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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

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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

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3

3

3

3

x

: 8

mass:

8

43:

4

x

8

4 0 7

. 4

Volume R

DensityR

Rfraction

R

fcccrystals

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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

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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)