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MCAT > G Chem > Flashcards

G Chem Flashcards

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MCAT 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
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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

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3
Q

Charge of e-

A

-1.6022 x 10^-19 C

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4
Q

Mass of proton

A

1.673 x 10^-27 kg

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5
Q

Mass e-

A

9.11 x 10^-31 kg

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6
Q

Relationship between c, λ, and v

A

c = λv

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7
Q

Principle quantum number

A

n

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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
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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)

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10
Q

Spin quantum number

A

ms
+1/2 spin up
-1/2 spin down

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11
Q

Common reducing agents (4)

A

LiAlH4
NaBH4
Metals
Fe2+

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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)
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13
Q

Lewis acid

A

e- pair acceptor

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14
Q

Lewis base

A

e- pair donor

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15
Q

Bronsted Lowry acid

A

H+ donor

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16
Q

Bronsted Lowry base

A

H+ acceptor

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17
Q

Electron and molecular geometry
Number lone pairs: 1
Number bonding pairs: 2

A

Trigonal planar (120), bent (sp2)

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18
Q

Electron and molecular geometry
Number lone pairs: 1
Number bonding pairs: 3

A

Tetrahedral (109.5), trigonal pyramidal (sp3)

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19
Q

Electron and molecular geometry
Number lone pairs: 2
Number bonding pairs: 2

A

Tetrahedral, bent (sp3)

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20
Q

Steric number (# hybridization orbitals in a molecule)

A

sigma bonds + # lone pairs

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21
Q

Electron and molecular geometry
Number lone pairs: 0
Number bonding pairs: 2

A

Linear (180)

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22
Q

Structural isomers

A

Same molecular formula, different compounds

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23
Q

Stereoisomers

A

Differ in 3-D arrangement (max = 2^n) n= # chiral centers

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24
Q

Enantiomers

A

Stereoisomers not superimposable mirror images

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25
Q

Diastereomers

A

Stereoisomers that are non-superimposable non-mirror images (differ at SOME chiral centers, not all)

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26
Q

Z configuration

A

Two top priority groups on same side of double bond

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27
Q

E configuration

A

Two top priority groups on opposite side of double bond

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28
Q

Sublimation

A

Solid to gas

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29
Q

Condensation

A

Gas to liquid

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30
Q

Freezing

A

Liquid to solid

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31
Q

Deposition

A

Gas to solid

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32
Q

Vaporization

A

Liquid to gas

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33
Q

Melting

A

Solid to liquid

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34
Q

Molality

A

(m) = 1 mole / 1000 g

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35
Q

Normality

A

(N) = one equivalent / 1 Liter

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36
Q

log base a of a

A

1

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37
Q

log base a (M^k)

A

KlogaM

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38
Q

log base a (MN)

A

loga(M) + loga(N)

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39
Q

log base a(M/N)

A

logaM - logaN

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40
Q

10^(log10(M))

A

M

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41
Q

First law of thermodynamics (equation)

A 
ΔE = Q - W
Q= heat
W = work
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42
Q

Q greater than 0, heat is….?

A

Absorbed

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43
Q

Q<0, heat is…?

A

Released

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44
Q

Work done on the system, W=…?

A

W less than 0

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45
Q

Work done by the system, W=…?

A

W is greater than 0

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46
Q

Calorimetry
Q=…?
Q=…?

A

Q=mcΔT

Q=mL

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47
Q

Change in Gibbs Free Energy (equation)

A

ΔG = ΔH - TΔS

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48
Q

Rate equation

A

Rate = k[A]^m[B]^n
A and B are reactants

Rate = Δ[concentration]/Δt

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49
Q

Relationship between Keq and ΔG (equation)

A

ΔG(naught) = -RTlnKeq

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50
Q

Hz

A

(1/s)

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51
Q

Wave period

A

T = 1/f (s)

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52
Q

Define: ionization energy

A

How much energy is required to remove an e- from a neutral atom in a gaseous phase (kJ or eV)

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53
Q

Relationship between E photon and BE (binding energy) and KE of photon

A

E = BE + KE

Therefore, hv = BE + KE

54
Q

Higher binding energies correspond with e- in ______ energy subshells

A

Lower, e- are closer to nucleus, require more energy to remove

55
Q

For gases: as temps increase, solubility…?

A

Decreases (gas molecules move too much)

56
Q

Gases: increasing pressure ______ solubility

A

Increases

57
Q

List forces in order of relative strength from strongest to weakest (polar covalent, non polar covalent, metallic, ionic bonds)

A

1 (strongest) metallic
2 ionic
3 polar covalent
4 (weakest) nonpolar covalent

58
Q

When adding an enzyme, does the overall energy released (ΔH) change?

A

NO (kinetics of rxn are not directly related to thermodynamics of rxn)

59
Q

What is an acid catalyst?

A

H+ ion

60
Q

What is a base catalyst?

A

OH- ion

61
Q

Define: heterogeneous catalyst

A

In a different phase than the reactants

62
Q

Define: surface catalyst

A

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
Q

Keq > 1

A

Reaction goes in forward direction (favors products)

64
Q

Le Chatelier’s principle, increase pressure favors which side?

A

Side w/ fewer molecules

65
Q

Define: neutralization reaction

A

Double replacement reaction between an acid and a base

Acid + base = H2O + salt

66
Q

Define: combustion reaction

A

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
Q

Define: disproportionation reaction

A

When a single reactant is both oxidized and reduced

ie: 3ClO- -> ClO3- + 2Cl-

68
Q

Reduction half reaction

A

X+ + e- = X

69
Q

Oxidation half reaction

A

X -> X+ + e-

70
Q

What does Keq describe? What is excluded?

A 
Molar concentrations (mol/L) at equilibrium for a specific temp
Pure solids and pure liquids are excluded
71
Q

Kp (equation) (2)

A

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
Q

What is the oxidation state of oxygen when bound to another oxygen (ie in H2O2)?

A

OS = 1

73
Q

What is the oxidation state of oxygen when bound to F?

A

+1 (F is more electronegative than O)

74
Q

Pressure (equation)

A 
P = F/A 
F = ma
75
Q

When temp increases, what happens to pressure?

A

P increases

76
Q

When increasing # particles, P does what?

A

Increases

77
Q

When volume decreases, P does what?

A

Increases

78
Q

What are the three conditions for a gas to be ideal?

A
  1. No IMF
  2. No volume for individual particles
  3. Collisions between particles are perfectly elastic (KE is conserved)
79
Q

Boyle’s law

A

PiVi = PfVf (at constant T)

80
Q

Charles’s law

A

(Vi/Ti) = (Vf/Tf) (at constant P)

81
Q

How much volume (L) does one mole of gas take up?

A

22.4L

82
Q

Avogadro’s law

A

(Vi/ni) = (Vf/nf) (at constant T & P)

83
Q

Due to Van de Walls forces between gas particles, what is the relationship of Preal to Pideal?

A

Preal < Pideal

84
Q

Due to Van de Walls forces between gas particles, what is the relationship of Vreal to Videal?

A

Vreal > Videal

85
Q

What are the 3 R values for ideal gas law?

A 
R = 8.3145 (m^3*Pa)/(mol*K)
R = 8.3145 (L*kPa)/(mol*K)
R = 0.082 (L*atm)/(mol*K)
86
Q

How do you calculate partial pressures (given total mass of total gas and percent compositions)

A

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
Q

Define: vapor pressure

A

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
Q

Define Specific heat and give the equation

A

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
Q

Heat of fusion/vaporization (equation)

A 
Q = mL
Q = amount of heat released or absorbed (J)
L = latent heat (how much heat energy needed to change phase)
m = mass
90
Q

At what T does the ideal gas law start to break down?

A

Low T

91
Q

If an inert gas is added, what happens to the Ptotal and partial pressures of other gases (Le Chatelier’s principle)?

A

Ptotal increases, but partial pressures of gas stay the same

92
Q

Q > Keq, favors …?

A

Reactants

93
Q

Q < Keq favors…?

A

Favors products

94
Q

CO(g) + 2H2(g) -> CH3OH (g)

If Q < Kc, reaction favors _____ , which favors side with (fewer, more) molecules, causing Ptotal to (increase/decrease)?

A

Favors products, favoring side with fewer molecules, Ptotal decreases

95
Q

Define: hydration

A

Water molecules break ionic bonds to dissolve

96
Q

Calculating solubility (g/L) requires what type of table?

A

ICE

97
Q

Solubility rules mnemonics (2)

A

All soluble: CASH-N-GIA
Chlorates, acetates, sulfates, halides (not F), nitrates, Group 1 metals, ammonium

All insoluble: HAPpy
Hg2 2+, Ag+, Pb2+)

98
Q

Why is water less dense when frozen than liquid?

A

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
Q

Why does it take a lot of heat to raise the T of water?

A

Because some of the heat must be used to break H bonds between the molecules (water has a high specific heat capacity)

100
Q

Define: evaporative cooling

A

As water molecules evaporate, the surface they evaporate from gets cooler because the molecules with the highest kinetic energy are lost to evaporation

101
Q

Define: cohesion

A

Attraction of molecules for other molecules of the same kind (responsible for surface tension)

102
Q

Define: adhesion

A

Attraction of molecules of one kind for molecules of a different kind (responsible for capillary action)

103
Q

Define: 0th Law of Thermodynamics

A

No heat is transferred between two objects in thermal equilibrium (they are the same T)

104
Q

Define: work

A

The energy required to move something against a force

105
Q

What is the relationship between W, P and V for expansion or compression of gases (work done by gases for expansion/compression)?

A

W = -PΔV (J)

**P is the external P

106
Q

When a system does work on its surroundings, the systems internal energy _____?

A

Decreases (negative work)

107
Q

Relate W, force, and displacement (equation)

A

W = F x d

108
Q

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?

A

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
Q

Define: endothermic

A

Heat absorbed or released from a system under constant P

110
Q

Breaking bonds between atoms requires (adding/removing) energy?

A

Adding

111
Q

Forming new bonds (releases/takes) energy

A

Releases

112
Q

Define: exothermic

A

Reactions where the products have lower PE than reactants

113
Q

Define: 2nd Law of Thermodynamics

A

The entropy of the universe is always increasing for a spontaneous process

114
Q

ΔG = ΔH + TΔS is true when what are held constant?

A

Temperature and pressure

115
Q

When ΔG is negative at constant T and P, the reaction is _____ and considered (exergonic/endergonic)?

A

Negative, exergonic

116
Q

When ΔH < 0 and ΔS > 0, ΔG is…?

A

ΔG < 0, spontaneous at all T

117
Q

When ΔH < 0 and ΔS < 0, ΔG is…?

A

Spontaneous at low T (when TΔS is small)

118
Q

When ΔH > 0 and ΔS > 0, ΔG is…?

A

Spontaneous at high T (when TΔS is large)

119
Q

When ΔH > 0 and ΔS < 0, ΔG is…?

A

Non-spontaneous at all T (ΔG > 0)

120
Q

Define: 1st Law of Thermodynamics

A

Energy can neither be created nor destroyed

121
Q

Define: open system

A

A system that freely exchanges both energy and matter with its surroundings

122
Q

Define: closed system

A

A system that exchanges only energy with its surroundings, not matter

123
Q

If ΔS > 0, the reaction is (spontaneous/non-spontaneous)?

A

Spontaneous

124
Q

If ΔS < 0, the reaction is (spontaneous/non-spontaneous)?

A

Non-spontaneous

125
Q

Water going from a solid to liquid is an (endothermic/exothermic) process?

A

Endothermic, heat absorbed, ΔH positive (chemical bonds broken)

126
Q

Water going from a liquid to gas is an (endothermic/exothermic) process?

A

Endothermic, heat absorbed, ΔH positive (chemical bonds broke )

127
Q

Water going from a gas to liquid is an (endothermic/exothermic) process?

A

Exothermic, heat released, ΔH negative (chemical bonds formed)

128
Q

Water going from a liquid to solid is an (endothermic/exothermic) process?

A

Exothermic, heat released, ΔH negative (chemical bonds formed)

129
Q

Equation for ΔH (relating to bonds formed/broken)

A

ΔH = ΔH (bonds broken in reactants) - ΔH (bonds formed in products)

130
Q

Enthalpy equation ( H, E, P, V)

A

H = E + PV

131
Q

Define: critical point

A

Point on phase diagram where phase can coexist (very high T and P)

132
Q

Define: triple point

A

Point on phase diagram where all three phases coexist (intersection point), equilibrium between phases

MCAT (17 decks)

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