A detailed overview of Laws of Chemical Combination

Chemistry, from its origin , is all about elements and their transformations. Understanding how elements interact with each other and combine to form new substances is studied under Laws of Chemical Combination.

In this comprehensive guide, we’ll discuss each law in detail, delving into its statement, exploring real-world examples, and discussing any exceptions that might exist in different laws of chemical combination.

1. The Law of Conservation of Mass:-

First point of our discussion is the Law of Conservation of Mass. This cornerstone principle was formulated by Antoine Lavoisier in 18th century. It states that:-

“Mass Can neither be created nor be destroyed , Only Rearrange or changes from one form to another”.

In a chemical reaction, “the total mass of reactants is always equal to total mass of products”.

So law of conservation of Mass can be represented as :- “Mass of Reactants =Mass of Products”

In simpler terms, during a chemical reaction, no mass is created or destroyed. It simply transforms from one form(From reactants) to another form (into products).

1.1 Examples of Law of Conservation of Mass :-

Let’s illustrate this law with some practical examples :-

1.1.1 Burning of Methane :-

Consider the burning of methane (CH₄) in oxygen (O₂) to form carbon dioxide (CO₂) and water (H₂O). The balanced chemical equation for this reaction is:

CH₄ + 2O₂ → CO₂ + 2H₂O

Imagine we burn 16 grams of methane (1 Mol methane) with 64 gram of dioxygen (2 mol dioxygen ) then According to the law of conservation of mass, the combined mass of reactants (CH₄ and O₂ = 16+ 64 =80 gm ) will be equal to the combined mass of products (CO₂ and H₂O =44 + 36 =80 gm).

1.1.2 Formation of water :-

Hydrogen and oxygen combine in presence of catalyst (and other experimental conditions) to give water according to the reaction :-

2H₂ + O₂ → 2H₂O

Mass of Reactants , H₂ =4 gm , O₂ = 32 gm , Total reactants = 4+32 = 36 gm ,

Mass of Products =mass of 2 moles of water = 2×18 = 36 gm

Here the combined mass of reactants is equal to combined mass of products .hence the law is confirmed.

1.1.3 Burning of Carbon :-

Carbon burns in sufficient supply of oxygen to give carbon dioxide gas according to the following reaction :-

C + O₂ → CO₂

Here also ,the mass of reactants , carbon = 12 gm , oxygen =32 gm ,Total = 12+32 =44 gm

Total mass of products=mass of CO₂ = 44 gm] , again the law of conservation of mass is confirmed

1.2 Exceptions of law of conservation of mass :-

While the Law of Conservation of Mass holds true for most classical chemical reactions, But it’s important to acknowledge a few exceptions:

Nuclear Reactions:

In nuclear reactions, where the nucleus of an atom undergoes a change, a small amount of mass can be converted into the energy according to Einstein’s famous equation, E=mc². so in these reactions , law is not followed(Exception)

2. The Law of Definite Proportions: An important law in Compound formation :-

The Law of Definite Proportions, also known as law of constant proportions (given by Joseph Proust), states that:

“A chemical compound always contains the same elements in a fixed /constant/ definite ratio by mass, regardless of its source or method of preparation”.

Essentially, a specific compound will always have the same proportion of its constituent elements, irrespective of how it’s formed or what is its source .

2.1 Examples of Law of Definite Proportions :-

here are some examples that clarify the law of definite proportions :-

2.1.1 Formation of Water :-

Take water (H₂O) as an example. No matter where water comes from – a lake, a river, or even produced through a chemical reaction – its composition will always remain constant that is two parts hydrogen by mass to one part oxygen by mass (That is 1:8).

This means all sources of water in the world will have same ratio of hydrogen to oxygen , which is 1:8 (Always a constant ratio)

2.1.2 Formation of CO₂ :-

In carbon dioxide , the mass ratio of carbon( 12 gm) to oxygen(32 gm) is always constant That is 12:32 or 3:8 , Whatever may be the source of CO₂ ( Be it evolved from volcano, or formed in chemical reaction or evolved by plants during respiration)

3. The Law of Multiple Proportions: A Crucial Law in compound formation :-

The Law of Multiple Proportions, states that:

“When two elements can combine to form two or more different compounds, the masses of one element that combine with a fixed mass of the other element will always be in a simple whole-number ratio.

In simpler terms, if elements A and B form multiple compounds, the ratio of the mass of element A that combines with a fixed mass of element B will always remain in ratio of small whole numbers.

3.1 Examples of Law of Multiple Proportions

Here are some examples to explore :-

3.1.1 Formation of CO and CO₂ :-

A classic example of this law is the Formation of carbon monoxide and carbon dioxide from carbon and oxygen . Carbon can combine with oxygen to form two different oxides: carbon monoxide (CO) and carbon dioxide (CO₂).

In CO, 12 gm of carbon combine with 16 gm of oxygen.(C:O=12:16)

In CO₂, 12 grams of carbon combine with 32 grams of oxygen.(C:O=12:32)

Here, the ratio of oxygen’s masses combining with a fixed mass of carbon (12 grams) is 16:32 , which simplifies to a 1:2 ratio, that is a simple whole number ratio.

3.2.2 Formation of H₂O and H₂O₂ :-

Another example of Law of Multiple Proportions is seen during the formation of Water (H₂O) and Hydrogen Peroxide (H₂O₂) from Hydrogen and oxygen. Hydrogen can combine with oxygen to give two different oxides: Water (H₂O) and Hydrogen Peroxide (H₂O₂).

In H₂O, 2 grams of Hydrogen combine with 16 grams of oxygen.(H:O=2:16)

In H₂O₂, 2 grams of Hydrogen combine with 32 grams of oxygen.(H:O=2:32)

Here, the ratio of oxygen’s masses combining with a fixed mass of Hydrogen (2 grams) is 16:32, which simplifies to a 1:2 ratio, again it is a simple whole number ratio.

3.2.3 Formation of Nitrogen oxides :-

Nitrogen and oxygen combine to form five oxides , which are as follow :-

(a) Nitrous oxide(N₂O) in which , N:O =28:16 = 14:8

(b) Nitric oxide (NO) in which , N:O =14:16

(d) Nitrogen Dioxide(NO₂) [or Dimer N₂O₄] in which , N:O =14:32

(e) Dinitrogen Penta oxide(N₂O5) in which N:O=28:80 = 14:40

In all five oxides of nitrogen , the mass of one element (Nitrogen) is fixed or constant that is 14 gram but the ratio of masses of oxygen in these compounds lie in the ratio 8:16:24:32:40 , which simplifies to 1:2:3:4:5 , again a simple whole number ratio or simple multiple ratio

4. Gay-Lussac’s Law of Combining Volumes :-

Gay-Lussac’s Law, also known as the Law of Combining Volumes, focuses on the volumes of gases involved in chemical reactions. The Law is as follow :

“At constant temperature and pressure, gases react with each other in simple whole number ratio by volume, simultaneously the volume of the gaseous products also follows a simple whole number ratio”.

In simpler terms, imagine a reaction between hydrogen gas (H₂) and oxygen gas (O₂) to form water vapour (H₂O).

2H₂ + O₂ → 2H₂O

According to Gay-Lussac’s Law, 2 volumes of hydrogen gas reacts with one volume of oxygen to produce 2 volumes of water vapor (H₂O).

This can be further simplified :- If 100 ml of hydrogen gas reacts with 50 ml of oxygen then they produce 100 ml of water vapor (H₂O). If we talk about volumes of reacting and producing gases as a whole then the ratio is 2:1:2 , which is a simple whole number ratio.

Avogadro’s Law :-

Avogadro’s Law builds upon Gay-Lussac’s findings and dives deeper into the connection between volume and the number of molecules. Here’s the concept:

Under identical(same) temperature and pressure conditions, equal volumes of all the gases contain the equal number of molecules.

At NTP , 1 mole of every gas has volume 22.4 Liters and number of molecules equals to Avogadro number (6.022 x 10²³ molecules)

Hence at NTP , I mole each of Nitrogen , Oxygen and Chlorine will have same number of molecules that is 6.022 x 10²³ molecules.

Conclusion :- Hence laws of chemical combination are obeyed in almost all chemical reactions except some exceptional reactions.