Chapter 12 Chemistry Notes PDF

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Chapter 12 Chemistry Notes...

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Chapter 12 Chemistry Notes: Solutions: Sec 12.1: Thirsty Solutions: Why You Shouldn’t Drink Seawater: Solution-a homogenous mixture of two or more substances or components Solvent-the majority component of a solution Solute-the minority component of a solution Sec 12.2: Types of Solutions and Solubility: Solubility-the amount of the substance that will dissolve in a given amount of solvent Entropy-a measure of energy randomization or energy dispersion in a system Sec 12.3: Energetics of Solution Formation: Three Steps of Solution Formation: 1.) Separating the solute into its constituent particles *Note: always positive ∆ H (endothermic) because energy is required to overcome the forces holding the solute together 2.) Separating the solvent particles from each other to make room for the solute particles *Note: Always positive ∆ H (endothermic) because energy is required to overcome the intermolecular forces that hold the solvent together 3.) Mixing the solute particles with the solvent particles *Note: Always negative (exothermic) because energy is released as the solute particles interact with solvent particles Enthalpy of Solution ( ∆ H soln)-the overall enthalpy change upon solution formation is the sum of the changes in enthalpy ∆ H soln = ∆ H solute + ∆ H solvent + ∆ H mix + + Heat of Hydration ( ∆ H hydration)- ∆ H solvent and ∆ H mix can be combine into this, this is the enthalpy change that occurs when 1 mole of the gaseous solute is dissolved in water *Note: ∆ H hydration is always largely negative (exothermic) for ionic compounds Sec 12.4: Solution Equilibrium and Factors Affecting Solubility: Dynamic Equilibrium-the rates of dissolution and recrystallization become equal Saturated Solution-A solution in which the dissolved solute is in dynamic equilibrium with the solid (dissolved) solute Unsaturated Solution-A solution containing less than the equilibrium amount of solute Supersaturated Solution-one containing more than the equilibrium amount of solute

Henrys Law-shows that the solubility of a gas in a liquid is directly proportional to the pressure of the gas above the liquid

Sgas = kHPgas Sgas = solubility of the gas usually in Molarity, kH = a constant of proportionality (henrys law constant) that depends on the specific solute and solvent and the temperature Pgas = partial pressure of the gas (usually in atm) Sec 12.5 Expressing Solution Concentration: Molarity (M)= moles of solute / liter of solution Molality (m)= amount of solute (Moles) / mass of solvent (kg) Percent by Mass= mass of solute/mass of solution X 100 ppm=mass solute/mass solution X 106 ppb=mass of solute/mass of solution X 109 Mole fraction= amount of solute (moles) / total amount of solute and solvent (moles) Sec 12.6: Colligative Properties: Vapor Pressure Lowering, Freezing Point Depression, Boiling Point Elevation, and Osmotic Pressure: Colligative Property-a property that depends on the number of particles dissolved in solution, not the type of particle Raoult’s Law-quantifies the vapor pressure of a solution Psolution = XsolventP ° solvent Psolution = the vapor pressure of the solution Xsolvent = mole fraction of the solvent P ° solvent = vapor pressure of the pure solvent at the same temperature Vapor Pressure Lowering ( ∆ P)- the difference in vapor pressure between the pure solvent and the solution ∆ P = P ° solvent -Psolution For a two component solution use: ∆ P = XsoluteP ° solvent *Note: this equation shows that the lowering of the vapor pressure is directly proportional to the mole fraction of the solute Ideal Solution-a solution that contains both a volatile solvent as well as a volatile solute *Note: both the solvent and the solute contribute to the overall vapor pressure of the solution For a two-component solution we can state: PA = XA P ° PB = XB P °

A B

The total pressure for a two-component solution is the sum of the partial pressures (equations above) Ptotal = PA + PB *Note: In a non-ideal solution, the solute-solvent interactions are either stronger or weaker than the solvent-solvent interactions Boiling Point Elevation-higher boiling point than the pure solvent Freezing Point Depression- Lower melting point than the pure solvent The amount that the freezing point decreases is given by: ∆ Tf = m × Kf ∆ Tf = the change in temp of the freezing point ( ℃ ) relative to the freezing point of the pure solvent (usually a + number) m = molality of the solution (Moles of solute / kg of solvent) Kf = the freezing point depression constant for the solvent For Water Kf = 1.86 ℃ /m The amount that a boiling point rises in solutions is given by: ∆ Tb = m × Kf ∆ Tb = the change in temp of the boiling point ( ℃ ) relative to the boiling point of the pure solvent

m = the molality of the solution (moles solute / kg solvent) Kb = the boiling point elevation constant for the solvent Osmosis-the flow of solvent from a solution of lower solute concentration to one of higher solute concentration Osmotic Pressure-the pressure required to stop the osmotic flow Π = MRT M = Molarity of solution (moles / Liter) T = temp in kelvin R = Ideal gas constant (.08206 L*atm/mol*K) Sec 12.7: Colligative Properties of Strong Electrolyte Solutions: Van’t Hoff Factor (i)-the ratio of moles particles in solution to moles of formula units dissolved

i=

moles of particles∈solution moles of formula units dissolved...