Biomolecules

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Biomolecules
A review of some of the chemical principals that
govern the properties of biological molecules.
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Learning Objectives
Identify the types of bonds, common in
biomolecules, formed between carbon and hydrogen,
oxygen, nitrogen, and other carbon
atoms. Identify and recognize by chemical
structure the common functional groups. Describe
the difference between molecular configuration
and molecular conformation. Define
stereochemistry and stereoisomers. Describe the
geometry of bonds in sp, sp2, and sp3 hybridized
carbon. Describe the geometry of a carboxylic
acid, a carbonyl group, and an amine.
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Chemical composition and bonding
Two atoms with unpaired electrons in their outer
shells can form covalent bonds with each other by
sharing electron pairs.
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Biomolecules are compounds of carbon
Versatility of carbon in forming covalent single,
double, and triple bonds.
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Shell a region of space around a nucleus where
electrons are found.
Electronic structure of the atom
five 3d orbitals
Nucleus
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3px, 3py, 3pz
3s
2px, 2py, 2pz
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2s
Energy
Orbital a region of space around the nucleus of
an atom that can hold up to two electrons.
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1s
Shells
Orbitals
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Valence shell the outermost (highest energy)
shell of an atom that contains electrons. Valence
electrons outer shell electrons.
Lewis structure the symbol of the chemical
element surrounded by a number of dots equal to
the number of electrons in the outer or valence
shell.
The atomic number of an atom indicates the total
number of protons in the nucleus of an atom the
atomic number also indicates the total number of
electrons (in the ground state) which surround
the nucleus.
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The Chemical Bond
According to the Lewis model, elements bond
together in such a way that each atom involved in
the chemical bond acquires a completed outer
shell electron configuration.
There are two ways this can occur 1. An atom
may lose or gain electrons to acquire a completed
outer shell. An atom that loses an electron
becomes a positively charged ion (a cation) an
atom that gains an electron becomes a negatively
charged ion (an anion). A chemical bond between
a positively charged ion and a negatively charged
ion is called an ionic bond.
NaCl (sodium chloride) is an example of a
compound in which the Na atom and the Cl atom are
bonded with an ionic bond.
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In the formation of an ionic bond, electrons are
transferred from the valence shell of an atom of
lower electronegativity to the valence shell of
an atom of higher electronegativity.
NaCl
Cl is much more electronegative than Na.
Therefore, Na transfers an electron to Cl Na
becomes Na, and Cl becomes Cl-.
An ionic bond is a chemical bond formed by the
attractive force between positive and negative
ions (electrostatic attraction).
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Sodium, Na
Atomic Number 11
Electron configuration 1s2 2s2 2p6 3s1
If Na loses one electron from the 3s orbital,
what is left is a completely filled outer shell
2s2 2p6 .
In losing one electron, Na acquires a positive
charge of one Na
Chlorine, Cl
Atomic Number 17
Electron configuration 1s2 2s2 2p6 3s2 3p5
If Cl gains an electron, then the outer shell
will be completely filled 3s2 3p6 In gaining an
electron, Cl becomes a negatively charged ion Cl-
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2. An atom may share electrons with one or more
other atoms to complete its outer shell. A
chemical bond formed by sharing electrons is
called a covalent bond.
The simplest example of a covalent bond is found
in the hydrogen molecule H2. When two hydrogen
atoms bond, the single electrons from each
combine to form an electron pair. This shared
pair of electrons completes the valence (or
outer) shell of each hydrogen.
H 1s1
H
H
Shared electron pair
H
H
This line between the two hydrogens symbolizes
the covalent bond formed by the sharing of a pair
of electrons.
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Electronegativity
Electronegativity is a measure of the force of an
atoms attraction for electrons (that it shares
in a chemical bond with another atom.)
Fluorine is the most electronegative element.
Electronegativity of the elements increases from
left to right in the Periodic Table, and
decreases from top to bottom. Oxygen and nitrogen
are considered to be highly electronegative in
comparison with carbon and hydrogen. In other
words, oxygen and nitrogen tend to attract
electrons much more strongly than either carbon
or hydrogen.
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Covalent bonds can be polar or nonpolar.
Although all covalent bonds involve the sharing
of electrons, they differ in the degree of
sharing the more electronegative atom involved
in the bond tends to attract the shared electrons.
A nonpolar covalent bond is one in which the
difference in electronegativities between boded
atoms is less than 0.5.
Example C H
C H
A polar covalent bond is one in which the
difference in electronegativities is between 0.5
and 1.9.
d-
d
H Cl
Example H Cl
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An important consequence of the unequal sharing
of electrons in a polar covalent bond is that the
more electronegative atom gains a greater
fraction of the shared electrons, and acquires a
partial negative charge indicated by the symbol
d-. The less electronegative element has a lesser
fraction of the shared electrons and acquires a
partial positive charge d
Example
d-
O
d
d
H
H
Water not only has polar bonds, water is a polar
molecule.
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Why is this electronegativity stuff and polar vs.
nonpolar bonds important? Because it tells you
something about the tendency of a particular bond
(or functional group) to dissociate in water to
come apart, leaving two charged species, one
positive and one negative. This is an important
concept in understanding acids and bases, and
hydrogen bonding.
Ionic bonds (salts like NaCl,
KCl, Na2SO4) Polar covalent
bonds(-OH, -NH) Nonpolar covalent bonds (C-C,
C-H, C-O, C-N)
Easily dissociate (and dissolve) in water
Somewhere in between
Do not dissociate in water
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Salts (compounds with ionic bonds) that are
important in human (vertebrate) biochemistry
and/or physiology
NaCl, KCl NaH2PO4, Na2HPO4
(electrolytes)
(buffers)
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3Ca3(PO4)2 Ca(OH)2
(hydroxylapatite occurs in bone)
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Hydrogen
Elemental symbol H Atomic Number 1
Electron Configuration 1s1
Number of valence electrons 1
Number of bonds hydrogen can form with another
atom 1
H
H
Lewis structure
Conventional designation of one single bond
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Carbon
Elemental symbol C Atomic Number 6
Electron Configuration 1s2 2s2 2p2
Number of valence electrons 4
Hybridization of orbitals sp3, sp2, sp
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Carbon
sp3 hybridization one s and three p orbitals
four equivalent, sp3 hybridized orbitals
C
C
In sp3 hybridization, carbon has four single
bonds to other atoms.
The sp3 orbitals (electron distribution) are
directed toward the corners of a regular
tetrahedron this geometry provides the maximum
separation of the electrons in each of the four
hybrid orbitals.
A regular tetrahedron
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The single bonds in sp3 hybridized carbon are
directed toward the corners of a regular
tetrahedron.
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sp3 hybridized carbon is tetravalent
it always forms
four bonds.
Carbon can form single bonds with hydrogen and
also with other carbon atoms.
Alkanes
Methane has four single bonds one single bond
from carbon to each of four hydrogens.
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Organic molecules have specific three-dimensional
shapes.
Normal lines bonds in the plane of the
paper Dashed lines (dashed wedges) bonds
receding behind the page Solid wedges bonds
coming out in front of the page
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Carbon
sp2 hybridization one 2s orbital and two 2p
orbitals three equivalent sp2 hybridized
orbitals
Carbon is sp2 hybridized when it forms a double
bond with another carbon, oxygen, or nitrogen
atom, and two single bonds with other atoms
(hydrogen, carbon, etc.)
ethylene
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Alkenes
Carbon-carbon double bond
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Ketones, carboxylic acids, esters
Carbon-oxygen double bond
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Carbon-carbon single bonds have freedom of
rotation
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Carbon-carbon double bonds are shorter than
single bonds, and do not allow free rotation.
The two doubly bonded carbons and the directly
attached atoms all lie in the same rigid plane.
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Covalently linked carbon atoms in biomolecules
can form linear chains, branched chains, and
cyclic structures. To these carbon skeletons are
added groups of other atoms, called functional
groups, which confer specific chemical properties
on the molecule.
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Most biomolecules can be considered as
derivatives of hydrocarbons. The carbon-carbon
backbone of hydrocarbons is very stable. One or
more of the hydrogens of these hydrocarbons can
be replaced by a variety of functional groups to
yield different families of organic compounds.
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Histidine
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Epinephrine
secondary amino group
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Acetyl-coenzyme A
(Acetyl-CoA)
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Three-dimensional Structure Configuration and
Conformation
Stereochemistry the arrangement of a molecules
constituent atoms in three-dimensional space.
Stereoisomers compounds that have the same
composition and the same order of atomic
connections, but different three-dimensional
molecular arrangements.
Isomers any two molecules with the same
molecular formula but different arrangements of
molecular groups.
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Stereochemical representations of molecular
structure
Solid wedge atom at the wide end (H2N) projects
out of the plane of the slide toward the viewer.
Dashed wedge atom at wide end (C) projects
behind the plane of the slide
Structural formula in perspective form.
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Stereochemical representations of molecular
structure
Space-filling model, in which each atom is
represented as a sphere with its correct relative
van der Waals radius.
Ball and stick model showing relative bond
lengths and bond angles.
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Configuration denotes the fixed spatial
arrangement of atoms in a molecule that is
conferred by the presence of either (1) double
bonds, around which there is no freedom of
rotation (2) chiral centers, around which
substituents are arranged in a specific sequence.
Configurational isomers cannot be interconverted
without temporarily breaking one or more covalent
bonds.
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An example of configurational isomers
Geometric or cis/trans isomers they cannot be
interconverted without breaking bonds.
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In the vertebrate retina, the initial event in
light detection is the absorption of visible
light by 11-cis-retinal. The energy of the
absorbed light converts 11-cis-retinal to the
all-trans-retinal, triggering events which lead
to nerve impulses.
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Molecular conformation refers to the spatial
arrangement of substituent groups that, without
breaking any bonds, are free to assume different
positions in space because of the freedom of bond
rotation.