Chemistry: Basic Concepts and Foundations

1. Introduction to Chemistry

What is Science?

Science is a continuous human effort to systematise knowledge for describing and understanding nature.

What is Chemistry?

It's the branch of science that studies the preparation, properties, structure, and reactions of material substances. Roald Hoffmann also defines chemistry as the science of molecules and their transformations, focusing on the infinite variety of molecules built from elements.

Historical Development:

Chemistry is not a very old discipline in its modern form. It originated from the search for two things:

Philosopher's stone (Paras): Believed to convert base metals (like iron, copper) into gold.
'Elixir of life': Supposed to grant immortality.

From 1300-1600 CE, chemistry developed mainly as Alchemy and Iatrochemistry.

Modern chemistry took shape in 18th-century Europe, building on alchemical traditions introduced by Arabs.

Contributions of Ancient India:

Ancient Indians had significant knowledge of scientific phenomena before modern science. Chemistry was known as Rasayan Shastra, Rastantra, Ras Kriya, or Rasvidya. It encompassed metallurgy, medicine, and the manufacture of cosmetics, glass, dyes, etc.

Archaeological evidence from Mohenjodaro and Harappa shows:

Use of baked bricks in construction.
Mass production of pottery, an early chemical process involving mixing, moulding, and heating materials.
Glazed pottery and gypsum cement (containing lime, sand, CaCO₃) were used.
Harappans made faience (a type of glass for ornaments) and worked with metals like lead, silver, gold, and copper, improving copper hardness with tin and arsenic.
Copper metallurgy dates back to chalcolithic cultures, with indigenous extraction technologies.

Rigveda mentions tanning leather and dyeing cotton (1000–400 BCE).

Kautilya’s Arthashastra describes salt production from the sea.

Sushruta Samhita explains the importance of Alkalies.

Charaka Samhita describes the preparation of sulphuric acid, nitric acid, oxides of copper, tin, zinc, sulphates of copper, zinc, iron, and carbonates of lead and iron. It also discusses reduction of particle size of metals (nanotechnology) and the use of bhasmas (which contain nanoparticles) in treatment.

Rasopanishada describes gunpowder mixture preparation. Tamil texts mention fireworks using sulphur, charcoal, saltpetre, mercury, camphor, etc.

Nagarjuna, a great Indian scientist, chemist, alchemist, and metallurgist, wrote Rasratnakar on mercury compounds and metal extraction methods (gold, silver, tin, copper).

Rsarnavam (c. 800 CE) discusses furnaces, ovens, crucibles, and methods for identifying metals by flame colour.

Chakrapani discovered mercury sulphide and is credited with inventing soap using mustard oil and alkalies. Indians made soaps using oil of Eranda, Mahua plant seeds, and calcium carbonate in the 18th century CE.

Paintings at Ajanta and Ellora demonstrate high levels of ancient Indian science.

Varāhmihir’s Brihat Samhita (6th century CE) describes preparing glutinous material from plant extracts for walls and roofs. It also references perfumes, cosmetics, and hair dyeing recipes from plants (indigo) and minerals.

Paper was known in India by the 17th century (I-tsing's account). Ink (from chalk, red lead, minimum) was used from the 4th century (Taxila excavations).

The process of fermentation was well-known, with Vedas and Kautilya’s Arthashastra mentioning various liquors. Charaka Samhita lists ingredients for Asavas.

The concept of matter being made of indivisible building blocks (atomic theory) appeared centuries BCE in India. Acharya Kanda (c. 600 BCE), also known as Kashyap, was the first proponent, naming these particles 'Paramānu' (comparable to atoms). He authored Vaiseshika Sutras, stating Paramānu are eternal, indestructible, spherical, suprasensible, and in motion. He conceptualized this 2500 years before John Dalton.

2. Importance of Chemistry

Central Role in Science: Chemistry is intertwined with other sciences.
Diverse Applications: Applicable in weather patterns, brain function, computer operation, and chemical industries.
Contribution to National Economy: Large-scale production of fertilisers, pesticides, insecticides, and essential chemicals like alkalis, acids, salts, dyes, polymers, drugs, soaps, detergents, metals, and alloys.
Healthcare: Provides life-saving drugs (e.g., cisplatin, taxol for cancer; AZT for AIDS).
New Materials: Helps design and synthesize materials with specific magnetic, electric, and optical properties (e.g., superconducting ceramics, conducting polymers, optical fibres).
Environmental Degradation: Successfully dealt with some issues like developing safer alternatives to CFCs (chlorofluorocarbons) for ozone depletion. However, challenges like Green House gas management remain.
Future Challenges: Understanding biochemical processes, large-scale chemical production using enzymes, and synthesizing new exotic materials are future intellectual challenges for chemists.

3. Nature of Matter

Definition of Matter: Anything that has mass and occupies space. Examples: book, pen, water, air, living beings.

States of Matter:

Matter exists in three physical states: solid, liquid, and gas.

Solids: Particles are held very close in an orderly fashion with limited movement. They have definite volume and definite shape.
Liquids: Particles are close but can move around. They have definite volume but no definite shape, taking the shape of their container.
Gases: Particles are far apart and move easily and fast. They have neither definite volume nor definite shape, completely occupying the container's space.

Interconversion: These states are interconvertible by changing temperature and pressure.

Classification of Matter:

Mixtures: Contain two or more pure substances in any ratio, so their composition is variable. Components can be separated by physical methods (e.g., hand-picking, filtration, distillation).
- Homogeneous Mixtures: Components completely mix and are uniformly distributed with uniform composition throughout (e.g., sugar solution, air).
- Heterogeneous Mixtures: Composition is not uniform, and different components may be visible (e.g., salt and sugar, grains with dirt).
Pure Substances: All constituent particles are same in chemical nature and have fixed composition. Their constituents cannot be separated by simple physical methods. Examples: copper, silver, gold, water, glucose.
- Elements: Particles consist of only one type of atoms (e.g., sodium, copper, hydrogen, oxygen). Atoms of different elements are different. Some elements exist as atoms (Na, Cu), others as molecules (H₂, N₂, O₂).
- Compounds: Formed when two or more atoms of different elements combine in a definite ratio. Properties of a compound are different from its constituent elements (e.g., H₂ and O₂ are gases, water (H₂O) is a liquid and fire extinguisher). Constituents cannot be separated by physical methods; chemical methods are required. Examples: water (H₂O), ammonia (NH₃), carbon dioxide (CO₂), sugar (C₁₂H₂₂O₁₁).

4. Properties of Matter and Their Measurement

Unique Properties: Every substance has unique or characteristic properties.

Types of Properties:

Physical Properties: Can be measured or observed without changing the identity or composition of the substance. Examples: colour, odour, melting point, boiling point, density.
Chemical Properties: Measurement or observation requires a chemical change to occur. Examples: composition, combustibility, reactivity with acids/bases.

Measurement of Physical Properties:

Quantitative measurement is essential for scientific investigation. A quantitative observation is represented by a number followed by units (e.g., 6 m).

The International System of Units (SI):

A common standard system established in 1960 by the General Conference on Weights and Measures (CGPM). Has seven base units for fundamental scientific quantities:

Length: metre (m)
Mass: kilogram (kg)
Time: second (s)
Electric current: ampere (A)
Thermodynamic temperature: kelvin (K)
Amount of substance: mole (mol)
Luminous intensity: candela (cd)

Other physical quantities (e.g., speed, volume, density) are derived from these base units.

Mass and Weight:

Mass: The amount of matter present in a substance. It is constant. SI unit is kilogram (kg); gram (g) is often used in laboratories (1 kg = 1000 g). Measured by an analytical balance.
Weight: The force exerted by gravity on an object. It may vary due to changes in gravity.

Volume:

The amount of space occupied by a substance. SI unit: (length)³, so m³. Commonly used smaller units: cm³ or dm³. A common non-SI unit for liquids is the litre (L): 1 L = 1000 mL = 1000 cm³ = 1 dm³.

Density:

Mass per unit volume. SI unit: kg m^-3. Commonly expressed as g cm^-3 in chemistry. Indicates how closely particles are packed; higher density means more closely packed.

Temperature:

Three common scales: °C (Celsius), °F (Fahrenheit), and K (Kelvin). Kelvin (K) is the SI unit. Conversion formulas:

°F = (9/5)°C + 32
K = °C + 273.15

Negative temperatures are possible in Celsius, but not in Kelvin.

5. Uncertainty in Measurement

Every experimental measurement has uncertainty due to instrument limitations and human skill.

Scientific Notation:

Used to handle extremely large or small numbers. Format: N × 10ⁿ, where N is a digit term (1.0 to 9.999...) and n is an exponent. Example: 232.508 = 2.32508 × 10²; 0.00016 = 1.6 × 10^-4.

Significant Figures:

Meaningful digits that are known with certainty, plus one estimated/uncertain digit. Uncertainty is typically ±1 in the last digit.

Rules for Determining Significant Figures:

All non-zero digits are significant. (e.g., 285 cm has 3 SF).
Zeros preceding the first non-zero digit are NOT significant. (e.g., 0.03 has 1 SF).
Zeros between two non-zero digits ARE significant. (e.g., 2.005 has 4 SF).
Zeros at the end or right of a number ARE significant if they are to the right of the decimal point. (e.g., 0.200 g has 3 SF).
Counting numbers (exact numbers) have infinite significant figures.
In scientific notation, all digits in 'N' are significant. (e.g., 4.01×10² has 3 SF).

Arithmetic Operations with Significant Figures:

Addition and Subtraction: The result cannot have more digits to the right of the decimal point than the number with the fewest digits after the decimal point.
Multiplication and Division: The result must be reported with no more significant figures than the measurement with the fewest significant figures.

Precision and Accuracy:

Precision: Refers to the closeness of various measurements to each other for the same quantity.
Accuracy: Refers to the agreement of a particular value to the true value of the result.

Dimensional Analysis (Factor Label Method):

A method used to convert units from one system to another. Uses "unit factors" (ratios equal to 1) to cancel out undesired units.

6. Laws of Chemical Combinations

Law of Conservation of Mass:

Proposed by Antoine Lavoisier in 1789. States that in any physical or chemical change, matter can neither be created nor destroyed.

Law of Definite Proportions:

Given by French chemist Joseph Proust. States that a given compound always contains exactly the same proportion of elements by weight (mass).

Law of Multiple Proportions:

Proposed by Dalton in 1803. If two elements can 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.

Gay-Lussac’s Law of Gaseous Volumes:

Given by Gay-Lussac in 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.

Avogadro’s Law:

Proposed by Avogadro in 1811. States that equal volumes of all gases at the same temperature and pressure should contain equal numbers of molecules.

7. Dalton’s Atomic Theory

Proposed by John Dalton in 1808 in 'A New System of Chemical Philosophy'.

Postulates:

Matter consists of indivisible atoms.
All atoms of a given element have identical properties, including identical mass. Atoms of different elements differ in mass.
Compounds are formed when atoms of different elements combine in a fixed ratio.
Chemical reactions involve the reorganisation of atoms; atoms are neither created nor destroyed in a chemical reaction.

8. Atomic and Molecular Masses

Atomic Mass:

Current system is based on Carbon-12 (¹²C) as the standard, assigned a mass of exactly 12 atomic mass units (amu). One atomic mass unit (amu) is defined as a mass exactly equal to one-twelfth (1/12) of the mass of one carbon-12 atom. 1 amu = 1.66056 × 10^-24 g. 'amu' is now replaced by 'u' (unified mass).

Average Atomic Mass:

Many elements exist as isotopes. When calculating atomic masses, the existence and relative abundance (percent occurrence) of isotopes are taken into account to compute the average atomic mass.

Molecular Mass:

The sum of the atomic masses of all elements present in a molecule. Example: Molecular mass of methane (CH₄) = (12.011 u) + 4(1.008 u) = 16.043 u.

Formula Mass:

Used for substances (like ionic compounds, e.g., NaCl) that do not contain discrete molecules. The formula mass is calculated by summing the atomic masses of the elements in the formula unit (e.g., NaCl: 23.0 u + 35.5 u = 58.5 u).

9. Mole Concept and Molar Masses

Mole:

A unit used to count entities at the microscopic level. In the SI system, mole (symbol mol) is the seventh base quantity. One mole contains exactly 6.02214076 × 10²³ elementary entities. This number is called the Avogadro constant (N_A).

Molar Mass:

The mass of one mole of a substance in grams. Numerically equal to its atomic mass, molecular mass, or formula mass expressed in grams. Example: Molar mass of water = 18.02 g mol^-1.

10. Percentage Composition

Provides information about the percentage of each element present in a compound. Formula: Mass % of an element = (mass of that element in the compound / molar mass of the compound) × 100.

11. Empirical Formula and Molecular Formula

Empirical Formula: Represents the simplest whole number ratio of various atoms present in a compound.
Molecular Formula: Shows the exact number of different types of atoms present in a molecule of a compound.

Determination:

Convert mass percent of elements to grams.
Convert grams of each element to moles.
Divide moles of each element by the smallest mole value to get the simplest ratio.
Write the empirical formula.
Calculate the empirical formula mass.
Divide the molar mass by the empirical formula mass to find 'n'.
Multiply subscripts in the empirical formula by 'n' to get the molecular formula.

12. Stoichiometry and Stoichiometric Calculations

Stoichiometry: Deals with the calculation of masses (and sometimes volumes) of reactants and products involved in a chemical reaction.

Balanced Chemical Equation:

A balanced equation has the same number of atoms of each element on both sides. It provides information about reactants, products, states of matter, and molar ratios.

Limiting Reagent (or Limiting Reactant):

In a reaction, the reactant that gets consumed first, thereby limiting the amount of product formed.