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CHEMISTRY 162A, SPRING 2011
SYLLABUS, POLICIES AND PROCEDURES
LECTURES: M, W, F 8:30 AM – 9:20 PM, KNE 120
LABS: By sections, Bagley 290291 (the new labs!!)
Web Address: http://depts.washington.edu/chem/courses/
Instructor: Dr. Andrea Carroll
Email: [email protected]
Office Hours: M 10-30-12:20
W 1:30-2:20
Quiz Sections and Labs scheduled back to back…new expectations and planning
See course syllabus for detailed information about times, locations, and policies.
MIDTERM EXAMS
There are 2 midterm exams during the quarter:
Midterm 1: Friday, April 22nd, 8:30 – 9:20 AM, KNE 120
Midterm 2: Wednesday, May 18 , 8:30 – 9:20 AM, KNE 120
FINAL EXAM
The final will be given on Tuesday, June 7, 8:30 – 10:20 AM, KNE 120
Chemistry 142 (text chapters 18)
 Properties of matter (examples: density, phases,
elemental composition, reactions)
 Elements and compounds  atoms and molecules
 Periodic properties, compound formulas, polyatomic ions
 Counting atoms/molecules using moles (n):
n = m/MM
(solids/liquids/gases)
n = PV/RT
(gases)
n = VM
(solutions)
 Reactions: dissociation/precipitation (solubility),
acid-base (H+ transfer), redox (e– transfer)
 Gases and atmospheric chemistry
 Equilibrium constants and LeChâtelier’s Principle
 Quantitative pH, solubility calculations (using Ka, Kb, Ksp)
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Chemistry 152 (text chapters 913)
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Chemistry 162 (text chapters 1419 and 21)

Energy flow in reactions: ΔE = q + w = ΔH  PΔV

Molecular orbitals and orbital hybridization

Spontaneity, entropy, and free energy:

Reaction rates, activation energy (k = AeEa/RT), catalysis

Properties of solids, liquids, and solutions: relation to
bond type

Specific chemistry of elements by periodic column:
ΔGo

=
ΔHo

TΔSo
Electrochemistry:
= RT ln Keq
ΔGo
nFEo
=
E = Eo  (RT/nF)ln Q

Light and electrons: atomic energy levels and orbitals
(electron spatial distributions). ΔEelectron = hphoton

Electronic configurations and periodic trends

Electronegativity; ionic and covalent bonds; Lewis
structures and molecular shapes
Main-group (representative) elements (s and p blocks)
Transition metals (d block)

Introduction to organic chemistry and biochemistry
Descriptive chemistry!!
Understanding the Periodic Table and what it can tell us!!
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3/26/2011
Chemistry 162
Atoms, Electrons, and Bonding
Chapter 12:
Atomic orbitals, Properties of electrons, Wave functions,
Electronic configurations, Aufbau principle, etc.
Chemical Principles, 6th Ed.
By Steven Zumdahl
Chapter 13:
General Concepts of Bonding in Molecules
Chapter #14 Covalent Bonding: Orbitals
Chapter #15 Chemical Kinetics
- Types of bonds: ionic, covalent, etc.
Chapter #16 Liquids and Solids
- Bond energies, lengths, polarities, etc.
Chapter #17 Properties of Solutions
Localized Electron Model
Chapter #18 The Representative Elements
- Lewis dot structures
Chapter #19 Transition Metals and Coordination Chemistry
- Resonance structures
Chapter #21 Organic and Biochemical Molecules
- The octet rule
- VSEPR model
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152 Review IONIC AND COVALENT BONDS
152 Review IONIC AND COVALENT BONDS
When 2 atoms bond, where the valence electrons end up:
 determines the properties of the compound;
 depends on how strongly each atom attracts electrons.

Polar covalent bond (example HCl). Atoms unequally
share valence e–. Bond energy is mainly from sharing,
somewhat from +/– attraction. Intermolecular forces are
between purely covalent and ionic, so compounds may
be liquids or solids, as well as gases.

Ionic bond (example KCl). One atom mostly transfers
valence e– to the other. Most bond energy is from +/– ion
attraction: same in all directions, so ion pairs always
coalesce to a solid of alternating + and – ions.
The 3 possibilities:
Purely covalent
Polar covalent
Ionic

Fig 13.12
Purely covalent bond (examples H2, Cl2).
Atoms equally share valence e–. Bond
energy is from simultaneous attraction of e–
to both nuclei. Intermolecular forces are
weak, so compounds are usually gases at
room temperature.
None of the 3 types is intrinsically the strongest.
Example
KCl, HCl, and H2 bond energies are
433, 432, and 436 kJ/mol, respectively.
No bonds are purely ionic. A few are purely covalent.
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152 Review ELECTRONEGATIVITY AND BOND TYPE
152 Review ELECTRONEGATIVITY AND BOND TYPE
Electronegativity, Bond Polarity, and Atomic Radius
Electronegativity of an element is the relative ability of an atom
to attract electrons in a bond.
 If ENB > ENA, AB bond polarity is A+B.
Greater ENBA  more ionic AB bond.
Fig 13.3
EN increases across a row and up a column.
Opposite of ratomic
Exception: ENH = ENP
 Of the non-radioactive elements that form compounds, Cs
has lowest and F highest EN  CsF is most ionic binary
compound. EN for CsF(g) = 3.3 (Cs = 0.7, F = 4.0)
 ENH = ENP, ENC = ENS  nonpolar bonds in PH3 and CS2 .
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152 Review ELECTRONEGATIVITY AND BOND TYPE
152 Review ELECTRONEGATIVITY AND BOND TYPE
Periodic Variation of Bond Type
Most compounds: metal-nonmetal or nonmetal-nonmetal
combinations. EN values mean that:
 s and f metals + nonmetals  mainly ionic compounds
(metal is +, nonmetal ). Examples: Na2O, NdF3.
Exception: Be is covalent except for BeF2 .
 d and p elements + nonmetals  mainly covalent
compounds (element further right and higher in table is ;
metals are + in all of their compounds).
Examples: Fe2O3, SF6.
In formulas, the more positive atom is usually written first
(exceptions: CH4, CH3CO2Na, NH3).
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Special Cases of Ion Formation
 Polyatomic Ions: covalent molecules stable with a
charge  form ionic compounds. Example: NH4NO3(s)
is like KCl(s): alternating NH4+ and NO3 ions.
 Covalent acids and bases may react with each other to
form dissolved ions (text page 234).
Example: HCl(g) + H2O(l)  H3O+(aq) + Cl (aq)
chemical reaction!
Different from:
H2O(l)
KCl(s)  K+(aq) + Cl (aq) …just a phase change!
 Redox reactions may produce ions from neutral, covalent
elements and compounds. Example:
2 K(s) + 2 HCl(aq)  H2(g) + 2 K+(aq) + 2 Cl (aq)
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HCl: oxidant, not acid!
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142 Review POLYATOMIC (COVALENT) IONS
142 Review POLYATOMIC (COVALENT) IONS
Some covalent molecules are stable with a net charge.
Most common examples:
 Know names, formulas, and charges of these ions, which
form ionic compounds with oppositely-charged ions.
Example: if you see “ammonium carbonate” on an exam,
you’re expected to know that it is (NH4)2CO3.
Name
(“… Ion”)
Formula
Example
Related
Compound Acid
Ammonium
Hydronium
Metal Aqua
Hydroxide
Acetate
Carbonate
Nitrate
Phosphate
Sulfate
Perchlorate
NH4+
H3O +
M(H2O)62,3+
OH 
CH3CO2
CO32
NO3
PO43
SO42
ClO4
NH4Cl
HCl(aq)
Cr(H2O)6Cl3
KOH
CH3CO2K
K2CO3
KNO3
K3PO4
K2SO4
KClO4
Name
Related
(“… Acid”) Oxide
 Negative ions covalently bond H+, always to O if available,
to form lower-charged ions and molecular oxyacids.
Example: HPO42, H2PO4, H3PO4
Know oxyacid names and formulas.
Hydrochloric
H2O
CH3CO2H
H2CO3
HNO3
H3PO4
H2SO4
HClO4
Acetic
Carbonic
Nitric
Phosphoric
Sulfuric
Perchloric
CO2
N2O5
P4O10
SO3
Cl2O7
 The related oxides form the oxyacids by reaction with H2O.
The central element's oxidation number does not change.
Example: SO3(l) + H2O(l)  H2SO4(l)
Practice writing and balancing these reactions for all the
oxides listed on the previous slide.
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142 Review POLYATOMIC (COVALENT) IONS
 The aqua ion is the form of most +2 and +3 metals in
aqueous solution (text pages 266-268).
Example: “Cr3+(aq)” is actually Cr(H2O)63+(aq)
Exceptions to M(H2O)62,3+: Be(H2O)42+; Mf(H2O)8, 93+ where
Mf = f-block metal.
 Solubility rules for common ionic compounds in water
include polyatomic ions.
Soluble: Column I, NH4+, acetates, nitrates, perchlorates,
most halides, some sulfates.
Insoluble: Sulfides, hydroxides, carbonates, phosphates
(except Column I, NH4+).
Exceptions: Soluble: Ba(OH)2
Insoluble: Ag+, Pb+2 halides and
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MSO4 where M = Ca, Sr, Ba, Pb
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152 Review MOLECULAR SHAPES
Unlike ionic bonds, covalent bonds are specific and
directional: each pair of atoms forms a bond of different
energy and e– distribution, and the bonds and unshared
electrons around each atom have a particular geometry.
The Valence Shell Electron Pair Repulsion (VSEPR) model
accurately predicts shapes of s/p-block covalent molecules. It
simply assumes that all e groups around each atom get as
far apart as possible. “Group”  all e in one region of
space, whether shared (single, double, or triple bonds) or
unshared (pairs or single unpaired e). Zumdahl calls groups
“effective pairs.”
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3/26/2011
152 Review MOLECULAR SHAPES
152 Review MOLECULAR SHAPES
Electron-group geometries around a central atom that
maximize group separations are:

Molecular shape (arrangement of the nuclei) depends on
(a) the total # of e groups; (b) how many are bonds.
Example: CH4, NH3, and H2O all have the same e-group
geometry around the central atom (tetrahedral), but 4, 3,
and 2 bonds, respectively – making their shapes
tetrahedral, trigonal pyramidal, and bent, respectively. Thus
each e-group geometry gives rise to several different
molecular shapes.

Bond angles in more complex molecules deviate from
those in AX2,3,4,5,6. Unshared pairs and multiple bonds
take up more room than single bonds.
# of e
groups
2
3
4
5
6
Geometry (and group
angular separation)
Examples
Line (180°)
Triangle (120°)
Tetrahedron (109.5°)
Trigonal Bipyramid
(2 axial 90°,
3 equatorial 120°)
Octahedron (90°)
CO2, HCN
H2CO, PbCl2, SO2
CH4, NH3, H2O
PF5, SF4, ClF3,
XeF2
SF6, IF5, XeF4
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152 Review MOLECULAR SHAPES
Examples:
 in H2O, H-O-H angle = 104.5°, not tetrahedral 109.5°
 in H2C=O, H-C-H angle = 116° and H-C-O angle = 122°,
not trigonal 120°;
 unshared pairs and multiple bonds prefer equatorial to
axial positions in a trigonal bipyramid and opposite
rather than adjacent positions in an octahedron.
So, to determine the shape of a molecule:
1. Draw the correct Lewis structure.
2. Figure out the geometry of all the e– groups around the
central atom.
3. See what shape the nuclei alone form.
4. Predict the direction of any deviation of bond angles from
the ideal angles.
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