Laws of Chemical Combination
There are certain rules or laws in chemistry that all chemical reactions follow when they occur. These laws are called the laws of chemical combination. These laws are exclusively for chemical reactions in which electrons of elements are involved. The nuclear reactions do not follow these laws as the nucleus is involved in the nuclear reactions.
There are mainly five laws of chemical combinations that are described as follows:
1) Law of Conservation of Mass or (law of indestructibility of Matter)
This law was given by Antoine Lavoisier in 1789 and tested by Landolt. It states that the mass can neither be created nor be destroyed in a chemical reaction. However, one type of mass can be converted into another type of mass when reactants are converted into products.
So, according to this law, the reactants when consumed completely or converted completely into products, the sum of masses of reactants will be equal to the sum of masses of products, which means the total mass of reactants = total mass of products. Due to this reason, we are required to balance a chemical equation.
If the reactants are not converted into products completely, Total mass of reactants = total mass of products + mass of unreacted reactants
For example, in the following chemical reaction, one molecule of hydrogen reacts with one molecule of chlorine to form 2 molecules of hydrogen chloride (HCl).
H2 + Cl2 → 2HCl
In the above chemical reaction, the atomic mass of hydrogen is 2 g as there are 2 atoms of hydrogen, and the atomic mass of chlorine is 71 as there are two atoms of chlorine. So, reactants' total mass is 2 + 71 = 73 g.
Now, in this reactions, 2 molecules of hydrogen chloride (HCl) are formed, whose total mass is 2 (1x 35.5), which is equal to the total mass of reactants. It shows this reaction follows the law of conservation of mass as mass is neither created nor destroyed, it is just converted from one form to another form. Besides this, in this reaction, the reactants are completely consumed to form the products, so reactants' total mass = products' total mass.
2) Law of Definite Proportion or Law of Constant Composition
This law was given by Joseph Proust in 1806. It states that the ratio of constituting elements of a compound remains fixed. It means compound always contains the same elements in a fixed proportion by mass irrespective of the method of its preparation or the source from which it is obtained.
For example, you can take water from any source or prepare it by any method, the ratio of hydrogen and oxygen atoms in a molecule of water by weight remains fixed or the same (H2O). For example, the atomic weight of hydrogen is always 2 gm and oxygen is 16 gm in the molecule of water. So, the ratio of hydrogen and oxygen is always 2: 16 or 1: 8 in water molecules. So, according to this law, for every 1 gm hydrogen there should be 8 gm oxygen to form water.
3) Law of Multiple Proportions
This law was given by John Dalton. According to this law, when two elements combine and form more than one compound then the different masses of one element that combine with the fixed mass of another element always bear a simple whole-number ratio. Let us take few examples to understand this law:
i) Hydrogen (H) and Oxygen (O)
Hydrogen and oxygen form two different compounds water (H2O) and hydrogen peroxide (H2O2). The mass of hydrogen remains the same (2 gm) in these two compounds. whereas, the mass of oxygen is different; in water, it is 16 gm and in hydrogen peroxide, it is 32. So, according to the law of multiple proportions, different masses of oxygen is combining with the fixed mass of hydrogen. So, the oxygen in these two compounds is supposed to bear a simple whole-number ratio as per this law. In H2O oxygen is 16 gm, whereas in H2O2 the oxygen is 32 gm. Now, we can see, the different masses of oxygen bears a simple whole-number ratio (16: 32 or 1:2) to one another.
ii) Carbon (C) and Oxygen (O)
Carbon and oxygen react with each other to form carbon dioxide (CO2) and carbon monoxide (CO). In this case, oxygen combines with the fixed mass of carbon to form two different compounds. Here, in CO2, the oxygen is 32 gm by mass and in CO the oxygen is 16 gm by mass. Now, the different masses of oxygen again bears a simple ratio to one another when it combines with a fixed mass of another element. The ratio of oxygen, in this case, is 32: 16 or 2: 1, which is a simple whole-number ratio.
iii) Nitrogen (N) and Oxygen (O)
Nitrogen combines with oxygen to form the following compounds.
If we take two moles of NO and NO2 in ii and iii compounds, the weight of N in both molecules becomes 28, however, the weight of oxygen becomes 32 and 64 respectively.
Now, we have a fixed weight of N (28 gm) in all molecules and O is reacting with a fixed weight of Nitrogen. Now according to this law, the ratio of different masses of oxygen to each other in these different molecules should be a simple whole-number ratio, which is 16: 32: 64: 48: 80 or 1: 2: 4: 3: 5.
4) Law of Reciprocal Proportions
It was first formulated by Jeremias Richter in 1791. It is also a basic law of stoichiometry as it deals with the proportions of elements that take part in a chemical reaction. According to this law, the ratio of masses X and Y elements that combine separately with a fixed mass of the third element Z will be either the same or some multiple of the ratio of masses of X and Y when they combine with each other.
Let us understand this law with the following example:
Suppose there are three elements P, Q, and R.
X gm of P combines with Y gm of Q to form PQ
Z gm of R combines with Y gm Q to form ZQ
Here, the ratio of the masses of P and R that combines with the fixed mass (Y gm) of Q = X: Z
Now, U gm of P combines with V gm of R, to form PR, so their ratio when they combine with each other = U: V
Now according to this law, the ratio of P and R (X: Z) when they combine with a fixed mass of the third element is equal to or multiple of their ratio (U: V) when they combine with each other.
So, either X: Z = U: V or X: Z = n times of U: V where, n can be 1, 2, 3, etc.
Let us take three elements hydrogen (H), oxygen (O), and sulphur (S) to understand it further:
Hydrogen forms water (H2O) with Oxygen and forms hydrogen sulphide (H2S) with sulphur. Sulphur forms sulphur dioxide (SO2) on reacting with oxygen. Now, in this case, sulphur and oxygen are combining with a fixed mass of hydrogen (2 gm) and they are also reacting with each other to produce SO2.
Now the ratio of masses of S: O when combine differently is 32: 16 = 2: 1
Now, the ratio of S: O when they combine with each other: 32: 32 = 1: 1
Now as per the law the ratio 2:1 should be equal to 1:1 or multiple of 1:1.
In this case it is a multiple of 1:1 as 2:1 = 1:1 x 2
5) Law of Gaseous Volume
This law was given by Gay Lussac in 1808. According to this law, in the gaseous reaction, the gases always combine in a simple ratio by volume and form products that also bear a simple ratio with reactants provided all gases are reacting with each other at the same temperature and pressure.
Let us understand it by a simple example:
In the following example, one volume of hydrogen gas combines with one volume of chlorine gas to form 2 volumes of hydrogen chloride gas.
H2 (gas) + Cl2 (gas) → 2HCl (g)
So, as per the law of gaseous volume, the ratio of their volumes that are combining would be simple a simple ratio, in this case, the ratio is 1: 1: 2.
Let us take another example to understand it:
N2 + 3H2 → 2NH3
Here, one volume of nitrogen gas combines with 3 volumes of hydrogen gas to form 2 volumes of ammonia.
Now, in this case, also, the ratios of volumes of reactants and product is a simple ratio (1: 3: 2), which follows the law of gaseous volume.