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Welcome to C341!! Chapter 1 & 2:
Review of General Chemistry What will we do today? 1. Review of the syllabus together.
2. Discuss course structure and textbook.
You will use the entire textbook between C341, C342 and C343.
3. Set realistic expectations for work load.
4. Start reviewing general chemistry concepts in chapter 1.
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Pertinent subjects in C117 that lead to success in C341:
1. Periodic table; charges on ions and atoms by group
2. Types of bonding; ionic versus covalent bonds, e.g. NaOCH3, NaH
3. Lewis structures; organic line drawings, understanding ball and stick drawings with wedges and dashes; bond polarity and molecular polarity; functional groups would be nice.
4. Hybridization, shape, and VSEPR.
5. Resonance structures with organic molecules, not just carbonate, sulfate and nitrate.
6. Intermolecular forces, physical properties and solubilities in water versus hydrocarbons.
7. Thermochemistry; bond energies, potential energy diagrams and enthalpy changes.
8. Kinetics; one‐step versus two‐step reactions, reaction mechanisms and their rate laws.
9. Equilibrium; equilibrium constants; Le Châtelier’s principle; product‐favored versus reactant‐favored.
10. Thermodynamics, ΔG and spontaneity; product‐favored versus
reactant‐favored.
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Chapter 1 & 2 outline:
1. Lewis Structures for organic compounds
Line Drawings ‐ drawing organic compounds
Formal charge
Resonance, resonance contributors & bond order
2. Bond Angles & Shapes
3. Hybridization
4. Bond Polarity & Molecular Polarity
Polar & Non‐polar compounds
5. Intermolecular forces
6. Predicting physical properties based on intermolecular forces
7. Predicting solubility based on intermolecular forces
Solubility of molecules
8. Functional Groups – may be new material for some people
You should do all the problems within the chapter and at the end of the chapter. But if you who want to do the minimal amount of work: Chapter 1: 1.34 (use line drawings), 1.39, 1.45, 1.46, 1.48, 1.52, 1.53, 1.56 (draw this as a line drawing), 1.61, 1.63, 1.64 Chapter 2: 2.43‐2.50, 2.53, 2.55‐2.58, 2.61, 2.63‐2.66
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1. Lewis Dot Structures from Condensed Structural Formulae
A Lewis Structure depicts the structure of a compound by its arrangement of atoms with its neighbors, bonds that exist, and the presence of lone pairs.
Molecules share electrons in order to achieve a closed shell or a total of 8 electrons surrounding each atom (octet rule).
Hydrogen is the only atom who is content with only 2 electrons.
Don’t forget that third row atoms or lower can exceed an octet, e.g. S and P.
CH3COCH2NH2
CH3CO2CHCHCH3
CH3CHOHCH2CONHCH2CH3
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Line Drawings Amoxicillin:
Estrogen (female steroid): Zingerone (pungent extract from ginger):
Capsaicin (you and a neighbor come up with the formula for this):
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Constitutional Isomers: Single bonds are axes of rotation to form conformers:
Constitutional isomers of hexanes:
Draw several constitutional isomers for C4H10O:
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Resonance Structures:
o Some molecules cannot be adequately represented by a single Lewis structure.
o Resonance structures are two Lewis structures having the same placement of atoms but a different arrangement of electrons.
o Resonance allows certain electron pairs to be delocalized over several atoms,
and this delocalization adds stability to the molecule. o A molecule with two or more resonance forms is said to be resonance
stabilized.
o Resonance structures are possible if electrons are in conjugation. Rule 1: All resonance structures must have the same number of valence e’s. Rule 2: The octet rule must be obeyed and not exceeded (esp. for C, O, N,& H). Rule 3: Nuclei do not change positions in space between resonance structures. Rule 4: Sigma bonds must not be broken and the skeletal structure rearranged.
Conjugation = A series of alternating single and multiple bonds with overlapping p orbitals.
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Organic Resonance structures (anions, cations & neutral compounds):
Recognizing allylic and vinylic positions:
Indicate all chlorides as either allylic and vinylic: Circle all the allylic lone pairs:
Cl
Cl Cl
Cl
Cl
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Common Mistakes
Localized vs. delocalize electrons. Generally, lone pars adjacent to a C=C double bond are capable of resonance, but not in this case.
Your turn to practice with a neighbor:
O
N
N
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2. Bond angles & shapes VSEPR Theory predicts the MOLECULAR SHAPE & ANGLES FACT: Molecules bond such that all the regions of electron density around a central
atom are as far away from each other as possible.
Determine the shape, not geometry, for the following atoms:
N
O O
OO
cocaine
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3. Hybridization, σ & π‐bonds
Single bonds –
End‐to‐end overlap of orbitals of two ‘s’ orbitals.
Electron density is found along the bond axis in between the atoms.
The first bond between any 2 atoms is always a σ‐bond.
Multiple bonds –
Multiple bonds arise from a combination of a σ ‐bond and π–bonds.
π ‐bonds have their electron density above and below the bond axis.
π ‐bonds arise from the parallel overlap of unhybridized p orbitals.
π ‐bonds are weaker than σ‐bonds.
Two π ‐bonds on the same atom are made in perpendicular planes.
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Determine the hybridization for the following atoms:
Co-enzyme Q-10
H3CO
H3CO
O
O
Hybridization of an amide?
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4. Bond Polarity & Molecular Polarity
Electronegativity = The relative ability of an atom to attract electron density to itself during a covalent bond (expressed on a scale relative to “F” being assigned a number of 4.0)
Expressions of bond polarity
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Which of the following bonds are polar? How do we illustrate that a bond is polar and which direction is the electron density flowing?
H—Br C—O C—H O—H B—H C—I
Bond polarity relates to acidity in chapter 3:
H2B—H H3C—H H2N—H HO—H H—F
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Polar molecules versus non‐polar molecules
A. Use electronegativity values to predict bond dipoles (assign δ+ and δ ‐). B. Use the VSEPR method to predict the molecular shape. C. From the molecular shape, determine whether bond dipoles cancel to give a
non‐polar molecule or combine to produce a resultant dipole moment for the molecule.
Start to learn these organic solvents below (they will be tested on quiz 1/exam 1). Which are polar and non‐polar?
O
acetone
O
H N
DMF(dimethyl
formamide)
O
OHacetic acid
CH3
toluenehexanediethyl ether
S
O
DMSOdimethly sulfoxide
H3C C N
acetonitrilemethanol
CH3OH
O
O
THF(tetrahydrofuran)
methylene chlorideCH2Cl2
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Where do you draw the line between polar vs. non‐polar?
Dielectric constant = the property of a dielectric that determines the electrostatic energy per unit volume for unit potential gradient. WHAT does that mean?
What is considered a polar solvent based on dielectric constant?
Name Structure bp, oC dipole momentdielectric constant
water H‐OH 100 1.85 80
formic acid 100 1.41 58
diemthyl sulfoxide (DMSO) 189 3.96 47.2
N,N‐dimethylformamide (DMF)
153 3.82 38.3
acetonitrile 81 3.92 36.6
methanol CH3‐OH 68 1.70 33
ethanol CH3CH2‐OH 78 1.69 24.3
acetone 56 2.88 20.7
1‐propanol CH3CH2CH2‐OH 97 1.68 20.1
methyl ethyl ketone 80 2.78 18.5
1‐butanol CH3CH2CH2CH2‐OH 118 1.66 17.8
methylene chloride (DCM) CH2Cl2 40 1.60 9.08
tetrahydrofuran (THF) 66 1.63 7.52
acetic acid 118 1.74 6.15
ethyl acetate 78 1.78 6.02
diethyl ether CH3CH2OCH2CH3 35 1.15 4.34
benzene 80 0 2.28
carbon tetrachloride CCl4 76 0 2.24
hexane CH3(CH2)4 CH3 69 ‐‐‐‐ 2.02
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5. Intermolecular Interactions
Intramolecular forces (covalent & ionic forces) are strong.
Intermolecular forces are comparatively weak. But these intermolecular forces add up to make an “overall” significant contribution to the overall physical property of a substance.
Review the types of intermolecular forces: Relative strengths
o H‐bonding (HBA & HBD) 10‐40 kJ/mol o Dipole‐dipole forces (DD) 5‐25 kJ/mol o London dispersion forces (LDF) 0.5‐40 kJ/mol o Induced dipole 2‐10 kJ/mol
Compare to: Relative strengths
o Ionic 400‐4000 kJ/mol o Covalent 150‐1100 kJ/mol o Metallic 75‐1000 kJ/mol
Why is understanding intermolecular forces (IMF) important?
Solubility in solvents (hydrophobic vs. hydrophilic)
Physical properties like boiling and melting points (e.g. the stronger the
intermolecular forces the higher the boiling point)
Three dimensional structures (consider the tertiary structure of a protein)
Reactivity between species
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Non‐polar molecules tend to experience predominately:
London Dispersion Forces (a.k.a. van der Waals forces) WEAK, attractive forces between molecules with TEMPORARY DIPOLE. What dictates extent of LDF?
What is polarizability? Which are more polarizable: cations or anions? Compare the isomers of hexane – they have the same number of atoms and electrons. Which has higher LDF and why?
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Polar organic molecules tend to experience:
Dipole‐dipole forces: attractive forces between molecules with a PERMANENT DIPOLE. Demonstrate how acetone exhibits a DD interaction (you must understand polar bonds and dipoles first).
H‐bonding: an especially strong dipole‐dipole attraction forces between molecules that have X—H bonds (X = O, N, F only). What are HBA and HBD? Demonstrate an H‐bonding interaction:
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6. Predicting physical property trends in organic chemistry
Work with your neighbors to put the following compounds in order of boiling point.
Also put in an example of the methyl butyl amine.
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7. Polarity Effects on Solubility (“Like dissolves like”)
Which would you predict to be soluble in water?
Which would you predict to be NOT water soluble?
Which are protic?
Which are aprotic?
O
acetone
O
H N
DMF(dimethyl
formamide)
O
OHacetic acid
CH3
toluenehexanediethyl ether
S
O
DMSOdimethlysulfoxide
H3C C N
acetonitrilemethanol
CH3OH
O
O
THF(tetrahydrofuran)
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What type of solvent would you use to dissolve cholesterol? (Ask me now about the 5‐carbon rule…)
Micelles:
HO
H
H
H
cholesterol
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Predict the solubility of the following molecules:
NHO
HO OH
NH2
ON
O
Vitamin B6 Novocain (procaine)
OO OH
OH
HO OH
O
O
5
Vitamin C Vitamin K
O
HO
HO Vitamin E Vitamin D
O N
Cl
N
H3CHN Cl
Cl
Valium (diazepam) Zoloft – antidepressant
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8. Functional Groups:
Organic Chemistry is the study of carbon and the compounds that contain carbon (also contain H, N, O & S)
Organic compounds are organized by classes called functional groups.
A functional group is an atom or a group of atoms with characteristic chemical and physical properties. It is the reactive part of the molecule.
The different functional groups refer to compounds that have similar properties with similar chemical bonds.
Functional groups differ based on their component heteroatoms or pi‐bonds.
Heteroatoms have lone pairs and create electron‐deficient sites on carbon.
Many molecules have several functional groups in one molecule!
Capsaicin:
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Hydrocarbons
Hydrocarbons contain carbons and hydrogens; generic formula of R—H (ask me what the R means).
Functional group
Characteristic Examples
Alkanes Contains all C—C single bonds
Alkenes Contains at least one C=C double bond
Alkynes Contains at least one C≡C triple bond
Arenes Contains at least one benzene ring
CH3
NO2
NO2O2N
CH3
Toluene
TNT = trinitrotoluene
Naphthalene = moth ballsOH
NO2
NO2O2N
OH
Phenol
Picric acid = insecticide
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FG Characteristic Examples
Alcohol R—OH
Primary (1o) Secondary (2 o) Tertiary (3 o)
Thiol R—SH
Primary (1o) Secondary (2 o) Tertiary (3 o)
Ether R—O—R
Thioether R—S—R
SSAmines R—NH2
Primary (1o) Secondary (2 o) Tertiary (3 o)
Nitriles R—C≡N
Alkyl Halides
R—X, X = halogen
Primary (1o) Secondary (2 o) Tertiary (3 o)
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Carbonyl containing compounds: C=O is called a “carbonyl group”
RC
O
H RC
O
R Aldehyde Ketone
RC
O
OH RC
O
OR RC
O
NH2
Acid chloride Carboxylic acid Ester Amide
Identify correct functional groups in the following molecules.
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Working with a neighbor/after class, determine the following: Functional groups? Soluble (S) or insoluble (N) in water?
IMF? LF (London Forces), DD (dipole‐dipole), HBA (hydrogen bond
acceptor), and HBD (hydrogen bond donor) as appropriate.
CH3
OHO
O
OH
S or N S or N S or N S or N
LF LF LF LF
DD DD DD DD
HBA HBA HBA HBA
HBD HBD HBD HBD
NH
N
O
O
O
S or N S or N S or N S or N
LF LF LF LF
DD DD DD DD
HBA HBA HBA HBA
HBD HBD HBD HBD