Unit 1: Some Basic Concepts of Chemistry

Definition and Scope of Chemistry

General Definition: Science is a continuing human effort to systematise knowledge for describing and understanding nature. Chemistry is the specific branch of science that studies the preparation, properties, structure, and reactions of material substances.
Perspective by Roald Hoffmann: Chemistry is the science of molecules and their transformations. It is focused not just on the 118 elements but on the infinite variety of molecules built from them.
Core Focus: Chemistry deals with the composition, structure, properties, and interaction of matter. It is often referred to as the science of atoms and molecules.

Historical Development of Chemistry

Origins: Modern chemistry emerged from the search for two things:
- Philosopher’s Stone (Paras): A substance intended to convert base metals like iron and copper into gold.
- Elixir of Life: A substance intended to grant immortality.
Timeline and Traditions:
- Chemistry developed as Alchemy and Iatrochemistry between 1300–1600 CE.
- Modern chemistry took shape in 18th-century Europe after alchemical traditions were introduced by the Arabs.

Chemistry in Ancient India

Terminology: In ancient India, chemistry was known as Rasayan Shastra, Rastantra, Ras Kriya, or Rasvidya.
Archaeological Evidence:
- Mohenjodaro and Harappa: Excavations prove ancient knowledge of chemical processes. Baked bricks and mass production of pottery are the earliest chemical processes involving mixing, moulding, and heating.
- Construction: Gypsum cement used in construction contained lime, sand, and traces of $CaCO_3$ .
- Glass and Metals: Harappans made Faience (a sort of glass) for ornaments. They forged objects from lead, silver, gold, and copper. They used tin and arsenic to harden copper.
- South and North India: Glass objects found in Maski (1000–900 BCE) and Hastinapur/Taxila (1000–200 BCE) were colored using metal oxides.
Chemical Technologies:
- Copper and Iron: Extraction technologies were developed indigenously.
- Rigveda: Mentions tanning of leather and dyeing of cotton (1000–400 BCE).
- Kautilya’s Arthashastra: Describes the production of salt from the sea.
- Sushruta Samhita: Explains the importance of Alkalies.
- Charaka Samhita: Mentions the preparation of sulphuric acid, nitric acid, oxides of copper, tin, and zinc; sulphates of copper, zinc, and iron; and carbonates of lead and iron. It also discusses reducing particle size for medicine (termed Bhasmas), which modern science identifies as nanotechnology.
- Nagarjuna: A great Indian scientist, alchemist, and metallurgist. His work Rasratnakar deals with mercury compounds and the extraction of gold, silver, tin, and copper.
- Rsarnavam (c. 800 CE): Discusses furnaces, ovens, crucibles, and identifying metals by flame color.
- Chakrapani: Discovered mercury sulphide and is credited with inventing soap using mustard oil and alkalies.
- Varahmihir’s Brihat Samhita (6th Century CE): An encyclopedia containing recipes for glutinous material for roofs/walls made from plant extracts, fruits, seeds, and barks.
Atomic Theory: Acharya Kanda (born 600 BCE, also known as Kashyap) was the first proponent of atomic theory. He named indivisible particles Paramanu. He described them as eternal, indestructible, spherical, and in motion. He authored the Vaiseshika Sutras.

Importance of Chemistry in Modern Life

National Economy: Chemistry contributes significantly to healthcare, food, and manufacturing.
Healthcare:
- Cisplatin and Taxol: Effective in cancer therapy.
- AZT (Azidothymidine): Used for assisting AIDS patients.
Environment: Chemists synthesized safer alternatives to CFCs (chlorofluorocarbons) to protect the ozone layer. Management of Greenhouse gases ( $CH_4$ , $CO_2$ ) remains a current challenge.
Materials Science: Development of superconducting ceramics, conducting polymers, and optical fibers.

Nature and Classification of Matter

Definition: Matter is anything that has mass and occupies space.
States of Matter:
- Solids: Particles held very close in an orderly fashion; definite volume and definite shape.
- Liquids: Particles are close but can move; definite volume but no definite shape (take the shape of the container).
- Gases: Particles are far apart with fast movement; no definite volume or shape (completely fill the container).
- Interconvertibility: $Solid \rightleftharpoons Liquid \rightleftharpoons Gas$ (affected by temperature and pressure).
Classification of Matter:
1. Mixtures: Contain two or more substances in variable ratios.
  - Homogeneous: Uniform composition (e.g., air, sugar solution).
  - Heterogeneous: Non-uniform composition; components are sometimes visible (e.g., salt and sugar, grains and dirt).
  - Separation: Components can be separated by physical methods (hand-picking, filtration, crystallization, distillation).
2. Pure Substances: Fixed composition (e.g., Gold, Water, Glucose).
  - Elements: Consist of only one type of atom (e.g., $Na, Cu, H_2, O_2$ ).
  - Compounds: Atoms of different elements combined in a fixed ratio (e.g., $H_2O$ , $CO_2$ .). Properties of a compound differ from those of its constituent elements (e.g., $H_2$ and $O_2$ are gases, while $H_2O$ is a liquid).

Properties of Matter and Measurement

Physical Properties: Measured without changing identity (colour, odour, melting point, density).
Chemical Properties: Require a chemical change to observe (composition, combustibility, acidity, basicity).
The International System of Units (SI):
- Established by the 11th General Conference on Weights and Measures (CGPM).
- Seven Base Units:
  1. Length: metre ( $m$ )
  2. Mass: kilogram ( $kg$ )
  3. Time: second ( $s$ )
  4. Electric Current: ampere ( $A$ )
  5. Thermodynamic Temperature: kelvin ( $K$ )
  6. Amount of Substance: mole ( $mol$ )
  7. Luminous Intensity: candela ( $cd$ )
Mass vs. Weight:
- Mass: Amount of matter; constant.
- Weight: Force exerted by gravity; varies with location.
Density: $Density = \frac{\text{Mass}}{\text{Volume}}$ . SI unit: $kg\,m^{-3}$ . Commonly expressed as $g\,cm^{-3}$ .
Temperature Scales:
- Conversion between Fahrenheit ( ${^\circ}F$ ) and Celsius ( ${^\circ}C$ ): ${^\circ}F = \frac{9}{5}({^\circ}C) + 32$
- Conversion between Kelvin ( $K$ ) and Celsius ( ${^\circ}C$ ): $K = {^\circ}C + 273.15$
- Note: Red temperatures in Kelvin cannot be negative.

Uncertainty in Measurement

Scientific Notation: Represented as $N \times 10^n$ , where $N$ is between $1.000...$ and $9.999...$ and $n$ is an exponent.
Significant Figures:
- Rules:
  1. All non-zero digits are significant.
  2. Zeros preceding the first non-zero digit are not significant ( $0.0052$ has two).
  3. Zeros between non-zero digits are significant ( $2.005$ has four).
  4. Terminal zeros are significant only if they are on the right of a decimal point ( $0.200$ has three; $100$ has one).
  5. Exact numbers (e.g., 2 balls) have infinite significant figures.
Precision vs. Accuracy:
- Precision: Closeness of various measurements for the same quantity.
- Accuracy: Agreement of a particular value to the true value.
Dimensional Analysis: Also called the Factor Label Method or Unit Factor Method. It uses unit factors to convert one unit to another.

Laws of Chemical Combination

Law of Conservation of Mass (Antoine Lavoisier, 1789): Matter can neither be created nor destroyed in a chemical reaction.
Law of Definite Proportions (Joseph Proust): A given compound always contains exactly the same proportion of elements by weight (e.g., Cupric carbonate samples always contain $51.35\%$ Copper, $9.74\%$ Carbon, and $38.91\%$ Oxygen).
Law of Multiple Proportions (Dalton, 1803): If two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in a ratio of small whole numbers (e.g., in $H_2O$ and $H_2O_2$ , oxygen masses 16g and 32g combine with 2g hydrogen, a $1:2$ ratio).
Gay Lussac’s Law of Gaseous Volumes (1808): Gases combine or are produced in a simple ratio by volume at constant temperature and pressure (e.g., $100\,mL$ hydrogen + $50\,mL$ oxygen → $100\,mL$ water vapour; ratio $2:1$ ).
Avogadro’s Law (1811): Equal volumes of all gases at the same temperature and pressure should contain equal number of molecules.

Atomic and Molecular Masses

Atomic Mass Unit (amu): Defined as a mass exactly equal to one-twelfth of the mass of one carbon-12 atom.
- $1\,amu = 1.66056 \times 10^{-24}\,g$
- Mass of $H$ atom: $1.0080\,amu$ .
Average Atomic Mass: Calculated using the relative abundance of isotopes.
- Example Carbon: $({^{12}C}, 98.892\%) + ({^{13}C}, 1.108\%) + ({^{14}C}, 2 \times 10^{-10}\%) = 12.011\,u$ .
Molecular Mass: Sum of atomic masses of elements in a molecule.
- Example $CH_4$ : $12.011 + 4(1.008) = 16.043\,u$ .
Formula Mass: Used for substances like $NaCl$ that do not contain discrete molecules.
- Example $NaCl$ : $23.0\,u + 35.5\,u = 58.5\,u$ .

Mole Concept and Molar Masses

Mole (mol): The SI unit for amount of substance.
Avogadro Constant ( $N_A$ ): One mole contains exactly $6.02214076 \times 10^{23}$ elementary entities.
Molar Mass: The mass of one mole of a substance in grams. It is numerically equal to the atomic/molecular/formula mass in $u$ .
- Molar mass of $H_2O$ : $18.02\,g\,mol^{-1}$ .

Chemical Formulas and Stoichiometry

Mass Percent: $\text{Mass \%} = \frac{\text{Mass of element}}{\text{Molar mass of compound}} \times 100$ .
Empirical Formula: Simplest whole-number ratio of atoms in a compound.
Molecular Formula: Exact number of atoms of each element in a molecule ( $Molecular\,Formula = n \times Empirical\,Formula$ ).
Stoichiometry: Deals with calculating masses or volumes of reactants and products.
Limiting Reagent: The reactant that is consumed first and limits the amount of product formed.
Concentration Terms in Solutions:
- Mass Percent: $\frac{\text{Mass of solute}}{\text{Mass of solution}} \times 100$
- Mole Fraction ( $X_A$ ): $\frac{n_A}{n_A + n_B}$
- Molarity (M): $\frac{\text{Moles of solute}}{\text{Volume of solution in litres}}$
- Molality (m): $\frac{\text{Moles of solute}}{\text{Mass of solvent in kg}}$ (Note: Molality does not change with temperature).