Single displacement reaction

A single-displacement reaction, also known as single replacement reaction or exchange reaction, is a chemical reaction in which one element is replaced by another in a compound.[1][2][3]

It can be represented generically as:

${\displaystyle {\ce {A + BC -> AC + B}}}$

where either

• ${\displaystyle {\ce {A}}}$ and ${\displaystyle {\ce {B}}}$ are different metals (or any element that forms cation like hydrogen) and ${\displaystyle {\ce {C}}}$ is an anion;[2] or
• ${\displaystyle {\ce {A}}}$ and ${\displaystyle {\ce {B}}}$ are halogens and ${\displaystyle {\ce {C}}}$ is a cation.[2]

This will most often occur if ${\displaystyle {\ce {A}}}$ is more reactive than ${\displaystyle {\ce {B}}}$, thus giving a more stable product. The reaction in that case is exergonic and spontaneous.

In the first case, when ${\displaystyle {\ce {A}}}$ and ${\displaystyle {\ce {B}}}$ are metals, ${\displaystyle {\ce {BC}}}$ and ${\displaystyle {\ce {AC}}}$ are usually aqueous compounds (or very rarely in a molten state) and ${\displaystyle {\ce {C}}}$ is a spectator ion (i.e. remains unchanged).[1]

${\displaystyle {\ce {A(s) + \underbrace{B+(aq) + C^{-}(aq)}_{BC(aq)}-> \underbrace{A+(aq) + C^{-}(aq)}_{AC(aq)}+ B(s)}}}$
When a copper wire is dipped in a silver nitrate solution, copper displaces silver, turning the solution blue and solid silver precipitates out ("silver tree"): Cu + AgNO₃ → Cu(NO₃)₂ + Ag↓
NCSSM video on single displacement reaction
Formation of tin crystals as zinc displaces tin, seen under microscope.

In the reactivity series, the metals with the highest propensity to donate their electrons to react are listed first, followed by less reactive ones. Therefore, a metal higher on the list can displace anything below it. Here, is a condensed version of the same:[1]

${\displaystyle {\ce {K}}>{\ce {Na}}>{\ce {Ca}}>{\ce {Mg}}>{\ce {Al}}>{\color {gray}{\ce {C}}}>{\ce {Zn}}>{\ce {Fe}}>{\color {gray}{\ce {NH4^+}}}>{\color {gray}{\ce {H+}}}>{\ce {Cu}}>{\ce {Ag}}>{\ce {Au}}}$
(Hydrogen, carbon and ammonium — labeled in gray — are not metals.)

Similarly, the halogens with the highest propensity to acquire electrons are the most reactive. The activity series for halogens is: [1][2][3]

${\displaystyle {\ce {F2>Cl2>Br2>I2}}}$

Due to the free state nature of ${\displaystyle {\ce {A}}}$ and ${\displaystyle {\ce {B}}}$, single displacement reactions are also redox reactions, involving the transfer of electrons from one reactant to another.[4] When ${\displaystyle {\ce {A}}}$ and ${\displaystyle {\ce {B}}}$ are metals, ${\displaystyle {\ce {A}}}$ is always oxidized and ${\displaystyle {\ce {B}}}$ is always reduced. Since halogens prefer to gain electrons, ${\displaystyle {\ce {A}}}$ is reduced (from ${\displaystyle {\ce {0}}}$ to ${\displaystyle {\ce {-1}}}$) and ${\displaystyle {\ce {B}}}$ is oxidized (from ${\displaystyle {\ce {-1}}}$ to ${\displaystyle {\ce {0}}}$).

Cation replacement

Here one cation replaces another:

${\displaystyle {\ce {A + BC -> AC + B}}}$

(Element A has replaced B in compound BC to become a new compound AC and the free element B.)

Some examples are:

${\displaystyle {\ce {Fe + CuSO4 -> FeSO4 + Cu(v)}}}$
(Blue vitriol)____(Green vitriol)
${\displaystyle {\ce {Zn + CuSO4 -> ZnSO4 + Cu(v)}}}$
(Blue vitriol)___(White vitriol)
${\displaystyle {\ce {Zn + FeSO4 -> ZnSO4 + Fe(v)}}}$
(Green vitriol) (White vitriol)

These reactions are exothermic and the rise in temperature is usually in the order of the reactivity of the different metals.[5]

If the reactant in elemental form is not the more reactive metal, then no reaction will occur. Some examples of this would be the reverse.

${\displaystyle {\ce {Fe + ZnSO4 ->}}}$  No Reaction
External video
CuCl2 and Iron Part 2, 2011
Reaction of CuCl2 with Al, 2011
Copper Sulfate and Iron: Part 1, 2011

Anion replacement

Here one anion replaces another:

${\displaystyle {\ce {A + CB -> CA + B}}}$

(Element A has replaced B in the compound CB to form a new compound CA and the free element B.)

Some examples are: ${\displaystyle {\ce {Cl2 + 2NaBr -> 2NaCl + Br2(v)}}}$  ${\displaystyle {\ce {Br2 + 2KI -> 2KBr + I2(v)}}}$  ${\displaystyle {\ce {Cl2 + H2S -> 2HCl + S(v)}}}$

Again, the less reactive halogen cannot replace the more reactive halogen:

${\displaystyle {\ce {I2 + 2KBr ->}}}$  no reaction

Common reactions

Metal-acid reaction

Metals react with acids to form salts and hydrogen gas.

Liberation of hydrogen gas when zinc reacts with hydrochloric acid.
${\displaystyle {\ce {Zn(s) + 2HCl(aq) -> ZnCl2(aq) + H2 ^}}}$ [2][3]

However less reactive metals can not displace the hydrogen from acids.[3] (They may react with oxidizing acids though.)

${\displaystyle {\ce {Cu + HCl ->}}}$  No reaction

Reaction between metal and water

Metals react with water to form metal oxides and hydrogen gas. The metal oxides further dissolve in water to form alkalies.

${\displaystyle {\ce {Fe(s) + H2O (g) -> FeO(s) + H2 ^}}}$
${\displaystyle {\ce {Ca(s) + 2H2O (l) -> Ca(OH)2(aq) + H2 ^}}}$

Explosive reaction of sodium in water, shattering the glass vessel.

The reaction can be extremely violent with alkali metals as the hydrogen gas catches fire.[2]

Metals like gold and silver, which are below hydrogen in the reactivity series, do not react with water.

External video
Reaction of Potassium and Water, 2011
Reaction of Magnesium and Water, 2011

Metal extraction

Coke or more reactive metals are used to reduce metals by carbon from their metal oxides,[6] such as in the carbothermic reaction of zinc oxide (zincite) to produce zinc metal:

${\displaystyle {\ce {ZnO + C -> Zn + CO}}}$

and the use of aluminium to produce manganese from manganese dioxide:

${\displaystyle {\ce {3MnO2 + 4Al -> 3Mn + 2Al2O3}}}$

Such reactions are also used in extraction of boron and silicon.

${\displaystyle {\ce {3SiO2 + 4Al -> 3Si + 2Al2O3}}}$
${\displaystyle {\ce {B2O3 + 3Mg -> 2B + 3MgO}}}$

Thermite reaction

Using highly reactive metals as reducing agents leads to exothermic reactions that melt the metal produced. This is used for welding railway tracks.[6]

Thermite reaction proceeding for a railway welding: Shortly after this, the liquid iron flows into the mould around the rail gap
${\displaystyle {\ce {Fe2O3(s) + 2 Al(s) -> 2 Fe(l) + Al2O3(s)}}}$

a(Haematite)

${\displaystyle {\ce {3CuO + 2Al -> 3Cu + Al2O3}}}$

Silver tarnish

A tarnished silver coin

Silver tarnishes due to the presence of hydrogen sulfide, leading to formation of silver sulfide.[7][2]

${\displaystyle {\ce {4Ag + 2H2S + O2 -> 2Ag2S + 2H2O}}}$
${\displaystyle {\ce {3Ag2S + 2Al -> 6Ag + Al2S3}}}$

Extraction of halogens

Chlorine is manufactured industrially by the Deacon's process. The reaction takes place at about 400 to 450 °C in the presence of a variety of catalysts such as ${\displaystyle {\ce {CuCl2}}}$ .

${\displaystyle {\ce {4HCl + O2 -> 2 Cl2 + 2H2O}}}$

Bromine and iodine are extracted form brine by displacing with chlorine.

${\displaystyle {\ce {2HBr + Cl2 -> 2HCl + Br2 ^}}}$
${\displaystyle {\ce {2HI + Cl2 -> 2HCl + I2 ^}}}$

1. ^ a b c d "Single replacement reactions". Khan Academy.`{{cite web}}`: CS1 maint: url-status (link)
2. "Single Replacement Reactions". Chemistry LibreTexts. 2016-06-27.`{{cite web}}`: CS1 maint: url-status (link)
3. ^ a b c d "Types of Chemical Reactions: Single- and Double-Displacement Reactions". courses.lumenlearning.com.`{{cite web}}`: CS1 maint: url-status (link)
5. ^ "Exothermic metal displacement reactions". RSC Education. Nuffield Foundation.`{{cite web}}`: CS1 maint: others (link) CS1 maint: url-status (link)
6. ^ a b "Displacement reactions of metal oxides". BBC Bitesize.`{{cite web}}`: CS1 maint: url-status (link)