Some Basic Concepts of Chemistry

Introduction

Chemistry is the branch of science that deals with the study of matter—its composition, properties, and the changes it undergoes. It forms the foundation for understanding the structure of atoms and molecules, the laws governing their combination, and quantitative relationships in chemical reactions.

This chapter introduces basic concepts essential for studying chemistry at an advanced level.

1. Importance of Chemistry

Chemistry plays a vital role in our daily lives and in various scientific fields:

• Medicine – development of drugs and vaccines.
• Agriculture – fertilizers, pesticides, and soil testing.
• Industry – manufacture of polymers, fuels, cosmetics, etc.
• Environment – pollution control, green chemistry.
• Food and materials – preservatives, food chemistry, alloys.

2. Matter and Its Classification

Matter is anything that has mass and occupies space.
It can be classified as:

(a) Physical States

• Solid – definite shape and volume (e.g., ice).
• Liquid – definite volume but no fixed shape (e.g., water).
• Gas – neither definite shape nor volume (e.g., oxygen).

(b) Composition

1. Pure Substances
   • Have fixed composition.
   • Elements – composed of one type of atom (e.g., O₂, N₂).
   • Compounds – composed of two or more elements in a fixed ratio (e.g., H₂O, CO₂).
2. Mixtures
   • Combinations of two or more substances in any proportion.
   • Homogeneous mixture – uniform composition (e.g., air, salt solution).
   • Heterogeneous mixture – non-uniform composition (e.g., sand + water).

3. Properties of Matter and Their Measurement

Properties of matter are classified as:

Physical Properties

Can be measured without changing chemical composition, e.g., color, density, melting point, boiling point.

Chemical Properties

Can be observed only when a substance undergoes a chemical change, e.g., flammability, reactivity.

Measurement Units:

• The SI system (Système International d’Unités) is used universally.
• Base quantities and SI units:

4. Uncertainty in Measurement

Every measurement has some degree of uncertainty due to limitations of instruments or human error.

• Accuracy: Closeness of a measurement to the true value.
• Precision: Closeness of repeated measurements to each other.
• Significant Figures: Digits that carry meaningful information about precision.

Rules for Significant Figures:

1. All non-zero digits are significant.
2. Zeros between non-zero digits are significant.
3. Leading zeros are not significant.
4. Trailing zeros are significant only if a decimal point is present.

Example:
0.00456 → 3 significant figures
1.2300 → 5 significant figures

5. Laws of Chemical Combination

1. Law of Conservation of Mass

Mass can neither be created nor destroyed during a chemical reaction.

Example:
When 10 g of calcium carbonate decomposes, it produces 5.6 g of CaO and 4.4 g of CO₂.
→ Total mass before = Total mass after.

2. Law of Definite Proportions

A chemical compound always contains the same elements in a fixed ratio by mass.

Example:
Water from any source (river or rain) always contains hydrogen and oxygen in the ratio 1:8 by mass.

3. Law of Multiple Proportions

When two elements combine to form more than one compound, the masses of one element that combine with a fixed mass of the other are in simple whole-number ratios.

Example:
Carbon forms CO and CO₂.
In CO, 12 g C combines with 16 g O;
in CO₂, 12 g C combines with 32 g O.
Ratio = 16:32 = 1:2.

4. Law of Gaseous Volumes (Gay-Lussac’s Law)

Under similar conditions of temperature and pressure, the volumes of reacting gases and products bear a simple whole-number ratio.

Example:
2H₂ + O₂ → 2H₂O
Volumes: 2 : 1 : 2

5. Avogadro’s Law

Equal volumes of all gases under the same conditions of temperature and pressure contain an equal number of molecules.

This led to the concept of the Avogadro constant (Nₐ = 6.022 × 10²³ particles per mole).

6. Dalton’s Atomic Theory

Proposed by John Dalton (1808), this theory explains the laws of chemical combination.

Postulates:

1. Matter is made up of indivisible particles called atoms.
2. All atoms of a given element are identical.
3. Compounds are formed when atoms combine in simple whole-number ratios.
4. Chemical reactions involve rearrangement of atoms; no atoms are created or destroyed.

7. Atomic and Molecular Masses

Atomic Mass

• The mass of an atom compared to 1/12th of the mass of one carbon-12 atom.
• Atomic Mass Unit (amu or u): 1 u = 1/12 of the mass of a C-12 atom = 1.66056 × 10⁻²⁷ kg.

Molecular Mass

• Sum of atomic masses of all atoms in a molecule.
  Example:
  H₂O = (2 × 1.008) + 16.00 = 18.016 u

8. Mole Concept and Molar Mass

The mole is the amount of substance that contains as many entities (atoms, molecules, ions, etc.) as there are atoms in 12 g of carbon-12.

• 1 mole = 6.022 × 10²³ entities (Avogadro number)
• Molar mass (M): Mass of one mole of a substance (in g/mol)

Example:
1 mole of H₂O = 18 g
1 mole of CO₂ = 44 g

Relations:

Number of moles=Given mass (g)Molar mass (g/mol)\text{Number of moles} = \frac{\text{Given mass (g)}}{\text{Molar mass (g/mol)}}Number of moles=Molar mass (g/mol)Given mass (g)​ Number of particles=Number of moles×6.022×1023\text{Number of particles} = \text{Number of moles} \times 6.022 \times 10^{23}Number of particles=Number of moles×6.022×1023

9. Percentage Composition

% of element=Mass of element in compoundMolar mass of compound×100\% \text{ of element} = \frac{\text{Mass of element in compound}}{\text{Molar mass of compound}} \times 100% of element=Molar mass of compoundMass of element in compound​×100

Example:
For H₂O:
% H = (2×1.008 / 18.016)×100 = 11.19%
% O = 88.81%

10. Empirical and Molecular Formula

• Empirical formula – simplest whole-number ratio of atoms.
• Molecular formula – actual number of atoms in a molecule.

Molecular Formula=n×Empirical Formula\text{Molecular Formula} = n \times \text{Empirical Formula}Molecular Formula=n×Empirical Formula n=Molar massEmpirical formula massn = \frac{\text{Molar mass}}{\text{Empirical formula mass}}n=Empirical formula massMolar mass​

11. Stoichiometry and Stoichiometric Calculations

Stoichiometry deals with quantitative relationships between reactants and products in a chemical reaction.

Example:
2H₂ + O₂ → 2H₂O
→ 2 moles H₂ react with 1 mole O₂ to form 2 moles H₂O.

Steps in Stoichiometric Calculations:

1. Write balanced chemical equation.
2. Convert given data into moles.
3. Use mole ratio to find unknown.
4. Convert moles back to grams, if required.

12. Limiting Reagent

When reactants are not in stoichiometric proportions, one reactant is limiting (completely consumed first), and the other is in excess.

13. Concentration of Solutions

• Mass % = (Mass of solute / Mass of solution) × 100
• Mole Fraction (x) = Moles of component / Total moles
• Molarity (M) = Moles of solute / Volume of solution (L)
• Molality (m) = Moles of solute / Mass of solvent (kg)

Conclusion

The first chapter of chemistry lays the foundation for all upcoming topics. Understanding the laws of chemical combination, mole concept, stoichiometry, and measurement techniques is essential for mastering the subject.