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FinalCoggle - Anish Maram - Period 5 (Gases (Kinetic Molecular Theory…
FinalCoggle - Anish Maram - Period 5
Kinetics
Reaction Mechanisms
Valid Mechanism
Rate law of RDS is consistent w/ rate law of rxn
Elementary steps add up to stoichiometric equation
Elementary steps assumed to occur at equilibrium
Concentration of intermediates in RDS rate law can be expressed in terms of other reactants
Forward rate of any step = reverse step; k_1[A]^m=k_1[B]^n
Rate of Reaction
Concentration of Reactants: more reactants/volume --> more collisions --> higher rate
Temp --> Higher avg KE --> more molecules with KE> Ea --> more effective collisions --> higher rate
SA: Surface area --> more collisions --> higher rate
Pressure --> more collisions --> higher rate; no effect if no gaseous reactants
Collision Theory: reactions occur when molecules collide with sufficient energy
Rate Law
Rate constant
Arrhenius Equation: k=Ae^(-Ea/RT)
Order
Integrated Rate Law
Pseudo order integrated rate law: If [B]>>[A] and [C]>>A: Rate=k'[A]^n; allows to examine order of reactant in question
Laws by order
0th order: [A]_t=-kt + [A]_0
1st order: ln[A]_t= -kt + ln[A]_0
2nd order: 1/[A]_t = kt+ 1/[A]_0
Half-life
1st order: 0.693/k
0th order: [A]_0/(2k)
2nd order: 1/(k[A]_0)
Graphical analysis of order: Find Linear fit
1/[A] vs t --> 2nd order
ln[A] vs t --> first order
[A] vs t --> 1st order
Method of Initial Rates
Thermochemistry
State function
Property is independent of pathway, only dependent on initial and final state
Calorimetry
q_system = -q_surroundings
Heat capacity is energy/(mass*temp change) whereas specific heat is for specified amount, usually 1 g
q=mCΔT where C is heat capacity
STP for gas is 0 degrees C and 1 atm, for liquid or solid is 25 degrees C, for solution is concentration of 1 M
Enthalpy
Standard enthalpy of formation is energy required to form 1 mol of substance from reactants; is 0 for pure elements
Enthalpy of Fusion is enthalpy to melt substance into liquid
Bond enthalpy is energy of a bond; enthalpy of reaction can be determined from sum of bonds broken - sum of bonds formed
1st Law of thermodynamics: Energy of universe is constant
Hess's Law
Because enthalpy is a state function, the enthalpy of a process is the sum of the enthalpies of each component reaction
Enthalpy of combustion is enthalpy of substance being burned in oxygen
Heating curves
Periods of same temperature indicates phase change is occurring
3 periods of increasing temperature: while in solid, liquid, gaseous form
Atomic Structure
Quantum Numbers (electron position
m_s: Spin (-1/2, 1/2)
m_l: orbital orientation (p_x)
n: energy level (distance)
l: orbital subshell (s, p)
PES Diagram
y axis is relative # electrons removed
x axis is energy required to remove
Sometimes axes are switched
Trends
IE increases up -right - stronger charge + closer to nucleus
Atomic Radius increases down left - stronger charge and higher energy levels
Electron affinity increases up-right - stronger nuclear charge and closer to nucleus
Electronegativity increases up-right - stronger charge on electron
Z_eff is nuclear charge - number of electrons between nucleus and electron in question
Electron Rules
Pauli Exclusion Principle: no two electrons in an atom can have all same quantum numbers
Aufbau Principle: As protons added to nucleus, electrons are added
Hund's Rule: lowest energy config for atom is maximum unpaired electrons allowed by Pauli in particular set of degenerate orbitals
Size of radius vs. Strength of Nucleus
Shielding: Core electrons repel valence; best shielding to worst: s>p>d>f
EM Spectrum
Length: Gamma<xray<UV<visible<IR<radio
Wavelength*frequency = speed of light
Energy is Planck's constant* frequency
c=2.998E8 m/s
h=6.626E-34 J*s
Thermodynamics
Gibbs free energy: ΔG=ΔH-TΔS
ΔG=-RTlnK
Can plot lnK vs 1/T to find ΔH or ΔS
ΔG<0 means spontaneous, >0 means nonspontaneous
Spontaneity is irrelevant to speed of reaction
ΔG standard is when 1 mol of compound in standard state forms from constituent elements in their standard states
ΔG=ΔG+RTlnQ
ΔS is change in entropy
S of gas>S liquid> S solid
S is measure of number of energetically equivalent microstates
2nd law of thermodynamics: Any spontaneous process increases the entropy of the universe; alternatively, entropy of universe is always increasing
Third law of thermodynamics: entropy of perfect crystal at 0 K is 0
Molar Mass proportional to S
Molecular complexity proportional to S
S=klnW where k = R/Avogadro and W=#microstates
Equilibrium
Law of Mass Action: K=[C]^c[D]^d/([A]^a[B]^b)
Q is same formula as K, but at x given condition; Q=K at Eq
When adding reactions K=K_1
K_2
...
Used to solve w/ ICE table
H2O Production changes concentrations
At Dynamic Eq, forward=reverse
Eq reactions are reversible; K_f=K_r^-1
Le Chatelier's Principle: Disturbance of system causes system to shift in order to restabilize
Temperature causes rxn to shift depending on exothermic/endothermic
Increasing pressure or decreasing volume makes reaction shift to side with less moles of gas
concentration of reactants/products causes opposite shift; only works for gases, concentrations of liquids and solids are constant because amount/volume is same
Inert gases increase total pressure but don't react
K_p=K_c(RT)^Δn
Bonding
Ionic
metal-nonmetal, strong unless in solution, molten form is conductive
Covalent
nonmetal-nonmetal, strong in solution
Metallic
metal-metal
Lattice Energy: Energy associated with forming crystalline lattice of alternating ions from gaseous ions
Is the energy of the final step in formation of ionic compounds
lower magnitude means easier to break bonds of final compound
Strength based on Coulomb's Law q1q2/r; internculear distance lowers energy, higher ionic charges increases energy
Lewis Structures
Resonance structures: same formula but different arrangement
Actual bond length is intermediate between all resonance structures due to delocalized electrons
3rd energy level and above can have expanded octet
VSEPR
Geometry determined by repulsion between electron groups
lone electron pairs repel more than bonding pairs because not spread across two atoms
Difference between electron geometry and molecular geometry; lone pairs lead to different shapes and depress angles slightly
Hybridization is creation of new orbitals between bonding atoms; can be determined from steric number, which is X+E of AXE formula
Sigma and pi bonds
Sigma bond can be between any orbital subshell; linear overlap makes it stronger than pi bond
pi bond only formed between unhybridized p orbitals; small overlap = weaker
For all bonds between two atoms, first bond is sigma then each additional bond is pi
Triple bonds stronger, shorter than double, same for double and single
IMFs
London Dispersion Forces - occur in all molecules due to asymettric electron distributions (temp. dipole moment)
Hydrogen bonding - uniquely strong dipole-dipole interaction in which partially positive hydrogen atom bonds to N, O, or F - is very strong because the H is amost fully severed
Induced dipoles
Ion-dipole interactions - stronger thna hydrogen bonding
Covalent network - compound forms strong convalent bonds
Ionic compound - very strong
Metallic compound: delocalization
Solutions
Mole fraction: # moles of compound/ # total moles in solution
mass percent - percent of solution mass that is specified solute
Molality - moles solute/ kg solvent
Molarity: Moles/solution
Energies of Solution Formation
1: Separating solute into individual components (expanding solute). 2. Overcoming Intermolecular forces in solvent to make room for the solute (expanding the solvent)
Heat of Solution = Heat released/absorbed upon solution formation
Heat of Hydration = expansion of solvent + formation of solvent-solute interctions
Factors Affecting Solubility
Structure --> polarity --> Solubility
Pressure Effects
Solubility Rules
Salts w/ Group I elements or NH4+ are generally soluble
Salts w/ NO3- generally soluble
Salts w/ Cl-, Br-, I- generally soluble UNLESS w/ Ag+, Pb2+, Hg_2^2+
Salts w/ Ag insoluble UNLESS AgNO3 or Ag(C2H3O2)
Mosts sulfate salts soluble UNLESS BaSO4, PbSO4, Ag2SO4, SrSO4
Most hydroxide salts slightly soluble, Group I + hydroxide salts are soluble
Most sulfides of transition metals are insoluble
Carbonates usually insoluble
Chromates usually insoluble
Phosphates usually insoluble
Fluorides usually insoluble
Vapor Pressure
Raoult's Law: P_Total = X_A*P^0_A
Modified version: P_T = X_A*P^0_A + X_CP^0_C
Gases
Kinetic Molecular Theory
Gas particles are points
NO IMFs
Particles are in constant motion
Collisions are perfectly elasti
Boyle's Law: PV=k
Charles' Law: V=bT
Ideal Gas Law: PV = nRT
Modified Ideal Gas Law: P_obs = (P' - correction factor) = P_obs = nRT/(V-nb)
STP: 0 degrees Celsius, 1 atm
Dalton's Law of Partial Pressures: P_Total= P_1+P_2+P_3+...
Definition of gas: uniformly fills any container, easily compressed, mixes completely with any other gas
Units of Pressure: 1 atm = 760 torr = 760 mmHg = 101.325 kPa = 1.01325 bars
Root mean square velocity: u_rms = sqrt(3RT/M)
Diffusion - mixing of gases
Effusion - passage of gas through tiny orifice into evacuated chamber
Graham's Law of Effusion: Rate of effusion
gas 1/Rate of effusion
gas 2 = sqrt(M_2/M_1)
Electrochemistry
Electrochemical Cells
Galvanic Cells; Spontaneous, produce current
Cell Notation (Zn is anode): Zn(s)|Zn2+(aq)||Cu2+(aq)|Cu(s)
W/ same state on one side:
Fe2+(aq),Fe3+(aq)||Ag+(aq)|Ag(s)
Electrochemical Cells; Nonspontaneous, require current
deltaG = -nFEcell where n is electrons transferred/ mol rxn and F = Faraday's constant (96485 Coulombs)
Ecell = Ecathode - Eanode (Must be reduction potentials)
Voltage is an intensive property - doesn't change based on stoich
Redox rxn; cathode = site of reduction, anode = site of oxidation
OIL RIG - Ox is loss e- (inc. charge) , Reduction is gain e- (reduce charge)
Balancing redox rxns
Half-reaction method: Break into half-rxns, balance O with H2O, then H with H+ if acidic, add OH- to H+ side to neutralize and make H2O if basic, then charge with e-, then scale and add rxns
Ox number method
Electrolytic Cells: nonspontaneous, require current
Anode is positive instead of negative, cathode is negative
Oxidation State Rules (Ox state = charge of atoms in compound if were all ionic)
Metals can have more than one common ox state
H = +1 unless in metal hydride (-1 then)
F = -1
O usually = -2, is -1 in peroxide bc Os forn bond w/ each other
Group II = +2
Group I = +1
sum of ox states = total net charge of compound
ox of a lone ion = net charge
Free elements = 0
Acid Base
Buffers
Buffer Capacity
Titrations
Equivalence Point: titrant fully neutralized titrand, pH is 7 unless titrant is weak which would make conjugate relatively strong and therefore change pH
Midpoint: 1/2 titrant has reacted w/ titrand: use henderson-hasselbalch to find pH
Titrant and Titrand
Acid/Base is neutralized through titration
Strong: acid/base fully dissociated to begin with, rxn with titrant goes to completion
Weak: acid/base continually dissociates as titrant is added to solution, bc it reacts with small amount of H+ or OH- -> making Q = 0 --> causing rxn to keep occuring until completion
Titration Curve: pH as fxn of volume titrant added
Reactions
Dissociation Constant
Ka*Kb=Kw
pH = pKa + log([A-]/[HA])
Autionization constant of water (Kw = [H+][OH-])
@ 25 deg C = 1E-14
ICE Table Method
5% Rule (quadratic if otherwise)
Polyprotic Acids/Bases: Series of dissociation reactions (ex. H2SO4, different K values for each
Weak and Strong
Dissociation of Salts: different possible sol'ns
C from WB + A from SA = Acidic
C from SB + A from WA = Basic
C from WB + A from WB = Neutral
pH and pOH
pH+pOH = 14