Honors Organic Chemistry Chapter 1 Book Notes

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Department
Chemistry
Course
CHEM 2331H
Professor
T.Andrew Taton
Semester
Fall

Description
Chapter 1 Book Notes 1.1) The Origins of Organic Chemistry  Organic chemistry: the chemistry of carbon compounds o Carbon bonds strongly to carbon or other elements makes chains and rings to form variety of molecules  Organic: derived from living organisms o Organic chemistry originally studied natural products/compounds of living organisms (ex: sugar, urea, starch, waxes, plant oils) o Vitalism: belief that natural products need vital force to create them o Inorganic chemistry is about gases, rocks, minerals, etc.  But then later discovered that organic compounds can be synthesized from inorganic compounds  vital force theory o But vitalism still continues little bit since some believe that natural/plant-derived substances are more healthful than artificial/synthesized substances o Natural and artificial substances have same elements/molecules, but artificial ones have less C than natural because for artificial, it decayed more  can use C 14 dating to tell them apart  All organic compounds contain carbon but not all carbon compounds are organic (although most are organic) 1.2) Principles of Atomic Structure  1.2A) Structure of the Atom + - o Atoms are made of protons (p ), neutrons (neutral), and electrons (e ) o Protons and neutrons have similar masses together in the nucleus, where almost all of atom’s mass is o Protons and electrons have same magnitude of charge o Electrons determine chemical bonding and reactions o Elements are distinguished by number of protons (atomic number) o Isotopes: same # protons, different # neutrons  1.2B) Electronic Structure of the Atom o Electrons form bonds  determine structure of molecules o Electrons are so small show properties of particles and waves…more wavelike o Orbitals: probable (mathematical descriptions) locations of electrons bound to nuclei; allowed energy state of electrons  Heisenberg Uncertainty Principle  can’t tell exactly where electron is  Electron density: probability of finding electron in particular part of orbital o Most common elements of organic compounds are found in first 2 rows of periodic table their electrons are in the first 2 shells  Shells: defined by principle quantum number n (n = 1 is lowest energy level, closest to nucleus, holds less atoms…vice versa for high n values) o 1 shell has only 1s orbital (sphere)  nondirectional, only distance from nucleus matters  Electron density graph: o 2 shell: 2s (bigger, density shifted farther away (at node) sphere, higher energy than 1s) and 2p orbitals  2p: 3 perpendicular dumbells, more energy than 2s but same energy (degenerate orbitals) for all 2p nodes, nodal planes (perpendicular to dumbbell, 0 electron density) at nucleus o Pauli Exclusion Principle: each orbital can have 2 max electrons  S can hold 2  P can hold 6  D…10  1.2B) Electronic Configurations of Atoms o Aufbau principle: fill electrons from lowest to highest energy levels o Valence electrons: electrons in the outermost energy level/shell; same as the group (column) number of atom o Hund’s Rule: electrons go in different orbitals rather than pair up when possible because they repel each other 1.3) Bonding Formation: The Octet Rule  Octet rule: atoms transfer/share atoms to fill outermost electron shells (usually 8, but higher elements can have expanded octet)  1.3A) Ionic Bonding o Tranfer electrons  atoms have opposite charges  attracted to each other o Forms crystal lattice rather than individual molecule o More prevalent in inorganic compounds  1.3A) Covalent Bonding o Electrons shared, more common in organic compounds 1.4) Lewis Structures  Valence electrons = dots; bonding electron pair = dash  Nonbonding electrons: those not shared; if there are 2, it is a lone pair  Lone pairs not necessarily shown in organic molecules 1.5) Multiple Bonding  Single bond (1 pair of electrons shared), Double bond (2 pairs), triple bond (3 pairs)  1 dash, 2 dashes, 3 dashes  *Carbon usually forms 4 bonds (tetravalent), nitrogen forms 3 (trivalent), oxygen forms 2 (divalent)  Valence: # bonds an atom usually forms 1.6) Electronegativity and Bond Polarity  Nonpolar covalent bond: electrons shared equally  Polar covalent bond: electrons not shared equally; attracted to one of the nuclei; one side is partially positive and other is partially negative o Dipole moment: measures how polar the bond is (how separated the charges are)  Electrostatic potential map: red = electron-rich; blue = electron-poor  Electronegativity: the higher, the more electron attraction, and vice versa; increases from left to right in periodic table 1.7) Formal Charges  Helps see charge distribution in a molecule  Formal Charge = [group number] – [nonbinding electrons] – 1/2[shared electrons] 1.8) Ionic Structures  Some molecules have an ionic component; for others, either way works, but one is usually better 1.9) Resonance  1.9A) Resonance Hybrids o Resonance structures/forms: multiple structures of a molecule that may work equally as well  the real molecule is a hybrid of all those forms o Resonance stabilized: ion is more stable because charge is shared between multiple atoms instead of just 1 holding that charge  bond order becomes 1 ½ when between single and double bond o Occurs when there are double bonds o Expressed with ↔ o Some neutral molecules have equal positive and negative formal charges o Individual forms of resonance structures don’t exist…it’s always in hybrid form  1.9B) Major and Minor Resonance Contributors o If there is energy difference between resonance structures, molecule is most likely more like the stable, low energy structure  that is the major contributor o Unstable structure is the minor contributor o Best/most likely structures have low energy, max # octets, max # bonds, least charge separation o Only electrons can be delocalized o Rules for resonance structures  Must be valid Lewis structures  Only placement of electrons can be shifted  # unpaired electrons must remain the same (most stable ones have no unpaired electrons)  Major resonance contributor has lowest energy  Resonance stabilization is most important when it serves to delocalize a charge over 2 or more atoms 1.10) Structural Formulas  Structural formulas: chemical formulas that show which atom is bonded to which o Complete Lewis structures o Condensed structural formulas  1.10A) Condensed Structural Formulas
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