All Units Flashcards
(233 cards)
1
Q
Molecules vs Ions
Molecules
A
- Covalent
- All states
- Simplest = molecules
2
Q
Molecules vs Ions
Ions
A
- Electrostatic Forces of Attraction
- Solids with 3d crystal lattices
- No prefixes with exceptions
- can become hydrates
3
Q
Types of Reactions
Decomposition
A
- Breaking down compound to form simpler compounds/elements
- Endothermic
4
Q
Types of Rxns
Ionization
A
- Rxn with polar covalent ions to ions in H2O
5
Q
Types of Rxns
Dissociation
A
- Separation of ions in H2O
6
Q
Types of Rxns
Melting
A
- Solid ==> liquid
7
Q
Solubility
Soluability Rule
A
- Salts with Na+, K+, NH4+, NO3-
8
Q
Solubility
Insolubility Rule of Thumb
A
- max concentration less than .01 M, insoluable
9
Q
Conduction and Electrolytes
Strong Electrolytes
A
- More ions = More Strength
10
Q
Conduction and Electrolytes
Weak Electrolytes
A
- Weak acids and Bases
11
Q
Conduction and Electrolytes
Non-Electrolytes
A
- nonpolar componds/ org compounds except carboxylic acids and amines
12
Q
Thermodynamics Introduction
Heat
A
- speed or Energy of particles
13
Q
Thermodynamics Introduction
Energy
A
Capacity to do work
14
Q
Thermodynamics Introduction
Work
A
Action of Force through ∆x
15
Q
Thermodynamics Introduction
Total Energy of object
A
- KE: E associated with motion ==> Thermal Energy
- PE: (bond energy); energy associated with position or composition
16
Q
Energy Units
1 L*atm
A
101.325 J
17
Q
Energy Transfer
Energy Transfer between system and surroundings
A
- Sys decrease(-) = Increase surroundings(+)
- Sys increase(+) = Decrease surroundings(-)
18
Q
Definition of Specific Heat(c)
A
- Amount of energy required to raise 1 gram by 1 C
19
Q
Exothermic vs Endothermic
Exothermic
A
- Graph increases then ends lower than starting level(∆ PE = -)
20
Q
Exothermic vs Endothermic
Endothermic
A
- Graph increases then ends higher than starting level(∆ PE = +)
21
Q
Types of Systems
Open
A
- matter and Energy exchanged with surroundings
22
Q
Types of Systems
Closed
A
- Only Energy may exchange with surroundings but not matter
23
Q
Type of System
Insulated
A
- No energy or matter exchanged with surroudings
24
Q
Transfer between system and surroundings
Energy Transfer with work and heat
A
- E = q + w
25
# Transfer between system and surroundings
Heat/Q
* Driving Force = ∆T
26
# Transfer between system and surroundings
Work/W
* Driving Force = - P * ∆V
* W= Fd
* Compression = (+)
* Expansion = (-)
27
# Heat
Heat Capacity
* C
* q=C∆T
* slope of Q and ∆T
* units = J/C
* Extensive: Depends on mass
28
# Heat
Specific Heat
* c
* c = Q/m/∆T
29
# Calorimeter
Bomb
* Constant Volume
* qcal=Ccal* ∆T
* qcal = qrxn = ∆Erxn
30
# Calorimeter
Coffee-Cup
* Constant Pressure
* qcal=Ccal* ∆T
* qcal = qrxn = ∆Erxn
31
# Standard Enthalpy of Formation
Standard State
* State of Pure Substance at l atm P and Temp of Interest(25 C)
32
# Standard Enthalpy of Formation
Standard Enthalpy of change
* ∆H^o
* ∆H of reactants and products
33
# Standard Enthalpy of Formation
Standard Enthalpy of Formation
* ∆H when 1 mol substance formed from compund standard state elements
34
# Standard Enthalpy of Formation
Standard Enthalpy of Formation Equation
* ΔHoreaction=ΣΔHof(p)−ΣΔHof(r)
35
# Hess' Law
Hess' Law
* If equation can be explained as sum of 2+ equations, ∆ H for desired equation = ∆ H sum of other equation
36
# Phase Changes
Heat absorbed into system
* q > 0
* melting, vaporization, sublimation
37
# Phase Changes
Heat released into system
* q < 0
* freezing, condensation, deposition
38
# Gas Laws
Interaction of External and Internal Pressure
* No change: Internal Pressure = External Pressure
* Compression: Internal Pressure < External Pressure
* Expansion: Internal Pressure > External Pressure
39
# Gas Laws
Calculating Pressure under liquid
| Equations
* With Atmosphere: P(H2O)**["Column Pressure"]** + P(atm)[atmospheric/ barometric pressure]
* In Vaccum: P(H2O)**["Column Pressure"]**
40
# Properties of Gases
What is a Vapor?
* Gaseous state of substance usually a liquid or solid at room temp and pressure
41
# Properties of Gases
Gas Elements
| that are gases at Room Temp
* Noble gases(are isolated atoms)
* Diatomic Gases: *H2, N2, F2, O2 and Cl2*
* **HCl, NH3, CO2, N2O, CH4, HCN**
42
# Gas Characteristics
Gas Characteristics
1. Gases assume shape and volume of container
2. Move in constant, random motion but in straight line
3. Gas Density less than Liquid Density or Solid Density
* It is highly variable[*Increased T --> Decreased D vs Increased P --> Increased D] *
4. Gases form Homogeneous Mixture with each other in any proportion **[0 rxn = mutually miscible]**
43
# Pressure
Pressure equation
* F/a
* dgh**[in Pascals]**
* mmHg is a measure as well as a pressure unit
44
# Monometer
What is a Monometer
* Measures gas P
* P = difference of liquid height
* 2 types
1. Close-end(vaccum)
2. Open-end(atm)
45
# Pressure
Conversions
1 atm = 760 mmHg = 760 Torr = 101.325 kPa = 1.0325 bar = 14.7 lb/in^2 = 101,325 N/m^2
46
# Monometer
Monometer Equations
* For Close-Ended: Pgas = Δh(liquid)
* For Open Ended:
1. Pgas = Δh(liquid) + Patm[Pgas>Patm]
2. Pgas = Δh(liquid) - Patm[Pgas
47
# Ideal Gas Law Equation
What is STP?
* Standard Temperature and Pressure
* 273.15 K and 1 atm
48
# Ideal Gas Law Equation
What is an Ideal Gas?
* @ higher Temperature: More gases
* @ Lower pressure: More space
* Rules:
1. Gases move randomly
2. No attraction between particles
3. "Infinite" volume and and Volume of gas not important
4. Obeys simple gas laws: Boyle, Charles, Avagadro
## Footnote
MCQ?
49
# Ideal Gas Law Equation
Boyle, Charles, Avagadro Equations
* Boyle: V=1/P
* Charles: V=T
* Avagadro: V=n
## Footnote
MCQ?
50
# Ideal Gas Law Equation
Combined Gas Law
P1xV1/(n1xT1)=P2xV2/(n2xT2)
51
# Ideal Gas Law Equation
Units
* P = any
* V = any
* n= mol
* T = K
52
# Ideal Gas Law Equation
Types of R
* R = 0.08206 liter·atm/mol·K
* R = 62.36 L·Torr/mol·K or L·mmHg/mol·K
53
Density equation
D=MP/RT
54
Molarity Equation
M=mRT/(PV)=DRT/P
55
# Stoichiometry with Gases
Law of Combining Gas Volumes
* Ratios from balanced equations can be in L instead of mol when Same Temp and P
* If not use ideal gas law as part of stoic(Ex. 1/P to get rid/add Pressure)
56
# Dalton's Law of Partial Pressure
Equation for P.P.
* Ptot=P1+P2+...Pi
57
# Mole Fractions of Gases
Equations
* Xi=ni/ntot=Pi/Ptot
* Mole % = Mole Fraction * 100%
58
# Kinetic Energy Molecular Theory
Postulates
1. Gas particles so small and distance so large that individual volume is negligible
2. Particles in Constant Motion
3. No Forces between Particles(No Attraction or Repulsion)
4. Average Ke proportional to T(in K)
59
# Kinetic Energy Molecular Theory
Deviation of Ideal Gas Law Equation
* P=1/3(N/V)mū^2
* N= # of molecules
* m = mass of 1 mole
* ū^2 = Average of squared velocities
60
# Kinetic Energy Molecular Theory
Proportions
* KE proportional to T
* Urms proportional to √1/M
* Urms proportional to √T
61
# Kinetic Energy Molecular Theory
Equalities
* U1/U2 = √(M2/M1)
* t1/t2 = √(M1/M2)
62
# Kinetic Energy Molecular Theory
Root Mean Square Velocity Equation
* √(ū^2) = √(3RT/M)
1. R = 8.314 J/(mol·K)
2. M= **kg/mol**
* OR sum of speeds squared divided by number of molecules, then sqrted
63
# Kinetic Energy Molecular Theory
What does Volatile mean?
* Molecules that have weaker IMFS; easier to go from liquid to gas
64
# Kinetic Energy Molecular Theory
What does Effusion Mean?
* Gas escapes container through hole into an evacuated chamber
* Diffusion Properties still in play(Go from high to low pressure)
* Effusion inversely proportional to √M
65
# Non-Ideal Gas Behavior
When does Non-Ideal Gas Behavior happen?
* Low Temp, High Pressure
* As Temp increases, deviation decreases: Less IMFs because of less interaction
* As Pressure increases, deviation increases: Actual volume available will be greater than predicted because gas molecules take up more space
66
# Non-Ideal Gas Behavior
Van der Waals Equation
* P=nRT/(V-nb) - (n^2)(a)/(V^2)
1. a and b gas constants
2. nb = accounting for size of gas molecules
3. n^2a = accounting for IMFs
67
# Electrons
How Light Energy is Determined?
* wavelength
* frequency
* energy
68
# Electrons
When are wavelengths seen?
When electrons return to ground level
69
# Electrons
Planck's Constant
* h
* 6.022E-34 Jᐧs
* Slope of E∝V
70
# Properties of Light
What is Electromagnetic Radiation
* EMR
* Properties of Light
71
# Properties of Light
Wavelength
* Length of one wave
* Meters
* 𝝀
72
# Properties of Light
Frequency
* Number of wavelengths(or cycles) per sec passing a point
* 1/s or s^-1 or Hz
* 𝛎
73
# Properties of Light
Speed of Light
* 2.998E8 m/s
* c
74
# Properties of light
Amplitude
* Higher is Brighter while Lower is darker
75
# Properties of Light
Equations for Energy, and Light Parts
* c = 𝝀 ᐧ 𝛎
* E = h𝛎
* E = hᐧc/𝝀
* Is the Energy of 1 photon
76
What are Photons?
Particle side of electrons
77
# Planck's Work
What does delocalized mean?
the electron is not with a certain atom or nucleus
78
# Planck's Work
What is a Quantum
A packet of Energy for an electron to jump to the next electron level
79
# Einstein's work
What is Photon Energy related to?
* **E**photon ∝ 𝛎
* **E**photon ∝ 1/𝝀
80
PES equation
**E**photon = **KE**electron + **BE**electron
- BE = binding energy
81
# Bohr's Work
Energy Equation to move electron in Hydrogen
| NOT IMPORTANT FOR TEST
ΔE = (-2.178E-18 J )((1/*n*f^2) - (1/*n*i^2))
82
# de Broglie's Work
Mass Equation
h/(𝝀v)
83
# Quantum Numbers
Types
* n = Priniciple quantum number
* *l* = Angular quantum number
* m = Magnetic quantum number
* ms = Magnetic Spin
84
# Quantum Numbers
Nodes
No possibility for electron(white shell)
85
# Quantum Numbers
Principle Quantum Number
* n
* Integral Number
* Related to size and energy of orbital
* Corresponds to Bohr's energy level
* More n
1. Increased Orbital
2. Increased distance of electron from orbital
3. Increased Energy
4. Decreased energy between orbitals
86
# Quantum Numbers
Angular Momentum Quantum Number
* Shape
* *l* = 0--> (n-1) for each n val
* if n=3, possible orbitals s(0)-->d(2)
87
# Quantum Numbers
Magnetic Quantum Number
* m
* integers specifying orbital orinetation
* Values are from -*l*←→+ *l*
* Includes 0
* Example: *l* = 2: m=-2,-1,0,1,2 ⇒ 5 orbitals
88
# Quantum Numbers
Electron Spin
* Up = + .5 spin
* Down = - .5 spin
89
# Quantum Numbers
Difference between H and Multielectron atoms
* H = subshell E levels that are degenerate(same n-int at same level)
* Multi = lower orbital energies
* Subshell of prinicple shell at different energies
90
# Quantum Numbers
Rules for Electron Placement
1. Pauli Exclusion Principle
* No 2 electrons in atom has same 4 quantum numbers
* electrons in 1/2 filled orbitals have parallel spins
1. Hund's Rule
* One electron for each orbital before doubling up
3. Aufbau Principle
* Electron occupy lowest energy level possible
91
# Electron Configuration
How do you write Electron Configuration
* Removal Order
* Ex) Se: [Ar] 3d^10 4s^2 4p^4
92
# Electron Configuration
Valence Electrons
* Outermost principle shell
* Usually S or S and P
93
# Electron Configuration
Shortcut using Noble Gases
* Can use closest previous noble gas in brackets then build the rest of the electron config from there
* Cannot use this for noble gas in ground state
94
# Electron Configurations
Electron Configuration Exceptions
* Cr: [Ar] 3d^5 4s^1 (Same with Mo)
* Cu: [Ar]3d^10 4s^1(Same rule with Au and Ag)
95
# Ionization Energy
What is Ionization Energy
* Minimum Energy neede to remove electron from atom or ion
96
# Ionization Energy
Ionization Energy Requirements
1. Gas State
2. Endothermic
3. Valence electron first
4. Successive removed for 2nd etc IEs
97
# PES
Why are X-rays used?
* Can dislodge electrons
98
# PES
Relation between BE and KE
* Inversely Related
99
# Magnetic Properties of electrons
Paramagnetic
* one or more unpaired electron
* attracted by a magnetic field
100
# Magnetic Properties of electrons
Diamagnetic
* Paired electrons
* opposite spins that cancel out their magnetic fields
* Are not attracted to outside magnetic field
101
# Magnetic Properties of electrons
How are they detected
* weighing a substance in the presence of a magnetic field
102
# Periodic Trends
Atomic Radius
* Is the distance form the nucleus to the valence electrons
* Across Period: decreased radius due to higher effective nuclear charge
* Down Group: Increased radius due to higher n and a greater distance from the nucleus(higher n ==> higher V)
103
# Periodic Trends
Cations
* Decreased sized compared to original
* decreased electrons with same proton number
104
# Periodic Trends
Anions
* Increased size
* Increased electrons with same proton number
* Electron - Electron Repulsion
105
# Ionization Energy
What is the first IE equal to?
* Binding Energy
106
# Electron Affinity
Electron Affinity
* Energy released when neutral atoms gain electrons
1. Needs to be in gas state
2. M(g) + 1 electron --> M^1-(g) + EA etc(Successive)
* Exothermic
* Increased energy leads to Increased negative EA
* Outer electrons are delocalized: can move around which leads to sea of electrons
107
# Periodic Trends
Metallic
* Most = Bottom Left
* Least = Top Right
108
# Coulomb's law
Force
* (k x q1 x q2)/r^2
* q = charges
* Repulsion Force = (+)
* Attraction Force = (-)
109
# Coulomb's law
Energy
* (k x q1 x q2)/r
* q = charges
* E released or required to make bond
110
# Lattice Energy
Lattice Energy
* Energy required to seperate 1 mole of an ionic solid into gaseous ions
* Higher LE ==> Higher Ionic Bond Strength
111
# Lattice Energy
Lattice Energy Calculations
* Born-Haber Cycle = Applying Hess' Law to its maximum
112
# Lewis Structures
Lewis Structures
* Used to explain that atoms combined to acheieve more stable electrong configuration
113
# Lewis Structures
Ions
* = [ ] + charge
* Ex) .Cä. --> [Ca] 2+
114
# Lewis Structures
Coordinate Covalent Bond
* Electrons donated from one atom to another
115
# Formal Charge
What are Formal Charges used for?
* Determines most possible lewis structure
116
# Formal Charge
Formal Charge Equation
* FC= V.E. - (non-bonding e-) - (# of bonds)
117
# Formal Charge
Formal Charge Rules
1. FC= 0 for neutral or Ion FC= Ion charge
2. FC = 0 more favorable than FC >/< 0
3. Less FC > Greater FC
4. Best strucure = Lewis Structure with FCs similar to ENs
* Increased EN = Increased - FC
118
# Resonance
Resonance Theory
* Molecule/ Ions with 2+ plausible Lewis Structures with diff e- distribution
* Hybrids joined with double-headed arrows
119
Delocalized Electrons (Again)
* *bonding* electrons spread out over several atoms
120
# Exceptions to the Octet Rule
Expanded Octet
* More than 8 electrons
* Happens in elelments that have n > 2
* Have s-->d sublevels
121
# Valence Bond Theory
Sigma (𝞂) Bond
* 1 covalent bond where orbitals overlap
122
# Valence Bond Theory
Domain of Double Bond
* 1
123
# Valence Bond Theory
Hybridization
* Promotion of electrons and mixing of orbitals
* Only central atoms
124
# Valence Bond Theory
Pi (𝛑) bond
* For bonds after single bonds
* ex) Double or Triple
* Bonding pairs are parallel to each other
125
# Resonance Strutures
Relation to delocalized atoms
* Having them makes a structure more stable and the more you have the more stable because the lectrons are able to spread out over a larger volume
126
# Polarizability
What is Polarizability?
* Ability to induce another nonpolar molecule to have a momentary dipole between instantaneuous and indused dipole
* Happens with Covalent Compounds
127
# IMF Trends
LDF Trends
* Increased = Increased size
* Increased = Increased polarizability
* Increased = Greater Charge Separation
128
# Dipole-Dipole forces
What are they?
* Polar molecules with positive "end" and negative "end" which have a permanent dipole which attracts other dipoles when closer
* Stronger than LDF
129
# Important Note!
What to do with LDF on Test
* Write full form: London Dispersion Forces
130
# Hydrogen bonds
Conditions
1. H has to be attached to F,O,N
2. Other end is also F,O,N with/without H
131
# Hydrogen bonds
What are they?
* Strong and special d-d
* Have to write both d-d + h-b
* bet H proton and negative side of F,O,N
132
# Surface Tension
Definition
* Amount of energy required to stretch
* Increased surface of liquid by 1 unit area
* Higher IMFs = Higher ST
133
# Surface Tension
Adhesion
IMFs between unlike molecules
134
# Surface Tension
Cohesion
IMFs between like molecules
135
# Viscocity
Definition
* Measure of a fluid's resistance to flow
* Increased T = Decreased V
* Increased V = Increased IMFs
* Tangling ==> Increased V
136
# H2O
Properties
* Excellent Solvent
* High Specific heat
* Ice Density is less than Water because struct for spread out
137
# Liquid-Vapor Equilibrium
Vapor Pressure
* Partial pressure of liquids once equilbrium is established between gases and liquids
* Book Definition: Gaseous molecules pressure form evap. liquid
138
# Liquid-Vapor Equilibrium
Dynamic Equilbrium
* Rate of forward process(l-->g) = Rate of backward process(g-->l)
* Constant Pressure
139
# Liquid-Vapor Equilibrium
When is Boiling Point?
* When Vapor Pressure is equal to External pressure
140
# Phase Changes
Critical Point
* Temperature that gas cant become a liquid no matter the pressure
141
# IMF Trends
How are physical characteristics affected by IMFs?(only writing abt. high)
| BP, VP, Viscocity, ST, Cohesive Forces
* Higher BP
* Lower VP
* Higher viscocity
* Higher ST
* High Cohesive Forces
142
# Hydrogen Bonding
Salicylic Acid
* H-Bonding within molecule
143
# Chromatography Lab
Rf
solute Distance/ Solvent Distance
144
# VP Curves
Equation
* ln(Pvap1, T1/Pvap2, T2)=(∆Hvap/R)((1/T2)-(1/T1))
145
# Alloys
Definition
* Solid Soln consisting of 2+ metals or metals with 1 + nm
146
# Covalent Network Solids
Definition
* Network of covalently-bonded atoms into 2D and 3D network, holding it firmly together
147
# Covalent Network Solids
Allotropes
* Elements exist in 2 diff formes in same physical state
* Ex) diamond and graphite for carbon
148
# Semiconductors
Definitions
* Elements that are normally not conductors but are at high Temperatures or when coombined with other elements
149
# Semicoductors
N-type
* Have an extra electron from bonding which can be used to create a voltage
* **Donor Impurities**
150
# Semiconductors
P-type
* Have 1 less electron which creates "positive holes" that constatnly shift and allow current to flow through holes
* Acceptor Impurities
151
# Physical Properties of Solns
Solvents
* determine soln's state of matter
* usually majority component
152
# Physical Properties of Solns
Solutes
* Substance dissolbed/dispersed in the solvent
153
# Physical Properties of Solns
Solution Concentration
* Amnt of solute/ amount of solvent/soln
154
# Physical Properties of Solns
Molality(m)
* moles of solute/ kg of solvent
* varies with mass of solvent and is *independent of temperature*
155
# Physical Properties of Solns
Enthalpy of Soln
* Combination of Heat required breaking IMFs of solute and solvent plus the heat released for letting them mix
156
# Physical Property of Solutions
Non-Ideal Solutions
* Not additive solutions(adding two solutions doesn't result in a solutions with both that is not equal than combined amount: Ex) 50mL +50mL≠100mL)
* Happens due to unequal IMFS
* example of *Ideal Mixture* = Benzene and Methylbenzene
157
# Physical Properties of Solutions
How Heat of Solution impacts ideality
* If = 0, Ideal
* If < 0, stronger solute-solvent forces; exothermic
* If > 0, weaker solute-solvent forces; endothermic
158
# Formation of a Saturated Solution
What is a Saturated Solution
* Maximum amount of solid or gas allowed in a liquid with constant Temperature
159
# Formation of Saturated Solution
How temperature affects saturation
* Increased ==> Increased in solids
* Opposite in gases
160
# Formation of Saturated Solution
Henry's Law
* Increased Solubility = Increased Pressure
* s = P * k
161
# Formation of Saturated Solution
K in Henry's Law
* Units: mg[gas]/100g[H2O]/atm
162
# Colligative vs Noncolligative Properties
What are Colligative Properties?
* Physical Properties that depends only on concentration of solute particles, not their identity
* Decreased solvent vapor pressure, Freezing Point Depression, Boiling Point Elevation
163
# Colligative vs Noncolligative Properties
Examples of Noncolligative Properties
* Color, odor, etc.
164
# Volatile Solvents
What if solute is non-volatile?
* Lower solvent volatility because less on surface
* **Raoult's Law:** Vapor Psolv = Xsolv ·P°solv
165
# Volatile Solvents
What if solute is volatile?
* Have to consider both Vapor Pressures
* Pressure[Tot] = P[solvent A] + P[solvent B]
= (XsolvA · P°solvA) + (XsolvB · P°solvB)
166
# Fractional Distillation
How does Fractional Distillation work?
1. Boil 2 liquids
2. Lower Boiling Point will boil out
3. Rise to the condenser
4. distillate
* Example: Water(50mL) + Ethanol(50mL) = mixture (< 100mL)
* Distillation = 50 mL each again
167
# Chemical Kinetics
What is it?
How fast rxns take place, facts that affect rates, and mechanisms
168
# Chemical Kinetics
Factors affecting Kinetics
1. Concentration
2. Temp
3. Surface
4. Catalyst
169
# Chemical Kinetics
what does Rate of Rxn equal?
-(1/a)(Rate of Reactant A) = -(1/b)(Rate of Reactant B) = (1/c)(Rate of Reactant C) = (1/d)(Rate of Reactant D)
170
# Rate Law
What does Rate Law equal?
= k[A]^m[B]^n
* for aA + bB --> cC + dD
* m≠a and n≠b
* m = order of rxn wrt A
* n = order of rxn wrt B
171
# Rate Law
Overall Order of Rxn
m+n
172
# Rate Law
K
* Proportionality constant
* rate constant
* function of T
173
# Reaction Rate
Methods to monitor change in reactant/product concentration
1. Change in color using spectrophotometer
2. Change in Pressure with manometer
3. Change in Electrical Conductance
174
# Reaction rates
Zero Order
* Rxn Rate: r=k
* Integrated Rxn Rate: [A]t = -kt + [A]0
* Half Life: ([A]0)/2k
175
# Reaction Rate
First Order
* Rxn Rate: r=k[A]
* Integrated Rxn Rate: ln[A]t= -kt + ln[A]0
* Half Life: .693/k
176
# Reaction Rate
Second Order
* Reaction Rate: r=k[A]^2
* Integrated Reaction Rate: 1/[A]t = kt + 1/[A]0
* Half Life: 1/([A]0* k)
177
# Activation Energy
What is it?
Energy barrier that prevents less energetic molecules from reacting
Separatesineffective and effective collisions wrt energy
178
# Activation Energy
Area under Maxwell-Boltzmann
* Represents fraction of collisions that are effective with correct orientation
* A= e ^(-Ea/RT)
* R = 8.314 J/mol/k
* T = K
179
# Activation Energy
Rate constant equation with Activation Energy
* k = A * e ^(-Ea/RT)
* A = frequency factor = constant of soln to find rate constant
* R = 8.314 J/mol/k
* T = K
180
# Activation Energy
To compare K values at different temperatures...
ln(k1/k2) = Ea/R(1/(T1) - 1/(T2))
181
# Reaction Mechanism
What is it?
* Sequence of steps(each step = elementary step) and sum of all steps
* In the ES, rxn order for each reactant = stoic coeff.
182
# Reaction Mechanism
Rate determining step
* Slowest ES = Speed of overall rxn
183
# Reaction Mechanism
How to determine plausible mechanism
1. ES must sum up to overall rxn
2. Rate-determining step must have same rate law as overall exp.aly-det. rate law
184
# Reaction Mechanism
Intermediate
* Produced in ES and consumed in later ES
185
# Reaction Mechanism
Rate determining step senarios
* If the 1st step slow: 1st = rate of rxn and rest fast
* If 2 + steps slow: Step right before slow step = equilibrium and slow step = RDS
186
# Activation Energy - Catalyst
What do they do?
* Increases rate of rxn(faster rxn) by decreasing Ea
* Present at beginning and end of rxn
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# Activation Energy - Catalyst
RDS relation to Ea
Highest Ea = RDS
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# Equilibrium
Chemical Equilibrium
Occurs when opposing rxn proceeds at equal rates
* Concentrations dont change with time making rxn appear stopped
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# Equilibrium
What is K?
* Equilibrium Constant
* = [A]eq/[B]eq = Pressure of A eq/ Pressure of B eq
* At Constant Temperature
* K >>1(>10): Equilibrium favors product formation
* K <<1( <.10): Equilibrium favors reactant formation
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# Equilibrium
Law of Mass Action for general rxn
* Kc = [C]^c * [D]^d/([A]^a * [B]^b)
* Kp same just with P instead of [ ]
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# Equilibrium
Rules
* Solids not considered
* Liquid water when in excess not considered
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# Equilbrium
Converting from Kc-->Kp
* Kp=Kc(RT)^∆n
* R = .08206
* delta n = moles of product gas - moles of reactant gas
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# Modifying Equilibrium Expressions
Types of Modification
1. Reversing Equation: K'c = 1/Kc
2. Multiplying Equation by x: K'c = Kc^x
3. Adding 2 equations: K'c = Kc1 * Kc2
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# Rxn Quotient
What is it?
* Qc/ Qp
* Concentration ratio for Kc/Kp at non equilibrium conditions
* Not constant but allows us to predict direection of net change
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# Rxn Quotient
How it helps us predict
* Q > K ==> backward rxn favored
* Q < K ==> forward rxn favored
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# ICE Table
What are they?
1. I: Initial [ ] or P
2. C: Change over time
3. E: When at Equilibrium
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# Le Chat's principle
What is it?
* Predicts how eq'm restored
* System will react to stress
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# Le Chat's Principle
Adding or Removing a Substance in rxn
* Adding: system will shift and increased rate of rxn that decreases [substance]
* Removing: system will shift and decrease rate of rxn that decreases [substance]
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# Le Chat's Principle
Stress sys by changing T
* Exo:
1. Increased T = Increased R and Decreased P
2. Decreased T = Increased P and Decreased R
* Endo:
1. Increased T = Increased P and Decreased R
2. Decreased T = Increased R and Decreased P
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# Le Chat's Principle
Change P or V of Sys
* As P increases or V decreases(decrease space)
* eq net shift to side with less number of moles of gas
* As P decreases or V increases(increase space)
* Shift to side with more moles of gas
* If number of moles reactant = number of moles product, then there is no shift
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# Le Chat's Priniciple
Special Cases
1. Partial Pressures of reacting changes does change with adding or removing inert gases
2. In a rigid container: Total Pressure will increase but Partial Pressures don't(no shift)
3. Constant External Pressure[balloon/ moveable piston]
4. Same as ∆Pressure
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# Le Chat's Principle
Changes that will not cause Eq. Shift
1. Adding Inert Gas to constant volume container
2. Catalyst
3. Pure solids or Liquids
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# pKw, pH, pOH
Equations
* pH= -log([H3O+])
* Also H+ = 10^-pH
* pOH= -log([OH-])
* OH- = 10^-pOH
* pKw = pH + pOH = -log([OH-] [H3O+])
* Ka * Kb = Kw
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# pH, pOH
How to decide concentration of H3O+ or OH-
* if in strong acid just that from the acid because self-ionization of water is negligible
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# Conjugate Base/Acid
Hydrolysis Rxns
* Acid - base rxns between ions(from salts) and H2O molecules
* May react with water in a/b rxn
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# Conjugate Base/Acid
pH vs H3O+ / OH- concentration
* 7: [H30+] = [OH-]
* <7: [H30+] `>` [OH-]
* >7: [H30+] < [OH-]
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# Conjugate Base/Acid
Determining Conjugate Base/Acid
* If acid, then find with extra electron (goes from donator to acceptor) ⇒conj base
* If base, then find with less electron (goes from acceptor to donator) ⇒ conj acid
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# Acid, Bases, and Acid-Base Equilibria
Acid
* Proton Donor
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# Acid, Bases, and Acid-Base Equilibria
Base
* Proton Acceptor
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# Acid, Bases, and Acid-Base Equilibria
Strong Acids
1. HCl
2. HBr
3. HI
4. HNO3
5. H2SO4
6. HCLO4
* Complete Ionization: →
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# Acid, Bases, and Acid-Base Equilibria
Weak Acids
1. CH3COOH
2. NH4+
* Partial Ionization: ⇌
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# Acid, Bases, and Acid-Base Equilibria
Strong Bases
* Group 1A and 2A with Hydroxides
* Complete Ionization: →
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# Acid, Bases, and Acid-Base Equilibria
Weak Base
* NH3
* Partial Ionization: ⇌
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# Acid, Bases, and Acid-Base Equilibria
Ka/ Kb
* Acid/Base Ionization Constant
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# Acid, Bases, and Acid-Base Equilibria
Relation bet a/b and conj a/b
* Stronger the original, weaker the conjugate and vice versa
* Favored in the direction = stronger to weaker
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# Acid, Bases, and Acid-Base Equilibria
Kw
* Ion product of H2O
* = [H3O+][OH-] = 1.0 * 10^-14 M
* both H3O+ and OH- have concentrations of 1.0 * 10^-7
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# Acid, Bases, and Acid-Base Equilibria
How to deal with weak Acid/ Base equilibria
* Use ICE Table
* If M(a/b)/K(a/b) > 100, then [A/B]≈ Eq. value
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# Acid, Bases, and Acid-Base Equilibria
Which ions hydrolyze
* Weak acids or bases hydrolyze appreciably
* Strong a/b form neutral solutions
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# Acid, Bases, and Acid-Base Equilibria
Relationship between weak and strong a/b
* Weak Acids and Strong Bases = basic
* Strong Acids and Weak Bases = acidic
* Weak a + b = any of them
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# Common Ion Effect
Common Ion Effect
| ASK ABT THIS
* Supression of ionization of weak acids or weak bases by presence of common ion from a strong electrolyte
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# Buffer Solutions
Definition
Solution that changes pH only slightly when small amounts of strong acid or base added
* Usually contain: weak acid/base with salt
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# Buffer Solutions
Buffer Rules
1. ratio of [conjugate base/ acid] / [weak acid/ base] between .10 and 10
2. both [conjugate base/ acid] and [weak acid/ base] exceed Ka by factor of 100+
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# Buffer Solutions
Henderson–Hasselbalch equation
pH = pKa + log([conj base]/ [weak acid])
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# Buffer Solutions
Heylman equation
pOH = pKb + log([conj acid]/ [weak base])
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# Buffer Solutions
Buffer Action
* best buffer is when [conj acid / base] = [weak base / acid] → best buffer!
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# Buffer Solutions
Effective Buffer Range
* Approximately when pH = pKa ± 1 pH-unit
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# Acid-Base Titrations
Analyte
* solution of unknown concetration
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# Acid-Base Titrations
Titrant
Solution with known concentration
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# Acid-Base Titrations
Equivalence point
| AKA Stoichiometric pt
* # of moles of H3O+ = # of moles of OH-
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# Indicators
What are they
A weak organic acid that has a different color than conjugate base
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# Indicators
Relation bet. [HIn] and [In-]
* equal means intermediate color
* [In-]/[HIn] >10: solution will be color of In-
* [In-]/[HIn] < .10: solution will be color of HIn
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# Indicators
Endpoint
* Point where indicator changes color
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# Indicators
Equation
* Ka/[H3O+] = [In-]/[HIn]
