1. Physical Pharmacy Solutions Khalid T Maaroof
MSc. Pharmaceutical sciences
School of pharmacy – Pharmaceutics department
Online access: bit.ly/physicalpharmacy
10/31/2015

2. Two types: Heterogeneous Homogeneous Mixtures
a combination of two or more substances that do not combine chemically, and can be separated by physical means
Two types:
Heterogeneous
Homogeneous

3. Heterogeneous Mixture
“Hetero” means different
consists of visibly different substances or phases (solid, liquid, gas)
a suspension is a special type of heterogeneous mixture of larger particles that eventually settle
Example:

4. Homogeneous Mixture “Homo” means the same
has the same uniform appearance and composition throughout; maintain one phase (solid, liquid, gas)
Commonly referred to as solutions
Example:
Notice the
uniform
appearance
Salt Water

5. Solution
a mixture of two or more substances that is identical throughout
can be physically separated
composed of solutes and solvents
Salt water is considered a solution. How
can it be
physically separated?
the substance in the smallest
amount and the one that dissolves in the solvent
the substance in the larger
amount that dissolves the solute
Iced Tea Mix
(solute)
Water
(solvent)
Iced Tea
(solution)

6. Solutions

7. Types of solutes Dissociation in to ions in solution Eg: NaCl
Electrolytes (Conductive):
Dissociation in to ions in
solution
Eg: NaCl
Nonelectrolytes (no conductivity):
no dissociation
Eg: suger
Na+
Cl-

8. Concentration
the amount of solute dissolved in a solvent at a given temperature
described as dilute if it has
a low concentration of
solute
described as saturated if it
has a high concentration of
described as supersaturated if
contains more dissolved solute
than normally possible

9. Concentration expressions for solutions
Molarity?
Normality?
Molality?
Mole fraction?
Mole percent?
Percent
Percent by weight % w/w
Percent by volume %v/v
Percent weight in volume % w/v

10. Concentration expressed as percentage
Percent weight-in-weight (w/w) is the grams of solute in 100 grams of the solution.
Percent weight-in-volume (w/v) is the grams of solute in 100ml of the solution.
Percent volume-in-volume (v/v) is the milliliters of solute in 100ml of the solution.
A 20 % w/w solution contains 20g solute how many grams the solvent is?

11. Molarity, Normality, and Molality
Molarity and normality both depend on the volume of the solvent, so their values are affected by change of volume caused by factors such as change in temperature.
Molality doesn’t has this disadvantage.

12. Mole fraction
In a solution containing 0.01 mole of solute and mole of solvent, the mole fraction of the solute is 0.2 and for solvent it is 0.8
Mole percent = mole fraction X 100

13. An aqueous solution of ferrous sulfate was prepared by adding 41
An aqueous solution of ferrous sulfate was prepared by adding g of FeSO4 to enough water to make mL of solution. The density of the solution is and the molecular weight of FeSO4 is
Calculate
(a) molarity
(b) molality
(c) mole fraction of FeSO4, mole fraction of water, and the mole percent of the two constituents
(d) % w/w of FeSO4.

14. Solutes Change Solvents
The amount of solute in a solution determines how much the physical properties of the solvent are changed
Examples:
Lowering the Freezing Point
Raising the Boiling Point
The freezing point of a liquid solvent
decreases when a solute is dissolved in it.
Ex. Pure water freezes at (00C), but when salt is
dissolved in it, the freezing point is lowered.
This is why people use salt to melt ice.
The boiling point of a solution is higher
than the boiling point of the solvent.
Therefore, a solution can remain a liquid at
a higher temperature than its pure solvent.
Ex. The boiling point of pure water is (1000C),
but when salt is dissolved in it, the boiling
point is higher. This is why it takes salt water
longer to boil than fresh water.
Those physical properties of solution that are changed by the amount of solute are called colligative properties

15. Colligative properties

16. Colligative properties
Colligative properties of solutions are those that affected (changed) by the presence of solute and depend solely on the number (amount of solute in the solutions) rather than nature of constituents.
Examples of colligative properties are:
Vapor pressure
Boiling point
Freezing point
Osmotic pressure

17. Colligative vs Non-colligative
Compare 1.0 M aqueous sugar solution to a 0.5 M solution of salt (NaCl) in water.
both solutions have the same number of dissolved particles
any difference in the properties of those two solutions is due to a non-colligative property.
Both have the same freezing point, boiling point, vapor pressure, and osmotic pressure
Compare 1.0 M aqueous sugar solution to a 0.5 M solution of table salt (NaCl) in water.
Despite the concentrations, both solutions have precisely the same number of dissolved particles because each sodium chloride unit creates two particles upon dissolution - a sodium ion, Na+, and a chloride ion, Cl-.
Therefore, any difference in the properties of those two solutions is due to a non-colligative property.
Both solutions have the same freezing point, boiling point, vapor pressure, and osmotic pressure because those colligative properties of a solution only depend on the number of dissolved particles.

18. Non-Colligative Properties
Sugar solution is sweet and salt solution is salty.
Therefore, the taste of the solution is not a colligative property.
Another non-colligative property is the color of a solution.
Other non-colligative properties include viscosity, surface tension, and solubility.

19. Pure solvent > solutions
vapor pressure
Vapor pressure:
Pure solvent > solutions

20. Vapor pressure of solvent in solution containing non-volatile solute is always lower than vapor pressure of pure solvent at same T
At equilibrium rate of vaporization = rate of condensation
Solute particles occupy volume reducing rate of evaporation. (They limit the number of solvent molecules at the surface)
The rate of evaporation decreases and so the vapor pressure above the solution must decrease to recover the equilibrium

21. Boiling point
Boiling point elevation is a colligative property related to vapor pressure lowering.
The boiling point is defined as the temperature at which the vapor pressure of a liquid equals the atmospheric pressure.
Due to vapor pressure lowering, a solution will require a higher temperature to reach its boiling point than the pure solvent.

22. Freezing Point
Every liquid has a freezing point - the temperature at which a liquid undergoes a phase change from liquid to solid.
When solutes are added to a liquid, forming a solution, the solute molecules disrupt the formation of crystals of the solvent.
That disruption in the freezing process results in a depression of the freezing point for the solution relative to the pure solvent.

23. How this is done?

24. 1. salting roads in winter FP BP water 0oC (NFP) 100oC (NBP)
water + a little salt
–11oC
103oC
water + more salt
–18oC
105oC
2. antifreeze (AF) /coolant FP BP
water
0oC (NFP)
100oC (NBP)
water + a little AF
–10oC
110oC
50% water + 50% AF
–35oC
130oC
Why do you think some towns use calcium chloride on roads in the winter versus sodium chloride?
CaCl2 yields three ions while NaCl yields only two ions. Calcium chloride will work in colder weather. Calcium chloride will work better than sodium chloride. Secondly, calcium chloride doesn’t kill grass like sodium chloride.

25. What happens to the triple point?

26. Which is more effective for lowering the freezing point of water?
NaCl or CaCl or Glucose
Ionic solutes produce two or more ion particles in solution.
They affect the colligative properties proportionately more than molecular solutes (that do not ionize).
The effect is proportional to the number of particles of the solute in the solution.

27. Osmotic Pressure
When a solution is separated from a volume of pure solvent by a semi-permeable membrane that allows only the passage of solvent molecules, the height of the solution begins to rise.
The value of the height difference between the two compartments reflects a property called the osmotic pressure of a solution.
If you add more solvent to a solution, the two mix together to form a more dilute solution.
The same forces allowing that mixing serve to force solvent molecules from the pure solvent compartment across the membrane into the solution compartment causing the change in volume.
The amount of osmotic pressure is directly related to the concentration of the solute.
That is because more concentrated solutions have greater potentials for dilution.

28. Osmotic Pressure and Cells
In the figure, red blood cells are placed into saline solutions.
In which case (hypertonic, isotonic, or hypotonic) does the concentration of the saline solution match that of the blood cells?
In which case is the saline solution more concentrated than the blood cells?
Crenation
Hemolysis

29. Ideal solutions Two types of molecules are randomly distributed
Typically, molecules are similar in size and shape
Intermolecular forces in pure liquids & mixture are similar
Examples: benzene & toluene, hexane and heptane
Ideality in solutions means complete uniformity of attractive forces

30. Ideal Solutions pA = pA◦ XA pB = pB◦ XB P solution = pA + pB
Raoult's Law States that, in an ideal solution, the partial vapor pressure of each volatile constituent is equal to the vapor pressure of the pure constituent multiplied by its mole fraction in the solution. Thus, for two constituents A and B,
pA = pA◦ XA
pB = pB◦ XB
P solution = pA + pB

31. Real Solutions
Cohesion and Attraction ??
Ideality in solutions presupposes complete uniformity of attractive forces.
Many examples of solution pairs are known, however, in which the “cohesive” attraction of A for A exceeds the “adhesive” attraction existing between A and B.
Similarly, the attractive forces between A and B may be greater than those between A and A or B and B.
Such mixtures are real or nonideal; that is, they do not adhere to Raoult’s law
Two types of deviation from Raoult’s law are recognized, negative deviation and positive deviation.

32. Negative deviation Positive deviation
When the “adhesive” attractions between molecules of different species exceed the “cohesive” attractions between like molecules, the vapor pressure of the solution is less than that expected from Raoult’s ideal solution law, and negative deviation occurs.
Adhesion > Cohesion
Positive deviation
When the “adhesive” attractions between molecules of different species are weaker than “cohesive” attractions between like molecules, the vapor pressure of the solution is more than that expected from Raoult’s ideal solution law, and positive deviation occurs.
Adhesion < Cohesion

33. Questions !
10/31/2015