Organic Chemistry--Week 2 Material
Organic Chemistry--Week 2 Material Chem 2410
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This 13 page Class Notes was uploaded by Georgia King on Sunday August 28, 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 25 views. For similar materials see Organic Chemistry in Chemistry at Saint Louis University.
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Date Created: 08/28/16
Organic Chemistry—Week 2 Material Chapter 2: Acids and Bases Section 2.1: Bronsted-Lowry Acids and Bases • Bronsted-Lowry acid: a proton donor, must contain a hydrogen atom (HA) o Net charge may be 0, ( - ), or ( + ) • Bronsted-Lowry base: a proton acceptor, must be able to form a bond to a proton o Must contain an “available” electron pair o Includes lone pairs and pi bonds o Net charge may be 0 or ( - ) o Usually do not protonate sigma bonds • Salts: a combination of charged species such as OH- and NH2- and cations such as LI+, Na+, or K+ to balance the negative charge • Counterions: cations that balance the negative charge of anions in a salt (aka “spectator ions”) ü Morphine: an addictive pain reliever that can act as both a Bronsted-Lowry acid and base because it contains many hydrogen atoms, but also has lone pairs on O and N, and four pi bonds • Means definition ü Means notes Section 2.2: Reactions of Bronsted-Lowry Lowry Acids and Bases ü A Bronsted-Lowry acid-base reaction results in transfer of a proton from an acid to a base ü In an acid-base reaction, one bond is broken and one is formed o The electron pair of the base forms a new bond to the proton of the acid o The acid loses a proton, leaving the electron pair in the H—A bond on A o A bond MUST be present for an acid to give off a proton ü Loss of a proton from an acid forms its conjugate base. ü Gain of a proton by a base forms its conjugate acid. ü Electron-rich species react with electron-deficient species!!! ü When only one starting material contains a hydrogen, it must be an acid. If only one starting material has a lone pair or a pi bond, it must be a base. ü A starting material with a net positive charge is usually the acid. A starting material with a negative charge is usually the base. Section 2.3: Acid Strength and pKa ü Acid strength is the tendency of an acid to donate a proton. o The more readily a compound donates a proton, the stronger the acid • Means definition ü Means notes ü Acidity constant (Ka): o The stronger the acid, the further the equilibrium lies to the right, and the larger the Ka. o The smaller the pKa, the stronger the acid o Ka and pKa are inversely proportional ü An inverse relationship exists between acidity and basicity. ü pKa provides useful information: o The relative acidity of 2 acids o The relative basicity of the 2 acids’ conjugate bases Section 2.4: Predicting the Outcome of Acid-Base Reactions ü Equilibrium always favors formation of the weaker acid or base. o However, the system is still dynamic at equilibrium, so equilibria should always be considered in BOTH directions. o Just because one side is favored doesn’t mean there isn’t some of the less-favored present. ü An acid can be deprotonated by the conjugate base of any acid having a higher pKa o Can convert neutral atom to an ion o Look at pKa values to determine the strength of acids, and therefore which the reaction favors. § Exploiting pKa differences allows us to control what molecules are soluble in what phases. • Means definition ü Means notes Section 2.5: Factors That Determine Acid Strength ü Anything that stabilizes a conjugate base makes the starting acid more acidic. ü 4 Factors That Affect the Acidity of H—A: o Element effects § Conjugate bases where negative charge is placed on a more electronegative element tend to be favored. o Inductive effects o Electronegative elements tend to be stabilized near negative charge. o Resonance effects o Delocalization of charge by a resonance effect tends to stabilize o Hybridization effects: o S orbitals tend to place electron density closer to the nucleus (which is more stable) than p orbitals ü The more stable the conjugate base, the more acidic the acid. ü Across a period, the acidity of HA increases as the electronegativity of A increases. ü Positive or negative charge is stabilized when it is spread over a larger volume. ü Down group, the acidity of H—A increases as the size of A increases. o Size is more important than electronegativity. ü The identity of A is the most important factor in determining the acidity of the H—A bond. o The most acidic proton in a compound is the one removed first by a base. • Inductive effect: the pull of electron density through sigma bonds caused by electronegativity differences of nearby atoms ü More electronegative atoms stabilize regions of high electron density by an electron-withdrawing inductive effect. • Means definition ü Means notes ü The acidity of H—A increases with the presence of electron- withdrawing groups in A. ü The acidity of H—A increases when the conjugate base is resonance stabilized. ü The higher the percent s-character of the hybrid orbital, the more stable to conjugate base. Section 2.6: Common Acids and Bases ü Organic Acidsà o Acetic acid o P-toluenesulfonic acid (TsOH) • Means definition ü Means notes ü Organic Basesà o Butyllithium o Pyridine o Triethylamine Section 2.7: Aspirin • Sacilylates: a well known group of compounds that includes aspirin ü Aspirin is a synthetic compound. • Means definition ü Means notes Section 2.8: Lewis Acids and Bases • Lewis acid: an electron pair acceptor o Means the molecule must possess an empty, energetically- accessible orbital • Lewis base: an electron pair donor o Structurally the same as Bronsted-Lowry bases o Must have lone pairs or pi bonds ü All Bronsted-Lowry acids are also Lewis acids, but the reverse is not always true. ü Electron-rich species react with electron-deficient species. ü A Lewis acid is also called an electrophile. ü When a Lewis base reacts with an electrophile other than a proton, the Lewis base is called a nucleophile. • Means definition ü Means notes Chapter 3: Introduction to Organic Molecules and Functional Groups 3.1: Functional Group • Heteroatoms: atoms other than carbon or hydrogen o Generally have higher electronegativity • Functional group: an atom or a group of atoms with characteristic chemical and physical properties o The reactive part of the molecule o Functional groups on different molecules tend to behave similarly o Molecules usually have multiple functional groups ü Heteroatoms have lone pairs and create electron-deficient sites on carbon. • Carbon backbone: (or skeleton) the C—C and C—H sigma bond framework that makes up the skeleton of an organic molecule (“R”); bonded to the functional groups o A C bonded to something else • Hydroxy group: a functional group, OH • Means definition ü Means notes • Alkane: an aliphatic hydrocarbon having only C—C and C—H sigma bonds o This is an example of an alkane, not the only molecule. • Alcohol: a compound having the general structure ROH, a hydroxyl group (OH) bonded to an sp3 hybridized carbon atom Section 3.2: An Overview of Functional Groups • Hydrocarbons: compounds made up of only the elements carbon and hydrogen o Aliphatic hydrocarbons § Alkanes: have only C—C sigma bonds and no functional group (maximum number of hydrogens)(saturated) § Alkenes: have a C—C double bond as a functional group (unsaturated) § Alkynes: have a C—C triple bond as a functional group (unsaturated) o Aromatic hydrocarbons: a planar, cyclic organic compound that has p orbitals on all rings and a total of 4n and 2 pi electrons in the orbitals • Means definition ü Means notes ü A functional group determines all of the following properties of a molecule: ü Name ü Bonding and shape ü Type and strength of intermolecular forces ü Physical properties ü Chemical reactivity Section 3.3: Intermolecular Forces • Intermolecular forces: the interactions that exist between molecules o Functional groups present determine type and strength of these interactions ü Ionic compounds contain oppositely charged particles held together by extremely strong electrostatic interactions ü Ion-ion interactions ü Strongest intermolecular forces ü “Opposites attract” • Van der Waals forces: very weak interaction caused by the momentary changes in electron density in a molecule o Include dipole-dipole, dipole-induced dipole, and induced dipole- induced dipole o Only in nonpolar compounds • London forces: attractive forces arising from induced dipoles that form when molecules approach each other o Also called dispersion forces • Means definition ü Means notes o NOT interchangeable with Van der Waals forces, though London forces are a type of Van der Waals forces o Typically weak ü The larger the surface area, the larger the attractive force between two molecules, and the stronger the intermolecular forces ü More space to come in contact with electrons • Polarizability: a measure of how the electron cloud around an atom responds to changes in its electronic environment o Measure of its susceptibility to the inducement of momentary dipoles ü Larger atoms, which have more loosely held valence electrons, are more polarizable than smaller atoms, which have more tightly held electrons. • Dipole-dipole interactions: the attractive forces between the permanent dipoles of two polar molecules o Tend to be strong o More polar = stronger bond • Hydrogen bonding: when a hydrogen atom is bonded to O, N, or F is electrostatically attracted to a lone pair on an O, N, or F in another molecule o A strong dipole-dipole interaction o Strongest besides ion-ion interactions Section 3.4: Physical Properties • Boiling point: the temperature at which a liquid is converted to a gas o Must overcome boiling point for conversion to occur o The temperature where its vapor pressure equals the atmospheric pressure ü The stronger the intermolecular forces, the higher the boiling point § Ionic compounds, therefore, have very high boiling point § For covalent molecules, the boiling point depends of the identity of the functional group ü The larger the surface area, the higher the boiling point. • Means definition ü Means notes ü The more polarizable the atoms, the higher the boiling point. ü Liquids having different boiling points can be separated in the lab using a distillation apparatus. • Melting point: the temperature at which a solid is converted to its liquid phase ü The stronger the intermolecular forces, the higher the melting point. o More energy (heat) is required to overcome stronger intermolecular forces between molecules in the crystal lattice. ü Given the same functional group, the more symmetrical the compound, the higher the melting point. • Solubility: the extent to which a compound (solute) dissolves in a liquid (solvent) to form a homogeneous solution ü “Like dissolves like”: compounds dissolve in solvents having similar kinds of intermolecular forces ü Most organic solvents are either nonpolar or weakly polar. o Most ionic compounds are soluble in water, but are in soluble in organic solvents. • Ion-dipole interactions: when ionic compounds dissolve in polar solvents ü An organic compound is water soluble ONLY if it contains one polar functional group capable of hydrogen bonding with the solvent for every 5 carbon atoms it contains. • Miscible: when two compounds form solutions in all proportions with each other • Hydrophobic: water-fearing o The nonpolar part of a molecule that is not attracted to water • Hydrophilic: water-loving o The polar part of a molecule that can hydrogen bond to water • Means definition ü Means notes Section 3.5: Application: Vitamins • Vitamins: organic compounds needed in small amounts for normal sell function o Cannot be synthesized by our bodies, so they must be obtained in the diet • Fat soluble: dissolves in organic media • Water soluble: dissolves in water • Vitamin A: (retinol) an essential component of the vision receptors in the eyes o Water insoluble o Soluble in any organic medium o Very hydrophobic o Stored in fatty tissue • Vitamin C: a vitamin that the body can synthesize, that is known as a deterrent for all kinds of diseases o Water soluble o Hydrophilic o Excreted quickly • Means definition ü Means notes
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