Chemistry Fundamentals: Concepts and Laws

Posted on Apr 11, 2025 in Chemistry

Unit 1: Concepts and Fundamental Laws of Chemistry

Pure Substance

A pure substance is a form of matter that has a constant composition and defined, distinctive properties. It cannot be separated into other substances without losing its properties. Examples include gold and water. Pure substances can be classified into elements and compounds.

Element

Elements are simple or pure substances that cannot be decomposed into simpler substances by chemical processes. Examples include gold, silver, and oxygen.

Compound

Compounds are pure substances that can be decomposed into simpler ones by chemical processes. For example, water can be decomposed by electrolysis into oxygen and hydrogen (simple substances or elements).

Physical Properties

Physical properties are those that can be observed without changing the composition of the material. Examples include hardness, solubility, and color. Physical properties can be further classified as:

  1. Extensive Properties: These properties depend on the amount of matter (mass dependent). For example, volume and kinetic energy.
  2. Intensive Properties: These properties do not depend on the amount of matter. Example: temperature and density.

Chemical Properties

Chemical properties describe how a substance changes its composition through a chemical reaction. An example is combustion or reaction with acids.

Chemical Reaction

A chemical reaction is the transformation of initial substances, called reagents, into products with different compositions and properties.

Ponderal Laws

Ponderal laws relate to mass:

  1. Lavoisier’s Law (Law of Conservation of Mass, 1789): In every chemical reaction, the total mass of the reacting substances is equal to the total mass of the products formed in the reaction.
  2. Proust’s Law (Law of Constant Proportions, 1799): Different samples of one pure compound always contain the same elements in a constant mass ratio, regardless of the process followed in its preparation.
  3. Dalton’s Law (Law of Multiple Proportions, 1802): The masses of the same element that combine with a fixed mass of another element to form a compound are in a ratio of whole numbers.
  4. Richter’s Law (Law of Equivalent Proportions, 1802): The masses of different elements that combine with a fixed mass of one element indicate the mass ratios in which they combine when they react, or simple multiples or sub-multiples of those ratios.

This law introduces the concept of equivalent mass or gram equivalent. The gram equivalent of an element is its mass that combines with 8.0 grams of oxygen or with 1.00 grams of hydrogen.

This law can also be stated as: when two elements combine, they do so in amounts equal to their equivalent masses or proportional to them.

Law of Combining Volumes (Gay-Lussac’s Law, 1808)

At constant temperature and pressure, the volumes of gases involved as reagents or products in a chemical reaction are in a ratio of whole numbers.

Ions

Ions are chemical species endowed with an electric charge, consisting of one or more identical or different atoms. Negatively charged ions are called anions, and positively charged ions are called cations.

Dalton’s Atomic Theory

Dalton theorized that matter is made of atoms that are indivisible and indestructible. These atoms are solid spheres (at the time, protons, electrons, and neutrons were unknown). He explained his theory of weightings but not the volumetric ones.

Avogadro’s Hypothesis

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

Avogadro’s Number

The number of molecules in one mole. Its value is approximately 6.022 x 1023.

Mole (Mol)

A mole is the chemical unit corresponding to Avogadro’s number of the substance in question. So, one mole of gold atoms would correspond to 6.022 x 1023 gold atoms. In the case of one mole of sulfuric acid molecules, it corresponds to 6.022 x 1023 sulfuric acid molecules. The mass corresponding to one mole of atoms or molecules is equal to the atomic or molecular mass expressed in grams.

By extension, one mole of ions, electrons, etc., is Avogadro’s number (6.022 x 1023) of ions, electrons, and so on.