Chemistry Flashcards
(178 cards)
0
Q
How many molecules are in a mole?
A
6.02x10^23
1
Q
1 AMU = ____ kg
A
1.66x10^-27 kg
2
Q
What are diamagnetic and paramagnetic orbitals?
A
Diamagnetic - contains only spin paired electrons
Paramagnetic - 1 or more electrons are not spin paired
Note: odd number of electrons must be paramagnetic but even number of electrons can be either one
3
Q
Quantum number table
A
Page 247
4
Q
Degenerate orbitals
A
Orbitals that have identical energies (ie 2p orbitals)
5
Q
How are electrons placed into orbitals?
A
From lowest to highest energy
6
Q
What is an excited state electron?
A
Not a ground state configuration
Contains the same number of electrons as the element and does not violate filing rules (number of e per subshell)… Fill diff orbitals
7
Q
Isoelectronic ion
A
Ion that has same number of electrons as another
8
Q
Periodic trends
- Electronegativity
- Atomic radius
- Ionization energy
- Metallic character
- Electron affinity
A
- F>O>N>Cl>Br>I>S>C~>H»>Fr
- Decreases L-R, increases T-B
- Increase L-R
- Decreases L-R
- Increases L-R
9
Q
Ionization energy and bonds
A
Difference in E
E>1.7 = ionic bond
1.7>E>0 = polar covalent
0 = nonpolar covalent
10
Q
Equation for a quanta
A
[]E=hf Or []E=h(c/wavelength) h is planck's constant c is speed of light f is frequency of electromagnetic radiation used to increase energy of the electron
11
Q
Heisenberg uncertainty principle
A
Position and momentum cannot be determined simultaneously at all times for an election and therefore it’s orbit is not spherical (p 251)
- combined with schrodinger wave equation, gives rise to probability density orbitals
12
Q
De broglie hypothesis
A
Any particle with momentum should have a wavelength
Wavelength=h/p or w=h/mv
h is planck’s constant
P 252
13
Q
Radioactive decay process (what is it)
Decay type-symbol-change in mass number-change in atomic number
-antimatter electron?
A
Page 252/253
14
Q
Einsteins equation for nuclear biding energy and mass deficit
A
E=mc^2
15
Q
Types of intermolecular bonds by strength
A
ii>id>dd>iid>did>idid (LDF)
H bonding is from ii to did
16
Q
What must be present for h bonds to occur
What does h bonding do to bp?
A
H bonded to F,O, or N
Lone pair of electrons
-increase it
17
Q
Ammonium
A
NH4+
18
Q
Nitrate
A
NO3-
19
Q
Carbonate
A
CO3^2-
20
Q
Sulfate
A
SO4^2-
21
Q
Phosphate
A
PO4^3-
22
Q
Strength of repulsive forces in VSEPR model
A
Lone pair-lone pair>LP-BP>BP-BP
23
Q
1BP 0LP
A
Linear
24
2bp, 0lp
Linear
25
3lp, 0bp
Trigonal planar
26
4bp, 0lp
Tetrahedral
27
3bp, 1lp
Trigonal pyramidal
28
2bp,2lp
Bent
29
1bp,3lp
Linear
30
5bp,0lp
Trigonal bipyramidal
31
4bp,1lp
Seesaw
32
3bp,2lp
T-shaped
33
2bp,3lp
Linear
34
6bp,0lp
Octahedral
35
5bp,1lp
Square pyramidal
36
Equation for moles ad mola mass, etc
n=m/MM
MM is molar mass
m is mass
n is moles
37
Does the limiting reagent limit the rate of the reaction?
No, it limits the extent... Ie how much products will form
| - see how to calculate limiting reagent p 280
38
Explain oxidation rules
1. Group 1A +1
2. Group 2A +2
3. Group 3B +3
4. Fluorine -1
5. Oxygen -2 (-1 if peroxide)
6. Hydrogen +1 when bonded to nonmetals, -1 bonded to metals
39
Oxidizing reducing agent and oxidant reductant and oxidized reduced
How to balance oxidation/reduction reactions
Page 281
40
What is kinetics
Study of chemical reactions focusing on SPEED and MECHANISM.
| Irrelevant - amount of products formed, equilibrium amount of products/reactants, thermodynamics, stoichiometry
41
What determines the reaction mechanism that prevails in a chemical reaction?
```
Concentrations
Pressure
Temperature
Absence or presence of a catalyst
Also microscopic reversibility
```
42
Unimplecular vs bimecular process
An elementary process having one or two reactant molecules
43
Activation energy and rate-determining step
The rate determining step is the slowest step in the reaction due to highest activation energy.
For any mechanism the step with the highest Ea is the slowest step.
44
How to calculate reaction rate
For aA + bB -> cC + dD
Rate = -[][A]/a[]t = -[][B]/b[]t = [][C]/c[]t = [][D]/d[]t
Note the - sign!!
45
How to calculate rate due to concentration
```
From aA + bB -> cC + dD
Rate = k[A]^x[B]^y
Overall order = x+y
x and y not related to coefficient
k is temperature dependent and does not change with time
- see p 288 for info about K
```
46
How to convert log
logx=y
| x=10^y
47
What do catalysts do?
Alter activation energy of reaction
Allows reaction to occur via alternate pathway
Inhibits (inhibitory catalyst) or promotes reaction (slows it down or speeds it up)
See p 290 for homo vs heterogenous catalysts
48
What can change a reactions equilibrium constant?
A change in temperature and in some cases pressure and volume
Position of equilibrium is changed in other cases (Le Chatelier's principle) to fix the change made but Keq stays the same.
Note: increase in T, p or v does not necessarily mean increase in Keq
49
Equilibrium constant equation
```
Keq = [C]^c[D]^d/[A]^a[B]^b
In molarity (M)
Pure liquids and pure solids don't appear in equation because they tend to have concentration of 1M
```
50
Finding equilibrium constant for gasses
Kp=Kc(RT)^[]n
Kp is equilibrium constant calculated from partial pressures
Kc is Keq calculated from molar concentrations
R gas constant
T absolute temp
n change in total number of moles of gas from reactants to products
51
Ratio of products to reactants and Keq
Keq>>1 products favoured
Keq=1 neither
Keq<<1 reactants favoured
52
Reaction quotient
Q = (C)^c(D)^d/(A)^a(B)^b
| Brackets indicating concentrations
53
Relationship between Q and Keq
Q>Keq products in excess, decrease in products and increase in reactants as reaction approaches eq
Q=Keq at equilibrium
Q
54
What would be he equilibrium for the reverse reaction?
Keq(reverse) = 1/Keq(forward)
55
How to determine Keq for the reactions
aA+bB->cC+dD = K1
cC+dD->eE+fF = K2
aA+bB->eE+fF
| Keq = K1(K2)
56
How does addition or removal of product or reactant affect equilibrium? Addition of non-reactive product?
Le Chatelier's principle (pg 300)
Note this does not chance value of Keq!!
If unreactive substance is added, no change in pp or Keq, but increase in total pressure
57
Describe pH and equilibrium wrt:
| HA H+ + A-
Increase H+ lowers ph, shifts equation left and decreases A- (the conjugate base)
Decrease H+ raises ph, shift equation right and increases A-
58
What happens when volume is increased or decreased in a reaction?
If volume is decreased (p increased) reaction is pushed in direction with less moles of gas
Opposite for increased v
If both sides contain equal moles of gas, no change in eq if pressure is changed
-wrt adding a gas, a nonreactive gas will have no affect on equilibrium and adding a reactive gas will follow le chataliers principle
59
Explain the effects of temperature on equilibrium
Determines whether it's endothermic or ectothermic and then use Le Chatelier's principle
60
Open vs closed vs isolated system
Open - energy and matter exchanged
Closed - energy only
Isolated - neither
Note universe = surroundings + system
61
Name the state and nonstate functions
Nonstate - work and heat
State - everything else
Note: nonstate functions depend on the path and therefore deal with the concepts of reversibility and irreversibility (ie can't unable a cake)
-p 309
62
What are the 2 principal ways a system can gain or lose energy?
By heat transfer or by work
63
What is the equation for work in thermodynamics?
W=P[]V
P is pressure
V is internal volume change by system
Note: this is mechanical work done by expanding gasses - reversible process always related than irreversible process (like heat)
64
Describe positive and negative heat and work
- Q system (+Q surroundings) when heat is transferred from system to surroundings (and vice versa)
- W when work is done by system, +W when it's done on system
65
What is greater (wrt heat and work) reversible or irreversible process?
Q - reversible process>irreversible process (whether given off or gained)
W - performed by reversible process>performed by irreversible process (does not matter if work is done on system or surroundings)
66
Internal energy of a system
```
Total energy within a system (kinetic and potential energy) expressed in relative terms
[]U=Q-W and U=Q+W
U is change in internal energy
Q is heat GAINED by system
W is work done by or on system
```
67
What is entropy? When is it increasing?
```
Measure of disorder
Entropy increases when:
1. Temp increases
2. In a reaction if the reaction produces more product molecule than it contained reactant molec.
3. Pure liquids or solids form solutions
4. S(universe) always increases
```
68
Equation for enthalpy and it positive/negative values
Measure of heat release when pressure is held constant
[]Hsys=Qsys,p
Qsys,p is heat absorbed by the system at constant pressure
[]H = []U + [](PV)
+ []H means endothermic
- []H means exothermic
69
Converting between Kalvin and degrees
0 C = 273.15 K
70
Heat of formation
Enthalpy required for the formation of one mole of that substance from its elements
0 for all elements in their naturally occurring state
71
Standard heat of formation
Heat of formation in standard conditions (T=298.15K, p=1atm)
72
STP (standard temperature and pressure)
T=273.15, p=1atm
73
Heat of reaction
[]H^o={Hf products- {Hf reactants
Endothermic is +
Exothermic is -
74
Why is heat of formation of liquid water different from that of gaseous water?
Extra enthalpy is given off in he Hf of liquid water as result of transition from higher to lower energy state.
Known as enthalpy of liquefaction or enthalpy of condensation.
75
Gibbs free energy
```
[]G=[]H-T[]S
T is constant
All thermodynamic quantities refer to the system
Measure of energy available to that system for the performance of useful work
-[]G is exergonic (spont. Forward)
+[]G is endergonic (non-sp. forward)
0 means reaction is at equilibrium
Know the chart!!
[]H - T []S
Endothermic - high (+) = -
Exothermic - high (-) = +
```
76
3 laws of thermodynamics
1. For any isolated system, energy is constant ([]U=Q+W)
2. Entropy is always increasing - entropy in the forward direction does not equal entropy in the reverse direction - heat is never completely converted to useful work in a cyclic process (it's possible in a linear one) ie if the forward process does so, the reverse will not
3. Absolute zero is unattainable
77
Know how to draw reaction diagrams
| Activation energy, transition state, reaction intermediate, sn1/sn2 reactions, rate determining step, catalyst
Pg 320
78
Note: limiting reagent is not necessarily rate determining step!
Because LR has little to do with kinetics ...it just gets used up first because there is less of it, it doesn't effect the speed
79
Does a catalyst affect []H or []G?
No, it just changes the activation energy. It also does not change equilibrium. It lowers the transition state.
80
Finding []G^o at equilibrium
```
[]G=[]G^o + RTlnQ
But []G at equilibrium is 0 so:
[]G^o = - RTlnKeq
[]G^o is at 298K and 1atm
R is universal gas constant
T is temperature
Q is the equation from before (calculated at diff temperatures)
+ []G^o favours reactants
- []G^o favours products
0 []G^o favours neither
```
81
Heat of solution
Enthalpy change that occurs when a solute dissolves in a solvent.
- exothermic solvation (heat released)
+ endo, heat absorbed
82
Henry's law (solubility of gases in a liquid solvent)
p=kC
C is solubility of gas
k is Henry's law constant for the gas
p is partial pressure of gas solute over the solution
Note: solubility of gas is proportional to pressure of gas above the solution
83
How to know if solid is soluble in liquid
If the solid attains:
>0.05M - soluble
<0.01M - insoluble
Btw .05 and .01 - moderately soluble
84
Molarity
M=moles of solute/liters of solution
85
Normality
Used almost exclusively for acids and bases
| N=(n of potential H+ or OH-)/liters of solution
86
Molality
m=(moles of solute)/(kg of solvent)
87
Mole fraction
Usually used with gases
| Xs=(moles of solute)/(total miles of solution)
88
Molarity of water
55M
89
Equilibrium constant equation for solubility (solubility constant)
Keq=Ksp=[A]^a[B]^b
From: AaBb(s)aA(aq)+bB(aq)
Ksp will never have a denominator because it is a pure liquid/solid and thus does not appear in Ksp
90
Examples of weak electrolytes
Weak acids and bases
91
Examples of strong electrolytes
Salts of alkali metals and ammonium
| Some salts of alkaline earth metals, all nitrates, chlorates, perchlorates, and acetates
92
What substances are solvated (not dissociated)
Methanol, ethanol, glucose and fructose
93
Van't Hoff factor
(i) how many species a substance dissociates into
i(NaCl)=2 i(CaCl2)=3
Glucose=1
94
Raoult's law (how vapor pressure of a pure liquid is lowered by the addition of a nonvolatile solution)
```
pi = Xi(pi^o)
pi is vapor pressure of solvent i above the solution i
Xi is the mole fraction of solvent I
pi^o is vapor pressure of pure solvent i
or
Psoln=Xsolv(P*solv)
In the case of a nonvolatile solute dissolved in a solvent
P 339/360
```
95
What happens to boiling point of a pure liquid when nonvolatile solute is added?
```
It increases
[]Tb=Kb(i)(m)
[]Tb is boiling point elevation
Kb is bp elevation constant
i is Van't Hoff factor for solute
m is molality of solute
P 340
```
96
What happens to freezing point of a pure liquid when a solute is added?
```
It decreases
[]Tf=Kf(i)(m)
[]Tf is freezing point depression
Kf is fp depression constant
i is Van't Hoff factor for solute
m is molality of solute
Page 341
```
97
Osmotic pressure
Pressure exerted on the membrane when equilibrium is reached
O = nRT/V or O=iMRT
n is number of moles of particles dissolved
R is the gas constant
T is the absolute temperature
V is the volume of the solution
98
Air pressure conversions
```
P=F/A=Pa
1atm=
760mmHg=
760torr=
1.013x10^5Pa
```
99
One mole of ideal gas at STP occupies what volume?
STP= 0C at 1atm
22.4L or 0.0224 cubic meters
(Note that this value is 25C in the standard thermodynamic state)
100
5 premises of kinetic molecular theory
For IDEAL gas behaviour
1. Gases are composed of molecules in rapid, random, translational motion (Ek=1/2mv^2)
2. These particles undergo perfectly elastic collisions
3. Space occupied by particles is negligible
4. There are no attractive or repulsive forces btw particles
5. Ek=3/2nRT -n being for moles of gas particles
- if deviation from any of these laws occurs, it is not an ideal gas
101
Boyle's law
PV=constant
or
P1V1=P2V2
For ideal gases under constant temperature and with constant moles
102
Charles's law
```
V/T=constant
or
V1/T1=V2/T2
or
V1T2=V2T1
For ideal gases under constant n and P
```
103
Gay-Lussac's law
```
For ideal gasses under constant n and V
P/T=constant
or
P1/T1=P2/T2
or
P1T2=P2T1
```
104
Combined gas law
```
For ideal gases at constant n
PV/T=constant
or
P1V1/T1=P2V2/T2
or
P1V1T2=P2V2T1
```
105
Avogadros hypothesis
```
For ideal gases at constant P and T
V/n = constant
or
V1/n1=V2/n2
or
V1n2=V2n1
```
106
Ideal gas law/gas constant
R=PV/nT
or
PV=nRT
R is in J/mol•K
107
What conditions favour ideal gases?
Low pressure, high temperature
| High P and low T lead to deviations from ideal gas law
108
What are the deviations from ideal gas law wrt P and V?
| - excluded volume
Actual pressure is always less than or equal to pressure calculated by ideal has law
Actual volume occupied by a gas is always grater than or equal to the volume calculate by ideal gas law
- the actual volume of the gas is equal to the volume of the container plus the excluded volume (volume of gas molecules themselves)
109
Vanderwalls equation
(P+an^2/V^2)(V-nb)=nRT
What does this signify?
Adjustment for real gasses wrt ideal gas law
Ideal P = realP + an^2/V^2
Ideal V = realV - nb
The sections after the +&- in each equation is the correction factors
110
Law of partial pressure
```
Ptotal=Pa+Pb
Pt=(na+nb)RT/V
Pa=Xa(Pt)
Xa=na/nt
Xa is the mole fraction of gas a
```
111
Graham's law of diffusion (effusion)
v1^2/v2^2=m2/m1
or
v1/v2=_/(m2/m1)
v is velocity, m is mass
-the lighter molecule will have a higher velocity (rate of diffusion) than the heavier one
-diffusion is movement of gas through space and effusion is movement if gas under pressure through a small hole
112
The conversion from solid to gas is called
Sublimation
113
The direct conversion from gas to solid is called
Deposition
114
What is the formula for heat absorbed by a system when it isn't undergoing a phase change?
```
q=mc[]T
q is heat a absorbed
m is mass of substance
c is specific heat capacity (cal/g•C) - water is 1 cal/g•C
[]T is temp change IN C!!
```
115
Formula for heat of a phase change
q=m[]H
q is heat
m is mass
[]H is the enthalpy of the phase change in units of energy per mass
116
How to find C (specific heat capacity for exact amount of substance being considered) from a phase change diagram
Slope=1/C
mc=C and 1/slope=mc
P 375
117
How to find heat of fusion for phase change diagram
Length of horizontal line = []Hfus•m
| m is mass of substance
118
Plot a phase diagram
Pg 379
119
What is the diff for a pt diagram of water?
Ph 380
120
Draw a pt diagram for a substance with more than one solid form
Pg 381
121
What happens if the triple point of a substance occurs at pressure >1atm?
It will sublimate under atmospheric conditions and appropriate temp change (example co2 at 216.8 K and 5.11 atm)
122
Describe he critical point
Beyond this temperature, there is no way to distinguish btw liquid and gas (no matter what the pressure is)
Defined at critical temp and critical pressure
123
Equilibrium expression for auto ionization of water
Kw = [H+][OH-] = 1x10^-14
and
[H+] = [OH-] = 1x10^-7
This Kw value is valid for any aqueous solutions at 25C (not just pure water)
124
Formula for pH/pOH of a solution
```
pH=-log[H+]
pOH=-log[OH-]
and
pH+pOH=14
(-log[H+])+(-log[OH-])=(-logKw)
```
125
How to transform log values
[H+]=n•10^-x
if n=1, pH = x
if n=3.17, pH=x-0.5 (bc root 10=3.17)
if n is higher or lower than 3.17, use x-0.5 and evaluate (pg 389)
126
Arrhenius, bronsted-Lowry, and Lewis definition of acids/bases
1. Substance that produces H+ or OH- ions respectively
2. Proton donor or accepter respectively
3. Electron acceptor or donor respectively
- water is neither acid or base in Arrhenius terms but in BL it can be either an acid or a base
127
Define strong acid/base
Strong A/B fully undergo forward dissociation in aqueous solution. Dissociation equilibrium lies entirely to the right and reverse reaction doesn't occur to an appreciable degree. No appreciable amount of original A/B left in solution.
128
Examples of strong acids (6)
```
HCl hydrochloric acid
HBr hydrobromic acid
HI hydroiodic acid
HNO3 nitric acid
HClO4 perchloric acid
H2SO4 sulfuric acid
```
129
Examples of strong bases (6)
```
LiOH lithium hydroxide
NaOH sodium hydroxide
KOH potassium hydroxide
RbOH rubidium hydroxide
(NH2-) amide ion
(H-) hydride ion
```
130
What makes acids more acidic?
Electron withdrawing groups, which stabilize their conjugate bases (making the conjugate base less willing to protonate and thus, weaker)
131
Acid/base ionization constant expression
HA + H2O H3O+ + A-
Ka = [H3O +][A-]/[HA]
Because water is a pure liquid.
Same thing for base
132
What is the pKa/Ka relationship?
pKa=-logKa and pKb=-logKb
Strong a/b have large Ka/Kb and small pKa/pKb values
Weak a/b have small Ka/Kb and small pKa/pKb values
133
2 other equations used for any acid/bad pair in dilute aqueous solution
KaKb=Kw
pKa + pKb = 14
For any conjugate acid/base pair in a dilute aqueous solution
134
Hendeson Hasselbalch equation
Useful in titrations
pH = pKa + log[A-]/[HA]
pOH = pKb + log[BH+]/[B]
135
What happens when you mix a strong acid and a strong base?
They neutralize. This process is always ectothermic. If there is excess acid or base, solution will be acidic or basic respectively and not completely neutral.
HA + BOH -> H2O + B+ + A-
Creates water and salt
136
What is the equivalence point?
The point at which the acid has been neutralized by the base added. Exactly one equivalent of base has been added for one equivalent of acid. This is 7 for strong acid/base titration.
137
What is the equivalence point for strong/weak a/b?
Equivalence point of weak acid titrated with strong base will be pH >7. For weak base titrated with strong acid, <7.
Recall the equivalence point is when equal amounts of a/b have been added. This effect is due to the excess of conjugate a/b, with that of strong a/b act as spectators (unreactive salts).
138
What is the half-equivalence point?
Point when the amount of a/b added is exactly half of the base needed to reach equivalence point.
At this point, referring back to HH equation, the fraction =1 and therefore:
pH=pKa or pOH=pKb
139
What is log 1?
0
140
Where is the buffering region an what is always found within it?
It is the first part of the weak/strong a/b titration curve and always includes the half-equivalence point.
141
What is a polyprotic acid and what does its titration curve look like? What are indicators?
Pg 404/405
142
A more concentrated buffer is stronger or weaker?
Stronger
143
What is a buffer composed of?
A weak acid/base and it's corresponding conjugate
144
Describe the pH and pKa of a buffer
For a buffer, pH=pKa
145
How do you write redox couples?
Oxidized form/reduced form
Ie: Cu++/Cu for the reaction:
Cu++ + 2e- -> Cu^0
Think of the redox couples as similar to acid/conjugate base pairs
146
Describe the anode and cathode
- electrode and inert electrode
- direction of current
Cathode is the electrode at which reduction occurs
Anode is the electrode at which oxidation occurs
Electrons flow from anode to cathode and therefore current moves from cathode to anode
-p 415
147
How to calculate electromotive force and cell emf
| -what can alter emf?
Emf is E^o at standard conditions of 1M for solutes, 1atm for gasses, and 25C. Standard half cell is for H+
Cell emf is the sum of all the half reactions. Don't multiply it by stoihiometric coefficient!
P 418
148
How to draw cell diagrams
Oxidation reactant, oxidation product || reduction reactant, reduction product
Fe2+(aq), Fe3+(aq) || Mn2+(aq), Mn7+(aq, 0.5M)
If conditions are not standard, concentration may be shown in parenthesis
Note that with a salt bridge, (Pt) is added to each end to represent the platinum, unreactive conductor (electrodes)
149
Explain a galvanic cell
Redox reaction is spontaneous
The cell creates e- flow
Oxidation is at anode & reduction is at cathode
Cathode is the positive electrode (anode is -)
Electrons flow from anode to cathode
150
Describe electrolytic cell
Redox reaction is nonspontaneous
The cell requires e- flow
Oxidation is at anode & reduction is at cathode
Cathode is the - electrode (anode is positive)
Electrons flow from anode to cathode
151
Explain what a concentration cell is
Page 425
152
Equation relating []G to E^o
[]G^o=-nFE^o
n is moles of e-
F is Faradays constant
[]G=-nFE for nonstandard conditions
153
Nernst equation
```
E = E^o - 2.303(RT/nF)logQ
R and F are constants
also
E^o = (0.0592/n)logKeq
For equilibrium under standard conditions (298K and 0E)
```
154
What direction is the arrow pointing in a dipole moment?
The head points towards the negative charge
155
What is a formal charge?
The overall charge on a molecule (ie +1). Partial charge is the dipole moment, the uneven distribution of electrons.
156
Energy and bond formation
Breaking bonds require energy, forming bonds release energy
157
What is (x)(2x)^2?
4x^3
158
Explain how to represent where atomic mass, atomic weight, charge, and atomic number are located
P 242
159
What is a period/group?
Periods are horizontal and groups vertical
Periods = same principle quantum number (shell number)
Groups = same valence electrons and similar physical properties
160
Describe Pauli exclusion principle, hunds rule, and the aufbau principle
```
P.e. = no electron in any one atom will have the same 4 quantum numbers as any other electron in that atom
H.r. = no 2 electrons will become spin paired unless there are no empty orbitals at that energy level
A.p. = how electrons are placed Ito orbitals from lowest to highest energy
```
161
What are representative (main block elements) vs transition metals? What are actinides/transuranium elements?
251
162
What are the 5 main types of reactions?
P 275
163
Describe how to make hybridized orbitals
Chapter 33 (731)
164
Describe the types of isomerism/properties
Chapter 34 (731)
165
Describe energy wrt breaking and forming bonds
Breaking bonds requires an input of energy and forming bonds an output
166
What is microscopic reversibility?
The principle that all elementary steps are reversible in exactly the reverse manner (even though the overall reaction is not exactly reversible)
167
Solubility general info
P 329
168
Look at solubility rules
P 332
169
What is solubility affected by?
Temperature, common ions, pH
170
What are colligative properties?
P 338
171
Pressure depends on what?
Number of collisions (increase with increasing number of molecules or decreasing volume)
Velocity of colliding gas molecules (increases with temperature)
Mass of colliding gas molecules
172
Under which conditions are the solid/liquid/gas phase favoured? How are each of these phases characterized?
P 369
173
Heat of fusion, vaporization and sublimation (endo or exothermic)
P 371
174
Relate heat and phase changes to kinetic energy and temperature
P 372
175
What is a calorie? What is the specific heat of liquid water?
The amount of heat energy that must be supplied to one gram of liquid water at 25C to raise the temperature of the water by 1C
-specific heat of water is 1cal/g•C
176
What is a ligand?
Lewis bases that form bonds with transition metals
177
How to use reduction potential chart
416
