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Chemistry Flashcards

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0
Q

How many molecules are in a mole?

A

6.02x10^23

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1
Q

1 AMU = ____ kg

A

1.66x10^-27 kg

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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

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3
Q

Quantum number table

A

Page 247

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4
Q

Degenerate orbitals

A

Orbitals that have identical energies (ie 2p orbitals)

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5
Q

How are electrons placed into orbitals?

A

From lowest to highest energy

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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

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7
Q

Isoelectronic ion

A

Ion that has same number of electrons as another

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8
Q

Periodic trends

  1. Electronegativity
  2. Atomic radius
  3. Ionization energy
  4. Metallic character
  5. Electron affinity
A
  1. F>O>N>Cl>Br>I>S>C~>H»>Fr
  2. Decreases L-R, increases T-B
  3. Increase L-R
  4. Decreases L-R
  5. Increases L-R
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9
Q

Ionization energy and bonds

A

Difference in E
E>1.7 = ionic bond
1.7>E>0 = polar covalent
0 = nonpolar covalent

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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
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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

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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

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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

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14
Q

Einsteins equation for nuclear biding energy and mass deficit

A

E=mc^2

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15
Q

Types of intermolecular bonds by strength

A

ii>id>dd>iid>did>idid (LDF)

H bonding is from ii to did

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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

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17
Q

Ammonium

A

NH4+

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18
Q

Nitrate

A

NO3-

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19
Q

Carbonate

A

CO3^2-

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20
Q

Sulfate

A

SO4^2-

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21
Q

Phosphate

A

PO4^3-

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22
Q

Strength of repulsive forces in VSEPR model

A

Lone pair-lone pair>LP-BP>BP-BP

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23
Q

1BP 0LP

A

Linear

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24
Q

2bp, 0lp

A

Linear

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25
Q

3lp, 0bp

A

Trigonal planar

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26
Q

4bp, 0lp

A

Tetrahedral

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27
Q

3bp, 1lp

A

Trigonal pyramidal

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28
Q

2bp,2lp

A

Bent

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29
Q

1bp,3lp

A

Linear

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30
Q

5bp,0lp

A

Trigonal bipyramidal

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31
Q

4bp,1lp

A

Seesaw

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32
Q

3bp,2lp

A

T-shaped

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33
Q

2bp,3lp

A

Linear

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34
Q

6bp,0lp

A

Octahedral

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35
Q

5bp,1lp

A

Square pyramidal

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36
Q

Equation for moles ad mola mass, etc

A

n=m/MM
MM is molar mass
m is mass
n is moles

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37
Q

Does the limiting reagent limit the rate of the reaction?

A

No, it limits the extent… Ie how much products will form

- see how to calculate limiting reagent p 280

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38
Q

Explain oxidation rules

A
  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
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39
Q

Oxidizing reducing agent and oxidant reductant and oxidized reduced
How to balance oxidation/reduction reactions

A

Page 281

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40
Q

What is kinetics

A

Study of chemical reactions focusing on SPEED and MECHANISM.

Irrelevant - amount of products formed, equilibrium amount of products/reactants, thermodynamics, stoichiometry

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41
Q

What determines the reaction mechanism that prevails in a chemical reaction?

A 
Concentrations 
Pressure
Temperature
Absence or presence of a catalyst
Also microscopic reversibility
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42
Q

Unimplecular vs bimecular process

A

An elementary process having one or two reactant molecules

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43
Q

Activation energy and rate-determining step

A

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.

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44
Q

How to calculate reaction rate

A

For aA + bB -> cC + dD
Rate = -[][A]/a[]t = -[][B]/b[]t = [][C]/c[]t = [][D]/d[]t
Note the - sign!!

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45
Q

How to calculate rate due to concentration

A 
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
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46
Q

How to convert log

A

logx=y

x=10^y

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47
Q

What do catalysts do?

A

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

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48
Q

What can change a reactions equilibrium constant?

A

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

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49
Q

Equilibrium constant equation

A 
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
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50
Q

Finding equilibrium constant for gasses

A

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

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51
Q

Ratio of products to reactants and Keq

A

Keq»1 products favoured
Keq=1 neither
Keq«1 reactants favoured

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52
Q

Reaction quotient

A

Q = (C)^c(D)^d/(A)^a(B)^b

Brackets indicating concentrations

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53
Q

Relationship between Q and Keq

A

Q>Keq products in excess, decrease in products and increase in reactants as reaction approaches eq
Q=Keq at equilibrium
Q<Keq reactants in excess, increase in products and deceease in reactants as reaction approaches eq

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54
Q

What would be he equilibrium for the reverse reaction?

A

Keq(reverse) = 1/Keq(forward)

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55
Q

How to determine Keq for the reactions
aA+bB->cC+dD = K1
cC+dD->eE+fF = K2

A

aA+bB->eE+fF

Keq = K1(K2)

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56
Q

How does addition or removal of product or reactant affect equilibrium? Addition of non-reactive product?

A

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

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57
Q

Describe pH and equilibrium wrt:

HA H+ + A-

A

Increase H+ lowers ph, shifts equation left and decreases A- (the conjugate base)
Decrease H+ raises ph, shift equation right and increases A-

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58
Q

What happens when volume is increased or decreased in a reaction?

A

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

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59
Q

Explain the effects of temperature on equilibrium

A

Determines whether it’s endothermic or ectothermic and then use Le Chatelier’s principle

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60
Q

Open vs closed vs isolated system

A

Open - energy and matter exchanged
Closed - energy only
Isolated - neither
Note universe = surroundings + system

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61
Q

Name the state and nonstate functions

A

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

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62
Q

What are the 2 principal ways a system can gain or lose energy?

A

By heat transfer or by work

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63
Q

What is the equation for work in thermodynamics?

A

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)

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64
Q

Describe positive and negative heat and work

A
  • 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
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65
Q

What is greater (wrt heat and work) reversible or irreversible process?

A

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)

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66
Q

Internal energy of a system

A 
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
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67
Q

What is entropy? When is it increasing?

A 
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
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68
Q

Equation for enthalpy and it positive/negative values

A

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 +
+ []H means endothermic
- []H means exothermic

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69
Q

Converting between Kalvin and degrees

A

0 C = 273.15 K

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70
Q

Heat of formation

A

Enthalpy required for the formation of one mole of that substance from its elements
0 for all elements in their naturally occurring state

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71
Q

Standard heat of formation

A

Heat of formation in standard conditions (T=298.15K, p=1atm)

72
Q

STP (standard temperature and pressure)

A

T=273.15, p=1atm

73
Q

Heat of reaction

A

[]H^o={Hf products- {Hf reactants
Endothermic is +
Exothermic is -

74
Q

Why is heat of formation of liquid water different from that of gaseous water?

A

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
Q

Gibbs free energy

A 
[]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
Q

3 laws of thermodynamics

A
  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
Q

Know how to draw reaction diagrams

Activation energy, transition state, reaction intermediate, sn1/sn2 reactions, rate determining step, catalyst

A

Pg 320

78
Q

Note: limiting reagent is not necessarily rate determining step!

A

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
Q

Does a catalyst affect []H or []G?

A

No, it just changes the activation energy. It also does not change equilibrium. It lowers the transition state.

80
Q

Finding []G^o at equilibrium

A 
[]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
Q

Heat of solution

A

Enthalpy change that occurs when a solute dissolves in a solvent.
- exothermic solvation (heat released)
+ endo, heat absorbed

82
Q

Henry’s law (solubility of gases in a liquid solvent)

A

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
Q

How to know if solid is soluble in liquid

A

If the solid attains:
>0.05M - soluble
<0.01M - insoluble
Btw .05 and .01 - moderately soluble

84
Q

Molarity

A

M=moles of solute/liters of solution

85
Q

Normality

A

Used almost exclusively for acids and bases

N=(n of potential H+ or OH-)/liters of solution

86
Q

Molality

A

m=(moles of solute)/(kg of solvent)

87
Q

Mole fraction

A

Usually used with gases

Xs=(moles of solute)/(total miles of solution)

88
Q

Molarity of water

A

55M

89
Q

Equilibrium constant equation for solubility (solubility constant)

A

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
Q

Examples of weak electrolytes

A

Weak acids and bases

91
Q

Examples of strong electrolytes

A

Salts of alkali metals and ammonium

Some salts of alkaline earth metals, all nitrates, chlorates, perchlorates, and acetates

92
Q

What substances are solvated (not dissociated)

A

Methanol, ethanol, glucose and fructose

93
Q

Van’t Hoff factor

A

(i) how many species a substance dissociates into
i(NaCl)=2 i(CaCl2)=3
Glucose=1

94
Q

Raoult’s law (how vapor pressure of a pure liquid is lowered by the addition of a nonvolatile solution)

A 
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
Q

What happens to boiling point of a pure liquid when nonvolatile solute is added?

A 
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
Q

What happens to freezing point of a pure liquid when a solute is added?

A 
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
Q

Osmotic pressure

A

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
Q

Air pressure conversions

A 
P=F/A=Pa
1atm=
760mmHg=
760torr=
1.013x10^5Pa
99
Q

One mole of ideal gas at STP occupies what volume?

A

STP= 0C at 1atm
22.4L or 0.0224 cubic meters
(Note that this value is 25C in the standard thermodynamic state)

100
Q

5 premises of kinetic molecular theory

A

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
Q

Boyle’s law

A

PV=constant
or
P1V1=P2V2
For ideal gases under constant temperature and with constant moles

102
Q

Charles’s law

A 
V/T=constant 
or
V1/T1=V2/T2
or
V1T2=V2T1
For ideal gases under constant n and P
103
Q

Gay-Lussac’s law

A 
For ideal gasses under constant n and V
P/T=constant
or
P1/T1=P2/T2
or
P1T2=P2T1
104
Q

Combined gas law

A 
For ideal gases at constant n
PV/T=constant
or
P1V1/T1=P2V2/T2
or
P1V1T2=P2V2T1
105
Q

Avogadros hypothesis

A 
For ideal gases at constant P and T
V/n = constant
or
V1/n1=V2/n2
or
V1n2=V2n1
106
Q

Ideal gas law/gas constant

A

R=PV/nT
or
PV=nRT
R is in J/mol•K

107
Q

What conditions favour ideal gases?

A

Low pressure, high temperature

High P and low T lead to deviations from ideal gas law

108
Q

What are the deviations from ideal gas law wrt P and V?

- excluded volume

A

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
Q

Vanderwalls equation
(P+an^2/V^2)(V-nb)=nRT
What does this signify?

A

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
Q

Law of partial pressure

A 
Ptotal=Pa+Pb
Pt=(na+nb)RT/V
Pa=Xa(Pt)
Xa=na/nt
Xa is the mole fraction of gas a
111
Q

Graham’s law of diffusion (effusion)

A

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
Q

The conversion from solid to gas is called

A

Sublimation

113
Q

The direct conversion from gas to solid is called

A

Deposition

114
Q

What is the formula for heat absorbed by a system when it isn’t undergoing a phase change?

A 
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
Q

Formula for heat of a phase change

A

q=m[]H
q is heat
m is mass
[]H is the enthalpy of the phase change in units of energy per mass

116
Q

How to find C (specific heat capacity for exact amount of substance being considered) from a phase change diagram

A

Slope=1/C
mc=C and 1/slope=mc
P 375

117
Q

How to find heat of fusion for phase change diagram

A

Length of horizontal line = []Hfus•m

m is mass of substance

118
Q

Plot a phase diagram

A

Pg 379

119
Q

What is the diff for a pt diagram of water?

A

Ph 380

120
Q

Draw a pt diagram for a substance with more than one solid form

A

Pg 381

121
Q

What happens if the triple point of a substance occurs at pressure >1atm?

A

It will sublimate under atmospheric conditions and appropriate temp change (example co2 at 216.8 K and 5.11 atm)

122
Q

Describe he critical point

A

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
Q

Equilibrium expression for auto ionization of water

A

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
Q

Formula for pH/pOH of a solution

A 
pH=-log[H+]
pOH=-log[OH-]
and
pH+pOH=14
(-log[H+])+(-log[OH-])=(-logKw)
125
Q

How to transform log values

A

[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
Q

Arrhenius, bronsted-Lowry, and Lewis definition of acids/bases

A
  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
Q

Define strong acid/base

A

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
Q

Examples of strong acids (6)

A 
HCl hydrochloric acid
HBr hydrobromic acid
HI hydroiodic acid
HNO3 nitric acid
HClO4 perchloric acid
H2SO4 sulfuric acid
129
Q

Examples of strong bases (6)

A 
LiOH lithium hydroxide
NaOH sodium hydroxide
KOH potassium hydroxide
RbOH rubidium hydroxide
(NH2-) amide ion
(H-) hydride ion
130
Q

What makes acids more acidic?

A

Electron withdrawing groups, which stabilize their conjugate bases (making the conjugate base less willing to protonate and thus, weaker)

131
Q

Acid/base ionization constant expression

A

HA + H2O H3O+ + A-

Ka = [H3O +][A-]/[HA]

Because water is a pure liquid.
Same thing for base

132
Q

What is the pKa/Ka relationship?

A

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
Q

2 other equations used for any acid/bad pair in dilute aqueous solution

A

KaKb=Kw

pKa + pKb = 14

For any conjugate acid/base pair in a dilute aqueous solution

134
Q

Hendeson Hasselbalch equation

A

Useful in titrations

pH = pKa + log[A-]/[HA]

pOH = pKb + log[BH+]/[B]

135
Q

What happens when you mix a strong acid and a strong base?

A

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
Q

What is the equivalence point?

A

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
Q

What is the equivalence point for strong/weak a/b?

A

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
Q

What is the half-equivalence point?

A

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
Q

What is log 1?

A

0

140
Q

Where is the buffering region an what is always found within it?

A

It is the first part of the weak/strong a/b titration curve and always includes the half-equivalence point.

141
Q

What is a polyprotic acid and what does its titration curve look like? What are indicators?

A

Pg 404/405

142
Q

A more concentrated buffer is stronger or weaker?

A

Stronger

143
Q

What is a buffer composed of?

A

A weak acid/base and it’s corresponding conjugate

144
Q

Describe the pH and pKa of a buffer

A

For a buffer, pH=pKa

145
Q

How do you write redox couples?

A

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
Q

Describe the anode and cathode

  • electrode and inert electrode
  • direction of current
A

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
Q

How to calculate electromotive force and cell emf

-what can alter emf?

A

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
Q

How to draw cell diagrams

A

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
Q

Explain a galvanic cell

A

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
Q

Describe electrolytic cell

A

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
Q

Explain what a concentration cell is

A

Page 425

152
Q

Equation relating []G to E^o

A

[]G^o=-nFE^o
n is moles of e-
F is Faradays constant
[]G=-nFE for nonstandard conditions

153
Q

Nernst equation

A 
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
Q

What direction is the arrow pointing in a dipole moment?

A

The head points towards the negative charge

155
Q

What is a formal charge?

A

The overall charge on a molecule (ie +1). Partial charge is the dipole moment, the uneven distribution of electrons.

156
Q

Energy and bond formation

A

Breaking bonds require energy, forming bonds release energy

157
Q

What is (x)(2x)^2?

A

4x^3

158
Q

Explain how to represent where atomic mass, atomic weight, charge, and atomic number are located

A

P 242

159
Q

What is a period/group?

A

Periods are horizontal and groups vertical
Periods = same principle quantum number (shell number)
Groups = same valence electrons and similar physical properties

160
Q

Describe Pauli exclusion principle, hunds rule, and the aufbau principle

A 
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
Q

What are representative (main block elements) vs transition metals? What are actinides/transuranium elements?

A

251

162
Q

What are the 5 main types of reactions?

A

P 275

163
Q

Describe how to make hybridized orbitals

A

Chapter 33 (731)

164
Q

Describe the types of isomerism/properties

A

Chapter 34 (731)

165
Q

Describe energy wrt breaking and forming bonds

A

Breaking bonds requires an input of energy and forming bonds an output

166
Q

What is microscopic reversibility?

A

The principle that all elementary steps are reversible in exactly the reverse manner (even though the overall reaction is not exactly reversible)

167
Q

Solubility general info

A

P 329

168
Q

Look at solubility rules

A

P 332

169
Q

What is solubility affected by?

A

Temperature, common ions, pH

170
Q

What are colligative properties?

A

P 338

171
Q

Pressure depends on what?

A

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
Q

Under which conditions are the solid/liquid/gas phase favoured? How are each of these phases characterized?

A

P 369

173
Q

Heat of fusion, vaporization and sublimation (endo or exothermic)

A

P 371

174
Q

Relate heat and phase changes to kinetic energy and temperature

A

P 372

175
Q

What is a calorie? What is the specific heat of liquid water?

A

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
Q

What is a ligand?

A

Lewis bases that form bonds with transition metals

177
Q

How to use reduction potential chart

A

416

MCAT (6 decks)

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