CHEMISTRY The Molecular Nature of Matter and Change

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1 CHEMISTRY The Molecular Nature of Matter and Change Third Edition Chapter 13 The Properties of Mixtures: Solutions and Colloids Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

2 Definitions Solutions Homogeneous Mixtures Particles are individual atoms, ions, or small molecules. Colloids Heterogeneous Mixtures Particles are either macromolecules or aggregations of small molecules that are not large enough to settle out.

3 Solute vs. Solvent Solution Terms Concepts to use for Solvent Miscible Most abundant component Physical state matches solution physical state Mix in any proportion Solubility Maximum amount of solute dissolved in a fixed amount of solvent at a specified temperature, given that excess solute is present

4 Solubilities S, (NaCl) = g / 100 ml 100 o C S, (AgCl) = g / 100 ml 100 o C S, {(NH 4 ) 2 SO 4 } = 931 g / 100 ml 100 o C

5 Table 13.1

6 What makes different substances mix? Like Dissolves Like Intermolecular Forces must be Similar!

7 Intermolecular Interactions in Mixtures Attraction Energies in kj/mol

8 Fig. 13.1b Intermolecular Interactions in Mixtures [Part 2] Attraction Energies in kj/mol

9 Solutions: Solvent vs. Solute

10 Like Dissolves Like Intermolecular Interactions

11 Table 13.2 As the percentage of the molecule which doesn t interact favorably with water increases, the solubility decreases.

12 Prob CH 3 CH 2 CH 2 CH 2 OH H C H C H C H C H H H H O H H HOCH 2 CH 2 CH 2 CH 2 OH O C H C H C H C H H H H O H H H

13 Fig B13.1 Amphipathic Molecules: Surfactants {surface active agents} Soaps

14 Types of Solutions o Solid in Liquid o Liquid in Liquid o Solid in Solid o Gas in Liquid o Gas in Gas o Gas in Solid

15 Solid Solutions: Alloys

16 Table 13.3 Gas in Liquid Solutions

17 Enthalpy Changes in Solution Processes Separation of solute Separation of solvent Mixing of solute and solvent

18 Lattice energy H latt : H soln = H latt + H hydr Lattice energy is energy gained when ions form a solid structure From the gas phase. This causes H solute to be positive and equal in magnitude to H latt Heats of hydration are always negative, so a dissolution can be Exothermic or endothermic

19 Components of Solvation Enthalpies exothermic endothermic

20 M (g) H 2 O M (aq) H <0 {always} X (g) H 2 O X (aq) H <0 {always} MX (S) H 2 O M (aq) X (aq) H > 0 {always} Charge density (ratio of charge to size) 2) 2+ ions attract more than +1 of same size 3) Small 1+ attracts more than large +2 ion

21 Fig Why do endothermic solution processes occur at all? Systems that increase in the degrees of freedom of constituents are favorable. ENTROPY measure of a system s degrees of freedom: Generally a larger number of pieces produced is entropically favored. Heats of Solution

22 Fig Dissolving Substances in Water: An EQUILIBRIUM Process

23 Fig Crystallization Saturated Supersaturated

24 Fig Solubility and Temperature Effects Generally, increasing temperature increases solubility. Gas solubility always decreases with temperature.

25 Fig Gases Dissolved in Liquids Henry s Law: Solubility of a gas is directly proportional to the partial pressure of the gas above the liquid.

26 Prob Henry s Law S (gas) = k gas x P gas [N 2] kh, (N ) P 2 N2 PN atm 0.78atm -4 [N 2] (7 x 10 mol/l atm) 0.78 atm 5 x 10-4 mol/l

27 Page 500 Solubility of a Gas: [Gas] = k H P Gas

28 Table 13.5 Solute Concentrations are Quantitative. Concentration Terms

29 Other Related Concentration Terms Parts per Thousand (ppt) Parts per Million (ppm) Parts per Billion (ppb) Mass Percent = mass solute /mass solution x 100% Volume Percent = volume solute /volume solution x 100% Mole Percent = mole solute /(mole solute + mole solvent ) x 100%

30 Once we are able to quantify solution composition, we can now predict solution properties.

31 Prob a & (a) Calculate the molarity of a solution of 0.82 g of ethanol (C 2 H 5 OH) in 10.5 ml of solution. C: 2 x 12 H: 6 x 1.0 O: 1 x 16 = 46 g/mol M = mol solute /L solution mol ethanol = 0.82 g / 46 g/mol = 1.8 x 10-2 mol L solution = 10.5 ml x (1L/1000mL) = 1.05 x 10-2 L M = 1.8 x 10-2 mol/1.05 x 10-2 L = 1.7 M m = mol solute /kg solvent = g glucose g 180 mol 563 g 1000 g/kg = 2.40 x 10-2 m g glucose = 2.43 g

32 Prob mass % = (mass solute / mass solution ) x 100% mass % = [35.0 g / (35.0 g g)] x 100% mass % = 18.9% Mole fraction = mol solute / mol solution = X solute X PrOH = mol PrOH / (mol PrOH + mol EtOH )

33 Prob M = mol solute /L solution mass % = (mass solute / mass solution )) x 100% m = mol solute / kg solvent X = mol solute / (mol solute + mol solvent )

34 Fig Electrical Conductivity Strong Electrolyte Weak Electrolyte Non-electrolyte

35 Fig Vapor Pressure of Solutions Raoult s Law P Solvent = X solvent P o Solvent

36 Fig Colligative Properties: 1. Vapor Pressure Lowering 2. Freezing Point Depression 3. Boiling Point Elevation

37 Colligative Properties Vapor Pressure Lowering P Solvent = X solute P o Solvent P Solvent = X solvent P o Solvent Boiling Point Elevation T b = k b m solute Freezing Point Depression T f = k f m solute

38 Prob P Solvent = X solute P o Solvent mol solute = 2.00 g / g/mol X solute mol molsolute mol solute solvent & mol solvent = 50.0 g / 32.0 g/mol X solute = 7.06 x 10-3 P Solvent = 7.06 x torr P Solvent = torr

39 Table 13.6 Examples of BP Elevation & FP Depression Constants

40 Prob Freezing Point Depression: T f = m solute k f, H 2O m solute = T f / k f, H 2O T f = 32.0 o F 0.00 o F = 32.0 o F (5 o C/9 o F) = 17.8 o C m solute = 17.8 o C / (1.86 o C/m) m solute = 9.56 m

41 Fig Osmotic Pressure: Π = MRT {M = n solute /V solvent }

42 Prob. 13.8b Osmotic Pressure: Π = MRT Π = 0.30 M (L atm)/(k mol) ( )K Π = 7.6 atm

43 Fig Tyndall Effect Colloids True Solution Colloidal Dispersion Particles >> wavelength of light scatters the light Brownian Motion

44 Table 13.7 Colloidal Dispersions Will not settle out by gravity. Remain suspended.

45 Fig

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