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CHM 2110
Professor Ketcha

CHAPTER 1 TEXTBOOK SUMMARY/NOTES th Source: Organic Chemistryby John E. McMurry 8 Edition ISBN-13:9780840054531 -- CHAPTER 1 OVERVIEW –  OChem is relevant because all living organisms surrounding you are made of organic chemicals.  OChem is a significant part of medicine and biological sciences.  Foundation of OChem began in the mid-1700‟s when chemistry evolved from alchemy to modern science.  Early chemists believed that the difference between organic and inorganic compounds was that organic chemicals had some sort of vital force due to the face that they were in living organisms.  In 1816 Michel Chevreul started to break down this ideal of a “vitalistic theory” showed that soap could be broke down into several organic compounds (fatty acids).  In 1828 Friedrich Wohler showed that it was possible to turn salt ammonim cyanate into urea.  By the mid-1800‟s there weren‟t so many believers in this „vitalistic theory‟ because there was more proof that there is no difference between organic/inorganic compounds.  All organic chemicals contain carbon. Organic Chemistry is the study of carbons.  50 million + chemical compounds (most of them with carbon)  Carbon is able to form many compounds because of how easy it is for carbon to bond to itself.  Much of this chapter is an overview of some important elements of general chemistry, that are necessary to know for organic chemistry. There will also be an introduction to some organic chemistry concepts. --CHAPTER 1.1 ―Atomic Structure: The Nucleus‖—  Nucleus: contains subatomic particles called protons (+ charge) and neutrons (no charge, neutral). -10  Electrons: surround neutron in electron cloud, about 2x10 in diameter, 200pm.  Atomic Number (Z): # of protons or electrons.  Mass Number (A): # protons and nucleus.  Isotopes: Same Z but different A. --CHAPTER 1.2 ―Atomic Structure: Orbitals‖—  An orbital is the solution to a wave equation. The volume of the space around the nucleus. Electrons spend 90%+ of its time in an orbital. There are s, p, d, and f orbitals. S and p orbitals are the most common; s orbital is a sphere; p orbital is dumbbell-shaped; d orbitals are the shape of a “lucky” 4 leaf clover. All of the nucleus‟ are right in the middle of these orbitals.  Electron shells surround the nucleus and increase 1s to 2s to 2p to 3s to 3p to 3d, etc.  3 different p orbitals: p , p , and p . x y z "Section 1.5: Valence Bond Theory: Sp, Sp2, and Sp3 Hybrid Orbitals - ChemWiki." Section 1.5: Valence Bond Theory: Sp, Sp2, and Sp3 Hybrid Orbitals - ChemWiki. N.p., n.d. Web. 14 Sept. 2013. .Introd  Nodes have zero electron density and separate different lobes. --CHAPTER 1.3 ―Atomic Structure: Electron Configurations‖—  Ground state electron configuration: orbitals of increasing energy with a specific number of electrons in them. o Aufbau Principle: lowest-energy orbitals fill up first. Order: 1s, 2s, 2p, 3s, 3p, 4s, 3d. o Pauli Exclusion Principle: Max of 2 electrons (must have opposite spins) in orbital, one is indicated by an up arrow and the other by a down arrow. o Hund’s Rule: one electron fills one orbital until all are filled up, and then a second set of electrons fill up the orbitals until they . The first set of electrons are generally spun upward, the second set are usually downward. --CHAPTER 1.4 ―Development of Chemical Bonding Theory‖—  1858 August Kekule and Archibal Couper: Proposed the idea that carbon is tetravalent (forms 4 bonds), so carbon can bond to itself.  1865 Kekule: carbon can form rings of atoms.  1874 Jacobus Van’t Hoff and Joseph Le Bel: introduced 3d look that carbon has specific spatial orientations.  Atoms will always strive to achieve the most stable configuration, that of a noble gas.  Bonds can be made or broken by adding or releasing energy.  The number of electrons left in the valence shell (outside shell) are called valence electrons and are important in determining the stability and energy of an atom.  Covalent Bond: sharing of electrons between molecules.  Ionic Bond: Donating or gaining of electrons between two molecules.  Molecule: all the atoms held together by bonds.  Electron dot structures: valence shell electrons are shown as dots.  Line-bond (or Kekule) structure: bonds are shown as lines, with the understanding that at each end of the bond there are 2 electrons. Lone pairs are represented as dots.  Lone pair electrons: electrons that are not bonded to another atom. --CHAPTER 1.5 ―Describing Chemical Bonds: Valence Bond Theory‖—  Valence Bond Theory: w
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