Wright State University
CORE Scholar
Computer Science and Engineering Faculty
Publications
Computer Science and Engineering
2003
A Gentle Introduction to (or Review of )
Fundamentals of Chemistry and Organic
Chemistry
Dan E. Krane
Wright State University - Main Campus, [email protected]
Michael L. Raymer
Wright State University - Main Campus, [email protected]
Follow this and additional works at: http://corescholar.libraries.wright.edu/cse
Part of the Computer Sciences Commons, and the Engineering Commons
Repository Citation
Krane, D. E., & Raymer, M. L. (2003). A Gentle Introduction to (or Review of) Fundamentals of Chemistry and Organic Chemistry. .
http://corescholar.libraries.wright.edu/cse/385
This Presentation is brought to you for free and open access by Wright State University’s CORE Scholar. It has been accepted for inclusion in Computer
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[email protected].
CS 790 – Bioinformatics
A Gentle Introduction to
(or review of)
Fundamentals of Chemistry
and Organic Chemistry
Square one…
CS790 – Bioinformatics
Fundamentals of Chemistry
• Reading the periodic table
• Neutrons and isotopes
Isotopes of Chlorine
Isotope
35Cl
37Cl
Protons Neutrons
17
18
17
20
Atomic
mass
34.97
36.97
Natural
abundance
76%
24%
6
C
Carbon
12.01
• Electron shells, subshells and orbitals
• Each orbital can hold at most 2 electrons
• In the ground state orbitals are filled from lower to
higher energy
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Electron shells and orbitals
• Quantum numbers
• n = First quantum number = shell
• l = Second quantum number = orbital type
• Golden rule: l < n
Types of Orbitals
Second
quantum
number
Letter
denoting
orbitals
Number
of
orbitals
Maximum
number of
electrons
0
1
2
3
s
p
d
f
1
3
5
7
2
6
10
14
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Know these
two.
3
Subshells and valence
• All orbitals of the same type (same l and n) are
# electrons in
called a subshell
the subshell
• Subshell
Electron
5
Type of
2p
notation:
shell
orbitals
Electron Subshells
1st Quantum 2nd Quantum Notation for
number
number
subshells
1
2
3
4
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0
0,1
0,1,2
0,1,2,3
…
1s
2s,2p
3s,3p,3d
4s,4p,4d,4f
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Electronic configurations
• Since the subshells
are filled from lowest
to highest energy,
we can specify only
the outermost shell.
• Atoms tend to lose or
gain electrons such
that the outermost
subshell is full:
valence
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Covalent Bonds
• For almost all of the
elements that we will
deal with, 8 valence
electrons is an
electronically stable
configuration.
• Covalent bonds are
formed when atoms share
electrons to fill the
valence shell
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Covalent bonds: Lewis diagrams
• How many covalent bonds will an atom form?
• Flourine: Atomic number = 9,
Electron configuration: 1s2,2s2,2p5
F
• Oxygen:
O
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FF
or
F F
Atomic number = 8
Electron configuration: 1s2,2s2,2p4
O O
or
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O O
7
How many covalent bonds?
• Note the common
valences for the
elements most
common in
proteins and
DNA:
•
•
•
•
•
Carbon
Oxygen
Nitrogen
Hydrogen
Sulfur
• Note the similarity
between S and O.
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Ions and ionic bonds
• Formation of ions
• Conflicting goals: neutral charge vs. stable
electronic configuration
• Some atoms have a strong tendency to gain or lose
electrons:
Sodium (Na): Atomic # = 11: 1s2,2s2,2p6,3s1  Na+
 Chlorine (Cl): A# = 17: 1s2,2s2,2p6,3s2 ,3p5  Cl–

• Complete electron transfer, no sharing
q+ q−
• Coulombs law: force = 2
d
• Ionic bond or salt bridge
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Polar Bonds
• In reality, some atoms will attract
shared electrons more strongly.
That is, the shared electrons will
be “off center”.
• The tendency to attract electrons is
called electronegativity.
• There is a continuum between
covalent bonds and ionic bonds.
K I
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K
δ+
I
δ–
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The Hydrogen Bond
• When hydrogen forms a polar bond, the nucleus
is left without any unshared electrons
• It can make a secondary bond with another negative
ion, called a hydrogen bond
• Very common in water:
δ+ δ–
H
O
• Weaker than polar and
δ+
covalent bonds
H
• Donor: covalent/polar bond to H
• Acceptor: ionic attraction to H
O N
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Van der Waals bonds
• Nonspecific – when any two atoms at ~3 to 4 Å
apart
• Å = angstrom units = 10−10 meters = 0.1 nm
• Low energy interaction
• Significantly smaller than
h-bonds or ionic attraction
• Adds up over many atoms
• When two atoms have very
similar shapes, the Van der
Waals contacts can become significant
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Energy of molecular interactions
• 1 calorie = the amount of energy to raise the
temperature of 1g of water from 14.5 to 15.5°C
• Molecules have about 0.6 kcal/mole of energy
from heat/vibration
• Molecular interactions:
• C–C : 83 kcal/mole
• Electrostatic and hydrogen bonds: ~3 – 7 kcal/mole
• Van der Walls interaction: ~1 kcal/mole
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Looking at chemical structures
Propane:
Benzene:
H H H
H
H
C
H C C C H
C
C H
H C
H H H
C
CH3 CH2 CH3
H
C
H
C C C
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A hydrocarbon isomer
• Carbon can make 4 covalent bonds
• There are more carbon-based compounds present on
earth than the total of all compounds lacking carbon
• We could spend an entire course examining the
properties of hydrocarbons: molecules made up
only of carbon and hydrogen.
• Example: Isomers of C4H10
• Butane: CH3 CH2 CH2 CH3
• Isobutane: CH3 CH
CH3
CH3
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Double Bonds
• Double bonds can force a molecule or
functional group to be planar:
• Geometric isomers
• cis = on the same side
• trans = on the opposite side
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Some Common Functional Groups
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Concentration
• 1 mole of a substance = 6.02 × 1023 atoms or
molecules of that substance
• C – atomic weight = 12, one mole = 12 grams
• We express concentration in molarity or
moles/liter.
• Denoted [x].
• Example – If we take 1 mole of sodium sulfate
(142.1g of Na2SO4) and add enough water to make
1 liter of solution: M = [Na2SO4] = 1.0
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Acids and Bases
• Acids give off protons in solution
• HCl  H+ + Cl−
• In water, the H+ ion often binds with water to form
a hydronium ion H3O+
• Strong acids dissociate completely
• Weak acids do not dissociate completely
• pH of a solution
• pH = −log[H+]
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More on pH
• A simple example:
• Suppose we add 0.001 moles of HCl to 1.0 L of H20
• [H+] = 10−3 moles/liter, so pH = 3
•0
7
←acidic
basic→
• Bases accept H+ ions
14
• pOH = −log[OH −]
• pH + pOH = 14
• Water: pH = 7, pOH = 7
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pKa
• For a weak acid, the pKa is a measure of the
tendency of the acid to dissociate (give of an H+
ion)
• Key rule:
• pH = pKa : protonated and unprotonated forms are
at equilibrium
• pH < pKa : more protonated
• pH > pKa : less protonated
• Biological pH varies but is generally close to
neutral (7.0) or slightly acidic
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Properties of Water
• The polarity of water makes it highly cohesive:
• Water solvates & weakens
ionic and hydrogen bonds:
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Hydrophobic Attraction
• Nonpolar (hydrophobic atoms), are driven
together
• Hydrophobic interactions
• Driven by water’s affinity for itself
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