Chapter Two
Water: The Solvent for
Biochemical Reactions
Page 33-48
Paul D. Adams • University of Arkansas
cengage.com/chemistry/campbell
Chapter Outline
2-1 Water and Polarity
2-2 Hydrogen Bonds
2-3 Acids, Bases and pH
2-5 Buffers
Water and Polarity
Water is principal component of most cells.
Geometry of water molecule and its properties as a
solvent play major roles in determining the properties of
living systems.
Electronegativity
Electronegativity
• Electronegativity: The tendency of an atom to
attract electrons to itself in a chemical bond.
• Oxygen and Nitrogen,
more electronegative
than carbon and
hydrogen
• Fluorine is most
electronegative (4)
Polar and Nonpolar bonds
When two atoms with the same electronegativity form a
bond, electrons are shared equally, such as in C-H
bond in methane (CH4). Such bond is non-polar
If atoms with different electronegativity form a bond, the
electrons are not shared equally.
In O-H bond in water, oxygen is more electronegative,
so bonding electrons are closer to oxygen. This
difference give rise ẟ+ and ẟ-. Such bonds are called
polar bond
Polar Bonds & Molecules
• Molecules such as CO2 have polar bonds but, given
their geometry, are nonpolar molecules; that is, they
have a zero dipole moments
What makes water polar
Water is a bent molecule and uneven sharing of electrons is not cancelled out
as in CO2
• Bonds with positive and negative ends are called dipoles
Electrostatic Attraction
• In a covalent bond, the electrostatic attraction
between the positively charged nuclei of the
bonded atoms and the negatively charged electrons
they share., the atoms are held together by
the electrostatic attraction between the positively
charged nuclei of the bonded atoms and the
negatively charged electrons they share ...
Solvent Properties of H2O
• Why do some chemicals dissolve in water while
others do not?
• Polar nature of water determines its solvent properties.
• Ionic compounds (KCl, NaCl) having full charges and
polar compounds (Ethyl alcohol, Acetone) having partial
charges tend to dissolve in water
• Electrostatic attraction is responsible
Noncovalent Bonds
• Ionic Bonds: Held together by positive and negative ions
• Salt Bridge: Attraction that occurs when oppositely
charged molecules are in close proximity.
• Ion-dipole interactions: When ions in solution interact
with molecules with dipoles, e.g., KCl dissolved in H2O
Lysine
Hydration Shells Surrounding Ions in Water
Noncovalent Bonds
• van der Waals Forces: bonds that do not involve
electrostatic interactions. These forces are of three types:
1-Dipole-dipole interactions:
Forces that occur between molecules with dipoles, with
partial positive side of one molecule attracting partial
negative side of another molecule
Ion-dipole and Dipole-dipole Interactions
• Ion-dipole and dipole-dipole interactions help
ionic and polar compounds dissolve in water
2-Dipole-Induced Dipole Interaction
Dipole- induced dipole interactions: A permanent dipole in a molecule, when it
comes into close contact with any molecule have no dipole, can induce a transient
dipole in the other molecule
3-Induced Dipole-Induced Dipole
Interactions or London Dispersion Forces
When two molecules lacking dipoles bump into each other, they distort
each other’s electron cloud, thereby creating a brief interaction between
these induced dipoles
That explain, why non polar molecules have attraction for one another
Biochemically-Relevant Bond Energies
Why some substances dissolve in water?
• Water being polar molecule, can make ion-dipole
bonds with ionic compounds or dipole-dipole bonds
with polar neutral compounds.
• Ionic and polar substances tend to dissolve in water
and they are referred to Hydrophilic (water-loving)
• Hydrocarbon are non polar. Ion-dipole and dipole-
dipole interactions donot occur for non polar, so
these compounds tend not to dissolve in water and
referred as Hydrophobic (water-fearing)
Hydrophobic and Hydrophilic Substances
Hydrophobic interactions: Attractions between non polar
molecules, also called hydrophobic bonds
Oil Slick
• a film or layer of oil floating on an expanse of water.
Amphipathic
• Amphipathic: A single molecule may have both polar
(hydrophilic) and nonpolar (hydrophobic) portion.
• A long chain fatty acid have polar carboxylic acid
(head) and a long nonpolar hydrocarbon portion (tail)
is an example of amphipathic
Italian dressing
Amphipathic molecules
Fatty acid
Micelle formation by amphipathic molecules
• Micelle: a spherical arrangement of organic
molecules in water solution clustered so that
• their hydrophobic parts are buried inside the sphere
• their hydrophilic parts are on the surface of the sphere
and in contact with the water environment
Tug of War
Hydrogen Bonds (Special dipole-dipole
interaction
• Hydrogen bond: the attractive interaction between
dipoles when:
• positive end of one dipole is a hydrogen atom bonded
to an atom of high electronegativity, most commonly O
or N, and
• the negative end of the other dipole is an atom with a
lone pair of electrons, most commonly O or N
• Hydrogen bond is non-covalent
Hydrogen Binding Sites
Each HF molecule has one hydrogen-bond
donor and three hydrogen-bond acceptors.
Each H2O molecule has two donors and two
acceptors (Four hydrogen bonds). Each NH3
molecule has three donors and one acceptor.
Biochemically-Relevant Bond Energies
Interesting and Unique Properties of Water
• Each water molecule can be involved in 4 hydrogen
bonds: 2 as donor, and 2 as acceptor
• Due to the tetrahedral arrangement of the water
molecule.
Difference between liquid water and ice
crystals
•In liquid water, H-bonds are constantly breaking and
new ones are constantly forming
•Some molecules breaking off and others joining
cluster.
• A cluster can break up and reform in 10-10 to 10-11
seconds
• Ice crystal has more stable arrangement of hydrogen
bonds
Strength of Hydrogen Bonding
• Even though hydrogen bonds are weaker (20 kj/mol)
than covalent bonds (O-H, 460 Kj/mol), they have a
significant effect on the physical properties of water
such as melting point, boiling point and density
Expansion of Water
• Ice has lower density than liquid water because ice
crystal is less densely packed. Thus, Ice cubes and
icebergs float
• Water expand on cooling, while others contract
• Cooling system of cars require antifreeze to prevent
freezing and expansion of water, which could crack
the engine.
• Aquatic organisms can survive in cold climates
because of density difference between ice and liquid
water
Hydrogen Bonding Between Polar Groups
and Water
Hydrogen bonding play a role in the behavior of water as a solvent
If polar solute can serve as donor or acceptor of hydrogen bonds, it can form H-bond
and dipole dipole interactions with water
Alcohols, amines, carboxylic acids, esters, aldehydes and ketones can form hydrogen
bond with water and are soluble in water
Other Biologically Important Hydrogen bonds
• Hydrogen bonding is important in stabilization of 3-D
structures of biological molecules such as: DNA
(complementary bases), RNA, proteins (α-helix,β-
pleated)
Acids, Bases and pH
• Acid: a molecule that acts as a proton (hydrogen ion)
donor. This is also known as Bronsted acid
• Strong base: a molecule that acts as a proton
acceptor
The degree of dissociation of acids in water ranges from
complete dissociation for strong acid to no dissociation for a
very weak acid
Acid Strength (Acid dissociation constant Ka)
• One can derive a numerical value for the strength of an
acid (amount of hydrogen ion released when a given
amount of acid is dissolved in water).
• Describe by Ka:
• [ ] = molar concentration
• Greater Ka, stronger the acid
• Written correctly,
Ionization of H2O and pH
• Extend of self dissociation of water is small
• Lets quantitatively examine the dissociation of water:
•Molar concentration of water is quite large compared to any possible
concentration of solute and considered as constant. Numerical value
is (55.5 M)
Ionization of H2O and pH
• Kw is called the ion product constant for water.
• [H+] = [-OH] =10-7
Must define a quantity to express hydrogen ion concentrations…pH
• When pH = 7 , solution is neutral
• pH < 7 , it is acidic, pH > 7 is alkaline (or basic)
• One pH unit represents a ten-fold change in H+ concentration
In biochemistry, most acids are weak acids, Their Ka is well below 1.
pKa = -logKa
Smaller Pka, stronger the acid. This is reverse situation with ka
Buffers
• buffer: a solution whose pH
resists change upon addition
of small to moderate amounts
of a strong acid or base
• consists of a weak acid
and its conjugate base
• Examples of acid-base
buffers are solutions
containing
• CH3COOH and
CH3COONa
• H2CO3 and NaHCO3
• NaH2PO4 and Na2HPO4
