Atoms and Molecules

1. Law of Conservation of Mass

Proposed by Antoine Lavoisier in 1789, this law states:

Classic experiment: Lavoisier burned mercury in a closed vessel and showed that the mass of mercury calx (HgO) formed equalled the mass of mercury plus the mass of oxygen consumed -- nothing was lost or gained.

Example: When 16 g of methane (CH4) burns with 64 g of oxygen (O2), it produces 44 g of carbon dioxide (CO2) and 36 g of water (H2O).

The masses balance perfectly -- no matter is created or destroyed.

NCERT activity: If 5.3 g of sodium carbonate (Na2CO3) reacts with 6 g of ethanoic acid (CH3COOH), and the products are 2.2 g CO2 + 8.2 g sodium ethanoate + 0.9 g H2O, verify conservation: 5.3 + 6 = 11.3 g on the left; 2.2 + 8.2 + 0.9 = 11.3 g on the right. Mass is conserved.

2. Law of Constant (Definite) Proportions

Stated by Joseph Proust in 1799:

Example -- Water (H2O): Whether water is taken from a river, the sea, rain, or made in a lab, the mass ratio of hydrogen to oxygen is always 1 : 8. That is, 18 g of water always contains 2 g of H and 16 g of O.

Example -- Carbon dioxide (CO2): Whether CO2 comes from burning carbon, respiration, or decomposition of limestone, the ratio C : O by mass is always 3 : 8.

Why this matters: This law shows chemical compounds are not random mixtures -- they are fixed, definite entities. Mixtures (like air or alloys) do NOT obey this law; compounds (like NaCl, H2O, CO2) always do.

Remember the difference: Conservation of mass deals with the total mass before and after a reaction; constant proportions deals with the ratio in which elements combine.

3. Dalton's Atomic Theory -- Postulates

In 1808, John Dalton published A New System of Chemical Philosophy and gave a theoretical basis for both chemical laws. His postulates are:

1. All matter is made of extremely small, indivisible particles called atoms.
2. Atoms of the same element are identical in all respects -- mass, size, and chemical properties. Atoms of different elements differ in these properties.
3. Atoms can neither be created nor destroyed in a chemical reaction. This explains the Law of Conservation of Mass.
4. Atoms of different elements combine in simple whole-number ratios to form compounds. This explains the Law of Constant Proportions.
5. The relative number and kinds of atoms in a given compound are always the same.

Successes: Dalton's theory explained both laws and provided the first quantitative framework for chemistry -- for the first time, chemists could think about reactions in terms of atoms rearranging.

Limitations (discovered later):

• Atoms are NOT indivisible -- they contain subatomic particles: electrons, protons, neutrons.
• Atoms of the same element can differ in mass -- these are called isotopes (e.g. C-12 and C-14).
• Atoms of different elements can have the same mass -- these are called isobars (e.g. Ar-40 and Ca-40).
• The ratio of atoms in some compounds is not a simple whole number (e.g., sucrose is C12H22O11).

4. Atoms -- Symbol, Atomic Mass, and Atomic Mass Unit

An atom is the smallest particle of an element that takes part in a chemical reaction. Atoms are extremely small -- a hydrogen atom has a radius of about 1 x 10^-10 m (0.1 nm). To get an idea: if you magnified an apple to the size of the Earth, each atom in the apple would be roughly the size of the original apple.

Chemical symbols are short-hand representations of elements, first proposed by Berzelius. Rules:

• First letter of the English or Latin name, written in capital (e.g. O for Oxygen, N for Nitrogen, C for Carbon, S for Sulphur).
• When two elements start with the same letter, a second letter (in lowercase) is added (e.g. Ca for Calcium, Cl for Chlorine, Co for Cobalt).
• Some symbols come from Latin names: Na (Natrium = Sodium), K (Kalium = Potassium), Fe (Ferrum = Iron), Ag (Argentum = Silver), Au (Aurum = Gold), Pb (Plumbum = Lead), Hg (Hydrargyrum = Mercury), Cu (Cuprum = Copper).

Atomic mass unit (u): Since actual atomic masses are incredibly small (H atom weighs 1.673 x 10^-24 g), chemists use a relative scale. The atomic mass unit (u) is defined as:

Relative atomic mass = (mass of 1 atom of the element) / (1/12 x mass of 1 C-12 atom). This gives a pure number (no units), but by convention we write it as "u".

Key atomic masses to memorise (from NCERT Table 3.1):

5. Molecules -- Elements and Compounds

A molecule is the smallest particle of a substance that can exist independently and shows all the properties of that substance. Molecules are formed by the combination of two or more atoms held together by chemical bonds.

Atomicity = number of atoms present in one molecule of an element.

Molecules of elements -- atoms of the SAME element join together:

Molecules of compounds -- atoms of DIFFERENT elements join together in fixed ratios:

• Water (H2O): 2 hydrogen atoms + 1 oxygen atom
• Ammonia (NH3): 1 nitrogen atom + 3 hydrogen atoms
• Carbon dioxide (CO2): 1 carbon atom + 2 oxygen atoms
• Hydrochloric acid (HCl): 1 hydrogen atom + 1 chlorine atom
• Glucose (C6H12O6): 6 carbon + 12 hydrogen + 6 oxygen atoms
• Ethanol (C2H5OH): 2 carbon + 6 hydrogen + 1 oxygen atom

Key point: Not all substances exist as molecules. Ionic compounds like NaCl, CaCO3 exist as giant lattices of ions -- we represent them by their formula units, not molecules.

6. Ions -- Cation and Anion

An ion is a charged species (atom or group of atoms) formed when an atom loses or gains electrons.

• Cation = positively charged ion, formed by loss of electrons. Examples: Na+ (loses 1e-), Ca2+ (loses 2e-), Al3+ (loses 3e-).
• Anion = negatively charged ion, formed by gain of electrons. Examples: Cl- (gains 1e-), O2- (gains 2e-), N3- (gains 3e-).

A polyatomic ion is a group of atoms that carries a net charge and behaves as a single unit. Examples: OH- (hydroxide), SO4^2- (sulphate), NO3- (nitrate), CO3^2- (carbonate), NH4+ (ammonium), PO4^3- (phosphate).

Common ions (NCERT Tables 3.3 and 3.4):

7. Writing Chemical Formulae Using Valency

The valency of an element or radical is its combining capacity. For ions, valency = magnitude of the charge. To write a chemical formula, use the criss-cross (swap) method:

1. Write symbols side by side: cation first, anion second.
2. Write their valencies as superscripts above each symbol.
3. Swap the valencies -- the cation's valency becomes the anion's subscript and vice versa.
4. Reduce to simplest whole-number ratio if possible (e.g. Ca2O2 becomes CaO).
5. Enclose polyatomic ions in brackets if their subscript is greater than 1.

Worked examples:

Naming when multiple valencies: Fe2+ = Iron(II) or Ferrous; Fe3+ = Iron(III) or Ferric; Cu+ = Copper(I) or Cuprous; Cu2+ = Copper(II) or Cupric.

8. Molecular Mass Calculations

Molecular mass of a compound = sum of the atomic masses of all atoms present in one molecule, expressed in atomic mass units (u).

9. Formula Unit Mass

Ionic compounds like NaCl and MgO do NOT exist as individual molecules -- they exist as large three-dimensional lattice structures of ions. For these, we calculate formula unit mass (not molecular mass), which is the sum of atomic masses of all atoms in the simplest formula unit.

Rule of thumb: Use "molecular mass" for covalent compounds; use "formula unit mass" for ionic compounds. The calculation method is identical.

10. The Mole Concept and Avogadro's Number

Atoms and molecules are so tiny that even a pinhead contains billions of atoms. We need a counting unit for such vast numbers. Just as we use "dozen" for 12 items, chemists use the mole for 6.022 x 10^23 items.

The mole is defined so that the mass of 1 mole of a substance in grams equals the atomic/molecular mass in u:

• Atomic mass of C = 12 u, so 1 mole of C atoms = 12 g of carbon (gram atomic mass).
• Molecular mass of H2O = 18 u, so 1 mole of H2O = 18 g of water (gram molecular mass).
• Atomic mass of Na = 23 u, so 1 mole of Na atoms = 23 g of sodium.
• Molecular mass of O2 = 32 u, so 1 mole of O2 = 32 g of oxygen.

Physical intuition of N_A = 6.022 x 10^23: If you were to count atoms at 10 million atoms per second, it would take about 2 billion years to count just one mole. That is how large Avogadro's number is!

Important relationships:

11. Mole-Mass-Number Conversions

Three quantities are interlinked. Master these formulas -- the mole is the bridge:

Where: M = gram atomic mass (for atoms) OR gram molecular mass (for molecules); N_A = 6.022 x 10^23.

12. Common Mistakes to Avoid

• Atomic mass vs molar mass: Atomic mass of H = 1 u (for one atom); molar mass of H = 1 g/mol (for 6.022 x 10^23 atoms). Numerically the same -- units differ.
• Wrong term for ionic compounds: Never say "molecular mass" of NaCl -- always say formula unit mass.
• Missing atoms in polyatomic groups: In Ca(OH)2, there are 2 O and 2 H atoms (not 1 each). Expand brackets first!
• Avogadro's number: The correct value is 6.022 x 10^23, NOT 6.023 x 10^23.
• Criss-cross errors: MgCl2 not MgCl; Al2O3 not AlO; write brackets for polyatomic ions when subscript is > 1.
• Cation vs anion: Cation is positive (less electrons), anion is negative (more electrons). Cation written first in formula.
• Law mix-up: Conservation of mass = total mass unchanged; constant proportions = ratio of elements in a compound is fixed.
• Dalton's theory -- atomicity of elements: Do not write H as monoatomic in reactions -- hydrogen exists as H2 under normal conditions.