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Chapter One

by: Maggie Bruce

Chapter One CHM 1240

Maggie Bruce
GPA 3.86

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These notes cover topics in chapter one.
Organic Chemistry 1
Stanislav Groysman
Class Notes
Organic Chemistry
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This 3 page Class Notes was uploaded by Maggie Bruce on Sunday September 4, 2016. The Class Notes belongs to CHM 1240 at Wayne State University taught by Stanislav Groysman in Fall 2016. Since its upload, it has received 223 views. For similar materials see Organic Chemistry 1 in Chemistry at Wayne State University.


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Date Created: 09/04/16
Ch. 1 Atoms and Molecules: Orbitals and Bonding  Octet Rule o Most main group elements will have eight valence electrons o There are exceptions: hydrogen, boron o Example:  Fluorine: 9 electrons  Electron configuration: F 1s 2s 2p 5  Energy levels: 2p valence 2s valence 1s core  Chemistry reactions only use valence electrons  Fluorine has seven valence electrons but it want eight, so it undergos a favorable reduction to become F-  All valence orbitals are occupied; it has a closed shell  Covalent Bonds o Sharing of a pair of electrons between two atoms  F2  Both fluorines have eight electrons because they are sharing  Lewis Structures o Non-bonding electrons are dots o Bonding electrons are lines o Non-polar bonds  No dipole, same electronegativities  F2 o Polar bonds  Dipole, different electronegativities  HF o Dipole Moment: electrons are drawn to the more electronegative atom o Methane CH 4  Carbon has four valence electrons so it will form four bond to make an octet o Ammonia NH 3  Nitrogen has five valence electrons so it will form three bonds to make an octet o Ethane C 2 6  Both carbons have eight electrons o Ethylene C 2 4  A double bond is needed for the carbons to have eight electrons o Acetylene C 2 2  A triple bond is needed for the carbons to have eight electrons  Constructing Lewis Structures o 1. Draw atoms with valence electrons o 2. Connect valence electrons to make octet  Formal Charge (charge on atom) o Hydrogen gains an electron to become H- o Carbon can gain an electron to become C- o Carbon can lose an electron to become C+  Formal Charge (charge on molecules/ions) o Take normal number of valence electrons, subtract bonding pairs, subtract non-bonding electrons, and then you have the formal charge on that atom. o CH 3  Carbon usually has four valence electron  Subtract the three bonding pairs  Subtract the two non-bonding electrons  The formal charge on carbon is 1-  Resonance Forms and curved arrow formalism o Sometimes more than one Lewis Structure is possible for a molecule o Lewis Structures may not be complete on their own, multiple resonance forms may be needed to best describe a molecule o Resonance helps to describe delocalization of electrons o Curved arrows are used to move pairs of electrons o Double headed arrows mean resonance o Resonance forms are usually equivalent meaning they contribute equally  Resonance does not equal equilibrium o Equilibrium is like Dr. Jekyl and Mr. Hyde, sometimes it is one, sometimes the other o Resonance is like Frankenstein, it is always a human and a monster  Not all resonance forms contribute equally o Example: Formaldehyde o There is a major and minor form  Stability of resonance forms o More bonds means more stable o Less separation of charge means more stable o Delocalization of electrons over many atoms (spreading the charge) means more stable o Electronegativity is important, resonance forms with more electrons on the more electronegative atoms are more stable o Equivalent resonance forms contribute equally to the resonance hybrid o All resonance forms must have the same number of paired and unpaired electrons, do not break pairs o Do not break single bonds  Molecular orbitals o Lewis structures and resonance describe chemical bonds but, o Molecular orbital theory better describes chemical bonds o Atomic orbitals combine to form molecular orbitals o Molecular orbitals must always equal atomic orbitals  Bond Strength o Energy that holds molecules together can be quantified o Combining atoms leads to enthalpy change o H + H = H 2eltaH = -104 Kcal/mol o Negative enthalpy (exothermic) means the product is more stable than the reactants o Positive enthalpy (endothermic) means the product is less stable than the reactants o H = H + H deltaH= 104 Kcal/mol 2  Reaction Profile o Plot of energy of products/reactants vs. reaction progress o Exothermic starts higher and gets lower o Endothermic starts lower and gets higher  Bond Dissociation Energy (BDE) o Amount of energy that must be applied for homolytic bond cleavage o Homolytic bond cleavage; atoms split electrons equally o Fish hook arrows move one electron


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