Chemical decomposition

Chemical decomposition , or chemical dissolution , is the process or effect of simplifying a chemical unit (common molecule, reaction intermediate , etc.) into two or more pieces. [1] Chemical decomposition is generally considered and defined as the exact opposite of chemical synthesis . In short, a chemical reaction in which two or more products are formed from the same reactant is called a decomposition reaction.

The details of a decomposition process are not always well defined but some process is understood; Breaking bonds requires a lot of energy. Since all decomposition reactions break the bonds holding it together to produce it in its simplest basic parts, the reactions will require some form of this energy in varying degrees. Because of this fundamental rule, it is known that most of these reactions are endothermic, although exceptions exist.

The stability of a chemical compound is ultimately limited when exposed to extreme environmental conditions such as heat , radiation , moisture , or acidity of a solvent . Because of this chemical decomposition is often an unwanted chemical reaction . However chemical decomposition is being used in an increasing number of ways.

For example, this method is employed for many analytical techniques, notably mass spectrometry , conventional gravimetric analysis , and thermogravimetric analysis . Additionally decomposition reactions are used today for many other reasons in the production of a variety of products. One of these is the explosive breakdown reaction of sodium azide [(NaN3 )2 ] into nitrogen gas (N 2 ) and sodium (Na). It is the process that powers the life-saving airbags present in almost all automobiles today.

Decomposition reactions can generally be classified into three categories; Thermal, electrolytic and photolytic decomposition reactions.

Reaction formula

In the breakdown of a compound into its constituent parts, the reaction generalized to chemical decomposition is:

AB → A + B

An example is the electrolysis of water to hydrogen and oxygen gases :

2H 2 O ( L ) → 2 H 2 ( G ) + O 2 ( G )

Additional examples

An example of a spontaneous ( without an external energy source ) decomposition is that of hydrogen peroxide which slowly decomposes into water and oxygen:

2 H 2 O 2 → 2 H 2 O + O 2

This reaction is one of the exceptions to the endothermic nature of decomposition reactions.

Other reactions involving decomposition require the input of external energy. This energy can be in the form of heat, radiation, electricity or light. The latter is the reason some chemical compounds, such as many prescription drugs, are placed and stored in dark bottles that reduce or eliminate the possibility of light reaching them and decomposition.

When heated, the carbonates will decompose. A notable exception is carbonic acid , ( H2CO3 ) . [4] Commonly seen as a “fizz” in carbonated beverages, carbonic acid will automatically decompose over time into carbon dioxide and water. The response is written as:

2 CO 3 → H 2 O + CO 2

Other carbonates will decompose upon heating to produce their respective metal oxides and carbon dioxide. [5] The following equation is an example, where M represents the given metal:

M CO 3 → M O + CO 2

A typical example is that calcium carbonate contains:

CaCO 3 → CaO + CO 2

Metal chlorates also decompose on heating. In this type of decomposition reaction, metal chloride and oxygen gas are the products. Here, again, M represents the metal:

M ClO 3 → 2 M Cl+ 3 O 2

A common decomposition of chlorate occurs in the reaction of potassium chlorate where oxygen is the product. It can be written as:

2 KClO 3 → 2 KCl + 3 O 2