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Chapter 1

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University of Florida
CHM 2210

Chapter 1 Tuesday, June 25, 2013 6:03 PM 1) Nomenclature A. Alkanes i) ii) B. Naming complex substituents i) Place numbers on the substituent going away from the parent chain on the longest straight chain ii) 4 carbon chain = butyl iii) Group below is called a (2-methylbutyl) iv) C. Common alkyl group names i) Propyl  ii) Butyl  iii) Pentyl  D. Common suffixes i) Aldehyde = "-anal" ii) Ketone = "-anone" iii) Alcohol = "-anol" iv) Amine = "-amine" v) Carboxylic Acid = "-oic acid" vi) 2) Bonding and Molecular Orbitals A. Lewis Dot structures i) For every valence electron, a dot is placed around the atom ii) Boron has three valence electrons, so it can make only three bonds B. Bonding model i) A bond is considered the sharing of two electrons between two adjacent nuclei ii) Ligand bond = Lewis acid-base bond i. http://www.youtube.com/watch?feature=player_detailpage&v=EufPZFAwWco iii) In most cases, with the exception of ligand bonds, one electron from each atom goes into forming the bond iv) Equal sharing = covalent v) Electron transfer = ionic bond vi) Difference in electronegativity between two atoms <1.5, covalent vii) Difference in electronegativity between two atoms <2.0, ionic viii) Difference in electronegativity between two atoms >1.5 and <2.0, polar-covalent C. Covalent bonds i) Two types of covalent bonds, sigma and pi ii) Sigma bonds have electron density shared between the two nuclei iii) Pi bonds have no electron density shared between the two nuclei, just above and below iv) Sigma bonds can be made from many types of orbitals, including atomic and hybridized  v) Pi bonds can only be made from parallel p-orbitals    vi) Only about 80-90% of electron density lies between the nuclei, NOT ALL OF IT vii) Sigma bond is usually the strongest, except with molecular fluorine, F I2.is weak  This is due to The small size of F. It wants to create a sigma bond but there is too much repulsion at the distance it needs to be. So it has to resort to a Pi bond.  This is why the bond dissociation energy of F is2less than the bond dissociation energy of Cl , 2ven though chlorine is below fluorine in the periodic table D. Molecular Orbitals i) An anti-bond is a molecular orbital that results in bond breaking when coupled with a bonding orbital E. Molecular bonds i) Single, double and triple bonds ii) Triple>double>single iii) The longer the bond, the less electron density overlaps between nuclei, and thus the weaker the bond iv) Single bonds are composed of only one sigma bond between two atoms v) Double bonds are composed of one sigma and pi bond vi) Triple bonds are composed of a sigma bond and two pi bonds F. Molecular Structures i) In a neutral compound, Hydrogen makes one bond, oxygen makes two, nitrogen makes three, carbon makes four ii) Neutral structures always obey the octet rule iii) If a structure does not satisfy this rule, there MUST be a charge present iv) Too many bonds = cation… TOO MANY CATS!  If oxygen makes three bondsinstead of two, it carries a positive charge v) Too few bonds= anion  If nitrogen makes two bonds instead of three, it carries a negative charge vi) If carbon makes three bonds,the charge depends on the presence or absenceof a lone pair  Presence yields anion  Absence yields cation vii) G. Hybridization of Atomic Orbitals i) Electron density is relocated in AO PRIOR to bonding ii) sp, sp , sp ii) sp, sp , sp3 iii) 2 iv) Boranes (BH3, BF3, BR3…) are an exception to this table (no pi bonds for sp )  Boron has three valence electrons, so a neutral boron cannot satisfy the octet rule  The result is boron has sp hybridization for its three sigma bonds and has a p orbital with no electrons present  H. VSEPR shapes i) The atoms themselves are not set up in the shape, but the orientation of the electron pairs about the central atom are I. Bond energy i) Bond dissociation energy   If the enthalpy of a reaction is known (H), the bond dissociation energies for bonds that are formed and broken during the rxn can be determined ii) Ionic bonds  The strength of an ionic bond can be determined by using coulombs law   Ionic bonds are typically stronger than covalent bonds, but they can often be cleaved more readily by the addition of a protic solvent 3) Intramolecular Features A. Encompass anything that affects the stability of a molecule and the sharing of electron density beyond the localized region between the two neighboring bonding atoms B. Five factors that dictate the chemical reactivity of a compound: i) Resonance ii) Inductive effect iii) Steric interactions iv) Aromaticity v) Hybridization C. Resonance i) Electron density shifts through molecule via pi bonds ii) Overlapping p orbitals - Pi bonds and adjacent lone pairs iii) iv) Resonance outweighs inductive effect (more stabilizing) D. Inductive effect i) Induces charge separation in a molecule ii) A highly electronegative atom pulls electron density from its neighbor which in turns pulls electron density from its neighbor iii) Most often considered when molecule has a halogen iv) Electronegative atoms thus increase a compounds acidity (because it makes the conjugate base more stable), increase its electrophilicity (because it makes the compound want electrons), decrease its basicity, or decrease its nucleophilicity i. An electronegative atom that stabilizes a charge (or lone pair) will make it a worse nucleophile, and i. An electronegative atom that stabilizes a charge (or lone pair) will make it a worse nucleophile, and thus a shitty base ii. The greatest reaction rate (the fastest reaction) is observed with the best nucleophile. v) When lone pairs exist on neighboring atoms, we call this the alpha effect. The lone pairs are repelling each other, and this is a destabilizing effect. vi) ALKYL GROUPS ARE ELECTRONDONATING! vii) viii) F>0>N>Cl>Br>I>S>C>H "Fonclbrisch" E. Steric Hindrance i) Substituents with axial orientation experience greater steric hindrance i. F. Aromaticity i) Huckle Rule: 4n+2π electrons in a continuous, overlapping ring of p orbitals ii) Antiaromatic: 4n 4) Fundamental reactivity A. Brønsted-Lowry Acid-Base reaction (Proton transfer) i) Acid = proton donor
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