CHEM1982 Winter 2025 PDF - General Applied Chemistry

CHEM1982 (Winter 2025) General Applied Chemistry Part I Gases; Thermochemistry; Chemical kinetics and Chemical Equilibrium Instructor: Dr. Mason Lawrence © McGraw Hill LLC Exercise Predicting shifts in Equilibrium For the reaction in which direction will the equilib...

CHEM1982 (Winter 2025) General Applied Chemistry Part I Gases; Thermochemistry; Chemical kinetics and Chemical Equilibrium Instructor: Dr. Mason Lawrence © McGraw Hill LLC Exercise Predicting shifts in Equilibrium For the reaction in which direction will the equilibrium shift when (a) Cl2(g) is removed, (b) the temperature is decreased, (c) the volume is increased, (d) PCl3(g) is added? © McGraw Hill LLC How a System at Equilibrium Responds to Disturbances 3) Effect of Catalysts Catalysts increase the rate of both the forward and reverse reactions. Equilibrium is achieved faster, but the equilibrium composition remains unaltered. Activation energy is lowered, allowing equilibrium to be established at lower temperatures. © McGraw Hill LLC How a System at Equilibrium Responds to Disturbances 3) Effect of Catalysts Exothermic reactions N2 g + 3 H2 g ⇌ 2 NH3 g ∆𝑟 𝐻 𝑜 = −91.8 kJ © McGraw Hill LLC How a System at Equilibrium Responds to Disturbances 3) Effect of Catalysts Exothermic reactions N2 g + 3 H2 g ⇌ 2 NH3 g ∆𝑟 𝐻 𝑜 = −91.8 kJ The Haber process for producing ammonia from the elements is exothermic, cooling down the reactants should result in more product. However, the activation energy for this reaction is high! It is slow of lower temperature This is an instance where a system in equilibrium can be affected by a catalyst, by allowing reactions to reach equilibrium at a lower temperature. © McGraw Hill LLC  Attractive Forces Intramolecular or bonding forces are found within a molecule. The chemical behavior of each phase of matter is the same because the same constituent particle is present in each case. H2O molecules are present whether the substance is in the solid, liquid, or gas phase. Intermolecular or nonbonding forces are found between molecules. The physical behavior of each phase of matter is different because the strength of these forces differs from state to state. © McGraw Hill LLC 6  Intramolecular Forces vs. Intermolecular Forces © McGraw Hill LLC 7  The Nature of Intermolecular Forces Intermolecular forces arise from the attraction between molecules with partial charges, or between ions and molecules. Intermolecular forces are relatively weak compared to bonding forces because they involve smaller charges that are farther apart. © McGraw Hill LLC 8  Covalent and Van Der Waals Radii The van der Waals distance is the distance between two nonbonded atoms in adjacent molecules. The van der Waals radius is one-half the closest distance between the nuclei of two nonbonded atoms. The VDW radius is always larger than the covalent radius. © McGraw Hill LLC 9  Periodic Trends In Covalent And Van Der Waals Radii Like covalent radii (blue quarter-circles and numbers in the figure), van der Waals radii (red quarter-circles and numbers) decrease across a period and increase down a group. © McGraw Hill LLC 10  Comparison of Bonding Forces Basis of Energy Force Model Attraction (kJ/mol) Example Ionic Cation–anion 400–4000 NaCl Nuclei–shared Covalent 150–1100 e⁻ pair Cations– Metallic delocalized 75–1000 Fe electrons © McGraw Hill LLC 11  Comparison of Nonbonding Forces Basis of Energy Force Model Attraction (kJ/mol) Example Ion charge– Ion-dipole 40–600 dipole charge Polar bond to H– dipole charge H bond 10–40 (high EN of N, O, F) Dipole-dipole Dipole charges 5–25 Dispersion Polarizable 0.05–40 (London) e⁻ clouds © McGraw Hill LLC 12  Ion-Dipole and Dipole-Dipole Forces Ion-dipole forces result when an ion and a nearby polar molecule attract each other. The most important example of an ion-dipole force takes place when an ionic compound dissolves in water. Dipole-dipole forces are the attractive forces between the positive pole of one polar molecule and the negative pole of another polar molecule. Dipole-dipole forces only exist for polar molecules. For compounds of similar molar mass, the greater the molecular dipole moment, the greater the dipole-dipole forces. Dipole-dipole forces are not as strong as ion-dipole forces. © McGraw Hill LLC 13  Polar Molecules and Dipole-Dipole Forces © McGraw Hill LLC 14  Dipole Moment and Boiling Point © McGraw Hill LLC 15  The Hydrogen Bond Hydrogen bonding is possible for molecules that have a hydrogen atom covalently bonded to a small, highly electronegative atom with lone electron pairs, specifically N, O, or F. An intermolecular hydrogen bond is the attraction between the H atom of one molecule and a lone pair of the N, O, or F atom of another molecule. Hydrogen bonds are generally stronger than other dipole-dipole forces. Substances that exhibit hydrogen bonding have exceptionally high boiling points. © McGraw Hill LLC 16  Hydrogen Bonding and Boiling Point © McGraw Hill LLC 17  Problem Drawing Hydrogen Bonds Between Molecules of a Substance PROBLEM: Which of these substances exhibits H bonding? Draw examples of the H bonds between two molecules of each substance that does. (a) C2H6 (b) CH3OH (c) © McGraw Hill LLC 18  Dispersion (London) Forces Dispersion forces or London forces arise when an instantaneous dipole in one particle induces a dipole in another, resulting in an attraction between them. Dispersion forces exist between all particles, increasing the energy of attraction in all matter. They are the only force existing between nonpolar particles. Dispersion forces are stronger for more polarizable particles. In general, larger particles experience stronger dispersion forces than smaller ones. As molar mass increases, dispersion forces increase in strength, and so do boiling points. A molecular shape that has more area allows stronger attractions. © McGraw Hill LLC 19  Dispersion Forces Among Nonpolar Particles When atoms are far apart they do not influence one other. When atoms are close together, the instantaneous dipole in one atom induces a dipole in the other. The process occurs throughout the sample. © McGraw Hill LLC 20  Molar Mass and Trends in Boiling Point Dispersion forces are stronger for larger, more polarizable particles. Polarizability correlates closely with molar mass for similar particles. © McGraw Hill LLC 21  Molecular Shape, Intermolecular Contact, and Boiling Point © McGraw Hill LLC 22  Determining the Intermolecular Forces In a Sample © McGraw Hill LLC 23  Problem Identifying the Types of Intermolecular Forces PROBLEM: For each substance, identify the key bonding and/or intermolecular force(s), and predict which substance of the pair has the higher boiling point: (a) MgCl2 or PCl3 (b) CH3NH2 or CH3F (c) CH3OH or CH3CH2OH (d) Hexane (CH3CH2CH2CH2CH2CH3) or 2,2- dimethylbutane. © McGraw Hill LLC 24  Phases of Matter Each physical state of matter is a phase, a physically distinct, homogeneous part of a system. The properties of each phase are determined by the balance between the potential and kinetic energy of the particles. The potential energy, in the form of attractive forces, tends to draw particles together. The kinetic energy associated with movement tends to disperse particles. © McGraw Hill LLC 25  A Comparison of Gases, Liquids, and Solids Gas Liquid Solid Kinetic energy dominates the potential Intermolecular attractions between particles are strong Intermolecular forces clearly dominate the kinetic energy of intermolecular forces. enough to pull the particles closer together. energy of the particles. On average, gas particles are far apart Liquid particles touch each other (higher density) and The particles are fixed in position next to each (low density) and move randomly have limited freedom of motion. The particles move other in a regular three-dimensional pattern; solids throughout the container. randomly around each other. have relatively high density and very low freedom of motion; particles in a solid just jiggle in place. The large distance between gas particles Liquids resist an applied force and thus compress only The particles are so close together that a solid means that a gas is highly compressible. slightly. compresses even less than a liquid. Since the gas particles have freedom of The particles still have enough kinetic energy to move Since the particles have very little freedom of motion, a gas fills its container, taking randomly around each other, so a liquid conforms to motion, a solid does not take the shape nor volume both the volume and shape of the the shape of its container, but it has a surface (it does of its container. container. not fill the container). A gas flows easily through another gas. A liquid does flow, but much more slowly than a gas. A solid does not flow significantly. © McGraw Hill LLC 26  Types of Phase Changes and Their Enthalpies Solid to liquid, and vice versa: As the temperature increases, the particles in a solid gain kinetic energy and move out of their fixed positions in the process of melting, or fusion; the opposite change is called freezing. Liquid to gas, and vice versa: As the temperature increases further, the molecules in the liquid phase gain sufficient kinetic energy to separate from each other completely and form a gas in the process of vaporization; the opposite process, changing from a gas to a liquid, is called condensation. Solid to gas, and vice versa: Under certain conditions, as the temperature of a solid increases and the particles gain kinetic energy, they move directly to the gas phase in a process called sublimation; the opposite process is called deposition. © McGraw Hill LLC 27  Heats of Vaporization and Fusion for Several Common Substances © McGraw Hill LLC 28  Phase Changes and Their Enthalpy Changes © McGraw Hill LLC 29  Quantitative Aspects of Phase Changes Within a phase, heat flow is accompanied by a change in temperature, since the average Ek of the particles changes. q = ( amount )  ( heat capacity )  T During a phase change, heat flow occurs at constant temperature, as the average distance between particles changes. q = (amount ) ( H of phase change) © McGraw Hill LLC 30  A Heating-cooling Curve For The Conversion Of Gaseous Water To Ice Figure 12.13 Access the text alternative for slide images. © McGraw Hill LLC 31  Sample Problem Finding the Heat of a Phase Change Depicted by Molecular Scenes PROBLEM: The scenes below represent a phase change of water. find the heat (in kJ) released or absorbed when 24.8 g of H2O undergoes this change. © McGraw Hill LLC 32  Solution © McGraw Hill LLC 33