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Chem 1000 Chemical Bonds - exam notes.docx

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CHEM 1000
Michael Hempstead

Lewis Structures Steps: V - valence e's needed N - e's needed T - (N-V) TAKES B - (T/2) BONDS L - (V-T)/2 - Lone pairs Oxidation method is another method for accounting where the electrons are Lee Allen's method is most accurate method. Bond order in 3/2 for lewis structure of "real life" Octahedral = 3 bonds Tetrahedral = 2 or 1 bonds (mostly 2 bonds) Trigonal pyramidal does not have an atom in the vertices whereas the tetrahedral does. Trigonal pyranmidal makes a triangular square prism kind of a base when drawn in 3D shape Trigonal pyramidal will have smaller angles between elements = 107 degrees Tetrahedral will have larger angle = 109.5 degrees Linear molecule = Molecular orbital theory - Today o End bonding o Starting of solids and liquids  November 12 finishing solids and liquids (Finish lectures ALL MATERIAL!!!)  November 19 quiz 3  November 26 LAST CLASS!!!! (review quiz 3 + overview)  December 3 - Study day  December 10 - Exams Begin Beyond Lewis Structures  Useful in explaining bonding in simple molecules + predicting molecular shape  Don't explain why electrons in bond pairs bring nuclei together  Can't estimate Valence Bond Theory Atomic orbitals in valence shell orbitals to an adjacent atom (overlapping) Ex. Overlap x - z and overlap is y Bond forming = two nuclei together= interaction in molecule that's not in atom Blend atomic orbitals = overlap to create bond Ex. Coordinate Covalent Bond - One atom w/ two electrons, other with none, overlap with two electrons (Ammonium N-H [NH ]) 3 Ex. Covalent Bond- One atom w/ one electron, other has one, overlap w/ two electrons (Ammonia N-H + NH 3) (share common region of space and create bond) Chemically methane is stable (from VESPR, it is tetrahedral) - but not 90 , carbon needs to have 4 unpaired electrons to be 90 Two nuclei= effect way of electrons= blend atomic orbitals= hybridization (hybrid orbitals, ex. sp3) Energy and Hybrid orbitals Energy is conserved sp3 as energy between s and orbitals NH3 vs. PH3 N- sp3 bonding at 109 degrees P- bond angle (93 degrees) looks like it doesn’t hybridize  Hybridizes at 93 degrees  Lone pair on P > than on N  Bond pair-bond pair angle is smaller than in PH3 H2O (104 degrees) H2S (90 degrees) Sp3 hybrid doesn't explain Boron or Beryllium (insufficient electron, can create sp2[Boron] or sp[Beryllium] instead) TOTAL ENERGY IS CONSERVED! # of hybrids must = # of valence electrons Multiple Bonds Consider C2H4  Suggests sp2 hybrids involved (still remaining 2p orbitals  Carbon-carbon bond from overlap of sp2 hybrids (basically since there are 2p orbitals, they are vertical, the 3 planar sp3 orbitals can overlap to form a sigma bond C-C, the remaining 4sp3 bonds bond to H as sigma bonds C-H, then 1 pi bond forms through the 2p orbitals)  Sigma bond (single bond, overlap along the inter-nuclear axis(straight line between nuclei))  Pi bond (two bonds [above and below molecule], stretches over and under (side-to-side: loop))  Pi bond stronger than single bond but more reactive Lone electron that must react, therefore, will most likely create the pi bond Sigma determines shape, pi restricts rotation about the axis Orbital overlap is more extensive for the sigma bonds so the pi bonds are weaker than the sigma bond Triple bonds (suggests sp orbitals are involved, there are still electrons in the 2p orbitals) Ex. Acetylene (ethyne):  3 sigma bonds (2 2C-H bonds and 1 C-C bond)  2 pi bonds (2p orbitals of each to the other, the one on the opposite side)  appears to have 4 bonds but are actually 2 D orbitals Ex. Sulphate:  Sum of all charges is 2- (two O carrying one -ve charge each)  S 6 valence electrons, O 6 or 7 valence electrons (based on charge)  4 sigma bonds (S 3sp3-O 2sp2)  2 pi bonds (S 3d-O 2p)  Resonance structure (8 electrons, 6 orbitals [3d S and all 2p O]) Delocalization Benzene:  Suggest sp2 hybrid (planar)  12 Sigma bonds (6 C-H and 6 C-C bonds)  6 2p orbitals (form ring): 6 electrons are delocalized over 6 atoms Van der Waal Forces  Force b/w molecules  Inter/intra-molecular (b/w vs. w/i)  Responsible for Real Gas behaviour, Condensation London Dispersion forces  -ve for a given time  Finite chance that electronic charges are not uniformly distributed  Charge at a point  Forms an instantaneous dipole (influences neighbouring molecule, induces them to create a dipole [induced dipole], = two dipoles that can attract}  Partial change in something = lowercase delta (S)  Change in something = upper case delta ( /_\ )  Normal => instantaneous dipole => instantaneous dipole induced dipole  Strongest w/ many electrons or elongated molecule  Polarizability (molecules has the ability or tendency to have charge separation to occur)  Often seen in boiling points of similar compounds Ex. He and Rn  BPT (Boiling point): He 4K vs. Rn 211K …..Rn is larger  Another ex. n-pentane is more elongated (36.1 degrees BP) vs. smaller molecule (9.5 degrees BP) Dipole-Dipole  Elements that have different electronegativity have a permanent charge separation, creates a polar bond  Can contain induced dipole as well (addition to L
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