disclaimer reviews do not cover all the material midterm 2 review
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Reviews do not cover all the material
MIDTERM 2 Review
CHEMISTRY
Most chemistry is in the electrons
(the valence electrons)
Atomic orbitals
Atomic orbitals
Representation of the 2p orbitals.
Atomic orbitals
Wolfgang PauliAtomic orbitals
The orbitals filled for elements in various parts
Atomic orbitals
CHEMISTRY
Full shells make the most stable
atoms
intra molecular bonding
intra molecular bonding
The HCL molecule has a dipole moment
intra molecular bonding
The Pauling electronegativity values as updated by A.L. Allred in 1961. (cont’d)
Arbitrarily set F as 4
intra molecular bonding
Skeletal Structure
• Hydrogen atoms are always terminal atoms.• Central atoms are generally those with the
lowest electronegativity.• Carbon atoms are always central atoms.• Generally structures are compact and
symmetrical.
molecular structure
Exceptions to the Octet Rule
• Molecules with an odd number of electrons.
• Molecules in which an atom has less than an octet of electrons.
• Molecules in which an atom has more than an octet of electrons.
molecular structure
Resonance Forms• Lewis structures that differ only in the placement of
electrons are resonance forms. For O3:
• Experimentally, it is found that both bonds are 0.128 nm long.
• The Lewis structure of O3 must show both resonance forms.
O O O · ·
· ·
· · · ·
· · ·· ·· O O O
· ·
· ·
· · · ·
· ·
·· ==
molecular structure
Molecular ShapesAB2
Linear
AB3
Trigonal planar AB4
Tetrahedral
AB5
Trigonal bipyramidal
AB6
Octahedral
AB3EAngular or Bent AB3E
Trigonalpyramidal
AB3E2
Angular or Bent
AB4EIrregular tetrahedral(see saw)
AB3E2
T-shaped
AB2E3
Linear
AB6ESquare pyramidal
AB5E2
Square planar
molecular structure
Dipole Moment
Nonpolar
Polar
....
H H
O
C OO
Bond dipoles
Overall dipole moment = 0
Bond dipoles
Overall dipole moment
The overall dipole moment of a moleculeis the sum of its bond dipoles. In CO2 thebond dipoles are equal in magnitude butexactly opposite each other. The overall dipole moment is zero.
In H2O the bond dipoles are also equal inmagnitude but do not exactly oppose eachother. The molecule has a nonzero overall dipole moment.
221
dqqk
F Coulomb’s law
m = Q r Dipole moment, m
Bond Enthalpies and Bond Lengths
As bond order increases, the bond enthalpy increases and the bond length decreases.
D(C-C) = 348 kJ 0.154 nm
D(C=C) = 614 kJ 0.134 nm
D(CºC) = 839 kJ 0.120 nm
D(C-O) = 358 kJ 0.143 nm
D(C=O) = 799 kJ 0.123 nm
D(CºO) = 1072 kJ 0.113 nm
Hydrogen, H2
Hydrogen fluoride, HF
Fluorine, F2
Molecular orbitals
(a) Lewis structure of the methane molecule (b) the tetrahedral molecular geometry
of the methane molecule.
Molecular orbitals
Hybrid Orbitals
sp sp2 sp3 sp3d sp3d2
Types of Hybrid Orbitals
Shapes: linear triangular tetrahedral trig. bipyram. Octahedral# orbitals: 2 3 4 5 6
Molecular orbitals
The relationship among the number of effective pairs, their spatial arrangement,
and the hybrid orbital set required
Molecular orbitals
(a) Orbitals predicted by the LE model to describe (b) The Lewis structure for carbon dioxide
Molecular orbitals
The combination of hydrogen 1s atomic orbitals to form MOs
Molecular orbitalsenergies
(a) The MO energy-level diagram for the H2 molecule (b) The shapes of the Mos are obtained
by squaring the wave functions for MO1 and MO2.
Molecular orbitalsenergies
The expected MO energy-level diagram for the combustion of the 2P orbitals on two boron atoms.
Molecular orbitalsenergies
The MO energy-level diagrams, bond orders, bond energies, and bond lengths for the
diatomic molecules, B2 through F2.
Molecular orbitalsenergies
16a–26
Intermolecular Forces• The covalent bond holding a molecule together is an
intramolecular force.• The attraction between molecules is an intermolecular
force.• Intermolecular forces are much weaker than
intramolecular forces (e.g. 16 kJ/mol vs. 431 kJ/mol for HCl).
• When a substance melts or boils the intermolecular forces are broken (not the covalent bonds).
• When a substance condenses intermolecular forces are formed.
Intermolecular forces
16a–27
Larger INTERmolecular forces →
• Higher melting point• Higher boiling point• Larger enthalpy of fusion
Intermolecular forces
16a–28
Larger INTERmolecular forces →
• Higher melting point• Higher boiling point• Larger enthalpy of fusion
•Larger viscosity•Higher surface tension•Smaller vapor pressure
Intermolecular forces
Table of Force Energies
Type of Force Energy (kJ/mol) Ionic Bond 300-600Covalent 200-400
Hydrogen Bonding 20-40Ion-Dipole 10-20Dipole-Dipole 1-5Instantaneous Dipole/Induced Dipole 0.05-2
Intermp;ecular forcesIntermolecular forces
Intermolecular Forces
London Dispersion Forces
• London dispersion forces increase as molecular weight increases.• London dispersion forces exist between all molecules.• London dispersion forces depend on the shape of the molecule.• The greater the surface area available for contact, the greater the dispersion forces.• London dispersion forces between spherical molecules are lower than between sausage-
like molecules.
Intermolecular forces
H-Bonding
Occurs when Hydrogen is attached to a highly electronegative atom (O, N, F).
N-H… N- O-H… N- F-H… N-
N-H… O- O-H… O- F-H… O-
N-H… F- O-H… F- F-H… F-
d+ d- Requires Unshared Electron Pairs of Highly Electronegative Elements
Intermolecular forces
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Intermolecular Forces Summary
Intermolecular Intramolecular
Intermp;ecular forcesIntermolecular forces
16a–33
Which forces?London Dipole H-bond ionic
Xe
CH4
CO2
CO
HBr
HF
CH3OH
NaCl
CaCl2
X
X
X
XX
XX
XX
XX
X
X
Intermolecular forces
16a–34
Relative forcesI2 Cl2
H2S H2O
CH3OCH3 CH3CH2OH
CsBr Br2
CO2 CO
SF2 SF6
>LargerLondon
< H-bond
< H-bond
< polar
>polar
>ionic
Intermolecular forces
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Bonding in Solids SOLIDS
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Examples of Three Types of Crystalline Solids
SOLIDS
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crystals
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Figure 16.11: Reflection of X rays of wavelength
n λ = 2 d sin θ
crystals
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1
½
¼
1/8
Atoms in unit cell
crystals
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Cubic Unit Cells of Metals
Simple cubic (SC)Simple cubic (SC)
Body-centered cubic (BCC)
Face-centered cubic (FCC)
1 atom/unit cell
2 atoms/unit cell
4 atoms/unit cell
crystals
Copyright © Houghton Mifflin Company. All rights reserved. 16a–41
crystals
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Ion Count for the Unit Cell: 4 Na+ and 4 Cl- Na4Cl4 = NaCl
Can you see how this formula comes from the unit cell?
Your eyes “see” 14 Cl- ions and 13 Na+ ions in the figure
crystals
Copyright © Houghton Mifflin Company. All rights reserved. 16a–43
3
3
3
3
x
: 8
mass:
8
43:
4
x
8
4 0 7
. 4
Volume R
DensityR
Rfraction
R
fcccrystals
Copyright © Houghton Mifflin Company. All rights reserved. 16a–44
crystals
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Molarity = Moles of solute/Liters of Solution (M)
Molality = Moles of solute/Kg of Solvent (m)
Mole Fraction=Moles solute/total number of moles
Mass %=Mass solute/total mass x 100
Concentration solutions
solutions
Figure 16.55: The phase diagram for water
phase cahnges
Qtotal = q1 + q2 + q3 + q4 + q5phase cahnges
Copyright © Houghton Mifflin Company. All rights reserved. 17a–49
Thermodynamics of Phase Changes
AB
Why does a liquid at A form a solid when the temperature is lowered to B
solutionsphase cahnges
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Gases:Large Entropy
Liquid:Smaller Entropy
Solids:Smallest Entropy
solutionsphase cahnges
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Thermodynamics for Phase Change
∆G = ∆H - T∆S
• liquid→solid• ∆H is negative (stronger intramolecular forces)• ∆S is negative (more order)• -T∆S is positive• As T decreases, -T∆S becomes smaller• ∆G goes to zero when ∆H = T∆S (at T = Tfusion)
• For T less than Tfusion, ∆G is negative, solid is stable.
Negative for spontaneous process Negative for
liquid to solidPositive for liquid to solid
solutionsphase cahnges
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Factors Affecting SolubilityGas – solvent: Pressure Effects
Henry’s Law:
Cg is the solubility of gas, Pg the partial pressure, k = Henry’s law constant.
Carbonated beverages are bottled under > 1 atm. As the bottle is opened, Pg decreases and the solubility of CO2 decreases. Therefore, bubbles of CO2 escape from solution.
gg kPC
solutions
Raoult’s Law
• Raoult’s Law: PA is the vapor pressure of A with solute PA is the vapor pressure of A alone A is the mole fraction of A
PA = XA PAo
PTotal = XA PAo + XB PBo
solutions
Figure 16.44: Behavior of a liquid in a closed container
solutionssolutions
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Colligative properties
• Vapor pressure – Mole fraction
• Freezing point depression – molality• Boiling point elevation – molality• Osmosis - Molarity
solutions
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Figure 16.24: A representation of the energy levels (bands) in a magnesium crystal
semiconductors
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Band structure of Semiconductorssemiconductors
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Silicon Crystal Doped with
(a) Arsenic and (b) Boron
semiconductors
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Figure 16.34: The p-n junction involves the contact of a p-type and an n-type semiconductor.
semiconductors
Semiconductors – key points to remember
• Band structure: Valence band – gap – conduction band
•DOPING: Group V n type, Group III p type
•n-p junctions
•Devices: (LED, laser, transistor, solar cell)
semiconductors
What is a transition metal?
“an element with valance d- or f-electrons”
ie. a d-block or f-block metal
d-block: transition elements
f-block: inner transition elements
3d
4d
5d
6d
l = 2
ml =
-2,-1,0,1,2
4f
5f
l = 3ml =-3,-2,-1,0,1,2,3
transition metalcomplexes
n+/-
What is a coordination complex?
•Central metal ion or atom surrounded by a set of ligands
•The ligand donates two electrons to the d-orbitals around the
metal forming a dative or coordinate bond
metal ion
ligands
charge on complex
X+/-
n
counterion
transition metalcomplexes
transition metalcomplexes
2+-1
Common Coordination Numbers of Transition Metal Complexes
transition metalcomplexestransition metalcomplexes
Classes of isomers transition metalcomplexestransition metalcomplexes
Isomers I and II
transition metalcomplexestransition metalcomplexes
Energy of 3d orbitals
t2g
eg
transition metalcomplexestransition metalcomplexes
Strong/weak fields, d6 Configuration
Paramagnetic – 4 Unpaired Electron Spins
Diamagnetic – No Unpaired Electron Spins
transition metalcomplexestransition metalcomplexes
Correlation of High and Low Spin Complexes With Spectrochemical Series
t2g4eg
2 t2g3eg
3
t2g6 t2g
5eg1
transition metalcomplexestransition metalcomplexes
Name calling
• CH4 methane• C2H6 ethane• C3H8 propane• C4H10 butane• C5H12 pentane• C6H14 hexane• C7H16 heptane• C8H18 octane
Figure 22.3: Structures of (a) propane (b) butane
CH
HH
HCH
HCH
CH
Naming Branched Alkanes
CH3 methyl branch
CH3CH2CH2CHCH2CH3
6 5 4 3 2 1 Count
3-Methylhexane
on third C CH3 six carbon chain group
Cycloalkanes with Side GroupsCH3
CH3
CH3
CH3
CH3
CH3
methylcyclopentane
1,2-dimethylcyclopentane
1,2,4-trimethylcyclohexane
Calling names
• ALKANES
• ALKENES
• ALKYNES
• CYCLO-
• ALKYL-
isomers
• Structural – chain
• Structural - position
• Structural – function
• Stereo - geometrical
• Stereo - optical
butanemethyl propane
2methylhexane3methylhexane
cistrans
Cis and Trans Isomers
Double bond is fixed Cis/trans Isomers are possible
CH3 CH3 CH3
CH = CH CH = CH
cis trans CH3
alkan-OL
alkan-AL
alkan-ONE
Amino Acids
• Building blocks of proteins• Carboxylic acid group• Amino group• Side group R gives unique characteristics
R side chain I
H2N—C —COOH I
H
Most Amino Acids Have
Non-Superimposable Mirror Images
What is the exception?
NH2 COOH1 NH2 COOH2
NH2 C N COOH
O
H21
Amino acids are connected head to tail
Formation of Peptide Bonds by Dehydration
Dehydration-H2O
Juang RH (2004) BCbasics
HIERARCHY OF PROTEIN STRUCTURE
Tertiary
1. 2.
3. 4.
Cytoplasm
Nucleus
DNA
DNA is the genetic material within the nucleus.
Central Dogma
RNA
Protein
Replication
The process of replication creates new copies of DNA.
TranscriptionThe process of transcription
creates an RNA using
DNA information.
TranslationThe process of translation
creates a protein using
RNA information.
Translation• The process of reading the RNA sequence of an mRNA
and creating the amino acid sequence of a protein is called translation.
Transcription
Codon Codon Codon
Translation
DNA
T T C A G T C A G
DNAtemplatestrand
mRNA
A A G U C A G U C MessengerRNA
Protein Lysine Serine ValinePolypeptide(amino acidsequence)