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Chapter 1-6

Organic Chemistry Chapters 1-6 .docx

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York University
CHEM 2020
Sadia Mariam Malik

Chapter 1 Intro [Organic Chemistry the chemistry of carbon compounds] - the only 14y to tell synthesized compounds apart from plant-derived compounds is through C dating. - Plant derived compounds have higher levels of C because they have just been synthesized from the CO in2the air - All organic molecules have one or more carbon atoms, however not all molecules with carbon are organic (ie. Diamond, graphite, carbon dioxide, sodium carbonate, ammonium cyanate. Basic Principles: 1-2A Structure of the Atom - atoms made of : protons(+), neutrons, and electrons (-) - only electrons take part in bonding, number of protons distinguish the molecule - atoms with same number of protons but different number of neutrons are isotopes 1-2B Electron Structure of the Atom - elements chemical properties are determined by the number of protons and the corresponding number of electrons - electrons have both properties of particles and waves, and many times behave more like waves - electrons bound to nuclei are found in orbitals - Heisenberg uncertainty principle we can never determine exactly where the electron is - We can determine the electron density however - Orbitals are allowed energy states w. an associated function that defines the distribution of electron density in space - Atomic orbitals are grouped into shells each identified by a quantum number (n) increasing n = increasing energy - 2p orbitals are slightly higher in energy than 2s orbitals because the average location of the electron is further away from the nucleus - Each p orbital has 2 lobes, with a nodal plane at the nucleus - Nodal plane is a flat region of space w. zero electron density - Orbitals with identical energies are called degenerate orbitals (ex. All 2p orbitals) - Pauli exclusion Principle each orbital can hold a maximum of 2 electrons provided that their spins are paired 1-2C Electron Configuration of Atoms - Aufbau principle shows how to build up the electron configuration of an atoms ground state - Valence electrons are those in the outermost shell - Hunds Rule when there are 2 or more orbitals of the same energy electrons will go in different orbitals rather than pair up in the same orbital. Bond Formations the Octet Rule: 1-3A Ionic Bonding - transferring of electrons to give each of 2 elements a stable orbit is called ionic bonding ; resulting ions have opposite charges and attract eachother - usually resulting in the formation of a large crystal lattice - this is common in inorganic molecules but relatively uncommon in organic molecules *(covalent bonding is the most common) Lewis Structures: - Lewis structures are used to symbolize covalent bonding - Vlance electron is symbolized by a dot, shared pair of e-s are symbolized as a dash - Non-bonding electrons/ lone pair electrons (not shared) usually found in O, N & halogens; F, Cl, Br, I (halogens often have 3 lone pairs) Multiple Bonding: - The number of bonds an atom usually forms is called its valence (ie. Carbon is tetravalent) Electronegativity and Bond Polarity: - Non-polar covalent bond e-s are shared equally between 2 atoms - Polar covalent bond unequal sharing of e-s - Bond polarity is measured by a molecules dipole moment (charge separation and bond length) - Electronegativities are used to predict whether a bond will be polar and the direction of the dipole moment - Higher electronegative atoms generally have higher attraction for the bonding electrons - Electronegativities increase from left to right and decrease going down on the periodic table Formal Charges: - Calaculation: FC= group # - lone pairs shared e-s Ionic Structures - Bonds btwn atoms w. very large Electronegativity differences are usually drawn as ionic Resonance: 1-9A Resonance Hybrids - Resonance structures move only the placement of electrons - Charges can ne delocalized; spread out over two atoms (you get a partial double bond) - This makes the bond more stable and is called a resonance-stabilized cation - ***individual resonance structures do not exist molecule does not resonate btwn forms. It is a hybrid molecule with some characteristics of both 1-9B Major and Minor Resonance Contributors - resonance structures are not equal in estimated energy - the more stable structure is the major contributor (lower E) and the less stable the minor contributor (higher E) - the structure of the actual compound will resemble the major contributor more however minor contributor structures can help to explain the properties of molecules. - Lower energy usually has max bonds, max octects, and min charge seperation - Rules for drawing resonace: -Must be valid lewis structures -Only electrons in bonds and lone pairs can be moved -Nuclei cannot be moved -Bond angles must stay the same -# of unpaired e-s must stay the same - * negative charges are more stable on electronegative atoms (O,N,S) - *resonance is most important when its purpose is to delocalize a charge Structural Formulas: 1-10A Condensed Structural Formulas - Condensed structural formulas are written without showing bonds; atoms bonded to central atom are listed after it - Multiple bond groups may actually be different then suggested in condensed form 1-10B Line Angle Form - Skeletal structure/ stick figure; often used for cyclic & non-cyclic - Nonbonding electrons rarely shown, carbon atom where ever 2 lines meet, oxygen nitrogen and halogens shown Molecular Formulas and Empirical Formulas: - Calculate Empirical Formula: - Assume Sample is 100g (therefore %=grams) - Divide # of grams of each element by its atomic weight (answer in moles) - Divide each # of moles by smallest # to get empirical # molecular formula can be any multiple of empirical formula - Molecular formula can be determined by dividing the molecular weight of the compound by the molecular weight of the empirical formula molecule Arrhenius Acid and Bases: + - Acids dissociate in water to give H O 3ons - Bases disassociate in water to give hydroxide ions OH - Stronger acids and bases disassociate completely + -14 - K =w[ H O 3 [ -OH ] = 1.00 x 10 + -7 - These concentrations are equal in neutral solutions [ H O 3 = [ -OH ] = 1.00 x 10 - Acidicy or basicity is measured by pH =-long [ H 10] 3 + Bronsted-Lowry Acids and Bases: - Acids donate a proton - Bases accept a proton (hydroxide accepts proton to form H O) 2 -these include bases with no hydroxide ions - Conjugate acid and bases
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