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Mebane; Organic Chem ;Aug 22-28 Notes

by: Jasmine Nord

Mebane; Organic Chem ;Aug 22-28 Notes CHEM 3010

Marketplace > University of Tennessee - Chattanooga > CHEM > CHEM 3010 > Mebane Organic Chem Aug 22 28 Notes
Jasmine Nord
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Notes for Prof. Mebane Chem 3010 class. There are a few pictures, and some urls that will give you examples. There are also key words that you can google. This week was sort of a review from Gen ch...
Organic Chemistry I
Class Notes
Organic Chemistry, Chemistry, Chem, 3010, Mebane




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This 3 page Class Notes was uploaded by Jasmine Nord on Thursday August 25, 2016. The Class Notes belongs to CHEM 3010 at University of Tennessee - Chattanooga taught by Dr.Mebane in Summer 2016. Since its upload, it has received 68 views. For similar materials see Organic Chemistry I in CHEM at University of Tennessee - Chattanooga.


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Date Created: 08/25/16
Organic Chem Chap 1 Notes Week 8/22­26 1) Periodic table a) groups vs. periods i) groups (1) vertical columns (2) groups 1­8 are most important ii) periods (1) horizontal rows iii) For a really good periodic table go to; .  (1) it’s all color coded and everything is  marked thoroughly b) atomic number i) Atomic numbers are based on the number of protons an element has. ii) Usually in the upper right hand corner  (1) iii) Hydrogen (H) and Helium (He) are the most abundant c) Mass number i) (Protons)+ (Neutrons) = mass number ii) Electrons don’t really matter when it comes mass number iii) However the atoms size in dependent on electrons. NOT on the nucleus d) Isotopes i) Elements that have the same atomic number but different mass number 12 13 14 (1) 6C    6 C    6C ii) There is more C­12 than C­13 or C­14 (1) C­14 is very radioactive iii) Valance electrons (1) Outer most electrons (2) During chemical bonding these are the only ones involved 2) Quantum mechanics a) Schrodinger wave equation i) Also known as orbital (1) Google for examples (a) orbitals periodic table (b) orbitals shapes ii) There are 4 orbitals (1) (S,P,D,F) orbital (2) We won’t deal much with the D and F orbital. Since carbon is our main  focus, we’ll mostly deal with S and P orbital iii) S orbital is spherical (1) 1s is closest to the nucleus iv) P orbital starts at 2p. Which is more dumbbell shape v) Allows us to find how far electrons are from the nucleus vi) A node is a place where there no electrons b) Electron Configurations i) Describes the number of atomic orbital occupied by electrons at its lowest  energy state. (1) Carbons electron configuration (a) 1s 2s 2p (hopefully this looks familiar, because this is very important) c) Aufbau principle i) The German word for building up ii) This principle stated that, an electron will ALWAYS go into lowest most  available orbital (1) So an electron would is more likely to go to a (1s) orbital than a (3d)  orbital because (1s) is a lower energy state iii) Google “Aufbau principle” to see examples  d) Pauli exclusion principle i) This principle states that no more than 2 electrons can be in an orbital. ii) The electrons also need to have opposite spins e) Hund’s rule i) An electron would rather occupy an empty orbital than share with another  electron. ii) This only applies if there is 2 or more orbital (1) This rule ALWAYS applies when you go to 2s and so on. f) Closed shell electron configuration i) All elements want to look like noble gases (group 8) and have all their orbital  full.  So they will share electrons with each other in hopes to look like their  closest noble gas ii) Noble gases don’t like to share electrons because they have full orbital g) Octet rule i) This is when an atom is most stable because they 8 valance electrons h) Valance vs. Core electrons i) Valance electrons are the outer most electrons ii) Core electrons are the inner electrons 3) Ionic and covalent bonding a) Ionic bonding i) Metals generally react with nonmetals to form ionic compounds (1) Example;  NaCl ii) A chemical bond by transfer of electrons iii) Elements do this in hopes to fill their octet iv) Ionic compounds have very high melting point because it’ll take a lot of  energy for these compounds to come apart. b) Electrostatic bonding i) Creating opposite charged partials ii) Electropositive partials are ready to lose electrons iii) Electronegative partial are ready to gain electrons iv) Example;  Calcium is will to lose electrons to look like Argon. Making  calcium electropositive c) Covalent bonding i) Elements sharing electrons to be bonded to each other ii) Elements can be either polar or non­polar covalent bonds (1) Polar bonds are when each side partically posivily or negattivly vharged (a) Example; HCL (2) Non­polar bonds are when the charge from each element is equal (a) H2 or Br2 d) Electronegativity  i) The pull on a nucleus on a bonded pair of electrons ii) Fluorine is the most electronegative iii) As you go across the periodic table electronegativity increases iv) As you go down the periodic table electronegativity decreases 4) Practice problems a) Pg 5 i) #1,2 b) Pg 7 i) #4­6 5) Important Pages in the textbook a) Pg 5 i) Table 1.1 (orbital shells) b) Pg 6 i) Table 1.2 (examples of electron configurations) c) Pg 10  i) Table 1.3 (electronegativity chart)


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