G Chem Flashcards
1
Q
Energy of a photon (equation)
A
E = hf h = Planks constant = 6.63x10^-34 Js E = hv (v in Hz) E = cλ c = 3.00 x 10^8 m/s
2
Q
EM Spectrum (order, E and f levels)
A
Gamma X-ray UV VIBYOR IR Microwave Radio
E high; f high, λ low E low; f low, λ high
3
Q
Charge of e-
A
-1.6022 x 10^-19 C
4
Q
Mass of proton
A
1.673 x 10^-27 kg
5
Q
Mass e-
A
9.11 x 10^-31 kg
6
Q
Relationship between c, λ, and v
A
c = λv
7
Q
Principle quantum number
A
n
8
Q
Angular quantum number
A
l (n-1) l = 0 s orbital l = 1 p orbital l = 2 d orbital l = 3 f orbital
9
Q
Magnetic quantum number
A
ml
-l to +l
l = 0 ml =0 s orbital
l = 1 ml = -1, 0, +1 p orbital (py, px, pz)
10
Q
Spin quantum number
A
ms
+1/2 spin up
-1/2 spin down
11
Q
Common reducing agents (4)
A
LiAlH4
NaBH4
Metals
Fe2+
12
Q
Common oxidizing agents (7)
A
Halogens Permanganate-(MnO4) salts Peroxide compounds (ie H2O2) Ozone (O3) Osmium tetroxide (OsO4) Nitric acid (HNO3) Nitrous oxide (N2O)
13
Q
Lewis acid
A
e- pair acceptor
14
Q
Lewis base
A
e- pair donor
15
Q
Bronsted Lowry acid
A
H+ donor
16
Q
Bronsted Lowry base
A
H+ acceptor
17
Q
Electron and molecular geometry
Number lone pairs: 1
Number bonding pairs: 2
A
Trigonal planar (120), bent (sp2)
18
Q
Electron and molecular geometry
Number lone pairs: 1
Number bonding pairs: 3
A
Tetrahedral (109.5), trigonal pyramidal (sp3)
19
Q
Electron and molecular geometry
Number lone pairs: 2
Number bonding pairs: 2
A
Tetrahedral, bent (sp3)
20
Q
Steric number (# hybridization orbitals in a molecule)
A
sigma bonds + # lone pairs
21
Q
Electron and molecular geometry
Number lone pairs: 0
Number bonding pairs: 2
A
Linear (180)
22
Q
Structural isomers
A
Same molecular formula, different compounds
23
Q
Stereoisomers
A
Differ in 3-D arrangement (max = 2^n) n= # chiral centers
24
Q
Enantiomers
A
Stereoisomers not superimposable mirror images
25
Diastereomers
Stereoisomers that are non-superimposable non-mirror images (differ at SOME chiral centers, not all)
26
Z configuration
Two top priority groups on same side of double bond
27
E configuration
Two top priority groups on opposite side of double bond
28
Sublimation
Solid to gas
29
Condensation
Gas to liquid
30
Freezing
Liquid to solid
31
Deposition
Gas to solid
32
Vaporization
Liquid to gas
33
Melting
Solid to liquid
34
Molality
(m) = 1 mole / 1000 g
35
Normality
(N) = one equivalent / 1 Liter
36
log base a of a
1
37
log base a (M^k)
KlogaM
38
log base a (MN)
loga(M) + loga(N)
39
log base a(M/N)
logaM - logaN
40
10^(log10(M))
M
41
First law of thermodynamics (equation)
```
ΔE = Q - W
Q= heat
W = work
```
42
Q greater than 0, heat is....?
Absorbed
43
Q<0, heat is...?
Released
44
Work done on the system, W=...?
W less than 0
45
Work done by the system, W=...?
W is greater than 0
46
Calorimetry
Q=...?
Q=...?
Q=mcΔT
| Q=mL
47
Change in Gibbs Free Energy (equation)
ΔG = ΔH - TΔS
48
Rate equation
Rate = k[A]^m[B]^n
A and B are reactants
Rate = Δ[concentration]/Δt
49
Relationship between Keq and ΔG (equation)
ΔG(naught) = -RTlnKeq
50
Hz
(1/s)
51
Wave period
T = 1/f (s)
52
Define: ionization energy
How much energy is required to remove an e- from a neutral atom in a gaseous phase (kJ or eV)
53
Relationship between E photon and BE (binding energy) and KE of photon
E = BE + KE
Therefore, hv = BE + KE
54
Higher binding energies correspond with e- in ______ energy subshells
Lower, e- are closer to nucleus, require more energy to remove
55
For gases: as temps increase, solubility...?
Decreases (gas molecules move too much)
56
Gases: increasing pressure ______ solubility
Increases
57
List forces in order of relative strength from strongest to weakest (polar covalent, non polar covalent, metallic, ionic bonds)
1 (strongest) metallic
2 ionic
3 polar covalent
4 (weakest) nonpolar covalent
58
When adding an enzyme, does the overall energy released (ΔH) change?
NO (kinetics of rxn are not directly related to thermodynamics of rxn)
59
What is an acid catalyst?
H+ ion
60
What is a base catalyst?
OH- ion
61
Define: heterogeneous catalyst
In a different phase than the reactants
62
Define: surface catalyst
Reactant molecules are adsorbed onto a solid surface before reacting with the catalyst to form product
* rate increases with the surface area of the catalyst in contact
63
Keq > 1
Reaction goes in forward direction (favors products)
64
Le Chatelier’s principle, increase pressure favors which side?
Side w/ fewer molecules
65
Define: neutralization reaction
Double replacement reaction between an acid and a base
Acid + base = H2O + salt
66
Define: combustion reaction
Redox reaction between a compound and molecular oxygen (O2) to for, oxygen-containing products (when one of the reactants is a hydrocarbon, the products include CO2 and H2O)
67
Define: disproportionation reaction
When a single reactant is both oxidized and reduced
ie: 3ClO- -> ClO3- + 2Cl-
68
Reduction half reaction
X+ + e- = X
69
Oxidation half reaction
X -> X+ + e-
70
What does Keq describe? What is excluded?
```
Molar concentrations (mol/L) at equilibrium for a specific temp
Pure solids and pure liquids are excluded
```
71
Kp (equation) (2)
Kp = (Pproducts)^coefficient / (Preactants)^coeff.
Kp = Kc(RT)^Δn
Δn = mole of product gas - mole of reactant gas
Describes the ratio of product and reactant concentrations at equilibrium in terms of partial pressures
72
What is the oxidation state of oxygen when bound to another oxygen (ie in H2O2)?
OS = 1
73
What is the oxidation state of oxygen when bound to F?
+1 (F is more electronegative than O)
74
Pressure (equation)
```
P = F/A
F = ma
```
75
When temp increases, what happens to pressure?
P increases
76
When increasing # particles, P does what?
Increases
77
When volume decreases, P does what?
Increases
78
What are the three conditions for a gas to be ideal?
1. No IMF
2. No volume for individual particles
3. Collisions between particles are perfectly elastic (KE is conserved)
79
Boyle’s law
PiVi = PfVf (at constant T)
80
Charles’s law
(Vi/Ti) = (Vf/Tf) (at constant P)
81
How much volume (L) does one mole of gas take up?
22.4L
82
Avogadro’s law
(Vi/ni) = (Vf/nf) (at constant T & P)
83
Due to Van de Walls forces between gas particles, what is the relationship of Preal to Pideal?
Preal < Pideal
84
Due to Van de Walls forces between gas particles, what is the relationship of Vreal to Videal?
Vreal > Videal
85
What are the 3 R values for ideal gas law?
```
R = 8.3145 (m^3*Pa)/(mol*K)
R = 8.3145 (L*kPa)/(mol*K)
R = 0.082 (L*atm)/(mol*K)
```
86
How do you calculate partial pressures (given total mass of total gas and percent compositions)
For each gas:
1. multiply percent composition by total mass (g)
2. divide by molar mass of gas to get mol of each gas (n)
3. Plug into PV=nRT to get P, which is total P
4. Multiply mol of individual gas as percent (ie 3 mol gas = 3%) by total pressure to get partial pressure of each gas
87
Define: vapor pressure
Pressure at which vapor and liquid state are at equilibrium (pressure is high enough for H2O molecules in liquid to reach high enough KE to change to gas, but low enough to force vapor H2O molecules to convert back to liquid)
88
Define Specific heat and give the equation
Amount of heat energy needed to raise 1g of water by 1C
```
Q = mcΔT
Q = amount of heat released or absorbed (J)
c = specific heat (amount of heat energy stored before T increases)
m = mass
```
89
Heat of fusion/vaporization (equation)
```
Q = mL
Q = amount of heat released or absorbed (J)
L = latent heat (how much heat energy needed to change phase)
m = mass
```
90
At what T does the ideal gas law start to break down?
Low T
91
If an inert gas is added, what happens to the Ptotal and partial pressures of other gases (Le Chatelier’s principle)?
Ptotal increases, but partial pressures of gas stay the same
92
Q > Keq, favors ...?
Reactants
93
Q < Keq favors...?
Favors products
94
CO(g) + 2H2(g) -> CH3OH (g)
If Q < Kc, reaction favors _____ , which favors side with (fewer, more) molecules, causing Ptotal to (increase/decrease)?
Favors products, favoring side with fewer molecules, Ptotal decreases
95
Define: hydration
Water molecules break ionic bonds to dissolve
96
Calculating solubility (g/L) requires what type of table?
ICE
97
Solubility rules mnemonics (2)
All soluble: CASH-N-GIA
Chlorates, acetates, sulfates, halides (not F), nitrates, Group 1 metals, ammonium
All insoluble: HAPpy
Hg2 2+, Ag+, Pb2+)
98
Why is water less dense when frozen than liquid?
Due to the way H bonds are oriented as it freezes, the water molecules are pushed further apart than they are in liquid water
99
Why does it take a lot of heat to raise the T of water?
Because some of the heat must be used to break H bonds between the molecules (water has a high specific heat capacity)
100
Define: evaporative cooling
As water molecules evaporate, the surface they evaporate from gets cooler because the molecules with the highest kinetic energy are lost to evaporation
101
Define: cohesion
Attraction of molecules for other molecules of the same kind (responsible for surface tension)
102
Define: adhesion
Attraction of molecules of one kind for molecules of a different kind (responsible for capillary action)
103
Define: 0th Law of Thermodynamics
No heat is transferred between two objects in thermal equilibrium (they are the same T)
104
Define: work
The energy required to move something against a force
105
What is the relationship between W, P and V for expansion or compression of gases (work done by gases for expansion/compression)?
W = -PΔV (J)
**P is the external P
106
When a system does work on its surroundings, the systems internal energy _____?
Decreases (negative work)
107
Relate W, force, and displacement (equation)
W = F x d
108
What is the work done on a gas (or by a gas) in a rigid container (no change in volume is possible)? How do changes in energy occur?
No possible change in volume means it is not possible for gas to do work because ΔV =0, so
W =0
Changes in energy must occur through other methods such as heat
109
Define: endothermic
Heat absorbed or released from a system under constant P
110
Breaking bonds between atoms requires (adding/removing) energy?
Adding
111
Forming new bonds (releases/takes) energy
Releases
112
Define: exothermic
Reactions where the products have lower PE than reactants
113
Define: 2nd Law of Thermodynamics
The entropy of the universe is always increasing for a spontaneous process
114
ΔG = ΔH + TΔS is true when what are held constant?
Temperature and pressure
115
When ΔG is negative at constant T and P, the reaction is _____ and considered (exergonic/endergonic)?
Negative, exergonic
116
When ΔH < 0 and ΔS > 0, ΔG is...?
ΔG < 0, spontaneous at all T
117
When ΔH < 0 and ΔS < 0, ΔG is...?
Spontaneous at low T (when TΔS is small)
118
When ΔH > 0 and ΔS > 0, ΔG is...?
Spontaneous at high T (when TΔS is large)
119
When ΔH > 0 and ΔS < 0, ΔG is...?
Non-spontaneous at all T (ΔG > 0)
120
Define: 1st Law of Thermodynamics
Energy can neither be created nor destroyed
121
Define: open system
A system that freely exchanges both energy and matter with its surroundings
122
Define: closed system
A system that exchanges only energy with its surroundings, not matter
123
If ΔS > 0, the reaction is (spontaneous/non-spontaneous)?
Spontaneous
124
If ΔS < 0, the reaction is (spontaneous/non-spontaneous)?
Non-spontaneous
125
Water going from a solid to liquid is an (endothermic/exothermic) process?
Endothermic, heat absorbed, ΔH positive (chemical bonds broken)
126
Water going from a liquid to gas is an (endothermic/exothermic) process?
Endothermic, heat absorbed, ΔH positive (chemical bonds broke )
127
Water going from a gas to liquid is an (endothermic/exothermic) process?
Exothermic, heat released, ΔH negative (chemical bonds formed)
128
Water going from a liquid to solid is an (endothermic/exothermic) process?
Exothermic, heat released, ΔH negative (chemical bonds formed)
129
Equation for ΔH (relating to bonds formed/broken)
ΔH = ΔH (bonds broken in reactants) - ΔH (bonds formed in products)
130
Enthalpy equation ( H, E, P, V)
H = E + PV
131
Define: critical point
Point on phase diagram where phase can coexist (very high T and P)
132
Define: triple point
Point on phase diagram where all three phases coexist (intersection point), equilibrium between phases
