UNIT 1 — Some Basic Concepts of Chemistry

DEVELOPMENT OF CHEMISTRY (HISTORICAL & CULTURAL CONTEXT)

  • Ancient objective of Alchemy:
    • Philosopher’s Stone (Paras) to transmute base metals to gold.
    • Elixir of Life for immortality.
  • Indian tradition (Rasayan Shastra / Rasvidya):
    • Indus–Saraswati sites (Mohenjodaro, Harappa) show baked bricks, glazed pottery, faience, metallurgy (Cu, Pb, Ag, Au) and gypsum cement.
    • Mastery of kiln temperature → black-polished ware (chemical mystery even today).
    • Vedic & post-Vedic texts (Rigveda, Atharvaveda, Kautilya’s Arthashastra, Charaka & Sushruta Samhita, Rasopanishada, Rasratnakar, Rsarnavam): methods for dyes, tanning, salt-making, gun-powder, mineral acids, oxides, sulphates, carbonates, bhasma-nanotechnology, furnaces, flame–colour tests, soap (mustard oil + alkali) etc.
    • Nagarjuna, Chakrapani, Varahamihira, Acharya Kanda (Paramanu atomic theory, >2500 yrs before Dalton).
    • Decline during colonial influx; revival with modern science (19th-20th cent.).

ROLE & RELEVANCE OF CHEMISTRY

  • Central science—links physics, biology, geology, material science.
  • Contributes to: food (fertilisers, pesticides), health (drugs e.g., cisplatin, taxol, AZT), materials (polymers, alloys, superconductors, optical fibres), environment (CFC alternatives, greenhouse management), computing, weather, brain studies.
  • Generates employment & national economy; India needs creative chemists.

MATTER: NATURE & CLASSIFICATION

  • Definition: Anything that has mass & occupies space.

States of Matter

  • Solid: particles closely packed, fixed shape & volume.
  • Liquid: close but mobile; fixed volume, variable shape.
  • Gas: particles far apart, free motion; no fixed volume/shape.
  • Inter-conversion by temperature/pressure: Solid⇌Liquid⇌Gas\text{Solid}\rightleftharpoons Liquid\rightleftharpoons Gas (melting, vaporisation, condensation, freezing).

Classification at Macroscopic Level

  • Pure Substance vs Mixture.
    • Pure: identical particles; fixed composition; cannot be separated by physical means.
    – Elements: one kind of atoms (Na, Cu) or molecules O<em>2, N</em>2\text{O}<em>2,\ \text{N}</em>2.
    – Compounds: atoms of ≥2 elements in fixed ratio; properties differ from constituents (e.g., H<em>2+O</em>2→H<em>2O\text{H}<em>2 + \text{O}</em>2 \rightarrow \text{H}<em>2\text{O} – water extinguishes fire though H(2) burns & O(_2) supports combustion).
    • Mixtures: ≥2 pure substances, variable composition, separable physically.
    – Homogeneous (uniform; sugar solution, air).
    – Heterogeneous (non-uniform; salt–sugar mix, grains + stones).

PROPERTIES OF MATTER & MEASUREMENT

Physical vs Chemical Properties

  • Physical: colour, odour, MP, BP, density; measured without changing identity.
  • Chemical: composition, reactivity, acidity/basicity, combustibility; require chemical change.

Measurement & Uncertainty

  • Quantitative datum = number + unit.
  • Historical systems: English & Metric → modern SI (Le Système International d’Unités).
SI Base Quantities
QuantityUnitSymbol
Lengthmetremm
Masskilogramkgkg
Timesecondss
Electric currentampereAA
Thermodynamic temp.kelvinKK
Amount of substancemolemolmol
Luminous intensitycandelacdcd
  • Definitions linked to physical constants (speed of light cc, Planck constant hh, cesium frequency ΔνCs\Delta\nu_{Cs}, etc.).
  • National Physical Laboratory (NPL-India) maintains standards.
Prefixes for multiples/sub-multiples
  • 10310^{3} kilo kk, 10−310^{-3} milli mm, 10−610^{-6} micro μ\mu, 10610^{6} mega MM, etc.

Derived Physical Quantities

  • Volume: m3m^{3} (common lab units: cm3, dm3, L (=10−3 m3)cm^{3},\ dm^{3},\ L\,(=10^{-3}\,m^{3})).
  • Density: ρ=massvolume\rho = \dfrac{\text{mass}}{\text{volume}} units kg m−3kg\,m^{-3} or g cm−3g\,cm^{-3}.
  • Temperature scales:
    • Celsius: ∘C^\circ C ; K=∘C+273.15K = ^\circ C + 273.15.
    • Fahrenheit: ∘F=95(∘C)+32^\circ F = \dfrac{9}{5}(^\circ C) + 32.
    • Kelvin has no negatives.
  • Mass vs Weight: mass constant; weight depends on gravity.

Scientific Notation & Significant Figures

  • Any number N=a×10nN = a \times 10^{n} where 1\le a<10.
  • Operations follow exponent rules.
  • Significant figures (s.f.): certain digits + one uncertain; rules for zeros provided; rounding conventions.
  • Precision (repeatability) vs Accuracy (closeness to true value) illustrated via students A, B, C.

Dimensional Analysis

  • Unit-factor method to convert units (e.g., 3 in×2.54 cm1 in=7.62 cm3\,in \times \dfrac{2.54\,cm}{1\,in}=7.62\,cm).

LAWS OF CHEMICAL COMBINATION

  1. Conservation of Mass (Lavoisier, 1789): matter neither created nor destroyed.
  2. Definite Proportions (Proust): a given compound has constant elemental mass ratio (natural vs synthetic CuCO3CuCO_3 identical composition).
  3. Multiple Proportions (Dalton, 1803): if two elements form >1 compounds, mass ratios are simple integers (C & O in CO vs CO(_2)).
  4. Gay-Lussac’s Law of Gaseous Volumes (1808): reacting gas volumes at same T,PT,P are in simple integer ratio (100 mL H<em>2H<em>2 : 50 mL O</em>2O</em>2 → 100 mL H2OH_2O vapour).
  5. Avogadro’s Law (1811): equal gas volumes at same T,PT,P contain equal molecules; introduced molecule concept, supported diatomic gases; paved way for atomic masses.

DALTON’S ATOMIC THEORY (1808)

  • Matter = indivisible atoms.
  • Atoms of same element identical in mass & properties; different elements differ.
  • Compounds = fixed atom ratios.
  • Chemical change = rearrangement of atoms; atoms conserved.

ATOMIC, MOLECULAR & FORMULA MASSES

  • Standard: 12C=12 u^{12}C = 12\,u exactly.
  • 1 u=1121\,u = \dfrac{1}{12} mass of one 12C^{12}C atom =1.66056×10−24 g=1.66056\times10^{-24}\,g.
  • Average atomic mass incorporates isotopic abundance (Cl example yields 35.5 u35.5\,u).
  • Molecular mass = Σ atomic masses in molecule: e.g., CH4CH_4 → 16.043 u16.043\,u.
  • Formula mass for ionic lattices (NaCl =23+35.5=58.5 u=23+35.5=58.5\,u).

MOLE CONCEPT & MOLAR MASS

  • 1 mol=6.02214076×10231\,mol = 6.02214076\times10^{23} entities (Avogadro constant NAN_A).
  • Molar mass (g mol⁻¹) numerically equals atomic/molecular mass in u.
  • Visual: one mole each of atoms, molecules, ions occupies different masses/volumes but same count.

PERCENTAGE COMPOSITION, EMPIRICAL & MOLECULAR FORMULAE

  • Mass % element =mass of element in formulamolar mass compound×100=\dfrac{\text{mass of element in formula}}{\text{molar mass compound}}\times100.
    • Water: %H=11.2, %O=88.8\%H = 11.2,\ \%O = 88.8.
  • To derive empirical formula:
    1. Convert % → grams (assume 100 g).
    2. Convert g → moles (divide by atomic mass).
    3. Divide by smallest mole value → simplest ratio.
    4. If necessary multiply to whole numbers.
  • Molecular formula =(empirical)n=\text{(empirical)}_n where n=molar massempirical massn=\dfrac{\text{molar mass}}{\text{empirical mass}}.

STOICHIOMETRY & CALCULATIONS

  • Balanced equation gives stoichiometric coefficients—molar, mass & volume ratios.
  • Example combustion: CH<em>4+2O</em>2→CO<em>2+2H</em>2OCH<em>4 + 2O</em>2 \rightarrow CO<em>2 + 2H</em>2O.
    • 1 mol CH<em>4CH<em>4 reacts with 2 mol O</em>2O</em>2 → 1 mol CO<em>2CO<em>2 + 2 mol H</em>2OH</em>2O.
    • 16 g CH<em>4CH<em>4 → 36 g H</em>2OH</em>2O (mass) or 22.7 L CH<em>4CH<em>4 → 45.4 L H</em>2OH</em>2O at STP (volume).

Limiting Reagent

  • Reactant consumed first limits product yield; identify by mole comparison against stoichiometry.

Concentration Units for Solutions

  1. Mass % (w/w): mass solutemass solution×100\dfrac{\text{mass solute}}{\text{mass solution}}\times100.
  2. Mole Fraction x<em>A=n</em>An<em>A+n</em>Bx<em>A = \dfrac{n</em>A}{n<em>A+n</em>B}.
  3. Molarity M=nV<em>solution (L)M = \dfrac{n}{V<em>{solution}\,(L)}; temperature-dependent; dilution obeys M</em>1V<em>1=M</em>2V2M</em>1V<em>1 = M</em>2V_2.
  4. Molality m=nmasssolvent (kg)m = \dfrac{n}{mass_{solvent}\,(kg)}; temperature-independent.
Laboratory Glassware
  • Graduated cylinder, burette, pipette, volumetric flask for precise volume measurement.
  • Analytical vs platform balance for mass; first provides higher accuracy.

SCIENTIFIC & ETHICAL INSIGHTS

  • Chemistry’s evolution from mysticism (alchemy) to quantitative science mirrors humanity’s quest for understanding nature.
  • Ethical use: drug discovery, environmental stewardship (CFC alternatives, GH gases management).
  • Philosophical: concept of atom (Paramanu) demonstrates universality of ideas across cultures.

QUICK REFERENCE EQUATIONS & CONSTANTS

  • ρ=mV\rho = \dfrac{m}{V} (Density)
  • % w/w=w<em>solutew</em>solution×100\%\,w/w = \dfrac{w<em>{solute}}{w</em>{solution}}\times100
  • K=∘C+273.15K = ^\circ C + 273.15; ∘F=95(∘C)+32^\circ F = \dfrac{9}{5}(^\circ C)+32
  • NA=6.022×1023 mol−1N_A = 6.022\times10^{23}\,mol^{-1}
  • M<em>1V</em>1=M<em>2V</em>2M<em>1V</em>1 = M<em>2V</em>2 (Dilution)
  • PV=nRTPV = nRT (Ideal Gas; relevant for molar volume 22.7 L mol−122.7\,L\,mol^{-1} at STP 2023 convention)

EXEMPLAR PROBLEMS (SELECTED)

  • Combustion of 16 g CH<em>4CH<em>4 → 36 g H</em>2OH</em>2O (see stoichiometric steps).
  • 50 kg N<em>2N<em>2 + 10 kg H</em>2H</em>2; limiting reagent H<em>2H<em>2; produces 56.1 kg56.1\,kg NH</em>3NH</em>3.
  • Preparing 0.2 M NaOH from 1 M stock: take 200 mL stock and dilute to 1 L.

COMMON PREFIXES TABLE (SI)

MultiplePrefixSymbol
102410^{24}yottaYY
102110^{21}zettaZZ
101810^{18}exaEE
101510^{15}petaPP
101210^{12}teraTT
10910^{9}gigaGG
10610^{6}megaMM
10310^{3}kilokk
10210^{2}hectohh
10110^{1}decadada
10−110^{-1}decidd
10−210^{-2}centicc
10−310^{-3}millimm
10−610^{-6}microμ\mu
10−910^{-9}nanonn
10−1210^{-12}picopp
10−1510^{-15}femtoff
10−1810^{-18}attoaa
10−2110^{-21}zeptozz
10−2410^{-24}yoctoyy

KEY TAKE-AWAYS

  • Master SI units, prefixes, significant-figure rules.
  • Understand mole concept as bridge between atomic scale & macroscopic lab scale.
  • Balanced chemical equation is essential for quantitative (stoichiometric) predictions.
  • Historical & cultural contributions (esp. Indian) highlight chemistry’s deep roots and ethical considerations for its modern practice.