WATER and INTERMOLECULAR FORCES
Water has many roles in living systems:
 solvent
 reactant
 structural component
 thermal regulation
Water is very polar. electronegativity – ability of atoms of an element to attract electrons in a covalent bond (fluorine is most electronegative)
For any molecule to be polar (has a dipole moment) it must meet both requirements:
 contains one or more polar bonds (atoms of different electronegativity are bonded to each other
 the geometry must not cause bond polarities to cancel

Examples: methane carbon dioxide water nonpolar nonpolar polar no polar bonds bond dipoles cancel by geometry
Polarity is an important factor in determining how and how strongly molecules attract each other.
Weak Interactions
 weaker than covalent bonds
 exist between different molecules (intermolecular forces) or
 exist between “distant” parts of the same molecule

There are 4 main types:
 dipole-dipole attractions
- hydrogen bonding is a particularly strong type
 ionic interactions (charge-charge interactions)
- including ion-dipole interactions
 van der Waals forces
 hydrophobic interactions
1.
Dipole-dipole attractions – electrostatic attractions between polar molecules (molecules with
dipoles)

Hydrogen bonds are a particularly strong type of dipoledipole interaction found when H is covalently bonded to
O, N, or F.
Water molecules attract each other by hydrogen bonds.
These attractions give ice its open structure and give water its unusually high heat of vaporization and heat capacity.
Water-water
Water-ammonia
Water-formaldehyde

2.
Ionic interactions (charge-charge interactions) – electrostatic attraction between charged particles (full or partial charges)
Types:
 ion-ion
 ion-dipole
Ion-ion - Ions with full charges attract each other.
Example: carboxylate ion and ammonium ion (ion pairing)
Ion-dipole
Example: Solvation (surrounding of a solute by the solvent)
Hydration (solvation with water as the solvent)
Dissolving of NaCl (note the gain in entropy)
Na + Cl -

3.
London forces (van der Waals forces) – weak electrostatic interaction between all types of atoms
2 ways to think about them:
Way 1: Repulsion and attraction between protons and electrons, and minimization of the potential energy. repulsion
Energy
0 attraction

Way 2: Transient (non-permanent) dipoles caused by the constant motion of electrons.
4.
Hydrophobic Interactions – aggregation of non-polar species or parts driven by entropy increase in surrounding solvent hydrophilic (“water loving”) – polar or ionic species able to attract water hydrophobic (“water hating”) – nonpolar, unable to favorable attract water
Recall: "Like dissolves like"
Nonpolar gases dissolve poorly in water
O
2
N
2
CO
2
O
2
- has soluble protein carriers (hemoglobin & myoglobin)
CO
2
- dissolves as bicarbonate ion (HCO
3
)

Polar gases dissolve easily in water: NH
3
H
2
S amphipathic – adjective for molecules that contain both polar and nonpolar parts (e.g. soap) a soap – sodium oleate (salt of a fatty acid) a detergent – sodium dodecyl sulfate (SDS)
What happens to the soap anion when dissolved in water?
Carboxylate end - easily solvated
Non-polar tail (R group) - "cage" of ordered water molecules surround it (no hydrogen bonding possible between water and non-polar tail)
Clathrate - highly ordered water molecules surrounding non-polar molecule(s)
Water loses entropy when it's in a clathrate.
Systems seek to minimize this loss.

Non-polar - aggregate to minimize clathrate formation.
The forces causing the aggregation are called hydrophobic interactions.
Some structures whose formation is caused by hydrophobic interactions when amphipathic molecules are in contact with water:
 micelle
 monolayers at air-water interface
 bilayers
 liposomes
Weak interactions - weak individually, strong with many
Maintain structures (maintains function)
Examples: membranes
Proteins
Enzyme-substrate interactions
Water is often very important in maintaining structure.
Inside proteins
Associated with DNA and RNA
Etc.
Can detect by X-ray crystallography, NMR, etc.

Colligative Properties
Properties of a solution that depend on the solute molecules dissolved.
of
 Vapor pressure (lowering)
 Boiling point (elevation)
 Freezing point (melting point) (depression)
 Osmotic pressure (raising)
Osmosis - 2 solutions separated by a semi-permeable membrane will have solvent (water) move from low concentration of solute to high concentration of solute
Osmolarity =
Where
And
is the van't Hoff factor ( i = 1 for non-electrolytes)
= solute's molar concentration
Isotonic - refers to solutions of equal osmolarity
Often compare to osmolarity of the cytosol of cells.
Hypertonic solution - has higher osmolarity than cytosol
Hypotonic solution - has lower osmolarity than cytosol
Buffers used in laboratories often contain about 0.2 M sucrose to protect cells from osmotic lysis.

Acid-Base Properties
Water has the ability to self-ionize (auto-ionize).
H
2
O + H
2
O H
3
O + + OH -
hydronium hydroxide ampholytic or amphoteric – adjective which describes any molecule that can act as either an acid or a base
Lowry-Bronsted definitions acid – proton (H + ) donor H
3
O +  H + base – protons (H + ) acceptor
Equilibrium expression (all aqueous solutions at 25 o C; 1 atm)
K

[ H
3
O

][ OH

]

1 x 10

14 w
Recall pK w
= -log
w
= _________
For pure water [H
3
O + ] = [OH ] = 1 x 10 -7 M pH = -log[H + ] = -log[H
3
O + ] = _________ pH + pOH = __________
The pH scale is a logarithmic scale. This means a change in a factor of  1 in pH changes [H
3
O + ] by a factor of 10.

pH < 7 acidic [H
3
O + ] > [OH ] pH = 7 neutral [H
3
O + ] = [OH ] pH > 7 basic [H
3
O + ] < [OH ]
Acids
Strong acids – ionize completely (100%) in water
HCl, HNO
3
, H
2
SO
4
HCl + H
2
O (
) H
3
O + (
) + Cl (
)
Weak acids - ionize incompletely (<100%) in water
all organic acids are weak acids
HA + H
2
O (
) H
3
O + (
) + A (
) weak conjugate acid base
Equilibrium expression
K a

[ H
3
O

][
[ HA ]
Recall pK a
= -log K a
Example: Acetic acid
A

]

Values for K a
and/or pK a
can be found on tables and used to compare relative strengths of weak acids.
The larger the K a
the _______________________ the weak acid.
The larger the pK a
the ______________________ the weak acid.
Bases
Strong bases – ionize completely (100%) in water
NaOH, other soluble metal hydroxides
NaOH (
) Na + (
) + OH (
)
Weak bases - ionize incompletely (<100%) in water
all organic bases are weak bases
B + H
2
O (
) HB + (
weak conjugate
) + OH (
)
base acid
Equilibrium expression
K b
[ HB

][ OH


]
[ B ]
Recall pK b
= -log K b pK a
+ pK b
= __________

Example: Ammonia
Values for K b
and/or pK b
can be found on tables and used to compare relative strengths of weak bases.
The larger the K b
the _______________________ the weak base.
The larger the pK b
the ______________________ the weak base.
Le Chatelier’s Principle – When a system at equilibrium is disturbed, the system adjusts to a new equilibrium in a way that partially counteracts the disturbance.
Acetic acid dissociation:
Which way will this equilibrium for acetic acid shift when: pH is lowered (e.g. adding HCl) _________________ add sodium acetate _________________ increase concentration of acetic acid______________ dilute with water _______________

Neutralization reaction: acid + base  a salt (ionic compound) + water (usually)
Recall monoprotic acid – has only 1 H + to donate (HCl, HAc) polyprotic acid - has 2 or more H + ito donate (H
3
PO
4
)
Titration curves:
Strong acid titrated with a strong base: HCl with NaOH
Weak acid titrated with a strong base: HAc with NaOH
Weak base titrated with a strong acid: Ammonia with HCl
Important parts of the curves:
 Initial pH – determined by concentration and pK a
or pK b
 Buffering region – [weak acid]  [conjugate base]
- here the curve is relatively flat, indicating that the
- pH changes little when extra acid or base is added
 Half-titration point – midpoint of the buffering region
[weak acid] = [conjugate base]
pH = pK a

 Equivalence point (stoichiometric point)
- complete neutralization; stoichiometric amounts of
acid and base are present (watch with polyprotic)
occurs at inflection point of the curve (steepest slope)
occurs at pH = 7 only for strong acid – strong base titration
occurs at pH > 7 for weak acid – strong base titration*
occurs at pH < 7 for weak base – strong acid titration
pH depends on the cations and anions produced when titrating a weak acid with a weak base (or vice versa)
* Most important to know at this point.
Buffers – Solutions which resist pH changes upon addition of acid or base
This is very important in maintaining homeostasis.
 even relatively small changes in pH can have dramatic effects
 alteration of 3-d shape of proteins can cause loss of function; shows importance of structure  function
 ideal pH is called the biological activity
; pH of greatest

Recall – Henderson-Hasselbach equation (this is just a rearrangement of the expression for K a
)
Derivation

How to make a Buffer
1. For pH less than 7
a. Mix a weak acid with a salt of the anion of the
weak acid
Example: acetic acid:sodium acetate b. Make a solution of the salt of the anion of a weak
acid and adjust to desired pH with HCl
Example: sodium acetate; HCl
2. For pH greater than 7 a. Mix a weak base with a salt of the cation of the
weak base
Example: ammonia:ammonium chloride b. Make a solution of the salt of the cation of a weak
base and adjust to desired pH with NaOH
Example: ammonium chloride; NaOH

Other properties of good buffers:
 pK a
(or pK b
) of weak acid (or base) should be
+1 from pH desired
 concentrations of components relatively large compared to expected possible added acid or base
 concentrations of conjugate pairs
 approximately equal
Blood Buffering system – main one  ______________________ system

__________________ occurs when blood pH is too low
Treatment for mild cases  _______________________
Note shifts in equilibria.
___________________ occurs when blood pH is too high
Treatment for mild cases  _______________________
Note shifts in equilibria.
Cytoplasm Buffering system: phosphate
H
3
PO
4
(polyprotic) pK a1
= 2.14 pK a1
= 6.86 pK a1
= 12.4 pH of cytoplasm is near 7
Which species are most important?

Water in reactions
Product - in condensation reactions
ADP + P i
(inorganic phosphate)  ATP + H
2
O
Reactant - in hydrolysis reactions
Pyrophosphate + H
2
O  2 Phosphate + H
2
O
Oxidation-Reduction Reactions
C
6
H
12
O
6
+ 6 O
2
 6 CO
2
+ 6 H
2
O
CO
2
+ H
2
O ⇋ HCO
3
+ H +
Suggested problems for chapter 2:
2, 4, 5, 7, 8, 10, 11, 12, 13, 15