Exam 1 Study Guide PDF

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Hans Mentzen ISB 341D [email protected] Office Hours: M W 11:15-noon Tu F 1-3:00 PM Or Appt HW: Due Sun and Thurs nights (Incl. this Thurs) Owl Book + Mastery Section: recommended to read and do ques. before class LRC library 10th floor SI Sessions ExSEL (in library) DO OWL GO TO CLASS Topics We’re Covering ● Kinetics ● Equilibrium: foundation of other topics ● Solutions and Intermolecular Forces ● Acids and Bases - Buffers ● Thermodynamics ● Electrochemistry (smartphones, iPods, batteries, fuel cells) ●

Chap. 14 Kinetics ● ● ● ● ●

Chemical Kinetics Rates of Reactions Concentration Change Over Time Activation Energy Reaction Mechanisms

14.1 Intro to Kinetics ● Molecules must collide (come in contact) ● With enough energy to break and form bonds (collide hard enough) ● In a proper orientation that leads to atom rearrangement (specific atoms in molecules must collide with each other to break bonds) ● Can you control any of these factors? ○ Yes - more molecules, incr. concentration, temperature ○ Can’t control orientation ● As the concentration is increased, the number of collisions also increases ● High temp. = High KE = Higher rate of reaction ● Kinetics: Studying the rate of chemical reactions ○ Simply, how fast the reaction proceeds ■ Average Rate of Reac = △ [Concentration] / △ Time ■ △ = final - initial ■ In mol/L/time ○ The rate is most helpful in det. a reaction mechanism ○ The individual steps of how a reaction proceeds ○ Temp. = constant ● 2 types of reaction rates: ○ Average ○ Instantaneous (usually to find initial rate of reaction) ● Ex. Write the “rate law” for decomp. reac. ○ 2N2O5 → 4NO2 + O2 ○ Rate = -△[N2O5] / △Time ○ Final [N2O5] - Initial [N2O5] ■ Final is smaller ■ Answer is negative, compensate with negative in equation ● Rates over the course of a reaction change ○ Concentration decreases over time in rxn, therefore so does rate ○ Reactions take time, nothing is truly instantaneous ● Using stoichiometry - the relative average rate ○ The rate of rxn compared to reactants and products ○ = The decrease of reactants and the increase of products

■ Amol + Bmol -> Cmol + Dmol ■ Rate of decrease of reactants = (1/a) - △[A]/△t = (1/b) - △[B]/△t ■ Rate of decrease of products = -(1/c) - △[c]/△t = (1/d) - △[D]/△t ■ You can compare any of them to one another ● Concentration effects Instantaneous Rate of rxn ○ Need given experimental data ■ Exp 2/ Exp 1 M gives factor (k) ■ k= rate constant ● Relates rate proportional to concentration ● Concentration independent ● Temperature dependent ● Value of k is unique for every rxn ● Can manipulate units of k to make overall units mol/L/unit time ○ Oth order, k is in mol/L x units time ○ 1st order, k is in 1/time ○ 2nd order, k is L/mol x units time ○ Knowing units of k gives us overall rxn order ■ Apply that factor to rate ○ Rate law ■ Rate of rxn = k [Conc] ■ Rate of ran = k [A]^x[B]^y ■ No products, A and B are reactants ● Not all reactants may appear in the rate law, some other molecule not in reactants or products could though ■ The rate is proportional to the [ ] of each reactant raised to some power ○ Reaction orders (exponents) ■ Helps determine how concentration effects reaction rate ■ Must be determined experimentally ■ NOT from stoichiometric coefficients ● 2NO + Cl2 = 2NOCl ● Do experiments get concentrations of both reactants and rates ● (factor)^x = rate factor ● (Exp2/Exp1)^2 = rate2/rate1 ● NO rate is proportional to []^2, we say it is 2nd order in this reactant ● Cl2 rate is proportional to []^3, we say it is 1st order ● Rate = k [NO]^2[Cl2] ● Overal order is 3 (1 +2 ) ● Predicting 4th experiment, 8 ■ If you know overall rxn order, you know rate law ● Rxn order = 0th, exponent is 0, rate = k [R]^0

○ RO = 0, rate = k, rate of rxn does not depend on concentration of reactant ● Rxn order = 1st, rate = k [R]^1 ● Rxn order = 2nd, rate = k [R]^2 or k[R1][R2] ○ Integrated Rate Laws ■ An equation used to calculate [conc] at any point in time ■ Is different for all reaction orders ■ 1st Order Reactions ● Rate = -△ [R] / △ Time ● OR Rate = k [R] ○ Make equal to each other (integrate) ● Ln ([R]t / [R]0) = -kt ○ [R] = initial [R] ○ [R]t = [R] at any time ■ 2nd Order Reactions ● Now rate is proportional to [conc]^2 ● -△ [R] / △ Time = k[R]^2 ○ Integrate ○ 1/ ■ 0th order Reactions ● Now rate does not depend on the conc of a reactant ● -△ [R] / △ Time = k ○ Integrate ● [R]0 x [R]t =kt ● Graphing can be convenient, only 0th order gives a straight line ○ Solve for [R]t ○ [R]t = -kt + [R]0 ○ Y = mx + b ○ Reaction Half Lives (only responsible for 1st order) ■ [R]sub t ½ = ½ [R]sub 0 ■ Or 1 ● Increasing surface area increases rate of reaction ○ Exposing more molecules to react ● Increasing temperature increases the number of molecules with increased energy ● Activation Energy- Ea ○ The energy needed to overcome bond energies, endothermic process initially ■ = Transition state energy - initial reactant energy ○ A chemical hill that must be climbed for the rxn to occur ○ High Ea - A slow reaction ■ Most molecules do not possess enough kinetic energy to climb the hill

○ Low Ea - A fast reaction ■ Most molecules have enough kinetic energy to climb the hill ○ Graph that describes the reaction progress with respect to energy of the reactants and products is a reaction coordinate diagram ■ Transition State ● Energy maximum between products and reactants ● Non- isolatable, happens in microseconds ● Can only guess what the rxn looks like at this point ■ Enthalpy- Delta H ● Energy required to break old and form new chemical bonds ● Negative value indicates Exothermic ● Positive value indicates Endothermic ■ Reaction Intermediates ● Energy minimum between two transition states ● A very short-lived species ● Can be observed in special cases ■ Effect of Temperature ● As temp increases, starting molecular energy increases ● Able to lower the gap between the reaction energy and transition state ● Temp affects initial point of reactants on y-axis ● Not changing magnitude of Ea or transition state ● Arrhenius Equation ○ Relates k to frequency, energy, and temperature ○ As temp increases, k increases ○ K = Ae^(Ea/Rt) ■ A = frequency factor, number of atoms interacting in correct orientation ● Smaller the value, the number of collisions ■ Ea/Rt = Fraction of molecules with sufficient Energy to overcome Ea ■ R = Universal gas constant (check units are same) ■ T = kelvin temp ○ If k is known at only two temperatures ■ 2 equations can be written, 1 for each value of krc ■ As temp changes, only t changes, and k ○ Dont flip k1 and k2 or T1 and T2 ■ Reaction Mechanisms and Catalysts ● Effect of a catalyst on rate ○ Catalysts speed up reactions by lowering activation energy ○ Is not consumed in reaction, but it is involved

■ Reactant in step 1 and then product in one of next steps. none lost or gained ■ Does not appear in overall rxn, but could in rate law ○ Can change the selectivity of the reaction ■ Orientation of molecule slightly different ■ Catalysts help chemists choose orientations or isomers ● Mechanisms ○ How a reaction proceeds ○ Can be a 1 step reaction, called a concerted process ○ More often many elementary steps ■ Each step has a rate law and a k value, but the slow step is the only one that matter ■ Reaction can only proceed as fast as its slowest step ○ MOlecularity of Elementary steps ■ Molecularity = number of reactant molecules in an elementary step ■ Unimolecular = 1 molecule ■ Bimolecular = 2 molecules ■ Termolecular = 3 molecules , rare ○ A reaction is only as fast as its slowest step ○ Mechanisms are supported, never proven by data ○ A RDS (Rate-determining step) is the slow elementary step Rate law units = mol/Lxt-1 Rate constant units = varying

Chap. 15 Equilibrium ● The nature of the equilibrium state ○ All reactions are reversible and dynamic ○ When the RATE of forward and reverse reactions are equal ■ Equilibrium does not mean equal concentrations of reactants and products ○ Can go on in favor of reactants or products ■ Favoring reactants, product concentration small compared to reactants, most reactants have not reacted

■ Vise versa ○ Equilibrium is an expression and a constant (K) ■ Products over reactants ■ coefficients= exponents ■ Solids and pure liquids are never included ● The equilibrium constant, K ○ Never changes unless temperature changes, because temp changes rate ○ Rate of K1 does not have to equal K2 ● Reaction Quotient (Q) ○ Q= [Prod]^coeff./[React] ○ Determine is the reaction is at equilibrium ○ If Q does not equal K, not at equilibrium ○ If Q = K, at equilibrium ■ Concentrations remain constant ○ If Q < K, let reaction run in forward direction ■ Reactants must be converted to more product ○ If Q > K, shift backwards, ■ More reactants, less products ● Disturbing Equilibrium: Le Chatelier’s Principle ○ Disturbed by ■ 1) Changing temperature ● K changes ● Cannot predict the actual value of K ● Can only predict shift in reaction ○ If I know delta H of the reaction (endo or exothermic) ■ 2) Changing concentration ● Remove/add a reactant/product ● In all cases, K does not change ○ Remove products, Q>K - always shifts towards reactants (left) ○ Remove reactants, Q...