BIOCHEMISTRY
The Chemical Context of Life
Matter consists of chemical elements in pure
form and in combinations called compounds.
• Element: a substance that cannot be broken down
into other substances by chemical reactions.
• There are 92 naturally occurring elements
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Matter consists of chemical elements in pure
form and in combinations called compounds.
• Compound: consists of 2 or more different elements
combined in a fixed ratio.
• A compound has characteristics different from its
element.
• Na (soft metal, explodes in water) + Cl (poisonous gas)  NaCl (a seasoning we
sprinkle on food without fear!)
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The Elements of Life
C. HOPKINS CaFe
• About 20–25% of the 92 elements are
essential to life
• Carbon, hydrogen, oxygen, and nitrogen
make up 96% of living matter
• Most of the remaining 4% consists of
calcium, phosphorus, potassium, and sulfur
• Trace elements are those required by an
organism in minute quantities
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An element’s properties depend on the
structure of its atoms
• An atom is the smallest unit of matter that still
retains the properties of an element.
• Atoms are composed of subatomic particles
• Relevant subatomic particles include
–
–
–
–
–
Neutrons (no electrical charge)
Protons (positive charge)
Electrons (negative charge)
p+ and n0 reside in a very dense nucleus
e− reside in the electron cloud
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One of these diagrams is fraught with danger!
Cloud of negative
charge (2 electrons)
Electrons
Nucleus
(a)
(b)
More annoying Atomic Structure Vocabulary
Atomic Number and Atomic Mass
• Atoms of the various elements differ in number of
subatomic particles
• An element’s atomic number is the number of protons
in its nucleus
• An element’s mass number is the sum of protons +
neutrons in the nucleus
• Atomic mass, the atom’s total mass, can be
approximated by the mass number but is actually
represented by an AVERAGE molecular mass based on
the abundance of various isotopes.
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Speaking of Isotopes…
• All atoms of an element have the same number of
protons but may differ in number of neutrons
• Isotopes are two atoms of an element that differ in
number of neutrons
• Radioactive isotopes decay spontaneously, giving off
particles and energy
• Some applications of radioactive isotopes in
biological research are
– Dating fossils (C-14)
– Tracing atoms through metabolic processes (I-131)
– Diagnosing medical disorders
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On to the electron cloud…
• The Energy Levels of Electrons
• Energy is the capacity to cause change, perhaps by
doing work.
• Potential energy is the energy that matter has because
of its location or structure, there are many kinds…not just
gravitational PE!
• The electrons of an atom differ in their amounts of
potential energy
• An electron’s state of potential energy is called its energy
level, or electron shell*
* “Shell” is fraught with misconception—but biologists often use this description.
“Energy level” is a much better phrase since the region is not “hard” like a shell.
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Electrons farther from the nucleus have more
potential energy
• A ball bouncing down a
flight of stairs provides
an analogy for energy
levels of electrons,
because the ball can
come to rest only on
each step, not between
steps.
• It’s a quantized event!
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Energy must be absorbed in order to
promote an electron to a higher E-level
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What’s the big deal?
Electron Distribution and Chemical Properties
• The chemical behavior of an atom is determined by
the distribution of electrons in electron energy levels
and sublevels
• The periodic table of the elements shows the
electron distribution for each element—think of it as
a giant BINGO card
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This diagram focuses on the “shells”
which can be misleading!
Hydrogen
1H
Mass number
First
shell
2
He
4.00
Atomic number
Helium
2He
Element symbol
Electron
distribution
diagram
Lithium
3Li
Beryllium
4Be
Boron
5B
Carbon
6C
Nitrogen
7N
Oxygen
8O
Fluorine
9F
Neon
10Ne
Silicon
14Si
Phosphorus
15P
Sulfur
16S
Chlorine
17Cl
Argon
18Ar
Second
shell
Sodium Magnesium Aluminum
11Na
12Mg
13Al
Third
shell
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Stuff you probably already know!
• Valence electrons are those in the outermost energy
level (or sublevel) , or valence sublevel
• The chemical behavior of an atom is mostly
determined by the valence electrons
• Elements with a full valence shell are chemically inert
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Out with the old model…
First shell
Neon, with two filled
Shells (10 electrons)
Second shell
(a) Electron distribution diagram
…in with a better one—electron orbitals
• An orbital is the three-dimensional space where an
electron is found 90+% of the time
• Each electron energy level consists of a specific
number of orbitals
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First E-level
Second E-level
y
x
1s orbital
2s orbital
z
Three 2p orbitals
(b) Separate electron orbitals
Heard of s, p, d and f?
1s, 2s, and
2p orbitals
(c) Superimposed electron orbitals
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The formation and function of molecules depend
on chemical bonding between atoms
• Atoms with incomplete valence shells can share or
transfer valence electrons with certain other atoms
• These interactions usually result in atoms staying close
together, held by attractions called chemical bonds
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Covalent Bonds
• A covalent bond is the sharing of a pair of valence
electrons by two atoms
• In a covalent bond, the shared electrons count as part
of each atom’s valence shell
• A covalent bond is formed between shared pairs of
electrons:
 1 pair—a single bond
 2 pairs—a double bond
 3 pairs—a triple bond
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Hydrogen atoms (2 H)
Hydrogen atoms (2 H)
Hydrogen atoms (2 H)
Potential Energy Curve
for Formation of H2
Hydrogen molecule (H2)
Covalent Bonds:
To share or not to share, that is the question!
• The notation used to
represent atoms and
bonding is called a
structural formula
– For example, H—H
• This can be abbreviated
further with a
molecular formula
– For example, H2
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Name and
Molecular
Formula
(a) Hydrogen (H2)
(b) Oxygen (O2)
(c) Water (H2O)
(d) Methane (CH4)
Electron
Distribution
Diagram
Lewis Dot
Structure and
Structural
Formula
SpaceFilling
Model
Electronegativity
• Atoms in a molecule
attract electrons to varying
degrees
• Electronegativity is an
atom’s attraction for the
electrons in a covalent
bond
• The more electronegative
an atom, the more
strongly it pulls shared
electrons toward itself
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Not all sharing is created EQUAL!
• In a nonpolar covalent bond,
the atoms share the electron
equally
• In a polar covalent bond, one
atom is more electronegative,
and the atoms do not share
the electron equally
–
O
+
H
H
H2O
• Unequal sharing of electrons
causes a partial positive or
negative charge for each atom
or molecule
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+
Not all sharing is created EQUAL!
• In a nonpolar covalent bond,
the atoms share the electron
equally
• In a polar covalent bond, one
atom is more electronegative,
and the atoms do not share
the electrons equally
Net dipole moment
–
O
+
H
H
H2O
• Unequal sharing of electrons
causes a partial positive or
negative charge for each atom
or molecule
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+
Not all molecules with polar bonds are polar!
• In a nonpolar covalent bond,
the atoms share the electron
equally
• In a polar covalent bond, one
atom is more electronegative,
and the atoms do not share
the electrons equally
• Unequal sharing of electrons
causes a partial positive or
negative charge for each atom
or molecule
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Ionic Bonds: A case study involving greed!
• Atoms sometimes strip
electrons from their
bonding partners
• An example is the transfer
of an electron from
sodium to chlorine
• After the transfer of an
electron, both atoms have
charges
• A charged atom (or
molecule) is called an ion
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Ionic Bonds: A case study involving greed!
• A cation is a positively charged ion
• An anion is a negatively charged ion
• An ionic bond is an attraction between an anion and
a cation
• Compounds formed by ionic bonds are called ionic
compounds, or salts
• Salts, such as sodium chloride (table salt), are often
found in nature as crystals
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Here’s where the trouble starts…
• There is a rift between biology and chemistry text books.
• The biology book often speaks of “weak bonds” when they really
mean intermolecular forces (forces of electrostatic attraction
“between molecules”) which are not at all the same as sharing a
pair of electrons within a molecule
• IMFs are intermolecular whereas chemical bonds are
intramolecular
• Inter—means between molecules (think interstate highway, one
between states, connecting states)
• Intra—means within the molecule (actual chemical bonds)
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Here’s where the trouble continues…
• Furthermore, the biology books often speak of van der
Waals interactions which “lumps” all of the different
IMFs together
• Older chemistry books spoke of van der Waals but
specifically meant London Dispersion forces.
More on that coming up…
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There are many types of IMFs
 London dispersion forces (LDFs)—every molecule has
these since every molecule has moving valence
electrons. Essentially the electrons are in constant
motion and not always evenly distributed (think
traffic jam).
If the electrons pile up on one portion of the three
dimensional molecule, then we get a temporary
concentration of negating charge, creating a
temporary negative pole, if you will, since the
electrons are not dispersed evenly we now refer to
the molecule as a temporary dipole.
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LDFs increase with increasing numbers of electrons
on a given molecule
Temporary Dipole
has formed
The larger the molecule, the more likely this will happen
since the valence electrons are farther from the mother
nucleus, thus less tightly held! What wasn’t polar at all is
now “sort of” polar, thus +/- attractions now exist.
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“Peer pressure”
The dotted lines represent attractive forces. This is also happening in all THREE DIMENSIONS!
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What’s the big deal?
• Molecules that were not originally
attracted to one another, now find each
other quite attractive, thus more energy
is required to separate them!
• In other words, molecules become
“sticky” or adhere to one another.
• Collectively, such interactions can be
strong, as between the molecules of a
gecko’s toe hairs and the surface of a
wall. He’s not really defying gravity!
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Induced Dipole- Induced Dipole IMFs
a.k.a. LDFs, or Dispersion Forces
• Induced-Dipole, Induced-Dipole is
just another name for what we
just described as LDFs. (Don’t you
just love grown-ups that can’t
agree on terminology??)
• A spontaneous traffic jam of
electrons created a slight (-)
charge on the end of the
molecule with the most electrons
leaving the other end slightly (+)
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Dipole - Induced Dipole IMFs
• In Dipole-Induced Dipole, there is a
permanent dipole
(electronegativity difference is
enough to make a permanent (+)
and (-) end of the molecule) that
induces a nonpolar molecule to
become a dipole.
• Now the two are more attracted to
each other than they were before
the induction occurred.
• Ever induced behavior in another
human? Got siblings??
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Dipole – Dipole IMFs
• In Dipole-induced dipole,
there is a permanent dipole
(electronegativity difference is
enough to make a permanent
(+) and (-) end of the molecule
that induces a nonpolar
molecule to become a dipole.
• Now the two are more
attracted to each other than
they were before the
induction occurred.
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Hydrogen Bonding:
A Special Case of Dipole-Dipole IMFs
• A hydrogen bond is not the
same as a bonded
hydrogen!
• It’s a special case of DipoleDipole IMFs
• A bonded hydrogen is
within a water molecule
• A hydrogen bond is
between molecules!
Bonded
Hydrogens
(actual chemical bonds
consisting of a shared
pair of electrons)
Hydrogen
bonds (IMFs)
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Hydrogen Bonding:
A Special Case of Dipole-Dipole IMFs
• Hydrogen bonding occurs when
the H of one molecule is
attracted to a highly
electronegative molecule on an
adjacent molecule.
• THE CATCH: For an H to be a
candidate for hydrogen bonding,
it must itself be attached to a
highly electronegative element
such as F, O or N (think
“phone”…call me!)
Bonded
Hydrogens
(actual chemical bonds
consisting of a shared
pair of electrons)
Hydrogen
“bonds” (IMFs)
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Hydrogen Bonding:
A Special Case of Dipole-Dipole IMFs
• Hydrogen bonding is an IMF
that makes molecules more
attracted to each other, thus
more tightly held to each
other, thus more energy is
required to separate them!
• Higher MP, BP, heat of
vaporization, etc. occur as a
result.
• Also, enhanced solubility of
substances such as ammonia.
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Time to think in 3-D!
• A molecule’s shape is
very important to its
function
• A molecule’s shape is
determined by the
positions of its atoms’
valence orbitals
• In a covalent bond, the s
and p orbitals may
hybridize, creating
specific molecular shapes
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Why is Linus Pauling famous??
• He developed the electronegativity scale.
• In the 1930’s Linus Pauling introduced the concept of
hybridization to explain chemical bond formation.
Hybridization is the mixing of atomic orbitals in an atom to
generate a set of new atomic orbitals called hybrid orbitals.
• Mixing an s orbital with one of the p orbitals generates two
equivalent sp hybrid orbitals. Note that the number of
hybrid orbitals is equal to the number of atomic orbitals that
are hybridized. The set of two sp hybrid orbitals has a linear
arrangement. The angle between the orbitals is 180˚.
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Step 1: 1-s and 3-p atomic orbitals
blend to make 4 sp3 molecular orbitals
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Water & Methane are both sp3 hybridized
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The puzzle pieces have to fit!
• Biological molecules recognize and interact with each
other with a specificity based on molecular shape
• Molecules with similar shapes can have similar
biological effects
• Long ago you accepted that the 4 DNA bases pair
A-T and G-C, but WHY??
• Why can’t the A pair with the C?
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The puzzle pieces have to fit!
• Why does a molecule have shape in the first place?
• That has to do with IMFs holding things in “place” defining
its 3-D shape. The big people word for “shape” is
“conformation”.
• In the case of A-T vs. G-C, it’s all about lining up that IMF
called “hydrogen bonding”; 2 sites for A-T and 3 for G-C
Ever heard of “a runner’s high”?
• Our body manufactures endorphins which are made
by the pituitary gland and bind to the receptors in
the brain that relieve pain and produce euphoria
during times of stress, such as intense exercise.
• Opiates such as morphine and heroin are structured
similarly, thus can bind with the receptors . These
“bindings” are actually those electrostatic attractions
we call IMFs.
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The boxed regions are shaped similarly!
Carbon
Hydrogen
Natural endorphin
Nitrogen
Sulfur
Oxygen
Morphine
(a) Structures of endorphin and morphine
Pain killer!
• The brain receptors “bind” with either with similar
results.
Natural
endorphin
Brain cell
Morphine
Endorphin
receptors
(b) Binding to endorphin receptors
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Chemical Reactions Make and/or
Break Chemical Bonds
Lets get two things straight before going
further…
1. Energy must be added to a system to BREAK a
chemical bond.
2. Energy is released when chemical bonds are MADE.
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Chemical Reactions Make and/or
Break Chemical Bonds
• Chemical reactions are the making and breaking of
chemical bonds
• The starting molecules of a chemical reaction are
called reactants
• The final molecules of a chemical reaction are called
products
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Chemical Reactions Make and/or
Break Chemical Bonds
• Chemical reactions are the making and breaking of
chemical bonds
• The starting molecules of a chemical reaction are
called reactants
• The final molecules of a chemical reaction are called
products
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Photosynthesis is a mighty important
chemical reaction!
• Photosynthesis is an important
chemical reaction
• Sunlight powers the conversion
of carbon dioxide and water to
glucose and oxygen
6 CO2 + 6 H2O → C6H12O6 + 6 O2
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Chemical Equilibrium
• All chemical reactions are reversible: products of the
forward reaction become reactants for the reverse
reaction
• Chemical equilibrium is reached when the rate of the
forward reactions is equal to the rate of the reverse
reaction
• Chemical equilibrium does NOT mean “equal”
amounts of reactants and products are present, but
rather that their concentrations have stabilized in a
constant ratio!
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