Topic 2: WATER (lectures 3-4)
OBJECTIVES:
1. Understand the fundamental physical properties of water in relation to hydrogen
bonding and how this impacts water’s cohesion, adhesion, surface tension, thermal
and solvent properties.
2. Know the physical properties of substances that render them hydrophilic or
hydrophobic.
3. Define what a mole is and be able to calculate how to make M, mM and/or M
solutions.
4. Understand the operational definitions of weak and strong acids and weak and
strong bases and how pH is used as a convenient way of specifying [H +] (actually
[H30+]). Be able to quantitatively compare the [H+]’s in various solutions.
5. Understand the concept of buffers and how buffers work.
Most organisms are 65-80% water. Water covers a large fraction of the earth. Water
has unique properties which make it useful as a medium for life and transport.
Fig. 3.1- charged portions of water molecule are attracted to oppositely charged
portions of other water molecules. Leads to unique physico-chemical properties.
A. Cohesion- tendency of a substance to stay together; permits water to be transported
in bulk (blood, sap in plants etc)
B. Adhesion- tendency of a substance to attach to another substance; water adheres
to surfaces (respiratory epithelia)
C. Surface tension- ability to resist deformation at a surface
D. Thermal properties
(a) heat- measure of the total quantity of kinetic (movement) energy in a body of
matter (unit is calorie; amount of heat needed to raise temp. of 1 g water 1 C)
(b) temperature- measures the intensity (or concentration) of heat (unit is oC)
( c ) specific heat (heat capacity)- amount of heat that a substance must lose or
absorb for its temperature to change; water has a very high specific heat vs. air
which is much lower. Bodies lose heat to water much more rapidly.
(d) evaporative cooling- when water is converted from liquid to vapor the latent
heat of vaporization must be invested; this cools the surface where
evaporation is taking place.
(e) ice formation- for water to be converted from liquid to solid (ice) the latent
heat of fusion must be removed; upon ice formation there is an expansion of
volume so that the density (g/ml) is less than water, so ice floats.
Fig. 3.5E. Solvent properties of water
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Solution- a homogeneous mixture of two substances
Solvent- the dissolving agent of a solution
Solute- the material that is dissolved
Water has very good solvent properties for hydrophilic (water-loving) substances
Hydrophilic- molecules that are polar or ionic; charges interact with water
Hydrophobic- molecules that are apolar, lack charges and thus cannot interact with
water
Fig. 3.7- dissolving NaCl in water; Na+ and Cl- interact with polar water molecules
Fig. 3.8- surfaces of proteins are covered with positive and negative charges which
interact with water (proteins are polar [have separation of charges]; form hydrogen
bonds with water); proteins go into solution when they form polar interactions with
water.
Concentration: the amount of solute dissolved in a solvent
Mole (mol): the molecular weight of a substance scaled to grams (g); This is often
referred to as the GRAM FORMULA WEIGHT; thus 58.44
g represents one mole of NaCl
1 mole of NaCl is 58.44 g
1 millimole (mmole) is 1 x 10-3 mole; therefore 1 mmole NaCl is 58.44/1,000
= 0.05844 g
1 micromole (mole) is 1 x 10-6 mole; therefore 1 mole is 58.44/1,000,000 =
0.00005844 g
Molar (M): an expression of concentration; the number of moles of solute per liter of
solution; to make a 1 M solution of NaCl, add 58.44 g to water and bring up
to 1 liter (this makes 1 mole in a liter). When we talk about solutions in organisms
we usually deal with concentrations much lower than 1 M:
milliMolar (mM) = 10-3 M; 1 mmole in a liter = 0. 05844 g in a liter
microMolar (M)= 10-6 M; 1 mole in a liter = 0.00005844 g in a liter
(50 mM solution; 50 mmoles in one liter; 50 mM NaCl solution; 0.05 x 58.44g =
2.922g when brought to a liter yields 50 mM ( or 0.05 M) solution)
Solubility (expressed as g/liter): the amount of solute that a solvent can dissolve before
the solute no longer goes into solution; because of its polar nature, water is a
superb solvent. Solubility is a function of the nature of the solute, temperature
and other factors. (typically, increases in temperature increase solubility of
solutes [except for gases])
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Dissociation of water and acids/bases (dissociation: breakdown into ionic species)
Water occasionally reacts with itself:
2 H2O  H3O+ (hydronium ion) + OH- (hydroxide ion)
Both ions are very reactive; in living systems major effort is made to keep the
concentrations of these ions at very low levels. Normally, biologists refer to the
hydronium ion as the hydrogen ion (H+) or proton since only a proton remains when the
electron is pulled from the hydrogen atom. We will use H+ (proton) throughout this
course.
Acids & bases.
Acid: any substance that will dissolve (dissociate) in water yielding H+ and an anion (or
an acid is a substance that can donate protons); increases H+ concentration.
HCl + H2O  H+ + ClBase: any substance that can accept protons; decreases H+ concentrations.
H+ + NH3 +  NH4+
NaOH + H2O  Na+ + OH- OH- + H+  H2O
Strong acids (bases)- completely dissociate in water to H+ and an anion; HCl example
above.
Weak acids (bases)- do not completely dissociate in water; good example of acetic
acid, a solution of acetic acid contains the following:
acetic acid (free acid) + acetate- (anion, also known as conjugate base) + H+
( note: the relative concentrations of free acid and conjugate base are a function of [H+])
Many biologically important molecules are weak acids and bases.
The pH scale.
Note: in any solution [H+] x [OH-] = 10-14 M2 (this is known as the ion product of water)
at 25 C, both [H+] and [OH-] are equal to 1 x 10-7 M; the solution is said to be neutral;
if acid is added the relative concentrations change, increasing the [H +] and decreasing
the [OH].
In biological systems the H+ concentration of falls in the 10-6 though 10-8 M range.
These are extremely small numbers, yet very small changes in concentration ( say
going from 1 x 10-7 to 1.1 x 10-7 M) may cause profound changes in biological
processes. A more convenient way stating the H+ concentration is by the pH scale.
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(note: log of a number < 1.0 is negative!!!!!)
pH = - log [H+]
-log 1 x 10-7 [H+] = 7
-log 1 x 10-8 [H+] = 8
-log 1 x 10-6 [H+] = 6
The lower the pH, the higher the [H+]; the higher the pH the lower the [H+]. A very acidic
solution has a low pH. Each pH unit represents a 10-fold difference in [H+] and [OH-].
pH 2 solution (1 x 10-2 M) has 105 (100,000) times higher proton concentration than a
pH 7 (1 X 10-7 M) solution (and so on)
Fig. 3.9- concrete examples of pH values
Buffers: buffers are substances which tend to minimize changes in [H+] and [OH-].
These substances gain or lose protons depending on pH and their chemical properties
(they are always weak acids or bases). These are very important in living organisms.
For instance, in the blood of mammals there is a buffering system based on carbon
dioxide (CO2) produced by metabolism. There is a mixture of H2CO3 (actually CO2) ,
HCO3- (bicarbonate) and H+ which are related to each other by the following reaction:
H2CO3  HCO3- + H+
When acid is added to the solution, HCO3- + H+ combine to H2CO3 ; this keeps pH
relatively constant.
When base is added to the solution, H2CO3 dissociates into HCO3- + H+; H+ reacts with
the base to keep pH relatively constant.
A number of other buffers are present in cells.
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