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Orgo 275 Week One Notes

by: Andrea Scota

Orgo 275 Week One Notes CHE 275

Marketplace > Syracuse University > Organic Chemistry > CHE 275 > Orgo 275 Week One Notes
Andrea Scota
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These notes cover in class lectures for the Week of August 29th and also the chapter one textbook notes.
Organic Chemistry 1
J. Hougland
Class Notes
Organic Chemistry, orgo, Chemistry, Chem, notes




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This 8 page Class Notes was uploaded by Andrea Scota on Saturday September 3, 2016. The Class Notes belongs to CHE 275 at Syracuse University taught by J. Hougland in Fall 2016. Since its upload, it has received 4 views. For similar materials see Organic Chemistry 1 in Organic Chemistry at Syracuse University.

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Date Created: 09/03/16
Orgo Notes Week of August 29 th th Class notes August 29 Chapter One: Structure Determines Properties  Organic compounds: o Solvents o Dyes o Synthetic fibers….  Organic reactivity depends on electrons; how they move, change, interact o Think about electrons in waves to help picture them in space Wave Functions  Quantum numbers: o Principle number (n): whole number that specifies shell; related to energy o Angular momentum (s, p, d, f): describes shape  Each orbital (where electrons are) is a solution to wave function  2 Electron Configurations  The nature of science is “lazy”, wants to do what easiest and has lowest energy  Lowest energy orbitals are filled first  Electrons spin up or down  2 electrons per orbital to parallel spin  2 or more empty, fill all singly until all have one and then begin to put 2 in an orbital o to minimize repulsions  Common atoms: H, C, Be, N, O, F, Si, P, S, Cl, Br, I o First two rows and halogens  S orbitals: o Spheres o Symmetric o 1s, 2s o n = 1  P orbitals o 3 different perpendicular orbitals along x, y, and z axis o 1s, 2s, 2p o n = 2  We only care about valence electrons because they are those that interact with one another Nature of Chemical Bond  Form because a compound is more stable  Ionic are the simplest in which the elctron transfers between elements (NaCl)  Organic compounds are rarely ionic o They are commonly covalent in which pairs of electrons are shared  Max stability results when atom isoelectronic with noble gas  Carbon, 4 dots o Electron configuration is 2s 2p 2 o Don’t include 1s because it is not valence Covalent Bonds  Depend on number of valence electrons  Atoms want to get 8 electrons/a full shell  Nonbonding = lone pair = non sharing Chapter One Textbook Notes Structure Determines Properties  Structure is key to everything in chemistry  Properties of substance are determined by atoms contained and the way they are connected 1.1 Atoms ,Electrons, and Orbitals  Review of atomic structure o Atomic number – number of protons o Element (He, Si, N, etc..) o Amu – number of protons and neutrons  Protons have positive charge  Neutrons have no charge  Electrons have negative charge  Wave functions: mathematical descriptions of the energies of an electron wave   According to Heisenberg uncertainty principle we cant tell exactly where electron is, but by squaring the wave function we can tell where it is likely to be  Wave functions can also be called orbitals, which are descrbied by size, shape and directional properties  Orbitals are represented by boundary surfaces, where an electron is most likely to be found  Eleectrons also spin o Spin quantum number is either +1/2 or -1/2 o They may occupy the same orbital only when they have opposite spins o Pauli Exclusion principle  No orbital can have more than 2 electrons in it  P orbitals are “dumb bell” shaped, each contains 2 lobes on the x, y, or z plane; perpendicular to each other o Each lobe may be separated by nodal surfaces where wave function changes and the probablilty of finding electrons is 0  Each orbital is singly filled until theyre all full and then start to put 2, Hunds rule  Valence electrons are the outermost electrons involved in chemical bonding- the most important in orgo  For main group elements, the number of valence elctrons is the group number  After 2p orbital is 2sx, 2sy, and 2sz  The closer an orbital is to a nucleus the less ebergy 1.2 Ionic Bonds  Atoms combine with one another to form compounds  Ionic bonds are gorces of attraction between oppositely charged ions o Cations (+) o Anions(-)  Review: o Elements of left of periodic table lose electrons easier than those on the right o Ionization energy is the energy required to remove an electron o Electron affinity os the energy needed to gain an electron  To find energy change of ionic bond add ionization energy and electron affinity  Forces between charges particles are electrostatic 1.3 Covalent bonds, Lewis Formulas, the Octet Rule  Covalent/shared electron pair allows each atom to have a closed-shell electron configuration  Bond dissociation enthalpy is the energy required to dissolve covalent bonds  Covalent bonding only applies to valence shell electrons  Lone pairs are unshared  Octet rule: in forming compounds, they gain lose, or share electrons to achieve stable electron configurations with 8 valence elctrosn o Compounds can have double/triped bonds to achieve octet rule 1.4 Polar Covalent Bonds, Electronegativity, and Bond Dipoles  Electrons in covalent bonds are not necessarily shared equally- if one has tendancy to attract electrons more than the other, the electron distribution is polarized o “polar covalent bond” o can use electrostatic potential map to show polarization where blue is positive and red is negative  Electronegativity increases from left to right across periodic table and decreases foing down o The bigger the difference the more polar o Flouorine is most, oxygen is second most  Bond dipole moments exist whenever opposite charges are separated from each other, dipole moment  is the product of o Charge (e) x distance (d) 1.5 Formal Charge  If a molecule as a while is neutral, the sum of its positive charges must equal the sum of its negative charges  Formal charges correspond to the difference between number of valence electrons in the neutral free atom vs its bonded state o ( ½ (# covalent bond electrons) + # lone pair electrons ) - # of electrons atom has alone 1.6 Structural Formulas of Organic Molecules: Isomers  STEP 1: Moleculat formula tells us which atoms and how many of each are present  STEP 2: # valence elctrons  STEP 3: connectivity sets out order in which atoms are connected (determined experimentally) o Can have same molecular formula but different connectivity; called isomers  Constitutional (structural): differ in connectivity  Stereoisomer: differ in arrangement in space  Condensed formula: omit all bonds, they are written in sequence  Bond-line formulas: fprmulas in which labels for individual carbons are omitted and hydrogens attached are only shown when necessary 1.7 Resonance and Curved Arrows  Sometimes more than one lewis structure can be written for a molecule, especially if theres a double or triple bond o EX: ozone, O3  The lewis formulas differ in placement of electrons so according to resonance concept, no one Lewis structure is correct – instead the true structre os a resonance hybrid of various formulas called contributing structures  Lewis formulas show atoms ar being localized o Either shared or unshared  Resonance fixes that issue, meaning electrons can be delocalized or shared by several nuclei o Can use curved arrows to show delocalization  Some resonance structires are mire stable/favored than others  Rules of Resonance o 1. Only electrin positions may vary amond various structures o 2. Each structure must have the same number of electrons and the same net charge and same number of lone pairs o 3. The contributing structure that contributes most is the structure with the MOST covalent bonds o 4. Structure with the smallest dipole contributes more o 5. Negative charge resides on most electronegative atom and positive in the least electronegative o 6. Delocalisation stabilized a molecule 1.8 Sulfur and Phosphorus- containing Organic compounds  The octet rule can be exceeded for third-row elements  Resonance contributes in which sulfur contains 10 or 12 electrons in its valence shell are allowed  Phospoorus are allowed 10 in valence shell 1.9 Molecular Geometries  Shapes of molecules can be predicted by the valence shell electron- pair repulsions (VSEPR) 1.10 Molecular Dipole Moments  Molecular dipole moment is the resultant of all individual bond dipoles o Some molecules have polar bonds, but lack molecular dipole because of symmetry  EX: CO 2as no dipole moment but H O do2s 1.11 Curved Arrows, Arrow Pushing, and Chemical Reactions  More common use of curves arrows are to track electron flow in chemical reactions o Double barbed arrows shows the movement of a pair of electrons o Single barbed shows the movement of one electron 1.12 Acids and Bases: The Bronsted-Lowry View  An acid is a proton donor and a base is a proton acceptor  We are mostly concerned wih acids abd bases as reactants, products, and catalysts in organic chemistry  Conjucate acids and bases always differ from their acids and bases by a single proton  In B.L. view, acid doesn’t dissociate in water, water acts as a base and accepts protons, turning it to hydronium ion  Strength of acid is measured by the acididy constant (Ka) o Ka = [H O3] [A-] / [HA] OR pKa = -log Ka 10  The stringer the acid the weaker its conjugate base and vise versa 1.13 How Structure Affects Acid Strength  Depend on: o 1. The strength of the bond to the atom from which the proton is lost o 2. The electronegativity of an atom from which the proton is lost o 3. Electron delocalization in the conjugate base  Bond strength decreases going down group  Strength of an acid depends on atom to which proton is bonded o Strength of H – X bond is important as well as the electronegativity of X  Bond strength more important for atoms in same group  Electronegativity more important for atoms in same row  Electronegative atoms elsewhere in molecule can increase acidity in inductive effects which are structural effects transmitted through bonds  Elctron delocalization is the conjugate base, usually expresses via resonance resonance between contributing lewis formulas o Increases acidity by stabilizing the conjugate base 1.14 Acid Base Equilibria  The position of equilibrium in an acid-base reaction lies to the side of the weaker acid  Useful relationship 1.15 Acids and Bases: The Lewis View  An acid is an electron-pair acceptor, and a base is an electron-pair donor  Atom accepts electron pair in Lewis acid is a hydrogen Class Notes September 2 nd Resonance Hybrids  Structure with resonance forms does not alternate between forms o Instead its hybrid of all resonance forms, called a “resonance hybrid”  EX: benzene (C6H6) has 2 resonance in which all C – C bonds are equilvalent  Which resonance structures contribute more? o Least separation of charge o Full octets  Minor resonance contributers tell you a lot about the chemistry of the system  Exoanded octets of third row atoms o Many phosphorus/sulfur due to hypervalency  Hypervalent : not full octet VSEPR Model  Atoms wish to be far away from each other  Most stable = maximum separation  Electron groups o Lone electrons: more electrostatic density (leads to them appearing “larger”) o Bonding  EX:  Methane, tetrahedral, bond angle 109.5 degrees  Water, bent geometry, is “actually” tetrahedral with lone pair electrons  Ammonia, trigonal pyramidal, electron lone pairs still tetrahedral  Boron triflouride, trigonal planar, max separation  Carbon Dioxide, linear geometry Dipole Moment  Posses dipole moment if centers of positive and negative charge do not coincide o Separation of charge over molecule o Net dipole is total dipole over entire molecule Curved Arrows / Electron Bookeeping  Used to show movement of electrons in chemical reactions  Full head is a pair, fishhook is one  Can be used to show bond breaking amd forming Acids and Bases  Arrhenius, Bronstead-Lowry, Lewis o Arrhenius: acid ionizes in water to give protons; base ionizes to hive hydroxide ions o Brosntead: Acid is proton donor; base is proton acceptor o Lewis: Acid electron pair acceptor; base electron pair donor  Strength measured by pKa o Low # is good = stronger acids  Stronger base is conjugate of weaker base Bond Strength  As go down table, weaker bond stronger acidity  More electronegative = stronger acid


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