12 class

Transition metals exhibit strong metallic bonding due to delocalized d-electrons, leading to dense and stable lattice structures. Most transition metals adopt body-centered cubic (BCC), face-centered cubic (FCC), or hexagonal close-packed (HCP) structures. BCC metals (e.g., vanadium, tungsten) are harder but less dense, while FCC metals (e.g., copper, gold) are more ductile and malleable. HCP metals (e.g., titanium, zinc) offer high strength but limited plasticity. Lattice structures influence m...

In transition metals, atomic and ionic sizes show a gradual decrease across a period due to increasing nuclear charge, which pulls electrons closer. However, the decrease is less pronounced than in main-group elements due to electron shielding by inner d-electrons. Down a group, atomic and ionic sizes increase due to the addition of electron shells, though lanthanoid contraction causes 4d and 5d series elements to be nearly the same in size. Ionic radii decrease with higher oxidation states as e...

Lanthanoid contraction refers to the gradual decrease in atomic and ionic radii of the lanthanide elements (from lanthanum to lutetium) as the atomic number increases. This occurs due to the poor shielding effect of the 4f electrons, which leads to a stronger attraction between the nucleus and outer electrons. As a result, elements after the lanthanides in the periodic table, such as transition metals, also exhibit smaller atomic sizes than expected. Lanthanoid contraction influences chemical pr...

Chemical kinetics questions and numericals focus on reaction rates, rate laws, half-life, and activation energy. Common topics include: 1. Rate Law Calculations – Determining reaction order and rate constant () from experimental data. 2. Integrated Rate Equations – Finding concentration over time for zero, first, and second-order reactions. 3. Half-life () – Calculating decay times for different reaction orders. 4. Arrhenius Equation – Finding activation energy () using...

Collision theory explains how chemical reactions occur based on molecular collisions. For a reaction to take place, reactant molecules must collide with: 1. Sufficient energy (equal to or greater than the activation energy, Ea). 2. Proper orientation to break and form bonds effectively. Only a fraction of collisions, called effective collisions, lead to product formation. The reaction rate increases with higher temperature (more energetic collisions) and concentration (more freque...

A catalyst increases the rate of a reaction by providing an alternative pathway with a lower activation energy (Ea), without being consumed. By lowering Ea, more reactant molecules have enough energy to reach the transition state, increasing the reaction rate. Catalysts do not affect the equilibrium position but help achieve equilibrium faster. Types include homogeneous catalysts (same phase as reactants, e.g., enzyme-catalyzed reactions) and heterogeneous catalysts (different phase, e.g., me...

The frequency factor (A) in the Arrhenius equation (k = A e^(-Ea/RT)) represents the number of times reactant molecules collide with the correct orientation per unit time. It reflects the likelihood of a successful reaction before considering activation energy (Ea). A higher frequency factor indicates more frequent and properly oriented collisions, increasing the reaction rate. Unlike Ea, which depends on energy barriers, A is influenced by molecular size, shape, and reaction mechanism. Though o...

Molecularity of a reaction refers to the number of reactant molecules involved in an elementary step of a reaction. It is always a whole number and can be: Unimolecular: A single molecule undergoes decomposition or rearrangement (e.g., radioactive decay). Bimolecular: Two reactant molecules collide and react (e.g., SN2 reactions). Termolecular: Three molecules collide simultaneously (rare due to low probability). Molecularity is different from reaction order, which is determined ex...

Integrated rate equations describe how reactant concentrations change over time for different reaction orders. Zero-order: (linear decrease, rate independent of concentration). First-order: (exponential decay, rate proportional to concentration). Second-order: (inverse relation, rate proportional to the square of concentration). Each equation helps determine reaction kinetics by analyzing concentration-time data. Half-life () varies with order: constant for first-order, but depe...

In a zero-order reaction, the reaction rate is independent of the concentration of reactants. This means the rate remains constant until the reactant is depleted. The rate law is expressed as rate = k, where k is the rate constant. The integrated rate equation is [A] = [A]₀ - kt, showing a linear decrease in concentration over time. The half-life (t₁/₂) depends on the initial concentration, given by t₁/₂ = [A]₀ / 2k. Zero-order reactions often occur in enzyme-catalyzed processes or s...