CHEM 2 Exam 1 Study Guide
CHEM 2 Exam 1 Study Guide CHEM 1123
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This 8 page Study Guide was uploaded by Wade Carter on Wednesday February 10, 2016. The Study Guide belongs to CHEM 1123 at University of Arkansas taught by Lorraine Brewer in Spring 2016. Since its upload, it has received 186 views.
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Date Created: 02/10/16
Chemistry II Exam 1 Study Guide Bonding: A chemical bond is a force that holds groups of atoms together and makes them function as a unit. o Ionic Bonding electrons are transferred o Covalent Bondingelectrons are shared equally Electronegativity: The ability of an atom in a molecule to attract shared electrons to itself. If the difference in electronegativity is less than 0.4, the bond is polarized. The strength of a bond increases as the difference in electronegativity values increases. Electronegativity increases up a column and left to right across a row on the periodic table. Geometry: Number of Electron Electron Geometry Number of Lone Molecular Geometry Groups Pairs 2 Linear 0 linear 3 Trigonal Planar 0 Trigonal planar 3 Trigonal Planar 1 Bent 4 Tetrahedral 0 Tetrahedral 4 Tetrahedral 1 Trigonal pyramidal 4 Tetrahedral 2 Bent 5 Trigonal bipyramidal 0 Trigonal bipyramidal 5 Trigonal bipyramidal 1 Seesaw 5 Trigonal bipyramidal 2 Tshaped 5 Trigonal bipyramidal 3 Linear 6 Octahedral 0 Octahedral 6 Octahedral 1 Square pyramidal 6 Octahedral 2 Square planar 1 Chemistry II Solids, liquids and gases at the molecular level: Strength of State Density Shape Volume Intermolecular Forces Gas Low Indefinite Indefinite Weak Liquid High Indefinite Definite Moderate Solid High Definite Definite Strong Intermolecular Forces: Intramolecular Bonding o Bonds that form “within” the molecule o Molecules are formed by sharing electrons between atoms. Intermolecular Bonding o Forces that occur between molecules. o Dipoledipole forces, hydrogen bonding, London dispersion forces, ect… o London Dispersion Forces Instantaneous dipoles that occur spontaneously and fleetingly in a given atom and induce similar dipoles in neighboring atoms. o Significant in large atoms/molecules o Larger surface area facilitates the formation of these induced dipoles o Present in all molecules Magnitude of an induced dipole depends on: o Polarizability of the electrons Volume of the electron cloud Larger molar mass = more electrons = larger electron cloud = increased polarizability = stronger induced dipoles = stronger attractions Shape of the molecule o More surfacetosurface contact = larger induced dipole = stronger attraction DipoleDipole Forces o Molecules with dipole moments can attract each other electrostatically. o Only about 1% as strong as covalent or ionic bonds. 2 Chemistry II Hydrogen Bonding o Very strong dipoledipole forces o Hydrogen bonded to a very electronegative atom (Oxygen, Nitrogen, and Fluorine) IonDipole Forces o The positive charged end of a polar molecule is attracted to negative ions, and the negatively charged end of the molecule is attracted to positive ions. Attractive Forces and Solubility: Solubility depends, in part, on the attractive forces of the solute and solvent molecules. o Like dissolves like Capillary Action – The ability of a liquid to flow up a narrow tube vs. gravity o Cohesive forces – intermolecular forces among the molecules of the liquid o Adhesive forces – forces between the liquid molecules and their container Surface tension – resistance of a liquid to an increase in its surface area Viscosity – measure of a liquid’s resistance to flow Vaporization and Vapor Pressure: Vaporization the process by which thermal energy can overcome intermolecular forces and produce a state change from liquid to gas The rate of vaporization increases with: o Increasing temperature o Increasing surface area o Decreasing strength of intermolecular forces Liquids that vaporize easily (weak intermolecular forces) are called volatile. Vaporization = endothermic The heat required to vaporize one mole of a liquid to gas is its heat (or enthalpy) of vaporization (ΔHvap) Vapor Pressure & Dynamic Equilibrium: When the rate of condensation and the rate of vaporization are equal, dynamic equilibrium is reached. Both processes still occur, but their rates are equal. 3 Chemistry II When a system in dynamic equilibrium is disturbed, the system responds so as to minimize the disturbance and return to a state of equilibrium. The boiling point of a liquid is the temperature at which the liquid’s vapor pressure equals the external pressure. Once the boiling point of a liquid is reached, additional heating only causes more rapid boiling; it does not raise the temperature of the liquid above its boiling point. The temperature dependence of vapor pressure is given by the Clausius—Clapeyron equation: P 2 −∆H vap 1 1 o lnP = R (T − T ) 1 2 1 Transitions Between States of Matter: Sublimation is the transition from solid directly to gas. Deposition is the transition from gas directly to solid. Fusion or melting is the transition from solid to liquid. Freezing is the transition from liquid to solid. Energetics of Melting and Freezing: The heat required to melt 1 mole of a solid is the heat (or enthalpy) of fusio).( The heat of fusion is positive because melting is endothermic. ΔH ΔH fus is positive and generally much smaller than vap. ΔH crystallization fusion sublimation=¿∆ H +∆ H fusion vaporization ∆H ¿ Unique Properties of Water: Water is a liquid at room temperature. o Water’s high boiling point is due to hydrogen bonding between water molecules. Water is an excellent solvent, dissolving many ionic and polar molecular substances. 4 Chemistry II o It has a large dipole moment. Water has a very high specific heat for a molecular substance. Ice is less dense than water. o Water expands about 9% when it freezes at a pressure of 1 atm. Structures and Types of Solids: Amorphous Solids disorder in the structures o Glass Crystalline Solids has a regular, ordered arrangement of its components. o The arrangement is usually represented by a lattice, a 3D system of points giving the positions of the components. o The components can be atoms, ions, or molecules. o Unit cell the smallest repeating unit of a lattice structure. Types of Crystalline Solids o Molecular Solids – discrete covalently bonded molecules at each of its lattice points. o Ionic Solids – ions at the points of the lattice that describes the structure of the solid. o Atomic Solids – atoms at the lattice points that describe the structure of the solid. Molecular solids are solids whose composite particles are molecules. o The molecules are held together by intermolecular attractive forces: dispersion forces, dipole–dipole attractions, and hydrogenbonds o Because the attractive forces are weak, they tend to have low melting points (generally < 300 °C) Ionic solids are solids whose composite particles are ions. o They are held together by attractions between oppositely charged ions and each ion attracts all oppositely charged ones around it. Atomic solids are solids whose composite particles are atoms. o Nonbonding atomic solids are held together by London dispersion forces o Metallic atomic solids are held together by metallic bonds. o Network covalent atomic solids are held together by actual covalent bonds. Metals: o Properties Malleability, ductility, conduct electricity and heat, high melting points. o Strong and nondirectional bonding. 5 Chemistry II Network Covalent Solids: o Atoms attach to their nearest neighbors by covalent bonds. o Very high melting points (often greater than 1000°C) o Examples include graphite, diamond, glass, buckyballs, and quartz. Solutions: Solutions = homogeneous mixtures. o Two or more substances make up a mixture o A solution may be composed of a solid and a liquid, a gas and a liquid, or other combinations o Solution formation is the result of the interaction of the intermolecular forces of solute and solvent particles o Nature has a tendency toward spontaneous mixing o LIKE DISSOLVES LIKE Solubility of one substance in another depends on: o The tendency towards mixing o The types of intermolecular forces o Temperature o Pressure The Enthalpy of Solution To make a solution: 1. Overcome all attractions between the solute particles; therefore, H soluteis endothermic. Δ H 2. Overcome some attractions between solvent molecules; therefore, solven is endothermic. 3. Form new attractions between solute particles and solvent molecules; therefore, Δ H mix is exothermic. o The overall ΔH for making a solution depends on the relative sizes of the ΔH for these three processes. For aqueous solutions of ionic compounds, the energy added to overcome the attractions between water molecules and the energy released in forming attractions between the water molecules and ions are combined into a term called the heat of hydration. Attractive forces between ions = lattice energy 6 Chemistry II Attractive forces in water = Hydrogen bonds Attractive forces between ion and water = iondipole ΔHhydration = heat released when 1 mol of gaseous ions dissolves in water = ΔHsolvent + ΔHmix When ions dissolve in water, they become hydrated. Each ion is surrounded by water molecules. The formation of these ion–dipole attractions causes the heat of hydration to be very exothermic. At equilibrium: rate of dissolution = the rate of recrystallization The solution is saturated with solute and no more solute will dissolve A solution that has the solute and solvent in dynamic equilibrium is said to be saturated. o If you add more solute, it will not dissolve. o Saturation concentration depends on the temperature and pressure of gases. A solution that has less solute than saturation is said to be unsaturated. o More solute will dissolve at this temperature. A solution that has more solute than saturation is said to be supersaturated. For most solids, the solubility of the solid increases as the temperature increases. Gases Gases generally have lower solubility in water than ionic or polar covalent solids because most are nonpolar molecules. For all gases, the solubility of the gas decreases as the temperature increases. The larger the partial pressure of a gas in contact with a liquid, the more soluble the gas is in the liquid. Henry’s Law: Sgas P H gas where Sgas is solubility and gas is partial pressure. The k H is Henry’s law constant for that gas Solution Concentration: Molarity (M or mol/L) is defined as the moles of solute per 1 liter of solution. molesof solute Molarity= liter of solution 7 Chemistry II Molality (m or mol/kg) is defined as the moles of solute per 1 kilogram of solvent. molesof solute Molality=kilogramsof solvent Vapor Pressure A nonvolatile solute lowers vapor pressure of solvent The vapor pressure of pure solvent is greater than the vapor pressure of a solution The vapor pressure of the solution is directly proportional to the amount of the solvent in the solution. Raoult’s Law:P solvent in solutionχ solvent∙ P° Colligative Properties Boiling-Point Elevation o Nonvolatile solutes elevate the boiling point of the solvent. o ΔT = K m b solute Freezing-Point Depression o When a solute is dissolved in a solvent, the freezing point of the solution is lower than that of the pure solven. o ΔT = K m f solute Osmosis- flow of solvent into the solution through a semipermeable membrane. π o = M R T Van’t Hoff Factor o The relationship between the moles of solute dissolved and the moles of particles in solution is usually expressed as: 8
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