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Boiling-Point Elevation AT Physical Properties of Solutions 1 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display. A solution is a homogenous mixture of 2 or more substances The solute is(are) the substance(s) present in the smaller amount(s) The solvent is the substance present in the larger amount 2 We can have: Gaseous Solutions: gases can spontaneously mix in any proportion. Liquid Solutions: liquid solutions are the most common and they can be obtained by dissolution of a gaseous, liquid or solid solute. Solid Solutions: they are generally called alloys, where two or more metals are present. Mercury alloys are called amalgams, and they can be liquid or solid. Solid Solutions • Ruby is a solid solution of Cr2O3 in Corundum (Al2O3) • The variety called “pigeon blood” is one of the most precious stone. • Sapphire (less precious) color blu is due to the Fe and Ti impurities in the Al2O3 4 Ideal Solutions Ideal solutions have : ΔHsol = 0 The formation of ideal solutions is always spontaneous: ΔGsol < 0 © Dario Bressanini 5 A saturated solution contains the maximum amount of a solute that will dissolve in a given solvent at a specific temperature. An unsaturated solution contains less solute than the solvent has the capacity to dissolve at a specific temperature. A supersaturated solution contains more solute than is present in a saturated solution at a specific temperature. Sodium acetate crystals rapidly form when a seed crystal is added to a supersaturated solution of sodium acetate. 6 Three types of interactions in the solution process: • solvent-solvent interaction • solute-solute interaction • solvent-solute interaction Molecular view of the formation of solution ΔHsoln = ΔH1 + ΔH2 + ΔH3 7 “like dissolves like” Two substances with similar intermolecular forces are likely to be soluble in each other. • non-polar molecules are soluble in non-polar solvents CCl4 in C6H6 • polar molecules are soluble in polar solvents C2H5OH in H2O • ionic compounds are more soluble in polar solvents NaCl in H2O or NH3 (l) 8 Two driving forces are the basis of the solution formation (solubility of two substances): 1. Disorder tendency (entropic factor). It is the only factor presents in the case of (ideal) gases, which can be mixed in any proportion. 2. Intermolecular attraction among the molecules of the two substances (energetic factor). If the attractions A-A and B-B are stronger than those A-B, the two substances will not have the tendency to mix. The solubility of a solute in a solvent depends on the balance among these two factors. Molecular Solutions In this case the solute (liquid or solid) is a molecular species, characterized by low energy intermolecular forces. In the case of liquids the two compounds are soluble if these forces are similar, or if the heterogeneous interactions are stronger than the homogeneous one. Chloroform-acetone C7H16 – C8H18 London Forces Hydrogen Bonds Ionic Solutions In this case the solute is a ionic solid, and polar solvents are able to overcome the strong ionic bonds. The factors influencing the solution process of a ionic solid are: - the lattice energy of the solid (sum of the attraction energies of anions and cations): bigger this energy, lower the tendency of the solid to dissolve in the liquid phase - the attraction energy of dipole-ion among the ions of the solute and the electric dipoles of the solvent molecules: bigger this energy, bigger the tendency to the solution formation NaCl in water: NaCl(s) → Na+(aq) + Cl-(aq) + - - - + + + - + - + - - - + - + - + - + - + IONIC CRYSTAL SOLVATED IONS NaCl(s) → Na+(aq) + Cl-(aq) This process is called hydration and the + electrostatic interaction energy of the ion with water molecules is called hydration energy. The solubility of a ionic solid - depends on the balance among lattice and hydration energy. SOLVATED IONS Bigger the lattice energy of ionic compounds, lower will be their solubilities and vice versa. The lattice energy depends on both the ionic charges and their distances : - bigger the ionic charges, bigger will be the lattice energy - higher the ionic distances (bigger ions), lower will be the lattice energy The situation is complicated by the fact that the hydration energy is also bigger for small and highly charged ions. In general the lattice energy prevails, so we can predict that: - solids formed by ions with a single charge and big + + dimensions (K , NH4 ) are generally soluble - solids formed by ions with two a three charges and 2- 3- small dimensions (S PO4 ) are generally not soluble. Temperature and Solubility Solid solubility and temperature solubility increases with increasing temperature solubility decreases with increasing temperature 15 Temperature and Solubility O2 gas solubility and temperature solubility usually decreases with increasing temperature 16 Pressure and Solubility of Gases The solubility of a gas in a liquid is proportional to the pressure of the gas over the solution (Henry’s law). c is the concentration (M) of the dissolved gas c = kP P is the pressure of the gas over the solution k is a constant for each gas (mol/L•atm) that depends only on temperature low P high P low c high c 17 Molecular view: bigger the partial pressure of the gas, higher the number of the gas molecules colliding the liquid surface and consequently dissolved into the liquid phase Henry Law pA = xAK A • Knowing the Henry constants is important for a wide range of applications Gas (in H2O) K/(10 Mpa) CO2 0.167 Carbon dioxide dissolves H2 7.12 very well in water N2 8.68 O2 4.40 19 Henry Law • At high pressure molecular Nitrogen and Oxygen are dissolved in the blood. • Emerging too quickly can induce the formation of emboli 20 Blood Solubility At high pressure molecular Nitrogen and Oxygen are dissolved in the blood. Molecular Oxygen can be consumed, but Nitrogen will remain in the blood. Concentration Units The concentration of a solution is the amount of solute present in a given quantity of solvent or solution. Percent by Mass mass of solute % by mass = x 100% mass of solute + mass of solvent mass of solute = x 100% mass of solution Mole Fraction (X) moles of A XA = sum of moles of all components 22 Concentration Units Continued Molarity (M) moles of solute M = liters of solution Molality (m) moles of solute m = mass of solvent (kg) 23 Diluted solutions Addition of solvent M1V1 = M2V2 The moles of solute are constant Mixing of two solutions M1V1 + M2V2 = MfVf = Mf(V1 + V2) 24 Vapor Pressure and Ideal Solutions pA = xApA* The solutions following the Raoult law are called Ideal Solutions Francois Raoult (1830-1901) These solutions have ΔsolH =0 25 Raoult law • Assuming that the solvent- solvent and solute-solute interactions are identical to those solute-solvent, we can conclude that the vapor pressure is: pA = xApA* 26 Solution of two volatile liquids Both liquids have a vapor pressure 27 Raoult law for volatile liquids – The vapor pressure of each component can be calculated by the Raoult law – The total pressure is the sum of both partial pressures. pA = χA pA* pB = χB pB* ptot = pA + pB = χA pA* + χB pB* 28 Vapor-Pressure of a solution P = X P 0 Ideal Solution A A A 0 PB = XB P B PT = PA + PB 0 0 PT = XA P A + XB P B Raoult’s law 29 Ideal and not Ideal Solution ptot = χA pA* + χB pB* 30 Not Ideal Solutions • The not ideal behavior is the common case for real solutions • A and B interactions are different from those AA and BB Positive deviation Negative deviation 31 POSITIVE DEVIATION FROM THE RAOULT LAW The solution shows a vapor pressure higher than that calculated by the Raoult law; The cohesion forces A-A or B-B are stronger than the adhesion forces A-B Examples of these solutions are: ethanol-water, isopropanol-water, propanol-water NEGATIVE DEVIATIONS FROM RAOULT LAW The solution shows a vapor pressure lower than that calculated by the Raoult law; The cohesion forces A-A or B-B are weaker than the adhesion forces A-B Examples of these solutions are: pyridine-water, acetic acid-water, hydracid-water DISTILLATION OF IDEAL SOLUTIONS OF TWO LIQUIDS Distillation: separation process based on the vaporization of a liquid phase by heat application and subsequent condensation by cooling the vapor in a container different from that where the vaporization was carried out. When an ideal solution of two liquids is heated to the boiling point it is possible to obtain a curve similar to that reported below, where for each composition the vapor is more rich of the more volatile component, while the liquid is more rich of the other component. These diagrams are in the isobaric (constant pressure) form. The two curves represent the composition of the liquid (blue) and the vapor (red) phases, in equilibrium at a certain temperature. The blue line is called ebullition curve, the red one condensation curve. Fractional Distillation In this technique a column (fractioning column) is posed among the vaporization and condensation site. In this column the liquid-vapor phase change is repeated several times, thanks to the condensation over the cold surface of the column and the heat transferred by the hot vapor fluxing on the column. These liquid/vapor equilibria increase the efficiency of the liquids separation. From the figure it is possible to observe this phenomenon. DISTILLATION OF NON IDEAL SOLUTIONS OF TWO LIQUIDS - 1 Solutions having positive deviations from the Raoult law show the diagram reported below; In this case we observe a solution vapor pressure higher than that expected. The distillation diagram shows a minimum called minimum azeotrope. In this point liquid and vapor have the same composition and the azeotrope behaves as a pure liquid.
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