Introduction
The chapter "Atoms and Molecules" explains the fundamental concepts of chemistry, such as atoms, molecules, laws of chemical combination, and the concept of atomic and molecular mass.
1. Laws of Chemical Combination
The foundation of chemical reactions is based on certain laws, two of which are explained in this chapter:
Law of Conservation of Mass:
- This law states that mass can neither be created nor destroyed in a chemical reaction.
- The total mass of reactants is equal to the total mass of products in a chemical reaction.
- Example: When hydrogen and oxygen combine to form water, the total mass of hydrogen and oxygen before the reaction is equal to the mass of water produced.
Law of Constant Proportion (Definite Proportions):
- This law states that a chemical compound always contains the same elements in the same proportion by mass, irrespective of the source or method of preparation.
- Example: Water (H₂O) always contains hydrogen and oxygen in a 1:8 mass ratio.
2. Dalton’s Atomic Theory
John Dalton proposed a theory in 1808 based on the laws of chemical combination. The main postulates of Dalton’s atomic theory are:
- Matter is made up of indivisible particles called atoms.
- Atoms of the same element are identical in mass and chemical properties.
- Atoms of different elements have different masses and chemical properties.
- Atoms combine in simple whole-number ratios to form compounds.
- Atoms cannot be created or destroyed in a chemical reaction, only rearranged.
Although this theory was revolutionary, some modifications were made later as atoms were found to be divisible.
3. What is an Atom?
- Atoms are the basic units of matter.
- They are extremely small particles that cannot be seen with the naked eye.
- Atoms of different elements combine in fixed proportions to form compounds.
4. Atomic Mass
- Atomic mass is the mass of an atom, measured in atomic mass units (amu).
- The relative atomic mass of an element is based on the mass of a carbon-12 isotope, where 1 amu = 1/12th the mass of a carbon-12 atom.
- Example: Hydrogen has an atomic mass of 1 amu, and oxygen has an atomic mass of 16 amu.
5. Molecules
- Molecules are formed when two or more atoms combine chemically. They can be:
- Molecules of elements: Comprised of the same type of atoms (e.g., O₂, H₂).
- Molecules of compounds: Comprised of different types of atoms (e.g., H₂O, CO₂).
6. Ions
- Ions are charged particles that form when atoms gain or lose electrons.
- Cations are positively charged ions (formed by loss of electrons), and anions are negatively charged ions (formed by gain of electrons).
- Example: Sodium (Na) becomes Na⁺ after losing one electron, while Chlorine (Cl) becomes Cl⁻ after gaining one electron.
7. Writing Chemical Formulae
The chemical formula of a compound is a symbolic representation of its composition. To write the chemical formula:
- Identify the symbols of the elements.
- Write the valency (combining capacity) of each element.
- Cross-multiply the valencies to get the simplest ratio.
For example, the formula for water is H₂O because hydrogen has a valency of 1 and oxygen has a valency of 2, so two hydrogen atoms combine with one oxygen atom.
8. Molecular Mass and Formula Unit Mass
- Molecular mass is the sum of the atomic masses of all the atoms in a molecule.
- Example: The molecular mass of water (H₂O) is 18 amu (2 × 1 + 16).
- Formula unit mass applies to compounds with ionic bonds. It is the sum of the atomic masses of ions in the compound.
- Example: The formula unit mass of sodium chloride (NaCl) is 58.5 amu (23 + 35.5).
9. Mole Concept
- A mole is a unit used to measure the amount of substance. One mole contains 6.022 × 10²³ particles (Avogadro’s number).
- The mass of one mole of a substance is equal to its atomic or molecular mass expressed in grams.
- Example: 1 mole of water weighs 18 grams.
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