1. THEME: Solution. Colligative properties of biological liquids.
associate prof. Bekus I.R. prepared

2. 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

3. Solution Composition Dissolution:
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 3

5. 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.

6. TYPES OF SOLUTION Saturated Unsaturated solution Supersaturated
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

7. 3. Depending upon the physical states of the solute and solvent, the solution can be classified into the following nine type:

8. Selected Acids and Bases
Acids Bases
Strong Strong
Hydrochloric, HCl Sodium hydroxide, NaOH
Hydrobromic, HBr Potassium hydroxide, KOH
Hydroiodoic, HI Calcium hydroxide, Ca(OH)2
Nitric acid, HNO Strontium hydroxide, Sr(OH)2
Sulfuric acid, H2SO Barium hydroxide, Ba(OH)2
Perchloric acid, HClO4
Weak Weak
Hydrofluoric, HF Ammonia, NH3
Phosphoric acid, H3PO4
Acetic acid, CH3COOH
(or HC2H3O2)

9. 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.

14. Factors Affecting Solubility
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) 14

15. Factors Affecting Solubility of Gases
Structure Effects
Pressure Effects 15

16. 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.

17. 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.

18. According to Henry's Law, the solubility of a gas in a liquid 1) 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 18

19. An aqueous solution consists of at least two components, the solvent (water) and the solute (the stuff dissolved in the water).

20. Water is a chemical compound with the chemical formula H2O
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 co-exists 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.

23. Nonelectrolytes are substances such as sucrose or ethyl alcohol, which do not produce ions in aqueous solution.

24. 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 component =
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).
X 100%
mass of component
total mass of mixture

25. Molarity (Concentration of Solutions)= M
Moles of Solute Moles
Liters of Solution L
M = =
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.
In brass , Copper is the solvent (90%), and Zinc is the solute(10%)

26. Calculate the Molarity of a solution prepared by bubbling
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
17.03g
3.68g NH3
X = mol NH3
Change the volume of the solution into liters:
1 L
1000 mL
75.7 ml X = L
Finally, we divide the number of moles of solute by the volume
of the solution:
0.216 mol NH3 L
Molarity = = ____________ M NH3

27. Molarity Example Problem 1
NaCl
12.6 g of NaCl are dissolved in water making 344mL of solution. Calculate the molar concentration.

28. Molarity Example Problem 2
NaCl
How many moles of NaCl are contained in 250.mL of solution with a concentration of 1.25 M?
Volume x concentration = moles solute

29. Molarity Example Problem 3
NaCl
What volume of solution will contain 15 g of NaCl if the solution concentration is 0.75 M?
moles solute ÷ concentration = volume solution

30. 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

31. Molality
Calculate the molality of a solution consisting of 25 g of KCl in 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) = g of H2O (1 kg / 1000 g) = 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 / kg = 1.8 m (molal) solution

32. antifreeze is ethylene glycol C2H6O2
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?
antifreeze is ethylene glycol C2H6O2
1 qt antifreeze = 1053 grams
1 qt water = 946 grams
m =
= 13 m

33. 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
= CN = Number of gram equivalents of solute/Volume of solution in litres or
= 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

34. Mass and volume units must match.
% Concentration
% (w/w) =
% (w/v) =
% (v/v) =
Mass and volume units must match.
(g & mL) or (Kg & L)

35. % Concentration Example Problem 1 Example Problem 2
(Solid in a Liquid)
What is the concentration in %w/v of a solution containing 39.2 g of potassium nitrate in 177 mL of solution?
% (w/v) =
Example Problem 2
(Liquid in a Liquid)
What is the concentration in %v/v of a solution containing 3.2 L of ethanol in 6.5 L of solution?
% (v/v) =

36. What volume of 1.85 %w/v solution is needed to
% Concentration
Example Problem 3
What volume of 1.85 %w/v solution is needed to
provide 5.7 g of solute?
% (w/v) =
We know:
We want to get:
g solute ÷ concentration = volume solution

38. 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

39. 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.

41. The extent of vapor pressure lowering depends on the amount of solute.
Raoult’s Law quantifies the amount of vapor pressure lowering observed.
PA = XAPOA
where
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
Highlights
1886 Raoult's law , the partial pressure of a solvent vapor in equilibrium with a solution is proportional to the ratio of the number of solvent molecules to non-volatile solute molecules.
allows molecular weights to be determined, and provides the explanation for freezing point depression and boiling point elevation.

42. 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.

43. 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.

45. Freezing point depression
The freezing point of a solution is the temperature at which the first crystals of pure solvent begin to form.

47. Osmotic Pressure  = (n/v)RT = MRT
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

48. 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.

49. 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.

50. 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.

52. Difference between osmosis and diffusion

53. 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.

54. Van't Hoff

55. ΔT = − i m K J.H. van’t Hoff (1852-1901) van’t Hoff Factor (i)
The Person Behind the Science
J.H. van’t Hoff ( )
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)
ΔT = − i m K 55

56. 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.

57. 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

58. 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).

59. 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.

60. 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.

61. 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.

62. 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:
Reaction flasks
Multiple neck flasks
Schlenk flask
Distillation flasks
Reagent flasks
Volumetric flask

63. 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.

64. 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

65. 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.

66. 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

67. 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.

68. 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.

69. Thank you for attention