Organic Chemistry -- Week 1 Material
Organic Chemistry -- Week 1 Material Chem 2410
Popular in Organic Chemistry
Popular in Chemistry
This 16 page Class Notes was uploaded by Georgia King on Wednesday August 24, 2016. The Class Notes belongs to Chem 2410 at Saint Louis University taught by Paul Bracher in Fall 2016. Since its upload, it has received 88 views. For similar materials see Organic Chemistry in Chemistry at Saint Louis University.
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Date Created: 08/24/16
Organic Chemistry—Week 1 Material Chapter 1: Structure and Bonding Section 1.1 The Periodic Table • Atoms: the building blocks of all matter o Nucleus: contains positively charged protons and uncharged neutrons o Electron cloud: composed of negatively charged electrons ü In a neutral atom, the number of protons in the nucleus equals the number of electrons. • Atomic number: the number of protons in an atom • Charged ions: o Cations: positively charged ions, have fewer electrons than its neutral form o Anions: negatively charged ions, have more electrons than its neutral form • Isotopes: two atoms of the same element having a different number of neutrons • Mass number: the total number of protons and neutrons in the nucleus of an atom • Deuterium (D): the isotope of hydrogen o Most common has one proton and no neutrons o .02% have one proton and one neutron ü means notesition • Atomic weight: the weighted average of the mass of all isotopes of a particular element, reported in atomic mass units (amu) • Periodic table: an arrangement of elements in groups of similar properties o Elements in the same row are similar in size o Elements in the same column have similar electronic and chemical properties ü Most elements that are routinely seen in organic compounds are located in the first and second row of the periodic table. ü Electrons are first added to the shells closest to the nucleus. • Orbitals: subshells contained within each electron shell; a region of space that is high in electron density o s, p, d, f… o The redistribution of electrons among nuclei is what controls chemical reactions o Most reactions occur when a filled orbital overlaps with an empty orbital to form a new bonding interaction ü An s-orbital has a sphere of electron density. o Lower in energy ü A p-orbital has a dumbbell shape and contains a node of electron density. o Higher in energy • means definition ü means notes • Valence electrons: the outermost electrons that are more loosely held than the electrons closer to the nucleus and participate in chemical reactions Section 1.2: Bonding • Bonding: the joining of two atoms in a stable arrangement ü Bonding always leads to lowered energy and increased solubility. • Compounds: two or more elements joined ü Through bonding, atoms attain a complete outer shell of valence electrons. ü Through bonding, atoms attain a stable noble gas configuration of electrons o Have full octets o Atoms will gain, lose, or share electrons to get this configuration ü A first-row element like hydrogen can accommodate two electrons around it. • Octet rule: a second-row element is most stable with eight valence electrons around it. • means definition ü means notes o Shared or unshared Ex. • Ionic bonding: results from the sharing of electrons between nuclei o “opposites attract” o electrostatic attraction between ions o form between atoms with a large difference in electronegativity (although there is no specific turning point value) • Covalent bonding: results from the sharing of electrons between nuclei ü Sometimes a compound can have both covalent and ionic character. ü Ionic bonds usually occur between elements on the far left side of the periodic table with elements on the far right side of the periodic table. o Resulting ions are held together by extremely strong electrostatic interactions o Salts (ex. NaCl) ü The transfer of electrons forms stable salts composed of cations and anions. ü Covalent bonding occurs with elements like carbon in the middle of the periodic table, which would otherwise have to gain or lose several electrons to form an ion. ü A covalent bond is a two-electron bond o Form molecules ü For second row elements…. Predicted # of bonds = 8 – # of valence electrons Nonbonded pair of electrons = unshared pa ir – lone pair • means definition ü means notes • Lone pairs: unshared electrons o Represented by a pair of dots • Lewis structures: electron dot representations for molecules (communicate composition and structure/shape of molecules) 1. Draw only valence electrons 2. Give every second row element no more than 8 electrons 3. Give each hydrogen 2 electrons ü After placing all electrons in bonds and lone pairs, use a lone pair to form a multiple bond if an atom does not have an octet. • Formal charge: the charge assigned to individual atoms in a Lewis structure Formal charge = # of valence electrons - # of electrons an atom “owns” *An atom “owns” all of its unshared electrons and half of its shared electrons ü The sum of the formal charges on the individual atoms equal the net charge of the molecule or ion. ü If no formal charge is labeled, it is assumed to be zero. Section 1.4: Isomers • Isomers: different molecules having the same molecular formula and charge ü Nature favors states of the lower potential energy. ü Nature does not like to build charge. • means definition ü means notes • Constitutional isomers: isomers that have the same molecular formula, but the connectivity of their atoms is different • Stereoisomers: isomers that differ only in the special arrangements of their atoms o Not superposable Section 1.5: Exceptions to the Octet Rule ü Most of the common elements in organic compounds (C, N, O, and the halogens) follow the octet rule. ü Exceptions: o Hydrogen o Boron o Beryllium o Phosphorus o Sulfur Section 1.6: Resonance • Resonance structures: two Lewis structures having the same placement of atoms but a different arrangement of electrons o Each contributes some character to the resonance hybrid o Are NOT two “resonating” structures of a molecule • Resonance hybrid: the true structure, a composite of both resonance structures • means definition ü means notes o More stable than any resonance structure because it delocalizes electron density over a larger volume o Each resonance structure contributes equally to the hybrid • Resonance stabilized: when a molecule has two or more resonance structures ü Resonance structures are NOT isomers o Resonance structures only differ in the arrangement of electrons. ü Two resonance structures differ in the position of multiple bonds and nonbonded electrons. The placement of atoms and single bonds always stays the same. ü DO NOT give second-row elements (B, C, N, O) more than a full octet! ü Two resonance structures must have the same number of unpaired electrons. ü Resonance structures must be valid Lewis structures. • Curved arrow notation: a convention that shows how electron position differs between two resonance forms o “push” electrons into new arrangements o not just used for resonance structures o must then reassign formal charges • Major contributor: the “better” resonance structure that the hybrid looks similar to o ”Better” structures have more bonds and fewer charges • means definition ü means notes o Most stable – full octets o Have any negative charge associated with the most electronegative atoms • Minor contributors: all other resonance structures ü 2 Differences Between Each Resonance Structure: o position of multiple bond o site of a charge Section 1.7: Determining Molecular Shape • Bond length: the average distance between the centers of two bonded angles o decreases across a row of the periodic table as the size of the atom decreases o increases down a column of the periodic table as the size of an atom increases § longer = weaker & shorter = stronger • Bond angle: determines the shape around any atom bonded to two other atoms • VSEPR (valence shell electron pair repulsion) theory: electrons repel each other, so the most stable arrangement keeps groups as far away from each other as possible • Linear molecule: an atom surrounded by two groups o Bond angle: 180° ü Ignore multiple bonds when predicting geometry; count only atoms and lone pairs • Trigonal planar: an atom surrounded by three groups • means definition ü means notes o Bond angle: 120° • Tetrahedral: an atom surrounded by four groups o Bond angle: 109.5° • Square planar: an atom surrounded by four groups o Bond angle: 90° ü Rules for Drawing Molecular Shapes: o A solid line = bond in the plane o A wedge = bond in the front of plane o A dashed line = bond behind the plane • Trigonal pyramidal: an atom that has three groups attached as well as a lone pair o Bond angle: 107° • means definition ü means notes • Bent: a type of linear shape with a different bond angle o Bond angle: 105° ü The lone pairs take up slightly more space to ease electron-electron repulsion • Condensed structures: o All atoms are drawn in, but the two electron bond lines are generally omitted o Atoms are usually drawn next to the atoms to which they are bonded o Parentheses are used around similar groups bonded to the same atom o Lone pairs are omitted (applicable to atoms other than C and H) o Have to write in double bonds § Exception: NOT when C is bonded to O because it’s obviously a double bond Ex. • means definition ü means notes • Heteroatom: any atom that is not C or H • Skeletal structures: o Assume there is a carbon atom at the junction of any two lines or at the end of any line o Assume there are enough hydrogens around each carbon to make it tetravalent o Draw in a heteroatoms and the H’s directly bonded to them o Label formal charges o Included multiple bonds o Lone pairs implied Ex. ü The charge determines the number of lone pairs. Negatively charged carbon atoms have one lone pair and positively charged carbon atoms have none. ü Skeletal structures often leave out lone pairs on heteroatoms. Section 1.9: Hybridization • Sigma ( ) bond: a bond that concentrates electron density between nuclei o Cylindrically symmetrical because the electrons forming the bond are distributed symmetrically about an imaginary line connecting two nuclei • means definition ü means notes o A sigma bond concentrates electron density on the axis that joins two nuclei. All single bonds are sigma bonds. ü If two atoms are not able to form a bond together, they experience anti-bonding, which is actually weakening the bond between the atoms and increasing the energy of the molecule, encouraging them to separate • Excited state: the higher energy electron configuration • Hybridization: the combination of two or more atomic orbitals to form the same number of hybrid orbitals, each having the same shape and energy ü Sp3 hybrids are formed from one s orbital and three p orbitals (hybrid orbitals) • means definition ü means notes ü During hybridization, a given number of atomic orbitals hybridize to form an equivalent number of hybrid orbitals o One 2s orbital and three 2p orbitals form four sp3 orbitals o One 2s orbital and two 2p orbitals form three sp2 orbitals o One 2s orbital and one 2p orbital form two sp orbitals Section 1.10: Ethane, Ethylene, and Acetylene • Ethane: CH3CH3 o Each carbon is tetrahedral o Each carbon is sp3 hybridized o All bonds are sigma bonds o Rotation can occur around the central C—C sigma bond • Ethylene: C2H4 • means definition ü means notes o Each carbon is trigonal planar o Each carbon is sp2 hybridized o Each C—H bond is a pi bond and a sigma bond o Rotation cannot occur around the C=C bond • Acetylene: C2H2 o Each carbon is linear o Each carbon is sp hybridized o Each C—H bond is a sigma bond o The triple bond is two pi bonds and a sigma bond Section 1.11: Bond Length and Bond Strength ü As the number of electrons between two nuclei increases, bonds become shorter and stronger o Thus, triple bonds are shorter and stronger than double bonds, which are shorter and stronger than single bonds ü Double bonds consisting of both sigma and pi bonds are strong. However, the pi component of the double bond is much weaker than that of the sigma bond. • Percent s-character: the fraction of a hybrid orbital due to the 2s orbital used to form it o Sp = 1 2s orbital / 2 hybrid orbitals = 50% o Sp2 = 1 2s orbital / 3 hybrid orbitals = 33% • means definition ü means notes o Sp3 = 1 2s orbital / 4 hybrid orbitals = 25% Increased percent s-characterà increased bond strengthàdecreased bond length Section 1.12: Electronegativity and Bond Polarity • Electronegativity: a measure of an atom’s attraction for electrons in a bond o Increases across a period as the nuclear charge increases (excluding noble gases) o Decreases down a group as the atomic radius increases, pushing valence electrons farther from the nucleus ü Electronegativity is used as a guideline to indicate whether the electrons in a bond are equally shared or unequally shared between two atoms o Bonding between atoms of different electronegativity values results in the unequal sharing of electrons ü Usually a polar bond will be one in which the electronegativity difference between two atoms is ≥ .5 units • Electrostatic potential map: a color corded map that illustrates areas of high and low electron density 1.13: Polarity of Molecules ü Determining if a molecule is polar: o Identify all of the polar bonds and their directions • means definition ü means notes o Determine geometry o Decide if individual dipoles cancel or reinforce each other 1.14: L-Dopa—A Representative Organic Molecule • L-Dopa: a drug used to treat Parkinson’s disease • means definition ü means notes
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