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