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Chemistry 1061 Notes ch 12 & 15 for Final Exam

by: mandygh926

Chemistry 1061 Notes ch 12 & 15 for Final Exam chem 1061

Marketplace > University of Minnesota > Chemistry > chem 1061 > Chemistry 1061 Notes ch 12 15 for Final Exam
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This includes the last two chapters of material for the final exam: Phase Changes and Diagrams, Exothermic and Endothermic Reactions, Boiling Point, Special Properties of Water, Intermolecular Forc...
Chemical Principles I
Debra Salmon
Chemistry, polarizability, isomers
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This 4 page Bundle was uploaded by mandygh926 on Tuesday July 19, 2016. The Bundle belongs to chem 1061 at University of Minnesota taught by Debra Salmon in Fall 2015. Since its upload, it has received 14 views. For similar materials see Chemical Principles I in Chemistry at University of Minnesota.

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Date Created: 07/19/16
Chemistry Notes for Final Exam: Ch 12 & 15 on 12/22/15  **(Rest of Final exam is ch 1­11 on the previous Textbook chapter notes) Corresponding Textbook: Chemistry; The Molecular Nature of Matter and Change  Includes: Phase Changes and Diagrams, Exothermic and Endothermic Reactions, Boiling Point,  Special Properties of Water, Intermolecular Forces, Polarizability, Unit Cell, Isomers, and  Hydrocarbon Functional Groups Chapter 12: Intermolecular Forces: Liquids, Solids, and Phase Changes  Phase Changes: when a substance changes physical state. In a closed container phase  changes are reversible and reach static equilibrium. Include the positive and negative  enthalpy changes for sublimation, fusion, and vaporization.   Heating/cooling curves: show the phase changes and the heat they absorb/release  depending on the temperature. Higher temp=higher vapor pressure= weaker IMF  (intermolecular forces)  Phase diagrams: lines divide the physical states based on pressure and temperature. The  line between solid and liquid has a positive sloping line because of density for every  substance except H O. The critical points are where the densities are equal with one of  2 them being the triple point. o Exothermic­ (goes to the left) negative enthalpy values. Condensation (gas to  liquid), freezing (liquid to solid), deposition (gas to solid) o Endothermic­ (goes to the right) positive enthalpy values. Melting (solid to  liquid), vaporization (liquid to gas), sublimation (solid to gas) o  Boiling Point: the temp where vapor pressure = atmospheric pressure. Directly related to  pressure; Higher altitudes have lower atmospheric pressure and lower boiling points.  Stronger IMF = higher boiling point.  Vapor Pressure: kinetic energy needed to overcome IMF of liquid phase and escape to the gas phase. Not affected by pressure.  Properties of water:  Surface tension: downward force on surface molecules due to attractions to  interior molecules. Stronger IMF= higher surface tension  Viscosity: Resistance to flow. Stronger IMF= higher viscosity  Capillary Action: Hydrogen bonding because glass wall and H2O pulls up surface into tube.   Clausius­Clapeynon equation: quantifies the effect of temperature on vapor pressure. P’s  are vapor pressure, T’s are temperature in (K), includes the enthalpy of vaporization, and  R= 8.314 J/mol*K P 2 −∆H ° vap 1 1 ln( ) = ( − ) P 1 R T 2 T 1 Types of Intermolecular Forces:  Ion­dipole: between an ion and a polar molecule  Dipole­dipole: between two polar molecules  Hydrogen bonding: a type of dipole­dipole force (polar molecules) where Hydrogen  directly bonds to either N, O, or F.  Dipole­induced dipole: force between a polar and a nonpolar molecule where the  nonpolar molecule gets a temporarily induced dipole.  Dispersion forces: act on all atoms, ions, and molecules. Only force between two  nonpolar molecules.   Polarizability: how easily the electron cloud of an atom can be distorted. Higher mass=  more electrons= higher polarizability and greater dispersion forces. If masses are similar  then higher surface area has stronger dispersion forces.  *Cations are less polarizable than their parent atoms while anions are larger because of their size.  Solid state: crystalline solid (orderly arrangement packed in 3D array) or amorphous  solids ( no arrangement, poorly defined shape).  Atomic solids: held together by only dispersion forces. Only formed from noble gases  Molecular solids: dipole­dipole, dispersion, and H­bonding  Ionic solids: cations surround anions making ionic bonds much stronger than the van der  waals forces in atomic or molecular solids. High lattice energy. Four types of ionic solids  include” sodium chloride structure, zinc blende structure, fluorite structure, and  antifluorite structure.  Metallic solids: metallic bonding forces give it high electrical and thermal conductivity,  luster, and malleability  Network covalent solids: strong covalent bonds where separate particles are not present.  Extremely high melting and boiling points. Examples include graphite and diamond   Amorphous solids: noncrystalline, unordered structure. Ex. glass  Crystal lattice: arrangement of atoms (spheres) in regular pattern Unit Cell: smallest portion that is repeated in all directions. 3 types of cubic unit cells  Simple cubic: atom in each corner. 1 atom/ unit cell  Body­centered cubic system: one atom in the middle and 8 at the corners. 2 atoms/  unit cell  Hexagonal/Fare­centered cubic: atom at each face and corner. 4 atoms/ unit cell  Packing efficiency: the percentage of the total volume occupied by the spheres  themselves.  Polymers: chains of smaller molecules (monomers). Molar mass of polymer= molar mass of repeating unit x n (where n is the degree of polymerization.) Chapter 15: Organic Compounds and The Atomic Properties of Carbon  Isomers: same formula, different properties. o Constitutional isomer: aka structural isomer, same formula, different connectivity o Stereoisomers: same arrangement of atoms but different orientation of groups in  space.  Optical isomers (enantiomers): mirror images that cannot be  superimposed. Has many biological molecules.  Geometric isomers in alkenes: each C atom in the double carbon bonded  to a different group.qqq  Chiral: asymmetrical molecule. qowhen the central atom in an enantiomer is bonded to 4  different groups. Hydrocarbon Functional groups:  Alkane­ only single carbon­carbon bonds. Saturated molecule. 1. Find longest connected chain of carbon. Use root+ suffix “ane” 2. Name side chains by # of carbons. Root + “yl” 3. Number the carbons along main chain 4. Prefix to show how many similar outlier groups. Name the outliers alphabetically  Alkene­ double carbon­carbon bonds. Unsaturated molecule. Must label where double  carbon bond is. Put it with the lowest carbon #.  Alcohol­ hydroxide bonded to rest of hydrocarbon. Very polar because of the hydrogen  bonding.  Carboxylic acid­ Oxygen double bonded to a carbon which is bonded on one side to a  hydroxide and on the other side to the rest of the chain. Polar from the OH bond.  Sometimes the H+ falls off leading to acidic solution, example of a weak acid.  Amine­ Nitrogen bonded to 3 groups w one lone pair. Common weak base *Amide: peptide bonds. Reaction of carboxylic acid and amine


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