STRUCTURE Dr. Sheppard CHEM 2411 Spring 2015 Klein (2nd ed.) sections 1.8-1.10, 1.12-1.13, 2.7-2.12, 3.2, 3.4-3.5, 3.8-3.9, 4.6-4.13, 4.14, 8.5, 15.16, 21.3
Topics Structure Physical Properties Hybridization Resonance Acids and Bases Conformations of Alkanes and Cycloalkanes Unsaturation Alkene Stability
Structure Drawing organic structures Sigma (s) and pi (p) bonds Single bonds = 2e - = one sigma bond Double bonds = 4e - = one sigma bond and one pi bond Triple bonds = 6e - = one sigma bond and two pi bonds Which bond is shortest? Longest? Weakest? Strongest? Remember formal charges
Ionic Structures Be on the lookout for metals (cations) and ions Example: NaOCH 3 This is a Na + cation and a CH 3 O - anion Example: NH 4 Cl This is a NH 4+ cation and a Cl - anion
Classification of Atoms C atoms can be classified as: Primary (1º) = C bonded to 1 other C Secondary (2º) = C bonded to 2 other C Tertiary (3º) = C bonded to 3 other C Quaternary (4º) = C bonded to 4 other C
Classification of Atoms H atoms are classified based on the type of carbon to which they are attached
Classification of Alcohols and Alkyl Halides Alcohols and alkyl halides are classified based on the type of carbon to which the -OH or X is bonded Classify these alcohols/alkyl halides as 1º, 2º or 3º:
Alkyl Halide Structure In addition to primary, secondary, and tertiary, alkyl halides can be classified as: Geminal Vicinal Vinyl Aryl
Classification of Amines and Amides Amines and amides are classified based on the number of C atoms bonded to the N
Classification of Amines and Amides Classify these functional groups:
Electronegativity and Bond Polarity Electronegativity Ability of atom to attract shared electrons (in a covalent bond) Most electronegative atom = F Differences in electronegativity determine bond polarity Bond polarity How electrons are shared between nuclei Equal sharing of electrons = nonpolar; unequal = polar
Bond Polarity Example: C O What atom is more electronegative (C or O)? More EN atom has partial negative charge (d - ) Less EN atom has partial positive charge (d + ) Arrow shows direction of polarity Nonpolar bonds Any atom with itself C H
Molecular Dipole Moment Overall electron distribution within a molecule Depends on bond polarity and bond angles Vector sum of the bond dipole moments (consider both magnitude and direction of individual bond dipole moments) Lone pairs of electrons contribute to the dipole moment Symmetrical molecules with polar bonds = nonpolar
Intermolecular Forces Strength of attractions between molecules Based on molecular polarity Influence physical properties (boiling point, solubility) 1. Dipole-dipole interactions 2. Hydrogen bonding 3. London dispersions (van der Waals)
1. Dipole-Dipole Interactions Between polar molecules Positive end of one molecule aligns with negative end of another molecule Lower energy than repulsions Larger dipoles cause higher boiling points
2. Hydrogen Bonding Strongest dipole-dipole attraction H-bonded molecules have higher boiling points Organic molecule must have N-H or O-H The hydrogen from one molecule is strongly attracted to a lone pair of electrons on the other molecule
3. London Dispersion Forces van der Waals forces Exist in all molecules Important with nonpolar compounds Temporary dipole-dipole interactions Molecules with more surface area have stronger dispersion forces and higher boiling points Larger molecules Unbranched molecules CH 3 CH 2 CH 2 CH 2 CH 3 n-pentane, b.p. = 36 C CH 3 CH 3 CH CH 2 CH 3 isopentane, b.p. = 28 C H 3 C CH 3 C CH 3 CH 3 neopentane, b.p. = 10 C
Boiling Points and Intermolecular Forces CH 3 CH 2 OH ethanol, b.p. = 78 C CH 3 O CH 3 dimethyl ether, b.p. = -25 C H 3 C N CH 3 CH 3 CH 2 N CH 3 CH 3 CH 2 CH 2 N H CH 3 H H trimethylamine, b.p. 3.5 C ethylmethylamine, b.p. 37 C propylamine, b.p. 49 C CH 3 CH 2 OH CH 3 CH 2 NH 2 ethanol, b.p. = 78 C ethyl amine, b.p. 17 C
Solubility and Intermolecular Forces Like dissolves like Polar solutes dissolve in polar solvents Nonpolar solutes dissolve in nonpolar solvents Molecules with similar intermolecular forces will mix freely
Example Which of the following from each pair will have the higher boiling point? (a) CH 3 CH 2 CH 2 CH 3 CH 3 CH 2 CH 2 OH (b) CH 3 CH 2 NHCH 3 CH 3 CH 2 CH 2 NH 2 (c) CH 3 CH 2 CH 2 CH 2 CH 3 CH 3 CH 2 CH(CH 3 ) 2
Example Will each of the following molecules be soluble in water? (a) CH 3 CO 2 H (b) CH 3 CH 2 CH 3 (c) CH 3 C(O)CH 3 (d) CH 2 =CHCH 3
Example Draw the structure of the alkane with the molecular formula C 5 H 12 that has the lowest boiling point.
Structure of Organic Molecules Previously: Atomic/electronic structure Lewis structures Bonding Now: How do atoms form covalent bonds? Which orbitals are involved? What are the shapes of organic molecules? How do bonding and shape affect properties?
Linear Combination of Atomic Orbitals Bonds are formed by the combination of atomic orbitals containing valence electrons (bonding electrons) Two theories: Molecular Orbital Theory Atomic orbitals of two atoms interact Bonding and antibonding MO s formed Skip this stuff Valence Bond Theory (Hybridization) Atomic orbitals of the same atom interact Hybrid orbitals formed Bonds formed between hybrid orbitals
Let s consider carbon How many valence electrons? In which orbitals? So, both the 2s and 2p orbitals are used to form bonds How many bonds does carbon form? All four C-H bonds are the same i.e. there are not two types of bonds from the two different orbitals How do we explain this? Hybridization
Hybridization The s and p orbitals of the C atom combine with each other to form hybrid orbitals before they combine with orbitals of another atom to form a covalent bond Three types we will consider: sp 3 sp 2 sp
sp 3 hybridization 4 atomic orbitals 4 equivalent hybrid orbitals s + p x + p y + p z 4 sppp 4 sp 3 Orbitals have two lobes (unsymmetrical) Orbitals arrange in space with larger lobes away from one another (tetrahedral shape) Each hybrid orbital holds 2e -
Formation of methane The sp 3 hybrid orbitals on C overlap with 1s orbitals on 4 H atoms to form four identical C-H bonds Each C H bond strength = 439 kj/mol; length = 109 pm Each H C H bond angle is 109.5, the tetrahedral angle
Motivation for hybridization? Better orbital overlap with larger lobe of sp 3 hybrid orbital then with unhybridized p orbital Stronger bond Electron pairs farther apart in hybrid orbitals Lower energy
Another example: ethane C atoms bond by overlap of an sp 3 orbital from each C Three sp 3 orbitals on each C overlap with H 1s orbitals Form six C H bonds All bond angles of ethane are tetrahedral
Both methane and ethane have only single bonds Sigma (s) bonds Electron density centered between nuclei Most common type of bond Pi (p) bonds Electron density above and below nuclei Associated with multiple bonds Overlap between two p orbitals C atoms are sp 2 or sp hybridized
Bond rotation Single (s) bonds freely rotate Multiple (p) bonds are rigid
sp 2 hybridization 4 atomic orbitals 3 equivalent hybrid orbitals + 1 unhybridized p orbital s + p x + p y + p z 3 spp + 1 p = 3 sp 2 + 1 p Shape = trigonal planar (bond angle = 120º) Remaining p orbital is perpendicular to hybrid orbitals
Formation of ethylene (C 2 H 4 ) Two sp 2 -hybridized orbitals overlap to form a C C s bond Two sp 2 orbitals on each C overlap with H 1s orbitals (4 C H) p orbitals overlap side-to-side to form a p bond s bond and p bond result in sharing four electrons (C=C) Shorter and stronger than single bond in ethane
sp hybridization 4 atomic orbitals 2 equivalent hybrid orbitals + 2 unhybridized p orbitals s + p x + p y + p z 2 sp + 2 p Shape = linear (bond angle = 180º) Remaining p orbitals are perpendicular on y-axis and z-axis
Formation of acetylene (C 2 H 2 ) Two sp-hybridized orbitals overlap to form a s bond One sp orbital on each C overlap with H 1s orbitals (2 C H) p orbitals overlap side-to-side to form two p bonds s bond and two p bonds result in sharing six electrons (C C) Shorter and stronger than double bond in ethylene
Summary of Hybridization Hybridization of C sp 3 sp 2 sp Example Methane, ethane Ethylene Acetylene # Groups bonded to C 4 3 2 Arrangement of groups Tetrahedral Trigonal planar Linear Bond angles ~109.5 ~120 ~180 Types of bonds to C 4s 3s, 1p 2s, 2p C-C bond length (pm) 154 134 120 C-C bond strength (kcal/mol) 90 174 231
Hybridization of Heteroatoms Look at number of e - groups to determine hybridization Each lone pair will occupy a hybrid orbital Ammonia: N s orbitals (sppp) hybridize to form four sp 3 orbitals One sp 3 orbital is occupied by the lone pair Three sp 3 orbitals form bonds to H H N H bond angle is 107.3 Water The oxygen atom is sp 3 -hybridized The H O H bond angle is 104.5
Example Consider the structure of thalidomide and answer the following questions: a) What is the hybridization of each oxygen atom? b) What is the hybridization of each nitrogen atom? c) How many sp-hybridized carbons are in the molecule? d) How many sp 2 -hybridized carbons are in the molecule? e) How many sp 3 -hybridized carbons are in the molecule? f) How many p bonds are in the molecule?
Example Consider the structure of 1-butene: a) Predict each C C C bond angle in 1-butene. b) Which carbon-carbon bond is shortest? c) Draw an alkene that is a constitutional isomer of 1-butene.
Resonance Multiple Lewis structures for one molecule Differ only in arrangement of electrons Example: CH 2 NH 2+ ion These are resonance structures/forms Valid Lewis structures (obey Octet Rule, etc.) Same number of electrons in each structure Atoms do not move Differ only in arrangement of electrons (lone pair and p electrons)
Resonance Hybrid These structures imply that the C N bond length and formal charges are different Actually not true; these structures are imaginary Molecule is actually one single structure that combines all resonance forms Resonance hybrid Contains characteristics of each resonance form More accurate and more stable than any single resonance form Lower energy (more stable) because of charge delocalization
Electron Movement Electrons move as pairs Can move from an atom to an adjacent bond, or from bonds to adjacent atoms or bonds Use curved arrows to show e - motion (electron pushing) Start where electrons are, end where electrons are going Connect resonance forms with resonance arrow This is not an equilibrium arrow
Contribution to Hybrid Structure Resonance forms do not necessarily contribute equally to the resonance hybrid They are not necessarily energetically equivalent More stable structures contribute more 1. Filled valence shells 2. More covalent bonds (minimizes charges) 3. Least separation of unlike charges (if applicable) 4. Negative charge on more EN atom (if applicable) Which of these is the major contributor to the resonance hybrid?
Benzene Resonance structures: Curved arrows? Is one structure more stable (contribute more)? Resonance hybrid: All carbon-carbon bonds are the same length Somewhere between C C and C=C
Acetone Resonance structures: Curved arrows? Which structure is the major contributor? Which is(are) the minor contributor(s)? Are any structures not likely to form? Resonance hybrid:
Drawing Resonance Structures Rules: Never break a single bond Only lone pair or p-electrons can move Never exceed an octet for C, O, N, X (or 2e - for H) Patterns: 1. p bonds
Drawing Resonance Structures Patterns: 2. Allylic charges or lone pairs Electrons move towards positive charge!
Drawing Resonance Structures Patterns: 2. Allylic charges or lone pairs Electrons move away from negative charge!
Drawing Resonance Structures Patterns: 3. Lone pair next to positive charge
Summary of Resonance Structures
Examples CH 2 CH CH CH 3 CH 2 CH CH CH 3 O
Examples O NH 2 H H Br
Acids and Bases Two types in organic chemistry 1. Brønsted-Lowry Acid = proton (H + ) donor; base = proton acceptor Some molecules can be both (e.g. water) = amphoteric Reaction will proceed from stronger acid/base to weaker acid/base Acid strength measured by pk a Stronger acid = lower pk a
Brønsted-Lowry Acids and Bases Electron flow in acid-base (proton-transfer) reactions: The reaction mechanism Example:
Brønsted-Lowry Acids and Bases You can predict acid strength without a pk a value Strong acids have weak conjugate bases Weak conjugate bases are stable structures
Brønsted-Lowry Acids and Bases Weak conjugate bases are stable structures Have negative charge on EN atom (within a period) Have negative charge on a larger atom (within a group) Negative charge delocalized by resonance Negative charge stabilized by induction
Brønsted-Lowry Acids and Bases Weak conjugate bases are stable structures Negative charge on sp > sp 2 > sp 3
Example Which is the stronger acid in each pair? a) H 2 O or NH 3? b) HBr or HCl? c) CH 3 OH or CH 3 CO 2 H? d) CH 3 CO 2 H or Cl 3 CCO 2 H? Circle the most acidic H atom in this molecule:
Acids and Bases 2. Lewis Acid = electron pair acceptor, electrophile Base = electron pair donor, nucleophile Lewis acid react with Lewis base form a new covalent bond
Lewis Acids Incomplete octet (e.g. CR 3+, BX 3 ), or Polar bond to H (e.g. HCl), or Carbon with d + due to polar bond (e.g. CH 3 Cl)
Lewis Bases Nonbonded electron pair (anything with O, N, anions)
Lewis Bases If there is more than one possible reaction site (more than one atom with a lone pair), reaction occurs so that the more stable product is formed. Example: Which oxygen is protonated when acetic acid reacts with sulfuric acid?
Molecular Model Kits How to use Make a model for ethane Make a model for butane Make a model for cyclohexane Use 6 white hydrogens and 6 green hydrogens Put 1 green and 1 white hydrogen on each carbon atom The green and white hydrogen atoms should alternate (so as you look at the molecule from the top the H s should alternate greenwhite-green-white-green-white around the ring)
Alkane Three-dimensional Structure Methane: With 2 or more carbons, 3D arrangement can change due to C C bond rotation Conformations Same molecular formula Same atom connectivity Different 3D arrangement due to rotation around single bond Ethane:
Newman Projections Used to better visualize conformations View the C C from the end (look down the C C bond) Represent the C atoms as a dot (front carbon) and circle (back carbon) Show bonds coming out of the circle and dot Example:
Ethane Conformations Staggered vs. eclipsed Staggered is more stable (lower E) due to maximum separation of electron pairs in covalent bonds Eclipsed is less stable (higher E) due to electron repulsions
Dihedral Angle The degree of rotation between C-H bonds on the front and back carbons Torsional strain Accounts for energy difference between eclipsed and staggered Barrier to rotation Caused by electron repulsion Overcome by collisions of molecules
Butane Conformations Look down C2 C3 bond to draw Newman projections Each C has 2 H atoms and 1 CH 3 group Dihedral angle is angle between CH 3 groups There are six conformations of butane: How many staggered conformations? How many eclipsed?
Strain in Butane Conformations Torsional strain Barrier to rotation Example: eclipsed vs. staggered conformations Steric strain Repulsive interaction when atoms are forced close together (occupy the same space) Example: CH 3 -H eclipsed vs. CH 3 -CH 3 eclipsed conformations Example: Anti vs. gauche conformations So, which conformation is lowest in E? Highest in E? What would the plot of energy vs. dihedral angle look like?
Butane Conformations
Interpreting Newman Projections Which of the following Newman projections does NOT represent 2-methylhexane?
Cycloalkane Three-dimensional Structure C atoms in cycloalkanes are sp 3 Bond angles are not always 109.5º Bond angles are dictated by the number of atoms in the ring Angle strain = Forcing angles smaller or larger than 109.5º Cycloalkanes can also have torsional strain (eclipsed H s)
Strain in Cycloalkanes
Cycloalkane Conformations Cycloalkanes adopt more stable conformations to relieve strain Cyclopropane Bent bonds
Cycloalkane Conformations Cyclobutane Puckered conformation Cyclopentane Envelope conformation
Cyclohexane Most stable cycloalkane Most abundant in nature No angle strain (109.5º) No torsional strain (all H s staggered) Conformation = chair Drawing chairs: also see Klein p. 171
Cyclohexane Axial and equatorial hydrogens Axial = parallel to axis through ring Equatorial = perpendicular to axis Each C has one axial H and one equatorial H Look at molecular model
Cyclohexane
Ring Flip Interconversion of two chair conformations Try this with your molecular model If no substituents, these conformations are equal in energy
Monosubstituted Cyclohexanes Two conformations 1. Substituent in axial position 2. Substituent in equatorial position These conformations are not equal in energy Example: methylcyclohexane Steric strain = 1,3-diaxial interactions ( 1,3 refers to distance between groups) Larger groups have more steric strain
Disubstituted Cyclohexanes The most stable conformation has the most substituents in the equatorial position Conformational analysis Look at all chair conformations (cis and trans) and analyze stability Example: 1,4-dimethylcyclohexane
Additional Cyclohexane Conformations Boat No angle strain High torsional strain High steric strain Very unstable Twist-boat Relieves some torsional and steric strain No angle strain Lower E than boat Higher E than chair
Energy Diagram for Cyclohexane Conformations
Conformations of Polycyclic Molecules Fused rings Typically adopt chair conformations Norbornane and derivatives locked in boat conformation
Example Draw the most stable chair conformation for the following molecules: trans-1,2-dimethylcyclohexane trans-1-isopropyl-3-methylcyclohexane
Degree of Unsaturation Unsaturated compounds Have less than (2n+2) H atoms for (n) C atoms Contain elements of unsaturation p bonds Rings Calculating degree of unsaturation Index of Hydrogen Deficiency (IHD) IHD = C - ½ (H + X) + ½ (N) + 1 Ex: C 6 H 14 IHD = 6 - ½(14) + 1 = 0 Alkane Ex: C 6 H 12 IHD = 6 - ½(12) + 1 = 1 1 p bond or 1 ring Ex: C 6 H 10 IHD = 6 - ½(10) + 1 = 2 2 p bonds, 2 rings, or 1 of each C 6 H 14 C 6 H 12 C 6 H 10
Alkene Stability Which alkene is more stable, cis or trans? Cis has steric strain between R groups
Alkene Stability Stability determined by heats of hydrogenation or combustion Heat of hydrogenation = heat of reaction for addition of H 2 (with metal catalyst) to alkene Heat of reaction is proportional to energy of alkene Smaller magnitude DH = more stable alkene
Alkene Stability Heat of combustion= heat of reaction for combustion of alkene to CO 2 and H 2 O Heat of reaction is proportional to energy of alkene Smaller magnitude DH = more stable alkene
Alkene Stability Trends in alkene stability Trans is more stable than cis More substituted C=C is more stable Why? Hyperconjugation Stabilizing effect of adjacent orbital overlap Bond strengths sp 2 -sp 3 bond more stable than sp 3 -sp 3
Example There are four stereoisomers for 2,5-octadiene. Draw all of the stereoisomers and circle the structure that is the most stable.