Chemical Reactions Types

Types of Chemical Reactions According to Composition

On the basis of nature of reactants and products, chemical reactions have been divided into five common types:

1.	Combustion
2.	Decomposition Reactions or Thermal Decomposition
3.	Addition Reactions or Synthesis
4.	Single Displacement Reactions or Substitution Reactions
5.	Double displacement Reactions OR Double decomposition Reactions OR Metathesis
	(a)    Neutralization
	(b)   Hydrolysis
	(c)    Precipitation reactions
	(d)   Acid displacement reactions

1. Combustion Reactions

Definition

The rapid reaction of a chemical substance on ignition with either free oxygen or oxygen of the air to give CO₂ and water vapours, with the rapid evolution (release) of heat and flame is called Combustion Reactions. Combustion is the characteristic reaction of organic compounds especially of hydrocarbons. 

Organic compounds

+oxygen

→ CO2 +

H2O(g)

ΔH=- kJ/mol

Nature

Combustion is an exothermic process and evolves large amount of heat called heat of combustion and it is the fuel value of compounds.

Examples

1.   Carbon (coal) burns in air with smoky flame to produce CO₂ gas and heat.

C

+

O2

→

CO2

ΔH = - 393.7

kJ/mol

2.   Methane (Sui gas or marsh gas) burns in air with non-luminous flame forming CO₂ gas and water vapours releasing large amount of heat called heat of combustion.

	CH4	+	2O2	→	CO2	+	2H2O(l)	ΔH= - 890.4 kJ/mol
	CH4	+	2O2	→	CO2	+	2H2O(g)	ΔH = - 802kJ/mol

2.      Decomposition Reactions

Definition

A chemical reaction in which a single compound breaks down (or splits up) into two or more simpler substances by the application of heat (energy) is called Simple Decomposition Reactions. It is reverse of synthesis reactions, therefore, it is also known as Desynthesis Reactions.

AB

→

A

+

B

DH=+ kJ/mol

Nature

These reactions are always endothermic i.e. heat energy is required to bring about decomposition of compounds, so these reactions more correctly called Thermal Decomposition. [Only compounds may decompose, elements cannot].

Examples

1.  Decomposition of lime stone (calcium carbonate) at high temperature into quick lime and CO₂

CaCO3

→

CaO

+

CO2­

2.  Decomposition of potassium chlorate on heating into potassium chloride and oxygen gas. 

2KClO3

→

2KCl

+

3O2­

3.  Fermentation, the biochemical degradation of glucose by yeast into ethyl alcohol and carbon dioxide.

C6H12O6

→

2C2H5OH

+

2CO2­

General Rules for Decomposition of Some Compounds

1.  Decomposition of Carbonates Salts into metal oxide and carbon dioxide gas

	Metal carbonate	¾¾¾¾®	Metal oxide	+	CO2
	MCO3	¾¾¾¾®	MO	+	CO2   [M  =  Ca, Mg, Zn]

2.  Decomposition of Bicarbonates Salts into carbonate salt, water and carbon dioxide

	Metal bicarbonate	→	Metal carbonate	+	H2O	+	CO2­
	2MHCO3	→	M2CO3	+	H2O	+	CO2 ­	[M  =  Na, K]
	M(HCO3)2	→	MCO3¯	+	H2O	+	CO2 ­	[M  =  Ca, Mg]

3.  Decomposition of Group II Nitrates Salts into metal oxide, NO₂ gas and O₂ gas

	Metal nitrates	→	Metal oxide	+	NO2	+	O2­
	2M(NO3)2	→	2MO	+	4NO2	+	O2­	[M  =  Mg, Ca,Ba, Pb,Zn]
	4LiNO3	→	2Li2O	+	4NO2	+	O2­

4.  Decomposition of Group IA Nitrates Salts into metal nitrite and O₂ gas

	Metal nitrates	→	Metal nitrite	+	O2­
	2M(NO3)2	→	2MNO2	+	O2­	[M  =  Na, K, Rb, Cs]

3. Synthesis or Addition Reaction

A chemical reaction in which two or more substances (element or compound) combine to form a single compound is called an Addition Reaction or Synthesis Reaction or Combination Reaction or Composition Reaction or Formation Reaction”. These reactions are reverse of decomposition reactions. Addition reactions are usually exothermic. [Addition reactions may take place between metal and a non-metal, between two non-metals, between a compound and an element (non-metal) or between two compounds. In organic chemistry, addition reactions are characteristic of unsaturated organic compounds, which are of three types namely electrophilic addition, nucleophilic addition and polymerization].

Examples

Addition Reaction between Metals and Non-metals

2Na      +

Cl2 →

2NaCl

4Na      +

O2→

2Na2O

2Na      +    O2

→Na2O2

   Fe        +    S →FeS

Addition Reaction between Metals and Non-metals

	H2	+	Cl2	¾¾¾®	2HCl
	H2	+	S	¾¾¾®	H2S

Addition Reaction between compound and element

	C2H4	+	H2	→	C2H6
	C2H2	+	2H2	→	C2H6

Addition Reaction between two compounds

	CaO	+	CO2	→	CaCO3		CaO	+	SO2	→	CaSO3
	CaO	+	H2O	→	Ca(OH)2		CaO	+	SO3	→	CaSO4

4. Single Displacement Reaction

A chemical reaction in which an atom or group of atoms (or radical) of a molecule of a compound is displaced by another atom or group of atoms (or radical) is called Single Displacement Reaction or Simple Displacement or Substitution Reaction.

The tendency of an atom or radical to displace another depends upon its electropositive or electronegative nature. [The general principle of displacement is that a more electropositive element can displace a less electropositive element form its compound (i.e. a more reactive metal can displace a less reactive metal from its compound). Similarly a more electronegative element can replace a weak or less electronegative element form its compound. The activity series of metals helps in guiding the displacement of an electropositive element form its compound by another electropositive element. The general rule for metal displacement is that a metal lying above in activity series can displace a metal lying below it].

C More E.P atom	+	AB Less E.P atom	→	CB	+	A
D More E.N atom	+	AB Less E.N atom	→	AD	+	B

Displacement of less electronegative element from its compounds by more electronegative element

The displacement power of a non-metal depends upon its reactivity which in turn depends upon electronegativity (E.N). A more electronegative element can displace a less electropositive element from its compounds. The most electronegative element is F (E.N. = 4), oxygen is the 2^nd electronegative element (3.5) and Cl is the 3^rd most electronegative element3). The electronegativity of Br is 2.8 and that of I is 2.5. The electronegativity of S is 2.8 and that of H is 2.1. Chlorine replaces bromine or iodine from bromide or iodide salts to form respective chloride and bromine vapours or iodine vapours.

	Cl2	+	Br–	→	2Cl–	+	Br2
	Cl2	+	2KBr	→	2NaCl	+	Br2
	Cl2	+	2KI	→	2KI	+	I2

Displacement of less electropositive element from its compounds by more electropositive element

The displacement power of  a metal depend upon its reactivity which in turn depend upon its position in the reactivity series which is a convenient summary of the results of many possible metal or hydrogen displacement reactions. This series is based on metal-acid reaction and metal compound-metal reaction. Accroring to this series, any metal above hydrogen will displace it from water or from an acid, but metals below hydrogen will not react with either water or an acid.  [In fact, any species listed in the series will react with a compound containing any species listed below it]. A list of metals arranged in order of decreasing ease of oxidation is called an Activity Series. It ranks the elements in order of their reducing ability in aqueous solution. The metals at the top are most readily oxidized and are stronger reducing agents whereas the metals at the bottom of the series are less readily oxidized and are weaker reducing agents.

1. Some metals displace hydrogen from dilute acids, bases or even water or alcohol to produce hydrogen       gas along with respective metal salts.

Zn	+	2HCl	→	ZnCl2	+	H2
Zn	+	2NaOH	→	Na2ZnO2	+	H2
2Na	+	2H-OH	→	2NaOH	+	H2
2Na	+	2C2H5OH	→	2C2H5ONa	+	H2

2.      Magnesium replaces copper from copper sulphate solution to form magnesium sulphate and copper.

Mg	+	CuSO4	→	MgSO4	+	Cu
Zn	+	CuSO4	→	ZnSO4	+	Cu

5.      Double Displacement Reactions (DDR)

Definition

A chemical reaction in which two compounds (reactants) are decomposed to form two new compounds by exchanging (interchanging) their radicals is called Double Decomposition or Double Displacement Reaction (DDR). [Owing to formation of two new compounds, such reactions are also called metathesis reactions. These reactions involve the exchange of partners of two compounds, which is usually brought about by an exchange of their ionic radicals. That is why; these reactions are also referred as Exchange Partners Reactions.

AB

+

CD

→

AD

+

BC

Types of DDR

They are of following four types:

a).     Precipitation Reactions             (characterized by the formation of a solid compound)

b).     Acid displacement Reactions   (characterized by the formation of an acid in vapour state)

c).     Neutralization Reactions         (characterized by the formation of a salt and water)

d).     Hydrolysis                              (characterized by the formation of an acid and a base)

5.A    Precipitation Reactions

Definition                                                                                                             

A double decomposition reaction in which aqueous solution of two compounds (reactants) are decomposed to form two new compounds by exchanging (interchanging) their radicals one of which is in the form of solid particles called precipitate is called precipitation reaction. The common precipitating agents are NaOH, dilute H₂SO₄, dilute HCl, H₂S, NH₄OH, (NH₄)₂CO₃, Na₂HPO₄, AgNO₃, Na₃PO₄, Na₂CO₃,

Examples

Precipitation by using Sodium Carbonate

The mixing of aqueous solutions of calcium chloride with sodium carbonate, gives two new different compounds (sodium chloride and white ppt of calcium carbonate) by the exchange of their radicals.

CaCl2(aq)

+

Na2CO3(aq)

→

2NaCl(aq)

+

CaCO3(s)¯

(insoluble white ppt)

Precipitation by using Silver Nitrate

Silver nitrate is used to precipitate chlorides, bromides, iodides and sulphates radicals from their compounds. The respective ppt of silver chloride, silver bromide, silver iodide and silver sulphate are white, light yellow, deep yellow and white in colour respectively.

NaCl(aq)	+	AgNO3(aq)	→	NaNO3(aq)	+	AgCl(s)¯	(insoluble white ppt)
NaBr(aq)	+	AgNO3(aq)	→	NaNO3(aq)	+	AgBr(s)¯	(insoluble white ppt)
NaI(aq)	+	AgNO3(aq)	→	NaNO3(aq)	+	AgI(s)            ¯	(insoluble white ppt)
CuSO4(aq)	+	2AgNO3(aq)	→	Cu(NO3)2(aq)	+	Ag2SO4(s)¯	(insoluble white ppt)

5.B    Hydrolysis

Definition

The type of double decomposition reaction in which two reactants are salt and water which react to form acid and base resulting in acidic or basic solution is referred to as Hydrolysis and the salt is said to be hydrolyzed.  It is reverse of neutralization.

Salt(aq)

+

Water

→

Acid(aq)

+

Base(aq)

Examples

1. Hydrolysis of Salts of Weak Acids and Strong Bases

Salts of weak acids with strong bases undergo hydrolysis with water producing basic solution due to formation of strong base. Salts of weak acids and strong bases (such as CH₃COONa, Na₂CO₃, KCN, NaCN etc) hydrolyze in water to produce weak acid and strong base. Thus due to formation of strong base, concentration of OH¯ ions become greater which changes the pH of solution towards basic (pH > 7) and turn red litmus blue. Therefore, solution becomes basis. E.g.

Aqueous Solution of CH₃COONa is basic due to hydrolysis

	CH3COONa	+	H-OH	→	NaOH(aq)	+	CH3COOH(aq)
					Strong base		Weak acid
	CH3COO¯	+	H-OH	→	OH¯(aq)	+	CH3COOH (aq)

Na2CO3(aq)

+

2H-OH

→

2NaOH(aq)

+

H2CO3(aq)

2. Hydrolysis of salts of strong acids and weak bases

Salts of strong acids with weak bases undergo hydrolysis with water producing acidic solution due to formation of strong acid. Salts of strong acids and weak bases (such as CuSO₄, NH₄Cl, ZnCl₂, AlCl₃, FeCl₃ etc.) hydrolyze in water to produce strong acid and weak base. Thus due to formation of strong acid, concentration of H⁺ ions become greater which changes the pH of solution towards acidic (pH < 7) and it turns blue litmus red.  Therefore, solution becomes acidic.  e.g.

Aqueous Solution of CuSO₄ is acidic due to hydrolysis

         

	CuSO4 (aq)	+	2H-OH	→	Cu(OH)2(s)	+	H2SO4          (aq)
					Weak base		Strong acid
And	Cu2+	+	2H-OH	→	Cu(OH)2	+	2H+(aq)

         

	FeCl3(aq)	+	3H-OH	→	Fe(OH)3	+	3HCl(aq)
	NH4Cl(aq)	+	H-OH	→	NH4OH	+	HCl(aq)
	AlCl3 (aq)	+	3H–OH	→	Al(OH)3	+	3HCl(aq)
					Weak base		Strong acid

3. Hydrolysis of salts of strong acids and strong bases

Salts of strong acids and strong bases (such as NaCl, KCl, Na₂SO₄, NaNO₃, KNO₃) do not hydrolyze in water. Therefore pH of solution remains same and thus solution becomes neutral. e.g. Aqueous NaCl solution is neutral.

5.C    Acid Displacement Reaction

Definition

A double decomposition reaction in which dilute or concentrated solution of an acid displaces an acid from solid salt in the form of either vapours or solution is called acid displacement reaction.

Examples

A less volatile acid (such as dilute HCl or dilute H₂SO₄) can displace a more volatile acid (such as H₂CO₃, H₂S, H₂S₂O₃, H₂SO₃, HNO₂, CH₃COOH) from their salts. Such a double decomposition reaction is termed as Acid Displacement Reaction.     

The scheme of analysis for the acid radicals of group I is based upon this reaction. The vapours of the acid formed or its decomposition products give an indication of acid radicals (of group I) in the salts. The corresponding acids of acid radicals of group I are highly volatile and most of them are unstable. The vapours of the acid formed or its decomposition products possess a characteristic smell or colours.        

CO32¯	+	H2SO4	→	H2CO3	→	H2O + CO2­	® Effervescence of CO2
HCO3¯	+	H2SO4	→	H2CO3	→	H2O + CO2­	® Effervescence of CO2
S2¯	+	H2SO4	→	H2S	→		® Rotten egg smell
CH3COO¯	+	H2SO4	→	CH3COOH	→		® Vinegar smell
NO2¯	+	H2SO4	→	H2SO3	→	H2O + SO2­	® Smell of burning sulphur
SO32¯	+	H2SO4	→	HNO2	→	H2O+NO2­+ NO	® Brown gas
S2O32¯	+	H2SO4	→	H2S2O3	→	H2O +SO2­+ S¯	® Smell of burning sulphur with yellow  ppt.

The corresponding acids of acid radicals of group II are relatively less volatile and cannot be displaced by dilute acids. They can be displaced only by conc. H₂SO₄ on heating.

With the exception of HCl, HNO₃ and H₂C₂O₄ decomposed itself on heating while HBr and HI are oxidized by concentrated H₂SO₄ to Br₂ and I₂ respectively on heating. The vapours of the acid formed or its decomposition or oxidation products possess a characteristic smell or colour.  The vapours of the acid formed or its decomposition or oxidation products give an indication of acid radicals (of group II) in the salts.   

The scheme of analysis for the acid radicals of group II is based upon the fact that a less volatile acid (such as dilute H₂SO₄ or conc. H₂SO₄) can displace a more volatile acid (such as HCl, HBr, HI,HNO₃, H₂C₂O₄ )from their salts.

Heating is necessary in acid radical analysis of group II after adding conc.H₂SO₄ because in cold we may get colourless pungent smelling vapours of HCl, HBr, HI and HNO₃. After heating HBr and HI are oxidized by conc.H₂SO₄ to form brown vapours of Br₂ and violet vapours of I₂ respectively while HNO₃ is decomposed to give brown gas of NO₂.Thus without heating Br¯, I¯ and NO₃¯ radicals may be mistaken for Cl¯ radical.

Cl¯	+	H2SO4	®	HCl	®	Not decompose	® Colourless pungent smelling vapours
Br¯	+	H2SO4	®	HBr	®	H2O + SO2­+ Br2­	® Brown vapours
I¯	+	H2SO4	®	HI	®	H2O + SO2­ + I2­	® Violet vapours
NO3¯	+	H2SO4	®	HNO3	®	H2O + NO2­+ O2­	® Yellowish brown gas
C2O42¯	+	H2SO4	®	H2C2O4	®	H2O + CO2­+ CO­	® Colourless vapours with coughing odour

5.D. Neutralization

Definition

[Acids are the substances which give H⁺ ion in aqueous solution e.g. HCl, HNO₃ H₂SO₄ H₂CO₃ and CH₃COOH etc. Bases are compounds that yield OH¯ ions in water. The water-soluble bases are called Alkalis.  e.g. Caustic Soda (NaOH), Caustic Potash (KOH),NH₄OH,Ca(OH)₂, Ba(OH)₂. Salt is an ionic compound that is the neutralization product (other than water) of an acid and base, which is the aggregation of cation and anions. e.g. Na⁺Cl¯.] If we add an acid solution to a base solution drop by drop, the acidic character of the acid decreases gradually. A stage will reach at which the resultant solution becomes neutral to litmus. This stage is called neutralization.

 The type of double decomposition reaction in which equivalent quantities of two reactants    acid and base react to form salt and water is called neutralization”. i.e.

                                               

	Acid(aq)	+	Base(aq)	¾¾¾¾®	Salt(aq)	+	Water(l)
More simply,	H+Z¯	+	Y+OH¯ (aq)	¾¾¾¾®	Y+Z¯ (aq)	+	HOH (l)

Since naturalization reaction always involves the reaction between a H⁺ ion of an acid and a OH¯ ion of a base, therefore, water is the resultant product. During neutralization, cation of base and anion of acid remains in solution and do not react. These ions are called Spectator Ions. Thus neutralization may also be defined as:

          “The mutual chemical combination of hydrogen ion or proton (H⁺) of an acid and hydroxide (OH¯) ion of a base to form neutral water molecule is called Neutralization”. Thus neutralization reactions can be denoted by a single net ionic equation:                                 

	H+(aq)	+	OH¯(aq)	¾¾¾¾®	H2O(l)
OR	2H3O+(aq)	+	OH¯(aq)	¾¾¾¾®	2H2O(l)

Examples

	1.	HCl(aq)	+	NaOH(aq)	¾¾¾¾®	NaCl(aq)	+	H2O
	2.	HCl(aq)	+	KOH(aq)	¾¾¾¾®	KCl(aq)	+	H2O
	3.	HNO3(aq)	+	NaOH(aq)	¾¾¾¾®	NaNO3(aq)	+	H2O
	4.	HNO3(aq)	+	KOH(aq)	¾¾¾¾®	KNO3(aq)	+	H2O
	5.	2HCl(aq)	+	Ca(OH)2	¾¾¾¾®	CaCl2	+	2H2O
	6.	HCl(aq)	+	Ca(OH)2	¾¾¾¾®	Ca(OH)Cl	+	H2O
	7.	H2SO4(aq)	+	2KOH(aq)	¾¾¾¾®	K2SO4(aq)	+	2H2O
	8.	H2SO4(aq)	+	KOH(aq)	¾¾¾¾®	KHSO4(aq)	+	H2O
	9.	H2SO4(aq)	+	NaOH(aq)	¾¾¾¾®	NaHSO4(aq)	+	H2O
	10	H2SO4(aq)	+	2NaOH(aq)	¾¾¾¾®	Na2SO4(aq)	+	2H2O

Nature

The neutralization is an exothermic reaction releasing heat called heat of neutralization which is the amount of heat evolved during a neutralization in which 1 mole of water is formed. In other words, heat of neutralization is the amount of heat evolved when 1mole of H⁺ ions of an acid reacts with 1mole of OH¯ ions of a base to form salt and 1 mole of water. In fact, heat of neutralization is the heat of formation of water from H⁺ and OH¯ ions.

H+(aq)

+

OH¯(aq)

¾¾®

H2O(l)

DH neutralization  = - 13700 cal/mol Or - 57.24kJ/mol

The heat of neutralization for any strong acid against any strong base is approximately same i.e. - 13700 cal/mol or -57.3kJ/mol

HCl(aq)	+	NaOH(aq)	¾¾®	NaCl(aq)	+	H2O	DH=- 57.3kJ/mol
HNO3(aq)	+	KOH(aq)	¾¾®	KNO3(aq)	+	H2O	DH=- 57.3kJ/mol
H2SO4(aq)	+	2NaOH(aq)	¾¾®	Na2SO4(aq)	+	2H2O	DH = - 2x57.3kJ

The heat of neutralization for weak acid against weak base or strong acid against weak base (i.e. in case where either acid or base is not completely ionized and the neutralization reaction may not go to completion) may be less than 13700 cal e.g. when strong acid HCl reacts with a weak base Ca(OH)₂, the heat of neutralization is 24700cal or 12350cal/mol

2HCl(aq)

+

Ca(OH)2

¾¾¾¾®

CaCl2

+

2H2O

DH = -24700cals

Types of Neutralization

Neutralization is of following two types:

a) Complete Neutralization

A neutralization in which all H⁺ ions of the acid are neutralized by OH¯ ions of base or vice versa is known as Complete Neutralization. Salt obtained by Complete Neutralization do not have replaceable hydrogen atoms or hydroxide ions and normal (neutral) e.g. NaCl, NaNO₃, K₂SO₄, KNO₃, KCl etc.

HCl(aq)	+	NaOH(aq)	¾¾¾¾®	NaCl(aq)	+	H2O
HNO3(aq)	+	KOH(aq)	¾¾¾¾®	KNO3(aq)	+	H2O

b) Partial Neutralization

A neutralization in which all H⁺ ions of the acid are not neutralized by OH¯ ions of base or vice versa (i.e. all OH¯  ions of a base are not neutralized by the acid) is known as Partial Neutralization. Salt obtained by partial Neutralization have either replaceable hydrogen ions or hydroxide ions and are either acidic or basic. Salts formed by the partial neutralization of an acid by base containing replaceable hydrogen ion are acidic. They further react with bases to form normal salts. e.g. NaHSO₄, KHSO₄, NaHCO₃, KHCO₃ Na₂HPO₄, K₂HpO₄, NaH₂pO₄, KH₂PO₄. Salts formed by the partial neutralization of a base by an acid containing replaceable hydroxyl group are basic. They further react with acids to form normal salts. e.g. Ca(OH)Cl, Mg(OH)Cl, Zn(OH)Cl etc.

H2SO4(aq)	+	NaOH(aq)	¾¾¾¾®	H2O	+	NaHSO4(aq)	(Acidic salt)
H2CO3(aq)	+	KOH(aq)	¾¾¾¾®	H2O	+	KHCO3(aq)	(Acidic salt)

HCl(aq)	+	Ca(OH)2(aq)	¾¾¾¾®	H2O	+	Ca(OH)Cl (aq)	(Basic salt)
HCl (aq)	+	Ca(OH)2 (aq)	¾¾¾¾®	H2O	+	Mg(OH)Cl (aq)	(Basic salt)