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Chem 1220 Notes

by: Phillip Fishbein

Chem 1220 Notes Chem 1220(Chemistry, Dr. Clark, General Chemistry)

Phillip Fishbein
GPA 3.722

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General Chemistry
Dr. Clark
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This 11 page Bundle was uploaded by Phillip Fishbein on Sunday March 6, 2016. The Bundle belongs to Chem 1220(Chemistry, Dr. Clark, General Chemistry) at 1 MDSS-SGSLM-Langley AFB Advanced Education in General Dentistry 12 Months taught by Dr. Clark in Winter 2016. Since its upload, it has received 99 views. For similar materials see General Chemistry in Chemistry at 1 MDSS-SGSLM-Langley AFB Advanced Education in General Dentistry 12 Months.

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Date Created: 03/06/16
Chapter 13: Properties of Solutions  Solution o homogenous mixture of solute and solvent o gases, liquids, solids o largest component is solvent, others solutes  Concentration o Molarity: moles/L of solution o Molality: moles solute/1 kg solvent  Energetics o Entropy  natural tendency towards mixing  increases with increased mixing o Spontaneity  depends on enthalpy and entropy o Enthalpy  endothermic separating of solute particles (ΔH Solute  endothermic separating of solvent particles (ΔH ) Solvent  exothermic mixing of solute and solvent particles (ΔH mix  overall is sum of three (ΔHsoln o Spontaneous endothermic solutions are caused by increased entropy  Intermolecular Interactions o Ion-dipole  Cations with larger charge and/or smaller radius have stronger ion-dipole interactions  stronger means greater magnitude of ΔH mix  Solubility o Saturated  Equal rates of dissolution and crystallization  Dynamic equilibrium  Saturated solutions at the same temperature have the same concentration o Unsaturated  Less solute dissolved than max o Supersaturated  More solute dissolved than when solution is saturated  Unstable  Adding more solute makes crystallization occur around added solute, known as “seed crystal”  Factors affecting Solubility o “Like dissolves Like”  Intermolecular Forces  ΔH mixis greater with similar IMF o Miscible  A solution where two liquids mix together/dissolve in any proportion  Never supersaturated o Immiscible  Liquids do not mix/dissolve significantly o Temperature  Solubility generally increases with temperature  Each substance’s solubility in different substances change at different rates with temperature o Solubility of Gases  Increases with increasing mass (Dispersion Forces)  Decreases with increasing temperature  Directly proportional to partial pressure above liquid  Henry’s Law: S = kg g  k = proportionality constant  S g solubility  P = partial pressure g  Increased pressure increases rate of gas going into liquid  Concentration o Solute per solvent o Dilute and concentrated (qualitative) o Mass Percentage  mass solute/mass solution * 100% o Mole Fraction  moles solute/moles solution o Molarity  moles/L solution o Molality  moles/kg solvent o Density  Use density to get volume for concentration calculation  Use total mass, not solute mass  Properties of Solutions o Colligative Properties depend on number of particles  Vapor Pressure  Adding solute lowers vapor pressure o Reduces rate going into gas  Raoult’s Law: P solution solvensolvent o X is mole fraction (watch for dissociation) 0 o P is original vapor pressure  Boiling Point  Increases with added solute (elevation)  Follows reasoning for vapor pressure  Freezing Point  Decreases with added solute (depression)  Decreases rate of freezing  Osmotic Pressure  π=(n/V)RT=MRT o Modification of Ideal Gas Law  May include van’t Hoff Factor  Pressure necessary to equalize fluid levels on both sides of a semipermeable barrier  Ion Pairing  Change in property is less than predicted with ion solutes have higher charges and/or higher concentration  Electrostatics briefly pair up, reducing number of particles  van’t Hoff Factor “i” added to boiling point and freezing point calculations for predictions Chapter 14: Kinetics  Factors that Affect Reaction Rates o Concentration of reactants o Temperature o Physical state of reactants o Particle size o Catalysts increase rates of reactions  Reaction Rates o Change of Concentration over Time (ΔM/Δt) o Reaction is fastest at the start of the reaction o May be expressed in terms of any species involved in reaction stoichiometrically  Reactants have negative rates  Products have positive rates  Concentration and Rate Laws o Use data found experimentally by varying one reactant over many trials  Find Rate Law o Rate Law  Rate = k[A] [B] n  m and n are reactant order for respective reactants (may be zero)  m+n is overall order for entire reaction  k = rate constant, calculated after finding rate law using experimental data  Large value: reaction is fast  Small value: reaction is slow o Integrated Rate Laws  Equations give concentration as a function of time  Zero Order:  [A] = [A] 0 kt  First Order:  ln[A] = ln[A]0– kt  Second Order:  1/[A] = kt + 1/[A] 0  Rates and Temperature o Relationship between temperature and rates described by Arrhenius’ Model and Collision Theory o Arrhenius’ Mod-Ea/RTation  k = Ae  Rate constant is dependent on temperature, therefore rate is dependent on temperature  As temperature increases, the constant increases  Eais activation energy for the reaction  Larger corresponds to slower reaction and vice versa  R is gas constant (8.314 J/molK) and T is temperature in Kelvin  A is Frequency/Orientation Factor  Reactants need to collide in the correct orientation  Large A means easier to collide in the correct orientation  Reaction Mechanisms o Rate law is explained by reaction mechanism o Explains how reactions proceed o Mechanisms formed by combining elementary reactions  Must be consistent with balanced equation and rate law o Number of molecules present in an elementary step is molecularity  Unimolecular, bimolecular, termolecular  Rare to see termolecular steps or higher o Elementary steps typically form intermediates which are used in a later step  Elementary steps are represented graphically by transition states  Activate complex present at top of transition states o Transition state with highest potential energy is the slow/rate determining step o Rate Law of elementary steps determined by molecularity o Typically have either a slow first step or a fast, reversible first step and a slow second step (at least in this class)  If slow first step, rate law for elementary step is overall rate law for reaction  If other, substitute in rate law for first step into second step rate law in place of the intermediate to get the overall rate law (include k values for forward and backwards reactions)  Catalysis o Catalyst affects rate of reaction without undergoing permanent chemical change  Affects rate constant/reaction mechanism o Typically lowers energy of activation o Sometimes increases frequency factor o Be able to identify Chapter 15: Chemical Equilibrium  Equilibrium is where forward and reverse reactions occur at equal rates  Equilibrium constant = products / reactants o Powers for each species come from coefficients in stoichiometric balanced chemical reaction o Written K cr K p o Affected by temperature only  Arrows going both ways in balanced equation  Can use ideal gas law to change between K andcK p  Haber Process: o ReactiΔn to form NH f3om N and2H 2  K p K (cT)  Pure substances are not included in equilibrium expression  Adding reactions multiply their equilibrium constants  Equilibria that occur entirely in the same phase are called homogeneous equilibria  Lo’ and behold, different phases are heterogeneous equilibria  If given equilibrium constant and initial condition, set up equilibrium constant expression with x. o Solve for x  Larger constant means reaction tends towards products  Smaller means reaction tends towards reactants  Predicting Direction of Reactions o Q is calculation for equilibrium constant using initial conditions  Q > K means reaction will move towards reactants  Q < K means reaction will move towards products  Le Chatelier’s Principle o Describes what happens to a system in equilibrium that is disturbed  Adding/removing reactants or products (concentration)  Adding reactants or removing products o Shifts equilibrium right  Removing reactants or adding products o Shifts equilibrium left  Changing pressure by changing volume (changes concentration by volume)  Decreasing volume o Shifts equilibrium to side with fewer moles of gas  Increasing volume o Shifts equilibrium to side with more moles of gas  Temperature changes Think of heat as a reactant/product of the reaction  Exothermic or Endothermic?  Exothermic o Increasing temperature shifts equilibrium left o Decreasing temperature shifts equilibrium right  Endothermic o Increasing temperature shifts equilibrium right o Decreasing temperature shifts equilibrium left Chapter 16: Acid-Base Equilibria  Different definitions of Acids & Bases o Arrhenius Acid  Produces H in water (HCl) + +  H (aq) typically clusters into H3O (aq) (Hydronium) o Arrhenius Base  Produces OH in water (NaOH) o Bronsted-Lowry Acid  Donates a proton to another substance (NH ) 4+  Has a conjugate base (NH ) 3 o Bronsted-Lowry Base  Accepts a proton (NH )3  Has a conjugate acid (NH ) + 4 o Lewis Acid  Electron-pair acceptor (BF 3 o Lewis Base  Electron-pair donor  Hydration of ions can be considered acid-base through Lewis definition  Weak acid/base is defined by small K whCn written as reactant o Large K Cf written as product o Only acid-base reactions written in equilibrium  Strong Acids o HCl, H 2O 4 HNO ,3HClO ,3HClO , 4Br, HI o All others are typically weak o Typically not written in equilibrium  Strong acid has weak conjugates  Weak acids have strong conjugates  Hydroxide is not the strongest base in existence (tip for future chemistry classes)  Acids are typically written with H on one side  Autoionization of Water - + o Water sel-14onizes into OH and H O3 o K C 10 = K W o K Wncreases with temperature  Acid-Base Scales o pH  -log[H ]  7 for water o pOH  -log[OH ]  7 for water o pH + pOH = 14 at 25°C  This sum changes with temperature o Solutions are neutral when pH = pOH  Strong acids and bases are strong electrolytes  Strong Bases - o Alkaline and alkaline earth metals combined with OH - o Watch for number of OH that dissociate from the ion!!! o Highly charged ions may act as an acid and affect the pH (charge of 3+)  Weak Acids o Generic Form  HA + H 0  H O + A - 2 + 3-  K a [H O3][A ]/[HA]  pK a -log(K ) a o pH affected by presence of conjugate salts (EQUILIBRIUM!!!) o Polyprotic Acids  Have different dissociation constants for each proton that dissociates  Weak Bases o Generic -orm -  A + H O2 HA + OH  K b [HA][OH ]/[A ] -  pK b -log(K ) b o pOH affected by presence of conjugate salts (EQUILIBRIUM!!!) o K b1corresponds to the last K aalue for conjugates of polyprotic  Going Between K and Ka b o K a K /w b o K b K /w a o K w K *a b  Strong vs Weak o Bronsted-Lowry  Favorable for a strong acid to donate a proton and unfavorable to add a proton to its conjugate  Favorable for a strong base to pick up a proton and unfavorable to donate the proton from its conjugate o Energy  Easier to break apart strong acids and add a proton to a strong base  Oxyacids (O-Chem) o Acids containing one or more O-H bonds o Same number of O-H groups as oxygen atoms  Acid strength increases with increasing electronegativity of the central atom  Electronegativity spreads charge across molecule away from negative oxygen o Increasing strength with more oxygen atoms than OH groups  Helps spread charge across molecule to prevent rebonding  Organic Weak Acids o Has a proton that can be removed o Has carboxyl? group (COOH)  Leads to resonance with charge position, leading to more stable conjugate o Lower removal energy leads to stronger acid Chapter 17: Additional Aspects of Aqueous Equilibria  Henderson-Hasselbalch Equation - o pH = pK a log([X ]/[HX]) = pK + aog([base]/[acid]) o pOH = pK +bog([HX]/[X ])= pK + lbb([acid]/[base]) o Used for dealing with buffer solutions  Adding the conjugate base to an acidic solution of a weak acid  e.g. acetic acid with acetate  Adding conjugate base to acidic solution raises pH and vice versa for base and conjugate acid o EQUILIBRIUM!!!  A strong/weak acid/base interaction is a buffered solution  Buffers resist pH change o Optimal buffer has [base] = [acid]  pH is approximately pK a o Resists pH change due to the addition of both acid and base Chapter 17: Titrations  Basically acid-base equilibria and solubility investigations  K sptting introduced  Typical Titration Setup (Acid-Base) o Strong base in buret o Acid in beaker  Break titrations into 4 regions o Initial pH o Between initial pH and equivalence point o Equivalence point – equal moles of titrant and analyte (base and acid, respectively) o Post equivalence point  Weak Acid and Strong Base o Only the same calculations for pH in post equivalence o Use equilibriums constants and Henderson-Hasselbalch equation  K bor equivalence point  Indicators o Change color based on acid or base form of indicator species o


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