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UB CHE 102 - Final Exam Study Guide

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Chem 102 1st EditionFinal Study Guide- Ch. 13-23 (no 22)Chapter 13A. Units of concentration:Mass%= (Mass of component/ total mass of solution) x100Mole fraction (X)= moles of solute/ total moles of solventMolarity (M)= moles of solute/ liters of solutionMolality (m)= moles of solute/ Kg solvent Moles (mol)= mass(g)/ MMB. Energies of Solution1. Separation of solute molecules- endothermic solute= small amount2. Separation of solvent molecules- endothermic solvent= large amount3. Formation of solute-solvent interaction- exothermicEnthalpy Formula- ΔH= ΔH1 + ΔH2 + ΔH3C. Energy Changes & Solution ProcessHeat= specific heat x mass (solute and solvent) x ΔTRULE= Like dissolves like… such as, non-polar dissolves non-polar. Water is non-polar therefore it will be dissolved by the most non-polar substance. D. Hydrophilic vs. hydrophobic-Hydrophilic- “likes water” Has to have presence of ions, dipole moment and LDF. E. Boiling-Point Elevation ΔTb=i  m  Kb K=constant m=molality i= van’t Hoff factorThe higher the ΔTb the higher the boiling point. F. Freezing Point DepressionΔTf= i x m x Kf K= constant m=molalityThe higher the ΔTf the lower the freezing point. G. Osmosis and Osmotic Pressure/ Vapor pressureπ= MRT π= Osmotic pressure M=Molarity R=constant, .08206 T=Temp (K)π is measured in atm. Raoult’s Law- Vapor pressure equation Pa= Xa Pa° Pa- vapor pressure with solute Pa°-vapor pressure without solvent Xa- Mole fraction of A in solution H. Colligative Properties of ElectrolytesSome ions do not dissolve to their full extent. Van’t Hoff factor, i= moles particles in solution/ moles solute dissolvedi= 1 for non-electrolyteChapter 14A. Intro. To Reaction RatesAverage Rate= -ΔH[x[/ Δt= Δ[y]/Δt x= reactants (disappearance) y=products(appearance)Rate=k(reactant 1) ^m (reactant 2) ^n m and n refer to order k= rate constantZero order- Change in concentration of reactant produces no effectFirst order- Doubling concentration cause rate to doubleSecond order- Doubling concentration results in a 2^2 increase in rateFirst- Order Reactants: Ln [A]t= -kt + ln [A]0Half-Life: t1/2= -(ln1/2)/k= .693/kSecond- Order Reactants: 1/ [A]t =kt+ 1/[A]0Temp increase, rate of reaction increasesB. The Collision ModelIn order for molecules to react they must collide. The greater the number of collisions the fasterthe rate. The more molecules present, the greater the probability of collisions. The higher the temp, the more energy available to the molecules, therefore the faster the rate. C. Activation EnergyMolecules must present a minimum amount of energy to react, in order to break bonds. Activation Energy is the minimum energy required to initiate a chemical reaction. Arrhenius Equation: k=Ae ^(-Ea/RT)D. CatalystsOperate by increasing the number if effective collisions, increasing reaction rate. Intermediates may be added to these reactions. These are molecules that get cancelled out of the first equations and do not show up in the final equation. Chapter 15A. Chemical EquilibriumThe point at which the concentrations of all species are constant. The opposing rates are equal, forward and reverse reaction. These reactions are represented by a double arrow. Equilibrium constant expression:      baqpcKBAQP Kc is the equilibrium constant. The Kp is the equilibrium constant for reactions involving gases:     baqpPPPPPKBAQPChapter 16- ACIDS & BASES-Arrhenius Acid- Substances that produce H+ in water, Increase H+ (aq only)-Arrhenius Base- Substances that produce OH- in water, Increase OH- (aq only)-Bronsted Acid- Donates proton-Bronsted Base- Accepts Proton, must have lone pair to accept it. -Lewis Acid- Electron pair acceptor-Lewis Base- Electron pair donor*Not all Lewis are Bronsted acid/bases, but all Bronsted are Lewis acid/bases. -Conjugate Base- Removing a proton from an acid. HCl Cl- = AcidConjugate Base-Conjugate Acid- Adding a proton to a base. NH3 NH4 = Base Conjugate Acid-Kw= [H+][OH-] = 10-14 = (Ka)(Kb)-pH= -log [H+] and pOH=-log[OH-] *if given the pH need to use antilog to find M. For example what is M if pH=2.73, M=10-2.73. *-pOH+pH=14= pKw-Strong acids- HCl, HBr, HI, HClO3, HClO4, HNO3, H2SO4. 100% dissociation, therefore equilibrium=0. -Strong acids have weak bonds, large electronegativities (polar), and stable conjugate bases. -Weak Acids- Everything not on list. Partial dissociation, reversible, therefore, equilibrium exists.Kq>1- Strong acid, easy to lose a proton. Kq<1 Weak acid-Strong Bases- LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2, Sr(OH)2, Ba(OH)2. 100% dissociation, equilibrium=0. -Weak Base- Partially dissociate or reacts with water to form OH-. -Polyprotic acid- More than 1 donatable (reactive) proton. Only count the front 2 H’s, H2C6H6O6 has 2 reactive protons. Always easier to lose the 1st H because Ka is larger. H2SO4 is the only strong polyprotic acid. -% Ionization or Dissociation = (H equilibrium/ HA initial ) x 100 -Salts- Strong electrolytes=100% dissociation, can be acidic, basic, or neutral.-Anions=bases, negative ion attracted to H+ to form OH-. -Cation=acids, positive ion attracted to OH- to form H+. -Neutral salt= strong acid+ strong base-Acidic salt= strong acid+ weak base-Basic salt= weak acid+ strong baseChapter 17- ADITIONAL AQUEOUS EQUILIBRIA-Common Ion effect- Shift in equilibrium caused by adding an ion involved in the equilibrium.-Dissociation of a weak electrolyte decreases when a strong electrolyte containing a common ion is added. -Buffers- Weak acid/base and conjugate acid/base, resists a change in pH when small amount ofOH- or H+ is added. pH=pKa+ log [base/acid] Use ice table with Moles-Buffer capacity- Amount of acid/base added before a significant pH change occurs.-Titration Curve- pH vs. volume of titrant.Equivalence point- Straight vertical lineColor change- Upper curve, top of vertical line.Greatest Buffering Capacity- Halfway point to the equivalence point.-We want to use an indicator that has around the same pH as the equivalence point.-Precipitation Reactions- Ions, when combined, form insoluble salts.-Ksp- Equilibrium constant for a slightly soluble (mostly insoluble) ionic compound—No Solid orliquids!!!!!-Solubility- Grams dissolved to form a saturated solution. Grams/L= unit-Molar solubility- Moles dissolved to from a saturated


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