1 / 40

Solutions

Solutions. Chapter 13. Solution Process. Solution – homogeneous mixture of solute and solvent Can be gases, liquids, or solids Solvent is component present in the largest amount (rest are solutes) Intermolecular forces are rearranged with solutions of condensed phases.

shasta
Télécharger la présentation

Solutions

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Solutions Chapter 13

  2. Solution Process • Solution – homogeneous mixture of solute and solvent • Can be gases, liquids, or solids • Solvent is component present in the largest amount (rest are solutes) • Intermolecular forces are rearranged with solutions of condensed phases. • Ex. NaCl in Water – Water molecules orient themselves on NaCl; H-bonds between waters have to be broken; NaCl dissociates; ion-dipole forces form; interaction is called solvation or hydration

  3. The Solution Process

  4. Energy Changes • Three steps involving energy in the formation of a solution • Separation of solute molecules (∆H1) • Separation of solvent molecules (∆H2) • Formation of solute-solvent interactions (∆H3) • Enthalpy change in the solution process • ∆Hsoln = ∆H1 + ∆H2 + ∆H3 • ∆Hsoln can be + or – • Breaking attractive intermolecular forces is always endothermic (+) • Forming attractive intermolecular forces is always exothermic (-)

  5. Energy changes

  6. Energy Changes and Solution Formation • In general, solutions form if ∆Hsoln is negative. • If ∆Hsoln is too endothermic a solution will not form. • “Rule of Thumb” – Like dissolves Like • Polar solvents dissolve polar solutes • Nonpolar solvents dissolve nonpolar solutes • Ex. NaCl in gasoline: Only weak interactions are possible because gasoline is nonpolar; interactions do not compensate for separation of ions; NaCl does not dissolve to any great extent in gasoline.

  7. Solution Formation:Spontaneity and Disorder • Spontaneous process occurs without outside intervention • Energy of system decreases – spontaneous • Tend to be exothermic • Some are endothermic • Entropy is a measure of randomness or disorder of a system. • Solution formation is favored by the increase in entropy that accompanies solution formation (most cases) • A solution can spontaneously mix when Esoln is not significantly decreasing when entropy is increasing.

  8. Solution Formation:Chemical vs. Physical • Solutions can be created physically or chemically. • Ni(s) + 2HCl(aq)  NiCl2(aq) + H2(g) • Chemical form has changed (Ni  NiCl2) • When water is removed, no Ni is found (only NiCl2 remains) • NaCl(s) + H2O(l)  Na+(aq) + Cl-(aq) • When water is removed from the solution, NaCl is found.

  9. Solubility • Dissolution – formation of a solution • Crystallization – “destruction” of a solution • If these processes are in equilibrium, the solution is saturated. • Solubility is the amount of solute required to form a saturated solution. • Unsaturated • Supersaturated

  10. Factors affecting Solubility • Nature of the solute and solvent • Intermolecular forces – favorable? • Miscible vs. immiscible • Temperature • Generally, as temperature increases, solubility increases. • Sometimes not the cases – gases are less soluble at higher temperatures • Pressure • Solubility of a Gas is a function of the pressure of the gas over the solution. • Increase pressure, increase solubility • Henry’s Law – Sg = kPg

  11. Solubility: Pressure

  12. Concentration • Amount of solute per amount of solvent • Qualitative – dilute vs. concentrated • Quantitative – • Mass percentage • Mole fraction • Molarity • Molality

  13. Concentration

  14. Molality – Molarity

  15. Colligative Properties • Colligative properties depend on quantity of solute molecules. • (E.g. freezing point depression and melting point elevation.) • Lowering Vapor Pressure • Non-volatile solutes reduce the ability of the surface solvent molecules to escape the liquid. • Therefore, vapor pressure is lowered. • The amount of vapor pressure lowering depends on the amount of solute.

  16. Vapor Pressure

  17. Colligative Properties • Lowering Vapor Pressure • Raoult’s Law: PA is the vapor pressure with solute, PA is the vapor pressure without solute, and A is the mole fraction of A, then • Recall Dalton’s Law:

  18. Colligative Properties • Lowering Vapor Pressure • Ideal solution: one that obeys Raoult’s law. • Raoult’s law breaks down when the solvent-solvent and solute-solute intermolecular forces are greater than solute-solvent intermolecular forces. • Boiling-Point Elevation • Goal: interpret the phase diagram for a solution. • Non-volatile solute lowers the vapor pressure. • Therefore the triple point - critical point curve is lowered

  19. Colligative Properties: Boiling-Point Elevation • At 1 atm (normal boiling point of pure liquid) there is a lower vapor pressure of the solution. Therefore, a higher temperature is required to teach a vapor pressure of 1 atm for the solution (Tb). • Molal boiling-point-elevation constant, Kb, expresses how much Tb changes with molality, m:

  20. Colligative Properties:Freezing Point Depression • At 1 atm (normal boiling point of pure liquid) there is no depression by definition • When a solution freezes, almost pure solvent is formed first. • Therefore, the sublimation curve for the pure solvent is the same as for the solution. • Therefore, the triple point occurs at a lower temperature because of the lower vapor pressure for the solution.

  21. Colligative Properties:Freezing Point Depression • The melting-point (freezing-point) curve is a vertical line from the triple point. • The solution freezes at a lower temperature (Tf) than the pure solvent. • Decrease in freezing point (Tf) is directly proportional to molality (Kfis the molal freezing-point-depression constant):

  22. Osmosis • Semipermeable membrane: permits passage of some components of a solution. • Example: cell membranes and cellophane. • Osmosis: the movement of a solvent from low solute concentration to high solute concentration. • There is movement in both directions across a semipermeable membrane. • As solvent moves across the membrane, the fluid levels in the arms becomes uneven.

  23. Osmosis • Eventually the pressure difference between the arms stops osmosis.

  24. Osmosis • Osmotic pressure, , is the pressure required to stop osmosis: • Isotonic solutions: two solutions with the same  separated by a semipermeable membrane.

  25. Osmosis • Hypotonic solutions have a lower  than a more concentrated solution. • Hypertonic solutions have a higher  than a more dilute solution. • Osmosis is spontaneous.

  26. Ex. Red blood cells are surrounded by semipermeable membranes. • Crenation: • red blood cells placed in hypertonic solution (relative to intracellular solution); • there is a lower solute concentration in the cell than the surrounding tissue; • osmosis occurs and water passes through the membrane out of the cell. • The cell shrivels up.

  27. Ex. Red blood cells are surrounded by semipermeable membranes.

  28. Ex. Red blood cells are surrounded by semipermeable membranes. • Hemolysis: • red blood cells placed in a hypotonic solution; • there is a higher solute concentration in the cell; • osmosis occurs and water moves into the cell. • The cell bursts. • To prevent crenation or hemolysis, IV (intravenous) solutions must be isotonic.

  29. Colloids • Colloids are suspensions in which the suspended particles are larger than molecules but too small to drop out of the suspension due to gravity. • Particle size: 10 to 2000 Å. • Tyndall effect: ability of a Colloid to scatter light. The beam of light can be seen through the colloid.

  30. Types of Colloids • aerosol (gas + liquid or solid, e.g. fog and smoke), • foam (liquid + gas, e.g. whipped cream), • emulsion (liquid + liquid, e.g. milk), • sol (liquid + solid, e.g. paint), • solid foam (solid + gas, e.g. marshmallow), • solid emulsion (solid + liquid, e.g. butter), • solid sol (solid + solid, e.g. ruby glass).

  31. Hydrophobic and Hydrophilic Colloids • Focus on colloids in water. • “Water loving” colloids: hydrophilic. • “Water hating” colloids: hydrophobic. • Molecules arrange themselves so that hydrophobic portions are oriented towards each other.

  32. Hydrophobic and Hydrophilic Colloids • Typical hydrophilic groups are polar (containing C-O, O-H, N-H bonds) or charged. • Hydrophobic colloids need to be stabilized in water. • Adsorption: when something sticks to a surface we say that it is adsorbed. • If ions are adsorbed onto the surface of a colloid, the colloids appears hydrophilic and is stabilized in water.

  33. Removal of Colloid Particles • Colloid particles are too small to be separated by physical means (e.g. filtration). • Colloid particles are coagulated (enlarged) until they can be removed by filtration. • Methods of coagulation: • heating (colloid particles move and are attracted to each other when they collide); • adding an electrolyte (neutralize the surface charges on the colloid particles).

  34. Dialysis: Removal of Colloids • Semipermeable membrane is used to separate ions from colloidal particles. • Blood passes through a membrane immersed in a washing solution. • Washing solution is isotonic with ions that must be retained. • Washing solution does not contain waste products. • Washing solution removes (dialyzes) wastes and the “good” ions remain.

More Related