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IM Forces

IM Forces. Section 10.1. States of Matter. Forces Between Particles in Solids and Liquids. Ionic compounds Attractive forces between oppositely charged ions hold ionic compounds together. Ionic bonds are the strongest interparticle force.

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IM Forces

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  1. IM Forces Section 10.1

  2. States of Matter

  3. Forces Between Particles in Solids and Liquids • Ionic compounds • Attractive forces between oppositely charged ions hold ionic compounds together. • Ionic bonds are the strongest interparticle force. • Smaller the ion and the larger the charge on the ion the stronger the attractive forces among the ions

  4. Ionic Bonding

  5. Forces Between Particles in Solids and Liquids • Forces between molecular compounds • Intermolecular (IM) forces between molecules attract molecules to each other in the liquid and solid state. • IM forces are very weak as compared to ionic or covalent bonds

  6. IM Forces Three types of IM Forces • Dipole-dipole force • Hydrogen “bonding” • London dispersion forces See pages 440-442

  7. Interparticle Forces and Physical Properties • The stronger the attractive forces between particles in a liquid or solid, the • Higher the: • Melting point • Boiling point • Surface tension • Viscosity • Lower the: • Vapor pressure

  8. IM Forces • Dipole-dipole forces • Attractive forces between oppositely charged dipoles. • Dipole-dipole forces are found between polar compounds. • The more polar the compound the stronger the dipole-dipole force.

  9. IM Forces • Hydrogen “bonds” • Attractive force between a d+ H bonded to an O, N, or F and a d- O, N, or F generally on another molecule. • Really a relatively strong dipole-dipole force • Hydrogen bonding is the strongest of the IM forces. • H bonding is very important in water and in many biological molecules.

  10. Hydrogen “bond” is a weak attractive force between a d + hydrogen and a d-O, N, or F in a second polar bond

  11. London Dispersion Forces • London Dispersion force • Very weak and short-lasting attractive forces between temporary dipoles • See figure 10.5 • Weakest of the IM forces

  12. London Dispersion Forces • London Dispersion forces • Found between all molecules in liquid/solid state. • Of greatest significance in nonpolar molecules as it’s the only IM force between nonpolar molecules • The larger the molecule the stronger the dipersion forces.

  13. Dispersion Forces Occur between every compound and arise from the net attractive forces amount molecules which is produced from induced charge imbalances The magnitude of the Dispersion Forces is dependent upon how easily it is to distort the electron cloud. The larger the molecule the greater it’s Dispersion Forces are.

  14. Dispersion Forces and Molecular Shape • Elongated molecules have higher dispersion forces than compact molecules • Ringed structures have higher dispersion forces than straight chain molecules. • Consider: • Hexane • Cyclohexane • 2,2 – dimethyl butane

  15. Interparticle Forces • Weakest to Strongest: Intermolecular forces – all relatively weak London dispersion forces Dipole-dipole force Hydrogen Bonding Ionic bond - BY FAR THE Strongest: - not an IM Force

  16. Properties of Liquids • Freezing and boiling point • Surface tension • Capillary action • Viscosity • Which are directly related to the strength of the IM forces present between molecules?

  17. Change of State • Normal Freezing/Melting point • temperature at which the liquid and solid state co-exist at 1 atm pressure • Normal boiling point • temperature at which the liquid and gaseous state co-exist at 1 atm pressure • Predict the relative BP of: • Methane, acetone, methanol, ethanol, NaCl

  18. Surface Tension • Surface tension • Resistance of a liquid to increase its surface area • Measure of the energy needed to break the IM forces at the surface

  19. Capillary Action • Capillary action • Spontaneous rising of a liquid in a narrow tube • Related terms: • Cohesive forces – attractive forces among like molecules • Adhesive forces – attractive forces among dislike molecules

  20. See Figure 10.7, page 444 Concave meniscus Convex meniscus Adhesion > Cohesion Cohesion > adhesion

  21. Viscosity • Viscosity – resistance of a liquid to flow • Highly viscous liquids are thick (syrupy) • Consider relative viscosity of: • Propanol • Glycerol

  22. Graphite • Layers of ringed carbon structures • Each C is bonded to 3 other C • Each C is sp2 hybridized

  23. Diamond • A diamond is a gigantic molecule, each C atom is bonded to 4 other C atoms • Each C is sp3 hybridized

  24. A phase diagram summarizes the conditions at which a substance exists as a solid, liquid, or gas. Phase Diagram of Water 11.9

  25. 11.9

  26. CH 11: Properties of Solutions • Describing Solutions – concentration units • Energetics of solution formation • Colligative Properties of solutions • BP elevation • FP depression • Osmotic pressure • Vapor Pressure

  27. Terms • Solution – homogeneous mixture • Solvent – generally the larger component of the solution • Determines the physical state of the solution • Solute – generally the smaller component of the solution • Solute is dispersed in the solvent

  28. Solution Composition • Concentrated solution – relatively large amount of solute • Dilute solution – relatively small amount of solute

  29. Solution Composition • Unsaturated solution –solution with less than the maximum amount of solute that will normally dissolve at a given temperature • Saturated solution - solution with maximum amount of solute that will normally dissolve at a given temperature

  30. Solution Composition • Super-saturated solution - solution with more than the maximum amount of solute that will normally dissolve at a given temperature

  31. Concentration Units Molarity (M) = moles solute/Liters solution Molality (m) = moles solute/kg solvent Mass % = Mass solute/mass solution x100% Mole fraction (cA) = moles A/total moles

  32. Practice! • Start by writing definitions for the concentration units M = m = Mass % = Mole fraction =

  33. Starting with Molarity Solution: • 3.75 M H2SO4 solution with a density of 1.23 g/mL Calculate: • Mass % • Molality • mole fraction of H2SO4

  34. Starting with Masses Solution: • A solution is made by combining 66.0 grams of acetone (C3H6 O) with 146.0 grams of water. • Solution has a density of 0.926 g/mL Calculate: • Molarity – need volume of solution • Mass % • Molality • Mole fraction of acetone

  35. Starting with Mass % Solution: • 35.4 % H3PO4 • Density of 1.20 g/mL Calculate: • Molarity • Molality • Mole fraction of H3PO4

  36. Starting with Molality Solution: • 2.50 m HClsolution • Density of 1.15 g/mL Calculate: • Molarity – need _______ • Mass % • Mole fraction of HCl

  37. Solution Formation Formation of a solution involves 3 steps • Separate the solute particles • expand the solute • Separate the solvent particles • Expand the solvent • Form the solution • Solute and solvent interact

  38. Solution Formation • Each step of solution formation involves energy and has a DH. DH1 = energy needed to separate the solute DH2 = energy needed to separate the solvent DH3 = energy released when solution forms

  39. Solution Formation DHsolution = DH1 +DH2 +DH3 Solutions form when the DHsolution is a small value – see page 492

  40. Factors Impacting Solubility • Structure – like dissolves like • #38 on page 520

  41. Factors Impacting Solubility • Pressure • Pressure has little impact on the solubility of liquids and solids • Pressure has a significant impact on the solubility of gases in a liquid • The higher the pressure of gaseous solute above a liquid the higher the concentration of the gas in the solution

  42. Henry’s Law • Henry’s Law: C = kP C = Concentration of dissolved gas k = solution specific constant P = partial P of the solute gas above the solution • No calculations required. Page 494

  43. Temperature and Solubility • Temperature has variable effects on the amount of solid that will dissolve in an aqueous solution! • See figure 11.6 page 496 • Solutes do dissolve more rapidly at higher temperatures

  44. Temperature and Solubility • The solubility of a gas in water decreases as temperature increases. • See figure 11.7 on page 496 • Thermal pollution – read the story on page 497 when you get a chance

  45. Vapor Pressure of Solutions • See Raoult’s Law on page 498 • Psolution= csolvent P0solvent

  46. Colligative Properties • Colligative properties • properties of a solution that depend upon the amount of dissolved solute, not the identity of the solute. • Freezing point depression • Boiling point elevation • Osmotic Pressure • Note: I will be weaving section 11.7 and the van’t Hoff factor (i) into my consideration of these properties and not consider it separately.

  47. Colligative Properties • FP = Kf m i • BP = Kb m i See page 505 for needed constants

  48. Calculating the bp or fp of a solution • Calculating the molar mass of a solute from fp or bp data

  49. Osmotic Pressure • Osmotic Pressure (P) is often used to determine the molar mass of large biological molecules • See figure 11.17 on page 508 P = MRTi

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