Nature of Matter, States and Scientific Classification
Chemistry is the science of molecules and their transformations. Everything around us that has mass and occupies space is termed Matter.
Matter is classified physically into Solids, Liquids, and Gases, and chemically into:
- Pure Substances:
- Elements: Consist of only one type of atom (e.g., Na, Cu, H₂, O₂).
- Compounds: Consist of two or more atoms of different elements combined in a fixed ratio (e.g., H₂O, CO₂, NH₃).
- Mixtures:
- Homogeneous Mixtures: Uniform composition throughout (e.g., air, sugar in water).
- Heterogeneous Mixtures: Non-uniform composition (e.g., sand and salt, oil in water).
After studying this chapter, proceed to Class 11 Chemistry Chapter 2 Notes (Structure of Atom) or test your aggregate scores on our CBSE percentage calculator.
Uncertainty in Measurement: Scientific Notation, Significant Figures and Dimensional Analysis
In chemistry, measurements involve very large or extremely small numbers (such as Avogadro's number 6.022 × 1023 or electron mass 9.11 × 10-31 kg).
Rules for Determining Significant Figures:
- All non-zero digits are significant (e.g.,
285 cmhas 3 significant figures). - Zeros preceding the first non-zero digit are not significant (e.g.,
0.03 ghas 1 significant figure). - Zeros between two non-zero digits are significant (e.g.,
2.005 ghas 4 significant figures). - Zeros at the end or right of a number are significant if they are on the right of the decimal point (e.g.,
0.200 ghas 3 significant figures). - Exact numbers have an infinite number of significant figures (e.g., 20 apples = 20.000…).
The 5 Fundamental Laws of Chemical Combinations
Chemical transformations obey five fundamental empirical laws:
- Law of Conservation of Mass (Antoine Lavoisier, 1789): In any physical or chemical change, the total mass of reactants equals the total mass of products (Mreactants = Mproducts).
- Law of Definite Proportions (Joseph Proust, 1799): A given chemical compound always contains exactly the same proportion of elements by weight, regardless of its source or method of preparation.
- Law of Multiple Proportions (John Dalton, 1803): If two elements combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element are in the ratio of small whole numbers (e.g., CO and CO₂ where Oxygen ratio is 16:32 = 1:2).
- Gay-Lussac's Law of Gaseous Volumes (1808): When gases combine or are produced in a chemical reaction, they do so in a simple ratio by volume, provided all gases are at the same temperature and pressure (2H2(g) + O2(g) → 2H2O(g) with volume ratio 2:1:2).
- Avogadro's Law (Amedeo Avogadro, 1811): Equal volumes of all gases at the same temperature and pressure contain equal numbers of molecules (V ∝ n).
Dalton's Atomic Theory and Modern Atomic Postulates
In 1808, John Dalton proposed his revolutionary atomic theory:
- Matter consists of indivisible atoms.
- All atoms of a given element have identical properties and mass.
- Compounds are formed when atoms of different elements combine in a fixed ratio.
- Chemical reactions involve reorganization of atoms; atoms are neither created nor destroyed.
Atomic, Molecular and Formula Masses (Unified Mass 'u')
The atomic mass unit (amu or unified mass 'u') is defined as exactly 1/12th of the mass of one carbon-12 atom:
Average Atomic Mass: For elements possessing isotopes, the average mass is:
The Mole Concept, Avogadro's Number and Molar Volume at STP
The Mole is the SI unit for the amount of substance:
Where Avogadro Constant N_A = 6.02214076 × 10²³ entities/mol.
Percentage Composition, Empirical Formula and Molecular Formula
Mass Percentage of an Element:
Empirical vs Molecular Formula:
- Empirical Formula: The simplest whole-number ratio of various atoms present in a compound (e.g., CH for Benzene).
- Molecular Formula: The exact number of different types of atoms present in a molecule (C₆H₆).
- Relationship: Molecular Formula = n × (Empirical Formula), where
n = Molecular Mass / Empirical Formula Mass.
Chemical Stoichiometry and Limiting Reagent Algorithms
Stoichiometry deals with the quantitative relationships among reactants and products in a balanced chemical reaction.
Limiting Reagent Algorithm:
- Convert all given masses of reactants into moles:
n_A = m_A / M_Aandn_B = m_B / M_B. - Divide each mole value by its stoichiometric coefficient: Ratio
R_A = n_A / aandR_B = n_B / b. - The reactant with the smaller ratio is the Limiting Reagent (LR).
- Use the moles of the LR to calculate the theoretical yield of products.
Expressing Concentration of Solutions: Mass %, Mole Fraction, Molarity and Molality
| Concentration Term | Mathematical Formula | SI Units | Temperature Dependence |
|---|---|---|---|
| Mass Percentage (% w/w) | (Mass of Solute / Mass of Solution) × 100 | Unitless (%) | Temperature Independent |
| Mole Fraction (X_A) | n_A / (n_A + n_B) where X_A + X_B = 1 | Unitless | Temperature Independent |
| Molarity (M) | Moles of Solute / Volume of Solution in Litres | mol/L or M | Temperature Dependent (Expands with heat) |
| Molality (m) | Moles of Solute / Mass of Solvent in kg | mol/kg or m | Temperature Independent |
Molarity Dilution Equation: M₁ · V₁ = M₂ · V₂.
Step-by-Step Solved Exemplar Numericals for Board and Competitive Exams
Problem 1: Empirical Formula Determination
A compound contains 4.07% Hydrogen, 24.27% Carbon, and 71.65% Chlorine. Its molar mass is 98.96 g/mol. Determine its empirical and molecular formulas.
2. Moles of H = 4.07 / 1.008 = 4.038
3. Moles of Cl = 71.65 / 35.45 = 2.021
4. Atomic Ratio: C : H : Cl = (2.021/2.021) : (4.038/2.021) : (2.021/2.021) = 1 : 2 : 1
5. Empirical Formula: CH₂Cl (Empirical Mass = 12 + 2 + 35.5 = 49.5 g/mol)
6. n = 98.96 / 49.5 = 2
7. Molecular Formula: (CH₂Cl)₂ = C₂H₄Cl₂ (1,2-Dichloroethane).
Problem 2: Limiting Reagent in Ammonia Synthesis
50.0 kg of N₂(g) and 10.0 kg of H₂(g) are mixed to produce NH₃(g). Identify the limiting reagent and calculate the mass of NH₃ formed.
Moles of N₂ = 50,000 g / 28.0 g/mol = 1785.7 mol
Moles of H₂ = 10,000 g / 2.016 g/mol = 4960.3 mol
Ratio for N₂ = 1785.7 / 1 = 1785.7
Ratio for H₂ = 4960.3 / 3 = 1653.4
Since 1653.4 < 1785.7, H₂ is the Limiting Reagent.
Moles of NH₃ produced = (2/3) × 4960.3 = 3306.9 mol
Mass of NH₃ = 3306.9 mol × 17.03 g/mol = 56,316 g = 56.32 kg.
Frequently Asked Questions (FAQs)
One mole is the amount of substance that contains exactly 6.02214076 × 10²³ elementary entities (atoms, molecules, ions, or electrons). This number is known as the Avogadro Constant (N_A).
Molarity (M) is the number of moles of solute dissolved in 1 Litre of solution (M = n / V_L) and depends on temperature because liquids expand with heat. Molality (m) is the number of moles of solute dissolved in 1 kilogram of solvent (m = n / W_kg) and is temperature-independent.
The limiting reagent is the reactant that is completely consumed first in a chemical reaction. It limits the maximum amount of product that can be formed.
Under modern IUPAC standard temperature and pressure (STP: 273.15 K and 1 bar), 1 mole of an ideal gas occupies exactly 22.7 Litres. Under older standard conditions (273.15 K and 1 atm), it occupies 22.4 Litres.
Molecular Formula = n × (Empirical Formula), where n = (Molecular Mass) / (Empirical Formula Mass).
Significant figures are the meaningful digits in a measured quantity known with certainty plus one estimated digit. All non-zero digits are significant; zeros between non-zero digits are significant; leading zeros are never significant; and trailing zeros after a decimal point are significant.