Physical properties
Property | Metals | Non-metals |
|---|---|---|
State | solids (except mercury, a liquid) | solids, liquids (bromine) or gases |
Lustre | shiny | dull (except iodine and graphite) |
Hardness | generally hard (except sodium and potassium, which can be cut with a knife) | generally soft (except diamond, the hardest natural substance) |
Malleability | can be beaten into sheets; gold and silver are the most malleable | brittle |
Ductility | can be drawn into wires; 1 g of gold can make a wire about 2 km long | not ductile |
Conductivity | good conductors of heat and electricity; silver and copper are best | poor conductors (except graphite) |
Sonority | make a ringing sound when struck | not sonorous |
Melting point | generally high (except gallium and caesium, which melt in your hand) | generally low |
Lead and mercury are comparatively poor conductors of heat.
Chemical properties of metals
With oxygen: metals form basic oxides. 2Mg + O₂ → 2MgO.
Some oxides, such as aluminium oxide and zinc oxide, react with both acids and bases. These are amphoteric oxides.
Reactivity with oxygen varies. Sodium and potassium react so vigorously that they are stored under kerosene. Magnesium burns with a dazzling white flame. Iron burns only as fine filings sprinkled into a flame. Gold and silver do not react with oxygen even at high temperatures.
With water: metals form a hydroxide (or oxide) and hydrogen.
Sodium and potassium react violently with cold water.
Calcium reacts less violently; bubbles of hydrogen make it float.
Magnesium reacts with hot water.
Aluminium, iron and zinc react only with steam, forming oxides.
Lead, copper, silver and gold do not react with water at all.
With dilute acids: metals above hydrogen in the reactivity series give a salt and hydrogen. Copper, silver and gold do not. With nitric acid, hydrogen is usually not released, because nitric acid is a strong oxidising agent (except for Mg and Mn with very dilute HNO₃).
Displacement: a more reactive metal displaces a less reactive one from its salt solution.
Fe + CuSO₄ → FeSO₄ + Cu (the blue solution turns pale green, and copper coats the iron).
The reactivity series
Metals arranged in decreasing order of reactivity:
K > Na > Ca > Mg > Al > Zn > Fe > Pb > [H] > Cu > Hg > Ag > Au
A common memory line: "Please Stop Calling Me A Zebra, I Like Her Calling Me Smart Goat" (Potassium, Sodium, Calcium, Magnesium, Aluminium, Zinc, Iron, Lead, Hydrogen, Copper, Mercury, Silver, Gold).
How metals and non-metals react: ionic compounds
Noble gases have a complete outer shell (an octet) and are unreactive. Other atoms gain or lose electrons to reach this stable arrangement.
Sodium (2, 8, 1) loses one electron to become Na⁺ (2, 8). Chlorine (2, 8, 7) gains one to become Cl⁻ (2, 8, 8). The oppositely charged ions attract strongly, forming sodium chloride. Compounds formed by transfer of electrons are ionic (electrovalent) compounds.
Properties of ionic compounds:
hard, brittle solids,
high melting and boiling points, because the attraction between ions is strong,
soluble in water, insoluble in kerosene and petrol,
conduct electricity when molten or dissolved (the ions can move), but not as solids.
Extraction of metals
Metals occur in the Earth's crust as minerals; a mineral from which a metal can be extracted profitably is an ore. The unwanted earthy material in an ore is gangue, removed first by concentration (enrichment).
The method depends on the metal's reactivity:
Reactivity | Example metals | Method |
|---|---|---|
Low | mercury, copper | heating the ore alone: 2HgS + 3O₂ → 2HgO + 2SO₂, then 2HgO → 2Hg + O₂ |
Middle | zinc, iron, lead | convert to oxide by roasting (sulphide ores, heated in excess air) or calcination (carbonate ores, heated in limited air), then reduce the oxide with carbon: ZnO + C → Zn + CO |
High | sodium, magnesium, calcium, aluminium | electrolytic reduction of the molten chloride or oxide; carbon cannot reduce them |
Highly reactive metals are also used as reducing agents. In the thermit reaction, aluminium reduces iron(III) oxide: Fe₂O₃ + 2Al → 2Fe + Al₂O₃. So much heat is released that the iron is molten; it is used to join railway tracks.
Refining. Impure metals such as copper are purified by electrolytic refining: impure metal is the anode, a thin strip of pure metal is the cathode, and a salt solution of the metal is the electrolyte. Pure metal deposits on the cathode; insoluble impurities settle below the anode as anode mud.
Corrosion and its prevention
Corrosion is the slow eating away of metals by air, moisture and chemicals.
Silver turns black (silver sulphide forms).
Copper gets a green coating (basic copper carbonate).
Iron rusts, forming a brown flaky substance. Rusting needs both air and water.
Prevention: painting, oiling, greasing, galvanising (coating iron with zinc), chrome plating, anodising and making alloys.
An alloy is a homogeneous mixture of a metal with other metals or a non-metal. Steel is iron with a little carbon; stainless steel adds nickel and chromium and does not rust. Brass is copper and zinc; bronze is copper and tin; solder (lead and tin) has a low melting point, for joining wires. Pure gold (24 carat) is too soft, so jewellery uses 22-carat gold alloyed with silver or copper.
Summary
Metals are lustrous, malleable, ductile, sonorous conductors; non-metals generally are not, with notable exceptions.
Metals form basic oxides; Al₂O₃ and ZnO are amphoteric.
The reactivity series predicts reactions with water, acids and salt solutions.
Ionic compounds have high melting points and conduct when molten or dissolved.
Extraction depends on reactivity; corrosion is prevented by coatings and alloys.