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Semester 3: Chemical Dynamics & Coordination Chemistry
Chemical Kinetics: Reaction rates, order of reaction, experimental methods, collision and transition state theories
Chemical Kinetics
Reaction Rates
Reaction rate refers to the change in concentration of reactants or products over time. It can be influenced by various factors such as temperature, concentration, surface area, and the presence of catalysts. The rate is often expressed in mol/L/s.
Order of Reaction
The order of reaction indicates how the rate of reaction is affected by the concentration of reactants. It can be zero, first, second, or fractional order. The overall order is the sum of the individual orders with respect to each reactant.
Experimental Methods
Experimental methods to determine reaction rates include measuring the change in concentration, pressure, or volume over time. Common techniques are colorimetry for color change reactions, gas evolution measurements, and pH monitoring.
Collision Theory
Collision theory states that for a reaction to occur, reactant particles must collide with sufficient energy and proper orientation. The frequency and energy of collisions determine the reaction rate.
Transition State Theory
Transition state theory suggests that reactants must pass through a high-energy transition state before forming products. The energy required to reach this state is known as the activation energy, which affects the rate of reactions.
Chemical Equilibrium: Equilibrium constants, thermodynamic derivation, Le-Chatelier's principle
Chemical Equilibrium
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Equilibrium constant is a numerical value that expresses the ratio of concentrations of products to reactants at equilibrium.
K = [C]^c [D]^d / [A]^a [B]^b
1. Kc (concentration-based) 2. Kp (pressure-based)
Indicates the extent of reaction and direction based on concentration or pressure.
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At equilibrium, the change in Gibbs free energy, ΔG, is zero indicating a stable state.
ΔG = ΔG° + RT ln(Q) Where Q is the reaction quotient.
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Le-Chatelier's principle states that if an external stress is applied to a system at equilibrium, the system shifts to counteract that stress.
1. Concentration change 2. Temperature change 3. Pressure change
Helps predict the direction of shift when conditions are altered.
Phase Equilibrium: Gibbs phase rule, phase equilibria for one and two component systems
Phase Equilibrium
The Gibbs phase rule relates the number of phases in a system to the number of components and degrees of freedom. It is expressed as F = C - P + 2, where F is the degrees of freedom, C is the number of components, and P is the number of phases.
Used to predict the number of variables that can be changed independently in a system at equilibrium.
In a single component system, phase equilibrium can occur between solid, liquid, and gas phases.
The phase diagram of water showing the boundaries between solid, liquid, and gas phases at various temperatures and pressures.
Involving two components, phase diagrams can illustrate the interactions between two substances across various phases.
Phase diagrams of binary mixtures, such as the water-ethanol system, showcasing regions of liquid, vapor, and solid phases.
Understanding phase equilibrium is critical in various chemical processes, including distillation, crystallization, and the design of reactors.
Phase equilibrium knowledge aids in predicting and controlling the outcomes of chemical reactions and separations in industrial applications.
Kinetic theories of gases: Postulates, deviation from ideality, Van der Waals equation, molecular velocities
Kinetic theories of gases
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Gases consist of a large number of molecules that are in constant random motion.
The volume of the gas molecules is negligible compared to the volume of the container.
Collisions between gas molecules and between molecules and the walls of the container are perfectly elastic.
There are no attractive or repulsive forces between the molecules, except during collisions.
The average kinetic energy of gas molecules is directly proportional to the absolute temperature of the gas.
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Real gases deviate from ideal behavior under high pressure and low temperature conditions.
Factors like molecular size and intermolecular forces lead to non-ideal behavior.
The greater the complexity of a gas molecule, the more it deviates from ideality.
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Van der Waals modified the ideal gas law to account for the volume occupied by gas molecules and the attractive forces between them.
The equation is expressed as (P + a(n/V)²)(V - nb) = nRT, where a and b are Van der Waals constants.
This equation better predicts the behavior of real gases compared to the ideal gas law.
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The average molecular velocity can be derived from the kinetic theory, expressed as v_avg = √(3RT/M), where R is the gas constant, T is the temperature, and M is the molar mass.
Root mean square velocity (v_rms) is given by v_rms = √(3RT/M).
Molecular velocities help understand diffusion and effusion processes in gases.
Liquid State: Intermolecular forces, liquid crystals, gels
Liquid State: Intermolecular forces, liquid crystals, gels
Intermolecular Forces in Liquids
Intermolecular forces are attractions between molecules that play a critical role in determining the physical properties of liquids. Unlike in solids, where molecules are held in fixed positions, the arrangement of molecules in liquids is more dynamic. The primary types of intermolecular forces in liquids include: dipole-dipole interactions, hydrogen bonding, and London dispersion forces. The strength and nature of these forces affect boiling points, viscosity, and surface tension.
Liquid Crystals
Liquid crystals are substances that exhibit properties between those of conventional liquids and solid crystals. They can flow like liquids but have some degree of order in their molecular arrangement. Liquid crystals are categorized into three main phases: nematic, smectic, and cholesteric. These materials are widely used in display technologies, such as LCD screens, due to their ability to manipulate light.
Gels
Gels are a mixture of liquids and solids that exhibit unique properties. They are primarily characterized by a three-dimensional network structure that traps liquid. The solid component can be a polymer or silica, while the liquid phase allows for flexibility. Gels find applications in various fields including pharmaceuticals, food products, and cosmetic industries. The behavior of gels often depends on temperature, composition, and the nature of the interactions between the components.
Coordination Chemistry: Werner’s theory, ligands, coordination numbers, isomerism in complexes
Coordination Chemistry
Werner's Theory
Proposed by Alfred Werner, this theory emphasized the formation of coordination compounds and introduced key concepts such as coordination number and the distinction between primary and secondary valencies.
Ligands
Ligands are ions or molecules that can donate electron pairs to a central metal atom to form a coordination complex. They can be classified as monodentate, bidentate, and polydentate based on the number of bonds they can form.
Coordination Numbers
The coordination number refers to the number of ligand atoms that are bonded to the central metal atom in a complex. Common coordination numbers include 2, 4, and 6, which correspond to different geometries such as linear, tetrahedral, and octahedral.
Isomerism in Complexes
Coordination compounds can exhibit various types of isomerism, including structural isomerism (where the arrangement of atoms differs) and stereoisomerism (where the spatial arrangement differs). Geometric and optical isomers are key forms of stereoisomerism.
Theories of Coordination Chemistry: Metal-ligand bonding, crystal field theory, thermodynamics and kinetics of complexes
Theories of Coordination Chemistry
Metal-ligand Bonding
Metal-ligand bonding involves the interaction between a central metal atom and surrounding ligands. This bonding can be categorized into ionic, covalent, and coordinate covalent types. Ligands can be monodentate, bidentate, or polydentate depending on the number of donor atoms they use to attach to the metal.
Crystal Field Theory
Crystal field theory explains the electronic structure of transition metal complexes. It considers the effect of ligand fields on the energies of d-orbitals. Splitting of these orbitals occurs due to the presence of ligands, leading to distinct energy levels, which influences color, magnetism, and stability of the complexes.
Thermodynamics of Complexes
The thermodynamics of coordination complexes involves understanding the stability and formation constants. Factors affecting stability include ligand, metal identity, and steric factors. The free energy change during complex formation is essential for determining spontaneity and equilibrium.
Kinetics of Complexes
Kinetics in coordination chemistry evaluates the rates of reactions involving metal-ligand complexes. These rates can be influenced by factors such as ligand substitutions and inner-sphere vs outer-sphere mechanisms. Understanding kinetics helps in predicting reaction pathways and intermediates.
Inorganic Spectroscopy and Magnetism: Electronic spectra, selection rules, magnetic properties of transition metal complexes
Inorganic Spectroscopy and Magnetism
Electronic Spectra
Electronic spectra refer to the transitions between electronic energy levels of atoms or molecules. In the context of transition metal complexes, the spectra arise from d-d transitions and charge transfer transitions. The electronic spectra provide information about the oxidation states, coordination geometries, and ligand fields in transition metal complexes.
Selection Rules
Selection rules dictate the allowed transitions between energy levels in electronic spectroscopy. For d-d transitions in transition metal complexes, the primary selection rule is that the spin multiplicity must not change, meaning that transitions between states of different total spin are generally forbidden. Additionally, Laporte's rule states that for transitions to be allowed, there must be a change in symmetry.
Magnetic Properties of Transition Metal Complexes
The magnetic properties of transition metal complexes are determined by the arrangement of electrons in their d orbitals. The two main types of magnetism observed are paramagnetism, which arises from unpaired electrons, and diamagnetism, which arises when all electrons are paired. The magnetic moment can be quantitatively assessed using the effective magnetic moment formula, often denoted as μ_eff.
