Comprehensive University Chemistry Master Notes

SOME BASIC CONCEPTS OF CHEMISTRY

Chemistry is the scientific discipline that focuses on the preparation, properties, structure, and reactions of material substances. Major branches include:

Inorganic Chemistry
Organic Chemistry
Physical Chemistry
Analytical Chemistry
Polymer Chemistry
Biochemistry
Medicinal Chemistry
Industrial Chemistry
Hydrochemistry
Electrochemistry
Green Chemistry

Matter and its Classification

Matter occupies space and possesses a definite mass. In the crust of the earth, it exists primarily in three physical states, though seven are recognized:

Solid State: Particles are closely packed in an orderly arrangement. They cannot move freely, resulting in a definite shape and volume.
Liquid State: Particles are close but can move around. They have a definite volume but no definite shape.
Gaseous State: Particles are far apart and move quickly. They have neither a definite shape nor volume, expanding to fill their container completely.
Plasma State
Bose-Einstein Condensate
Fermionic Condensate
Quark-Gluon Plasma

By changing temperature and pressure, the three primary states of matter are interconvertible.

Chemical Composition of Matter

Classification based on chemical composition:

Pure Substances: Contain only one type of particle. They are divided into Elements and Compounds.
Elements: Pure substances containing one type of atom. Robert Boyle (the father of ancient Chemistry) introduced the term. There are 118 known elements, ranging from Hydrogen ( $1H$ ) to Oganesson ( $118Og$ ).
- Monoatomic: Noble gases (Helium, Neon) and metals (Sodium, Potassium).
- Diatomic: Hydrogen, Nitrogen, Oxygen.
- Polyatomic: Phosphorus ( $P_4$ ) and Sulphur ( $S_8$ ).
Compounds: Formed by combining two or more atoms of different elements in a fixed ratio. Their constituents can only be separated by chemical methods. Examples: $CO_2$ , $H_2O$ , $H_2SO_4$ .
Mixtures: Contain multiple particle types. Components can be separated by physical methods (filtration, distillation, crystallization).
- Homogeneous Mixtures: Uniform composition throughout (e.g., solutions, air).
- Heterogeneous Mixtures: Non-uniform composition (e.g., sea water, soil, muddy water).

Properties of Matter

Physical Properties: Measurable without changing composition or identity. Examples: color, melting point, density, and mass.
Chemical Properties: Measured only through a chemical change. Examples: reactivity with acids/bases, combustibility.

Units and Measurements

SI Base Units:

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

Comparison of Parameters:

Mass: Amount of matter present; constant. SI unit is $kg$ .
Weight: Gravitational force acting on a body; variable. SI unit is newton ( $N$ ).
Volume ( $V$ ): Space occupied. SI unit is $m^3$ . Conversions: $1\,m^3 = 10^6\,cm^3$ , $1\,L = 10^3\,cm^3\text{ (mL)}$ , $1\,dm^3 = 10^3\,cm^3$ .
Density ( $d$ ): Mass per unit volume ( $d = \text{mass/volume}$ ). SI unit is $kg/m^3$ ; commonly expressed as $g/cm^3$ .
Temperature ( $T$ ): Degree of hotness/coldness. SI unit is $K$ .
- Fahrenheit conversion: $^{\circ}F = \frac{9}{5}(^{\circ}C) + 32$
- Kelvin conversion: $K = ^{\circ}C + 273.15$

Measurement Quality:

Precision: Closeness of multiple measurements for the same quantity.
Accuracy: Agreement of a value with the true result.

Scientific Notation and Significant Figures

Numbers are represented as $N \times 10^n$ . Shifting decimals left makes $n$ positive; shifting right makes $n$ negative. Example: $368.9 = 3.689 \times 10^2$ .

Significant Figure Rules:

All non-zero digits are significant.
Zeros preceding the first non-zero digit are not significant ( $0.03$ has one).
Zeros between non-zero digits are significant ( $2.005$ has four).
Zeros at the end are significant only if on the right of a decimal ( $0.200\,g$ has three).
Exact numbers (2 balls) have infinite significant figures.
All digits in scientific notation are significant.

Rounding Rules:

If the removed digit is > 5, increase the preceding digit by 1.
If < 5, leave it as is.
If exactly 5, increase by 1 if the preceding digit is odd; leave it if even (e.g., $6.35 \rightarrow 6.4$ , but $6.25 \rightarrow 6.2$ ).

Laws of Chemical Combinations

Law of Conservation of Mass: Proposed by Antoine Lavoisier. Matter is neither created nor destroyed. The mass of reactants equals the mass of products ( $2H_2 + O_2 \rightarrow 2H_2O$ where $4\,g + 32\,g = 36\,g$ ).
Law of Definite Proportions: Proposed by Joseph Proust. A compound always contains the same proportion of elements by weight regardless of source.
Law of Multiple Proportions: Proposed by John Dalton. If two elements form multiple compounds, the masses of one element combining with a fixed mass of the other are in a small whole number ratio (e.g., $16\,g$ vs $32\,g$ Oxygen in $H_2O$ and $H_2O_2$ results in a $1:2$ ratio).
Gay Lussac’s Law of Gaseous Volumes: When gases react, their volumes are in a simple whole number ratio at constant temperature/pressure.
Avogadro’s Law: Equal volumes of all gases at the same temperature/pressure contain an equal number of moles or molecules.

Dalton’s Atomic Theory

Postulates:

Matter consists of minute, indivisible particles called atoms (from 'a-tomio').
Atoms are indestructible.
Atoms of the same element have identical mass and properties.
Atoms combine in fixed mass ratios to form compounds (molecules).

Atomic and Molecular Masses

Atomic Mass (Relative): How many times an atom's mass is greater than $\frac{1}{12}$ th the mass of a $C^{12}$ atom.
Atomic Mass Unit (amu/u): $\frac{1}{12} \times \text{mass of a } C^{12} \text{ atom} = 1.66 \times 10^{-24}\,g = 1.66 \times 10^{-27}\,kg$ .
Average Atomic Mass: Calculated using isotopic abundance. Chlorine isotopes $^{35}Cl$ and $^{^{37}}Cl$ in a $3:1$ ratio yield an average of $35.5\,u$ .
Formula Mass: Used for ionic compounds (e.g., $NaCl$ ) where no discrete molecules exist.

Mole Concept

Mole: Amount of substance containing as many particles as atoms in exactly $12\,g$ of $C^{12}$ .
Avogadro Constant ( $N_A$ ): $6.022 \times 10^{23}$ .
Molar Mass ( $M$ ): Mass of 1 mole in grams ( $n = w/M$ ).
Molar Volume: $22.4\,L\text{ (or 22400 mL)}$ for 1 mole of any gas at STP.

Percentage Composition and Formulas

Percentage Composition of an element = $\frac{\text{Mass of element in compound}}{\text{Molar mass of compound}} \times 100$ .

Empirical Formula: Simplest ratio of elements in a compound (e.g., Glucose = $CH_2O$ ).
Molecular Formula: Actual number of atoms ( $C_6H_{12}O_6$ ).
Relationship: $\text{Molecular Formula} = \text{Empirical Formula} \times n$ , where $n = \text{Molar Mass/Empirical Formula Mass}$ .

Stoichiometry and Solutions

Limiting Reagent: The reactant consumed completely that limits the product formed.
Concentration Terms:
- Mass percent: $\frac{\text{Mass of solute}}{\text{Mass of solution}} \times 100$ .
- Mole fraction ( $\chi_A$ ): $\frac{n_A}{n_A + n_B}$ ; $\sum \chi_i = 1$ .
- Molarity ( $M$ ): Moles of solute per litre of solution. Temperature dependent. Formula: $M_1V_1 = M_2V_2$ .
- Molality ( $m$ ): Moles of solute per kilogram of solvent. Temperature independent.

STRUCTURE OF ATOM

Sub-atomic Particles

Electron: Discovered by J.J. Thomson via cathode ray discharge tubes at low pressure ( $10^{-4}\,mm\,Hg$ ) and high voltage ( $10,000\,V$ ).
- Charge to mass ratio ( $e/m_e$ ): $1.758 \times 10^{11}\,C\,kg^{-1}$ .
- Charge ( $e$ ): $-1.6022 \times 10^{-19}\,C$ (determined by Millikan’s oil drop experiment).
- Mass ( $m_e$ ): $9.1 \times 10^{-31}\,kg$ .
Proton: Discovered by E. Goldstein using perforated cathodes (canal/anode rays). Characteristics depend on the gas used.
Neutron: Discovered by James Chadwick by bombarding beryllium with $\alpha$ -particles: $^9Be + ^4He \rightarrow ^{12}C + ^1n$ .

Atomic Terms

Atomic Number ( $Z$ ): Number of protons or electrons in a neutral atom.
Mass Number ( $A$ ): Protons ( $p$ ) + Neutrons ( $n$ ).
Isotopes: Same $Z$ , different $A$ (e.g., Protium $^1H$ , Deuterium $^2H$ , Tritium $^3H$ ).
Isobars: Same $A$ , different $Z$ (e.g., $^{14}C$ and $^{14}N$ ).

Rutherford’s Nuclear Model

Based on $\alpha$ -particle scattering on gold foil:

Observations: Most particles passed through, some deflected at small angles, 1 in 20,000 bounced back ( $180^{\circ}$ ).
Conclusions: Atoms are mostly empty space; positive charge and mass are in a tiny nucleus ( $10^{-15}\,m$ radius).
Drawbacks: Failed to explain stability (Maxwell's theory predicted electrons should spiral into the nucleus) and electronic structure.

Electromagnetic Radiation (EMR)

Wave Nature: Maxwell proposed that accelerating charges produce alternating electric and magnetic fields.

Velocity ( $c$ ): $3 \times 10^8\,m/s$ .
Relationship: $c = \nu\lambda$ .
Wave number ( $\bar{\nu}$ ): $1/\lambda$ .

Particle Nature: Planck’s Quantum Theory.

Energy is emitted in packets called quanta or photons.
$E = h\nu$ ( $h = 6.626 \times 10^{-34}\,Js$ ).

Photoelectric Effect: Ejection of electrons by light.

Einstein's equation: $h\nu = h\nu_0 + \frac{1}{2}m_ev^2$ , where $h\nu_0$ is the work function.

Atomic Spectra and Bohr’s Model

Hydrogen Line Spectrum: Lyman (UV), Balmer (Visible), Paschen, Brackett, Pfund (IR).
Rydberg Equation: $\bar{\nu} = 109677 \left(\frac{1}{n_1^2} - \frac{1}{n_2^2}\right)\,cm^{-1}$ .

Bohr’s Postulates:

Electrons orbit in stationary states with fixed energy.
Angular momentum is quantized: $m_evr = n\frac{h}{2\pi}$ .
Energy difference: $\Delta E = E_2 - E_1 = h\nu$ .

Quantum Mechanics

de Broglie’s Equation: $\lambda = \frac{h}{p} = \frac{h}{mv}$ .
Heisenberg’s Uncertainty Principle: $\Delta x \cdot \Delta p \ge \frac{h}{4\pi}$ . Significant only for microscopic objects.
Schrödinger Equation: $\hat{H}\psi = E\psi$ . $\psi^2$ represents probability density.

Quantum Numbers:

Principal ( $n$ ): Size and energy of orbit ( $n=1,2,3...$ ).
Azimuthal ( $l$ ): Shape of orbital ( $0$ to $n-1$ ). $l=0(s), 1(p), 2(d), 3(f)$ .
Magnetic ( $m_l$ ): Orientation ( $-l$ to $+l$ ).
Spin ( $s$ ): Orientation ( $+\frac{1}{2}$ or $-\frac{1}{2}$ ).

Filling Rules:

Aufbau Principle: Lowest energy fills first ( $1s < 2s < 2p < 3s...$ ).
Pauli Exclusion Principle: Max 2 electrons per orbital with opposite spins.
Hund’s Rule: Degenerate orbitals fill singly before pairing.
Half/Full Shell Stability: Chromium ( $3d^54s^1$ ) and Copper ( $3d^{10}4s^1$ ).

CLASSIFICATION OF ELEMENTS AND PERIODICITY

Historical Tables

Dobereiner’s Triads: Mid-element mass is average of others.
Newlands’ Law of Octaves: Properties repeat every 8th element (upto Ca).
Mendeleev’s Periodic Law: Properties are periodic functions of atomic weights. Predicted Eka-Aluminium (Galium) and Eka-Silicon (Germanium).

Modern Periodic Table

Modern Periodic Law: Properties are periodic functions of atomic numbers ( $Z$ ). Based on Henry Moseley's work.
Periods (7): Correspondence to principal quantum number ( $n$ ). Longest are VI and VII (32 elements).
Groups (18):
- Group 1: Alkali metals.
- Group 2: Alkaline earth metals.
- Group 3-12: Transition elements.
- Group 17: Halogens.
- Group 18: Noble gases.
Blocks:
- s-block: Reactive metals, low ionization enthalpy.
- p-block: Metals, non-metals, and metalloids.
- d-block: Transition metals; form colored ions, paramagnetic, variable oxidation states.
- f-block: Inner transition elements; Lanthanides and Actinides (radioactive).

Periodic Properties and Trends

Atomic Radius: Covalent, Metallic, or van der Waals. Decreases across a period (effective nuclear charge increases); increases down a group (shells increase).
Ionic Radius: Cations are smaller than parent atoms; Anions are larger.
Isoelectronic Species: Greater positive charge = smaller radius ( $Mg^{2+} < Na^+ < F^- < O^{2-}$ ).
Ionization Enthalpy ( $\Delta_i H$ ): Energy to remove an electron. Increases across a period, decreases down a group.
- Nitrogen ( $2p^3$ ) > Oxygen ( $2p^4$ ) due to half-filled stability.
Electron Gain Enthalpy ( $\Delta_{eg} H$ ): Heat change when adding an electron. More negative across a period.
- Chlorine is more negative than Fluorine due to less electronic repulsion in the larger $n=3$ shell.
Electronegativity: Ability to attract shared electrons. Fluorine (4.0) is highest.
Valency: Combining capacity ( $1 \rightarrow 4 \rightarrow 0$ in periods).

Anomalous Properties and Relationships

Diagonal Relationship: Similarities between Lithium/Magnesium, Beryllium/Aluminium.
Anomalous 2nd Period: Due to small size, high electronegativity, and lack of d-orbitals.

CHEMICAL BONDING AND MOLECULAR STRUCTURE

Octet Rule: Atoms react to attain 8 valence electrons.
Lewis Symbols: Valence electrons shown as dots.

Types of Bonds

Ionic Bond: Formed by complete electron transfer. Favored by low Ionization Enthalpy, high negative Electron Gain Enthalpy, and high Lattice Enthalpy.
Covalent Bond: Formed by mutual sharing. Multiple bonds (double/triple) occur when 2 or 3 pairs are shared.

Bond Parameters and Polarity

Bond Order: Number of bonds between atoms ( $N_2=3, O_2=2, H_2=1$ ).
Resonance: Stabilizes molecules (e.g., $O_3, CO_2, \text{benzene}$ ).
Dipole Moment ( $\mu$ ): $\mu = Q \times r$ . Vector quantity.
- $NH_3$ (1.47 D) > $NF_3$ (0.24 D) because lone pair dipole opposes bond dipoles in $NF_3$ .
Fajans’ Rules: Covalent character in ionic bonds increases with small cation, large anion, and high charge.

VSEPR and Hybridization

VSEPR Postulates: Electron pairs repel ( $LP-LP > LP-BP > BP-BP$ ).
- $NH_3$ : Pyramidal ( $107^{\circ}$ ).
- $H_2O$ : Bent ( $104.5^{\circ}$ ).
Orbital Overlap: Sigma ( $\sigma$ ) is axial (stronger); Pi ( $\pi$ ) is lateral (weaker).
Hybridization:
- $sp^3$ : Tetrahedral ( $109.5^{\circ}$ ), e.g., $CH_4$ .
- $sp^2$ : Trigonal Planar ( $120^{\circ}$ ), e.g., $BCl_3, C_2H_4$ .
- $sp$ : Linear ( $180^{\circ}$ ), e.g., $C_2H_2, BeCl_2$ .
- $sp^3d$ : Trigonal bipyramidal, e.g., $PCl_5$ .
- $sp^3d^2$ : Octahedral, e.g., $SF_6$ .

Molecular Orbital Theory (MOT)

Combination of AOs via LCAO method to form Bonding (BMO) and Anti-bonding (ABMO) orbitals.
Bond Order = $\frac{1}{2}(N_b - N_a)$ .
Para/Diamagnetism: Based on unpaired electrons (e.g., $O_2$ is paramagnetic).

Hydrogen Bonding

Attractive force between H and highly electronegative atoms (F, O, N).

Intermolecular: Between different molecules (leads to high B.P. in water).
Intramolecular: Within the same molecule (e.g., ortho-nitrophenol).

THERMODYNAMICS

Systems and Surroundings

Open System: Exchanges energy and matter.
Closed System: Exchanges only energy.
Isolated System: Exchanges neither.
Extensive Property: Depends on amount of matter (Volume, Enthalpy).
Intensive Property: Independent of amount (Temperature, Density).
State Functions: Depend on initial/final states ( $T, P, V, U, H, S, G$ ).

Laws and Energy

First Law: Energy is conserved. $\Delta U = q + w$ .
- Work ( $w$ ): $-P_{ex}\Delta V$ (Irreversible); $-2.303 nRT \log\left(\frac{V_2}{V_1}\right)$ (Reversible).
Enthalpy ( $H$ ): $H = U + PV$ . Relationship: $\Delta H = \Delta U + \Delta n_g RT$ .
Heat Capacity ( $C$ ): $C = q/\Delta T$ . Mayer’s relation: $C_p - C_v = R$ .
Hess’s Law: Total enthalpy is independent of the reaction path.
Entropy ( $S$ ): Measure of disorder.
- Spontaneity: $\Delta S_{total} = \Delta S_{syst} + \Delta S_{surr} > 0$ .
Gibbs Energy ( $G$ ): Maximum useful work. $\Delta G = \Delta H - T\Delta S$ .
- $\Delta G < 0$ : Spontaneous.
- $\Delta G = 0$ : Equilibrium.
- $\Delta G > 0$ : Non-spontaneous.
- Relationship: $\Delta G^{\circ} = -2.303 RT \log K$ .

EQUILIBRIUM

Physical and Chemical Equilibrium

Henry’s Law: Mass of gas dissolved is proportional to pressure.
Equilibrium Constant ( $K_c$ / $K_p$ ): Ratio of product to reactant concentrations/pressures raised to stoichiometric powers.
Relationship: $K_p = K_c(RT)^{\Delta n}$ .
Le Chatelier’s Principle:
- Conc: Increase reactant, shifts forward.
- Temp: Increase temp, favors endothermic.
- Pressure: Increase pressure, favors direction with fewer moles.
- Catalyst: Reaches equilibrium faster; no change in composition.

Ionic Equilibrium

Arrhenius: Acids give $H^+$ , Bases give $OH^-$ in water.
Bronsted-Lowry: Acids are proton donors, Bases are acceptors. Conjugate pairs differ by one $H^+$ .
Lewis: Acids are electron pair acceptors ( $BF_3$ ), Bases are donors ( $NH_3$ ).
pH: $-\log[H_3O^+]$ . $pH + pOH = 14$ at $298\,K$ .
Buffer Solutions: Resist pH change.
- Henderson-Hasselbalch: $pH = pK_a + \log\left(\frac{[Salt]}{[Acid]}\right)$ .
Solubility Product ( $K_{sp}$ ): Product of molar concentrations of ions in a saturated solution.
- Precipitation occurs if Ionic Product ( $Q_{sp}$ ) > $K_{sp}$ .

REDOX REACTIONS

Oxidation: Loss of electrons; addition of O; increase in Oxidation Number.
Reduction: Gain of electrons; addition of H; decrease in Oxidation Number.
Disproportionation: One element is simultaneously oxidized and reduced (e.g., decomposition of $H_2O_2$ ).
Electrochemical Cell (Daniel Cell): Converts chemical energy to electrical energy.
- Anode: Oxidation ( $Zn \rightarrow Zn^{2+} + 2e^-$ ).
- Cathode: Reduction ( $Cu^{2+} + 2e^- \rightarrow Cu$ ).
Salt Bridge: Maintains electrical neutrality.

ORGANIC CHEMISTRY BASICS

Unique Carbon: Tetravalency, catenation (self-linking), multiple bonds.
Representations: Structural, Condensed, Bond-line, Wedge (3D).
Classification: Acyclic/Aliphatic vs. Cyclic (Alicyclic vs. Aromatic).
Functional Groups: Determine chemical properties (e.g., $-OH$ alcohol, $-CHO$ aldehyde, $-COOH$ acid).
Isomerism:
- Structural: Chain, Position, Functional, Metamerism.
- Stereo: Geometrical ( $cis/trans$ ), Optical.
Reaction Fission:
- Homolysis: Forms Free Radicals ( $3^{\circ} > 2^{\circ} > 1^{\circ}$ stability).
- Heterolysis: Forms Carbocations ( $3^{\circ} > 2^{\circ} > 1^{\circ}$ ) or Carbanions ( $1^{\circ} > 2^{\circ} > 3^{\circ}$ ).
Electronic Effects:
- Inductive ( $I$ ): Permanent shift through $\sigma$ -bonds.
- Electromeric ( $E$ ): Temporary shift in multiple bonds with reagent.
- Resonance ( $R$ ): Interaction of $\pi$ -bonds or lone pairs.
- Hyperconjugation: No-bond resonance; stabilizes carbocations.
Purification: Sublimation, Crystallization, Distillation (Fractional, Reduced Pressure, Steam), Chromatography (TLC, Column, Paper).
Quantitative Analysis: Carbon/Hydrogen (Liebig), Nitrogen (Dumas, Kjeldahl), Halogens/Sulphur (Carius).

HYDROCARBONS

Alkanes ( $C_nH_{2n+2}$ ): Paraffins. Prepared via Wurtz reaction ( $R-X + Na \rightarrow R-R$ ) or Decarboxylation of sodium salts. Halogenation proceeds via free radical mechanism. Conformations: Staggered is more stable than Eclipsed.
Alkenes ( $C_nH_{2n}$ ): Olefins. Markovnikov Rule: negative part of reagent adds to carbon with fewer H atoms. Anti-Markovnikov occurs in peroxide presence. Ozonolysis forms aldehydes/ketones.
Alkynes ( $C_nH_{2n-2}$ ): 1-alkynes are acidic. Addition of water in $Hg^{2+}/H^+$ forms carbonyls. Ethyne undergoes cyclic polymerization above red hot iron at $873\,K$ to form Benzene.
Benzene ( $C_6H_6$ ): High stability due to resonance. Huckel Rule: Aromatic if cyclic, planar, and has $(4n+2)\pi$ electrons. Undergoes Electrophilic Substitution (Nitration, Sulphonation, Friedel-Crafts). substituents are ortho/para directing (activating) or meta directing (deactivating).