Laws of Chemical Combination

Introduction 

Chemical reaction

Let’s look at how solid calcium interacts with hydrobromic acid :

Ca + HBr → CaBr2 + H2

On the left side of the reaction, there is just one bromine and one hydrogen atom, but there are two atoms of each on the product’s side of the process. This is impossible because one atom of hydrogen and bromine cannot be created out of thin air. To demonstrate what occurs, we must balance the equation using coefficients, which are numbers placed in front of the compounds or components in the reaction. Let’s put this equation back together:

Ca + 2HBr → CaBr2 + H2

The coefficient in front of HBr balances the equation for mass. 

The chemical combinations of elements are governed by the following five main rules of chemical mixtures.

• Law of conservation of mass
• Law of definite proportions 
• Law of multiple proportions 
• Gay Lussac’s law of gaseous volumes 
• Avogadro’s law of chemical combination

1) The law of conservation of mass

In 1789, French scientist Antoine Lavoisier researched and studied this law. This rule states that “the total mass of the reactants is equal to that of the products in all physical and chemical reactions” or “mass can neither be created nor destroyed.”

This law is also known as the law of matter indestructibility. Although mass and energy are interconvertible, the total mass and energy stay constant during any physical or chemical change. 

As a result, the mass of the reactants is equal to the mass of the product created in every chemical reaction or chemical change. Dalton’s atomic theory helps explain this rule. Dalton’s atomic theory states that atoms are indivisible particles that cannot be formed or destroyed in a chemical process. 

By the law of conservation of mass,

The total mass of reactants = Total mass of products

In water formation, for example – 2H2 + O2 →2H2O 

Reactant mass = 4 + 32 = 36g Product mass = 2(2+16) = 2(18) = 36g 

As we can see, the mass of the reactants and the mass of the product are identical in the reaction. As a result, it demonstrates the law of mass conservation.

2. Law of definite/constant proportions

In 1799, Joseph Proust postulated a rule of fixed proportions. It is often referred to as the law of constant proportions. According to this law, the constituents of a chemical compound are always present in defined amounts by mass. 

John Dalton’s hypothesis also explained the law of constant proportions. According to John Dalton’s theory, the relative number and types of atoms in a particular molecule are constant. This statement supports the law of constant proportions. For example, in a water molecule, the mass ratio of hydrogen to a mass of oxygen is always the same, which is 1:8. This is because water molecules can come from everywhere, but the mass ratio of hydrogen to oxygen in a water molecule will always be 1:8.

The limitations

It is inapplicable if an element exists in several isotopes involved in the compound’s formation. Although the details may combine in the same proportion, the compounds generated may differ.

3. Gay Lussac’s law of gaseous volumes  

In 1808 Joseph Louis Gay-Lussac developed the Law of Gaseous Volumes. According to this rule, when measured at the same temperature and pressure, the ratio of the volumes of reacting gases are small whole numbers. This might be considered a variant of definite proportions because this law applies to volume, whereas the law of definite proportion applies to mass.

Example: The chemical interaction between hydrogen gas and chlorine gas, for example, is regulated by Gay Lussac’s law of gaseous volumes.

H2 + Cl2 → 2HCl

In this example, one volume of hydrogen gas and one volume of chlorine gas combine to form two volumes of HCl gas. Therefore, the volume ratios of hydrogen gas, chlorine gas, and HCl are 1:1:2.

4. Law of multiple proportions

In 1804, John Dalton introduced the law of multiple proportions. When the  two elements combine to make more than one compound, then the masses of one of the elements that combine with the  fixed mass of the other  exist in the ratio of smaller whole numbers, according to the law of multiple proportions.

Example: Consider the carbon and oxygen atoms. When carbon and oxygen atoms mix, they produce two different products: carbon monoxide and carbon dioxide. The quantity of oxygen in carbon monoxide for one carbon atom, i.e., 12.0 g, is 16.0 g. The amount of oxygen in carbon dioxide for one carbon atom 12.0 g is 32.0 g. Therefore, the carbon content of carbon monoxide is 32:16, or 2:1, to carbon in carbon dioxide.

5. Avogadro’s chemical combination law

In 1811, Avogadro proposed the fifth law of chemical combination.

At a fixed temperature and pressure, the volume of gaseous molecules is precisely proportional to the total number of moles, according to Avogadro’s equation of chemical combination. The following is an expression of Avogadro’s law.

V ∝ n

V denotes the volume of gaseous molecules, and n represents the total number of moles in the gas.

V = k × n

In this equation, k denotes the proportionality constant.

According to Avogadro’s law of chemical combination,

Vₗnₗ = V₂n₂

V1  denotes the volume, and n1 represents the number of moles of a single gas. And V2 is the volume of another gas, and n2 is the number of moles.

Blowing helium gas into a balloon is one example. The overall volume of the balloon grows as the total number of moles of helium gas blown into it increases.

Conclusion

These laws of chemical combination define the fundamental principles involved in the interaction of atoms and molecules and majorly interactions that can include a wide range of combinations that occur in a variety of ways. This incredible diversity of interactions allows for an astonishing range of chemical reactions and molecules.