THEME:
Solution. Colligative properties
of biological liquids.
associate prof. Bekus I.R. prepared
PLAN
1. The main concepts of solutions
2. Types of solutions
3. Heat effect of a dissolution
4. Methods for expressing the
concentration of a solution
5. Vapor pressure and Raoult’s law
6. Colligative properties
7. Factors Affecting Solubility
Solution Composition
The solute and solvent can be any
combination of solid (s), liquid (l),
and gaseous (g) phases.
Dissolution: Two (or more) substances mix at the level of
individual atoms, molecules, or ions.
Solution: A homogeneous mixture (mixed at level of atoms
molecules or ions
Solvent: The major component
Solute: The minor component
GENERAL PROPERTIES OF SOLUTIONS
1. A solution is a homogeneous mixture of two or
more components.
2. It has variable composition.
3. The dissolved solute is molecular or ionic in
size.
4. A solution may be either colored or colorless
nut is generally transparent.
5. The solute remains uniformly distributed
throughout the solution and will not settle out
through time.
6. The solute can be separated from the solvent
by physical methods.
TYPES OF SOLUTION
1. Depending upon the
total components
present in the
solution:
 Binary solution (two
components)
 Ternary solution (three
components)
 Quaternary solution
(four
components)…..etc.
2. Depending upon the
ability of the dissolution
some quantity of the
solute in the solvent:
• Saturated
• Unsaturated solution
• Supersaturated
3. Depending upon the physical states of the solute
and solvent, the solution can be classified into the
following nine type:
Selected Acids and Bases
Acids
Bases
Strong
Hydrochloric, HCl
Hydrobromic, HBr
Hydroiodoic, HI
Nitric acid, HNO3
Sulfuric acid, H2SO4
Perchloric acid, HClO4
Strong
Sodium hydroxide, NaOH
Potassium hydroxide, KOH
Calcium hydroxide, Ca(OH)2
Strontium hydroxide, Sr(OH)2
Barium hydroxide, Ba(OH)2
Weak
Hydrofluoric, HF
Phosphoric acid, H3PO4
Acetic acid, CH3COOH
(or HC2H3O2)
Weak
Ammonia, NH3
 Gas solution. Gaseous solutions have the structure that
is typical of all gases. (Air, the gaseous solution with which
we come in closest contact, is composed primarily of N2
(78 % by volume), O2 (21 %), and Ar (1 %), with smaller
concentrations of CO2, H2O, Ne, He, and dozens of other
substances at very low levels).
 Liquid solutions have the internal structure that is typical
of pure liquids: closely spaced particles arranged with little
order. Unlike a pure liquid, however, a liquid solution is
composed of different particles. Much of this chapter is
devoted to the properties of liquid solutions, and special
emphasis is given to aqueous solutions, in which the
major component is water.
 Two kinds of solid solutions are common. The first, the
substitutional solid solution, exhibits a crystal lattice that
has structural regularity but in which there is a random
occupancy of the lattice points by different species.
Factors Affecting Solubility
1. Molecular Interactions
 Polar molecules, water soluble, hydrophilic
(water loving)
 (Vitamins B and C; water-soluble)
 Non-polar molecules, soluble in non-polar
molecules, hydrophobic (water fearing)
 (Vitamins A, D, K and E; fat-soluble)
Factors Affecting Solubility of Gases
1. Structure Effects
2. Pressure Effects
Henry's law
At a constant temperature, the
amount of a given gas that
dissolves in a given type and
volume of liquid is directly
proportional to the partial pressure
of that gas in equilibrium with
that liquid.
An equivalent way of stating the law
is that the solubility of a gas in a
liquid is directly proportional to
the partial pressure of the gas
above the liquid.
Henry's law
Where
p is the partial pressure of
the solute in the gas
above the solution
c is the concentration of the
solute
kH is a constant with the
dimensions of pressure
divided by concentration.
According to Henry's Law, the solubility of a
gas in a liquid
1) depends on the polarity of the liquid
2) depends on the liquid's density
3) remains the same at all temperatures
4) increases as the gas pressure above
the solution increases
5) decreases as the gas pressure above the
solution increases
An aqueous solution consists of at least two
components, the solvent (water) and the solute
(the stuff dissolved in the water).
Water is a chemical compound with the
chemical formula H2O.
A water molecule contains one oxygen and
two hydrogen atoms connected by covalent
bonds. Water is a liquid at standard ambient
temperature and pressure, but it often coexists on Earth with its solid state, ice, and
gaseous state (water vapor or steam).
Water also exists in a liquid crystal state near
hydrophilic surfaces.
Nonelectrolytes are substances such as
sucrose or ethyl alcohol, which do not
produce ions in aqueous solution.
Concentration units of a solution
The concentration of a solution may be
defined as the amount of solute
present in the solution.
1. Mass percentage (weight percentage):
The mass percentage of a component in a
given solution is the mass of the
component per 100 g of the solution.
mass percentage of the mass of component
component
=
X 100%
total mass of mixture
2. Mole fraction: It is the number of
moles of the solute dissolved per
litre of the solution.
The amount of a given component (in
moles) divided by the total
amount (in moles).
n
m
CM 

V MV
Molarity (Concentration of Solutions)= M
M=
Moles of Solute =
Liters of Solution
Moles
L
solute = material dissolved into the solvent
In air , Nitrogen is the solvent and oxygen, carbon dioxide, etc.
are the solutes.
In sea water , Water is the solvent, and salt, magnesium chloride, etc.
are the solutes.
In brass , Copper is the solvent (90%), and Zinc is the solute(10%)
LIKE EXAMPLE
Calculate the Molarity of a solution prepared by bubbling
3.68g of Gaseous ammonia into 75.7 ml of solution.
Solution:
Calculate the number of moles of ammonia:
1 mol NH3
3.68g NH3 X
= 0.216 mol NH3
17.03g
Change the volume of the solution into liters:
75.7 ml X
1L
1000 mL
= 0.0757 L
Finally, we divide the number of moles of solute by the volume
of the solution:
Molarity =
0.216 mol NH3
0.0757 L
= ____________ M NH3
Molarity
NaCl
Molarity
Example Problem 1
12.6 g of NaCl are dissolved in water making
344mL of solution. Calculate the molar
concentration.
moles solute
M=
L solution
 1molNaCl 
12.6 g NaCl 

58.44
gNaCl


=
 1L 
344 mL 
 solution
 1000mL 
= 0.627 M NaCl
Molarity
NaCl
Molarity
Example Problem 2
How many moles of NaCl are contained in 250.mL
of solution with a concentration of 1.25 M?
moles solute
M=
L solution
 1L 
250. mL 
 = 0.250 L solution
 1000mL 
Volume
x concentration
therefore the
solution contains
1.25 mol NaCl
1 L solution
=
moles solute
 1.25 mol NaCl 
0.250 L solution 
 = 0.313 mol NaCl
 1 L solution 
Molarity
NaCl
Molarity
Example Problem 3
What volume of solution will contain 15 g of NaCl
if the solution concentration is 0.75 M?
moles solute
M=
L solution
 1 mol NaCl 
15 g NaCl 
 = 0.257 mol
 58.44 g NaCl 
therefore the
solution contains
0.75 mol NaCl
1 L solution
moles solute ÷ concentration = volume solution
 1 L solution 
0.257 mol NaCl 
 = 0.34 L solution
 0.75 mol NaCl 
3. Molality
It is the number of moles of the solute dissolved per 1000 g (or 1 kg)
of the solvent. It’s denoted by m or Cm
Cm = (m) = Moles of solute/Weight of solvent in kg
or
Cm = (m) = Moles of solute * 1000/Weight of solvent in gram
The unit of Molality is m or mol/kg
n solute
m solute
Cm 

m solvent M solutem solvent
Molality
Calculate the molality of a solution consisting of 25 g of KCl
in 250.0 mL of pure water at 20oC?
First calculate the mass in kilograms of solvent using the
density of solvent:
250.0 mL of H2O (1 g/ 1 mL) = 250.0 g of H2O (1 kg / 1000 g)
= 0.2500 kg of H2O
Next calculate the moles of solute using the molar mass:
25 g KCl (1 mol / 54.5 g) = 0.46 moles of solute
Lastly calculate the molality:
m = n / kg = 0.46 mol / 0.2500 kg = 1.8 m (molal) solution
Molal (m)
Example Problem 1
If the cooling system in your
car has a capacity of 14 qts,
and you want the coolant to be protected from freezing
down to -25°F, the label says to combine 6 quarts of
antifreeze with 8 quarts of water. What is the molal
concentration of the antifreeze in the mixture?
mol solute
m=
Kg solvent
m=
antifreeze is ethylene glycol C2H6O2
1 qt antifreeze = 1053 grams
1 qt water = 946 grams
 1053 g C 2 H 6O 2   1mol C2 H 6O 2 
6 Qts 
  62.1 g C H O 
1
Qt
C
H
O
2 6 2 
2 6 2 

 946 g H 2O 
8 Qts 

1
Qt
H
O

2

 1 Kg 


 1000 g 
= 13 m
4. Normality:
It is the number of gram equivalents of the solute
dissolved per litre of the solution. It’s denoted by N or
CN
(N)= CN = Number of gram equivalents of solute/Volume
of solution in litres
or
(N) = CN = Number of gram equivalents of solute *1000 /
Volume of solution in ml
Number of gram equivalents of solute = Mass of solute /
Equivalent mass of solute
% Concentration
mass solute
% (w/w) =
mass solution
x 100
mass solute
% (w/v) =
volume solution
x 100
volume solute
% (v/v) =
volume solution
x 100
Mass and volume units must match.
(g & mL)
or
(Kg & L)
% Concentration
Example Problem 1
What is the concentration in %w/v of a solution containing 39.2 g
of potassium nitrate in 177 mL of solution?
39.2 g
mass solute
 100 = 22.1 % w/v
 100
% (w/v) =
volume solution
177 mL
Example Problem 2
What is the concentration in %v/v of a solution containing 3.2 L of
ethanol in 6.5 L of solution?
volume solute
 100
% (v/v) =
volume solution
3.2 L
 100
6.5L
= 49 % v/v
% Concentration
Example Problem
What volume of 1.85 %w/v solution is needed to
3
provide 5.7 g of solute?
1.85 g solute
% (w/v) =
100 mL solution
We know:
g solute and
We want to get:
g solute
mL solution
 100 mL solution 
5.7 g solute 

1.85
g
solute


g solute ÷ concentration
mL solution
= 310 mL Solution
=
volume solution
Colligative Properties
Colligative properties depend only on
the number of solute particles present,
not on the identity of the solute
particles.
Among colligative properties are
Vapor pressure lowering
Boiling point elevation
Freezing point depression
Osmotic pressure
 The vapor pressure necessary to achieve equilibrium with
the pure solvent is higher than that required with the
solution.
 Consequently, as the pure solvent seeks to reach
equilibrium by forming vapor, the solution seeks to reach
equilibrium by removing molecules from the vapor phase.
A net movement of solvent molecules from the pure
solvent to the solution results. The process continues until
no free solvent remains.
The extent of vapor pressure lowering depends on the
amount of solute.
 Raoult’s Law quantifies the Highlights
amount of vapor pressure – 1886 Raoult's law , the partial
lowering observed.
pressure of a solvent vapor in
equilibrium with a solution is
PA = XAPOA
proportional to the ratio of the
number of solvent molecules to
where
non-volatile solute molecules.
PA = partial pressure of the solvent vapor
above the solution (ie with the solute)
XA = mole fraction of the solvent
PoA = vapor pressure of the pure solvent
– allows molecular weights to
be determined, and provides
the explanation for freezing
point depression and boiling
point elevation.
Ideal solutions are those that obey Raoult’s Law.
Real solutions show approximately ideal behavior
when:
1)The solution concentration is low
2)The solute and solvent have similarly sized
molecules
3)The solute and solvent have similar types of
intermolecular forces.
Boiling Point Elevation and Freezing
Point Depression
Solute-solvent interactions also cause solutions to
have higher boiling points and lower freezing
points than the pure solvent.
 For example, the addition of salt to water causes
the water to freeze below its normal freezing point
(0°C) and to boil above its normal boiling point
(100°C).
 At the normal boiling point of the pure liquid, the
vapor pressure of the liquid, Po = 1 atm.
Freezing point depression
The freezing point of a solution is the temperature at which the
first crystals of pure solvent begin to form.
Osmotic Pressure
 To stop osmosis, the chemical potential of the solvent in the
more concentrated solution can be increased by forcing the
molecules closer together under an externally applied pressure.
 The pressure required to stop osmosis, known as osmotic
pressure, , is
 = (n/v)RT = MRT





where n is number of moles of solute,
V volume of solution,
M is the molarity of the solution
T is thermodynamic temperature
R is gas constant
Classification of Solutions According to Their Osmotic Pressure:




Hypertonic
Hypotonic
Isotonic
Isotonic: The solutions being compared have equal
concentration of solutes.
 Hypertonic: The solution with the higher concentration of
solutes.
 Hypotonic: The solution with the lower concentration of
solutes.
Osmosis in Blood Cells
 If the solute
concentration outside
the cell is greater than
that inside the cell, the
solution is hypertonic.
 Water will flow out of
the cell, and crenation
(shrinking) results.
The effect of hypertonic and hypotonic solutions on animal cells.
а) Hypertonic solutions cause cells to shrink (crenation) - plasmolysis;
b) Hypotonic solutions cause cell rupture - hemolysis;
c) Isotonic solutions cause no changes in cell volume.
Plasmolysis is the process in plant cells where the cytoplasm
pulls away from the cell wall due to the loss of water through
osmosis. This occurs in a hypertonic solution. The reverse
process, cytolysis, can occur if the cell is in a hypotonic
solution resulting in a lower external osmotic pressure and a
net flow of water into the cell.
Difference between osmosis and diffusion
van 't Hoff factor
The van 't Hoff factor is a measure of the effect of a solute upon
colligative properties such as osmotic pressure, relative
lowering in vapor pressure, elevation of boiling point and
freezing point depression. The van 't Hoff factor is the ratio
between the actual concentration of particles produced when
the substance is dissolved, and the concentration of a
substance as calculated from its mass.
For most non-electrolytes dissolved in water, the van' t
Hoff factor is essentially 1. For most ionic
compounds dissolved in water, the van 't Hoff factor
is equal to the number of discrete ions in a formula
unit of the substance.
Van't Hoff
The Person Behind the Science
J.H. van’t Hoff (1852-1901)
Highlights
– Discovery of the laws of chemical
dynamics and osmotic pressure in
solutions
– Mathematical laws that closely
resemble the laws describing the
behavior of gases.
– his work led to Arrhenius's theory
of electrolytic dissociation or
ionization
– Studies in molecular structure laid
the foundation of stereochemistry.
Moments in a Life
– 1901 awarded first Noble Prize in
Chemistry
van’t Hoff Factor (i)
moles of particles in solution
i
moles of solute dissolved
ΔT = − i m K
Oncotic pressure
Oncotic pressure, or colloid osmotic pressure, is a form of
osmotic pressure exerted by proteins in a blood vessel's plasma
(blood/liquid) that usually tends to pull water into the
circulatory system. It is the opposing force to hydrostatic
pressure.
Throughout the body, dissolved
compounds have an osmotic pressure.
Because large plasma proteins cannot
easily cross through the capillary
walls, their effect on the osmotic
pressure of the capillary interiors
will, to some extent, balance out the
tendency for fluid to leak out of the
capillaries.
Raoult's law
 The vapour pressure of an ideal solution is
directly dependent on the vapour pressure of
each chemical component and the mole
fraction of the component present in the
solution.
 Where:
pi: pressure of component i
xi: mole fraction in the solution
: vapor pressure of the pure substance
CRYOSCOPY and EBULIOSKOPY
Determination of molecular weight substance
freezing temperature decrease or increase the
boiling point of solutions called according
cryoscopy (cryoscopic method) or ebulioskopy
(ebulioskopic method).
These methods are used to establish the composition of
compounds to determine the degree of dissociation of
electrolytes, the study of the polymerization agents
and associations in solutions.
Ware chemical
dishes chemical
 A Petri dish (or Petri plate
or cell culture dish) is a
shallow glass or plastic
cylindrical lidded dish that
biologists use to culture
cells or small moss plants.
Beaker
 A beaker is a simple
container for stirring,
mixing and heating liquids
commonly used in many
laboratories. Beakers are
generally cylindrical in
shape, with a flat bottom.
Most also have a small
spout (or "beak") to aid
pouring as shown in the
picture. Beakers are
available in a wide range of
sizes, from one millilitre up
to several litres.
Laboratory flasks

There are several types of laboratory
flasks, all of which have different
functions within the laboratory.
Flasks, because of their use, can be
divided into:
1. Reaction flasks
2. Multiple neck flasks
3. Schlenk flask
4. Distillation flasks
5. Reagent flasks
6. Volumetric flask
Volumetric flask
 A volumetric flask
(measuring flask or
graduated flask) is a piece of
laboratory glassware, a type
of laboratory flask,
calibrated to contain a
precise volume at a
particular temperature.
Volumetric flasks are used
for precise dilutions and
preparation of standard
solutions.
Graduated cylinder
 A graduated cylinder,
measuring cylinder or
mixing cylinder is a
piece of laboratory
equipment used to
measure the volume of a
liquid. Graduated
cylinders are generally
more accurate and
precise than laboratory
flasks and beakers
Burette
 A burette is a device used in
analytical chemistry for the
dispensing of variable,
measured amounts of a
chemical solution. A
volumetric burette delivers
measured volumes of liquid.
Piston burettes are similar to
syringes, but with precision
bore and plunger. Piston
burettes may be manually
operated or may be
motorized.
Funnels (laboratory)
 Laboratory funnels are
funnels that have been made
for use in the chemical
laboratory. There are many
different kinds of funnels
that have been adapted for
these specialized
applications. Filter funnels,
thistle funnels (shaped like
thistle flowers), and
dropping funnels have
stopcocks which allow the
fluids to be added to a flask
slowly. For solids, a powder
funnel
Glass tubes
 Glass tubes are hollow
pieces of borosilicate or
flint glass used primarily
as laboratory glassware.
Glass tubing is
commercially available in
various thicknesses and
lengths. Glass tubing is
frequently attached to
rubber stoppers.
chemical dropper
 A pipette, pipet, pipettor
or chemical dropper is a
laboratory tool
commonly used in
chemistry, biology and
medicine to transport a
measured volume of
liquid, often as a media
dispenser.
Thank you for
attention