Chemistry Midterm Marathon Half - Class 11 ISC Chemistry

Structure of Atom, Basic Concepts, Periodicity, and Chemical Bonding

Subatomic Particles and Atomic Models

Atoms are composed of three fundamental subatomic particles: electrons (negative charge), protons (positive charge), and neutrons (neutral). The discovery of these particles and their arrangement within the atom led to the development of various atomic models.

• Electron: Charge $-1.6 \times 10^{-19}$ C, mass $9.1 \times 10^{-31}$ kg, relative charge $-1$.
• Proton: Charge $+1.6 \times 10^{-19}$ C, mass $1.67 \times 10^{-27}$ kg, relative charge $+1$.
• Neutron: No charge, mass nearly equal to proton.

Discoveries:

• Electron: J.J. Thomson via discharge tube (cathode…

Structure of Atom, Basic Concepts, Periodicity, and Chemical Bonding

Subatomic Particles and Atomic Models

Atoms are composed of three fundamental subatomic particles: electrons (negative charge), protons (positive charge), and neutrons (neutral). The discovery of these particles and their arrangement within the atom led to the development of various atomic models.

• Electron: Charge $-1.6 \times 10^{-19}$ C, mass $9.1 \times 10^{-31}$ kg, relative charge $-1$.
• Proton: Charge $+1.6 \times 10^{-19}$ C, mass $1.67 \times 10^{-27}$ kg, relative charge $+1$.
• Neutron: No charge, mass nearly equal to proton.

Discoveries:

• Electron: J.J. Thomson via discharge tube (cathode rays), $e/m = 1.7 \times 10^{11}$ C/kg (constant for cathode rays).
• Proton: E. Goldstein via perforated cathode (anode/canal rays), $e/m$ for hydrogen $= 9.58 \times 10^7$ C/kg.
• Neutron: Chadwick using beryllium sheet and alpha particles.

Atomic Models:

• Thomson’s Model: Atom is a positively charged sphere with electrons embedded (plum pudding/watermelon model). Explained neutrality.
• Rutherford’s Model: Atom has a small, dense, positively charged nucleus; electrons revolve around nucleus. Could not explain atomic stability or line spectra.
  • Atomic radius to nuclear radius ratio: $10^5$.
  • Atomic volume to nuclear volume ratio: $10^{15}$.
• Bohr’s Model: Electrons revolve in fixed, discrete orbits (stationary states) with quantized angular momentum ($mvr = n\frac{h}{2\pi}$). Applicable only to single-electron systems (hydrogen-like atoms).

Isotopes, Isobars, and Related Terms

Atoms can be classified based on their atomic and mass numbers, as well as their electron count.

• Isotopes: Same atomic number, different mass number (e.g., $^{12}C$, $^{13}C$, $^{14}C$).
• Isobars: Same mass number, different atomic number (e.g., Argon and Calcium).
• Isotones: Same number of neutrons, different atomic and mass numbers.
• Isodiaphers: Same value of $A - 2Z$.
• Isoelectronic Species: Same number of electrons (e.g., $Na^+$, $Mg^{2+}$, $F^-$, $O^{2-}$ all have 10 electrons).
• Isosters: Isoelectronic species with the same number of atoms (e.g., $CO_2$ and $NO_2$).

Electromagnetic Radiation and Spectra

Electromagnetic radiation exhibits both wave and particle nature, explaining phenomena like black body radiation and photoelectric effect.

• Photoelectric Effect: Emission of electrons from metal surface when struck by radiation of minimum threshold frequency.
  • Formula: $h\nu = h\nu_0 + \frac{1}{2}mv^2$
• Black Body Radiation: Ideal body absorbs and emits all frequencies uniformly; temperature and wavelength are inversely related.
• Electromagnetic Spectrum Order: Cosmic rays, gamma rays, X-rays, UV, visible, IR, microwave, radio (frequency decreases, wavelength increases).

Spectra:

• Continuous Spectrum: White light through prism gives merged bands (VIBGYOR).
• Line Spectrum: Discrete lines for elements; explained by Bohr’s model.
• Balmer Formula (Visible Range): $\bar{\nu} = R\left(\frac{1}{n_1^2} - \frac{1}{n_2^2}\right)$, $R = 109,676$ cm$^{-1}$.

Quantum Mechanical Model and Quantum Numbers

The quantum mechanical model describes electrons in terms of probability, not fixed paths.

• Schrödinger Wave Equation: Solutions yield three quantum numbers (principal $n$, azimuthal $l$, magnetic $m$); spin quantum number ($s$) is not from this equation.
• Principal Quantum Number ($n$): Indicates shell/energy level; $n = 1,2,3...$ (cannot be zero or negative).
• Azimuthal Quantum Number ($l$): Subshell/shape; $l = 0$ to $n-1$; $l=0$ (s), $1$ (p), $2$ (d), $3$ (f).
• Magnetic Quantum Number ($m$): Orientation; $-l$ to $+l$.
• Spin Quantum Number ($s$): $+\frac{1}{2}$ or $-\frac{1}{2}$; two electrons per orbital with opposite spins.

Nodes:

• Radial Nodes: $n-l-1$
• Angular Nodes: $l$
• Total Nodes: $n-1$

Electronic Configuration and Filling Rules

Electrons fill orbitals in order of increasing energy, following specific rules.

• Aufbau Principle: Electrons fill lowest energy orbitals first.
• Order of Filling: $1s$, $2s$, $2p$, $3s$, $3p$, $4s$, $3d$, $4p$, $5s$, $4d$, $5p$, $6s$, etc.
• Maximum Electrons: $s$ (2), $p$ (6), $d$ (10), $f$ (14).
• Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers.
• Hund’s Rule: Orbitals of the same subshell are singly occupied before pairing.

Degenerate Orbitals: Orbitals with the same energy (e.g., $2s$ and $2p$ in hydrogen).

Fundamental Laws of Chemistry

Several laws govern chemical reactions and composition.

• Law of Conservation of Mass: Mass of reactants equals mass of products (Lavoisier).
• Law of Constant Composition: A compound always contains the same elements in the same proportion by mass (Proust).
• Law of Multiple Proportions: When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in simple whole number ratios.
• Law of Reciprocal Proportion: When two elements combine separately with a fixed mass of a third element, the ratio of the masses in which they combine is either the same or a simple multiple.
• Gay-Lussac’s Law of Combining Volumes: Gases combine in simple whole number ratios by volume under identical conditions.
• Avogadro’s Law: Equal volumes of gases at the same temperature and pressure contain equal numbers of molecules.

Mole Concept and Calculations

The mole is a counting unit for particles, atoms, molecules, etc.

• 1 mole: $6.022 \times 10^{23}$ entities (Avogadro’s number).
• Number of moles:
  • $n = \frac{\text{Given mass}}{\text{Molar mass}}$
  • $n = \frac{\text{Number of particles}}{N_A}$
  • For gases at STP: $n = \frac{\text{Volume}}{22.4\,L}$
  • From ideal gas law: $n = \frac{PV}{RT}$ (for gases only)
• Conversions:
  • Moles $\leftrightarrow$ Mass: Multiply/divide by molar mass.
  • Moles $\leftrightarrow$ Number: Multiply/divide by Avogadro’s number.
  • Moles $\leftrightarrow$ Volume (gas at STP): Multiply/divide by 22.4 L.

Concentration Terms

Concentration expresses the amount of solute in a given amount of solvent or solution.

• Mass Percent ($m/m$): $\frac{\text{Mass of solute}}{\text{Mass of solution}} \times 100$
• Volume Percent ($v/v$): $\frac{\text{Volume of solute}}{\text{Volume of solution}} \times 100$
• Mass/Volume Percent ($m/v$): $\frac{\text{Mass of solute}}{\text{Volume of solution}} \times 100$
• Parts Per Million (ppm): $\frac{\text{Mass of solute}}{\text{Mass of solution}} \times 10^6$
• Mole Fraction ($X_A$): $\frac{\text{Moles of A}}{\text{Total moles}}$; $X_A + X_B = 1$
• Molarity ($M$): $\frac{\text{Moles of solute}}{\text{Volume of solution in L}}$
• Molality ($m$): $\frac{\text{Moles of solute}}{\text{Mass of solvent in kg}}$
• Normality ($N$): $\frac{\text{Gram equivalents of solute}}{\text{Volume of solution in L}}$
  • Gram equivalent $= \frac{\text{Given mass}}{\text{Equivalent mass}}$
  • Equivalent mass $= \frac{\text{Molar mass}}{n\text{ factor}}$

Dilution and Mixing:

• $M_1V_1 = M_2V_2$ (dilution)
• $M = \frac{M_1V_1 + M_2V_2}{V_1 + V_2}$ (mixing)

Periodic Table and Periodicity

Elements are arranged in the periodic table based on atomic number, grouped into periods and blocks.

• Modern Periodic Law: Properties of elements are periodic functions of their atomic numbers.
• Table Structure: 18 groups (vertical), 7 periods (horizontal), 4 blocks (s, p, d, f).
• Blocks:
  • s-block: Groups 1-2
  • p-block: Groups 13-18
  • d-block: Groups 3-12 (transition elements)
  • f-block: Lanthanides and actinides (inner transition elements)
• Group Number: For s-block, equals number of valence electrons; for p-block, group number = valence electrons + 10.
• Period Number: Equals principal quantum number of valence shell.

Periodic Properties:

• Atomic Size: Increases down a group, decreases across a period.
• Ionization Energy: Increases across a period, decreases down a group.
• Electronegativity: Increases across a period, decreases down a group.
• Metallic Character: Decreases across a period, increases down a group.

Special Trends:

• Lanthanide Contraction: 4d and 5d series have nearly equal atomic sizes due to poor shielding by 4f electrons.
• Transition Contraction: Atomic size of gallium is less than expected due to poor shielding by d-electrons.

Chemical Bonding: Types, Parameters, and Theories

Chemical bonds are attractive forces holding atoms together in compounds.

• Ionic Bond: Complete transfer of electrons; favored by low ionization energy (metals), high electron gain enthalpy (non-metals), and high lattice energy.
  • Lattice energy ($U$): Directly proportional to product of charges, inversely to distance between ions.
• Covalent Bond: Sharing of electrons; explained by Lewis structures and valence bond theory.
• Bond Parameters:
  • Bond Length: Average distance between nuclei of bonded atoms; increases with atomic size, decreases with bond multiplicity.
  • Bond Angle: Angle between two bonds from the same atom.
  • Bond Enthalpy: Energy required to break one mole of bonds; higher for multiple bonds and smaller atoms.
• Resonance: Delocalization of electrons; actual structure is a hybrid of contributing forms.
• Hydrogen Bonding: Weak attraction between H and electronegative atom (F, O, N); affects boiling/melting points and solubility.

Molecular Geometry and Hybridization

Molecular shape is determined by the arrangement of electron pairs around the central atom.

• VSEPR Theory: Shape depends on number of bond pairs and lone pairs.
  • 2 pairs: Linear
  • 3 pairs: Trigonal planar
  • 4 pairs: Tetrahedral
  • 5 pairs: Trigonal bipyramidal
  • 6 pairs: Octahedral
• Steric Number: Sum of bond pairs and lone pairs; determines hybridization.
  • 2: $sp$
  • 3: $sp^2$
  • 4: $sp^3$
  • 5: $sp^3d$
  • 6: $sp^3d^2$
  • 7: $sp^3d^3$
• Hybridization: Mixing of atomic orbitals to form new hybrid orbitals of equal energy.

Molecular Orbital Theory and Bond Order

Molecular orbital theory explains bonding in terms of molecular orbitals formed by the combination of atomic orbitals.

• Bonding and Antibonding Orbitals: Constructive overlap forms bonding orbitals; destructive overlap forms antibonding orbitals.
• Filling Order (up to 14 electrons): $\sigma_{1s}$, $\sigma^*_{1s}$, $\sigma_{2s}$, $\sigma^*_{2s}$, $\pi_{2p}$, $\sigma_{2p}$, $\pi^*_{2p}$, $\sigma^*_{2p}$
• Bond Order: $\frac{\text{Number of bonding electrons} - \text{Number of antibonding electrons}}{2}$
  • Bond order $= 0$ or negative: molecule is unstable.
• Magnetic Properties: Unpaired electrons $\rightarrow$ paramagnetic; all paired $\rightarrow$ diamagnetic.
• Bond Order and Properties: Higher bond order means stronger bond and shorter bond length.

Key Exceptions and Exam-Focused Points

Certain exceptions and special cases are frequently tested.

• Half-filled and fully-filled subshells: Extra stability due to symmetry and exchange energy.
• Lanthanide and transition contraction: Atomic size anomalies due to poor shielding.
• Ionization energy jumps: Due to stable electronic configurations (e.g., Be, N).
• **Hydrogen and fluorine never central atom in Lewis structures.
• Formal Charge Calculation: $FC = V - L - \frac{B}{2}$, where $V$ = valence electrons, $L$ = lone pair electrons, $B$ = bonding electrons.
• Assignment: Write electronic configurations for atomic numbers 21 to 30.

Table: Maximum Electrons in Subshells

• **Bond order, geometry, and hybridization are directly linked to exam questions; memorize filling orders and steric number rules.