Organic Chem, Chapter 1, Week 1 Notes
Organic Chem, Chapter 1, Week 1 Notes CHEM141A
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This 3 page Class Notes was uploaded by Hannah Malcomson on Thursday September 1, 2016. The Class Notes belongs to CHEM141A at University of Vermont taught by Wurthmann in Fall 2016. Since its upload, it has received 7 views. For similar materials see Organic Chemistry 1 in Chemistry at University of Vermont.
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Date Created: 09/01/16
ORGO: Chapter 1 Notes Hannah Malcomson 1.1- Organic compounds are difficult to isolate and decompose more easily than inorganic compounds. o Scientists originally thought that organic molecules contained a ‘special vital force’- vitalism. o Vitalism was disproved, current definition of an organic molecule was created Organic molecule: a compound containing carbon atoms o Inorganic molecule: a compound with no carbon molecules 1.2- Constitutional Isomers- compounds with the same molecular formula, but have a different structure (atoms are connected in different orders) Each element will form a (generally) predictable number of bonds o Carbon- 4 electrons (e-), 4 bonds. Tetravalent o Nitrogen- 3 e-, 3 bonds. Trivalent o Oxygen- 2 e-, 2 bonds. Divalent o Hydrogen/ Halogens- 1 e-, 1 bond. Monovalent 1.3- Covalent bond- bond formed as a result of electrons being shared between atoms o Over the course of a covalent bond being formed, energy decreases, which is shown as a -∆H. Lewis structures: drawings to convey the electron interactions between molecules o Steps: Determine the number of valence e- in each atom Connect atoms that are bonded together Connect hydrogen atoms Octet rule- every atom (with a few exceptions), will be likely to form bonds in order to have 8 e- in its valence shell. o Lone pair- a pair of e- on an atom that is not involved in a bond. 1.4- Formal charge- a charge associated with any atom that does not exhibit the appropriate number of valence e-. 1.5- Bond types: o Covalent- diff. in electronegativity is <0.5 o Polar covalent- diff. in electronegativity is between 0.5 and 1.7 o Ionic- diff. in electronegativity is >1.7 Electronegativity: the measure of the ability of an atom to attract e-. Ionic bonds- not actually bonds. Atoms are held together by the forces of attraction between them. 1.6- Quantum Mechanics- theory that is based off the idea that e- act as both a particle and a wave o Wave equation- equation constructed to sum up the total energy in one hydrogen atom (one e- and one proton), that takes into account the idea that the e- is acting as a wave in the electric field of the proton. o Solving the equation creates wave functions, which demonstrate the different energy levels within an atom. o Each wave function relates to an incredibly small area around the nucleus, within which there may or may not be e-. Orbital- region of space that can be occupied by an e-. o Many orbitals are in an electron cloud, which is a single entity o Electron density- probability of finding an e- in a particular region of space around the nucleus Atomic orbital- all orbitals around an atom Node- the location on the wave function where the value is 0 The energy of an e- depends on which orbital it occupies. Orbitals with the same energy level are called degenerate orbitals. 1.7- Covalent bonds are formed by an overlap of atomic orbitals Constructive interference creates a wave with a larger amplitude, as opposed to destructive interference, which reduces the amplitude of the waves According to valence bond theory, a bond is the combined e- density of the atoms as a result of constructive interference. 1.8- Molecular orbital theory- describes a bond as the constructive interference between two overlapping atomic orbitals o Atomic orbitals, or the regions of space associated with one atom, are mathematically combined to produce new molecular orbitals, or regions of space associated with a molecule. The lower energy molecular orbital (bonding orbital) is the result of constructive interference, while the higher energy molecular orbital (antibonding orbital) is the result of destructive interference 1.9- Bond Strength o Single bond < Double bond < Triple bond 1.10- Steric number- sum of all sigma bonds and lone pairs o Geometry of the central atom will depend on the steric number VSEPR theory 1.11- Dipole moment- the center of negative charge and the center of positive charge are separated from one another by a certain distance o Indicative of polarity When dealing with a compound that has more than one polar bond, you have to take the vector sum off all of the different molecules in order to find the molecular dipole moment o Vectors take the magnitude and direction of forces into account when being calculated 1.12- Intermolecular forces (IMF)- attractive forces between molecules o All IMF are caused by the attraction between two opposite charges (electrostatic) o Three types of IMF: Dipole-dipole Fleeting dipole- dipole Also called London-dispersion forces At any moment in a molecule that may have a dipole, the partial positive and negative charges may not coincide, which allows for temporary, weak attractions to form Hydrogen bonding The difference in electronegativities between H and the atom it is bonded to will frequently lead to the formation of a dipole, which means that the H will have a partial positive charge, allowing it to bond with negatively charged atoms. H- bonding is very strong because hydrogen is a very atom, so the charges will be very close together Hydrocarbons- molecules that are made of carbon and hydrogen o Boiling point increases with molecular weight 1.13- “Like dissolves like”- polar compounds are soluble in polar solvents (dipole- dipole), non polar compounds are soluble in non polar solvents (London- dispersion forces) Hydrophobic- not soluble in water Hydrophilic- soluble in water
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