Chem 2 Notes 1Chemical kinetics - focuses on the rate at which chemical process occur, and on the reaction mechanismFactors affecting reaction rates:1. Physical state of the reactants gases faster than liquids(soln) faster than solids Stirring liquids and grinding solids will increase rr2. Concentration of reactants The greater the concentration of reactants, the higher the rr3. Temperature At higher temperatures, molecules have more KE which makes them move and collide more often with greater energy4. Presence of a Catalyst or an Inhibitor Catalysts speed up a reaction without being consumed Inhibitors slow down the reaction without being consumedReaction Rates: Units of Change in Concentration/time, M/s is the most common Defined as positive or 0 Rate =+-( [X]@t2 - [X]@t1 )/(t2 - t1) = +-(change in [X]/change in t) Use + for products, - for reactants The average rate of reaction over each interval is the change in concentration divided by the change in time Average rates decrease over time in a reaction Use a graph and lame 2 dimensional calculus to find the slope of a tangent line to the graph. Instantaneous rate - d[A]/dtReaction Rates with Stoichiometry: If ratio in a balanced equation of molecules is 1:1, then the magnitudes of their rates are equal If ratio is not 1:1, then the magnitudes of their rates are proportional to said ratio: 2HI(g) -> H2(g) + I2(g) Rate = -1/2 * [HI]/t = [I2] / t aA + bB = cC + dD -1/a * rate of A = -1/b * rate of B = 1/c * rate of C + 1/d * rate of DRate Laws: Shows the relationship between reaction rate and the concentrations of reactants Form of: Rate = k[reactant1]m * [reactant2]2 Overall reaction order can be found by adding the exponents With complex reactions, you may have to find the rate law experimentally Order - Units of k: 0 - M/s 1 - 1/s 2 - 1/(M*s) 3 - 1/(M2*s)Zero-Order Differential Rate Law: A-> products rate = k[A]0 = k*1 = k Average rate = - change in [A]/ change in t =~ k Instantaneous rate = -d[A]/dt = kZero-Order Integrated Rate Law: Integrate rate law to give: [A]t = [A]0 - kt This is a simple linear decay function, this is sometimes observed in reactions catalyzed by enzymes and by solidsFirst-Order Differential Rate Law: A -> products, rate = k * [A] Rate = -d[A]/dt = k[A] ln[A]t = ln[A]0 - kt Becomes ln ([A]t / [A]0) = -kt Or [A]t = [A]0 * e-kt CH3NC -> CH3CN is a first Order process

Second-Order Differential Rate Law: -d[A]/dt = k[A]2 Integrate to get: 1/[A]t = kt + 1 / [A]0 Or y = mx + b NO2 (g) -> NO (g) + 1/2 O2 (g) is an example

Half Life: Time required for one half of a reactant to react Half Life is independent of concentration [A]0 and is a constant for a first order process Half life depends on [A]0 for a second order process.Temperature and Rate: k is temperature dependent such that as temperature increases, so does the reaction rate Rate increases by 2 to 4 times when adding 10 degrees CelsiusCollision Model: In chemical reactions bonds are broken and new ones formed, molecules can only react if they collide with each other in the right orientation and energy There is a minimum activation energy required for a reaction Reaction profile shows energy of the reactants and products so delta E = delta H Middle is the transition state/activated complex, this is inherently unstable Difference in height is equal to the activation energyMaxwell-Boltzmann Distribution of Kinetic Energies in a Population of Molecules: As temperature increases, the curve flattens and broadens, leaving a larger population of molecules at higher energy, so reaction rate increases. Fraction of molecules with energy above Ea is: Arrhenius Equation: A is the frequency factor, represents the frequency of collisions in the right orientation y = mx + b for when you have two temperaturesReaction Mechanisms: Sequence of events that describes the process by which the reactants become products Can occur at once or in discrete steps Steps usually involve breaking a bond, forming a bond, transferring an atom, or transferring an electronElementary Reactions: The molecularity of an elementary reaction tells us how many molecules are involved Can predict rate law from stoichiometry, with the coefficients as the reaction orderMultistep Mechanisms One step will often be slower than the others Reaction cannot occur faster than this rate-determining stepSlow Initial Step Ex: NO2 (g) + CO (g) -> NO (g) + CO2 (g) Rate = k [NO2]2 CO is necessary, but the rate of the reaction doesn't depend on its concentration, meaning the reaction does not occur in one step Possible mechanism:1. NO2 + NO2 -> NO3 + NO (slow)2. NO3 + CO -> NO2 + CO2 (fast) NO3 is an intermediate in this reaction, it doesn't show up in the equation CO doesn't appear in the rate law because it isn't part of the slow step Determine rate from first equation: Rate(1) = k1 [NO2]2 as the overall rate law

Multi-Step Reaction: Fast Reversible Initial Step (Pre-Equilibrium)1. Fast Reversible Initial Step 2 NO (g) + Br2 (g) -> NOBr (g) Rate = k [NO]2 [Br2] Consistent with a one step termolecular reaction Alternatively, because termolecular reactions are rare, this is considered a two step mechanism1. NO + Br2 NOBr2 (fast)2. NOBR2 + NO -> 2 NOBr (slow) If the reaction rate of the forward reaction is greater than the reveres reaction, then NOBr2 will not accumulate, these intermediates often occur in small concentrations Rate of reaction determines on the rate of the slow step Rate = k2 [NOBr2][NO] Can find [NOBr2] by knowing it can react in two ways:1. With NO to form NOBr (slow)2. By decomposing to reform NO and Br2 (fast) The reactants and products of the first step are in fast dynamic equilibrium with each other k1 [NO][Br2] = k-1[NOBr2] k1/k2 [NO][Br2] = [NOBr2] Substitute this in the rate law expression: Rate = (k2k1)/k-1 [NO][Br2][NO] = [NO]2[Br2]Pre-equilibrium is a fast equilibrium that precedes the rate determining step in a mechanism Lie to the left side and produce a small amount of an intermediate in a rate determining step The concentration of the intermediate is defined by the pre-equilibrium alone Overall reaction rate equal to that of the rate-determining stepCatalysts: Increase rate of reaction but are unchanged at the end Homogeneous Catalysts are in the same phase as the reactants Heterogeneous catalysts are not in the same phase Enzymes are large molecules that have globules and their own surface, the exist as both homo and heterogenous catalysts They change the mechanism

Inhibitors Decrease the rate of a reaction, but found unchanged at the end of the reaction Shows up with a negative exponent in the rate law Products can be either Catalysts or inhibitors, but reactants cannotAdsorption Catalysts can speed up a reaction by adsorbing the reactants and bring them close to each other. Thus helping to break bondsChemical Equilibrium: At equilibrium, the forward and reverse reactions are proceeding at the same rate The amount of each reactant and product remains constant Can be reached from either direction Note that you cannot always rely on the change in concentration of products to check if something is in equilibrium: Some reactions just occur too slowlyEquilibrium Constant: Rateforward = Ratereverse k1[products] = k2[reactants] K = k2 / k1 = [products] / reactants With equation: aA + bB cC + dD For gasses, exchange concentration with pressure and use KpKc and Kp P = n/V * R * T Kp = Kc (RT)n n = (mol gas product) - (mol gas reactant)Equilibrium Constant and Units: If a + b = c + d, then Kc and Kp are unitless otherwise, Kc has units of Mn Kp would have units of atmn Keq without units are accepted, but are preferred with units

What does the Value of K Mean? If K >> 1, the reaction is product favored If k cC + dDHeterogeneous Equilibrium expressions: Rules 1 and 2 mean that concentrations of pure solids, pure liquids, and solvents are excluded from equilibrium expressions Ex: CaCO3 (s) CO2 (g) + CaO(s) Kp = P(CO2, g) Kc = [CO2] As long as some CaCO3 or CaO remains in the system, the amount of CO2 gas above the solids remain the same All solid or pure liquid components must be present for any heterogeneous equilibrium to occur Once CaCO3 decomposed, there is no more equilibriumEquilibrium Calculations:1. When K > 106 The reaction is practically complete, use stoichiometric calculations to calculate concentration of products. The tiny concentration of reactants can be calculated from known concentrations of the product and leftover reactants2. When K < 10-6 The reaction does not proceed significantly, concentrations of reactants at equilibrium stay close to their initial concentration. Use equilibrium expression to calculate any change3. When K is between 10-6 and 106 Both reactants and products are present with decent amounts at equilibrium, use a reaction table with unknowns to solve such problems.Reaction Quotient: To calculate Q, one substitutes the initial concentrations on reactants and products into the equilibrium expression Q gives the same ratio the equilibrium expression gives, but for a system not at equilibrium If Q = K, the system is at equilibrium If Q > K, there is too much product and the equilibrium shifts to the left If Q < K, there is too much reactant and the equilibrium shifts to the rightLe Chtelier's Principle: Adding a reactant shifts equil. towards products Removing a product shifts equil. to consume more reactants Reducing the volume of a gaseous mixture (compression) shifts equil. to the side with fewer moles of gas. Done with Haber-Bosch rocess If heat is added to an exothermic process, the equil. will shift left If heat is added to an endothermic process, the equil. will shift right Pressure: If system compressed or allowed to expand, the equilibrium shifts to the side with fewer gas moles Inert gas is added, no change in equilibrium A reactant or product is added/removed -> follow concentrationsAcids and Bases: Historically, acids were sour by tastes Historically, bases were known by neutralizing an acid, or creating a salt with an acid Arrhenius Definition (1880) Acids increase the concentration of hydrogen ions in water Bases increase the concentration of hydroxide ions in water Brnsted-Lowry Definition Acids are proton donors, must have an ionizable hydrogen Bases are proton acceptors, must have a lone electron pair Amphoteric species: Can either be proton donors or acceptors Water acts as a Bronsted Lowry base when you dissolve an acid in water, and forms hydronium Conjugate acids and bases differentiated by one H+, interconverted by proton transfer Spectator ions may exist in acid base reactionsAcids and Base Strength: Strong acids are completely dissociated in water (their conjugate bases have negligible basicity) Weak acids only dissociate partially in water (their conjugate bases are weak bases) If a molecule's conjugate base is strong, then it is not really an acidLeveling effect: Any acid stronger than hydronium in an acidic solution converts completely into hydronium Any base stronger than hydroxide converts completely into hydroxide Hydronium and hydroxide are the ultimate acid/baseAcid-Base Equilibria: The equilibrium is shifted to the side with the weaker conjugate acid and weaker conjugate base A system generally goes down in its chemical potentialAutoionization of Water: Water is amphiprotic, and normally 1 out of 109 water particle acts as a base, adn another as an acid Auto ionization: H2O H+ Ion-Product Constant: Kc = [H3O+][OH-] = [H+][OH-] Specifically called Kw Kw = 1.0 x 10-14 at 25 degrees CelsiuspH: pH = -log[H+] Concentration of H+ at 25 degrees Celsius is 1.0 x 10-7M -log[H+] = 7.00 for pure water If you dissolve an acid in water, pH decreases below 7. Bases dissolved in water increase the pH pOH and pKw exist as well, as the -log of each respective thingie Can be indicated by litmus test, or other papers

Strong Acids: Binary Compounds: HCl HBr HI Oxygen rich oxyacids HxEOy (y - x > 1) E is a non metal or transition metal. Four common ones are HNO3 H2SO4 HClO3 and HClO4 Exist as ions in solutions where [H+] = [acid] May form non metal oxides in solution (N2O5)Strong Bases: The 8 soluble metal hydroxides: W/ alkali metals (LiOH, NaOH, KOH, RbOH, CsOH) W/ heavier alkaline earth metals (Ca(OH)2 Sr(OH)2 Ba(OH)2 ) They dissociate completely in solution Anions like O2- H- NH2- and CH3- are even more basic than OH- and and turn into OH- in solutionSig fig rules for pH values: Only digits after the dots are signifigant DIgits before the dot are unimportant, like digits in an exponent pH is usually measured with 2 sig figs.Weak Acid Ionization For: HA + H2O -> A- + H3O+ Kc = [H3O+][A-] / [HA] Kc is called the acid dissociation constant Ka % ionization = Acid ionized/Acid dissolved x 100% = [A-] / [HA] + [A-] x 100%Other acidic things: Polyprotic acids have more than one acidic proton The first dissociation constant is the largest and usually defines the pH of the polyprotic acid's solutionSoluble weak bases: Inorganic or organic amines (include NH2) Anions of weak acids (like HS- or CO32-)

Weak bases: NH3 + H2O NH4+ + OH- Kb = [NH4+][OH-] / [NH3]Acid-Base Behavior of Salt Solutions: Salts are now defined as ionic compound (not an oxide or hydroxide) Salts are regarded as formed from a parent acid and parent baseSalt of Strong Acid and Strong Base: Salts that are formed by a strong base and strong acids are neutral Na+, Ca2+ are cations of strong bases and have negligible acidic properties Similarly, anions like Cl- and NO3- are anions of strong acids and have negligible basic propertiesSalt of Weak Acid and Strong Base: These salts are weakly basic The cation is neutral, but the anion is a weak baseHydrolysis of Salts: A salt formed by a strong base and weak acid is partially decomposed by water: NaClO + H2O NaOH + HClO ClO- + H2O OH- + HClO A decomposition reaction under water is known as hydrolysis Kb = [OH-][HClO]/ [ClO-] It is Kb because ClO- acts as a baseKa and Kb: Only tabulate Ka for weak acids Kw = Ka * Kb = 1.0*10-14 at 25 degreesSalt of Strong Acid and Weak Base: These salts are weakly acidic The cation is a conugate acid of a weak base, so the cation is a weak acid The anion is the anion of a strong acid and is thus neutralSalt of Weak Acid and Weak Base: If a salt is formed by a monoprotic weak acid and a monoprotic weak base: pH = (pKa1 + pKa2)/2 pKa1 is from the weak acid, pKa2 is from the weak baseKa and Kb for polyprotic acids: Polyprotic acids have several Kas, each anion has its own Kb Ka * Kb = Kw only worsk for a conjugate acid base pairSalts of Polyprotic Acids: Anions which are fully deprotonated (think S2-) are weakly basic. You can calculate their pH with only the basic properties of their anions. Anions of polyprotic acids which keep some protons are amphiprotic pH of a salt with amphiprotic anion: pH ~= (pKa1 + pKa2)/2 for anions HX- and H2X- pH ~= (pKa2 + pKa3)/2 for anions HX2- It is independent of salt concentrationFactors Affecting Acid Strength: In pure state, all acids are molecular. They have non metals only (HCl) or a transition metal in a high oxidation state (+4 or above) A molecule can donate H+ only if it has an H atom with a significant partial positive charge In all acids, H is bonded to a halogen or group 16 non metal Molecules with low polar H-X bonds like CH4 are neutral Metal hydrides (like NaH) are very strong basesBinary Acids The more polar the H-X bond or the weaker the H-X bond, the more acidic compound Acidity increases from left to right across a row, and from top to bottom down a group.Oxyacids (Ternary acids): An OH is bonded to an atom E. The more electronegative E is, the stronger the acids HOF is more acidic than HOCl Pauling's rule: an oxy acid HxEOy is strong if (y - x) > 1 For those with the same central atom, acidity increases with the number of oxygens Limited to: HEO3, HEO4, H2EO4Carboxylic Acids: RCOOH R is any organic group. COOH- is called the carboxylic group and it is weakly basic Alcohols (ROH) are neutral Acidity of carboxylic acids explained by:1. Electron withdrawing in the two Os which create a high partial charge on the H atom2. Resonance structures in RCOO- makes the anion more stableAmino Acids: Amphiprotic Molecules: Amino acids are NH2 - HCR - COOH Amine group (basic), R group, carboxyl group (acidic) The combination of both an acidic and a basic group makes amino acids amphiproticAmino Acids: Zwitterionic Structure: Normally amino acids are doubly ionized: The amino group is protonated while the carboxylate group is deprotonated

Lewis Acids and Bases:

Lewis acids are electron pair acceptors Atoms with an empty valence orbital can be Lewis Acids H+ is the standard example Lewis bases are electron pair donors All Bronsted-Lowry bases are also Lewis basesPolarization of Water Molecules by Cations: alkali and alkaline earth metals have small charges but large radii. So they do not interact much with water in solution and are neutral Other cations strongly polarize water molecules and create more positive charge in nearby hydrogenAcidic Metal Cations: Electrostatic attraction from a metal cation shifts electron density in nearby water molecules The O-H bond is made more polar, and bound water becomes acidic The greater the charge and smaller the size of a cation make it more acidic Cations of metals other than alkali and alkaline earth metals are weakly acidic