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CHAPTER 2
Water and Aqueous Solutions
Learning goals:
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What kind of interactions occur between molecules
Why water is a good medium for life
Why nonpolar moieties aggregate in water
How dissolved molecules alter properties of water
How weak acids and bases behave in water
How buffers work and why we need them
How water participates in biochemical reactions
Structure of the Water Molecule
• Octet rule dictates that there are four
electron pairs around an oxygen atom
in water
• These electrons are in four sp3 orbitals
• Two of these pairs covalently link two
hydrogen atoms to a central oxygen
atom
• The two remaining pairs remain
nonbonding (lone pairs)
• Water geometry is a distorted
tetrahedron
• The electronegativity of the oxygen
atom induces a net dipole moment
• Because of the dipole moment, water
can serve as both a hydrogen bond
donor and acceptor
Hydrogen Bonds
• Strong dipole-dipole interaction that arises between an acid
(proton donor) and a base (proton acceptor)
• Typically 4–6 kJ/mol for bonds with neutral atoms,
and 6–10 kJ/mol for bonds with one charged atom
• Can only form between H and N, O, F
• Hydrogen bonds are strongest when the bonded molecules are
oriented to maximize electrostatic interaction
• Ideally the three atoms involved are in a line
Hydrogen Bonding in Water
• Water can serve as both
– an H donor
– an H acceptor
• Up to four H-bonds per water molecule gives water its
– anomalously high boiling point
– anomalously high melting point
– unusually large surface tension
• Hydrogen bonding in water is cooperative
• Hydrogen bonds between neighboring molecules are weak
(20 kJ/mol) relative to the H–O covalent bonds (420 kJ/mol)
Water as a Solvent
• Water is a good solvent for charged and polar
substances
– amino acids and peptides
– small alcohols
– carbohydrates
• Water is a poor solvent for nonpolar substances
– nonpolar gases
– aromatic moieties
– aliphatic chains
Why is this? The interaction energies and
entropies are too low
Water dissolves many salts
• High dielectric constant reduces attraction
between oppositely charged ions in salt crystal;
almost no attraction at large (> 40 nm) distances
• Essentially, the water “shorts out” the
interaction b/w two ions
• Strong electrostatic interactions between the
solvated ions and water molecules lower the
energy of the system
• Entropy increases as ordered crystal lattice is
dissolved
Backup: Intermolecular Forces
Four Basic Types of Non Covalent Interactions
1.Ion-Dipole (z1m2/r3)
2.Dipole-Dipole (m1m2/r3)
3.Dipole-Induced dipole (m1a2/r6)
4.London Forces (a1a2/r6)
The strength of the interactions in each force is inversely
proportional to the distance between the interacting species
(That’s the 1/r term next to each name)
Coulomb’s Law and Interacting
Species
The force between two interacting ions is equal to the magnitude
of the charges divided by the square of the distance between them.
kq1q2
F=
2
Dr
k = proportionality
constant = 8.99x109
JmC-2
D is the dielectric
constant
The dielectric constant is a measure of the ability of the solvent
or medium containing the charges to keep them apart (D for a
vacuum is 1, D for water is ~80, D for a brick wall is ∞)
Interion and Intermolecular Forces
Physics of Noncovalent Interactions
Noncovalent interactions do not involve sharing a pair of electrons.
Based on their physical origin, one can distinguish between:
• Ionic (Coulombic) Interactions (Ionic and Ion/Dipole)
– Electrostatic interactions between permanently charged
species, or between the ion and a permanent dipole
• Dipole Interactions (Dipole/Dipole and Dipole/Ind. Dipole)
– Electrostatic interactions between uncharged, but polar
molecules
• van der Waals Interactions (London Forces)
– Weak interactions between all atoms, regardless of polarity
– Attractive (dispersion) and repulsive (steric) component
• Hydrophobic Effect (Rationale behind a lot of Biochemistry)
– Complex phenomenon associated with the ordering of water
molecules around nonpolar substances
Examples of Noncovalent Interactions
H-bonds are a
special type of
dipole/dipole
interaction
Between amino acids = Salt Bridges
“Hydrophobic Exclusions”
Hydrophobic, Dipole/Induced
Dipole and London Force
Interactions Combined make van
der Waals
Importance of Hydrogen Bonds
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Source of unique properties of water
Structure and function of proteins
Structure and function of DNA
Structure and function of polysaccharides
Binding of substrates to enzymes
Binding of hormones to receptors
Matching of mRNA and tRNA
“I believe that as the methods of structural chemistry are further applied to
physiological problems, it will be found that the significance of the hydrogen bond
for physiology is greater than that of any other single structural feature.”
–Linus Pauling, The Nature of the Chemical Bond, 1939
Hydrogen Bonds: Examples
Biological Relevance of Hydrogen Bonds
van der Waals Interactions
• van der Waals interactions have two components:
– Attractive force (London dispersion) depends on
the polarizability
– Repulsive force (Steric repulsion) depends on the
size of atoms
• Attraction dominates at longer distances (typically
0.4–0.7 nm)
• Repulsion dominates at very short distances
• There is a minimum energy distance (van der Waals
contact distance)
Origin of the London Dispersion Force
• Quantum mechanical origin
• Instantaneous polarization by fluctuating charge
distributions
• Universal and always attractive
• Stronger in polarizable molecules
• Important only at a short range
Biochemical Significance of
van der Waals Interactions
• Weak individually
– easily broken, reversible
• Universal
– occur between any two atoms that are near each other
• Importance
– determines steric complementarity
– stabilizes biological macromolecules (stacking in DNA)
– facilitates binding of polarizable ligands
The Hydrophobic Effect
• Refers to the association or folding of nonpolar
molecules in the aqueous solution
• Is one of the main factors behind:
– protein folding
– protein-protein association
– formation of lipid micelles
– binding of steroid hormones to their receptors
• Does not arise because of some attractive direct
force between two nonpolar molecules
Solubility of Polar and Nonpolar Solutes
Why are nonpolar molecules poorly soluble in water?
Low solubility of hydrophobic solutes
can be explained by entropy
• Bulk water has little order:
– high entropy
• Water near a hydrophobic solute is highly ordered:
– low entropy
Low entropy is thermodynamically unfavorable, thus
hydrophobic solutes have low solubility.
Remember to consider the Entropy of the Solute AND
the Entropy of the Solvent since both are part of the
System
•Seemingly unfavorable or unlikely Solute behaviour can often be explained when considering the changes in the
Solvent behaviour.
Water surrounding nonpolar solutes has
lower entropy
Clathrate
Cages
Origin of the Hydrophobic Effect (1)
• Consider amphipathic lipids in water
• Lipid molecules disperse in the solution; nonpolar
tail of each lipid molecule is surrounded by ordered
water molecules
• Entropy of the system decreases
• System is now in an unfavorable state
Origin of the Hydrophobic Effect (2)
• Nonpolar portions of the amphipathic molecule aggregate so that
fewer water molecules are ordered
• The released water molecules will be more random and the
entropy increases
• All nonpolar groups are sequestered from water, and the released
water molecules increase the entropy further
• Only polar “head groups” are exposed and make energetically
favorable H-bonds
Hydrophobic effect favors ligand binding
• Binding sites in enzymes and receptors are often hydrophobic
• Such sites can bind hydrophobic substrates and ligands such as steroid
hormones
• Many drugs are designed to take advantage of the hydrophobic effect
Ionization of Water
 H+ + OHH2O 
• O-H bonds are polar and can dissociate heterolytically
• Products are a proton (H+) and a hydroxide ion (OH–)
• Dissociation of water is a rapid reversible process
• Most water molecules remain un-ionized, thus pure water has
very low electrical conductivity (resistance: 18 M•cm)
• The equilibrium is strongly to the left
• Extent of dissociation depends on the temperature
Proton Hydration
• Protons do not exist free in solution.
• They are immediately hydrated to form hydronium (oxonium)
ions.
• A hydronium ion is a water molecule with a proton associated with
one of the non-bonding electron pairs.
• Hydronium ions are solvated by nearby water molecules.
• The covalent and hydrogen bonds are interchangeable. This
allows for an extremely fast mobility of protons in water via
“proton hopping.”
Proton Hopping
Water bound to proteins is
essential for their function
Water bound to proteins is
essential for their function
Water bound to proteins is
essential for their function
Ionization of Water:
Quantitative Treatment
Concentrations of participating species in an equilibrium process
are not independent but are related via the equilibrium constant:
+]•[OH-]
[H
 H+ + OHH2O 
Keq = ————
[H2O]
Keq can be determined experimentally, it is 1.8•10–16 M at 25C.
[H2O] can be determined from water density, it is 55.5 M.
• Ionic product of water:
K w  K eq  [H 2 O]  [H  ][ OH - ]  1 10 14 M 2
•In pure water [H+] = [OH–] = 10–7 M
What is pH?
pH = -log[H+]
• pH is defined as the negative
logarithm of the hydrogen ion
concentration
• Simplifies equations
K w  [H  ][OH- ]  11014 M 2 • The pH and pOH must always
add to 14
 log[ H  ]  log[ OH- ]  14 • In neutral solution, [H+] = [OH–]
and the pH is 7
pH  pOH  14
• pH can be negative ([H+] = 6 M)
pH scale is logarithmic:
1 unit = 10-fold
Dissociation of Weak Electrolytes:
Principle
O
H3C
+ H2O
O
Keq
H3C
OH
+
O-
H3O+
K a  Keq  [H 2O]
[H  ][CH 3COO - ]
Ka 
 1.74 10 5 M
[CH 3COOH]

[H ]  Ka 
[CH 3COOH ]
[CH 3COO  ]
• Weak electrolytes dissociate
only partially in water.
• Extent of dissociation is
determined by the acid
dissociation constant Ka.
• We can calculate the pH if the
Ka is known. But some
algebra is needed!
Dissociation of Weak Electrolytes:
Example
What is the final pH of a solution when 0.1 moles of
acetic acid is added to water to a final volume of 1L? The
Ka of acetic acid is 1.74x10-5
O
O
Ka
• We assume that the
+
H3C
H
H3C
+
only source of H+ is
OH
Othe weak acid
0.1 – x
x
x
Ka 
[x][ x]
 1.74 10 5 M
[0.1 - x]
x 2  1.74 10 6  1.74 10 5 x
x 2  1.74 10 5 x  1.74 10 6  0
x = 0.001310,
pH = 2.883
• To find the [H+], a
quadratic equation
must be solved
Dissociation of Weak Electrolytes:
Simplification
O
H3C
Ka
O
+ H+
H3C
O-
OH
0.1 – x
0.1
Ka 
x
x
x
x
[x][ x]
 1.74 10 5 M
[0.1]
x 2  1.74 10 6
• The equation can be
simplified if the amount of
dissociated species is much
less than the amount of
undissociated acid
• Approximation works for
sufficiently weak acids and
bases
• Check that x < [total acid]
x = 0.00132,
pH = 2.880
Nearly identical to the answer
determined by the quadratic
formula!
pKa measures acidity
• pKa = –log Ka (strong acid  large Ka  small pKa)
Buffers are mixtures of weak acids
and their anions (conjugate base)
• Buffers resist change in pH
• At pH = pKa, there is a 50:50 mixture of acid and
anion forms of the compound
• Buffering capacity of acid/anion system is greatest
at pH = pKa
• Buffering capacity is lost when the pH differs from
pKa by more than 1 pH unit
Acetic Acid-Acetate as a Buffer System
Weak acids have different pKas
Henderson–Hasselbalch Equation:
How to Find the pH of a weak acid during its titration
For a weak acid:
HA 
 H+ + A-
and, therefore
[HA]
[H ]  K a
[A - ]
+
[H  ][ A - ]
Ka 
[HA]
We want to know the pH, so solve for [H+]
[HA]
- log[H ]  -logK a  log
[A-]
-
[A ]
pH  pK a  log
[HA]
How to solve any weak acid/base
problem
1) Identify the Acid and the Base and the Ionization Reaction
2) If it is a titration:
(Eg: Find the pH at each of the following points in the titration of 25 mL of 0.3 M Acetic acid with 0.3 M NaOH. The Ka
value is 1.74x10-5: The initial pH, After adding 10 mL of 0.3 M NaOH, After adding 12.50 mL of 0.3 M NaOH,
After adding 25 mL of 0.3 M NaOH, After adding 26 mL of 0.3 M NaOH)
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Calculate how much of your starting material was reacted and
subtract that from the starting number of moles of acid or base
• Determine the new total volume and calculate the molarity of
the acid or base
• Use the Ka equation or the Henderson-Hasselbach equation to
determine the molarity of the remaining species in the beaker or
the pH
3) If it is not a titration (Eg: What is the concentration of acetate in 500 mL of a 0.05M Acetate buffer at pH 5.5?)
• Determine how many moles of acid or base you have
• Use the Ka equation or the Henderson-Hasselbach equation to
determine the molarity of the remaining species in the beaker or
the pH
Biological Buffer Systems
• Maintenance of intracellular pH is vital to all cells
– Enzyme-catalyzed reactions have optimal pH
– Solubility of polar molecules depends on H-bond donors and acceptors
– Equilibrium between CO2 gas and dissolved HCO3– depends on pH
• Buffer systems in vivo are mainly based on
– phosphate, concentration in millimolar range
– bicarbonate, important for blood plasma
– histidine, efficient buffer at neutral pH
• Buffer systems in vitro are often based on sulfonic acids of
cyclic amines
– HEPES
– PIPES
– CHES
HO
N
N
SO3Na
Chapter 2: Summary
In this chapter, we learned about:
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The nature of intermolecular forces
The properties and structure of liquid water
The behavior of weak acids and bases in water
The way water can participate in biochemical reactions