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Joyce Wu | 5th Period AP Chemistry (Acid Base (Strong acid (HCl, HBr, HI,…
Joyce Wu | 5th Period AP Chemistry
Thermochemistry
Enthalpy
Heat of a Reaction
Can vary with different temperatures
DeltaH
units = kJ/mol
if negative, then rxn exothermic
if positive, then rxn endothermic
Enthalpy of Formation
The heat of the reaction of the formation of an element at standard conditions
Must be only 1 product formed and only 1 mol of product
The heat of formation of any element in its natural state is 0
State makes a difference in the value of the heat of formation
Standard Conditions: 1 atm, 25 degrees Celcius
Enthalpy of a Reaction = [Sum of Enthalpy of Formation of Products] - [Sum of Enthalpy of Formation of Reactants]
State Function
Only care about the beginning and end, steps in between not considered
Hess's Law
Balance the equations to match the desired equation, add all of the enthalpies together.
if reactions reversed, reverse the sign of deltaH
If reaction multiplied by a factor, multiple deltaH by the same factor
q = mCDeltaT
q = joules released
m = mass of substance
C = specific heat
Specific heat is the amount of energy required to raise 1 gram of substance by 1 degrees Celcius
Heat capacity is the heat necessary to raise any amt of substance by 1 degrees Celcius
DeltaT = change in temp
Bond energies
Bonds break = absorb energy. Positive DeltaH
Bonds form = release energy. DeltaH = negative
Calorimetry
q = -q
Calorimeter constant is calculated by (qcoldwater-qhotwater)/change in temp cold water
energy released by exploded substances is energy absorbed by the calorimeter
Thermodynamics
Enthalpy
units: kJ/mol
Heat of a RXN (Energy)
if (+), enthalpically unfavorable
if (-) enthalpically favorable
Entropy
A measure of Randomness or Disorder
If (+), then entropically favorable.
If (-) then entropically unfavorable.
S (solid) < S (liquid) << S(gas)
Driving force for a spontaneous process is an increase in entropy, nature likes increased entropy
Standard states: the standard state of formation for an element/substance is always positive.
Micro State: a state of energy configuration
Macro State: The entire energy system itself
DeltaS of RXN = [Sum of DeltaS of products] - [Sum of DeltaS of reactants]
Gibbs Free Energy
DeltaG = DeltaH - TDeltaS
Gibbs Free Energy = Enthalpy - Absolute Temp(Entropy)
DeltaH = negative, DeltaS = positive; rxn always spontaneous and thermodynamically favorable
DeltaH = positive, DeltaS = negative; reaction never spontaneous
DeltaH = negative, DeltaS = negative; reaction spontaneous at low temp
DeltaH = positive, DeltaS = positive; reaction spontaneous at a high temp
units: kJ/mol
if negative, reaction is spontaneous
if positive, reaction is not spontaneous
DeltaG = DeltaGnot + RTLn(Q); when rxn not at equilibrium
DeltaGnot = -RTLnK; when rxn at equilibrium
IF equilibrium constant is product favored (K>1), then DeltaGnot is negative and rxn is spont.
IF equilibrium constant is reactant favored (K<1) then DeltaGnot is positive and rxn is not spont.
DeltaG becomes 0 when eqn is at equilibrium; Q = K
R constant: 8.314 J/K
elements like to go back .into their natural states; generally have a -DeltaG
Can be calculated DeltaG of RXN = [Sum of DeltaG products] - [Sum of DeltaG reactants]
Laws of Thermodynamics
1.) There is a constant amount of energy in the world.
2.) The entropy of the universe is always increasing.
3.) The only time when entropy is 0 is at the temperature of Absolute Zero.
Kinetics
Method of Initial Rates
Use the two trials that keep one concentration constant
Order of reactant
The degree that a change in concentration of said reactant would affect overall rate
Order of Reaction
The sum of all of the exponents (reactant orders) in the rate expression
r = k[A]^x[B]^y
r = rate
k = rate constant
[A] = concentration of reactant
[B] = concentration of another reactant
k1 = constant for forward reaction
k_-1 = constant for backwards reaction
Pseudo First-Order
A reaction can be pseudo 1st order if the concentration of one reactant is very very small while the concentration of all other reactants is very very large. This way, only a change in the small concentration reactant would have any noticeable effect on the overall rate.
Mechanisms
The slowest step is the Rate Determining Step (RDS)
the RDS has the exact same rate expression as the actual chemical reaction
Validating a mechanism
All steps must have a total sum of the original chemical equation
The slowest step must have the same rate expression as the overall chemical reaction
Intermediates
Species that are produced during the steps of the mechanism, but are not present in the overall reaction
Show up in the mechanisms as a product first, then a reactant second
Catalysts
Lower the activation energy by providing an alternative pathway for reactions to occur
Are not used up in the reaction itself
Show up as a reactant first, then as a product
Mechanisms with a Fast Initial Step
Assume the fast initial step is at equilibrium
therefore, rate forwards = rate backwards
write out the expression, solve for the intermediate, and substitute into the rate expression of the rate determining step
define k as need be -- there are different values of k that make up the overall k
Intermediates cannot be contained in the overall rate expression
Other factors that affect collisions and rate
Temperature;
higher temperatures increase rate; more energy when colliding
Concentration;
higher concentrations generally increase rate --> more collisions
Pressure/Volume:
lower volume + higher pressure can increase rate because there is a higher likelihood collisions will happen in a tighter space
Surface Area:
the more surface area the reactant has, the more places it can come in contact with other reactants, increasing the likelihood of an effective collision
Some species undergo similar reactions but at different rates b/c they have higher initial potential energy or their rxn has a lower activation energy
Unimolecular: most likely to happen
Bimolecular: less likely to happen
Termolecular: very unlikely to happen
For mechanisms, because they are single step reactions, the coefficients are the exponents in the rate expression
Integrated Rate Laws
Describes Concentration vs. Time
Can indicate whether the reaction is 0th, 1st, or 2nd order
0th order: [A] vs. time is linear
[A]t = -kt + [A]0
1st order: Ln[A] vs. time is linear
Ln[A]t = -kt + Ln[A]0
2nd order: 1/[A] vs. time is linear
1/[A]t = -kt + 1/[A]0
slope is -k (negative of the rate constant)
half life
1st order kinetics is the closest to actual half life (doesn't depend on concentration)
0th order: [A]0/2k
1st order: Ln(2)/k
2nd order: 1/k[A]0
Collision Theory
Effective Collisions
Kinetic Energy
Activation Energy: bonds must have enough potential energy
Orientation: there's only 1 correct orientation with enough energy that will start the reaction
Reactant Nature
Reactions happen faster in aqueous solutions than in pure substances (more particle contact)
Equilibrium
Equilibrium Constant / Law of Mass Action
Only changes with temperature changes
Represented by K
Only temperature changes K!!!! nothing else
temperature can change K depending on whether the reaction is endo or exothermic
in the reaction aA + bB --> cC + dD, K = ([A]^a[B]^b) / ([C]^c[D]^d)
can be Kc (concentration) or Kp (pressure)
Relationship between Kc & Kp: Kp = Kc * (RT)^(deltan)
deltan = (sum of product coefficients) - (sum of reactant coefficients)
Unitless
pure solids / liquids are not included in the Law of Mass Action
Relationship between K and the reaction
Reaction reversed = 1/K
Reaction multiplied by a constant n = (K)^n
Adding reactions together, you multiply the K constants
At equilibrium, rate forward = rate reverse
Therefore, K = kf/kb
kf is rate constant forward, kb is rate constant backwards
Concentrations not necessarily equal
K > 1 (product favored)
K < 1 (reactant favored)
K = 1 (concentration of products and reactants are the same)
Reaction Quotient
Represented by Q
Solved same way as K is solved, except the reaction is NOT at equilibrium
Q > K (reaction proceeds backwards to form more reactants in order to reach equilibrium)
Q < K (reaction proceeds forwards to form more products to reach equilibrium)
Q = K (reaction is at equilibrium)
Le Chatelier's Principle
When a system at equilibrium is placed under stress, it will move in the direction to reduce that stress.
Adding more reactants shifts equilibrium to the product side.
Adding more products shifts equilibrium to the reactant side.
Removing reactants would shift equilibrium to the reactant side.
Removing products would shift equilibrium to the product side.
Adding heat (increasing temp) would shift equil to the side with less heat, removing heat (decreasing temp) would shift equil to the side with more heat
Increasing pressure/decreasing volume would shift equil to the side with less moles. Decreasing pressure / increasing volume would shift equil to the side with more moles
Heterogeneous Equilibrium
When there's an equilibrium of a system with species in two or more states.
concentrations of pure solids/liquids don't change equilibrium
Adding water: bit trickier because water decreases the concentration of ions.
If there are ions, adding water would make equil shift to the side with ions.
Ice Tables
Initial (concentration or pressure) , change, equilibrium (concentration or pressure)
Sometimes need to solve a quadratic
Can possibly use the 5% rule
Temperature is constant at equilibrium. However, a system going TOWARDS equilibrium can change temperature.
k1/k-1 using rate constants is the same as the equilibrium constant
Atomic Structure
Shielding
Core electrons shield valence electrons, blocking the effective nuclear charge.
Foundation for other periodic trends/periodicity
Only s orbital and p orbital electrons can be valence
Valence doesn't have a significant shielding effect on other valence, but slightly repel
Effective Nuclear Charge
(Z effective)
Net positive charge that is attracting a particular electron
Calculated by subtracting the # of core electrons from the atomic number
trend of best shielding power to least shielding power: s > p > d > f
Reason: electrons can move around and penetrate into the core electron (penetration)
best penetration power to least penetration power: s > p > d > f
the closer an electron is to the nucleus, the increased attraction it experiences. The more penetrating, the more attraction from the nucleus.
Atomic Radius
Size of radius vs. Strength of Nucleus
Decreases across the period, increases down the group.
Quantum Mechanical Explanation:
Size of the atom is related to the distance the valence electrons are from the nucleus.
The larger the orbital the electron is in, the farther away it is from the nucleus due to feeling less attraction. (i.e. higher energy level orbitals)
More core electrons down the group causes more electron shielding, and thus lower Zeff, meaning lower attraction so size gets bigger
Across a period, Zeffective stays constant while protons/atomic number increases, so size decreases because the same amount of electrons shield a higher nuclear charge
Electron Affinity
The amount of energy released when a neutral atom gains an electron
usually exothermic, in some exceptions endothermic
Group 2 and Group 8 are endothermic
Group 2 is endo because it needs to add a new subshell, the p subshell
Group 8 is endo because it needs to add a new energy level (noble gas)
Always in the gaseous state, just like ionization energy
General Trends
Alkali metals decrease in electron affinity down the column
Generally increases across a period. (Group 5A generally has a lower EA because the extra e- needs to pair)
Halogens have the highest electron affinity
Ionization Energy
The energy it takes to remove one electron
increases across the period, decreases down the group
Exceptions: Oxygen and Nitrogen, because oxygen has 1 bonded pair which causes repelling, so ionization energy is lower as the repelling between the bonded electrons aids in the removal of that electron.
Electronegativity
How much attraction electrons feel towards the nucleus
Increases across a period, decreases down a group.
Higher electronegativity generally means smaller nucleus, since electrons are pulled tighter towards the nucleus.
Causes polarity / nonpolarity in bonding
Bonding
Covalent Bonds
Polarity
Dipole Moments / Polar Bonds
Molecular Geometric Symmetry
caused by a difference in electronegativity
electrons are shared between bonded atoms
can have multiple bonds
bond order indicates the energy associated with given bonds
bond order can be calculated by adding up all of the orders of a repeated bond, then dividing by the number of bonds
in order of increasing energy: single bond < double bond < triple bond
however, higher energy bonds have shorter lengths.
network covalent is the strongest type of bond, such as carbon in graphite and diamond
bonding forces
nucleus-electron attraction
electron-electron repulsion
proton-proton repulsion
Ionic Bonds
transfer of electrons
between a cation and anion
stronger bond than covalent; ionic structures have high melting points
this is because of its high lattice energy; one anion is bonded to multiple cations and vice versa
Lattice energy: the energy of the formation of one solid ion from its gaseous constituents. (Li+ (g) + F- (g) --> LiF (s))
Lattice energy is step 5 of the formation
Metallic Bonds
delocalized electrons
creates a sea of electrons around the block of protons and neutrons
creates a lot of the metallic qualities;
malleability
conducting electricity and heat
ductility
Molecular Geometry and Electronic Geometry
Electronic geometry deals only with the amount of electrons around the central atom
Molecular geometry deals with the entire shape of the molecule due to repulsion and interaction between lone and bonded pairs
VSEPR
Valence Shell Electron Pair Repulsion
Lone pair - Lone pair repulsion > Lone pair - Bonded pair repulsion > Bonded pair - Bonded pair repulsion
AXE Formula / Method
AX2 - Linear
AX3 Trigonal Planar
AX2E1 Bent
AX4 Tetrahedral
AX3E1 Trigonal Pyramidal
AX2E2 Bent
AX5 Trigonal Bipyramidal
AX4E1 Sawhorse/Seesaw
AX3E2 T-shape
AX2E3 Linear
AX6 Octahedral
AX5E1 Square Pyramid
AX4E2 Square Planar
AX3E3 T-shape
AX2E4 Linear
the subscript on X is the number of bonds and the subscript on E is the # of lone pairs
Hybridization
the blending of orbitals
can blend between s and p; d orbital hybridization is no longer supported
sp, sp2, sp3
depends on the steric # of the molecule (Steric # = the subscript on the X + the subscript on the E)
Basically how many electron pairs surround the central atom
Types of bonds
Pi bonds
are formed by parallel unhybridized p orbitals
Sigma bonds
are formed by standard atomic orbitals or hybrids in between the orbitals on the x axis
stronger than pi bonds
single bonds: 1 sigma bond
double bonds: 1 sigma bond, 1 pi bond
triple bonds: 1 sigma bond, 2 pi bonds
Resonance
when electrons move around the bonds so the bonds in the molecule are an average between 2 levels of bonds
the bond energy lies between the energy for the amount of bonds
example; if the molecule has resonance of a single and a double bond, the actual bond energy and length lies somewhere between the single and double energy/length
Formal Charge
the most ideal state of resonance can be determined through calculating formal charge
Formal Charge = (valence electrons in neutral atom) - (unshared valence + half of the shared)
the most ideal resonance state is where most of the atoms in the molecule have a 0 formal charge, while a negative formal charge is on the most electronegative atom
GAS
Ideal Gas
usually at high T low P
Elastic collisions
very light particles
always in constant, rapid motion
no IMFs
The average kinetic energy of a gas depends solely on the temperature, not the identity of the gas
KE = 1/2mv^2
KE has units in Joules; its a derived unit and the MASS MUST BE IN KG
PV = nRT
Boyle's Law: P is inversely proportional to V
Charles Law: V is directly related to T
on a Charles law graph, all lines must extrapolate to 0K
Avogadro's Law: V is directly proportional to n
Solving for density: D = PM/RT
Standard Temperature and Pressure
Temp = 0 degrees Celcius
Pressure = 1atm
Molar Volume @ STP: 22.42 L
Units of pressure
atm
1 atm = 760 torr = 760 mmHg = 101.3 kPa
kPa (kilopascal)
mmHg / torr
Nature of gases
expand to fill their container
fluid = they flow
low density
they're compressible
if you compress a liquid or a solid, IMFs will get stronger and they would change state
they effuse and diffuse
Effusion and Diffusion
if reactants and products are at the same conditions of T/P, then the mole ratio of the gases are the same as the volume ratio
Graham's Law:
Rate of Gas1/Rate of Gas2 = Sqrt(M2/M1)
Root mean square velocity
RMS = Sqrt(3RT/M)
IMFs
Hydrogen bonding
stronger IMFs
can only happen when an H is bonded to something extremely electronegative, such as O, N or F
Dipole - Dipole
happens due to a large difference in electronegativity
the more electronegative molecule attracts electrons from the less electronegative molecule, leaving a partial negative and a partial positive
London dispersion
weakest type of IMF
happens because of a temporary dipole; where electrons are sometimes more gathered on one side of the atom, creating a temporary partial negative/positive
Larger molecules have stronger london forces
forces between molecules, not within the molecule itself
Account for physical properties, such as BP, MP, etc.
BP is when the vapor pressure of the liquid becomes equivalent to atmospheric pressure
More/Stronger IMFs, = lower vapor pressure = more energy required to overcome IMFs = higher BP
Network Covalent
Strongest type of IMFs
Found in things like carbon chains or silicon
Scale of IMFs from strongest to weakest: Network Covalent > Metallic > Ionic > Hydrogen Bonds > Dipole > London
Larger charge = stronger attraction
Shorter distance = stronger attraction
Solutions
Molarity = mols/L
Molality: mols / kg
Heats of solution: the amount of heat energy absorbed / released when a specific amount of solute dissolves in a solvent
usually exothermic
if endothermic, that means entropy is the driving force
the magnitude / sign can be attributed to strength of IMFs
Making a solution
Overcome attractions between solute particles; DeltaH solute is endothermic
Overcome some attractions between solvent molecules (stretch the bonds) DeltaH solvent is endothermic
DeltaH solvent is more endothermic depending on the strengths of IMFs in the solvent
Form new attractions between solute particles and solvent molecules; DeltaH mix is exothermic
DeltaH solution = DeltaH solute + DeltaH solvent + DeltaH mix
If exothermic, that means DeltaH mix overwhelms the other two; therefore the solute-solvent bonds are very strong.
Generally, solids become more soluble with increasing temperature
Raoult's Law
the presence of a nonvolatile solute lowers the vapor pressure of the solvent
P(solution) = X(solvent)(Vapor pressure pure solvent)
The vapor pressure of a solvent in a solution is always lower than the vapor pressure of the pure solvent
Raoult's Law for Liquid-Liquid solutions
Ptotal = Xa(vapor pressure of pure solvent) + Xb(vapor pressure of pure solvent)
Ideal solutions
Negative deviation
lower than predicted vapor pressure
DeltaH soln is large and -
Ideal solutions perfectly follow Raoult's Law
Heat of solution = 0
Positive deviation
higher than predicted vapor pressure
DeltaH soln is small and +
The solute is just diluting the solvent in an I.S.
Factors that favor solution formation
negative DelatH soln
Increase in entropy
Solubility trends
more soluble as temp increases
increased surface area
gases are more soluble at low T, high P
Limits
unsaturated: less solute than saturation
saturated: solute/solvent in dynamic equilibrium
supersaturated: more solute than saturation
Acid Base
Strong acid
HCl
HBr
HI
H2SO4
HClO4
HClO3
HNO3
More electronegative/polar substances make a stronger acid because it weakens the H bond, making the proton easier to pop off
Strong base
NaOH
KOH
LiOH
RbOH
CsOH
Ca(OH)2
Sr(OH)2
Ba(OH)2
Strong acids and bases completely ionize
the K value is very high ( K >> 1); product favored
all other substances are weak acid/base
weak acid/base ionize less than 5%
K < 1, reactant favored
Arrhenius model
acids produce H+ ions in (aq) solution
bases produce OH- ions in (aq) solution
Bronsted-Lowry Model
acids donate protons
bases accept protons
Lewis acid model
acids are electron pair acceptors
bases are electron pair donors
Conjugates
Conjugate acid -- the product that received the proton in the products
Conjugate base -- the product that lost a proton from the original acid
Concentration and pH
-log[H+] = pH
[H+] = 10^-pH
-log[OH-] = pOH
[OH-] = 10^pOH
pKa = -log[Ka]
pKb = -log[Kb]
Acidic and Basic Salts
If anion from strong acid and cation from strong base, the salt is neutral.
If anion from strong acid and cation from weak base, salt is acidic.
If cation from strong base and anion from weak acid, salt is basic.
If cation and anion are both from weak substances, compare K values. the larger K value determines which is stronger
Equilibrium
KaKb=Kw
Kw = 1.0E^-14
Ksp = solubility constant
Solubility
Mols of substance able to be dissolved in 1L of solution
Larger the solubility constant, more soluble the substance
Molar solubility can be calculated with ICE table
Substance with the smallest solubility constant will precipitate first
Buffers
Resist pH change
Made from a weak acid and its salt or a weak base and its salt
Common ion effect: less ion wants to dissolve in a solution where that ion is already present.
H-H Equation: pH = pKa + log([A-]/[HA])
the best buffer is a 1:1 ratio of [base]:[acid]
Titration
Strong acid / strong base = pH is 7
If titrating a strong base with a weak acid the equivalence point pH is going to be higher than 7; vice versa for titrating a weak base with a strong acid
Equivalence Point: when moles of acid = moles of base and the solution is neutralized
At the half equivalence point, pH = pKa or pOH = pKb
buffer solution will mitigate pH change, the graph will not dip until the buffer is all used up
Electrochemistry
Reduction
When a substance gains electrons
It's oxidation number is reduced
Happens at the cathode
Standard reduction potential
Tendency for a substance to be reduced.
Usually for half-reactions
The higher the reduction potential, the more likely something will be reduced, and the better the oxidizing agent.
The lower the reduction potential, the more likely something will be oxidized, the better the reducing agent
Ecell = Eox + Ered
Eox can be found by switching the sign on the standard reduction potential.
Oxidation
When a substance loses electrons
Its oxidation number increases
Happens at the anode
Galvanic Cell
Electrons flow from the anode to the cathode
Spontaneous if Ecell is positive
Balancing redox equations
Half reaction method
Separate reaction into respective ions
Balance O first, using H2O
Then balance H's using H+ ions
Use electrons to balance the charges on each side.
Combine half reactions, cancel out substances that appear on both sides
If under basic conditions, add OH- to create water
Oxidation numbers
Any element in its standard/uncombined state is 0
Ions have an oxidation number of their charge
Oxygen is -2 unless in peroxide where its -1
Hydrogen is always +1
The sum of all oxidation numbers in a given compound is equal to the compound's charge
Fluorine is always -1
Electrolytic cell
Requires outside power source to drive a nonspontaneous electrolysis
The cathode and the anode switch, because the substance being reduced/oxidized switches
Metal plating
Use Faraday's constant and Amperes to do dimensional analysis