Metals and Reactivity

Reactivity Series and Trends:

The reactivity series arranges metals (and the non-metals carbon and hydrogen) in order from most
reactive to least reactive

. Metals high in the series lose electrons easily (forming positive ions readily) and thus react more
vigorously, whereas metals low in the series are very unreactive (they do not oxidize easily)

. The series from most reactive to least reactive is: potassium (K), sodium (Na), calcium (Ca),
magnesium (Mg), aluminium (Al), carbon (C), zinc (Zn), iron (Fe), hydrogen (H), copper (Cu), silver (Ag),
gold (Au)

. A metal above hydrogen in the series can displace hydrogen from acids (react with acids to produce H₂
gas), while metals below hydrogen (like Cu, Ag, Au) do not react with dilute acids

. Similarly, very reactive metals (K, Na, Ca) react even with cold water to produce a metal hydroxide and
hydrogen gas, while metals like magnesium only react with steam (hot water) to give metal oxide and H₂

. Metals low in the series (e.g. copper, silver, gold) do not react with water or steam at all. (Aluminium is
high in the series but appears unreactive because its surface is protected by a hard oxide layer.)

Evidence of reactivity: more reactive metals show more vigorous reactions:

 With water: Potassium and sodium react violently with cold water (moving on the surface,
producing H₂ which may ignite), calcium reacts steadily with cold water, magnesium reacts very
slowly with cold water but burns in steam to produce magnesium oxide and H₂.

-
 With acids: Magnesium, zinc, iron, etc. react with dilute acids to produce a salt and hydrogen
(observable as fizzing). Copper, silver, gold (below H) do not react with acids. The speed of
hydrogen gas evolution (and whether it occurs at all) indicates the metal’s place in the series.
Always test for hydrogen gas using a lit splint – it burns with a characteristic “pop” sound in the
presence of H₂.

O₂ → 2MgO). A metal high in the series may even tarnish in air at room temperature. Combustion
 With oxygen: Most metals will oxidize when heated in oxygen, forming metal oxides (e.g. 2Mg +

form iron(III) oxide: 4Fe + 3O₂ → 2Fe₂O₃ (iron(III) oxide). This tendency to form oxide also
in oxygen is an oxidation (addition of O₂) reaction. For example, heated iron reacts with oxygen to

correlates with reactivity.

Displacement Reactions:
A more reactive metal can displace a less reactive metal from its compounds. In a displacement
reaction, a solid metal is placed into a solution of a compound (usually a salt) of a less reactive metal.
The more reactive metal will take the place of the less reactive metal in the compound, forming a new
solution and releasing the less reactive metal. For example, iron (which is above copper in the series) will

 Fe (s) + CuSO₄ (aq) → FeSO₄ (aq) + Cu (s)

displace copper from copper(II) sulfate:

Here iron becomes iron(II) sulfate in solution and elemental copper is deposited. The iron, being more
reactive, loses electrons (gets oxidized to Fe²⁺), while Cu²⁺ gains electrons to form Cu metal. If you put a
strip of copper metal into iron(II) sulfate, no reaction occurs because copper is less reactive than iron.
Displacement reactions thus provide experimental evidence for the ordering of metals in the reactivity
series In general, “a more reactive metal will displace a less reactive metal from its compound”. This is
true for displacement of metals from their oxides as well: e.g. aluminium can remove oxygen from
iron(III) oxide (the thermite reaction), because aluminium is above iron in the series.

Why reactivity differs: Metals higher in the series have a greater tendency to form positive ions (they
lose their outer electrons more readily). This is often due to a weaker hold of the nucleus on outer
electrons (more electron shells, etc.). Thus, a metal that easily loses electrons (like potassium or
magnesium) will readily enter compounds (and thus react), whereas one that holds its electrons tightly
(like gold) resists reacting. Displacement reactions are essentially redox reactions where the more
reactive metal is oxidized (loses electrons) and the less reactive metal ion is reduced (gains electrons).

Extraction of Metals:

The method used to extract a metal from its ore depends on its position in the reactivity series.

In General
 Highly reactive metals (above carbon, e.g. potassium, sodium, calcium, magnesium, aluminium)
are very stable as compounds and cannot be reduced by carbon. These are extracted by
electrolysis of their molten salts or oxides For example, aluminium is extracted by electrolysing
molten aluminium oxide (alumina) dissolved in cryolite. (Aluminium’s main ore is bauxite
(Al₂O₃·H₂O); it is purified to alumina and then electrolyzed in the Hall-Héroult cell). During
aluminium electrolysis, Al³⁺ ions gain electrons at the cathode to form Al metal and oxide ions are
oxidized at the carbon anodes to form oxygen gas (which then reacts with the carbon anodes to
produce CO₂).
 Metals of medium reactivity (e.g. zinc, iron, tin, lead) are often extracted by reduction with
carbon or carbon monoxide in a furnace These metals are below carbon in the reactivity series so

iron ore (hematite, Fe₂O₃) by carbon monoxide reduction: Fe₂O₃ + 3CO → 2Fe + 3CO₂. In the blast
carbon can remove oxygen from their oxides. For instance, iron is extracted in a blast furnace from

furnace, coke (carbon) is burnt to CO₂, then CO₂ is reduced by hot carbon to CO, which in turn
reduces the iron(III) oxide ore to iron. Limestone (CaCO₃) is added to remove acidic impurities by
forming slag (CaSiO₃). The molten iron collects at the bottom of the furnace. This carbon reduction
method is cheap and works for metals below carbon (like ZnO to Zn as well).

 Less reactive metals (e.g. copper, silver, gold) occur in nature either as sulfides, oxides, or even
native (especially silver and gold). They can be obtained by simple heating or reduction by carbon

→ 2Cu + CO₂. Some unreactive metals like silver and gold are found in elemental form and just
without sophisticated methods. For example, copper(II) oxide can be reduced by carbon: 2CuO + C

need physical separation from ore.

  Overall, the ease of reduction is linked to reactivity: very reactive metals form very stable
compounds requiring powerful methods (electricity) to extract, whereas less reactive metals’
compounds are easier to reduce Extraction processes often involve redox reactions: the metal
ions are reduced to metal, while the reducing agent (carbon, CO, or electrons from electrolysis) is
oxidized. For example, in iron extraction, carbon monoxide is the reducing agent (getting oxidized
to CO₂) and Fe₂O₃ is reduced to Fe. In aluminium extraction, aluminum oxide is reduced to Al at
the cathode (gain of electrons) and oxide ions are oxidized at the anode (loss of electrons, forming
O₂).

Redox Reactions (Oxidation and Reduction):

Redox reactions involve simultaneous oxidation and reduction. Oxidation and reduction can be
defined in two useful ways:

example, in the extraction of iron: Fe₂O₃ + 3CO → 2Fe + 3CO₂, the iron(III) oxide is reduced (loses
 In terms of oxygen: Oxidation is the gain of oxygen, while reduction is the loss of oxygen. For

oxygen to become Fe), while carbon monoxide is oxidized (gains oxygen to become CO₂). This is a
redox reaction: iron(III) oxide acts as an oxidizing agent (giving oxygen to CO), and CO acts as a
reducing agent (removing oxygen from Fe₂O₃).

 In terms of electrons: Oxidation is loss of electrons, Reduction is gain of electrons. A helpful

reaction example CuO + Mg → MgO + Cu, magnesium gives up electrons (Mg → Mg²⁺ + 2e⁻, it is
mnemonic is OIL RIG: “Oxidation Is Loss, Reduction Is Gain (of electrons)”. During a displacement

oxidized) and copper(II) ions gain those electrons (Cu²⁺ + 2e⁻ → Cu, it is reduced). So Mg is the
reducing agent (causing Cu²⁺ to be reduced) and CuO is the oxidizing agent (causing Mg to be
oxidized).

In any redox reaction, one species is oxidized and another is reduced. These processes always occur
together. Common signs: if an element’s oxidation state increases, it’s oxidized; if it decreases, it’s
reduced. Addition of oxygen or removal of hydrogen are oxidation (in inorganic contexts), and the
reverse are reduction. Many reactions in metals topics are redox:

 Metal + acid (the metal atoms oxidize to metal ions, H⁺ ions are reduced to H₂ gas).
 Metal displacement (the more reactive metal oxidizes, the less reactive metal ion is reduced to
metal).

Oxidizing agents are substances that oxidize others (and get reduced themselves), often providing

oxygen or donating electrons. For instance, in rusting (iron + oxygen → iron oxide), oxygen is the
oxygen or accepting electrons. Reducing agents reduce others (and get oxidized), often removing

oxidizing agent; in a blast furnace, carbon monoxide is the reducing agent for iron ore.

Corrosion and Its Prevention:

Corrosion is the gradual destruction of metals by chemical reactions with their environment. For iron,
this is called rusting – the formation of hydrated iron(III) oxide. Rusting requires both oxygen and water
to be present.
4Fe + 3O₂ + xH₂O → 2Fe₂O₃·xH₂O (hydrated iron(III) oxide, rust).
The overall reaction can be represented (simplified) as:

Rust is a flaky, porous substance that easily crumbles off, exposing fresh iron beneath to further
corrosion. Other metals corrode to oxides too (e.g. aluminum forms Al₂O₃), but some oxide coatings (like
aluminium’s) are protective and stop further corrosion. Iron’s rust is not protective, causing continuous
corrosion in the presence of moisture and air. Salt (e.g. NaCl from road salt or sea water) accelerates
rusting by providing ions that facilitate electron transfer.

Preventing corrosion: The key is to prevent contact with oxygen and/or water. Common barrier
methods include painting, coating with oil/grease, and plastic coating. For example, painting iron gates
or applying grease to tools forms a physical barrier that keeps water/air out, thus preventing rust.
Another method is electroplating – coating iron with a thin layer of a less reactive metal (like tin or
chromium) by electrolysis, which physically protects it.

A more advanced method is sacrificial protection (cathodic protection). This involves attaching a piece
of a more reactive metal (like zinc or magnesium) to the iron object. The more reactive metal will oxidize
(corrode) in preference to the iron, “sacrificing” itself. For instance, iron ship hulls are protected by
bolting on magnesium or zinc blocks; these blocks slowly corrode instead of the iron. Similarly,
galvanizing is coating iron with zinc. The zinc coating not only serves as a barrier but, if scratched, will
continue to protect the iron by sacrificially corroding (since Zn is higher than Fe in the series). In
sacrificial protection, the iron stays as the cathode (protected, no oxidation) while the more reactive
metal (anode) oxidizes. As long as the sacrificial metal is present, the iron is safe from rust.

In summary, to prevent rust: keep oxygen and water away (by paint, oil, plastic, plating) and/or use a
sacrificial anode (a more reactive metal that corrodes instead of iron). Stainless steel is another solution
– it’s an alloy of iron with chromium/nickel that forms its own protective oxide layer and doesn’t rust
easily.\

Equations for Key Reactions:

Below are some important reactions involving metals and their equations:

Metal + Acid → Salt + Hydrogen: (only for metals above hydrogen in series)
Example: Zn + 2HCl → ZnCl₂ + H₂↑ – Zinc reacts with dilute hydrochloric acid to form zinc chloride
and hydrogen gas. (Test H₂ with a lit splint → “pop” sound).

Metal + Water → Metal Hydroxide + Hydrogen: (for very reactive metals, e.g. Group 1 and

Example: 2Na + 2H₂O → 2NaOH + H₂↑. Sodium reacts vigorously with water producing sodium
calcium with cold water)

instead: Mg + H₂O (g) → MgO + H₂.

hydroxide (alkaline solution) and hydrogen gas. Less reactive metals like Mg will react with steam

Metal + Oxygen → Metal Oxide: (oxidation)

Examples: 4Fe + 3O₂ → 2Fe₂O₃ (iron rusting/oxidation) ; 2Mg + O₂ → 2MgO (magnesium burning
in air). These are synthesis reactions forming oxides. Many metals need heating to burn in oxygen,
except very reactive ones which tarnish even at RT.

Example: Fe + CuSO₄ → FeSO₄ + Cu – iron displacing copper from solution (iron enters solution as
Displacement Reaction:

Fe²⁺, copper metal is released). If no reaction is observed when a metal is added to another’s salt
solution, it means the solid metal is lower in reactivity than the metal in solution.

2Al + Fe₂O₃ → 2Fe + Al₂O₃. Here aluminium (highly reactive) reduces iron(III) oxide to iron. This
Thermite Reaction (metal displacement of oxygen):

reaction produces molten iron (used for welding railway tracks). It’s an example of a more reactive
metal taking oxygen from a less reactive metal’s oxide.

ZnO + C → Zn + CO (at high temperature). Carbon (as coke) removes oxygen from zinc oxide to
Reduction of Metal Oxide with Carbon:

yield zinc metal. Similarly, in the blast furnace: Fe₂O₃ + 3CO → 2Fe + 3CO₂. Carbon (in form of CO)
is the reducing agent for iron oxide.

Aluminium: 2Al₂O₃ → 4Al + 3O₂ (overall electrolytic decomposition of alumina). At cathode: Al³⁺ +
Electrolysis for Metal Extraction:

3e⁻ → Al; at anode: 2O²⁻ → O₂ + 4e⁻.

Note: The oxygen reacts with the carbon anodes to produce CO₂, hence anodes burn away and
need replacement.