Download Hydrogen Bonding

New Biology
For Engineers and
Computer Scientists
Introduction & the
Chemistry of Life
Shu-Ping Lin, Ph.D.
Institute of Biomedical Engineering
E-mail: [email protected]
Website: http://web.nchu.edu.tw/pweb/users/splin/
Course Outline & Evaluation
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TA: 康毓珊 （BME）e-mail: [email protected]
Attendance 10％
Homework 10%
Quiz 1 10 %
Quiz 2 10 %
Mid Term 30％
Final Exam 30％
Office Hours: Thur.
10:00am~ 11:30am
(精密大樓3F）
Phone:22840734 X22
New Biology for Engineers and
Computer Scientists, by Aydin
Tozeren and Stephen W. Byers,
Publisher: Prentice Hall, 2003.
The Chemistry of Life -- Important
Elements in Living Organisms
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Six elements compose 98% of the mass of all living organisms.
Hydrogen
Carbon
Nitrogen
Oxygen
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Phosphorus
Sulphur
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Mineral ions, such as iron, calcium, and magnesium, account for as much as
2 percent of the human body weight.
The Chemistry of Life - Atoms
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The basic unit of each chemical element is the atom. Atoms have a large nucleus,
composed of protons and neutrons held together by the Strong Force. The
electrons "orbit" the nucleus, attracted by the Electrical Force.
The number of protons determines the chemical element, and the number of
neutrons determines the isotope of the element.
No net electrical charge  Ions -- when atoms gain or lose electrons
http://cass.ucsd.edu/public/tutorial/scale.html
Carbon
Orbitals
around a
nucleus
Figure1.1
Periodic Table
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Periodic table indicating the atomic properties of all elements found on earth.
Periodic Table Explorer is a simple Periodic Table software.
http://www.technosamrat.com/freewares/periodic-table-explorer/
Chemical Reactivity & Octet Rule
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Chemical Reactivity – The tendency of a substance
to undergo chemical changes in a system
The valence electrons are responsible for the
combining capacity of atoms.
Octet Rule – A tendency for eight electrons in the
outermost orbital to undergo chemical changes
It is the cardinal rule of bonding.
It is the gain in stability when atoms have a full
complement of eight electrons in their valence shells.
The bonding in carbon dioxide (CO2):
all atoms are surrounded by 8
electrons, fulfilling the octet rule
http://en.wikipedia.org/wiki/Octet_rule
Hydrogen (H)
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Has single valence electron
Form a stable compound 
hydrogen must either empty its
electron orbital by giving it up to
another atom (ionic bond) or fill
the orbital by sharing electrons
with another atom (covalent
bond)
Form ionic compounds with
metals
Form molecular compounds with
nonmetals
Carbon (C)
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An organic element with the second
smallest atomic number
Has six protons (Most carbon atoms
possess six neutrons.)
Has 4 valence electrons (half-filled
outermost orbital)  allows carbon
to form chemical bonds with various
other atoms
The resulting chemical versatility 
is essential for reactions of
biological metabolism and
propagation
Nitrogen (N)
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Has seven protons
Has five valence electrons
Outer electron orbit  only
slightly above half filled
The compounds of nitrogen,
though not as numerous as
those of carbon, are just as
varied in function.
Phosphorus (P)
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Has fifteen protons
Has five valence electrons
Abundant in rock formation
in the way of calcium
phosphate
Inorganic salt – Calcium
phosphate  60% of bones
in the human body
Oxygen (O)
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Has six valence electrons at
its outermost shell
Has eight protons
Abundant in the earth’s
crust
Gain 2 electrons to achieve
the stable octet
configuration
Sulfur (S)
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Has six valence electrons
Is found in a free-element
form in large beds several
hundred feet underground
Plays a fundamental role in
determining the threedimensional shapes of
many proteins  when
present in the amino acid
―cysteine‖
Calcium (Ca)
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Has 20 valence electrons
The most abundant mineral in the
human body
The average adult body contains in
total approximately 1 kg, 99% in the
skeleton in the form of calcium
phosphate salts.
~22.5 mmol in extracellular fluid (ECF)
~500 mmol of calcium is exchanged
between bone and ECF over a period
of 24 hours.
Ca2+ into and out of the cytoplasm
functions as a signal for many cellular
processes
Covalent Bonds
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Covalent Bonds – Formed when
electron overlaps  two atoms by
sharing one or more pairs of electrons
in order to have eight electrons in the
outermost orbitals (octet rule)
Only exception – Hydrogen ( has one
valence electron)  H2: two hydrogen
nuclei share the two electrons equally
and completely.  Each electron is
attracted to both protons, but the two
protons of the opposing hydrogen
atoms repulse each other and the
balance of these opposing forces are
the determinants of the resulting
molecular structure.
Sharing of electron pairs
in covalent bonds
Molecules
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Molecule – A combination of atoms held together by covalent
bonds
Identified by the symbols of elements that constitute it
Oxygen molecule (O2) – Composed of two atoms
Quantity of molecules – One mole per liter (mole/L , or M)
The clouds of unshared electrons are larger than those of
shared electrons.  Charge clouds of nonbonded electrons
push the bonded pairs toward each other.
Clouds of two unshared pairs are larger than the those of
bonded electrons in H2O.
H2O:
104.5 °
3D Structure of an
Organic Molecule
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Shared Pair – The pair of electrons involved in covalent bonding
Unshared Pair – The pairs of the outermost electrons not
involved in bonding
Each bond and unshared pair form a charge cloud that repels
all other charge clouds  repulsions between the charge
clouds determine the 3D shape of the molecules physical
rules:
Electron pairs spread as far apart as possible to minimize
repulsive forces.
Since the cloud of unshared pair occupies more space 
Repulsive Forces
between the unshared pairs of electrons >
Repulsive Forces
H2O:
104.5 °
between two shared pairs
Three Commonly Known Molecules Methane, Ammonia, Water
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Bond Length is not
fixed  acts as much
as a stiff spring
Methane  the shape of
a tetrahedron with bond
angles equal to 109.5 °
Ammonia  the shape of
a tetrahedron with the
unshared electron pair
occupying one corner of
tetrahedron (bond angle:
107°)
Water  triangular in
shape, with the bond
angle 104.5° between two
hydrogen-oxygen
covalent bonds
Double & Triple Bonds
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Single Bonds – Bonds in which a single pair of electrons is
shared between two atoms  straight-line segment
connecting two atoms
Double Bonds – Four electrons are shared between two
atoms  double-line segment
Single bond can rotate freely about the bond axis
Double bond occupies more space than a single bond  more
repulsive, Ex: H-C-H bond angle: 116°, H-C=O: 122°
Triple Bonds – Sharing of six electrons between two atoms
 rarely observed in organic materials
N-N (unstable)
N≡N
Organic Compounds –
Small Carbon Molecules
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Double bonds between carbon atoms occur
often in biological compounds.
Physical Strength of a Covalent Bond
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Important covalent bonds
Covalent bonds constitute stable links between
in biological systems:
atoms and are the strongest of bonds
connecting molecules.
Rupture of covalent bonds can occur in two ways:
(kcal/mol)
Linkage between two atoms are broken
symmetrically to provide a pair of free radicals
(derived from the unpaired spins of their electrons)
 often observed in the breaking of identical or
similar atoms
The bonds are broken asymmetrically to
produce a pair of ions  electron deficient and
electron rich
The energy required to break a covalent bond is
much greater than the internal (thermal) energy
Repulsive forces
available at body temperature (0.6 kcal/mol) 
The energy released by the formation of new bonds between the unshared
can break covalent bonds.
pairs of electrons make
N-N bonds less stable
Covalent bonds between carbon atoms are highly
than C-C bonds.
stable.
Electronegativity and Polar Bonds
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Electronegativity – The capacity of an atom to attract electrons from a
neighboring atom  measured on a scale from 4 (fluorine, the most
electronegative element) to a hypothetical 0
Covalent bonds between two atoms with comparable electronegativity correspond
to an equal sharing of electrons between the two nuclei.
Polar covalent bond or simply a polar bond – Highly electronegative atoms bond
with weaker electronegative atoms  partial charges (δ+ and δ-) in different parts
of the molecule
Unequal sharing of electrons between atoms in a molecule results in charge
polarity.
Polar molecules dissolve easily in water and interact with other polar molecules.
Nonpolar molecules dissolve better in a nonpolar environment, such as acetone or
ethanol.
Questions
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Why is sulfur less electronegative than
oxygen?
Ans: The electro negativity chart shows that oxygen is the most
electronegative atom of bioelement, that is to say, oxygen has an electro
negativity of 3.5 whereas sulphur has an electro negativity of just 2.5.
Therefore, oxygen will take two electrons to fill its s and p sub shells. Even
though the charge of four oxygen molecules is -8, each molecule has a -2
charge.
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What significant role does sulfur play in living
organisms?
Ans: Sulfur is found in the amino acids methionine, and cysteine. These
amino acids are known as the sulfur-bearing amino acids, which are
considered the building blocks of all proteins. Sulfur is important for the
regulation of plant growth and development since it is the main source
used by photosynthetic organisms.
Ionic Compounds and
Electrostatic Bonds
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Compound – One of the two interacting atoms is much more electronegative than
the other (one or more electrons in the less electronegative atom are transferred to
the more electronegative atom)  Two electrically charged particles are called ions.
Cation – Ion with a positive charge (Ca2+ or H+)
Anion – Ion with negative charge (OH-)
Ionic Bond – Electrostatic force holds two ions together due to their differing
charges.
Ionic Compounds – High melting points, conduct electricity in the molten state,
and tend to be soluble in water (Na+Cl- and Ca32+(PO43-)2)
Ammonia (NH3) – Nitrogen forms three covalent bonds with 3 hydrogen atoms 
one of the outermost electron pairs of nitrogen is not shared  dissolved in water,
ammonia picks up a hydrogen ion (H+) shared a previously unshared bond to
become an ammonium ion (NH4+), has a net positive charge of 1
Ethanol – Has no charge, but ionized ethanol has a charge of -1 due to acidity of
the medium
Water and Hydrogen Bonds
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At RT, oxygen and hydrogen: gaseous form; water molecule: liquid
state due to Hydrogen Bonding between water molecules
Water – Has important physical properties that make life possible on
the Earth, ability to dissolve many other substances, serves as a
medium in which a great variety of chemical changes occur
Hydrogen Bonding – Caused by the polar nature of covalent H—O
bonds that hold together water molecules  Hydrogen atoms of
water molecules are attracted to the unshared electrons of oxygen
atoms of adjacent water molecules.
Hydrogen bonds (3-dots):
Water molecules form transient
hydrogen bonds with several
others, creating a fluid network.
Hydrogen Bonds
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Water molecules form hydrogen bonds with other polar molecules including those
of nitrogen and oxygen compounds.
Nitrogen and oxygen are more electronegative than hydrogen  Covalent N—H
and O—H bonds are polar bonds, the H atoms in these bonds can participate in
hydrogen bonding.
Amino (—NH2) and hydroxyl (—OH) groups 
Two atomic groups that often engage in hydrogen bonding in living systems
The presence of these two groups makes many molecules soluble in water
Water molecules cluster around cations (Na+) and anions (Cl-) in solutions, blocking
their association into a solid.
Hydrogen Bonds – Also form between different parts of large polar molecules
such as proteins, much weaker than a covalent bond in physical strength and
duration, greatly influence the physical properties of biological substances
Hydrogen bonds
between methanol
(CH3OH) and water
Hydrogen bonds
between methylamine
(CH3NH2) and water
Van der Waals Attraction
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Hydrophobic or water fearing – Inability of nonpolar hydrocarbon (compounds of hydrogen
and carbon, nonpolar because of sharing bonding electrons equally) molecules to form
hydrogen bonds with water
Hydrophobic effect –
Drives a number of very important biological phenomen – Formation of cell membranes by lipid
bilayers 
 Hydrophobic nature of the hydrocarbon chains of the lipids make up the bilayer
 Lipid bilayers have a hydrophobic core with the hydrophilic heads facing the polar
environments outside or inside the cell
Drives the folding of proteins 
 Hydrophobic amino acids (building blocks of proteins) tend to fold away from more polar
regions
 Form a hydrophobic pockets or clefts  Provide a binding site for small hydrophobic
molecules such as steroid hormones
Van der Waals interactions – Brief and weak attraction, Van der Waals forces act on
nonpolar molecules brought together by a polar solvent.
Nonpolar molecules disturb
the hydrogen bonds
between water molecules,
forcing water to form a cage
around them
Hydrophobic tails
extend into the air
to avoid contact
with water
Types of Bonds between
Molecules, Atoms or Ions
*Four types of bonds between atoms or ions:
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Ionic Bonding – Electrons are transferred from one atom to another
producing a charge on both atoms (ions), in order to keep them together.
 NaCl
Covalent Bonding – Sharing of electrons between atoms resulting in an
overlap of their electron orbitals.  Diamond
Metallic Bonding – Formed when closely packed atoms bond by sharing
electrons from their inner electron shells, whereas the outer shell electrons
are free to form a sea of electrons  Giving metals their high electrical and
thermal conductivities
Van der Waals Bonds – Weak electrostatic interactions that hold adjacent
sheets of talc or graphite together, act on nonpolar molecules brought
together by a polar solvent.
*Bonds between molecules:
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Hydrogen bonding – Caused by the polar nature of covalent H—O bonds
that hold together water molecules
Acids and Bases
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Concentration of hydrogen ions in a solution: pH= -log[H+], pH of
cell fluid (cytoplasm)= 7.2~ 7.3
Acid – Donate protons (H+), pH<7, sour taste, and cause the purple dye
litmus to turn red, Ex: HCl  H+ and Cl-, —COOH (carboxyl group
molecules)  H+ and COO-, gastric fluids, vinegar, soda, lemon juice, and
contracting muscle cell fluid
Base – Accept protons, pH>7, taste bitter, feel slippery, and turn the
purple dye litmus blue, Ex: NH3  NH4+, —NH2 (amino group in biological
molecules)  —NH3+, sweat, human blood ([H+]=4×10-8M), and
household ammonia
Functional groups in biological molecules undergo acid-base reaction in
living systems  water  H+ and OH-; H+ affects the rates of chemical
reactions
Cellular pH play an important role in cell division and cell growth.
Many proteins denature in acidic conditions except gastric enzymes and
lysosomal enzymes.
Cellular membranes pump protons from one side to the other  Create a
pH gradient  An important component in the generation and
storage of energy in mitochondria
Chemical Reactions
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Occurs when atoms combine or change
binding partners
Arrow in the following equation  The
direction of the chemical reaction
The number of atoms on the left-hand side
of the equation = The number of atoms on
the right-hand side
C3 H8  5O2  3CO2  4H 2O
Reactants
Products
Energetics of Chemical Reactions
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In thermodynamics – ―System‖: the part of the universe that is
of interest, ―surroundings‖: the rest of the universe
Open system – Exchange matter and energy with its
surrounding, Ex: living cells (take up nutrients, release waste
products, and generate work and heat
1st law of thermodynamics: the conservation of energy law
and states that energy can be neither destroyed nor created
Energy gained by
an open system
U  q  w
Work done by the system on
the surroundings
Heat absorbed by the surroundings (release heat q<0: Exothermic process  Endothermic)
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2nd law of thermodynamics: some of the energy involved
irreversibly loses its ability to do work; S is entropy
Entropies of
the system
S1  S2  Su  0
Entropies of the
surroundings
Spontaneous reaction
Entropies of the universe
Chemical Reactions
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Combine the 1st and 2nd laws of thermodynamics 
H  G  TS
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△H: Overall change in bond energy due to reaction (kcal/mol)
△S: Change in entropy (measure of the energy lost to disorder in system)
△G: Change in free energy
Free energy (△G) – Energy can be used to do work
For spontaneous chemical processes occur at constant
temperature  S  q / T q: Heat imported into the system
1
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T: Absolute temperature
Spontaneity for constant pressure and temperature processes,
the criterion of  △G < 0
A + B  C + D
 △G < 0, exergonic or energetically favorable, △G products<△G reactants ,
spontaneously forward reaction
 △G > 0, endergonic, reverse reaction
 △G = 0, chemical equilibrium (at steady state), both forward and
reverse reactions occur at equal rates
Free Energy Change (△G) in a
Chemical Reaction
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2 parts to the free-energy change that occurs in a
chemical reaction –
 △G0 (Change in standard free energy),
 Concentrations of reactants and products
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G  G o  RT ln{([C ][ D]) /([ A][ B])}
R(gas constant)=1.978 cal/mol, T=273.15 °+ ℃
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△G increases with
increasing temperature and
increasing concentrations of
reactants
Chemical Reactions in
Biological Systems
Many chemical reactions in biological systems are
endergonic reactions △G>0  reverse reaction
 How to proceed in the forward direction?  Additive
property of free energy
 Chemical reaction with positive △G may be coupled to
a reaction with a negative △G of a larger magnitude
A  B + D △Go1 =12 kcal/mol
D  X + Y △Go2 =-16 kcal/mol
 △G3= △Go1+ △Go2=- 4 kcal/mol <0
 exergonic (reactions operate in the forward direction)
 This exergonic reaction coupled the hydrolysis of ATP
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ATP: The Standard
Energy Carrier
Nucleotide
monophosphate
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Nucleoside
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Pi
Pi
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H++PO43-
ATP: nucleotide adenine,
sugar ribose, and
triphosphate unit
Participates in most cellular
reactions and processes that
require energy
Very stable molecule
ATP hydrolysis reaction by
the enzyme – ATPase,
releases 7.2 kcal/mol
ATP + H2O  ADP+ Pi + H+
PO43-