Some Basic Concepts of Chemistry

Matter is anything that occupies space and has mass. It can be classified into two main ways: physical and chemical.

Physical Classification:

• Solids: Definite volume and definite shape. Particles are held very close to each other.
• Liquids: Definite volume but no definite shape. Particles are close but can move.
• Gases: Neither definite volume nor definite shape. Particles are far apart and move freely.

Chemical Classification:

Key Definitions:

• Atom: Smallest unit of an element that may or may not exist independently.
• Molecule: Smallest unit of a substance (element or compound) that exists independently.

• Law of Conservation of Mass (Lavoisier, 1789): Total mass of reactants = Total mass of products. (Exceptions: Nuclear reactions).
• Law of Definite Proportions (Proust, 1799): A compound always contains the same elements in a fixed ratio by mass, regardless of source.
• Law of Multiple Proportions (Dalton, 1803): If two elements form >1 compound, the masses of one element combined with a fixed mass of other are in a simple whole-number ratio. e.g., NO (14:16) and NO₂ (14:32). Ratio = 16:32 = 1:2.
• Gay-Lussac’s Law of Gaseous Volumes (1808): When gases combine at constant T and P, they do so in volumes which bear a simple ratio to one another and to the volume of gaseous products.
• Avogadro’s Law (1811): Equal volumes of all gases at the same temperature and pressure should contain equal number of molecules.

• Average Atomic Mass: Sum of (Isotopic Mass × Fractional Abundance).
• Molecular Mass: Sum of atomic masses of the elements present in a molecule.
• Formula Mass: Used for ionic compounds (like NaCl) where discrete molecules don't exist.

Empirical & Molecular Formula

• Empirical Formula: Simplest whole number ratio of atoms. e.g., CH₂O.
• Molecular Formula: Actual number of atoms. e.g., C₆H₁₂O₆.
• Relation: Molecular Formula = n × Empirical Formula (where n = Molar Mass / EF Mass).

Balancing Equations:

Essential for conservation of mass. Example: CH₄ + 2O₂ → CO₂ + 2H₂O.

Stoichiometric Calculations:

• Mole-Mole: 1 mole of CH₄ gives 1 mole of CO₂.
• Mass-Mass: 16g of CH₄ reacts with 64g (2×32) of O₂.
• Mass-Volume: At STP, 16g of CH₄ gives 22.4L of CO₂.

Significant Figures & Precision:

• Precision: Closeness of various measurements for the same quantity.
• Accuracy: Agreement of a particular value to the true value.
• Rule for + / -: Result should have same decimal places as the term with least decimal places.

9. Common Mistakes

• Not Converting to Moles: Attempting to use mass directly in stoichiometric ratios. Always convert grams to moles first!
• M vs m Confusion: Forgetting that Molarity (M) changes with temperature whereas Molality (m) does not.
• SF in Calculations: Rounding off too early in multi-step calculations. Keep extra digits until the final result.
• Solvent vs Solution: Using solution mass instead of solvent mass in molality calculations.

10. Example: Limiting Reagent

Problem: 50kg of N₂ and 10kg of H₂ are mixed to produce NH₃. Identify LR.

Solution:
- Moles of N₂ = 50000 / 28 = 1785.7 mol
- Moles of H₂ = 10000 / 2 = 5000 mol
- Stoic. Req (N₂ + 3H₂ → 2NH₃):
- N₂ index: 1785.7 / 1 = 1785.7
- H₂ index: 5000 / 3 = 1666.7
LR is H₂ (smaller index).

11. IAT Exam Focus Points

High Yield Areas:

• Concentration Conversions: Problems where density is given to convert M to m.
• Empirical Formula: Finding MF which is an 'n' multiple of EF.
• Limiting Reagent: Always the core of IAT stoichiometry questions.
• Atomic Mass Unit: Questions on the definition of 1 amu (based on C-12).

8. Practice Mock Test

Ready to test your knowledge?

Take a quick 15-question assessment specifically designed for Basic Concepts of Chemistry. Challenge yourself with IAT-level questions.

Start Practice Mock