Chapter 02
Lecture and
Animation Outline
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The Nature of Molecules and
the Properties of Water
Chapter 2
4
Nature of Atoms
• Matter has mass and occupies space
• All matter is composed of atoms
• Understanding the structure of atoms is
critical to understanding the nature of
biological molecules
Scanning Tunneling
Microscopy
Oxygen atoms on rhodium
crystal
5
Atomic Structure
• Atoms are composed of
– Protons
Bohr Atomic Model
• Positively charged particles
• Located in the nucleus
– Neutrons
• Neutral particles
• Located in the nucleus
– Electrons
• Negatively charged particles
• Found in orbitals surrounding the nucleus
6
http://d1jqu7g1y74ds1.cloudfront.net/wp-content/uploads/2010/02/c-atom_e1.gif
Atomic number
• Number of protons equals number of
electrons in a neutral atom
• Atomic number = number of protons
• Number of protons determines atom’s
identity
7
Atomic mass
• Mass or weight?
– Mass – refers to amount of substance
– Weight – refers to force gravity exerts
on substance
• Atomic mass = Mass number =
sum of protons and neutrons
• Each proton and neutron has a
mass of approximately 1 Dalton or
atomic mass unit (a.m.u.)
8
http://staff.norman.k12.ok.us/~cyohn/index_files/atom1notes.htm
Properties of Subatomic Particles
Subatomic Particle Summary
subatomic
particle
Mass
(in Daltons or
a.m.u.)
charge
proton
1
+
neutron
1
neutral
electron
almost 0
-
Elements
• Element
– Any substance that cannot be broken down to
any other substance by ordinary chemical
means
– A given element is made up one type of atom
Iron is an element because it
is made up of only iron atoms
10
http://www.periodictable.com/Samples/026.32/s9s.JPG
Elements
• 90 naturally occurring elements
• Only 12 elements are found in
living organisms in substantial
amounts
• Four elements make up 96.3% of
human body weight
– Carbon, hydrogen, oxygen,
nitrogen
– Organic molecules contain primarily
C, H, O & N
C
Carbon (C)
O Oxygen (O)
H Hydrogen (H)
N Nitrogen (N)
Na Sodium (Na)
Cl Chlorine (Cl)
Ca Calcium (Ca)
P Phosphorus (P)
K Potassium (K)
S
Sulfur (S)
Fe Iron (Fe)
Mg Magnesium (Mg)
• Some trace elements are very
important
11
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Periodic Table of the Elements
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
1
3
Li
11
8
Key
H
1
H
4
Be
atomic number
chemical symbol
B
13
Al
Na Mg
20
K
Ca
37
38
Rb Sr
55
56
21
22
Sc Ti
25
23
24
V
Cr Mn
41
42
Fe
14
Si
29
30
31
40
Y
Zr Nb Mo Tc Ru Rh Pd Ag Cd In
57
72
74
44
28
39
73
43
27
75
76
45
46
77
78
47
79
48
80
49
81
Ta
W
Re Os
Ir
Pt Au Hg Tl
105
106
107
109
110
89
108
32
111
112
113
50
9
7
F
N
10
(Lanthanide series)
Ne
a.
60
61
62
63
64
65
66
91
Th Pa
92
U
93
94
95
96
97
98
Hydrogen (H)
Nitrogen (N)
P
S
Cl
Ar
N
33
34
35
36
Na Sodium (Na)
51
52
83
84
Kr
53
54
I
Xe
85
86
Po At Rn
114
116
99
H
18
Pb Bi
67
Oxygen (O)
17
115
117
118
68
69
70
71
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
90
(Actinide series)
59
O
16
Fr Ra Ac Rf Ob Sg Bh Hs Mt Ds Rg Cn Uut Fl Uup Lv Uus Uuo
58
Carbon (C)
15
Sn Sb Te
82
C
He
Co Ni Cu Zn Ga Ge As Se Br
104
88
C
26
Cs Ba La Hf
87
6
5
12
19
2
O
100
101
102
103
Cl Chlorine (Cl)
Ca Calcium (Ca)
P
Phosphorus (P)
K Potassium (K)
S
Sulfur (S)
Fe Iron (Fe)
Mg Magnesium (Mg)
Np Pu Am Cm Bk Cf Es Fm Md No Lr
b.
Know the atomic numbers for:
H (1)
C (6)
N (7)
O (8)
12
Isotopes
• Atoms of a single element that possess
different numbers of neutrons
• Radioactive isotopes are unstable and
emit radiation as the nucleus breaks up
Carbon-12
Carbon-13
Carbon-14
6 Protons
6 Neutrons
6 Electrons
6 Protons
7 Neutrons
6 Electrons
6 Protons
8 Neutrons
6 Electrons
**
** radioactive
Electrons
• Negatively charged particles located in
orbitals
• # electrons = # protons in neutral atoms
14
http://www.universetoday.com/73207/what-are-electrons/
Electrons
• Ions are charged particles –
unbalanced
– Cation (H+, Na+, K+, Ca+2)
–more protons than electrons =
–net positive charge
– Anion (Cl-, OH-)
–fewer protons than electrons =
–net negative charge
15
http://www.chem4kids.com/files/atom_ions.html
Electron arrangement
• Key to the chemical behavior of an atom
lies in the number and arrangement of its
electrons in their orbitals
• Bohr model – electrons in discrete orbits
• Quantum model
 orbital as volume around a nucleus where an
electron is most likely to be found
 More realistic
16
Bohr Atomic Model
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b.
proton
(positive charge)
electron
(negative charge)
neutron
(no charge)
17
Quantum Atomic Model
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Hydrogen
Oxygen
1 Proton
1 Electron
8 Protons
8 Neutrons
8 Electrons
a.
proton
(positive charge)
electron
(negative charge)
neutron
(no charge)
18
Reactivity of Elements
• Valence electrons (e-) – electrons in outermost
energy level
• Octet rule – atoms with 8 e- (2 e- for small
atoms) in outer energy level (shell) are stable
• Inert (nonreactive) elements have full outer
shell
Li
How many valence
electrons in these
atoms?
Ne
Are they inert?
19
1 valence e-
8 valence e-
Electron Orbitals
Electron Shell Diagram
Corresponding Electron Orbital
Energy
Level K
Each orbital holds a
max of 2 electrons
Up to 2 electrons in first shell
One spherical orbital (1s)
Electron Shell Diagram
Corresponding Electron Orbitals
y
z
Energy
level L
x
One spherical orbital (2s)
Electron Shell Diagram
Three dumbbell-shaped orbitals (2kk)
Electron Orbitals
Up to 8 electrons in second shell
(4 x 2 orbitals)
y
z
x
20
Neon
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http://chemistrytextbookcrawl.blogspot.com/2012/07/inorganic-chemistry-shriver-and-atkins_31.html
Reactivity of Elements
• Octet rule –
 Atoms form chemical bonds to obtain
stable octet (8 paired e- in outer most
energy level)
 Valence e- used in chemical bonding
 Lose, gain or share electrons to from
stable chemical bonds
21
http://faculty.clintoncc.suny.edu/faculty/michael.gregory/files/bio%20101/bio%20101%20lectures/chemistry/chemistr.htm
Chemical Bonds
• Molecules are groups of atoms held together in
a stable association
• Compounds are molecules containing more than
one type of element
• Atoms are held together in molecules or
compounds by chemical bonds
O2, a diatomic molecule
H2O, a compound
22
http://en.wikipedia.org/wiki/File:Oxygen_molecule.png
https://commons.wikimedia.org/wiki/File:Water_molecule_3D.svg
Ionic bonds
• Formed by the attraction of oppositely
charged ions
• Gain or loss of electrons forms ions
Na
Sodium atom
a.
Na+
Sodium ion (+)
Cl
Cl–
Chlorine atom
Chloride ion (–)
Cl–
Na+
Cl–
Na+
Cl–
Na+
Cl–
Na+
Cl–
b. NaCl crystal
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23
Covalent bonds
• Form when atoms share 2 or more
valence electrons
• Results in no net charge, satisfies octet
rule, no unpaired electrons
• Strength of covalent bond depends on the
number of shared electrons
24
http://chemistrytextbookcrawl.blogspot.com/2012/07/inorganic-chemistry-shriver-and-atkins_31.html
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covalent bond
Single covalent bond
Hydrogen gas
H
H
O
N
O
O
N
N
O2
Triple covalent bond
Nitrogen gas
N
H
H2
Double covalent bond
oxygen gas
O
H
N2
25
Structural vs. Molecular Formulas
• Molecular formulas
 Show number and
types of elements
• Structural formulas
 Shows arrangement
of atoms and types of
bonds present
26
http://www.cliffsnotes.com/assets/276058.png
Covalent Bonding with H, O, & N
H  1 bond
O  2 bonds
• 2 single or…
• 1 double bond
N  3 bonds
NH3
(d) Ammonia
H
N
H
H
http://kentsimmons.uwinnipeg.ca/cm1504/Image55.gif
• 1 triple, or…
• 1 double + 1 single, or…
• 3 singles
Covalent Bonding with Carbon
C  4 covalent bonds
in some combination
• 4 single
• 1 double, 2
single
• 2 double
• 1 triple, 1 single
http://library.thinkquest.org/C005377/content/images/carbon1.gif
https://commons.wikimedia.org/wiki/File:D-glucose-chain-2D-Fischer.png
Carbon and Organic Molecules
• Carbon can form up to 4
single covalent bonds
(count them)
• Carbon atoms can form
linear chains or rings
• These carbon
“backbones” form basis
of organic molecules
https://commons.wikimedia.org/wiki/File:Glucose_Haworth.png
Glucose (C6H12O6)
29
http://www.3dchem.com/molecules.asp?ID=423
Electronegativity
• Atom’s affinity for electrons
• Differences in electronegativity dictate how
electrons are distributed in covalent bonds
– Nonpolar covalent bonds = equal sharing of
electrons
– Polar covalent bonds = unequal sharing of
electrons
30
http://chemistryfall2009.wikispaces.com/Covalent+bonding
http://www.chem.ubc.ca/courseware/pH/section10/eshift1.jpg
Relatively high
electronegativity
Relatively low
electronegativity
• When covalent bonds form between atoms with
similar electronegativity  nonpolar bonds
• Between atoms with different electronegativity 
polar covalent bonds
31
Electronegativity
• Since polar covalent bonds exhibit unequal sharing
of electrons…
• Electrons spend more time near the atom with the
greatest electronegativity…
• thus, there is a partial negative charge near this atom,…
• and, there is a partial positive charge near the atom with
less electronegativity
δ-
δ+
32
http://www.chem.ubc.ca/courseware/pH/section10/eshift1.jpg
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Chemical reactions
• Involve the formation or breaking of chemical
bonds
 Atoms shift from one molecule to another
without changing number or identity of atoms
 Reactants = original molecules
 Products = molecules resulting from reaction
6H2O + 6CO2
reactants
→
C6H12O6 + 6O2
products
34
http://www.tiem.utk.edu/~gross/bioed/webmodules/enzymes.htm
http://www.gcsescience.com/rc13-catalyst-hydrogen-peroxide.htm
Chemical reactions
• Extent of chemical reaction influenced by
1. Temperature (heat increase rx rate)
2. Concentration of reactants and products
3. Catalysts (presence speeds up rx)
• Many reactions are reversible
Water
• Life is inextricably tied to water
• Single most outstanding chemical property of
water is its ability to form hydrogen bonds
– Weak chemical associations that form between
the partially negative O atoms and the partially
positive H atoms of two water molecules
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a. Solid
b. Liquid
c. Gas
a: © Glen Allison/Getty Images RF; b: © PhotoLink/Getty Images RF; c: © Jeff Vanuga/Corbis
36
Polarity of water
• Within a water molecule, the bonds between
oxygen & hydrogen are highly polar
• Partial electrical charges develop
– Oxygen is partially negative δ+
– Hydrogen is partially positive δ–
δ
δ+
δ+
37
http://academic.brooklyn.cuny.edu/biology/bio4fv/page/image15.gif
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Bohr Model
Ball-and-Stick Model
+
+
–
+
104.5
8p
8n
+
+
–
–
–
+
a.
b.
Partial negative charge at
oxygen end;
Space-Filling Model
+
–
Partial positive charge at
hydrogen end
+
38
c.
Hydrogen bonds
• Cohesion – polarity of water
allows water molecules to be
attracted to one another
• Attraction produces hydrogen
bonds
• Each individual bond is weak
and transitory
• Cumulative effects are
enormous
• Responsible for many of water’s
important physical properties
39
https://www.boundless.com/biology/properties-of-water/the-chemistry-of-water/introduction-to-chemistry-of-water/
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Water molecule
Hydrogen atom
+
Hydrogen bond
a.
–
Oxygen atom
Note: H-bonds
can also form
between any
molecules with
polar covalent
bonds.
Hydrogen atom
Hydrogen bond
+
–
An organic molecule
b.
Oxygen atom
40
• Cohesion – water
molecules stick to
other water molecules
by hydrogen bonding
• Adhesion – water
molecules stick to
other polar molecules
by hydrogen bonding
41
Properties of water
1. Water has a high specific heat
– A large amount of energy is required to
change the temperature of water
2. Water has a high heat of vaporization
– The evaporation of water from a surface
causes cooling of that surface – evaporative
cooling
42
http://www.baltimoresun.com/media/photo/2012-05/70086876.jpg
Properties of water
3. Solid water is less dense than liquid
water
• Water expands when it freezes
• Bodies of water freeze from the top down
http://www.hdwallpapersplus.com/wp-content/uploads/2012/10/iceberg_2.jpg
43
http://sailboatdiaries.com/wp-content/uploads/iceHbonds.gif
44
4. Water is a good solvent
– Water dissolves polar molecules and ions
–
Water molecules
–
–
Na+
–
–
Hydration shells
Na+
Cl–
+
+
+
Cl–
+
+
Salt crystal
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45
45
4. Water is a good solvent
– When H2O dissolves a polar or ionic
compounds like NaCl
– Water is the solvent
–NaCl is the solute
• Both solvent and solute make up a
solution
46
http://bio1151.nicerweb.com/Locked/media/ch03/03_07DissolvingSalt-L.jpg
5. Water organizes nonpolar molecules
– Hydrophilic = “water-loving”
– Hydrophobic = “water-fearing”
– Water causes hydrophobic molecules to
aggregate or assume specific shapes
Hydrophobic molecules
http://www.uic.edu/classes/bios/bios100/lecturesf04am/hydrophobic.jpg
47
6. Water can dissociate to form ions
H 2O

OH–
hydroxide ion
+
H+
hydrogen ion
• H+ ions are also known as protons
48
http://www.kmacgill.com/lecture_notes/lecture_notes_17_files/image003.jpg
Molar Concentrations
A mole is the amount of a
substance in grams equal to
its molecular weight
• One mole of any given
compound contains 6.02 x
1023 molecules of that
compound (Avogadro’s
number)
http://cache.daylife.com/imageserve/07Zx6PxgkK8Ij/340x.jpg
2.4 What Properties of Water Make It So Important in Biology?
Molar Concentrations
A 1 molar (1 M or 1 mol/L) solution is one
mole of a compound dissolved in water to
make one liter
• Ex: A mole of NaCl is the molecular weight
(58.44), in grams. For a 1M solution…
 58.44 grams of NaCl are dissolved in water to
make one liter
The Periodic Table
Molar Concentrations
• The molar concentration of H+ or OH- ions
in a solution is the basis of the pH scale
H2O  H+ + OH(simplified reaction)
(actual reaction)
52
Acids and bases
• For pure water
– [H+] of 10–7 mol/L
– Considered to be neutral
– Neither acidic nor basic
– [H+] = [OH-]
• pH is the negative logarithm of the
hydrogen ion concentration of solution
pH = -log[H+]
pH= –log 10-7 = -(-7) = 7
53
Acids and Bases - pH
• Acid (pH < 7)
– Any substance that dissociates in water to
increase the [H+] (and lower the pH)
– The stronger an acid is, the more hydrogen
ions it produces and the lower its pH
• Base (pH > 7)
– Substance that combines with H+ dissolved in
water, and thus lowers the [H+]
54
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Hydrogen Ion
Concentration [H+]
pH Value
Examples of Solutions
100
10–1
0
1
Hydrochloric acid
10–2
10–3
2
3
Stomach acid, lemon juice
Vinegar, cola, beer
10–4
10–5
4
Tomatoes
5
Black coffee
10–6
6
Urine
Acidic
10–7
10–8
8
Seawater
10–9
9
10–10
10
Baking soda
Great Salt Lake
10–11
10–12
11
Household ammonia
10–13
10–14
Pure water
12
Household bleach
13
14
Basic
Sodium hydroxide
55
Buffers
• Substances that resist changes in pH
• Act by…
– Releasing hydrogen ions when a base is
added
– Absorbing hydrogen ions when acid is added
• Overall effect of keeping [H+] relatively
constant
56
Buffers Minimize pH Changes
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pH
9
8
7
6
5
4
3
2
1
0
Buffering range
0
1X
2X
3X
4X
Amount of base added
5X
57
Buffers
• Most biological buffers consist of a pair of
molecules, one a weak acid and one a
weak base…Ex: carbonic acid
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–
+
Water
(H2O)
+
+
Carbon
dioxide
(CO2)
Carbonic
acid
(H2CO3)
+
Bicarbonate Hydrogen
ion
ion
+
(H+)
(HCO3–)
58