CHE 102 1nd EditionExam # 1 Study Guide Chapters: 13, 14 and 15Chapter 13 (Properties of Solutions)Three types of Intermolecular Interactions:- Solute-Solute: Separation of solute molecules ; Endothermic Reaction - Solvent-Solvent: Separation of solvent molecules ; Endothermic Reaction- Solvent-Solute or Solvation/Hydration: Formation of solute-solvent interactions ; Exothermic Reaction Increase Interaction: ↑Charge and ↓Size- Solutions only form when the ∆H is negative. [Spontaneous and Exothermic]∆Hsolution ⁼ ∆Hsolute + ∆Hsolute + ∆Hsolute- If ∆H is positive . . . . .Entropy (∆S) or an increase in disorder and chaos compensates for an Non-spontaneous /Endothermic reaction.- “LIKE dissolves LIKE” Example: Polar/Ionic solutes will dissolve in Polar solvents. [NaCl in H2O]- Increasing the carbon chain, Decreases the solubility of liquids. Example: C5H12O (butanol) is MORE soluble than C4H10O (pentanol) because butanol has a longer carbon chain. Exception: C3H8 (propane) is MORE soluble than CH4(methane) because propane has more surface area.- Solubility of Gases Increases: ↑ Molecular Mass and ↑ Polarity- HENRY’S LAW: Sg ⁼ k P∙ g Increasing the partial pressure of the gas above the solution pushes the gas closer to the liquid for a stronger attraction. This allows for more dissolving and increases the solubility of the gas.Sg is in units of molarity (mol/L) ; Pg is the partial pressure of the gas above the solution. The Henry Law constant, k, units are mol/L atm.∙- Determining the Concentration of Solutions (Quantitatively) [NO UNITS]Mass %: Mass of solute ÷ Mass of solution ∙ 100Parts per Million (ppm): Mass of solute ÷ Mass of solution ∙ 1,000,000Part per Billion (ppb): Mass of solute ÷ Mass of solution ∙ 1,000,000,000Molarity(M): moles of solute ÷ Liters of solutionmolality (m): moles of solute ÷ Kilograms of solution- RAOULT’S LAW: Psolution ⁼ Xsolvent ∙ P⁰solvent Obeys Ideal solutions with Low solute concentrations, similar size and IMF’s as the solvent.Adding a nonvolatile solute →More ions dissociated in solution → Lower Vapor Pressure- Van’t Hoff Factor (i): # of particles dissolved [Beware of Polyatomic Ions] Example: K2SO4 has a Van’t Hoff Factor of 3. NaNO3 has a Van’t Hoff Factor of 2.- Boiling Point: Adding a nonvolatile solute → Increases Boiling PointA pure solvent will have a lower boiling point than a concentrated solution.Boiling Point Elevation: ∆Tb⁼ i K∙ b m ∆T∙ b = Tpure - Tsolution- Freezing Point: Adding a nonvolatile solute → Decreases Freezing Point Starts at the Triple Point; point where the solid, liquid and gas phases are equal or at equilibrium. A pure solvent will have a higher freezing point compared to a concentrated solution. Freezing Point Elevation: ∆Tf ⁼ i K∙ f m ∆T∙ f = Tsolution - Tpure∆T must always be POSITIVE!Boiling Point of Water: 100⁰C Freezing Point of Water: 0⁰C- Osmotic Pressure: Pressure that stops the movement of liquid ∏= i MRT R(Ideal Gas Constant)=.8206 L atm/mol K∙ ∙ ∙ Movement of the solvent will always go from low to more concentrated solutions!Chapter 14 (Chemical Kinetics) - Rate or the “speed” of a Reaction increases when . . . Temperature increases, Concentration increases, A catalyst is Present and similar physical states. Example: Gaseshave a far range oppose to solids which are more ridged.- Writing Reaction Rates: 3X + 2Y → 5Z Units: ∆concentration(final –initial)/∆time or M/sRate of Appearance (Products): ∆[Z]/5∆timeRate of Disappearance (Reactants): - ∆[X]/3∆time and - ∆[Y]/2∆timeObserve Coefficients!- ∆[X]/3∆time = - ∆[Y]/2∆time = ∆[Z]/∆time Rates of the reaction equal.- Rate Laws: equation of reactant concentrationsRate=k[A]X[B]Y[C]z Exponents are the “reaction orders” of the reactants. Exponents are NOT derived from the coefficients of the reaction!!Overall Order: SUM of all the reaction orders. (x+y+z)Rule out intermediates because of their low concentarations.- Half-Life: time it takes for a substance to decrease to half its original concentrationFirst Order: t1/2 = ln[2]/kSecond Order: t1/2 = 1/ k[A]0 [A]0 = Original Concentration- Average Reaction Rate VS. Instantaneous Reaction Rate Slope between 2 points on a curve Slope of a line Tangent to 1 point on a curve Used for Short Periods of time A Particular time- Rate Law Constant (k) Units: First Order: s-1Second Order: M-1s-1- Mathematical and Graphical Methods to determine Reaction OrdersMathematicalUsing a table, choose two experiments. Observe the change in reactant concentrationsand set it equal to the change in rate concentration. Solve for the exponent. If you have irregular or more complex concentrations, choose two experiments where one reactant remained constant and the other reactant changed in concentration. Graphical Integrated Rate Law relates concentration and time. ln[A]t = -kt +ln[A]0 First Order Reaction will have a straight line with the graph ; ln[A] vs. time. Second Order Reaction will have a straight line with the graph ; 1/[A] vs. time. Faster Reactions have lower Activation energies (EA) or lower “hills”. - Arrhenius Equation: Combines the Collision theoryLn[k2/k1] = Ea/R ∙ [(1/T1) - (1/T2)] is used to relate k to two different temperatures.Example of Molecularity: 2H2 + O2 → 2H2O is a termolecular reaction. Unimolecular reactions have one molecule and Bimolecular reactions have two molecules.The slow step is the Rate-Determining Step for the overall reaction!- Heterogeneous Catalyst are in different phases than the reactants so…. Adsorption occurs.Adsorption is when the molecules bind to the surface due to the difference in phase.Chapter 15 (Chemical Equilibrium)- What is Equilibrium? The rate of the forward reaction and the rate of the reverse reaction are EQUAL. - Kc (EQUILIBRIUM Constant for Concentration) A+2B 3C Kc = [C]3/[A][B]2ONLY use molecules in the gas or aqueous phase because their concentrations vary.Observe Coefficients!Example:3Fe(s) + 4H2O(g) Fe3O4(s) + H2(g) Kc = [H2]/[Fe]3 Kc is larger than 1, Products are favored and Equilibrium lies to the right of the