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by: Burnice Herman

GeneralChemistryI CHEM101

Marketplace > Drexel University > Chemistry > CHEM101 > GeneralChemistryI
Burnice Herman
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This 10 page Class Notes was uploaded by Burnice Herman on Wednesday September 23, 2015. The Class Notes belongs to CHEM101 at Drexel University taught by DanielKing in Fall. Since its upload, it has received 11 views. For similar materials see /class/212562/chem101-drexel-university in Chemistry at Drexel University.


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Date Created: 09/23/15
Chem Notes Chapter 899 11102010 45700 PM Resonance Structures o Molecule or ion represented by 2 or more plausible Lewis Structures o The difference bt resonance structures is the location or placement of the electrons Delocaized electrons 0 Spread out over several atoms 0 Part of resonance hybrid o True structure is single composite o Resonance hybrid CPI and rquot 1 39 sf 9 0 p p l 2 Exceptions to Octet Rule Free radicals molecules with odd number of valence electrons and at lease one of them unpairedvery reactive o Incomplete octets molecules with stable structure but atoms do not achieve octet Be B or Al BeCl2 even though we could create double bonds the molecule is more stable with incomplete o Expanded valence shellls a central atom has more than 8 e SP Te 0 They use the d forbital s to hold extra electrons o IF4 would be an exception to the octet rule because 2 extra electrons goes on I CHAPTER 9 ValenceShell ElectronPair Repulsion VSEPR simple method to determine geometry Basis pairs of valenece electrons in bonded atoms repel one another Molecules are most stable when atoms are farthest away from each other Electron groups o Valence e in region around central atom o Lone pair o Bond sdt Electron group Geometry o Orientation of electron groups based on electron repulsion o of electron pairs lone or shared Number of atoms VSEPR geometry 2 linear 3 Trigonal planar 4 tetrahedral 5 Trigonal bipyramidal 6 octahedral 0 Molecular Geometry o Tells of the shape of the molecule o Draw lewis electron dot structure o Count of bonding e around central atom double triple count as 1 pair o Count of lone pair o Match electron pair information to shapes VSEPR Notation Central atomA Terminal atomsX LoneE General notation AXnEm o Example PCL3 Bonded pairs 3 Lone pairs 1 AX3E1 Egeom tetrahedral o MGeom triangular pyramidal Molecular Polarity o NonPolar diatomic molecules no electronegativity difference o Polar electronegativity difference partial charges o Dipole moment p o p Ad deltacharge ddistance between posneg charged particles 0 O O O Intermolecular forces forces betweendifferent molecules phase changes determine melting freezing boiling points 0 0 London dispersion must have electrostatic forces for partially negative and partially positive molecules They occur for all molecules because at any given time a molecule can have a partial charges 0 DipoleDipoepermanent version of London dispersion forces same electrostatic forces DO NOT exist for nonpolar molecules 0 Hydrogen Bonding specific dipoledipole attractionhydrogen needs to be bonded to lone electron pair on small electronegative atom NOF also an example of an intramolecular force Intramolecular forces a single bond Melting and Boiling points Intermolecular forces involved Properties of Gases Molecules are far apart No fixed volume or shape Translucentlets light throught Constant random motion Compressible Exert pressure on container Expand into volume of container Higher energy than solids and liquids Kinetic Molecular Theory Molecules are much smaller than intermolecular distance Molecules in continuous random motion Pressure measured with a barometer more molecules equals higher pressure Since gases are so far apart forces bt molecules do not matter IDEAL Boyles Law o At constant temp and mass of gas PVconstant P1V1P2V2 Charles Law o At constant pressure and mass of gas 0 VconstantT Vtconstant V1T1V2T2 Avagadro s Law o At constant pressure and temperature o Vconstantn Ideal Gas Law PVnRT R08206 LatmmolK GAS DENSITY o p massvol to get mass use nvPRT o nmassgMW massVPMWRT o calculate density of CH4 at STP 1atm 0 deg Celcius o p 1atm1Ggmo 08206273 K 7lgL Lawof Combining Volumes o In a system where temperature and pressure are constant gas volumes can be calculated using stoichiometry o If temp and press Are not constant then must use the ideal gas law calculation Law of Partial Pressures c When gases come together but do not react they act independently of each other Therefore the total pressure in the system is the sum of the two gases in each system They also are present in the same amount of moles The number of collisions will still be the same Ptotal P1P2P3 o Partial pressure the pressure a gas would exert if it were alone in the container P1N1RTV 0 Mole Fraction o Ratio of moles of a component to total moles XN1Ntotal o XP1Ptotal Molecular Speed o The higher the molar mass the lower the most probable speed c As temperature is raised the speed increase o Lighter molecules move faster REAL GASES o Gases deviate from the Ideal gas law at low temp and high pressure o Volume of molecules a significant fraction of the total volume when molecules close together Intermolecular forces of attraction reduce pressure Van der Waals PaVA2VbnRT Aattractive forces bresidual volume Smog Automotive trafficemissions of contaminants o Sunlightinitiate chemical reactions o Little air movementallow build up of reactants thermal inversion Greenhouse effect c Absorption of outgoing radiation by gas molecules o Important natural greenhouse gas 0 Water vapor o Important anthropogenic greenhouse gasses 0 C02 CH4 N20 CFCs 11102010 45700 PM 11102010 45700 PM Most Important Information 0 Group lA charge of 1 0 Group ZA charge of 3 0 Group 3A charge of 3 0 Group 5A charge of 3 0 Group 6A charge of 2 0 Group 7A charge of 1 0 As n increases the energy of the electron increases and the electron is farther away from the nucleus and less tightly bound to it 0 Shell a collection of orbitals with the same value of the principal quantum number n 0 When n1 there is only one kind of atomic orbital possible and then there are two when n2 and so on o Subshells a group of atomic orbitals with the same n and quantum numbers 0 In 1 o The number of atomic orbital types within a principal energy level equals n 0 There are 2I1 atomic orbitals in a subshell o The total number ofatomic orbitals value 0 1 2 3 in a shell equals n2 Subshell s p d f o nrelates to the atomic orbital s size ml 0 101 21012 32 o I relates to the atomic orbital s 2390391392393 shape 0 m relates to the atomic orbital s orientation 0 mscan only equalor 12 o Pauli exclusion principle No more than two electrons can occupy the same atomic orbital in an atom and those electrons must have opposite spins 0 Each principle energy level n can accommodate a maximum of anelectrons o The ionization energy generally increases up and to the right 0 Electron affinity increases right and up a group on the periodic table 0 Isolated system does not exchange anything such as matter or energy with its surroundings 0 Closed system can exchange energy but not matter with its surroundings o Endothermic a process in which thermal energy must be transferred m a thermodynamic system in order to maintain constant temperature 0 Exothermic refers to a process in which thermal energy must be transferred of a thermodynamic system in order to maintain constant temperature I 0 Units of specific heat g C Units of heat capacity Atomic radius one half the distance between the nuclei centers of two like atoms in a molecule For main group elements atomic radii increase going down a group in the periodic table and decrease going across a period 0 There is a large increase in atomic radius going from any noble gas atom to the following Group 1A atom soeectronic refers to atoms and ions that have identical electron configurations Bomb Calolimetry O qreactioon qwater qbomb Noble gases have a ridiculously low electron affinity Orbital 0 An orbital is three dimensional 0 An electron shell consists of a collection of orbitals with the same principal quantum number 0 An orbital may be designated with the letters s p d f 0 An orbital describes the region of space in which one will most likely find an electron The d orbital occur in groups of 5 and hold up to 10 electrons If you are multiplying it is inversely proportional If you divide it is directly proportional First ionization energy increases up and to the right ofthe periodic table Group 1 elements have the largest second ionization energy Combustion will be a negative reaction Formation will be a positive reaction Bonds breaking is an endothermic reaction Bonds forming is an exothermic reaction Bond length increases with increase in atomic size Bond length decreases from single to double to triple bonds Bond Enthalpy increases from single to double to triple bonds Shorter BondStronger Bond Bond enthalpy decreases with increase in atomic size Halogens noble gases phosphorus and sulfur are all expanded octet rule Electronegativity increases up and to the right Polarity increases with the largest difference in electronegativity Formal chargevalence electrons number of bonds number of paired electrons o The most stable bond length represents a balance of attractive forces between nuclei and electrons and repulsive forces between nuclei 0 The strongest intermolecular force of attraction is hydrogen bonding o Symmetric geometry is nonpolar o If the geometry is symmetric it is polar o The higher the molecular weight the higher the melting point 0 The more electronegative atom gets its preference of oxidation number 0 AHkJmol 0 To find the molecule with the highest boiling point you look for the molecule that is going to overcome the strongest intermolecular forces 0 For hydrogen bonding the hydrogen needs to be bonded to a N O or F o Dipole dipole force the atom has to be polar o All molecules have London dispersion forces Number of Bonds Number of lone pair Electron pair Molecular Geometry Bond Angle electrons Geometry 2 0 linear Linear 180 3 O Trigonal planar Trigonal planar 120 2 1 Trigonal planar Bent Less than 120 4 O Tetrahedral Tetrahedral 1095 3 1 Tetrahedral Trigonal pyramidal Less than 1095 2 2 Tetrahedral Bent Less than 1095 5 O Trigonalbipyramidal Trigonalbipyramidal 90 120 and 180 4 1 Trigonalbipyramidal Seesaw 90 120 and 180 3 2 Trigonalbipyramidal T shaped 90 and 180 2 3 Trigonalbipyramidal Linear 180 6 O Octahedral Octahedral 90 and 180 5 1 Octahedral Square pyramidal 90 and 180 4 2 Octahedral Square planar 90 nd 180


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