Some Basic Concepts of Chemistry – Condensed Notes

Page 1

• Chemistry: science of preparation, properties, structure, reactions of substances.
• Historically driven by search for “philosopher’s stone” & “elixir of life”; modern chemistry shaped in 18th-century Europe.
• Unit goals include: role of chemistry, matter states, element–compound–mixture, scientific notation, significant figures, SI units, laws of chemical combination, atomic/molecular masses, mole concept, %, formulas, stoichiometry.

Page 2

• Ancient India (Rasayan Shastra) mastered metallurgy, glass, dyes, pottery, gypsum cement, faience, alloys (Cu-Sn/As).
• Evidence: Mohenjodaro, Harappa, Maski, Taxila artifacts; controlled kilns, glazed pottery.
• Texts: Arthashastra (salt), Sushruta & Charaka Samhita (alkalies, mineral acids, metal oxides), Rasopanishada (gun-powder), Tamil works (fireworks).
• Scientists: Nagarjuna (mercury compounds, metal extraction), Chakrapani (HgS, soap).
• Ajanta–Ellora paints, Varahamihira (plant polymers for glues) show chemical expertise.

Page 3

• Dyes in Atharvaveda (turmeric, madder etc.).
• Perfumes/cosmetics detailed in Brihat Samhita & Gandhayukti.
• Ink, paper known by 4th–17th centuries.
• Vedic liquors via fermentation (Asavas).
• Acharya Kanda (≈600 BCE): earliest atomic theory – indivisible $\textit{Param=anu}$, varied, combine in pairs/ triplets.
• Charaka Samhita: metal particle size reduction (nanotechnology).
• Decline of alchemy → iatrochemistry → western medicine.
• Modern chemistry in India grew with 19th-century European influence.

Page 4

Importance of Chemistry:
• Central to weather, brain, computers, industry (fertilisers, polymers, drugs, alloys).
• Key drugs: cisplatin, taxol, AZT.
• Enables novel materials (superconductors, conductive polymers, fibres).
• Addresses environmental issues (CFC alternatives, greenhouse gases).
• Foundation: matter concept.

Nature & States of Matter:
• Matter: has mass & occupies space.
• States: solid – fixed $V$ & shape; liquid – fixed $V$, variable shape; gas – variable $V$ & shape.
• Interconvert via T/P changes.

Page 5

Classification of Matter:
• Pure substance (identical particles) vs mixture (variable composition).
• Mixtures: homogeneous (uniform) vs heterogeneous. Components separable by physical methods.
• Pure substances → elements (one type of atom) or compounds (atoms of elements in fixed ratio, separable only chemically).
• Examples: $\text{Na},\, \text{O}<em>2,\, \text{H}</em>2\text{O},\, \text{CO}_2$.

Page 6

Properties:
• Physical – observed without composition change (colour, $T<em>m, T</em>b,\rho$).
• Chemical – composition change (reactivity, combustibility).
Measurement requires number + unit.

SI Base Units:
Length $m$, Mass $kg$, Time $s$, Current $A$, Temperature $K$, Amount $mol$, Luminous intensity $cd$.

Page 7

Derived & Prefixes:
• Common prefixes: $10^{3}=k,\,10^{-3}=m,\,10^{-6}=\mu,\,10^{9}=G$ etc.
Mass vs Weight: mass constant; weight varies with gravity.
Volume: $1\,L=10^{-3}\,m^{3}=10^{3}\,cm^{3}=1\,dm^{3}$.
Density $\rho=\dfrac{\text{mass}}{\text{volume}}$.

Temperature Scales:
$^\circ!F=\tfrac{9}{5}\,^\circ!C+32\qquad K=^\circ!C+273.15$.

Page 8

Uncertainty & Significant Figures:
• Scientific notation N\times10^{n},\;1\le N<10. • Sig-fig rules: non-zero; captive zeros; trailing zeros with decimal; leading zeros not significant. • Rounding: >5 up, <5 no change, =5 round to even.

Dimensional Analysis: use unit factors (e.g. $1\,in=2.54\,cm$) to convert.

Page 9

Laws of Chemical Combination:

1. Conservation of Mass – Lavoisier: mass unchanged.

2. Definite Proportions – Proust: fixed mass ratio in a compound.

3. Multiple Proportions – Dalton: if two elements form multiple compounds, mass ratios are simple integers.

4. Gay-Lussac (Volumes): gaseous reactants/products combine in simple volume ratios (T,P same).

5. Avogadro: equal $V$ of gases at same $T,P$ contain equal molecules (led to diatomic gas concept).

Page 10

Atomic Theory & Masses:
• Dalton (1808): atoms indivisible, identical in element, rearrange in reactions.
• Atomic mass unit $1\,u=\tfrac{1}{12}\,\text{mass of }{}^{12}!C$.
• Average atomic mass weighted by isotopic abundance.
• Molecular mass = sum of atomic masses; formula mass for ionic lattices.

Mole & Molar Mass:
• $1\,mol=6.022\times10^{23}$ entities ($N_A$).
• Molar mass (g mol$^{-1}$) numerically equals atomic/molecular mass in $u$.

Percentage Composition:
$\%\,\text{element}=\dfrac{\text{mass of element}}{\text{molar mass}}\times100$.

Empirical vs Molecular Formula:
• From mass % → moles → simplest ratio (empirical).
• $\text{n} = \dfrac{M<em>{\text{molar}}}{M</em>{\text{emp}}}$, molecular formula = (empirical)$\times n$.

Stoichiometry:
• Balanced equation gives mole ratios.
• Limiting reagent: reactant consumed first; determines product amount.
• Solution concentrations:
– mass % (w/w)
– mole fraction $x=\dfrac{n<em>i}{\sum n}$ – molarity $M=\dfrac{n</em>{\text{solute}}}{V<em>{\text{solution}}(L)}$ – molality $m=\dfrac{n</em>{\text{solute}}}{\text{mass solvent (kg)}}$.

Page 11 onwards (Summary & Exercises)

• Summary reiterates: matter classification, measurement (SI, sig-fig), laws, atomic masses, mole, formulas, stoichiometry.
• Practice problems reinforce calculations on molar mass, % composition, limiting reagent, concentration conversions.