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Chemistry 2 Week 3 Notes

by: Wade Carter

Chemistry 2 Week 3 Notes CHEM 1123

Marketplace > University of Arkansas > CHEM 1123 > Chemistry 2 Week 3 Notes
Wade Carter
GPA 3.73

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Properties of water, Phase diagrams, structures and types of solids, solutions, concentration
University Chemistry II
Lorraine Brewer
Class Notes
Chemistry, Chem II, chemistry 2
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This 5 page Class Notes was uploaded by Wade Carter on Friday February 5, 2016. The Class Notes belongs to CHEM 1123 at University of Arkansas taught by Lorraine Brewer in Spring 2016. Since its upload, it has received 46 views.


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Date Created: 02/05/16
Chemistry II 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. 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. Phase Diagrams:  The lines represent phase changes  The liquid/gas line is the vaporization curve  The solid/liquid curve is the fusion curve  The solid/gas curve is the sublimation curve  On the line, both phases are in equilibrium with each other  At the triple point, all three states exist in equilibrium. 1 Chemistry II 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. 2 Chemistry II  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.  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 3 Chemistry II The Enthalpy of Solution  To make a solution: Δ H solute 1. Overcome all attractions between the solute particles; therefore,  is  endothermic. 2. Overcome some attractions between solvent molecules; therefore,  Δ H 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   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. Solutions:  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. 4 Chemistry II  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. M olarity= molesof solute liter of solution  Molality (m or mol/kg) is defined as the moles of solute per 1 kilogram of solvent. molesof solute Molality= kilogramsof solvent 5


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