Chemistry 241
Chapter 1: Electronic Structure and Bonding, Acids and Bases
Homework: 68a-f, 69c-i, 71a-e, 73, 74, 76c-f, 77a-d, 78, 79, 82, 83, 84, 87, 89, 96
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IV.
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Introduction
A. What is an organic molecule?
B. # Ionic molecules (125K) vs # organic molecules (6 mil)
C. Vitalism
1. Friedrich Woehler: ammonium cyanate → urea
Electronic Structure
A. s, p, d, f orbitals
B. Von Pauli exclusion principle (max 2 electrons/orbital, opposite spin)
C. Hund’s Rule (empty orbitals fill first, then partially filled ones)
D. Remember fill order using periodic table (easiest) or chart:
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
E. Fill order is NOT the opposite of empty order
Ionic, Polar Covalent, Non Polar Covalent Bonds
A. Electronegativity
1. What is it?
2. Who is closest to F?
a. row is better than column
B. Ionic bonds: electron transfer, then attraction of opposites
1. between metal and non metal
C. Covalent is electron sharing
1. Electronegativity and sharing
2. Sharing equally (true covalent or non polar covalent)
3. Sharing unequally (polar covalent)
Lewis Structures
A. Valence electrons as dots
count up valence electrons
add or subtract for charge
figure out how many electrons needed to fill octets
use steps 1 and 2 to determine how many electrons that you have
subtract electrons you have (4) from electrons needed (3), result is electrons
shared
each shared electron pair equals one bond
arrange all outer atoms around the central atom (usually least electronegative)
put in shared pairs (bonds) – make sure everything has single bond before
going to double bonds or triple bonds
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9. Add lone pairs from leftover electrons to fill octets
10. do not violate octet rules (no more than 4 bonds or lone pairs to any atom
unless 3rd row or lower)
a. why? empty d shells to use
11. double check your electron count
Examples:
H2O
CCl4
CO2
CO32-
N2 O
H2CO
12. works for most things obeying octet rule
a. exceptions: PCl5, BF3, other inorganic cmpds
b. organic exceptions: may have 2+ central atoms (HCOOH, H4C2O2)
Examples:
B. Formal charge
1. assign 1 e-/bond and 2 e-/lone pair
2. # valence e- minus #assigned e- = formal charge
NO3N2O
CO32-
C. Bonds as lines instead of dots in Lewis Structures
1. Kekule structures are Lewis structures w/o lone pair
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D. Line diagrams
V.
VI.
Atomic and Molecular Orbitals
A. Shapes of atomic orbitals
B. Molecular orbitals and “orbital hybridization”
1. sp3, sp2, sp hybridization
2. bond angles, bond lengths, bond strengths (table, p41)
C. Dipoles in bonds, dipoles in molecules
Acids and Bases
A. terms: Arrhenius, Bronsted-Lowry, neutralization, strong and weak,
conjugate acid/base
B. pKa and pH
1. minus the log of
2. Henderson-Hasselbach Equation
3. Predicting whether or not the acid has given up the H+ at a
given pH if you know pKa
C. Substituents affect acid strength
1. electron withdrawing substituents stabilize the conjugate base,
making the acid stronger
2. Repeat: conjugate base (A- with negative charge) is made more
stable, NOT the original acid (HA)
3. The bigger the electronegativity on substituent→the bigger the
pull on the negative charge→ the more stable the conjugate
base is → the stronger the acid is
a. F > Cl > Br > I
4. The closer the electron withdrawing group is to the negative
charge on the conjugate base, the stronger the acid
5. The more electron withdrawers you have, the stronger the acid
6. Not a substituent effect, but the more resonance structures you
can draw, the stronger the acid
a. resonance moves lone pairs and pi bonds, but keeps
single bonds
D. Buffers
1. a weak acid plus conjugate base
2. or weak base plus conjugate acid
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Objectives
Knowledge
Remember the following terms from general chemistry: electronegative, valence
electrons, octet rule, ionic bond, covalent bond, polar and non-polar covalent bonds,
dipole moment, Lewis structures, formal charges, sigma and pi bonds, hybrid (molecular)
orbitals, Bronsted-Lowry acids and bases, Ka, pKa, pH, resonance
Comprehension
Use electronegativities to determine whether bond is ionic, covalent, or polar covalent
Predict acid strength based on stability of conjugate base
Application
Draw Lewis structures for molecules, including lone pairs
Analysis
Use Lewis structures to predict formal charge, bond angle, relative length and strength of
bonds, type of bond (sigma/pi, sp hybridization), whether molecule is polar
Use Henderson Hasselbalch eqn to determine relationship between pH and pKa
Compare the relative acidity of two (or more) molecules looking at substituents (# of
electron w/d atoms attached, distance of them to charge, strength of
electronegativity), and effect of resonance
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