Chemistry Half Yearly - Class 12 Haryana Board Chemistry

Chemistry Class 12: Half-Yearly Syllabus Rapid Revision

Exam Pattern and Syllabus Structure

The half-yearly exam consists of 35 compulsory questions divided into five sections: A, B, C, D, and E. The syllabus covers approximately 60% of the total content, focusing on six main chapters: Solutions, Electrochemistry, Chemical Kinetics, d- and f-Block Elements, Coordination Compounds, and Haloalkanes & Haloarenes.

• Section A: 18 objective-type questions, 1 mark each.
• Section B: 7 very short answer questions, 2 marks each.
• Section C: 5 short answer questions, 3 marks each.
• Section D: 2 case study questions, 4 marks each.…

Chemistry Class 12: Half-Yearly Syllabus Rapid Revision

Exam Pattern and Syllabus Structure

The half-yearly exam consists of 35 compulsory questions divided into five sections: A, B, C, D, and E. The syllabus covers approximately 60% of the total content, focusing on six main chapters: Solutions, Electrochemistry, Chemical Kinetics, d- and f-Block Elements, Coordination Compounds, and Haloalkanes & Haloarenes.

• Section A: 18 objective-type questions, 1 mark each.
• Section B: 7 very short answer questions, 2 marks each.
• Section C: 5 short answer questions, 3 marks each.
• Section D: 2 case study questions, 4 marks each.
• Section E: 3 long answer questions, 5 marks each.
• Chapters for Half-Yearly: Solutions, Electrochemistry, Chemical Kinetics, d- and f-Block Elements, Coordination Compounds, Haloalkanes & Haloarenes.

Solutions: Key Concepts and Formulas

Solutions chapter covers colligative properties, Henry’s law, Raoult’s law, and related numericals. The van’t Hoff factor ($i$) is crucial for understanding association/dissociation in solutions.

• Van’t Hoff Factor ($i$):
  • For non-electrolytes (e.g., urea): $i = 1$
  • $i < 1$ indicates association; $i > 1$ indicates dissociation.
• Molarity Calculation:
  • $M = \frac{\text{Number of moles of solute}}{\text{Volume of solution in liters}}$
• Mixing Solutions:
  • Total moles: $n_1 + n_2$
  • Total volume: sum of individual volumes (convert to liters).
• Henry’s Law:
  • Solubility of a gas in a liquid is directly proportional to the partial pressure of the gas.
  • $S \propto P$; $P = K_H x$ (where $x$ is mole fraction).
  • $K_H$ increases with temperature; solubility decreases with temperature.
• Raoult’s Law:
  • Vapor pressure of a component is directly proportional to its mole fraction.
  • $P = P_A^0 x_A + P_B^0 x_B$ for binary solutions.
• Colligative Properties:
  • Relative lowering of vapor pressure, elevation of boiling point, depression of freezing point, osmotic pressure.
  • Depression in freezing point: $\Delta T_f = K_f \cdot m$
  • Elevation in boiling point: $\Delta T_b = K_b \cdot m$
  • Osmotic pressure: $\pi = CRT$

Electrochemistry: Principles and Calculations

Electrochemistry includes concepts of electrode potential, electrochemical cells, Nernst equation, conductance, and Faraday’s laws.

• Electrode Potential Depends On:
  • Nature of metal, temperature, concentration of solution.
• Electrochemical Series:
  • More negative standard electrode potential = stronger reducing agent.
• Cell Representation:
  • Anode | Anode solution || Cathode solution | Cathode
  • Oxidation at anode (loss of electrons), reduction at cathode (gain of electrons).
• Salt Bridge:
  • Completes the circuit, maintains electrical neutrality, prevents mixing of solutions.
• Nernst Equation:
  • $E_{cell} = E_{cell}^0 - \frac{0.0591}{n} \log Q$
• Gibbs Free Energy:
  • $\Delta G^0 = -nFE_{cell}^0$
• Faraday’s Laws:
  • First Law: Mass deposited $\propto$ charge passed.
  • $m = ZQ$ where $Z$ is electrochemical equivalent.
  • Second Law: Masses deposited by same charge are proportional to equivalent masses.
• Conductivity:
  • Conductivity ($\kappa$) decreases with dilution.
  • Molar conductivity increases with dilution.
  • $\Lambda_m = \frac{\kappa \times 1000}{C}$

Chemical Kinetics: Order, Molecularity, and Rate Laws

Chemical kinetics focuses on reaction rates, order, molecularity, half-life, and related calculations.

• Order vs. Molecularity:
  • Order can be zero, fractional, or integral; molecularity is always a whole number.
• Half-Life Formulas:
  • Zero order: $t_{1/2} = \frac{[A]_0}{2k}$
  • First order: $t_{1/2} = \frac{0.693}{k}$
• First Order Rate Law:
  • $k = \frac{2.303}{t} \log \frac{[A]_0}{[A]}$
• Effect of Concentration:
  • For $n$th order: Rate $\propto [A]^n$
  • Increasing concentration affects rate as per order.
• Activation Energy:
  • Determined by rate constants at two different temperatures.
• Rate Determining Step:
  • The slowest step controls the overall reaction rate.

d- and f-Block Elements: Properties and Trends

This section covers transition elements, their properties, and trends in the periodic table.

• Transition Elements:
  • Exhibit variable oxidation states, colored compounds (due to $d$-$d$ transitions), paramagnetism (due to unpaired electrons).
  • Zn, Cd, Hg are not typical transition elements (fully filled $d$-orbitals).
• Interstital Compounds:
  • Small atoms (H, B, C, N) occupy spaces in metal lattices.
• Lanthanide Contraction:
  • Poor shielding by $4f$ electrons leads to similar radii for Zr and Hf.
• Common Oxidation States:
  • Lanthanides: +3 is most common.
• Preparation of Potassium Dichromate:
  • Chromite $\rightarrow$ sodium chromate $\rightarrow$ sodium dichromate $\rightarrow$ potassium dichromate (via KCl).
• Preparation of KMnO$_4$:
  • Pyrolusite ($MnO_2$) with KOH and O$_2$ forms potassium manganate, then oxidized to potassium permanganate.

Coordination Compounds: Structure, Nomenclature, and Isomerism

Coordination compounds involve central metal atoms/ions surrounded by ligands, with specific rules for naming and structure.

• Coordination Number:
  • Number of donor atoms directly bonded to central metal.
  • For $[Cr(C_2O_4)_3]^{3-}$, coordination number is 6 (oxalate is bidentate).
• Ligands:
  • Unidentate: one donor atom (e.g., $NH_3$, $Cl^-$).
  • Bidentate: two donor atoms (e.g., ethylenediamine).
  • Ambidentate: can coordinate through two different atoms (e.g., $NO_2^-$, $SCN^-$).
• Nomenclature:
  • Name cation first, then anion.
  • Ligands in alphabetical order, prefixes for multiples (di-, tri-, etc.).
  • Oxidation state of metal in Roman numerals.
• Isomerism:
  • Geometrical isomerism possible in square planar and octahedral complexes.
  • Ionization isomerism: exchange of ions inside and outside coordination sphere.
• Crystal Field Splitting:
  • $d$-orbitals split into $t_{2g}$ and $e_g$ in presence of ligands.
  • Crystal field splitting energy ($\Delta$) is the energy difference between these sets.
  • For tetrahedral: $\Delta_{tet} = \frac{4}{9} \Delta_{oct}$

Haloalkanes and Haloarenes: Reactions and Properties

This chapter covers the structure, reactions, and properties of haloalkanes and haloarenes, including mechanisms and isomerism.

• Reactions:
  • Chlorination of toluene (with $FeCl_2$) gives ortho and para chlorotoluene; para is major due to less steric hindrance.
  • Woods reaction: Alkyl halides with sodium in dry ether yield alkanes.
  • Gabriel phthalimide synthesis: Preferred for primary amines as it avoids secondary/tertiary amines.
• SN2 Mechanism:
  • Faster in primary halides, less steric hindrance, better leaving group (iodide $>$ bromide $>$ chloride).
• Boiling Points:
  • For same alkyl group: $RI > RBr > RCl > RF$
  • Para isomers have higher melting points due to symmetry.
• Dipole Moment:
  • Highest in $CH_3Cl$ among methyl halides.
• Coupling Reaction:
  • Diazotized aromatic amines react with activated aromatic compounds to form azo compounds (brightly colored, used in dyes).
• Gattermann Reaction:
  • Formation of haloarenes from diazonium salts with $Cu$ powder and $HX$.

Case Studies and Application-Based Questions

Case studies often focus on real-life applications of chemical principles, especially from the Solutions chapter.

• Henry’s Law Applications:
  • Scuba divers: Increased pressure increases gas solubility; rapid ascent causes nitrogen bubbles (bends).
  • Soft drinks: Bottled under high pressure to dissolve more $CO_2$.
• Osmosis and Osmotic Pressure:
  • RBCs in isotonic solution: No net movement, shape unchanged.
  • Osmotic pressure: $\pi = CRT$; higher for electrolytes due to van’t Hoff factor.
  • Reverse osmosis: Applied pressure greater than osmotic pressure forces solvent from solution to pure solvent side (used in desalination).
• Raoult’s Law and Azeotropes:
  • Ideal solutions obey Raoult’s law; non-ideal show positive/negative deviation.
  • Azeotropes: Binary mixtures with same composition in liquid and vapor phase, cannot be separated by fractional distillation.
  • Minimum boiling azeotropes show positive deviation; maximum boiling show negative deviation.

Important Numerical and Assertion-Reason Questions

Numerical problems and assertion-reasoning are frequent and test conceptual clarity and calculation skills.

• Numerical Types:
  • Calculation of molarity, molality, molecular mass from colligative properties.
  • Cell potential, Gibbs free energy, and number of electrons using Faraday’s laws.
  • Rate constant and half-life from kinetic data.
• Assertion-Reasoning:
  • Understand the relationship between colligative properties and solution behavior.
  • Distinguish between correct statements and correct explanations.
  • Recognize the effect of non-volatile solutes on boiling and freezing points.

Key Pitfalls and Exam Tips

Careful attention to units, conversions, and the specific wording of questions is essential. Many questions are repeated or have similar patterns across years.

• Units: Always convert volumes to liters, masses to kilograms as required.
• Formulas: Use the exact formula as per the question’s requirement; avoid unnecessary steps.
• Conceptual Traps: Do not confuse order with molecularity; remember the difference between conductivity and molar conductivity.
• Repetition: Focus on previous years’ questions for pattern recognition and efficient revision.