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Chapter 10 PowerPoint Presentation

Chapter 10

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Chapter 10

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  1. Chapter 10 Particle Forces

  2. States of Matter Solid- Particles moving about a fixed point Liquid-Particles moving about a moving point Gas-Particles filling the volume of the container with complete random motions.

  3. Particle Forces Affect Solubility Vapor Pressures Freezing Points Boiling Points

  4. Particle Forces • Intramolecular forces (Relative strength = 100) • Ionic bonding • Covalent bonding • Interparticle forces • Ion-dipole forces • Dipole-dipole (Polar molecules) • (relative Strength = 1) • London Forces (Dispersion forces)( Nonpolar molecules) • (relative strength = 1) • Hydrogen Bonding (Relative strength = 10)

  5. Ion-Ion Interactions • Coulomb’s law states that the energy (E) of the interaction between two ions is directly proportional to the product of the charges of the two ions (Q1 and Q2) and inversely proportional to the distance (d) between them.

  6. Predicting Forces of Attraction • Coulombs Law indicates the increases in the charges of ions will cause an increase in the force of attraction between a cation and an anion. • Increases in the distance between ions will decrease the force of attraction between them.

  7. Size of Ions

  8. Lattice Energy • The lattice energy (U) of an ionic compound is the energy released when one mole of the ionic compound forms from its free ions in the gas phase. M+(g) + X-(g) ---> MX(s)

  9. Comparing Lattice Energies

  10. Practice Determine which salt has the greater lattice energy. • MgO and NaF • MgO and MgS

  11. Lattice Energy Using Hess’s Law

  12. Electron Affinity • Electron affinity is the energy change occurring when one mole of electrons combines with one mole of atoms or ion in the gas phase. • Step 4 in diagram on the last slide. Cl(g) + e-(g) ---> Cl-(g) ΔHEa = -349 kj/mole

  13. Calculating U Na+(g) + e-(g) ---> Na(g) -HIE1 Na(g) ---> Na(s) -Hsub Cl-(g) ---> Cl(g) + e-(g) -HEA Cl(g) ---> 1/2Cl2(g) -1/2HBE Na(s) + 1/2Cl2(g) ---> NaCl(s) Hf Na+(g) + Cl-(g) ---> NaCl(s) U U = Hf - 1/2HBE - HEA - Hsub - HIE1

  14. Lattice energy for NaCl.

  15. Interactions Involving Polar Molecules • An ion-dipole interaction occurs between an ion and the partial charge of a molecule with a permanent dipole. • The cluster of water molecules that surround an ion in aqueous medium is a sphere of hydration.

  16. Illustrates of Ion-Dipole Interaction

  17. The Solution Process Bond Breaking Processes • Break solute particle forces (expanding the solute), endothermic • Break solvent particle forces (expanding the solvent), endothermic

  18. The Solution Process Attractive Forces • Energy released when solute solvent are attracted, exothermic • Energy is released due to new attractions • Ion dipole if the solute is ionic and the solvent polar. • London-Dipole for nonpolar solute and polar solvent • Dipole-dipole for polar solute and polar solvent

  19. The Solution Process Theromodynamics • Enthalpy • Entropy (Perfect crystal, assumed to be zero) • Gibbs free energy

  20. The Solution Process Oil dissolving in water • London forces holding the oil molecules together are large do to the large surface area of the oil • The hydrogen bonds holding water molecules together are large • The forces of attraction of between nonpolar oil and polar water are weak at best • Thus the overall process is highly endothermic and not allowed thermo chemically

  21. The Solution Process Oil dissolving in water • Entropy should be greater than zero • Free energy should be greater than zero, since the process is highly endothermic • Thus the overall process is nonspontaneous

  22. The Solution Process Sodium chloride dissolving in water • Large amount of energy is required to break the ionic lattice of the sodium chloride (expand solute) • Large amount of energy is required to separate the water molecules to expand the solvent breaking hydrogen bonds • Formation of the ion dipole forces releases a large amount of energy, strong forces (why?) • The sum of the enthalpies is about +6 kJ (slightly endothermic), which is easily overcome by the entropy of the solution formation.

  23. Water as a Solvent • Water most important solvent, important to understand its solvent properties • Most of the unusual solvent properties of water stem from it hydrogen bonding nature • Consider the following ∆S of solution KCl →75j/K-mole LiF→-36j/K-mole CaS→-138 j/K-mole

  24. Water as a Solvent • We would expect ∆S>0 for all solutions, right? • But two are negative, why? • Obviously, something must be happening for the increased order. • Ion-dipole forces are ordering the water molecules around the ions, thus causing more order in water i.e. less positions for water than in the pure liquid state

  25. Water as a Solvent • Smaller ions, have stronger ion dipole forces, thus pulling water closer, therefore less positions • Also, ions with a charge greater than one will attract to water stronger than a one plus charge, thus more order due to less space between particles

  26. Dipole-Dipole Interactions • Dipole-dipole interactions are attractive forces between polar molecules. • An example is the interaction between water molecules. • The hydrogen bond is a special class of dipole-dipole interactions due to its strength.

  27. Dipole-Dipole Forces Dipole-dipole (Polar molecules) • Alignment of polar molecules to two electrodes charged + and δ– • Forces compared to ionic/covalent are about 1 in strength compared to a scale of 100, thus 1% δ+ δ– δ– δ– δ+ δ+ H Cl H Cl H Cl

  28. Dipole Dipole Interactions

  29. Hydrogen Bonding • Hydrogen bonding a stronger intermolecular force involving hydrogen and usually N, O, F, and sometimes Cl • Stronger that dipole-dipole, about 10 out of 100, or 10 • Hydrogen needs to be directly bonded to the heteroatom • Since hydrogen is small it can get close to the heteroatom • Also, the second factor is the great polarity of the bond.

  30. Hydrogen Bonding in HF(g)

  31. Hydrogen Bonding in Water around a molecule in the solid in the liquid

  32. Boiling Points of Binary Hydrides

  33. Interacting Nonpolar Molecules • Dispersion forces (London forces) are intermolecular forces caused by the presence of temporary dipoles in molecules. • A instantaneous dipole (or induced dipole) is a separation of charge produced in an atom or molecule by a momentary uneven distribution of electrons.

  34. Illustrations

  35. Strength of Dispersion Forces • The strength of dispersion forces depends on the polarizability of the atoms or molecules involved. • Poarizability is a term that describes the relative ease with which an electron cloud is distorted by an external charge. • Larger atoms or molecules are generally more polarizable than small atoms or molecules.

  36. London Forces (Dispersion) • Induced dipoles (Instantaneous ) • Strength is surface area dependent • More significant in larger molecules • All molecules show dispersion forces • Larger molecules are more polarizable

  37. Instantaneous and Induced Dipoles

  38. Molar Mass and Boiling Point

  39. London vs Hydrogen Bonding

  40. The Effect of Shape on Forces

  41. Practice Rank the following compound in order of increasing boiling point. CH3OH, CH3CH2CH2CH3, and CH3CH2OCH3

  42. Practice Rank the following compound in order of increasing boiling point. CH3OH, CH3CH2CH2CH3, and CH3CH2OCH3 MM IM Forces CH3OH 32.0 London and H-bonding CH3CH2CH2CH3 London, only 58.0 CH3CH2OCH3 60.0 London and Dipole-dipole

  43. Practice Rank the following compound in order of increasing boiling point. CH3OH, CH3CH2CH2CH3, and CH3CH2OCH3 MM IM Forces CH3OH 32.0 London and H-bonding CH3CH2CH2CH3 London, only 58.0 CH3CH2OCH3 58.0 London and Dipole-dipole The order is: CH3CH2CH2CH3 < CH3CH2OCH3< CH3OH

  44. Polarity and Solubility • If two or more liquids are miscible, they form a homogeneous solution when mixed in any proportion. • Ionic materials are more soluble in polar solvents then in nonpolar solvents. • Nonpolar materials are soluble in nonpolar solvents. • Like dissolves like

  45. Polarity and Solubility • If two or more liquids are miscible, they form a homogeneous solution when mixed in any proportion. • Ionic materials are more soluble in polar solvents then in nonpolar solvents. • Nonpolar materials are soluble in nonpolar solvents.

  46. Polarity and Solubility How does polarity effect solubility? The thermodynamic argument, is that the lower the potential energy, the more stable the system. If subtracting the potential energy of the solute from the potential energy of the original solute and solvent is negative (exothermic) then solution is thermodynamically favored.

  47. Polarity and Solubility How does polarity effect solubility? Non polar solute and solvent: The forces holding these particles together are London Dispersion forces, the weakest of all of the inter-particle forces. The strength of these forces are relative to the surface area if solute and solvent are of similar size, then about the same amount of energy is required to separate solute and solvent particles from each other. And about the same amount of energy is released when solute and solvent are attracted to each other forming a solution. Thus we predict non polar solutes and solvents should dissolve

  48. Polarity and Solubility How does polarity effect solubility? Non polar solute and polar solvent: Considering solutes and solvents of similar surface area it should be noted that more energy is required to separate the polar solvent molecules from each other, since dipole-dipole interactions are stronger. The only interaction between a nonpolar solute and polar solvent would be London Dispersion forces, so the energy released is much less than required for separating the solvent and solute. Subtracting the potential energy of the products from reactants would give a positive (endothermic) result and the solution would be less stable than the dissolution.

  49. Practice Rank the following compound in order of increasing boiling point. CH3OH, CH3CH2CH2CH3, and CH3CH2OCH3

  50. Solubility of Gases in Water • Henry’s Law states that the solubility of a sparingly soluble chemically unreactive gas in a liquid is proportional to the partial pressure of the gas. • Cgas = kHPgas where C is the concentration of the gas, kH is Henry’s Law constant for the gas.