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Ruihua Yang Per: A Final Coggle (Thermodynamic (Enthalpy: H When ΔH is…
Ruihua Yang
Per: A
Final Coggle
Thermochemistry
q(calorimeter) = C(calorimeter) * ΔT
q(calorimeter) = -q(reaction)
ΔHrxn = qrxn / mol of reactants (heat realeased / mole of reactants)
Endo = +; Exo = -
q(solution) = m(solution)
C(solution)
ΔT
q(reaction) = -q(solution)
q=mcΔT
E= q + w (heat + work)
w = -PΔV (negative pressure * change in volume)
Kinetic Energy: Due to motion
Potential Energy: Due to position
1000 J = 1kJ
Thermodynamic
Enthalpy: H
When ΔH is positive, this reaction is endothermic (meaning it is gaining heat from the surroundings
When ΔH is negative, this reaction is exothermic (meaning it is losing heat to the surroundings
ΔHrxn = coe. of each proΔH(products) - coe. of each reactΔH(reactants)
Gibbs Free Energy: ΔG & Spontaneity
When ΔG°rxn is
negative
, then this reaction is
spontaneous
When ΔG°rxn is
positive
, then this reaction is
non-spontaneous
There are three methods to calculate ΔG°rxn:
ΔG°rxn = coe. of each proΔG°rxn(products) - coe. of each reactΔG°rxn(reactants)
Use Hess's law
ΔG°rxn = ΔHrxn - TΔSrxn
Entropy: S (how random is the particles after a reaction occur)
s < l < aq << g
This means the Entropy is increasing, vice versa
ΔSrxn = coe. of each proΔS(products) - coe. of each reactΔS(reactants)
When ΔSrxn is
positive
, the Entropy is
increasing
When ΔSrxn is
negative
, the Entropy is
decreasing
ΔSuniverse = ΔSsurrounding + ΔSsystem
ΔSsystem = ΔH / T
ΔSsurrounding = -ΔH / T
Spontaneous & Non-spontaneous
Spontaneous means the reaction occurs when outside factor ONLY occurs once, and then it just complete on its own
Non-spontaneous means the reaction occurs with FULL outside factors assists
Conditions When a Reaction is Spontaneous
ΔH is + & ΔS is +, it is
spontaneous at very HIGH Temperature
ΔH is - & ΔS is -, it is
spontaneous at very LOW Temperature
ΔH is + & ΔS is -, it is
NEVER spontaneous
ΔH is - & ΔS is +, it is
spontaneous at ALL Temperature
At Equilibrium
ΔGrxn = ΔG°rxn + RTln(Q), where Q = [Products] / [Reactants] = K
ΔGrxn = 0, therefore ΔG°rxn = -RTln(K)
Kinetic
Collision Theory
Effective Collision
must have three conditions: 1. Particles collide; 2. Sufficient Energy; 3. Correct orientation
Ineffective Collisions
does not have all the conditions above
Ex: If the particles lack the necessary KE may collide, but they will simply bounced off
Effect of Reaction Rate
Surface Area
- increase the surface area will cause the increase in particles collision; therefore, as more particles collide, the reaction rate will increase
Temperature
- when increasing the temperature, the particles collide faster; therefore, causing a faster reaction rate
Pressure
- when the pressure increases, more particles will collide to each other; therefore, causing the reaction rate increasing
Concentration
- increasing the concentration also increase the number of particles collision; therefore, a faster reaction rate will occur
Catalyst
- adding a catalyst into a reaction will lower the activation energy, and create an alternate faster pathway; therefore, the reaction rate will increase
Nature of the reactants
State: (solid < liquid (l, aq) (ions in solution) < gas); therefore, whenever two reactions (one contains solid state, the other are gases or liquid), then this solid state is the slowest
Bond type: reactions involving ionic-bonds (+,-) are more faster than molecular bonds
Bond strength: Reactions involving the breaking of weaker bonds proceed faster than reactions involving the breaking of stronger bonds
Number of bonds: Reactions involving the fewer ions are faster than involving more ions
Orders
0 order
[A]t = -kT + [A]o (integrated rate law)
M/s
t1/2 = [A]o/2k
1st order
ln[A]t = -kt + [A]o (integerated rate law)
1/s
t1/2 = 0.693/k (independent on initial rate)
2nd order
1/[A]t = kT + 1/[A]o (integrated rate law)
1/Ms
t1/2 = 1/k[A]o
Whenever the graph is very close to a linear line, then this graph's order is the overall order
Mechanism
Occurred as single steps
Two conditions:
When adding up the single steps, this has to equal the overall rate law
The rate law has to equal to the overall rate law
The slowest step is the rate-determining step, and use this reaction to determine the rate law
Pseudo First-Order Reaction
:
The fake 1st order reaction
In the Pseudo First-Order Reaction, we are basically isolating a reactant by increasing the concentration of the other reactants. When the other reactants are in excess, change in their concentrations does not affect the reaction much, Therefore, now the reaction only depends on the concentration of the isolated reactant. The concentrations of all the other reactants are taken as constant in the rate law. Thus, the order of reaction becomes one.
Ex: A + B => P
Where [B] >> [A], [B] = constant
r = k[A][B]
r = k'[A] where k = k'[B]
Equilibrium
Equilibrium Involving Gas
Kp = [(PC^c)(PD^d)] / [(PA^a)(PB^b)]
Kp = Kc(PT)^change in n (moles, products - reactants)
Reactions Quotient
-
Different as K, Q is under any conditions beside equilibrium
Qc = ([C]^c[D]^d) / ([A]^a[B]^b)
Qp = [(PC^c)(PD^d)] / [(PA^a)(PB^b)]
Q > K, shifts to the left; Q < K, shifts to the right; Q = K, the system is at equilibrium
Basic Ideas
An equilibrium reaction is when the rate of the forward reaction equals the rate of the reverse reaction and the concentration of products and reactants remains unchanged ( <=> sided arrows)
As the concentration of product increases and the concentrations of reactants decrease, the rate of forward reaction slows down, and the rate of the reverse reaction speeds up
Rates
are equal
doesn't mean everything else is equal
K: Equilibrium Constant
K > 1, product favor
0 < K < 1, reactants favor
Solid
&
Liquid
are not involved in the equilibrium expression
Law of Mass action: ([C]^c[D]^d) / ([A]^a[B]^b)
Kbackward = 1 / Kforward
Knew = K^n original; Knew = K1K2
ICE Table
Initial
Change
Equilibrium
Le Chatelier's Principle
Concentration: Increases in products shift to left (reactants), decreases in products shift to left. Increases in reactants shift to the right (products), decreases in reactants shift to the right.
Temperature: For Endothermic reaction, as the heat raised, it shifts to the right, vice versa. For Endothermic reaction, as the heat raised, it shifts to the left, vice versa.
Pressure: Increase the pressure favors less mol side, which means as the volume decreases, vice versa
Atomic Structure
Electron Affinity (Energy)
Energy is released when a neutral atom gains an electron
Halogen
has the highest EA
GAS
state
Ex: M(g) + e- <=> M-(g)
Exothermic Reaction
The more Energy released, the more EA
"Generally" increase from left to right BUT:
Group 5A lower (more positive) EA because now the electrons have to pair up = need more energy
Group 2A and 8A generally have low EA because after it adds an electron, it jumps to next energy orbital
Ionization Energy
Increased from left to right, and bottom to up
The minimum E that required to remove an electron from the nucleus
Endothermic Reaction
Always in the
GAS
state
Ex: M(g) <=> Mg1+ (g) + e-
Magnetic
Paramagnetic
- attract to magnetic field; unpaired electrons in electron configuration
Dimagnetic
- repelled to magnetic field; full paired electrons in electron configuration
Atomic Radius
Increase from right to left, up to bottom
(Zeff increases)
The larger the orbital an electron is in, the farther its most probable distance will be from the nucleus, and the less attraction it will have for the nucleus.
The stronger the Zeff, the more attraction between the electrons and the nucleus
Ionic Radius
Cations are much smaller than anions
Electronegativity
Increased from left to right, bottom to top
The attraction between electrons (bonding), the larger the difference in Electronegativity, the more polar of the bond
If the difference is 0, then it is equal sharing (O2)
If is greater than 2, then it is an ionic bond (NaCl)
If it is between 0.1-0.4, then it is nonpolar
If it is between 0.5-0.9, then it is polar
Dipole Moment
More electrons two atoms shared, and the atoms are large, then they have the larger dipole moment
Bonding
Triple bond < double bond < single bond (lengths)
Triple bond > double bond > single bond (energies)
Resonance
Resonance bonds are shorter and stronger than single bonds
Resonance bonds are longer and weaker than double bonds
Formal Charges
FC = Valence Electrons - (unpaired electrons + paired electron/2)
Atoms in molecules are trying to get closer to zero (FC)
All atoms' charges add up should equal its ion or molecule charge
Any negative FC's are expected to be on the most electronegative atoms
Ionic Bonds
Electrons are transferred; electronegativity is always greater than 1.7; exothermic
If conduct electricity, then it is ionic bond
Coulombs' Law
- describes the attraction and repulsion between charged particles
For like charges (same charge), the potential E is positive and will increase when r (distance) gets shorter
For opposite charges, the potential E is negative and become more negative when r (distance) gets farther
The strength of the interaction increases as the size of the charges increase
Covalent
Polar-Covalent Bond: unequal sharing (have lone pairs)
Non-polar Covalent Bond: equal sharing (no lone pairs)
Lattice Energy
Increased while charges of particles increase and atomic radius decreases