Chemistry 112 ACS Final

Chemistry 112 ACS Final Exam at 4 pm in ECTR 118 and you will be done by 6 pm Stop at question 56 Topics covered are below.honestly, the best way to study is to read through your notes multiple times making sure you know everything listed below and get a good night s sleep Solutions (8 Questions) 1. Intermolecular Forces (Dispersion Forces Induced-Dipole---Dipole-Dipole---Ion-Dipole-- Hydrogen Bonding) A. What do each require to exist (What causes the force?) B. What do the forces look like pictorially C. Which is stronger, strongest? D. Be able to draw Lewis Structures? E. Is there any way a compound with only a weaker force be overall stronger than a compound containing a stronger force F. How do intermolecular affect boiling point, melting point, vapor pressure and critical temperature G. What is miscibility and how are intermolecular forces involved 2. Phase Diagrams A. Where each phase is located B. What the lines represent C. Where is the normal freezing point, melting point, boiling point, vaporization point D. What is the triple point and critical point 3. Solutions A. What is a solution and how does miscibility factor in? B. Know the difference between an unsaturated, saturated and supersaturated solution C. Dilutions (M1V1 = M2V2) D. Ion concentrations (know your polyatomic ions) C. Expressing concentration 1. Molarity (M) moles solute/ 1 L of solution 2. molality (m) moles solute/ 1 Kg of sovent 3. Mass percent (%) grams solute/ 100 g solution 4. Mole fraction (χ) moles solute/ 1 mole of solution 5. Density is necessary to convert between Molarity and all of the others 6. The numerator and denominator are separate entities that only come together when Expressing a particular form of concentration 4. Henry s Law: [gas] = K H *Pgas (This equation will be given to you) 5. Colligative Properties Dependent on the particles in solution (the more particles the stronger the effect) A. Freezing point depression Δt = K f mi B. Boiling point elevation Δt = K b mi C. Osmotic pressure increases Π = TRiM D. Vapor Pressure Lowering P (solution) = χ solvent P (pure) Equilibrium (12 questions covering all of the following) General Equilibria: 1. Equilibrium is reached when the rate of the forward reaction = the rate of the reverse Reaction. Concentrations ARE NOT equal! Rate constants ARE not equal 2. Equilibrium constants are dependent on temperature 2. Know how to write an equilibrium equation. Solids and Liquids are NOT included 3. Know how to convert between Kp and Kc: Kp = Kc(RT) Δn where Δn is equal to the

Moles of gas products the moles of gas reactants 4. What does the magnitude of K mean? 5. Manipulations of K A. If you reverse a reaction you take the reciprocal of K B. If you multiply a reaction through by a value raise K to that value C. If you add reactions together you multiply their K values 6. Use ICE tables to solve for equilibrium constants and for equilibrium concentrations or partial pressures (Know how to use the quadratic formula) A. When solving for K always raise the concentrations by the stoichiometric Coefficient in the equilibrium equation B. When solving for x/s to get equilibrium concentration use the stoichiometric coefficient as a multiplier C. Think of words in terms of gone or still there so you get the right equilibrium concentration 7. Le Chatlier s Principles (When an equilibrium is disturbed it will find a way to reestablish equilibrium) A. Adding product or Removing reactants shift a reaction to the Left B. Adding Reactant or removing product shifts a reaction to the Right C. Increasing Pressure or decreasing volume shifts to the side with the fewer gaseous Compounds D. Decreasing Pressure or Increasing volume shifts to the side with more gaseous Compounds E. An increase in temperature favors and endothermic reaction (+ΔH) F. A Catalyst does not cause a shift 8. Reaction Quotient (Q) The ratio of products to reactants A. When Q < K reaction proceeds to the right B. When Q = K reaction is at equilibrium C. When Q > K reaction proceeds to the left Solubility Equilibria 1. Know polyatomic ions and charges 2. Know solubility rules (Or know how to use Ksp values to determine if a compound is insoluble 3. Always written as solid going into solution 4. solubility is in terms of grams/ ml 5. Molar solubility (s) which is used to solve for Ksp is in M (mols/ L) 6. Use ICE tables to solve for Ksp and molar solubility 7. What Q means: A. Q = Ksp the system is at equilibrium and is saturated B. Q < Ksp the system is not saturated and a ppt will not form C. Q > Ksp the system is supersaturated and a ppt will form D. When determining if a ppt will form solve for Q as you would for K using the concentrations given in the problem for the ions involved in the formation of the ppt. and then compare to Ksp 8. Common Ion Affect: A. A common ion is an ion that comes from more than one source and lowers the molar solubility of the insoluble salt B. You are only concerned with the ions involved in the formation of a ppt. However, this ion is found as part of a salt that has been added to the solution C. The concentration on the common ion serves as the initial concentration in the ICE table Acids and Bases 1. Definions a. Arrhehius b. Bronsted-Lowery (in terms of H + ) c. Lewis (in terms of electrons) 2. Know strong acids and strong bases 3. Be able to show how a reaction will take place using arrows 4. Be able to identify conjugate pairs 5. Be able to compare strength of weak acids using Ka and pka 6. Be able to compare strength of weak bases using Kb and pkb or pka of it s conjugate acid 7. Amphoteric substances 8. Self Ionization of water: [OH - ][H 3 O + ] = Kw at 25 C Kw = 1 x 10-14

9. ph A. ph = - log [H + ] conversely [H+] = 10 -ph B. poh = - log [OH - ] conversely [OH - ] = 10 -poh C. ph + poh = 14 D. ph < 7 solution is acidic E. ph = 7 solution is neutral F. ph > 7 solution is basic 10. Weak Acid and Weak base problems For weak acids the smaller the pka (the larger the Ka) the stronger the acid and the weaker it s conjugate base! For conjugate pairs Ka x Kb = Kw pka + pkb = 14 General Equation HA(aq) + H 2 O(l) H 3 O + + A - Ka = [H 3 O + ] [A - ] [HA] For weak bases the smaller the pkb (the larger the Kb) the stronger the base and the weaker the conjugate acid General Equation B(aq) + H 2 O(l) HB + + - OH Kb = [HB + ] [ - OH] [B] A. If you make OH - you need a Kb B. If you make H 3 O + you need a Ka C. Use ICE tables to solve D. pka = -log Ka conversely Ka = 10 -pka E. percent ionization = ( x/ Initial Conc) * 100 11. Salt solutions Salts can be neutral, acidic or basic How to tell Cation Anion Makes the solution acidic if not Group 1 or 2 metal Make the solution Basic if not Cl -, Br -, I -, NO - 3, SO 2- - 4,ClO 4 Are conjugate acids of weak bases Area conjugate bases of weak acids Cations: Metal cations such as Al 3+ which are small with a high charge will form a hydrate usually a hexahydrate and then undergo hydrolysis Al 3+ 3+ + 6 H 2 O Al(H 2 O) 6 + H 2 O Al(H 2 O) 5 (OH) 2+ + H 3 O + General Equation for cations derived from weak bases Cation + (aq) + H 2 O(l) H 3 O + (aq) + parent base(aq) (need Ka) (Where to get it from) To solve: You must get the Ka by finding the parent base and Ka = Kw/Kb Ka = [H 3 O + ] [parent base] [Cation] Solve for [H 3 O + ] and take the negative log Anions: General Equation for anions derived from weak acids Anion - (aq) + H 2 O(l) parent acid(aq) + - OH(aq) (Where to get it from) (need Kb) To solve: You must get the Kb by finding the parent base and Kb = Kw/Ka Kb = [ - OH] [parent acid] [Anion] Solve for [ - OH] and take the negative log then subtract from 14 A. What is affecting the ph of the solution the cation or the anion? B. Write a hydrolysis reaction for the ion affecting the ph of the solution

C. Determine if you need a Ka or Kb see A and B in Weak acid/ base (Always opposite of what s given) D. For conjugate pairs: Kw = Ka*Kb; pka + pkb = 14 12. Polyprotic Acids (H 2 A) A. loose H + in a stepwise fashion B. use multiple ICE tables the Equilibrium concentrations of H 3 O + and HA - are the initial concentrations in the second ICE table C. ph is determined from the first step 13. Buffers A. resists changes in ph with the addition of small amounts of strong base or acid B. made from a weak acid and it s conjugate base or a weak base and it s conjugate acid C. In making a buffer you want to choose a conjugate pair that has a pka closest to the desired ph D. When making a base from a polyprotic acid which K do you use? E. Henderson-Hasslebach equation: ph = pka + log Base/ Acid F. When [base] = [acid] in the Henderson Hasslebach equation the ph = pka G. Since you are often combining solutions it may be helpful to use an SRFC table even if you are only using the S row H. When adding a strong base to the buffer subtract it s moles from acid and add them to base I. When adding a strong acid to a buffer subtract it s moles from base and add them to acid 14. Titrations A. Be able to find volume of titrant to reach equivalence point of what you are titrating B. Titration of a strong acids and strong bases only dealing with OH - and H + so ph s based on concentrations of [H + ] or [OH - ] in solution C. Titration of a weak acid/ base with a strong acid/ base i. Be able to draw and label the curve ii. Find volume of strong needed to reach the equivalence point iii. Before strong is added you only have a weak (ICE table) iv. You have a buffer solution from the moment you add strong until equivalence point (SRFC table-----buffer SOLUTION) v. At equivalence point: SALT (you have just conjugate so you need to change from one K to the other using Kw) Use both SRFC and ICE vi. Past the equivalence point ph is dependent solely on the strong and need to find the concentration of the strong 15. Indicators are chosen by choosing an indicator with a Pk I closest to the ph of the equivalence point KINETICS (12 questions) 1. Expressing rate of appearance, rate of disappearance and rate of reaction in terms of reactants and products 2. When going from rate of one species to another use the mole to mole ratio 3. When finding the rate of reaction Remember to divide by stoichiometric coefficient 4. Finding rates from a graph----know when to use a tangent line 5. When given the mechanisms of reactants be able to identify intermediates and to give the rate law for the overall reaction 6. Using experimental data, be able to determine the rate law for the reaction using relative concentrations and relative rates don t forget the units for the rate constant k. 7. Be able to give the order of the reaction as well as order with respect to reactants 8. Zero order reactions A. know equations for ½ life and the integrated rate law B. Which plot gives you a straight line 9. First Order reactions A. Know what plot gives you a linear relationship between concentration and time B. Be able to solve for ½ life C. Know the integrated rate equation be able to use it for fraction remaining 10. Second order reactions

A. know equations for ½ life and the integrated rate law B. Which plot gives you a straight line 11. Factors that affect Reaction rates A. Concentration of reactants B. The use of a catalyst C. Temperature 12. Collision Theory A. What is the importance of molecular orientation? B. What would an increase in reactant concentration do? C. What would an increase in temperature do? 13. Transition State Theory A. Potential energy diagrams (how does an endothermic reaction differ from and exothermic reaction?) B. What is activation energy (E a )and how is it measured? C. Where is the heat of reaction considered? D. Be able to find the Ea for both the forward and reverse reaction 14. Affect of a catalyst A. Be able to identify a catalyst B. What is the difference between a heterogeneous and homogeneous catalyst 15. Affect of temperature on reaction rates---------must know the Arrhenius Equation A. k = Ae -Ea/RT B. ln k 2 = - Ea k 1 R 1-1 T 2 T 1 Thermochemistry (12 Questions) 1. Equations to know A. ΔSrxn, ΔHrxn and ΔG reaction are found by taking the sum of the products and subtracting the sum of the reactants from appendix 3 B. ΔS univ = ΔS system + ΔS Surroundings C. ΔS surroundings = -ΔH sys /T surroundings D. ΔG = ΔH - TΔS E. ΔG = ΔH - TΔS F. ΔG = ΔG + RTlnQ G. ΔG = - RTlnK 2. Enthalpy, Entropy and Gibbs Free Energy are all state functions which means they can be found by subtracting the standard enthalpy, entropy and Free energy of formation of reactants from the products. In chemical equations don t forget to take the stoichiometric coefficients into account. 3. Three laws of thermodynamics a. First law: The total energy of the universe is constant b. Second law: The total entropy of the universe is always increasing c. Third law: The entropy of a pure, perfectly formed crystalline substance at absolute zero is zero 4. Entropy increases as a solid goes to a liquid and as a liquid goes to a gas. Also entropy increases with the size of the molecule 5. Conventions for standard state (ΔH, ΔG and ΔS ) gas = 1 atm; liquids and solids are pure; Soln = 1M 6. During a phase change only ΔS sys = ΔH sys /T where T is temperature in Kelvin 7. The units for ΔG and ΔH are KJ/ mol and the units for ΔS are J/mol K 8. ΔSuniverse = ΔS sys + ΔS surrondings where ΔS surroundings = -ΔH system / T 9. ΔG = ΔH system TΔS system (How did we get this equation? And how does temperature affect it?) 10. ΔG = ΔH system TΔS system 11. ΔG = ΔG + RTlnQ at equilibrium ΔG = 0 and Q = K therefore: ΔG = -RTlnK a. R = 8.314 J/ mol K and T is in Kelvin b. As K increases G becomes more negative c. An increase in temperature favors the endothermic reaction (+ΔH)

12. Recap: Reactions are spontaneous in the forward direction when: a. K is large b. ΔG is negative c. ΔS of the universe is positive d. E is positive Electrochemistry (12 questions) 1. Equations to know A. ΔG = -nfe cell B. E = E - (RT/nF) ln Q at 25 C E = E - (0.0257/n) ln Q C. E = (RT/nF) ln K at 25 C E = (0.0257/n) ln K D. n = IT/ F 2. What is oxidation? What is Reduction? They go together hence REDOX reaction 3. Know how to assign oxidation numbers and how to balance REDOX equations 4. Standard Reduction Potential E (V) To find E cell or E rxn it is E red + E ox (flip the sign for the ½ rxn that gets oxidized) A. If the reaction is spontaneous in the forward direction E cell will be positive B. If the reaction is not spontaneous in the forward direction E cell will be negative 5. Galvanic cells are spontaneous (E is positive) A. The cathode is positive and is where reduction takes place B. The anode is negative and where oxidation takes place C. Electrons flow from the Anode to the Cathode D. A salt bridge allows for the flow of ions so there won t be a charge build up 6. What is Conventional cell notation? 7. ΔG = -nfecell where n is the moles of electron and F is faraday s constant 96500 J/mol V 8. The nerst Equation: Ecell = E cell - (0.0257/n ) ln Q 9. When at equilibrium Q = K and Ecell = 0 therefore E cell = (0.0257/n) ln K 10. Electrolysis (Ecell < 0) A. The anode and cathode remain the same but the charges on the electrodes change B. The electrolysis of water produces H 2 at the cathode and O 2 at the anode 11. Used in electroplating: n = It / F where n is moles of electrons and It is amps and t is time in seconds. It = Columbs