David A. Katz Department of Chemistry Pima Community College

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1 Solutions David A. Katz Department of Chemistry Pima Community College

2 A solution is a HOMOGENEOUS mixture of 2 or more substances in a single phase. One constituent t is usually regarded as the SOLVENT (usually water) and the others as SOLUTES.

3 Solutions In a solution, the solute is dispersed uniformly throughout the solvent.

4 Solutions: Gases mixed with gases

5 Solutions: Gas mixed with liquid

6 Solutions: Liquid mixed with liquid

7 Solutions: Solid mixed with liquid

8 Solutions: Gas mixed with solid Photomicrographs of Hydrogen on palladium

9 Solutions: Liquid mixed with solid (a mercury amalgam with gold)

10 Solutions: Solid mixed with solid (alloys) Brass: a substitution alloy Carbon steel: an interstitial alloy

11 Intermolecular Forces Why does a substance dissolve? The intermolecular forces between solute and solvent particles must be strong enough to compete with those between solute particles and those between solvent particles.

12 Intermolecular Forces: Why does a substance dissolve? The intermolecular forces between solute and solvent particles must tbe strong enough to compete with those between solute particles and those between solvent particles. (continued on next slide) stronger weaker

13 Intermolecular Forces: Why does a substance dissolve? (continued) stronger weaker

14 How Does a Solution Form? In this example, we have an ionic solid, NaCl, and a polar solvent, H 2 O. The solution forms because the solvent pulls solute particles apart and surrounds, or solvates, them. In water this is called hydration. Solute (NaCl) in The solute is Hydrated ions in water dissolving solution

15 How Does a Solution Form If an ionic salt is soluble in water, it is because the iondipole interactions are strong enough to overcome the lattice energy of the salt crystal. The ions are hydrated that prevents the ions from reforming the crystal lattice under normal conditions.

16 Solutions Just because a substance disappears when it comes in contact with a solvent, it doesn t mean the substance dissolved. Dissolution is a physical change you can get back the original solute by evaporating the solvent. If you can t, the substance didn t dissolve, it reacted.

17 Water as a Solvent How water dissolves molecular compounds: When the nonpolar part of an organic molecule is considerably larger than the polar part, the molecule no longer dissolves in water. For example ethanol, CH 3 CH 2 OH is soluble in water but butanol CH 3 CH 2 CH 2 CH 2 OH is not

18 Factors Affecting Solubility Chemists use the axiom like dissolves like : Polar and ionic substances tend to dissolve in polar solvents. Nonpolar substances tend to dissolve in nonpolar solvents.

19 Factors Affecting Solubility Solubility in water decreases as the nonpolar end of the alcohol molecules increases Methanol Butanol Heptanol

20 Factors Affecting Solubility The more similar the intermolecular attractions, the more likely one substance is to be soluble in another.

21 Factors Affecting Solubility The more similar the intermolecular attractions, the more likely one substance is to be soluble in another.

22 Factors Affecting Solubility Glucose (which has hydrogen bonding) is very soluble in water, while Cyclohexane (which only has dispersion forces) is not.

23 Factors Affecting Solubility Vitamin A is soluble in nonpolar compounds (like fats). Vitamin C is soluble in water.

24 Intermolecular Forces Why does a substance dissolve?

25 Energy Changes in Solution Three processes affect the energetics of the solution process: 1. Separation of solute particles 2. Separation of solvent particles 3. Interactions (attraction) between solute and solvent Note: Steps 1 and 2 are sometimes combined as a single step

26 Energy Changes in Solution Exothermic Endothermic The enthalpy change of the solution process depends d on H for each of these steps.

27 The Exothermic Solution We can break down the process into separate steps: 1. Ions break free from crystal lattice, H 1 2. Intermolecular forces between solvent, H 2 molecules are broken 3 H d ti ( l ti ) f 3. Hydration (solvation) of the ions by water molecules, H 3 H 3 H 1 + H 2 Process

28 The Exothermic Solution Process Energy (enthalpy) of hydration (solvation) of the ions Energy (enthalpy) of hydration (solvation) of the ions by water molecules, increases with increasing charge density of the ions.

29 The Exothermic Solution Process

30 Calculating the Enthalpy of an Exothermic Solution In this example, KF is dissolved in water. Step 2, the separation of solvent molecules, l is incorporated into the lattice energy

31 The Endothermic Solution The process occurs in the same steps as an exothermic process Process The energies are different:. H 3 H 1 + H 2

32 Enthalpy Is Only Part of the Picture The reason is that increasing the disorder or randomness (known as entropy) of a system tends to lower the energy of the system. So even though enthalpy may increase, the overall energy of the system can still decrease if the system becomes more disordered.

33 Types of Solutions Saturated Solvent holds as much solute as is possible at that temperature. To insure saturation, a small amount of solute remains undissolved on the bottom of the container Dissolved solute is in dynamic equilibrium with solid solute particles.

34 Unsaturated Types of Solutions Less than the maximum amount of solute for that temperature is dissolved in the solvent. The amount of solute in the solution can vary from a small amount to almost saturated

35 Supersaturated Types of Solutions Solvent holds more solute than is normally possible at that temperature. These solutions are unstable; crystallization can usually be stimulated by adding a seed crystal or scratching the side of the flask.

36 In general, the solubility of gases in water increases with increasing mass. Larger molecules have stronger dispersion forces. Gases in Solution

37 Gases in Solution The solubility of liquids ids and solids does not change appreciably with pressure. The solubility of a gas in a liquid is directly proportional to its pressure in contact with the liquid.

38 Henry s Law S g = kp g where S g is the solubility of the gas; k is the Henry s law constant t for that t gas in that solvent; P g is the partial pressure of the gas above the liquid.

39 Temperature Generally, the solubility of solid solutes in liquid solvents increases with increasing temperature.

40 Temperature The opposite is true of gases: Carbonated soft drinks are more bubbly if stored in the refrigerator. Warm lakes have less O 2 dissolved in them than cool lakes.

41 Some Henry s Law Constants for Common Gases at 25 C Gas k H (mol/atm) H x 10 4 N x 10 4 O 13x Cl x 10 2 Br x 10 1 I He 3.7 x 10 4 Ne 4.5 x 10 4 Ar 1.4 x 10 3 NH x 10 1 CO x 10 2 NO x 10 2 SO HCl 2.0 x 10 1 CH x 10 3 C 2 H x Reference: Sander, Rolf, Compilation of Henry s Law Constants for Inorganic and Organic Species of Potential Importance in Environmental Chemistry, 1999

42 Gases in Solution The Fizz-Keeper Claims to keep soft di drinks from going flat! Does it work? Why? Why not?

43 How does the solubility of a gas change with pressure?

44 Expressing Concentrations of Solutions

45 Can be expressed as: Percent, % Percent by mass, % (m/m) Percent by volume, % (v/v) Percent mass-volume, % (m/v) for solutions of liquids for solids in liquids % of A = amount of A in solution total t amount of solution 100 Most commonly, we use percent by mass

46 Parts per Million and Parts per Billion Parts per Million (ppm) ppm = Parts per Billion (ppb) ppb = mass of A in solution total mass of solution 106 mass of A in solution total mass of solution 109

47 Molarity (M) M = An alternate equation is mol of solute L of solution g 1000 ml M solute L MW ml solute solution Note: Because volume is temperature dependent, Molarity can change with temperature. t

48 Mole Fraction (X) X A = moles of A total moles in solution For a solution of two or more components, A, B, etc X A + X B + = 1

49 Molality (m) m = mol of solute kg of solvent An alternate equation is g 1000 g solute kg m MW g solute solvent Note: Because both moles and mass do not change with temperature, molality (unlike molarity) is not temperature t dependent. d

50 Concentrations: A Summary

51 Colligative Properties Changes in 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 Melting point depression Osmotic pressure

52 Vapor Pressure Due to both temperature effects and energy transfers from collisions, i molecules l on the surface of a liquid are able to gain sufficient kinetic energy to escape into the atmosphere

53 Vapor Pressure At any temperature, some molecules in a liquid have enough energy to escape. As the temperature rises, the fraction of molecules that have enough energy to escape increases.

54 Vapor Pressure If the container is open to the atmosphere, the molecules simply escape. This process is called evaporation. As molecules escape from the surface, they take energy with them resulting in a cooling effect on the liquid.

55 Vapor Pressure A desert water bag (left) g( ) A desert canteen (center) An Army canteen (right)

56 Vapor pressure increases with temperature. When the vapor pressure of a liquid equals the atmospheric pressure, the liquid boils. The normal boiling point of a liquid is the temperature at which its vapor pressure is 760 torr. Vapor Pressure

57 Vapor Pressure If the container is closed to the atmosphere, as more molecules l escape the liquid, id the pressure they exert increases.

58 Vapor Pressure In a sealed container, eventually, the air space in the container becomes saturated with vapor molecules. Th li id d h t t f d i The liquid and vapor reach a state of dynamic equilibrium: as liquid molecules evaporate, vapor molecules condense at the same rate. This is called the vapor pressure equilibrium

59 Vapor pressure of water at various temperatures Temperature Pressure Temperature Pressure Temperature Pressure Temperature Pressure CC mm Hg CC mm Hg CC mm Hg CC mm Hg

60 Vapor Pressure Because of solute-solvent intermolecular attraction, higher concentrations of nonvolatile solutes make it harder for solvent to escape to the vapor phase. Therefore, the vapor pressure of a solution is lower than that of the pure solvent.

61 Vapor Pressure There are less molecules of the solvent in the vapor phase above the solution

62 Raoult s Law The VP of H 2 O (or the solvent) over a solution depends on the number of H 2 O molecules per solute molecule. P solvent is is proportional to X solvent P solvent = X solvent P solvent where X solvent is the mole fraction of the solvent P solvent is the normal vapor pressure of the solvent at that temperature This equation is known n as Raoult s Law

63 Raoult s Law An ideal solution is one that obeys Raoult s law. P A = X A P A Because mole fraction of solvent, X A, is always less than 1, then P o A is always less than P A. A The vapor pressure of solvent over a solution is always LOWERED!

64 Vapor Pressure Lowering

65 Boiling Point Elevation and Freezing Point Depression Nonvolatile solute-solvent interactions cause solutions to have higher boiling points and lower freezing gpoints than the pure solvent.

66 Boiling Point Elevation The change in boiling point is proportional to the molality of the solution: T b = K b m where: K b is the molal boiling point elevation constant, a property of the solvent. T b is added to the normal boiling point of the solvent.

67 Freezing Point Depression The change in freezing point can be found similarly: T f = K f m Where: K f is the molal freezing point depression constant of the solvent. T f is subtracted from the normal freezing point of the solvent.

68 Boiling Point Elevation and Freezing Point Depression Note that in both equations, T does not depend on what the solute is, but only on how many particles are dissolved. T b = K b m T f = K f m

69 Colligative Properties of Electrolytes l t Since colligative properties depend on the number of particles dissolved, solutions of electrolytes (which dissociate in solution) should show greater changes than those of nonelectrolytes.

70 Electrolytes Strong electrolyte: a compound that dissociates completely to ions in an aqueous solution. Compound Dissociates to No. of ions per formula unit NaCl Na + and Cl - 2 CaCl 2 Ca 2+ and 2 Cl - 3 K 2 SO 4 2 K + and SO Mg 3 (PO 4 ) 2 3 Mg 2+ and 2 PO Ionic substances dissociate into the ions and polyatomic ions used din writing the chemical lformulas of fthe compounds Weak electrolyte: a compound that only partially dissociates to ions in an aqueous solution. An example is acetic acid, HC 2 H 3 O 2, which exists as HC 2 H 3 O 2 molecules, H + and C 2 H 3 O 2- in water solution

71 Colligative Properties of Electrolytes l t However, a1m solution of NaCl does not However, a 1 M solution of NaCl does not show twice the change in freezing point that a 1 M solution of methanol does.

72 The van t Hoff Factor One mole of NaCl in water does not really give rise to two moles of ions.

73 The van t Hoff Factor Some Na + and Cl reassociate for a short time, so the true concentration of particles is somewhat less than two times the concentration of NaCl.

74 The van t Hoff Factor Reassociation is more likely at higher concentration. ti Therefore, the number of particles present is concentration ti dependent. d

75 The van t Hoff Factor We modify the previous equations by multiplying l i by the van t Hoff factor, i T f = K f m i

76 Osmosis Some substances form semipermeable membranes, allowing some smaller particles to pass through, but blocking other larger particles. In biological systems, most semipermeable membranes allow water to pass through, but solutes are not free to do so.

77 Osmosis In osmosis, there is net movement of solvent from the area of higher solvent concentration (lower solute concentration) to the area of lower solvent concentration (higher solute concentration).

78 Osmosis The semipermeable membrane allows only the movement of solvent molecules. Solvent molecules move from pure solvent to solution in an attempt to make both have the same concentration of solute. The driving force is entropy

79 Process of Osmosis

80 Osmosis Dissolving the shell in vinegar Egg in pure water Egg in corn syrup

81 Osmotic Pressure The pressure required to stop osmosis, known as osmotic pressure,, is = ( n )RT = MRT V Osmotic pressure where M is the molarity of the solution If the osmotic pressure is the same on both sides of a membrane (i.e., the concentrations are the same), the solutions are isotonic.

82 Osmosis Osmosis of solvent from one solution to another can continue until the solutions are ISOTONIC they have the same concentration.

83 Normal red blood cells

84 Osmosis in Blood Cells If the solute concentration ti outside the cell is greater than that inside the cell, the solution is hypertonic. Water will flow out of the cell, and crenation results.

85 Osmosis in Cells If the solute concentration outside the cell is less than that inside the cell, the solution is hypotonic. Water will flow into the cell, and hemolysis results.

86 Dialysis Dialysis: the separation of larger molecules, dissolved substances, or colloidal particles from smaller molecules, substances, or colloidal particles by a semipermeable membrane.

87 Water Treatment

88 Reverse Osmosis Water Desalination Water desalination plant in Tampa

89 Reverse Osmosis

90 Water Filters: Ion Exchange

91 Colloids In true solutions, the maximum diameter of a solute particle is about 1 nm. Colloid: a solution in which the solute particle diameter is between 1nm and 1000 nm. Colloid particles have very large surface areas, which accounts for these two characteristics of colloidal systems; they scatter light and, therefore, appear turbid, cloudy, or milky. they form stable dispersions; that is, they do not settle out.

92 Types of Colloids Suspensions of particles larger than individual ions or molecules, but too small to be settled out by gravity.

93 Colloids John Tyndall ( ) Tyndall effect: a characteristic of colloids in which light passing through the colloid is scattered (i.e., reflected off of colloidal particles). Examples of colloids that exhibit the Tyndall effect are smoke, serum, and fog.

94 Tyndall effect Colloids

95 Why is the sky blue? Normal sky color Pale blue sky near horizon

96 Colloids Robert Brown ( ) In 1827 the English botanist Robert Brown noticed that pollen grains suspended in water jiggled about under the lens of the microscope, following a zigzag path. Brownian motion: the random motion of colloidsize particles.

97 Examples of Brownian motion are the motion of dust particles in the air; what we see are the dust particles due to scattered light. Colloids Joseph Perrin 1908

98 Colloids Why do colloidal particles remain in solution despite all the collisions due to Brownian motion? Most colloidal particles carry a large solvation layer; if the solvent is water, as in the case of protein molecules in the blood, the large number of surrounding water molecules prevents colloidal molecules from touching and sticking together. Because of their large surface area, colloidal particles acquire charges from solution; for example, they all may become negatively charged. When a charged colloidal particle encounters another particle of the same charge, they repel each other.

99

100 Colloids in Biological Systems Some molecules have a polar, hydrophilic (waterloving) end and a nonpolar, hydrophobic (water- hating) end.

101 Colloids in Biological Systems Sodium stearate is one example of such a molecule. hydrophilic (polar) end hydrophobic (nonpolar) end.

102 Micelle

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