ORGO EXAM #1 (Chapters 1-4)

Base Strength

-The conjugate base of a strong acid must be a weak base. On the other hand, if an acid is weak, its conjugate base is a strong base. -See Table 1-5 -In the reaction of an acid with a base, the equilibrium generally favors the WEAKER acid and base. -The strength of a base is measured just like the strength of an acid, by using the equilibrium constant of the hydrolysis reaction: A- (conjugate base) + H2O <=> (Kb) HA (conjugate acid) + OH- - Kb = base-dissociation constant for the base A- Kb = [HA] [OH-] / [A-] pKb = -logKb (Ka)(Kb) = ( [H3O+][A-] /[OH-] )([HA][OH-] /[A-] ) = [H3O+][OH-]= 1.0 x 10^-14 Ka x Kb = 10^-14 pKa + pKb = -log(10^-14)

1-1

1-1, 1-26 through 1-29

Structural Formulas

1-10: -Show which atoms are bonded to which

Condensed Structural Formulas

1-10A: -Written without showing all the individual bonds -See Table 1-2 -When a condensed structural formula is written for a compound containing double or triple bonds, the multiple bonds are often drawn as they would be in a Lewis structure. -Condensed structures are assumed to follow the octet rule even if the condensed notation does not show the bonding -See Table 1-3

Line Angle Formulas

1-10B: -AKA SKELETAL STRUCTURES or STICK FIGURES -See Table 1-4

Molecular Formulas and Empirical Formulas

1-11: -MOLECULAR FORMULA: gives the number of atoms of each element in one molecule of the compound => example: CH3CH2CH2CH2OH

Arrhenius Acids and Bases

1-12: -ARRHENIUS ACIDS: Substances that dissociate in water to give hydronium ions => stronger acids dissociate to a higher degree -ARRHENIUS BASES: Substances that dissociate in water to give hydroxide ions => stronger bases dissociate to a higher degree -The acidity or basicity of an aqueous solution is measured by the concentration of H3O+. This value also implies the concentration of OH- because these two concentrations are related by the water ion-product constant: Kw = [H3O+] x [OH-] = 1.00 x 10^-14 M^2 (at 25°C) -In a neutral solution, the concentrations of [H3O+] and [OH-] are equal: [H3O+] = [OH-] = 1.00 x 10^-7 M, in a neutral solution -Acidic and basic solutions are defined by an excess of H3O+ or OH-: acidic: [H3O+] > 10^-7 M and [OH-] < 10^-7 M basic: [H3O+] < 10^-7 M and [OH-] > 10^-7 M

Bronsted-Lowry Acids and Bases

1-13: BRONSTED-LOWRY ACID: any species hat can donate a proton BRONSTED-LOWRY BASE: any species that can accept a proton -These definitions also include all the Arrhenius acids and bases because compounds that dissociate to give H3O+ are proton donors, and compounds that dissociate to give OH- are proton acceptors. (Hydroxide ion accepts a proton to form water). -In addition to Arrhenius acids and bases, the Bronsted-Lowry definition includes bases that have no hydroxide ions, yet can accept protons-> example -When a base accepts a proton, it becomes and acid capable of returning that proton. When an acid donates its proton, it becomes a base capable of accepting that proton back. one of the most important principles of the Bronsted-Lowry definition is this concept of CONJUGATE ACIDS AND BASES. -Many compounds (water, for instance) can react either as an acid or a base

Acid Strength

1-13A: -The strength of a Bronsted-Lowry acid is expressed as it is in the Arrhenius definition, by the extent of its ionization of water. The general reaction of an acid (HA) with water is teh following: HA (acid) + H2O <=> (Ka) H3O+ + A- (conjugate base) - Ka = ACID-DISSOCIATION CONSTANT; its value indicates the relative strength of the acid => The stronger the acid, the larger the Ka. Strong acids are almost completely ionized in water, and their dissociation constants are greater than 1. Most organic compounds are extremely weak acids. pH = -log[H3O+] pka = -logKa

Structural Effects on Acidity

1-13C: -Look up reasons why -Electronegativity -> as electronegativity increases, acidity and basicity both increase -Size -> as size increases, acidity increases -Resonance Stabilization and Inductive Effects -> read about these

1-14

1-14 and 1-15

Lewis Acids and Bases

1-14: -LEWIS BASES: species with available electrons that can be donated form new bonds -LEWIS ACIDS: species that can accept these electron pairs to form new bonds -Since a Lewis acid ACCEPTS A PAIR OF ELECTRONS, it is called an ELECTROPHILE. A Lewis base is called a NUCLEOPHILE because it donates electrons to a nucleus with an empty (or easily vacated) orbital -The Lewis acid-base definitions include reactions having nothing to do with protons -A nucleophile donates electrons. An elecrophile accepts electrons. Acidic protons may serve as electron acceptors -CURVED-ARROW FORMALISM: used to show the flow of an electron pair FROM THE ELECTRON DONOR TO THE ELECTRON ACCEPTOR (see images)

1-17

1-17 and 1-18

1-2

1-2 through 1-5

1-20

1-20 through 1-25

1-21

1-21, 1-23, 1-25, 1-27, and 1-29

Electronic Structure of the Atom

1-2B: -Often electrons behave more like waves than like particles -Electrons that are bound to nuclei are found in ORBITALS -ELECTRON DENSITY: the probability of finding the electron in a particular pat of the orbital -Atomic orbitals are grouped into different "shells" at different distances from the nucleus => n = 1 -> lowest-energy shell closest to the nucleus; as n increases, shells are farther from the nucleus, higher in energy, and hold more electrons -NODE: a region of zero electron density (see accompanying graphs and images -The 2p orbitals are slightly higher in energy than the 2s, because the average location of the electrons in a 2p orbital is farther from the nucleus -NODAL PLANE: a flat (planar) region of space, including the nucleus, with zero electron density -DEGENERATE ORBITALS: orbitals with identical energies (e.g., Px, Py, Pz)

Electronic Configurations of Atoms

1-2C: -VALENCE ELECTRONS: electrons in the outermost shell

1-31

1-31 and 1-32

1-34

1-34 through 1-36

1-39

1-39 and 1-40

The Octet Rule

1-3: -Elements in the 3rd and higher rows (e.g., Al, Si, P) can have an "expanded octet" of 8+ electrons because they have low-lying d orbitals available

Ionic Bonding

1-3A: -Transfer of electrons

Covalent Bonding

1-3B: -Sharing of electrons -Most common type of bonding in organic compounds

1-42

1-42 through 1-45

Lewis Structures

1-4: -Nonbonding electrons = lone pairs => Help to determine the reactivity of their parent compounds

1-52

1-52 and 1-55

Multiple Bonding

1-5: -VALENCE: The number of bonds an atom usually forms (see following sentence for examples)

1-6

1-6 through 1-10

Electronegativity and Polar Bonding

1-6: -In a NONPOLAR COVALENT BOND, electrons are shared equally -We symbolize the bond polarity by an arrow with its head at the negative end of the polar bond and a plus sign at the positive end -ELECTROSTATIC POTENTIAL MAP (EPM): uses color to represent the calculated charge distribution in a molecule => red = electron-rich region; blue and purple = electron-poor regions; orange, yellow, and green = intermediate levels of electrostatic potential - C-H bonds are usually nonpolar -See Figure 1-6

Formal Charges

1-7: -FORMAL CHARGES: provide a method ofr keeping track, but they *may or may not correspond to real charges* (partial charges are real)

Resonance

1-9: -The different structures are called RESONANCE STRUCTURES or RESONANCE FORMS because they are not different compounds, just different ways of drawing the same compound. The actual molecule is said to be a RESONANCE HYBRID of its resonance forms. -If a charge is DELOCALIZED, it is spread out. -Spreading the positive charge over two atoms makes the ion more stable than it would be if the entire charge were localized only on one atom. We call this a RESONANCE-STABILIZED cation. Resonance is most important when it allows a charge to be delocalized over 2+ atoms. -Remember that individual resonance forms do not exist. The molecule does not "resonate" between these structures (see textbook for mule analogy)

Major and Minor Resonance Conributors

1-9B: -Two or more correct Lewis structures for the same compound may or may not represent electron distributions of equal energy. Although separate resonance forms do not exist, we can estimate their relative energies as if they did. More stable resonance forms are closer representations of the real molecule than less stable ones. -Many stable ions have a positive charge on a nitrogen atom with four bonds. We call the more stable resonance form the MAJOR CONTRIBUTOR, and the less stable form the MINOR CONTRIBUTOR. The structure of the actual compound resembles the major contributor more than the minor. -In drawing resonance forms, we try to draw structures that are as low in energy as possible. The best candidates are those that have the MAXIMUM NUMBER OF OCTETS AND THE MAXIMUM NUMBER OF BONDS. Also, we look for structures with the minimum amount of charge separation. -Rules for drawing resonance forms: 1. All the resonance structures must be valid Lewis structures for the compound 2. Only the placement of the electrons may be shifted from one structure to another (Electrons in double bonds and lone pairs are the ones that are most commonly shifted). ONLY electrons can be delocalized. Nuclei cannot be moved, and the bond angles must remain the same. 3. The number of unpaired electrons (if any) must remain the same. Most stable compounds have no unpaired electrons, and all the electrons must remain paired in all the resonance forms. 4. The major resonance contributor is the one with the lowest energy. Good contributors generally have all octets satisfied, as many bonds as possible, and as little charge separation as possible. Negative charges are more stable on more electronegative atoms, such as O, N, and S. 5. Resonance stabilization is most important when it serves to delocalize a charge over two or more atoms.

Intermolecular Forces

2-10: -When two molecules approach, they attract or repel each other. In general, the forces are attractive until the molecules come so close that they infringe on each other's van der Waals radius. When this happens, the small attractive force quickly becomes a large repulsive force, and the molecules "bounce" off each other. With complicated organic molecules, these attractive and repulsive forces are more difficult to predict. We can still describe the nature of the forces, however, and we can show how they affect the physical properties of organic compounds. Attractions between molecules are particularly important in solids and liquids. In these "condensed" phases, the molecules are continuously in contact with each other. The melting points, boiling points, and solubilities of organic compounds show the effects of these forces. Three major kinds of attractive forces cause molecules to associate into solids and liquids: 1. the dipole-dipole forces of polar molecules 2. the London dispersion forces that affect all molecules 3. the "hydrogen bonds" that link molecules having -OH and -NH groups

Dipole-Dipole Forces

2-10A: -DIPOLE-DIPOLE FORCES: attractive intermolecular forces resulting from the attraction of the positive and negative ends of the dipole moments of polar molecules (permanent dipole moments)

The London Dispersion Force

2-10B: -In nonpolar molecules such as carbon tetrachloride, the principal attractive force is the LONDON DISPERSION FORCE, one of the VAN DER WAALS FORCES. The force arises from temporary dipole moments that are induced in a molecule by other nearby molecules. -Molecules with higher surface area (and less branching) have higher boiling points

Hydrogen Bonding

2-10C: -A HYDROGEN BOND is not a true bond but a particularly strong dipole-dipole attraction. A hydrogen atom can participate in hydrogen bonding if it is bonded to oxygen, nitrogen, or fluorine. Organic compounds do not contain H-F bonds, so we only consider N-H or O-H hydrogens to be hydrogen bonded. -Although hydrogen bonding is a strong form of intermolecular attraction, it is much weaker than a normal C-H, N-H, or O-H covalent bond

Polarity Effects on Solubilities

2-11: -Good descriptions in text! -Polar Solute in a Polar Solvent (Dissolves) => example: sodium chloride and water -Polar Solute in a Nonpolar Solvent (Does Not Dissolve) => the attraction for the ions in the solid for each other are much greater than their attraction for the solvent -Nonpolar Solute in a Nonpolar Solvent (Dissolves) => example: paraffin wax dissolves in gasoline -Nonpolar Solute in a Polar Solvent (Does Not Dissolve)

Hydrocarbons

2-12: -We divide organic compounds into three classes: 1. Hydrocarbons 2. Compounds Containing Oxygen 3. Componds Containing Nitrogen

Alkanes

2-12A: -ALKANES: Hydrocarbons that only contain single bonds -Alkane names generally have the -ane suffix, and the first part of the name indicates the number of carbon atoms methane = 1 ethane = 2 propane = 3 butane = 4 pentane = 5 hexane = 6 heptane = 7 octane = 8 nonane = 9 decane = 10 -CYCLOALKANES: special class of alkanes in the form of a ring -Alkanes undergo few reactions because they have NO FUNCTIONAL GROUP, the part of the molecule where reactions usually occur. Functional groups are distinct chemical units, such as double bonds, hydroxyl groups, or halogen atoms, that are reactive. Most organic compounds are characterized and classified by their functional group. => alkyl group: an alkane portion of a molecule, with one hydrogen atom removed to allow bonding to the re rest of the molecule (see following pictures)

Alkenes

2-12B: -ALKENES: Hydrocarbons that contain carbon-carbon double bonds. A carbon-carbon double bond is the most reactive part of an alkene, so we say that the double bond is the FUNCTIONAL GROUP of the alkene. Alkene names end in the -ene suffix. If the double bond might be in more than one position, then the chain is numbered and the lower number of the two double-bonded carbons is added to the nae to indicate the position of the double bond. => Carbon-carbon double bonds cannot rotate, and many alkenes show geometric (cis-trans) isomerism => Cycloalkenes are also common. Unless the rings are very large, cycloalkenes are always the cis isomers, and the term cis is omitted from the names. In a large ring, a trans double bond may occur, giving a trans-cycloalkene (see following pics)

Alkynes

2-12C: -ALKYNES: hydrocarbons with carbon-carbon triple conds as their functional group. Alkyne names generally have the n-yne suffix, although some of their common names (acetylene, for example) do not conform to this rule. The triple bond is linear, so there is no possibility of geometric (cis-trans) isomerism in alkynes. -In an alkyne, four atoms must be in a straight line. These four collinear atoms are not easily bent into a ring, so cycloalkynes are rare. Cycloalkynes are stable only if the ring is large, containing 8 or more carbons.

Aromatic Hydrocarbons

2-12D: -These compounds may look like cycloalkenes, but their properties are different from those of simple alkenes. These AROMATIC HYDROCARBONS (also called ARENES) are all derivatives of benzene, represented by a six-membered ring with 3 double bonds. This bonding arrangement is particularly stable. -Just as a generic alkyl group substituent is represented by R, a generic aryl group is represented by Ar. When a benzene ring serves as a substituent, it is called a phenyl group, abbreviated Ph.

Organic Compounds Containing Oxygen

2-13: Many organic compounds contain oxygen atoms bonded to alkyl groups. The major classes of oxygen-containing compounds are alcohols, ethers, ketones, aldehydes, carboxylic acids, and acid derivatives.

Alcohols

2-13A: -ALCOHOLS: organic compounds that contain the HYDROXYL GROUP (-OH) as their functional group. The general formula for an alcohol is R-OH. Names of alcohols end in the -ol suffix (see pics)

Ethers

2-13B: -ETHERS: composed of two alkyl groups bonded to an oxygen atom. The general formula for an ether is R-O-R'. (The symbol R' represents another alkyl group, either the same as or different from the first). Ether names are often formed from the names of the alkyl groups and the word "ether" (e.g. diethyl ether).

Aldehydes and Ketones

2-13C: -The CARBONYL GROUP, C=O, is the functional group for both aldehydes and ketones. A KETONE has two alkyl groups bonded to the carbonyl group; an ALDEHYDE has one alkyl group and a hydrogen atom bonded to the carbonyl group. Ketone names generally have the -one suffix; aldehyde names use either the -al suffix or the -aldehyde suffix.

Carboxylic Acids

2-13D: -CARBOXYLIC ACIDS: contain the CARBOXYL GROUP, -COOH, as their functional group. The general formula for a carboxylic acid is R-COOH (or RCO2H). The carboxyl group is a combination of a carbonyl group and a hydroxyl group, but this combination has different properties from those of ketones and alcohols -Systematic names for carboxylic acids use the -oic acid suffix, but historical names are commonly used. => examples: formic acid, acetic acid, propionic acid, butyric acid

Carboxylic Acid Derivatives

2-13E: -Carboxylic acids are easily converted to a variety of acid derivatives. Each derivative contains the carbonyl group bonded to an oxygen or other electron-withdrawing element. Among these functional groups are ACID CHLORIDES, ESTERS, and AMIDES. All of these groups can be converted back to carboxylic acids by acidic or basic hydrolysis.

Organic Compounds Containing Nitrogen

2-14: -The most common "nitrogenous" organic compounds are amines, amides, and nitriles

Amines

2-14A: -AMINES: alkylated derivatives of ammonia -Becaues of their basicity, naturally occurring amines are often called alkaloids. Simple amines are named by naming the alkyl groups bonded to nitrogen and adding the word "amine."

Amides

2-14B: -AMIDES: acid derivatives that result from a combination of an acid with ammonia or an amine -High melting/boiling points

Nitriles

2-14C: -NITRILE: A compound containing the CYANO GROUP, -C=-N (sp hybridized bonding)

2-15

2-15 through 2-22

Wave Properties of Electrons in Orbitals

2-1: -Electrons in atoms are STANDING WAVES -The 1s orbital can be described by its WAVE FUNCTION (Ψ), the mathematical description of the shape of the wave as it vibrates. All of the wave is positive in sign for a brief instant; then it is negative in sign. The electron density at any point is given by the square of the wave function at that point. The plus sign and the minus sign of these wave functions are not charges. The plus or minus sign is the instantaneous phase of the constantly changing wave function. The 1s orbital is spherically symmetrical, and it is often represented by a circle (representing a sphere) with a nucleus in the center and with a plus or minus sign to indicate the instantaneous sign of the wave function (Figure 2-2). -If you gently place a finger at the center of a guitar string while plucking the string, your finger keeps the midpoint of the string from moving. The displacement (movement + or -) at the midpoint is always zero; this point is a node. The string now vibrates in two parts, with the two halves vibrating in opposite directions. We say that the two halves of the string are out of phase: When one is displaced upward, the other is displaced downward. -2p orbital -> two "lobes" separated by a node (nodal plane). The two lobes of the p orbital are out of phase with each other. Whenever the wave function has a plus sign in one lobe, minus in other

Linear Combination of Atomic Orbitals

2-1A: -Atomic orbitals can combine and overlap to give more complex standing waves. We can add and subtract their wave functions to give the wave functions of new orbitals. This process is called the linear combination of atomic orbitals (LCAO). The number of new orbitals generated always equals the number of starting orbitals. 1. When orbitals on DIFFERENT atoms interact, they produce MOLECULAR ORBITALS (MOs) that lead to bonding (or antibonding) interactions 2. When orbitals on the SAME atom interact, they give HYBRID ATOMIC ORBITALS that define the geometry of the bonds

Molecular Orbitals

2-2: See intro

The Hydrogen Molecule and Sigma Bonding

2-2A: -The hydrogen molecule is the simplest example of covalent bonding. As two hydrogen atoms approach each other, their 1s wave functions can add constructively so that they reinforce each other, or destructively so that they cancel out where they overlap. Figure 2-6 shows how the wave functions interact constructively when they are in phase and have the same sign in the region between the nuclei. The wave functions reinforce each other and increase the electron density in this bonding region. The result is a bonding molecular orbital (bonding MO). -The bonding MO in Figure 1-26 has most of its electron density centered ALONG THE LINE CONNECTING THE NUCLEI. This type of bond is called a cylindrically symmetrical bond or a sigma bond. Sigma bonds are the most common bonds in organic compounds. All single bonds in organic compounds are sigma bonds, and every double or triple bond contains one sigma bond. -When two hydrogen 1s orbitals overlap OUT OF PHASE with each other, an ANTIBONDING MOLECULAR ORBITAL results. The two 1s wave functions have opposite signs, so they tend to cancel out where they overlap. The result is a node (actually a nodal plane) separating the two atoms. The presence of a node separating the two nuclei usually indicates that the orbital is antibonding. The antibonding MO is designated σ* to indicate an antibonding (*), cylindrically symmetrical (σ) molecular orbital. -When the 1s orbitals are in phase, the resulting molecular orbital is a σ bonding MO, with lower energy than that of a 1s atomic orbital. When two 1s orbitals overlap out of phase, they form an antibonding (σ*) orbital with higher energy than a 1s atomic orbital.

Sigma Overlap Involving P Orbitals

2-2B: -When two p orbitals overlap along the line between the nuclei, a bonding orbital and an antibonding orbital result. Once again, most of the electron density is centered along the line between the nuclei. This linear overlap is another type of sigma bonding MO.

2-3

2-3 through 2-5

2-35

2-35 and 2-36

2-39

2-39 through 2-42, 2-44

Pi Bonding

2-3: See intro

Single and Double Bonds

2-3A: -A DOUBLE BOND requires the presence of four electrons in the bonding region between the nuclei. The first pair of electrons goes into the sigma bonding MO, forming a strong sigma bond. The second pair of electrons cannot go into the same orbital or the same space. It goes into a pi bonding MO, with its electron density centered above and below the sigma bond. -The combination of one sigma bond and one pi bond is the normal structure of a double bond. -MOs are not as common as bonds formed by hybrid atomic orbitals.

Hybridization and Molecular Shape

2-4: -A common way of accounting for bond angles is the VALENCE-SHELL ELECTRON-PAIR REPULSION THEORY (VSEPR THEORY): Electron pairs repel each other, and the bonds and lone pairs around a central atom generally are separated ye the largest possible angles. An angle of 109.5° is the largest possible separation for four pairs of electrons; 120° is the largest separation for three pairs; and 180° is the largest separation for two pairs. -To explain the shapes of common organic molecules, we assume that the s and p orbitals combine to form hybrid atomic orbitals that separate the electron pairs more widely in space and place more electron density in the bonding region between the nuclei.

sp Hybrid Orbitals

2-4A: -If we combine a p orbital and an s orbital on the same atom, the resulting orbital is called an SP HYBRID ORBITAL. Its electron density. Its electron density is concentrated toward one side of the atom. We started with two orbitals (s and p), so we must finish with two sp hybrid orbitals. The second sp hybrid orbital results if we add the p orbital with the opposite phase. -The result of this hybridization is a pair of directional sp hybrid orbitals pointed in opposite directions. These hybridized orbitals provide enhanced electron density in the bonding region for a sigma bond toward the left of the atom and for another sigma bond toward the right. They give a bond angle of 180°, separating the bonding electrons as much as possible. In general, sp hybridization results in this linear bonding arrangement.

sp2 Hybrid Orbitals

2-4B: -To orient three bonds as far apart as possible, bond angles of 120° are required. When an s orbital combines with two p orbitals, the resulting three hybrid orbitals are oriented at 120° angles to each other (Figure 2-14). These orbitals are called hybrid orbitals because they are composed of one sand two p orbitals. The 120°arrangement is called trigonal geometry, in contrast to the linear geometry associated with sp hybrid orbitals. There remains an unhybridized p orbital perpendicular to the plane of the three hybrid orbitals.

sp3 Hybrid Orbitals

2-4C: -Many organic compounds contain carbon atoms that are bonded to four other atoms. When four bonds are oriented as far apart as possible, they form a regular tetrahedron (109.5°bond angles), as pictured in Figure 2-15. This tetrahedral arrangement can be explained by combining the s orbital with all three p orbitals. The resulting four orbitals are called hybrid orbitals because they are composed of one sand three p orbitals.

Drawing Three-Dimensional Molecules

2-5: -Dashed lines indicate bonds that go backward, away from the reader. Wedge-shaped lines depict bonds that come forward, toward the reader. Straight lines are bonds in the plane of the page.

General Rules of Hybridization and Geometry

2-6: -Rule 1: Both sigma bonding electrons and lone pairs can occupy hybrid orbitals. The number of hybrid orbitals on an atom is computed by adding the number of sigma bonds and the number of lone pairs of electrons on the atom. => Because the first bond to another atom is always as sigma bond, the number of hybrid orbitals may be computed by adding the number of lone pairs to the number of atoms bonded to the central atom -Rule 2: Use the hybridization and geometry that give the widest possible separation of the calculated number of bonds and lone pairs => The number of hybrid orbitals obtained equals the number of atomic orbitals combined. Lone pairs of electrons take up more space than bonding pairs of electrons; thus, they compress the bond angles. -Rule 3: If 2 or 3 pairs of electrons form a multiple bond between two atoms, the first bond is a sigma bond formed by a hybrid orbital. The second bond is a pi bond, consisting of two lobes above and below the sigma bond, formed by two unhybridized p orbitals (see the structure of ethylene in Figure 2-17). The third bond of a triple bond is another pi bond, perpendicular to the first pi bond.

2-7

2-7 through 2-11

Bond Rotation

2-7: -If a bond rotates easily, each molecule can rotate through the different angular arrangements of atoms. If a bond cannot rotate, however, different angular arrangements may be distinct compounds (isomers) with different properties.

Rotation of Single Bonds

2-7A: -Structures differing only in rotations about a single bond are called conformations

Rigidity of Double Bonds

2-7B: -Not all bonds allow free rotation - We can make the following generalization: Rotation about single bonds is allowed, but double bonds are rigid and cannot be twisted -Because double bonds are rigid, we can separate and isolate compounds that differ only in how their substituents are arranged on a double bond

Isomerism

2-8: -ISOMERS: different compounds with the same molecular formula

Constitutional Isomerism

2-8A: -CONSTITUIONAL ISOMERS (or STRUCTURAL ISOMERS): Isomers that differ in their bonding sequence; that is, their atoms are connected differently (see example with C4H10) -Constitutional isomers may differ in ways other than the branching of their carbon chain. They may differ in the position of a double bond or other group or by having a ring or some other feature.

Stereoisomers

2-8B: -STEREOISOMERS: differ only in how their atoms are oriented in space. Their atoms are bonded in the same order, however. -Cis and trans isomers are only one type of stereoisomerism. The study of the structure and chemistry of stereoisomers is called STEREOCHEMISTRY. -CIS-TRANS ISOMERS (GEOMETRIC ISOMERS): differ in the geometry of the groups on a double bond. The cis isomer is always the one with similar groups on the same side of the double bond, and the trans isomer has similar groups on opposite sides of the double bond. -To have cis-trans isomerism, there must be two different groups on each end of the double bond-> see example

Bond Dipole Moments

2-9A: -Bond polarities can range from nonpolar covalent, through polar covalent, to totally ionic

Molecular Dipole Moments

2-9B: -A MOLECULAR DIPOLE MOMENT is the dipole moment of the molecule taken as a whole

Classification of Hydrocarbons (Review)

3-1: -Review first paragraph -A hydrocarbon with no double or triple bonds is said to be SATURATED because it has the maximum number of bonded hydrogens. Another way to describe alkanes, then, is as the class of SATURATED HYDROCARBONS.

Molecular Formulas of Alkanes

3-2: -The molecular formulas increase by two hydrogen atoms each time a carbon is added - -CH2- groups = METHYLENE GROUPS -CnH2n+2 -A series of compounds, like the unbranced alkanes, that differ only by the number of -CH2- groups, is called a HOMOLOGOUS SERIES, and the individual members of the series are called HOMOLOGS

Common Names

3-3A: -Alkanes end in the -ane suffix -If all alkanes had unbranched (straight-chain) structures, the nomenclature would be simple. Most alkanes have structural isomers, however, and we need a way of naming all the different isomers. Or example, there are two isomers for C4H10. The unbranched isomer is simply called butane (or n-butane, meaning "normal" butane), and the branched isomer is called isobutane, meaning an "isomer of butane". -The 3 isomers of C5H12 are called pentane (or n-pentane), isopentane, and neopentane (see pics) -Isobutane, isopentane, and neopentane are COMMON NAMES or TRIVIAL NAMES, meaning historical names arising from common usage. Common names cannot easily describe the larger, more complicated molecules having many isomers, however. The number of isomers for any molecular formula grows rapidly as the number of carbon atoms increases.

IUPAC or Systematic Names

3-3B: -IUPAC: International Union of Pure and Applied Chemistry => This international group has developed a detailed system of nomenclature that we call the IUPAC RULES. The IUPAC rules are accepted throughout the world as the standard method of naming organic compounds. The names that are generated using this system are called IUPAC NAMES or SYSTEMATIC NAMES -Rule #1: The Main Chain => The first rule of nomenclature gives the base name to the compound => Find the longest continuous chain of carbon atoms, and use the name of this chain as the base name of the compound => The longest chain is rarely drawn in a straight line => The groups attached to the main chain are called SUBSTITUENTS because they are substituted (in place of a hydrogen atom) on the main chain. When there are two longest chains of equal length, use the chain with the greatest number of substituents as the main chain. -Rule #2: Numbering the Main Chain => To give the locations of the substituents, assign a number to each carbon atom on the main chain => Number the longest chain, beginning with the end of the chain nearest a substituent => We start the numbering from the end nearest a branch so the numbers of the substituted carbons will be as low as possible (see pic) => Alkyl groups are named by replacing the -ane suffix of the alkane name with -yl. Methane becomes methyl; ethane becomes ethyl. (MEMORIZE FOLLOWING PIC) => The propyl and butyl groups are simply unbranched three- and four-carbon alkyl groups. These groups are sometimes named as "n-propyl" and "n-butyl" groups, to distinguish them from other kinds of (branched) propyl and butyl groups. The simple branched alkyl groups are usually known by common names. The isopropyl and isobutyl groups have a characteristic "iso" grouping, just as in isobutane. => LOOK AT FIGURE 3-2 => The names of the secondary-butyl (sec-butyl) and tertiary butyl (tert-butyl or t-butyl) groups are based on the DEGREE OF ALKYL SUBSTITION of the carbon atom attached to the main chain. In the sec-butyl group, the carbon atom bonded to the main chain is secondary (2°), or bonded to two other carbon atoms. In the tert-butyl group, it is tertiary(3°), or bonded to three other carbon atoms. In both the n-butyl group and the isobutyl group, the carbon atoms bonded to the main chain are primary (1°), bonded to only one other carbon atom. => Haloalkanes can be named just like alkanes, with the halogen atom treated as a substituent. Halogen substituents are named fluoro-, chloro-, bromo-, and iodo- (see pic) -Rule #4: Organizing Multiple Groups => When two or more substituents are present, list them in alphabetical order. When two or more of the same alkyl substituent are present, use the prefixes di-, tri-, tetra-, etc. to avoid having to name the alkyl group twice. => Complex Substituents =>=> Complex alkyl groups are named by a systematic method using the longest alkyl chain as the base alkyl group. The base alkyl group is numbered beginning with the carbon atom (the "head carbon") bonded to the main chain. The substituents on the base alkyl group are listed with appropriate numbers, and parentheses are used to set off the name of the complex alkyl group.

Solubilities and Densities of Alkanes

3-4A: -Alkanes are nonpolar, so they dissolve in nonpolar or weakly polar organic solvents. Alkanes are said to be hydrophobic ("water hating") because they do not dissolve in water. Alkanes are good lubricants and preservatives for metals because they keep water from reaching the metal surface and causing corrosion. -Alkanes have densities around 0.7 g/mL, compared with a density of 1.0 g/mL for water. Because alkanes are less dense than H2O and insoluble in H2O, a mixture of an alkane and water quickly separates into two phases, with the alkane on top.

Boiling Points of Alkanes

3-4B: -The boiling point increases smoothly with increasing numbers of carbon atoms and increasing molecular weights. -In general, a branched alkane boils at a lower temperature than the n-alkane with the same number of carbons.

Melting Points of Alkanes

3-4C: -Like their boiling points, the melting points increases with increasing molecular weight. The melting point graph is not smooth, however. Alkanes with even numbers of carbon atoms pack better into a solid structure, so that higher temperatures are needed to melt them. Alkanes with odd numbers of carbon atoms do not pack as well, and they melt at lower temperatures.

Reactions of Alkanes

3-6: -Alkanes are the least reactive class of organic compounds. Their low reactivity is reflected in another term for alkanes: paraffins.The name paraffincomes from two Latin terms, parum, meaning "too little," and affinis, meaning "affinity." Chemists found that alkanes do not react with strong acids or bases or with most other reagents. They attributed this low reactivity to a lack of affinity for other reagents, so they coined the name "paraffins."

Combustion

3-6A: -COMBUSTION: a rapid oxidation that takes place at high temperatures, converting alkanes to carbon dioxide and water. Little control over the reaction is possible, expect for moderating the temperature and controlling the fuel/air ratio to achieve efficient burning (see pic)

Halogenation

3-6C: -Alkanes can react halogens (F2, Cl2, Br2, I2) to form alkyl halides. For example, methyl reacts with chlorine (Cl2) to form chloromethane (methyl chloride), dichloromethane (methylene chloride), etc. (see pic) -Heat or light is usually need to initiate this HALOGENATION. Reactions of alkanes with chlorine and bromine proceed at moderate rates and are easily controlled. Reactions with fluorine are often too fast to control, however. Iodine reacts very slowly or not at all.

Structure of Methane

3-7A: -The simplest alkane is methane, CH4. Methane is perfectly tetrahedral, with the 109.5° bond angles predicted for an hybridized carbon. Four hydrogen atoms are covalently bonded to the central carbon atom, with bond lengths of 1.09 Å.

Conformations of Ethane

3-7B: -Ethane, the two-carbon alkane, is composed of two methyl groups with overlapping sp3 hybrid orbitals forming a sigma bond between them. -The two methyl groups are not fixed in a single position but are relatively free to rotate about the sigma bond connecting the two carbons.The bond maintains its linear bonding overlap as the carbon atoms turn. The different arrangements formed by rotations about a single bond are called conformations, and a specific conformation is called a conformer("conformational isomer").* Pure conformers cannot be isolated in most cases, because the molecules are constantly rotating through all the possible conformations. -In drawing conformations, we often use NEWMAN PROJECTIONS, a way of drawing a molecule looking straight down the bond connecting two carbon atoms. The front carbon atom is represented by three lines (three bonds) coming together in a Y shape. The back carbon is represented by a circle with 3 bonds pointing out from it. -An infinite number of conformations are possible for ethane, because the angle between the hydrogen atoms on the front and back carbons can take on an infinite number of values. SAWHORSE STRUCTURES picture the molecule looking down toward the carbon-carbon bond. -Any conformation can be specified by its DIHEDRAL ANGLE (theta), the angle between the C-H bonds on the front carbon atom and the C-H bonds on the back carbon in the Newman projection. Two of the conformations have special names. The conformation with is called the eclipsed conformationbecause the Newman projection shows the hydrogen atoms on the back carbon to be hidden (eclipsed) by those on the front carbon. The staggered conformation,with has the hydrogen atoms on the back carbon staggered halfway between the hydrogens on the front carbon. Any other intermediate conformation is called a skew conformation. (see next paragraph and pic)

Conformations of Butane

3-8: -Propane is a 3-carbon alkane, with formula C3H8 -Butane is the four-carbon alkane, with molecular formula C4H10. We refer to n-butane as a straight-chain alkane, but the chain of carbon atoms is not really straight. The angles between the carbon atoms are close to the tetrahedral angle, about 109.5°. Rotations about any of the carbon-carbon bonds are possible. Rotations about either of the end bonds just rotate a methyl group like in ethane or propane. Rotations about the central bond are more interesting, however. (see pic) -When the methyl groups are pointed in the same direction (theta = 0 degrees), the eclipse each other. This conformation is called totally eclipsed, to distinguish it from the other eclipsed conformations like the one at At the butane molecule is staggered and the methyl groups are toward the left and right of each other. This 60° conformation is called gauche. -Another staggered conformation occurs at theta = 180 degrees, with the methyl groups pointing in opposite directions. This conformation is called ANTI because the methyl groups are "opposed".

Torsional Energy of Butane

3-8A: -All the staggered conformations (anti and gauche) are lower in energy than any of the eclipsed conformations. The anti conformation is lowest in energy because it places the bulky methyl groups as far apart as possible. The gauche conformations, with the methyl groups separated by just 60°, are 3.8 kJ (0.9 kcal) higher in energy than the anti conformation because the methyl groups are close enough that their electron clouds begin to repel each other.

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Go through PPT Chapter 1 and Notes (plus discussion notes)

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Go through PPT Chapter 2 and Notes (plus discussion notes)

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Go through PPT Chapter 3 and Notes (plus discussion notes)

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Go through PPT Chapter 4 and Notes (plus discussion notes)

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T or F: ALKANE: a hydrocarbon that contains only single bonds (read rest of intro to Chapter 3)

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T or F: Atomic orbitals on an atom mix to form HYBRID ATOMIC ORBITALS and orbitals on different atoms combine to form MOLECULAR ORBITALS

Organic Chemistry

The chemistry of carbon compounds. Carbon, unlike other elements, is very diverse.