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Chem 1040, week 2 of notes

by: Olivia Hammond

Chem 1040, week 2 of notes CHEM 1040 - 003

Marketplace > Auburn University > Chemistry > CHEM 1040 - 003 > Chem 1040 week 2 of notes
Olivia Hammond
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Includes all notes so far over the material in Chapter 13: Physical Properties of Solutions.
Fundamental Chemistry II
Ria Astrid Yngard
Class Notes
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This 9 page Class Notes was uploaded by Olivia Hammond on Friday January 29, 2016. The Class Notes belongs to CHEM 1040 - 003 at Auburn University taught by Ria Astrid Yngard in Spring 2016. Since its upload, it has received 37 views. For similar materials see Fundamental Chemistry II in Chemistry at Auburn University.


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Date Created: 01/29/16
Solutions - homogeneous mixture of two or more pure substances - the solute is uniformly dispersed throughout - A solution is made up of a solute that is disolved in the solvent - A solution that has water as its solvent is considered to be an aqueous solution - Unsaturated Solution: contains less solute than the solvent has the chapatis to dissolve at specific temperature - Saturated Solution: contains the maximum amount of solute that will dissolve in a solvent at a specific temperature - Solubility: amount of solute dissolved in a given volume of a saturated solution at a a specific temperature - Supersaturated Solution: contains more dissolved solute than is present in a saturated solution and generally unstable Solution Formation Process There is an attraction between solute-solute particles, and solvent-solvent particles. These particles will be pulled apart created solute-solvent particle attraction in addition to the other attractions as well. The attraction between Na and H2O will be Iron Dipole Interaction. - Hydration: when the solvent is water - Solvation: solute molecules are separated from one another and surrounded by solvent molecules - salvation depends on the relative strengths of these interaction between particles: solute-solute, solvent-solvent, or solute-solvent interactions Energy Changes in Solution Formation Both of these first two steps require an put 1) Solute-Solute Interactions [the separation of solute of energy to overcome intermolecular molecules from one another] = ∆H1 > 0 attraction, so they are endothermic 2) Solvent-Solvent Interactions [the separations of solvent molecules from one another] = ∆H2 > 0 However the last step is usually exothermic 3) Solvent-Solute Interactions [the mixing of solvent because it releases energy. solute molecules] = ∆H3 < 0 A system will have a tendency to loose energy to reach its most stable form. If something requires energy it is not in its most stable form. This idea is the driving force behind whether or not something will occur spontaneously. Enthalpy Dissolution of MgSO4: ∆Holution < 0 = exothermic [release of energy] Dissolution of NH4NO3: ∆Hsolution > 0 = endothermic [takes energy] Solution Process depends on: enthalpy decrease & entropy increase Entropy: a measure of how dispersed or spread out energy is - There is a natural tendency for entropy to increase or for the energy of a system to become more dispersed or spread out (unless something, like a barrier, is preventing that dispersal) - ex: farting in class —> disorder increases as the barrier is released and the gases will mix Solubility of a Substance in a Solvent “Like Dissolves Like”: molecules that have similar intermolecular forces will mix well; however, different intermolecular forces do not mix well [ex: oil and water] - Polar substances tend to dissolve in polar substances - Nonpolar substances tend to dissolve in nonpolar substances - Miscible: mixing in all proportions - Immiscible: does not mix in all proportions Vitamin A: has one polar end due to the -OH hydroxyl group and large non polar body, therefore it is fat-soluble [nonpolar dissolves in nonpolar] Vitamin C: is a polar group with -OH hydroxyl groups surrounding main structure, therefore it is water soluble [polar dissolves polar] Problem 1 [in coordination with slides on canvas] Concentration Units Molarity (M mol/L) = Moles of solute Molarity is temperature dependent; therefore, the molarity changes as temperature changes. L of Solution Molality (m or mol/kg) = moles of solute Molality is temperature independent; therefore mass of solvent (kg) it remains constant as temperature changes. What does this mean? 1.00 molal NaCl solution = 1.00 mol NaCl 1.00 molar NaCl solution = 1.00 mol NaCl 1kg water 1L solution Mole Fraction A = Xa = Moles A Percent by Mass, Parts per Million, and Parts moles total per Billion are all temperature independent Percent by Mass = mass of solute total mass of solution Parts per Million (ppm) = mass of solute X 10^6 total mass of solution Parts per Billion (ppb) = mass of solute X 10^9 total mass of solution Problem 2 [in coordination to the problems in slides] Factors Affecting Solubility - Structure “like dissolves like” - Temperature - Pressure Temperature Effects on Solubility - Gases in water: solubility will go down as temperature increases [added temperature causes added energy; therefore, gases are more likely to escape from the liquid into the gas phase] - Solids in water: generally solubility increases as temperature increases [however, this is not always true, see figure 13-4 in the book] Problem 3 [coordinates to problem on slides] Pressure Effects on Solubility Pressure: - Solubility os solids/liquids is hardly affected - Solubility of gases increases with increasing pressure - Henry’s Law: the solubility of a gas in a liquid is proportional to the pressure of the gas over the solution C = KP C = molar concentration K = Henry’s Law constant (mol/L atm) P = pressure of gas over solution (atm) Problem 4 [coordinates to problem on slides] Colligative Properties Colligative Properties: properties that depend on the number of solute particles in a solution; depend on the concentration of solute particles regardless of whether those particles are atoms, molecules, or ions. - vapor-pressure lowering - boiling-point elevation - freezing-pont depression - osmotic pressure Vapor-Pressure Lowering (non electrolyte solution = dissociate without formation of charged particles) When a non-volatile solute is disolved in a liquid, the vapor pressured exerted by the liquid decreases. This is due to the fact that the solvent has to compete at the surface with solute in order to escape into the vapor phase; therefore there is less solvent in the vapor phase which leads to lower vapor pressure. Raoult’s Law: the partial pressure of a solvent over a solution, P1, is given by the vapor 0 pressure of the pure solvent, P1 , times the mole fraction of the solvent in the solution X1: P(solvent) = X(solvent) P (solvent) X (solute) + X(solvent) = 1 —> the sum of the mole fractions of the solvent and the solute should equal to 1 - The larger the amount of solute means there is a smaller X(solvent) which equals to a lesser vapor pressure - The larger the amount of solvent is means there is a larger X(solvent) which equals to a greater vapor pressure - The decrease in vapor pressure, ∆P, is directly proportional to the solute concentration 0 0 expressed as the mole fraction. (P 1 - P1) = ∆P = (X2)(P 1) - The smaller difference in entropy between the solution and gas phases, results in a decreased tendency for solvent molecules to enter the gas phase. This lowers the vapor pressure. - The solvent in a solution will always exert a lower vapor pressure than the pure solvent. - 0 0 Solution containing solvent (A) and volatile solute (B): P(total) = (Xa)(P a) + (Xb)(P b) [due to Dalton’s law of partial pressure you can take the sum of to the total pressure] - This only holds true in an ideal solution: a solution that obeys Raoult’s Law Problem 5 [corresponds to problems in slides] Boiling-Point Elevation (non electrolyte solution) - The boiling point of the solution is greater that the boiling point of a pure solvent - This is because a higher temperature is needed to make the solvent’s vapor pressure equal to atmospheric pressure - 0 ∆Tb = Tb -T b = (Kb)(m) m = molality Kb = the molal boiling point elevation constant of the solvent Freezing-Point Depression (non electrolyte solution) - The freezing point of the solution is lesser than the freezing point of the pure solvent - When a solution freezes, the solid that separates out is actually pure solvent. the solute remains in the liquid solution. - Freezing-point depression occurs regardless of the solute’s volatility - ∆Tf = T f -Tf = (Kf)(m) m = molality Kf = the molal freezing point depression constant of the solvent Figure 13.9 The solution has greater entropy than the pure solvent. The bigger difference in entropy between the solution and the solid means that more energy must be removed from the solution for it to freeze. Thus, the solution freezes at a lower temperature than the pure solvent. (pg 545) Osmotic Pressure (non electrolyte solution) Osmosis: the selective passage of solvent molecules through a semipermeable membrane from a more dilute solution to a more concentrated one. Osmotic Pressure of a solution (∏): the pressure required to stop osmosis ∏= MRT ∏ = osmotic pressure M = molarity R = ideal gas constant T = absolute temperature ex: potatoes, red blood cells Colloids [between heterogeneous mixture and homogeneous mixture] Colloids: a dispersion of particles of one substance throughout another substance - Colloid particles are much larger than the normal solute molecules, but too small to be settled out by gravity Tyndall Effect: used to distinguish between homogeneous solution and a colloid. When a beam of light passes through a colloid, it scatters light; however a homogeneous solution does not and there is no beam of light visible through the mixture. - ex: how fog scatters light emitted from car headlights - Colloidal particles in water: - Hydrophobic = water loving - Hydrophilic = water fearing Van’t Hoff Factor (i) - in electrolyte solutions,the number of dissolved particles is increased by dissociation or ionization - The magnitudes of colligative properties are thus increased by the van’t Hoff factor (i) i = actual number of particles in solution after dissociation number of formula units initially disolved in solution - non electrolytes: i = 1 - electrolytes: i depends on the electrolyte [table 13.3 in book] - When i is measured it is usually a little less than when it is calculated because of ion pair formation (figure 13.3 in book) - Iron pair formation is when ion collide with each other and are held together by electrostatic forces for a brief period in time the number of particles in the solution is reduced, thus reducing the observed colligative properties - The more diluted the solution is the closer you will be to the calculated value of i because there is less iron pair formation and less collisions because there is more space to spread out - The more concentrated the solution is the more collisions there will be and, therefore, a greater chance for iron pair formation. - 1M NaCl —> 2M particles - ∆Tb = i (m)(Kb) Problem 6 [corresponds to problems on slides] Problem 7 [corresponds to problems on slides]


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