Unit 1: Some Basic Concepts of Chemistry Comprehensive Notes

Some Basic Concepts of Chemistry

Science and Chemistry

• Science: A systematic human effort to understand nature.

• Chemistry: The study of the preparation, properties, structure, and reactions of material substances.

Development of Chemistry

• Ancient pursuits:

  • Philosopher’s stone (Paras): To convert base metals into gold.

  • Elixir of life: To grant immortality.

• Ancient India:

  • Knowledge of scientific phenomena applied in various aspects of life.

  • Chemistry developed as Alchemy and Iatrochemistry (1300-1600 CE).

• Modern Chemistry: Took shape in 18th century Europe, influenced by alchemical traditions introduced by Arabs.

Contributions of India

• Ancient names for chemistry: Rasayan Shastra, Rastantra, Ras Kriya, or Rasvidya.

• Included: Metallurgy, medicine, cosmetics, glass, dyes.

• Archaeological Evidence:

  • Mohenjodaro and Harappa: Baked bricks, mass production of pottery (early chemical process).

  • Glazed pottery, gypsum cement (lime, sand, CaCO_3).

  • Faience (glass) used in ornaments.

  • Metalwork: Lead, silver, gold, copper; hardening of copper with tin and arsenic.

• Glass objects: Found in Maski (South India, 1000–900 BCE) and Hastinapur & Taxila (North India, 1000–200 BCE).

• Copper metallurgy: Evidences from chalcolithic cultures.

• Textile Arts: Tanning of leather and dyeing of cotton (Rigveda, 1000–400 BCE).

• Kiln Technology: Mastery in controlling temperatures for black polished ware.

• Salt Production: Described in Kautilya’s Arthashastra.

• Vedic Literature: Agreement with modern scientific findings.

• Materials Found: Copper utensils, iron, gold, silver ornaments, terracotta discs, painted grey pottery.

• Sushruta Samhita: Importance of Alkalies.

• Charaka Samhita: Preparation of sulphuric acid, nitric acid, oxides of copper, tin, zinc; sulphates of copper, zinc, iron; carbonates of lead and iron.

• Rasopanishada: Preparation of gunpowder mixture.

• Tamil texts: Fireworks using sulphur, charcoal, saltpetre (KNO_3), mercury, camphor.

• Nagarjuna: A reputed chemist, alchemist, and metallurgist.

  • Rasratnakar: Mercury compounds.

  • Extraction methods: Gold, silver, tin, copper.

• Rsarnavam (c. 800 CE): Furnaces, ovens, crucibles; metal identification by flame color.

• Chakrapani: Discovered mercury sulphide and invented soap (mustard oil and alkalies).

• Indian Soaps (18th century CE): Oil of Eranda, seeds of Mahua plant, calcium carbonate.

• Ajanta and Ellora Paintings: Testify to high level of science in ancient India.

• Varähmihir’s Brihat Samhita: Encyclopaedia with recipes for glutinous material for walls and roofs (extracts of plants, fruits, seeds, barks, resins).

• Atharvaveda (1000 BCE): Dye stuffs (turmeric, madder, sunflower, orpiment, cochineal, lac, kamplcica, pattanga, jatuka).

• Varähmihir’s Brihat Samhita: Perfumes and cosmetics.

• Hair Dye Recipes: Indigo, iron powder, black iron/steel, acidic extracts of sour rice gruel.

• Gandhayukli: Scents, mouth perfumes, bath powders, incense, talcum powder.

• Paper & Ink: Known in India by the 17th century; ink used from the 4th century (chalk, red lead, minimum).

• Fermentation: Well-known to Indians; liquors mentioned in Vedas and Kautilya’s Arthashastra (barks, stems, flowers, leaves, woods, cereals, fruits, sugarcane for making Asavas).

• Atomic Theory: Concept of indivisible building blocks appeared in India centuries before BCE.

• Acharya Kanda (600 BCE): First proponent of atomic theory.

  • Theory: Substances are aggregates of smaller units called atoms (Paramãnu).

  • Paramãnu: Eternal, indestructible, spherical, suprasensible, in motion.

  • Varieties of atoms, forming pairs/triplets; interaction caused by unseen forces.

  • Conceptualized 2500 years before John Dalton (1766-1844).

• Charaka Samhita: Oldest Ayurvedic epic describing treatment of diseases.

  • Reduction of particle size of metals (nanotechnology).

  • Use of bhasma of metals in treatment (nanoparticles of metals).

• Decline: Alchemy and Iatrochemistry declined with Western medicine in the 20th century.

• Pharmaceutical Industry: Ayurveda-based industry declined gradually.

• Modern Science: Appeared in India in the late 19th century with European scientists.

Chemistry - The Science of Atoms and Molecules

• Deals with the composition, structure, properties, and interactions of matter.

• Essential for human beings in daily life.

• Basic constituents of matter: Atoms and molecules.

• Quantitative description: Physical properties described using numerical values with units.

Importance of Chemistry

• Central role in science, intertwined with other branches.

• Applications: Weather patterns, brain function, computers, chemical industries, fertilizers, alkalis, acids, salts, dyes, polymers, drugs, soaps, detergents, metals, alloys, new materials.

• National Economy: Large-scale production of fertilizers, pesticides, insecticides.

• Healthcare: Isolation and synthesis of life-saving drugs (cisplatin, taxol for cancer; AZT for AIDS).

• National Development: Design and synthesis of new materials (superconducting ceramics, conducting polymers, optical fibres).

• Industries: Manufacturing utility goods (acids, alkalies, dyes, polymers, metals), contributing to economy and employment.

• Environmental Issues: Safer alternatives to CFCs for ozone depletion; managing greenhouse gases (methane, carbon dioxide).

• Future Challenges: Biochemical processes, enzymes for chemical production, synthesis of new exotic materials.

Nature of Matter

• Matter: Anything with mass that occupies space.

• Examples: Book, pen, pencil, water, air, living beings.

States of Matter

• Three physical states: Solid, liquid, gas.

• Solids: Particles close, orderly arrangement, limited movement; definite volume and shape.

• Liquids: Particles close, can move; definite volume, no definite shape (takes container's shape).

• Gases: Particles far apart, easy and fast movement; no definite volume or shape (occupies container's space).

• Interconversion: Changing temperature and pressure can convert between states (Solid \rightleftharpoons Liquid \rightleftharpoons Gas).

Classification of Matter

• Macroscopic level: Mixture or pure substance.

• Pure substance: All constituent particles are the same in chemical nature.

• Mixture: Contains particles of two or more pure substances in any ratio; variable composition; pure substances are components.

• Homogeneous mixture: Components completely mix, uniform distribution, uniform composition (e.g., sugar solution, air).

• Heterogeneous mixture: Non-uniform composition, different components visible (e.g., salt and sugar, grains and pulses with dirt).

• Separation: Components of a mixture can be separated by physical methods (hand-picking, filtration, crystallisation, distillation).

• Pure substances: Fixed composition (e.g., copper, silver, gold, water, glucose); constituents cannot be separated by physical methods.

• Elements: Particles of one type of atom (e.g., sodium, copper, silver, hydrogen, oxygen); atoms of different elements are different.

• Molecules: Two or more atoms (e.g., hydrogen, nitrogen, oxygen gases).

• Compounds: Atoms of different elements combine in a definite ratio; constituents cannot be separated by physical methods but can by chemical methods (e.g., water, ammonia, carbon dioxide, sugar).

• Properties: Compound properties differ from constituent elements (e.g., hydrogen and oxygen gases form liquid water; hydrogen burns, oxygen supports combustion, water extinguishes fire).

Properties of Matter and Their Measurement

Physical and Chemical Properties

• Physical properties: Measured without changing identity or composition (e.g., color, odor, melting point, boiling point, density).

• Chemical properties: Require a chemical change to occur (e.g., composition, combustibility, reactivity with acids and bases).

• Chemists: Describe, interpret, and predict substance behavior based on physical and chemical properties through measurement and experimentation.

Measurement of Physical Properties

• Quantitative measurement: Required for scientific investigation (length, area, volume).

• Representation: Number followed by units (e.g., 6 m).

• Measurement Systems:

  • English System and Metric System exist.

  • Metric system (France, late 18th century) more convenient due to decimal system.

  • International System of Units (SI) established in 1960 as a common standard.

The International System of Units (SI)

• Established by the General Conference on Weights and Measures (CGPM).

• Seven base units: Listed in Table 1.1 (length, mass, time, electric current, thermodynamic temperature, amount of substance, luminous intensity).

• Derived quantities: Speed, volume, density, etc. can be derived from base quantities.

• National Standards of Measurement:

  • Maintained by National Metrology Institutes (NMI) in each industrialized country, including India (National Physical Laboratory, NPL, New Delhi).

  • Establish experiments, maintain standards, and inter-compare with other NMIs and the International Bureau of Standards in Paris.

• Unit Definitions: Evolve with improved measurement accuracy.

• Table 1.1: Base Physical Quantities and their Units:

  • 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)

• Table 1.2: Definitions of SI Base Units (length, mass, time, electric current, thermodynamic temperature, amount of substance, luminous intensity).

• Prefixes: Used to indicate multiples or submultiples of a unit (Table 1.3, yocto to yotta).

Mass and Weight

• Mass: Amount of matter in a substance (constant).

• Weight: Force exerted by gravity on an object (varies with gravity).

• Measurement: Analytical balance used for accurate mass determination.

• SI unit: Kilogram (kg), but gram (g) is used in laboratories (1 kg = 1000 g).

Volume

• Volume: Amount of space occupied by a substance; units of (length)^3.

• SI unit: m^3, but cm^3 or dm^3 often used in chemistry laboratories.

• Litre (L): Common unit for liquid volume (1 L = 1000 mL, 1000 cm^3 = 1 dm^3).

• Measurement: Graduated cylinder, burette, pipette, volumetric flask (for known volumes).

Density

• Density: Mass per unit volume.

• Formula: Density = \frac{Mass}{Volume}

• SI units: kg m^{-3} (often expressed as g cm^{-3}).

• Indication: Density indicates how closely particles are packed.

Temperature

• Scales: °C (Celsius), °F (Fahrenheit), K (Kelvin; SI unit).

• Thermometers: Calibrated from 0° to 100°C (Celsius) and 32° to 212°F (Fahrenheit).

• Relationships:

  • °F = \frac{9}{5}°C + 32

  • K = °C + 273.15

• Kelvin Scale: No negative temperatures.

Uncertainty in Measurement

Scientific Notation

• Used for very large or very small numbers.

• Form: N × 10^n, where N is between 1.000… and 9.999…, and n is an exponent.

• Examples: 232.508 = 2.32508 × 10^2, 0.00016 = 1.6 × 10^{-4}.

• Mathematical Operations:

  • Multiplication and Division: Use exponential number rules.

  • Addition and Subtraction: Adjust numbers to have the same exponent before adding or subtracting coefficients.

Significant Figures

• Meaningful digits with certainty plus one estimated/uncertain digit.

• Indicate uncertainty by writing certain digits and the last uncertain digit (e.g., 11.2 mL means 11 is certain, 2 is uncertain, uncertainty of ±1).

• Rules for Determining:

  1. All non-zero digits are significant (e.g., 285 cm has 3 significant figures, 0.25 mL has 2).

  2. Zeros preceding the first non-zero digit are not significant (e.g., 0.03 has 1 significant figure, 0.0052 has 2).

  3. Zeros between non-zero digits are significant (e.g., 2.005 has 4 significant figures).

  4. Zeros at the end or right of a number are significant if on the right side of the decimal point (e.g., 0.200 g has 3 significant figures). Otherwise, terminal zeros are not significant if there is no decimal point (e.g., 100 has only 1 significant figure, but 100. has 3, and 100.0 has 4).

  5. Counting numbers have infinite significant figures (e.g., 2 balls = 2.000000…). Exact numbers, and defined constants could be considered to have an infinite number of significant figures.

• Scientific Notation: All digits are significant (e.g., 4.01 × 10^2 has 3 significant figures, 8.256 × 10^{-3} has 4).

• Precision: Closeness of various measurements for the same quantity.

• Accuracy: Agreement of a particular value to the true value of the result.

• Table 1.4: Data to Illustrate Precision and Accuracy (example of measurements by students A, B, and C).

• Addition and Subtraction:

  • Result cannot have more digits to the right of the decimal point than either of the original numbers (e.g. 12.11 + 18.0 + 1.012 = 31.122 reported as 31.1).

• Multiplication and Division:

  • Result must be reported with no more significant figures than in the measurement with the fewest significant figures (e.g. 2.5 \times 1.25 = 3.125 reported as 3.1).

• Rounding off the Numbers:

  1. If the rightmost digit to be removed is more than 5, the preceding number is increased by one (e.g., 1.386 rounded to 1.39).

  2. If the rightmost digit to be removed is less than 5, the preceding number is not changed (e.g., 4.334 rounded to 4.33).

  3. If the rightmost digit to be removed is 5, then the preceding number is not changed if it is an even number, but it is increased by one if it is an odd number (e.g., 6.35 rounded to 6.4, 6.25 rounded to 6.2).

Dimensional Analysis

• Method to convert units from one system to another (factor label method/unit factor method).

• Example 1: Convert 3 inches to cm (1 in = 2.54 cm, 3 in = 7.62 cm).

• Example 2: Convert 2 L to m^3 (1 L = 1000 cm^3, 1 m = 100 cm, 2 L = 0.002 m^3).

• Example 3: Convert 2 days to seconds.

  • 2 \text{ days} = 2 \text{ days} \times \frac{24 \text{ h}}{1 \text{ day}} \times \frac{60 \text{ min}}{1 \text{ h}} \times \frac{60 \text{ s}}{1 \text{ min}} = 172800 \text{ s}

Laws of Chemical Combinations

• Combination of elements to form compounds governed by five basic laws.

Law of Conservation of Mass

• Proposed by Antoine Lavoisier in 1789.

• In all physical and chemical changes, there is no net change in mass.

• Matter can neither be created nor destroyed.

Law of Definite Proportions

• Given by Joseph Proust.

• A given compound always contains exactly the same proportion of elements by weight.

• Example: Cupric carbonate samples (natural and synthetic) have the same composition.

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.

• Example: Hydrogen and oxygen form water and hydrogen peroxide (16:32 or 1:2 ratio of oxygen).

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.

• Example: 100 mL of hydrogen combines with 50 mL of oxygen to give 100 mL of water vapour (2:1 ratio).

Avogadro’s Law

• Proposed by Avogadro in 1811.

• Equal volumes of all gases at the same temperature and pressure should contain equal numbers of molecules.

• Distinction between atoms and molecules.

• Example: Two volumes of hydrogen combine with one volume of oxygen to give two volumes of water.

Dalton’s Atomic Theory

• Proposed in 1808 in ‘A New System of Chemical Philosophy’.

  1. Matter consists of indivisible atoms.

  2. All atoms of a given element have identical properties, including identical mass. Atoms of different elements differ in mass.

  3. Compounds are formed when atoms of different elements combine in a fixed ratio.

  4. Chemical reactions involve reorganization of atoms. These are neither created nor destroyed in a chemical reaction.

• Explained laws of chemical combination but not laws of gaseous volumes.

Atomic and Molecular Masses

Atomic Mass

• Mass of an atom is very small.

• Determined relative to another by experimental means.

• Present system: Based on carbon-12, assigned a mass of exactly 12 atomic mass units (amu) in 1961.

• One atomic mass unit: Mass exactly equal to one-twelfth of the mass of one carbon-12 atom.

• 1 \text{ amu} = 1.66056 × 10^{-24} \text{ g}

• Mass of hydrogen atom: 1.0078 amu ≈ 1.0080 amu.

• Oxygen-16 atom: 15.995 amu.

• ‘amu’ replaced by ‘u’ (unified mass).

Average Atomic Mass

• Many elements exist as more than one isotope.

• Average atomic mass: Takes into account the existence of isotopes and their relative abundance.

• Example: Carbon isotopes and average atomic mass calculation. \text{Average atomic mass} = (0.98892)(12 \text{ u}) + (0.01108)(13.00335 \text{ u}) + (2 × 10^{-12})(14.00317 \text{ u}) = 12.011 \text{ u}

Molecular Mass

• Sum of atomic masses of elements present in a molecule.

• Obtained by multiplying the atomic mass of each element by the number of its atoms and adding them together.

• Example: Methane (CH4) molecular mass = 12.011 u + 4(1.008 u) = 16.043 u.

• Water molecular mass (H2O) = 2(1.008 u) + 16.00 u = 18.02 u.

Formula Mass

• Substances like sodium chloride (NaCl) do not contain discrete molecules.

• Ions arranged in a three-dimensional structure.

• Formula mass: Sum of atomic masses of ions in the formula unit (e.g., NaCl formula mass = 23.0 u + 35.5 u = 58.5 u).

Mole Concept and Molar Masses

Mole

• Unit to count entities at the microscopic level (atoms, molecules, particles, electrons, ions).

• SI system: Mole (mol) is the seventh base quantity for the amount of a substance.

• 1 \text{ mole} = 6.02214076 × 10^{23} \text{ elementary entities}. Avogadro constant (
  N_A).

• 1 mol of hydrogen atoms = 6.022 × 10^{23} atoms.

• 1 mol of water molecules = 6.022 × 10^{23} water molecules.

• 1 mol of sodium chloride = 6.022 × 10^{23} formula units of sodium chloride.

Molar Mass

• Mass of one mole of a substance in grams.

• Numerically equal to atomic/molecular/formula mass in u.

• Molar mass of water = 18.02 g mol^{-1}.

• Molar mass of sodium chloride = 58.5 g mol^{-1}.

Percentage Composition

• Information regarding the percentage of a particular element present in a compound.

• Formula: \text{Mass % of an element} = \frac{\text{mass of that element in the compound}}{\text{molar mass of the compound}} × 100

• Example: Water (H2O) mass % of hydrogen = 11.18, mass % of oxygen = 88.79.

• Ethanol mass % of carbon = 52.14%, mass % of hydrogen = 13.13%, mass % of oxygen = 34.73%.

Empirical Formula and Molecular Formula

• Empirical formula: Simplest whole number ratio of atoms in a compound.

• Molecular formula: Exact number of different types of atoms in a molecule.

• Determination sequence: Mass per cent → empirical formula → molecular formula (if molar mass is known).

• Conversion of mass percent to grams

• Convert into number of moles of each element

• Divide each of the mole values obtained above by the smallest number amongst them

• Write down the empirical formula by mentioning the numbers after the symbols of respective elements

Stoichiometry and Stoichiometric Calculations

Stoichiometry

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

• Balanced chemical equation: Provides information about the reaction (reactants and products).

• Example: CH4(g) + 2O2(g) \rightarrow CO2(g) + 2H2O(g)

• Stoichiometric coefficients: Numbers representing the number of molecules (and moles) involved.

• Relationships: Molar ratios, molecular ratios, volume ratios, mass ratios.

Limiting Reagent

• Reactions often carried out with amounts of reactants that are different from the amounts required by a balanced chemical reaction.

• Reactant in least amount gets consumed first, limiting the amount of product formed.

Reactions in Solutions

• Majority of reactions carried out in solutions.

• Concentration of a solution expressed in various ways:

• Mass per cent or weight per cent (w/w %)

  • \text{Mass percent } = \frac{\text{Mass of solute}}{\text{Mass of solution}} \times 100

• Mole fraction

  • \text{Mole fraction } = \frac{\text{Moles of component}}{\text{Total moles in solution}} (
    nA and nB)

• Molarity (M)

  • \text{Molarity } = \frac{\text{Moles of solute}}{\text{Volume of solution in liters}}

• Molality (m)

  • \text{Molality } = \frac{\text{Moles of solute}}{\text{Mass of solvent in kg}}

• Dilution formula: M1V1 = M2V2.