: neither hydrochloric acid nor nitric acid can dissolve it alone, yet mix them in a 3:1 ratio and the resulting aqua regia, Latin for royal water, can dissolve gold and even platinum, a reminder that gold's famous unreactiveness has limits after all, just very narrow ones.

Dip an iron nail into a blue copper sulphate solution and, after twenty minutes, a reddish-brown coating of copper appears on the nail while the blue colour of the solution visibly fades. Try the reverse, a copper wire dipped into iron sulphate solution, and nothing at all happens. Iron has displaced copper from its compound; copper cannot do the same to iron. Fe + CuSO4 -> FeSO4 + Cu. This kind of reaction, a more reactive metal pushing a less reactive metal out of its salt solution, Metal A + Salt solution of B -> Salt solution of A + Metal B, is called a displacement reaction, and it offers something the earlier oxygen and water tests could not: a clean, direct, side-by-side comparison between any two specific metals, without needing either one to react with oxygen or water at all. Copper and lead, for instance, barely react with water or dilute acid, which makes them hard to rank using those tests alone, but a displacement reaction between them settles the question immediately and unambiguously.

Every reaction explored so far, with oxygen, with water, with dilute acids, with each other's salt solutions, was quietly answering the same underlying question about each metal: how badly does it want to react? Assembled together, that evidence produces the reactivity series, or activity series, a single list of common metals ranked from most to least reactive: potassium, sodium, calcium, magnesium, aluminium, zinc, iron, lead, hydrogen (included as a reference point, though it is not a metal), copper, mercury, silver, gold. Everything from earlier sections now reads as a direct consequence of this one ordering. Potassium and sodium sit at the very top, and it shows: violent with water, vigorous with oxygen, stored in kerosene out of sheer necessity. Gold and silver sit at the very bottom, and that shows too: barely touched by oxygen, untouched by water, dissolved only by the narrow special case of aqua regia. Any metal above hydrogen in this list can displace hydrogen from a dilute acid; any metal below it cannot, which is exactly why copper produced no bubbles in dilute HCl. And a more reactive metal will always displace a less reactive one from its salt solution, precisely the iron-and-copper result observed moments earlier. One table, built from four completely different kinds of evidence, and it predicts all four kinds of behaviour at once.

All of this reactivity has one underlying cause, and it goes back to electronic configuration. Noble gases, with a completely filled outermost (valence) shell, show almost no chemical activity at all, and that stability is the target every other element is quietly trying to reach. Look at sodium: 2, 8, 1, eleven electrons across three shells, one lone electron sitting in the outermost shell. Lose that single electron and the shell beneath, already a stable, complete 2, 8, becomes the new outermost shell. The nucleus still holds 11 positively charged protons, but now only 10 electrons remain, leaving a net single positive charge, a sodium cation, Na+. Chlorine, meanwhile, sits at 2, 8, 7, just one electron short of its own stable octet, and readily accepts the exact electron sodium is trying to lose, becoming Cl-, a chloride anion. Na -> Na+ + e-, and Cl + e- -> Cl-. These oppositely charged ions do not simply sit near each other; they are held together by strong electrostatic attraction into sodium chloride, and it is worth being precise here: NaCl does not exist as individual molecules at all, but as a vast, repeating aggregate of Na+ and Cl- ions locked in a fixed pattern. Magnesium chloride forms the very same way, just with magnesium losing two electrons instead of one, Mg -> Mg2+ + 2e-, needing two chlorine atoms to each accept one. Compounds built this way, through the outright transfer of electrons from a metal to a non-metal, are called ionic compounds, or electrovalent compounds, and this electron-losing tendency is precisely what the entire reactivity series has been measuring all along.

Take common salt, or any similar ionic compound, and heat a small amount directly in a flame: it does not melt easily, and a considerable amount of heat is required before it does. This is the first clue to the properties ionic compounds share. Because positive and negative ions are held together by strong electrostatic attraction throughout the entire structure, not just at one bond, ionic compounds are solids that are somewhat hard, since real force is needed to pull the ions apart, yet also brittle, since a hard enough blow shifts one layer of ions just enough to bring like charges into sudden contact, and the whole structure cracks apart. That same strong inter-ionic attraction is why ionic compounds have high melting and boiling points, sodium chloride does not melt until 1074 K, calcium oxide not until a striking 2850 K, since breaking apart a lattice of strongly attracted ions takes real, considerable energy. Ionic compounds are generally soluble in water but insoluble in kerosene or petrol. Most tellingly of all: a solid ionic compound does not conduct electricity, its ions are locked too rigidly in place to move, but dissolve that same compound in water, or melt it, and it conducts electricity readily, because now its ions are finally free to move toward oppositely charged electrodes. Every one of these properties, hardness with brittleness, high melting points, solubility, conduction only in solution or when molten, traces directly back to the same electron transfer just described.

The earth's crust, along with the dissolved salts in seawater, is where nearly every metal starts out. An element or compound occurring naturally in the crust is a mineral; where a mineral holds a high enough percentage of a particular metal to make extraction worthwhile, it is called an ore. And the reactivity series, already built, turns out to predict exactly how a metal will be found. Metals at the very bottom of the series, gold, silver, platinum, copper, are unreactive enough to occur as free elements, uncombined with anything else. Metals at the very top, potassium, sodium, calcium, magnesium, aluminium, are so reactive that they are never found free in nature at all, only locked inside compounds. Metals in the middle, zinc, iron, lead, are moderately reactive and typically occur as oxides, sulphides or carbonates, oxides especially, since oxygen itself is abundant and reactive. Before any of this ore becomes metal, it must be cleaned of gangue, the sand, soil and rock mixed in with it, using separation methods matched to how gangue and ore actually differ. Extracting metals low in the activity series is comparatively easy: their oxides give up their metal simply by heating, no chemical trickery required. Cinnabar, mercury's ore, converts to mercuric oxide on heating in air, then straight to liquid mercury on further heating, 2HgS + 3O2 -> 2HgO + 2SO2, then 2HgO -> 2Hg + O2. Copper's sulphide ore behaves almost as obligingly. Metals in the middle of the series need one extra step first: sulphide ores are roasted, heated strongly in excess air, and carbonate ores are calcined, heated strongly in limited air, both converting the ore into an oxide, which is then reduced to the metal using carbon, ZnO + C -> Zn + CO. Sometimes, though, carbon is not the reducing agent of choice, and a spectacularly different method takes over: a highly reactive metal like aluminium, mixed with a less reactive metal's oxide and ignited, reacts so exothermically that the freed metal comes out molten, Fe2O3 + 2Al -> 2Fe + Al2O3 + heat. This thermit reaction, hot enough to weld railway tracks and cracked machine parts together on the spot, is precisely the reactivity series turned into raw industrial force.

Metals high in the reactivity series will not give up their oxygen to carbon at all, since they hold onto oxygen more strongly than carbon itself does, which is exactly why heating alone, the trick that worked for mercury and zinc, does nothing for sodium, magnesium, calcium or aluminium. These metals are obtained instead by electrolytic reduction, passing electricity through the molten compound. Sodium, magnesium and calcium are extracted by electrolysing their molten chlorides: at the cathode, Na+ + e- -> Na, metal quietly depositing; at the anode, 2Cl- -> Cl2 + 2e-, chlorine gas bubbling away. Aluminium is obtained the same way, by electrolysing molten aluminium oxide. Even after extraction, though, most metals are not yet pure enough to use; they carry impurities that must be removed, and the standard method for this final step, electrolytic refining, is itself another use of the very same electrolysis. The impure metal becomes the anode, a thin strip of pure metal becomes the cathode, and a solution of the metal's own salt serves as the electrolyte. Pass current through, and the anode's metal dissolves into solution while an equivalent amount of pure metal deposits onto the cathode; soluble impurities drift off into the solution, while insoluble impurities simply sink to the bottom of the anode as a residue called anode mud. Copper, zinc, tin, nickel, silver and gold are all commonly purified this exact way. From ore in the ground to a pure metal ingot, the entire journey, low reactivity needing only heat, middle reactivity needing roasting or calcination plus a reducing agent, high reactivity needing electrolysis, and finally refining, is one continuous story, and the reactivity series was the map for all of it from the very first step.

Silver darkens to black over time as it reacts with sulphur in the air, forming silver sulphide. Copper loses its shine and gains a green coat as it slowly reacts with moist carbon dioxide, forming basic copper carbonate. Iron acquires a brown, flaky coating, rust, when left in moist air for long enough, confirmed by sealing iron nails in three test tubes, one exposed to both air and water, one to water alone with air kept out by a floating oil layer, one to dry air alone with moisture absorbed by anhydrous calcium chloride: only the nail exposed to both together actually rusts. All three are examples of corrosion, the gradual surface damage any metal can suffer from air, water or other substances, and it is a genuinely expensive problem, an enormous sum is spent every year in India alone repairing or replacing structures damaged by rusting. Prevention takes several forms: painting, oiling and greasing all work by physically keeping air and moisture off the metal surface; galvanising coats iron or steel with a thin layer of zinc, and remains protective even if that coating is scratched, since zinc, being more reactive, corrodes preferentially and sacrifices itself before the iron underneath is touched; chrome plating and anodising add their own protective layers; and alloying changes a metal's properties outright. One iron structure has needed none of this. The Iron Pillar near the Qutub Minar in Delhi, over 1,600 years old, 8 metres tall, weighing 6 tonnes, the very same pillar mentioned back in Class 7, has stood exposed to sun, rain and humidity for over a millennium and a half and still shows barely any rust, its surface composition and iron-working process still studied by scientists from around the world for the corrosion resistance it achieves without any of the modern methods listed above.

Pure iron is soft and stretches out of shape too easily when hot to be much use on its own, yet mixed with a mere 0.05 percent carbon, it becomes hard and strong, and mixed further with nickel and chromium, it becomes stainless steel, hard and resistant to rusting altogether. This deliberate mixing, a homogeneous blend of two or more metals, or a metal with a non-metal, made by melting the primary metal and dissolving the others into it in fixed proportions before cooling, is called an alloy, and it exists because pure metals are frequently not quite right for the job on their own. Brass, copper mixed with zinc, and bronze, copper mixed with tin, are both markedly poorer conductors of electricity than pure copper, which is exactly why electrical wiring stays pure copper while brass and bronze are reserved for fittings and statues instead. Solder, an alloy of lead and tin, melts at an unusually low temperature, precisely the property that makes it useful for welding electrical wires together without damaging them. If one of the metals involved happens to be mercury, the resulting alloy has its own special name: an amalgam. Even gold jewellery is rarely pure gold: 24-carat gold is far too soft to hold its shape as a ring or bangle, so Indian jewellery is conventionally made from 22-carat gold instead, 22 parts pure gold alloyed with 2 parts silver or copper, hard enough to wear daily, gold enough to still be called gold. Every alloy in this chapter, steel, stainless steel, brass, bronze, solder, amalgam, 22-carat gold, is the same underlying idea: take a metal's natural properties, and engineer them deliberately into something more useful than the pure element ever was alone.