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Exam one Study Guide

by: Lauren Dennis

Exam one Study Guide Chemistry 1212

Lauren Dennis

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An overview of points to know for Exam one of Chem 1212. Describes bonds and properties of compounds described in Chapter 11 and 12.
Chemistry 1212
Dr. Doyle
Study Guide
hydrogen bonds, Chemical Bonding, molarity, molality, properties
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This 6 page Study Guide was uploaded by Lauren Dennis on Tuesday September 27, 2016. The Study Guide belongs to Chemistry 1212 at Georgia State University taught by Dr. Doyle in Fall 2016. Since its upload, it has received 12 views. For similar materials see Chemistry 1212 in Chemistry at Georgia State University.


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Date Created: 09/27/16
Chemistry Study Guide Bonds  Electrostatic: the ability of an object to acquire an electron o Holds matter together o Trend on the Periodic Table:  Increases going down the columns  Increases going across the rows  Phases: solid, liquid, and gas o Spacing between molecules  Gas > Solid > Liquid  Forces: o London (induced dipole)  Every molecule have these forces  Very weak  Larger molecular weight = more electrons  Larger the molecule = Stronger the force  BP determined by molecular weight  Ex. CH4, CH3CH, CH3CH2CH3  CH3CH2CH3 > CH3CH > CH4 o If they are similar Molecules  As the force increases so does the boiling point, bp  Lowers the density  Lowers the melting point  Dipole – Dipole Interactions o Occurs only between polar molecules  Water or H2O o Ex. CH4, CH3F, CF4  CH4: NONPOLAR  CH3F: POLAR  Dipole- Dipole  High BP and MP  CF4: NONPOLAR  London  Polar bonds, but are symmetrical  Van der Waals o The category of forces  Include: London and Dipole-Dipole o Short Range of Forces o Very Weak o Very few interactions between molecules  Hydrogen Bonding o Similar to dipole- dipole o Not a “real” bond o Need: H-F, C-H, C-O, O-H  Hydrogen Bond Donor o Lone Electron Pair  Hydrogen Bond Acceptor H o Ex. CH3FOH to NH3  N H H H C O H H  Ion Dipole o Strongest forces o Only found in solutions o The molecule “dissolves”  Breaks down to their original components  Ex NaCl dissolved into H2O o Na+ and Cl- o Ionic Solution o “Like dissolves Like”  Polar with Nonpolar: Oil/H2O o Need an ion/dipole: ionic substance o Larger the charge = the greater the interaction  ION > HYDROGEN > DIPOLE- DIPOLE > DISPERSION (LONDON)  PROPERTIES o Boiling point: the vapor pressure = the atmospheric pressure o Melting point: the temperature at which a given solid will melt o Surface tension: energy required to increase surface area o Heat of vaporization: the energy required to change a liquid into a gas o Vapor pressure: the pressure of a vapor when in contact with a liquid or a solid o Visocity: resistance to float  Lower molecular forces = tend to float  Motor Oil o Capillary action: ability to flow upward  Forces between tube and substances  Example of the Heat of Vaporization: o Heat (q) in KJ to vaporize 2.58 kg of H2O  Bp = 100 C  Heat of Vaporization = 40.7 KJ/mol q= Heat of Vaporization x moles(n) (40.67KJ/mol)(143mol) = q q= 5820 KJ  Temperature Change: o Change in Temperature = (q)(SH)(m)  q= heat  SH= specific heat  m= mass  Change in Phase o Change in Vaporization = - Change in Condensation o Change in Fusion - - Change in Freezing o Change in Sublimation = (Change in Fusion) + (Change in vap.) o Change in Disposition = _____ - (Change in Sublimation)  Phase Diagrams  Ga Properties MP BP Surface Visocity Vapor Heat of s Tension Pressure Vaporization London Lowest Lowest Lowest Lowest Highest Lowest Dipole – Lower Lower Lower Lower Higher Lower Dipole Hydrogen Low Low Low Low High Low Bonding Ion Dipole Highest GaHighest Highest Highest Lowest highest s Liqui d Soli d  Dynamic Equilibrium: a flat line; no change in temperature o Solid – Liquid o Liquid – Gas  Solid is found at lower temperatures  Gas is found at higher temperatures  Solid – Liquid: melting point  Liquid – Gas: boiling point  Liqui Soli d d Ga s o Where the three lines meet: TRIPLE POINT o Where it goes from solid to liquid: N. MELTING POINT o Where it goes from liquid to gas: N. BOILING POINT o CRITICAL POINT: the liquid to solid phase does not exist; the molecule becomes cloudy  It is at the tip of the line dissecting from liquid to gas  Clausius – Clapeyron Equation: o Ln(Pvap) = ((-Heat of Vaporization)/(R))(1/T) + lnB  What is the equation?  Y = mx + b o Ln(P2/P1) = ((-Heat of vap.)/(R))((1/T2)-(1/T1))  All pressures in either torr or atm  All temperatures in Kelvin  All Heat of vaporization in J/mol  R = 8.314 J/molK  Cubics: o Simple/ Primitive  1 atom total  1/8 of the atom is in each corner of the square o Body Center  2 atoms total  1 in the center and 1/8 in each corner o Face Center  4 atoms total  ½ atom on each face and 1/8 in each corner  Energetically favorable  Compounds in a solution dissociate into the molecular components o Must disrupt the network; to break the bonds = to form new interactions with H2O = Solution o It is unfavorable to disrupt the bonds = higher energy  2 types of favorable forming solutions: o Change in enthalpy increase o Change in entropy decrease  Disorder  More disorder = very favorable  2 substances that form a homogensis mixture: o Solid/liquid: NaCl and H2O o Gas/gas: Air (Nitrogen and Oxygen) o Liquid/ liquid: Vodka and H2O o Gas/ liquid: Soda (CO2 and HO) o Solid/ solid: brass  2 parts to a Solution: o Solute: substances present in lesser amounts o Solvent: a substance in the greatest amount in the solution  Solubility: whether one substance dissolves in another o Miscible o “like dissolves like”  intermolecular forces at work o Example:  Which compounds are soluble in Hexane (CH3CH2CH2CH2CH3)?  London forces present o (1): H2O “hydrogen bonds” = NOT o (2) CH3CH2CH3 “London” = SOLUBLE o (3) HCl “dp-dp” = NOT  Equilibrium: point of even; both substances are dissolved with no excess o Saturated: no more solute in a solution o Unsaturated: more solute can go into the solution o Supersaturated: unstable; solution can fall out easily  Things to Know: o Most of the time solubility of solids increase with temperature o Most of the time solubility of gases decrease with temperature o Most of the time solubility of gases increase with pressure  Henry’s Law: o Solubility of a gas is dependent on temperature o S gas= K h gas  Khis a constant that every gas has o Example:  O2 at 25C with a pressure of air at 1atm  P oxygen.21 atm  K h 1.3 x 10 M/atm -3 o S o2(1.3 x 10 )(.21)  = 2.7 x 10 M-4 o Expressing Concentrations:  Molarity = (the mol of the solute)/(liters of the solution) = M  Temperature dependent o Volume changes with temperatures  There can be multiple molarities in a solution  Molality = (mol of solute) / (kg of solvent) = m  (X) Mole Fraction = (mol of solute) / (tot. mol of substances)  Mole % = Mole Fraction x 100%  Mass % = (Mass of solute)/ (total mass of solution) x 100  Density = mass/ volume o Rouault’s Law (Colligative Properties)  Vapor Pressure of a liquid when adding to another substance  Adding a NON-volatile to a volatile o Vapor Pressure will never go up o Will always go down o P solutionX x P solvent  Adding volatile to volatile o P = X P + X P solution a a b b o Boiling and Freezing Point Depressions:  Change in Temperature = m x K f or b  m= molality  K = freezing or boiling point constant o Osmotic Pressure:  Travel through a semipermeable membrane without stress  = MRT o M= molarity o R= .0821 o T= Temperature (kelvin)


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