Chemical decomposition

Chemical decomposition, analysis or breakdown is the separation of a single chemical compound into its two or more elemental parts or to simpler compounds.[1] Chemical decomposition is usually regarded and defined as the exact opposite of chemical synthesis.

The details of a decomposition process are not always well defined, but some of the process is understood. A lot of energy is needed to break bonds. Since all decomposition reactions break apart the bonds holding it together in order to produce into its simpler, basic parts, the reactions would require some form of this energy in varying degrees. Because of this fundamental rule, it's known that most of these reactions are endothermic, although exceptions do exist.

The stability of a chemical compound is eventually limited when exposed to extreme environmental conditions like heat, radiation, humidity, or the acidity of a solvent. Because of this, chemical decomposition is often an undesired chemical reaction.However, chemical decomposition, is being used in a growing number of ways.

For example, this method is employed for several analytical techniques, notably mass spectrometry, traditional gravimetric analysis, and thermogravimetric analysis. Additionally, decomposition reactions are used today for a number of other reasons in the production of a wide variety of products, some of which may be surprising. One of these is the explosive breakdown reaction of sodium azide [(NaN3)2] into nitrogen gas (N2) and sodium (Na). It is this process which powers the life-saving airbags present in virtually all of today's automobiles.[2]

Of the six known basic decomposition reactions, this discussion will focus on what are referred to as the 'three broad types' and considered to be the most common,These three are the thermal, electrolytic, and catalytic decomposition reactions.[3]

Reaction formula

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

AB → A + B

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

2 H2O(I) → 2 H2 + O2

Additional examples

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

2 H2O2 → 2 H2O + O2

It should also be noted that this reaction is one of the exceptions to the endothermic nature of decomposition reactions. Because it occurs spontaneously, it is, by definition, an exothermic reaction.

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

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

H2CO3 → H2O + CO2

Other carbonates will decompose when heated to produce their corresponding metal oxide and carbon dioxide.[5] The following equation is an example, where M represents the given metal:

MCO3 → MO + CO2

A specific example is that involving calcium carbonate:

CaCO3 → CaO + CO2

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

2 MClO3 → 2 MCl+ 3 O2

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

2 KClO3 → 2 KCl + 3 O2

See also

References

  1. ↑ Helmenstine, Anne Marie (June 11, 2016). "What Is a Decomposition or Analysis Reaction in Chemistry?". Chemical Decomposition Reaction (Website). Retrieved 2017-05-01.
  2. ↑ "Chemical reactions in Everyday life". prezi.com. Retrieved 2017-05-01.
  3. ↑ Lindsay, K (May 1, 2014). "Types of Decomposition Reactions". Learning Tools For Free. Retrieved January 5, 2017.
  4. ↑ ibburke (2011-03-27). "Decomposition of Carbonic Acid Culminating by Elizabeth Burke". ibburke. Retrieved 2017-03-04.
  5. ↑ Walker, MS (2016) [Available now]. "Synthesis and Decomposition Reactions". Quizlet.com/MSWalker22 (Audio-Video Online Lecture). Online Series in Organic Chemistry. Retrieved 2017-03-04. (Registration required (help)).
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