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Chapter 17. Liquid and Solution. Intermolecular Forces Dipole-Dipole Forces. Molecules with dipole moments can attract each other electrostatically by lining up so that the positive and negative ends are close to each other. Dipole-Dipole Forces.
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Chapter 17 Liquid and Solution
Intermolecular ForcesDipole-Dipole Forces • Molecules with dipole moments can attract each other electrostatically by lining up so that the positive and negative ends are close to each other.
Dipole-Dipole Forces • Dipole-dipole forces are typically only about 1% as strong as covalent or ionic bonds.
Hydrogen Bonding • Particularly strong dipole-dipole forces are seen among molecules in which hydrogen is bound to a highly electronegative atom, such as nitrogen, oxygen, or fluorine. • The hydrogen bond has only 5% or so of the strength of a covalent bond.
Relatively large electronegativity value of the lightest elements Nonpolar tetrahedral hydrides
Intermolecular ForcesLondon Dispersion Forces • Atoms can develop a momentary nonsymmetrical electron distribution that produces a temporary dipolar arrangement of charge. • This instantaneous dipole can then induce a similar dipole in a neighboring atom, leading to an interatomic attraction .
The Effect of Molecular Size for London dispersion forces • The electrons distribution around an atom or molecule can be distorted is called the polarizability (極化度). • Larger and heavier atoms and molecules exhibit stronger dispersion forces than smaller and lighter ones. • In a larger atom or molecule, the valence electrons are, on average, farther from the nuclei than in a smaller atom or molecule. They are less tightly held and can more easily form temporary dipoles.
The Effect of Molecular Shape for London dispersion forces • At room temperature, neopentane (C5H12) is a gas whereas n-pentane (C5H12) is a liquid. • London dispersion forces between n-pentane molecules are stronger than those between neopentane molecules. • The cylindrical shape of n-pentane molecules allows them to come in contact with each other more effectively than the spherical neopentane molecules.
n-pentane neopentane
Normal melting point: the temperature at which the solid and liquid states have same vapor pressure under conditions where the total pressure is 1 atm. • Normal boiling point: the temperature at which the vapor pressure of liquid is exactly 1 atm.
Supercooled and Superheated • Supercooled: supercooling occurs because, as it is cooled, the water may not achieve the degree of organization necessary to form ice at 0oC; thus it continues to exist as the liquid. • Superheated: superheating can occur because bubble formation in the interior of the liquid requires that many high-energy molecules gather in the same vicinity.
Phase Diagram of Water Supercritical phase 超臨界態 P=218 atm T=374oC
The Thermodynamics of Solution Formation • The cardinal rule of solubility is like dissolves like. Three distinct steps for the formation of solutions Step 1: Breaking up the solute into individual components. Step 2: Overcoming intermolecular forces in the solvent to make room for the solute. Step 3: Allowing the solute and solvent to interact to form the solution.
Step 1 and 2 require energy, are endothermic, since forces must be overcome to expand the solute and the solvent. • Step 3 is usually exothermic. • The overall enthalpy change associated with the formation of the solution, called the enthalpy of solution (∆Hsoln). • ∆Hsoln= ∆H1+ ∆H2+ ∆H3
The Solubility of Sodium Chloride in Water • ∆H1 is large and positive because of the strong ionic forces in the crystal that must be overcome. • ∆H2 is expected to be large and positive because of the hydrogen bonds that must be broken in water. • ∆H3 is expected to be large and negative because of the strong interactions between the ions and the water molecules.
Why is NaCl so Soluble in Water? • Consider ∆G= ∆H-T∆S • ∆H is positive and thus unfavorable. Therefore, ∆S must be positive and large enough to make ∆G negative. • ∆S1 and ∆S2 are positive since the solute and solvent are expanded. • ∆S3 would be expected to be positive in general case.
Ionic Compounds in Water • The assembling of a group of water molecules around the ion is an order producing phenomenon and would be expected to make a negative contribution to ∆S. • The more charge density (Z/r) an ion possesses, the greater this hydration effect will be. • The smaller ions presumably are able to bind the hydrating water molecules more firmly and thus show a more negative value to ∆S.
Nonpolar Compounds in Water • The dispersal of nonpolar solute particles in water can also produce negative value to ∆S. • The polar water molecules will not strongly hydrate the nonpolar molecules. • Water forms a cage to isolate the nonpolar solute from the water bulk.
Factors Affecting SolubilityStructure Effects Fat soluble hydrophobic Water soluble hydrophilic
Henry’s Law • P=kHx P: the partial pressure of the gaseous solute x: the mole fraction of the dissolved gas. kH: constant The amount of gas dissolved in a solution is directly proportional to the pressure of the gas above the solution.
The Vapor Pressure of Solutions • Nonvolatile Solute/Solvent Raoult’s Law
Raoult’s Law with Multiple Volatile compounds • Ideal solution: A solution, obeys Raoult’s law, formed with no accompanying energy change, when the intermolecular attractive forces between the molecules of the solvent are the same as those between the molecules in the separate components.
1.What is the total vapor pressure in a mixture of 50.0 g CH3OH (P° = 93.3 torr) and 25.0 g H2O (P° = 17.5 torr)? 2.In a mixture of 86.0 g C6H6 (P° = 93.96 torr) and 90.0 g C2H4Cl2 (P° = 224.9 torr), what is the total vapor pressure?
Nonideal Solution • Solutions that do not obey Raoult’s Law are called nonideal solutions. • Solute-solvent interactions are significantly different from solute-solute and solvent-solvent interactions, the solution is likely to be a nonideal solution. • Intermolecular forces between components in a dissolved solution can cause deviations from the calculated vapor pressure.
Negative Deviation from Raoult’s Law • Both components have a lower escaping tendency in the solution than in the pure liquids. d- d+
Positive Deviation from Raoult’s Law • If two liquids mix endothermically, this indicates that the solute-solvent interactions are weaker than the interactions among the molecules in the pure liquids. • More energy is required to expand the liquid than is released when the liquids are mixed. The molecules in the solution have a higher tendency to escape than expected.
ideal system positive deviation negative deviation
Colligative Properties In the dilute solution, this change in solvent chemical potential leads to a change in the vapor pressure, the normal boiling point and the normal freezing point and causes the phenomenon of osmotic pressure. They depend only on the number of the solute particles in an ideal solution
Molal Freezing Point Depression Constant Molal Boiling Point Elevation Constant
=MRT : the osmotic pressure in atmospheres M: molarity of the solute R: gas law constant T: Kelvin temperature Osmotic Pressure
