Understanding the Laws of Chemical Reactions: A Comprehensive Review

Posted on Mar 14, 2025 in Chemistry

Laws Governing Chemical Reactions

Laws are pondered and are based on experiments where quantities are measured in the masses of the substances involved in chemical reactions. The systematic use of the balance allowed us to obtain quantitative data, providing an overview of chemical reactions that led to the laws of weight and the birth of chemistry as a science.

A) Lavoisier’s Law of Conservation of Mass

In a chemical reaction, the sum of the masses of the reactants equals the sum of the masses of the products.

Etching + Soda -> Salt + Water

HCl + NaOH -> NaCl + H2O

36.5 g + 40 g -> 58.5 g + 18 g = 76.5 g -> 76.5 g

Another more general statement of the law of conservation of mass is: all mass-energy in a process is neither created nor destroyed, only transformed. The inclusion of energy in the law of conservation is a consequence of the theory of relativity, which relates mass and energy by the equation: E = mc2 -> m = E / c2. Where: E -> is energy, m -> is the mass, and c -> is the speed of light and equals: C = 300,000 km / s. In chemical reactions, the transformation of mass into energy is negligibly small; therefore, we can use the first statement of the law.

B) Law of Definite Proportions (Proust’s Law)

The elements that combine to form a compound always do so in the same proportion by mass, regardless of how the compound is obtained or from where it originates. For example, potassium chloride (KCl) always satisfies the following relation: m(K) / m(Cl) = 20 / 18.205. This means that if 20 g of potassium come into contact with 20 g of chlorine, only 18.205 g of chlorine will react with the 20 g of potassium, leaving the rest of the chlorine unreacted.

Proust’s Law is not always exactly satisfied for two reasons:

  • The average atomic mass depends on the isotopic composition of the element (and this may vary according to its origin).
  • Some solid ionic compounds do not comply with Proust’s Law due to defects present in the crystal lattice (such as vacancies or fractures). These are called non-stoichiometric or berthollide compounds.

Atomic Hypothesis of Dalton

Dalton tried to explain Lavoisier’s and Proust’s laws using a theory, currently called Dalton’s atomic theory, whose most important principles are:

  1. Chemical elements are formed by atoms. Dalton considered atoms to be indivisible and very small, indestructible particles.
  2. Atoms of the same element are all identical and have the same properties.
  3. Chemical reactions involve the union or separation of whole atoms.
  4. Two or more atoms can join to form molecules. Molecules are the smallest particles that form a compound.


C) Dalton’s Law of Multiple Proportions

If a mass mA of element A combines with a mass mB of element B to form one compound, and another mass m’A of element A combines with the same mass mB of element B to form another compound, then mA and m’A are in a simple whole number ratio.

D) Richter’s Law or Law of Reciprocal Proportions

The masses of two different elements, mA and mB, that combine with the same mass mC of another element, are the masses relative to each other with which they combine, or multiples or fractions of those masses.

Volumetric Laws

Gay-Lussac’s Law of Combining Volumes

This law concerns the volumes of substances and applies only to gases. It states: “Volumes of gaseous substances involved in a chemical reaction (measured under the same conditions of pressure and temperature) are in a simple whole number relationship.” Gay-Lussac also observed that the sum of the volumes of the reactants is always greater than the sum of the volumes of the products.

Avogadro’s Hypothesis

Equal volumes of different gases, under the same conditions of pressure and temperature, contain the same number of molecules.

The Mole, Avogadro’s Number, and Atomic/Molecular Masses

Avogadro’s number is a constant with the value: n = 6.022 x 1023.

The mole is the SI unit of quantity and is defined as follows: one mole is the amount of substance that contains as many elementary entities as there are atoms in 12 g of the carbon-12 isotope. Precisely, 12 g of the carbon-12 isotope contains 6.022 x 1023 atoms.

The mass of a single atom is extremely small. For example, the mass of a carbon-12 atom is 1.99 x 10-23 g. To measure atomic masses, it is necessary to use a unit of measurement.

The mass of atoms is measured using the atomic mass unit (amu). One amu is equal to one-twelfth of the mass of a carbon-12 atom: