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Flashcardlet > MCAT Chemistry 2 > Flashcards

MCAT Chemistry 2 Flashcards

1
Q

ex of a closed syst

A

steam radiator

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2
Q

____________ is a meas of the averag ke of the particles in a syst

A

temperature

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3
Q

G is ?

A

free energy

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4
Q

bond e or bond dissoc e is?

A

averag of the energy req to break a part type of bond in one mole of gaseous molec

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5
Q

give an ex of an isolated system?

A

insulated bomb reactor

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6
Q

T or Fit is not poss to meas H directly

A

T

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7
Q

system

A

partic part of universe being studied

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8
Q

open system is?

A

can exch matter and e with surroundings.

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9
Q

ΔG <0 means a proc can?

A

occur spont

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10
Q

qrev/T =

A

ΔS

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11
Q

conversion of C (diamond) to C (graphite) is __________ but its rate is slow

A

spont

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12
Q

bond ____________ is always endothermic

A

breakage

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13
Q

stand heat of format of a compound is?

A

enthalpy ch that would occur if one mole of a comp were formed direct from its elem in their standard states

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14
Q

T or Fheat and temp are different

A

t

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15
Q

boiling has a ____________ in entropy

A

increase

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16
Q

state functions

A

prop that dep only on initial and final states of the system

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17
Q

E or U is?

A

internal energy

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18
Q

the react quotient Q is=

A

Q=[C]^c[D]^d/[A]^a[B]^b

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19
Q

units of heat?

A

joules

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20
Q

ΔH ΔS Outcome ?- +

A

spontaneous at all temperatures

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21
Q

heat changes at const press, is

A

enthalpy

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22
Q

isothermal process is?

A

temperat of systm is const

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23
Q

a system undergoes a process when?

A

when one or more of its prop change

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24
Q

ΔH ΔS + +

A

spontaneous only at high temperatures

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25
Q

-RTlnKeq =

A

ΔG

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26
Q

system can be?

A

isolatedclosedopen

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27
Q

ΔH ΔS - -

A

spontaneous only at low temperatures

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28
Q

t or fentropy is a state function

A

T

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29
Q

t or frate of a react dep on ΔG

A

f

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30
Q

for hesses law, if you multiply the prod and react by 3 ( or a cert number) then what must you do to ΔH?

A

also multiply it by 3

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31
Q

endothermic react?

A

absorb e

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32
Q

q=mc chTc is?

A

specific heat

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33
Q

q=mc chTq is?

A

heat

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34
Q

T or FΔHreact=Hprod-Hreact

A

T

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35
Q

enviroment

A

everthing outside system

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36
Q

freezing has a ___________ in entropy

A

decrease

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37
Q

a spontaneous react may or may not?

A

proceed to completion

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38
Q

bomb calorimeter, the overall system is?

A

adiabatic

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39
Q

standard free energy of format of a comp is?

A

free energy ch that ocurs when 1 mol of a comp in its stand form

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40
Q

ΔHf of an element in its standard state is?

A

zero

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41
Q

second law of thermodyn says that?

A

all spont proceed so that entropy of the systm plus its surr increases

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42
Q

temp is?

A

a measure of average ke of particles in a systm

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43
Q

name 7 state funct?

A

tempvolenthalpyentropyfree energyinternal energy E and U

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44
Q

H is?

A

enthalpy

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45
Q

what condit are heat changed measured under?

A

const vol calorimetryconst press calorimetry

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46
Q

enthalpy of a process does not depend on the?

A

path

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47
Q

rate of a reac depends on?

A

activat e

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48
Q

standard condit in thermody must not be confused with?

A

standard temperature and pressure in gas laws STP

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49
Q

closed system is?

A

can exch e but not matter

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50
Q

ΔH ΔS + -

A

nonspontaneous at all temperatures

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51
Q

isobaric process is?

A

press of systm is const

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52
Q

q react + q water + q steel=0 in a ?

A

bomb calorimeter

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53
Q

Goose hunters take shotguns reminds you of?

A

ΔG=ΔH-TΔS

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54
Q

discuss when water boils in terms of ch in G

A

TchS> chH only when T is above 373 K

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55
Q

reverse reaction has the same __________ as that of the forward react, but its ________ is oppos

A

enthalpy,but its sign is oppos

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56
Q

for simple diatomic molec such as H2, bond dissoc e can be easily measured?

A

spectroscopically using react E=hv E=bond dissoc e and v is freq of light absorbed

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57
Q

ΔS= ?/T

A

qrev/T

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58
Q

when a react commences ΔG=?

A

ΔG= ΔG° + RTlnQ

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59
Q

calorimetry measures?

A

heat changes

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60
Q

spectroscopy can be used to measure the bond dissoc e of?

A

simple diatomic molecules

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61
Q

standard free energy ΔG is ?

A

ΔG of a proc occur at 25 C and 1 atm . concentr of sol are 1 M

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62
Q

ΔS=Sfin- Sinit

A

Sfin- Sinit

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63
Q

standard heat of react is?

A

ΔHreact= (sum of ΔHf of prod)-(sum of ΔHf of react)

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64
Q

exothermic react?

A

release e

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65
Q

a system undergoes a ____________ when one or more of its properties change

A

process

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66
Q

standard state of a subs is the form a subst takes?

A

at 25 C and 1 ATM

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67
Q

Hess law says?

A

that enthalpies of react are additive

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68
Q

heat absorbed by a system is?

A

pos

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69
Q

Thermodynamics

A
  • Is the study of energy and it’s relationship to macroscopic properties of chemical systems. It’s functions are based on probabilities and are only valid for systems that are composed of a large number of molecules.* I.e The rules of thermodynamics gov
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70
Q

System vs. Surroundings

A

A system is a part of a universe that is the macroscopic body under study, and the surroundings is everything else.Systems are based on mass and energy exchange with the surroundings. There are three systems:1) open: exchange both mass and energy with surroundings.2) closed: Exchange energy but not mass3) isolated: Do not exchange energy or mass.

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71
Q

State Function

A

A state is the physical condition of a system described by a specific set of thermodynamic property values. Such properties that describe the state of a system are called State Functions.Two types of properties used to describe the state of a system: 1) Extensive: properties are propotional to the size of the system (eg. V + n)2) Intensive: are independent of the size of the system. ( eg. P + T )Seven State Functions: U (internal Energy), T, P, V, H (Enthalpy), S (entrophy) and G (Gibbs energy).

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72
Q

Heat

A

(q)Heat has three forms:1) Conduction2) Convection3) RadiationAlways the movement of energy from hot to cold.

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73
Q

Work

A

work = P∆V

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74
Q

Convection

A
  • Is the thermal energy transfer via fluid movements.- differences in density or pressure drive warm fluid to cold fluid*Oceans and air currents are commone examples of convection.
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75
Q

Radiation

A
  • is thermal energy transfer via electromagnetic waves. (eg. heated metal red, orange, white, blue-white)- All objects above 0K radiate heat.P= σ∈AT⁴
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76
Q

The First Law of Thermodynamics

A

∆E = q + w*Warning: work ON the system is positive for “convention” a passage on MCAT may define work done BY the system as a postive in which case you use this formula.

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77
Q

The Second Law of Themodynamics

A

*See heat engines

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78
Q

Internal Energy Types

A

MCAT may refer to internal energy as “heat energy”, “thermal energy,” or even “heat.”“Heat energy and thermal energy” are really the vibrational, rotational and translational parts of interanl energy. Called this because they effect temperature.”Heat” is a transfer of energy. Don’t mistake.

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79
Q

Temperature

A

0 K = 273°C

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80
Q

Enthalpy

A

(H) = is a man- made property that accounts for this extra capacity to do PV work. Unlike functions such as pressure, volume , and temperature, enthalpy is not a measure of some intuitive property. Defined more so as an equation then a property:H≡ U + PV

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81
Q

Standard State

A
  • Don’t confuse with STP. STP is at 0°C whereas standard state is at 25°C and is arbritrarily assigned an enthalpy value of 0 J/mol
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82
Q

Reference Form

A

Define

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83
Q

Standard Enthalpy of Formation

A

Define

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84
Q

Exothermic vs. Endothermic

A

Define

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85
Q

Activation Energy

A

Define

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86
Q

Transition state

A

Define

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87
Q

Intermediates

A

Define

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88
Q

Catalyst

A

Define

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89
Q

Irreversible vs. reversible reactions.

A

Define

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90
Q

Third law of Thermodynamics

A

Define

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91
Q

Spontaneity

A

Define

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92
Q

Gibbs Free Energy

A

Define

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93
Q

Thermodynamics

A
  • Is the study of energy and it’s relationship to macroscopic properties of chemical systems. It’s functions are based on probabilities and are only valid for systems that are composed of a large number of molecules.* I.e The rules of thermodynamics gov
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94
Q

What are two ways to transfer energy between systems

A

NAME?

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95
Q

Conduction

A
  • is thermal energy transfer via molecular collisions.- Requires direct physical contact- an objects ability to conduct heat is called it’s thermal conductivity (k) **On page 46 go over again.Q / t = kA [ T(h) - T(c) ) / L ]
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96
Q

Stephan- Boltzman Law

A

P= σ∈AT⁴Where:A is the surface area of the objectT is the temperatureσ is the Stefan- Boltzman constant (5.67 x 10⁻⁸ )∈ is the emissivity of the object’s surface.

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97
Q

Zeroth Law of Thermodynamics

A

Discovered after 1st, 2nd, and 3rd Law of Thermodynamics. All three rely on this Law, as it is based off of temperature.States: Two systems in thermal equilibrium with a third system are in thermal equilibrium with each other. The law declares that the two bodies in thermal equilibrium share a termodynamic property, which must be a state function. This is temperature.

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98
Q

Standard Temperature and Pressure

A

(STP)

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99
Q

Mean free path

A

Define

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100
Q

Ideal gas

A

Define

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101
Q

Kinetic Molecular theory

A

Define

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102
Q

Ideal gas law

A

PV = nRTK.E.(avg) = 3/2 (RT)

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103
Q

Standard Molar volume

A

22.4L —> At STP one mole of any gas (behaving ideally will occupy this amount.

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104
Q

Partial Pressure

A

Pa= Xa Ptotal

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105
Q

Dalton’s Law

A

P (total) = P1 + P2 + P3 …

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106
Q

Graham’s Law

A

v1/v2 = √M2 / √M1

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107
Q

Effusion

A

effusion rate 1/ effusion rate 2 = √M2 / √M1

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108
Q

Diffusion

A

Define

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109
Q

Chemical Kinetics

A

Define

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110
Q

Collision Model

A

Define

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111
Q

Activation Energy

A

Define

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112
Q

The effect of temperature on the rate of reaction

A

The rate of a reaction increases with temperature

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113
Q

Intermediates

A

Define

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114
Q

Rate Law

A

Define

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115
Q

Order of each respective reactant

A

Define

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116
Q

Overall Order of a reaction

A

Define

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117
Q

Rate Determining Step

A

Define

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118
Q

Catalyst

A

Define

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119
Q

Heterogeneous catalyst

A

Define

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120
Q

Homogeneous Catalyst

A

Define

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121
Q

Chemical Equilibrium

A

Define

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122
Q

Equilibrium Constant

A

K

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123
Q

The law of Mass Action

A

Define

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124
Q

Reaction Quotient

A

Define

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125
Q

Le Chatelier’s Principle

A

Define

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126
Q

Metals

A
  • large atoms that tend to lose electrons and form positive ions (cations) to obtain a noble gas configuration.- Can bond with other metal atoms (metallic bonding)- can also form ion bonds with negative ions (anions) *Usually non-metals.-Metals are
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127
Q

Ductile

A
  • easily stretched.- occurs in metals- if you stretch or hammer e- the metals can adapt (which is why you can stretch metal without breaking them)
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128
Q

Luster

A

NAME?

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129
Q

Non- metals

A
  • form negative ions (anions) to obtain a noble gas configuration.- form covalent bonds with other non-metals- form ionic bonds with positive ions (usually metals)- poor conductors of heat and electricity- non-malleable, brittle and possess low to moder
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130
Q

Metalloids

A

NAME?

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131
Q

Transition metals

A
  • wont be focused on for MCAT, just know that they exist and that this might be a passage based question.- 1/2 vs fully filled orbitals occur here often.
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132
Q

Alkaline Earth Metals

A
  • Harder metallic sollids have 2 valence electrons (can easily form +2 cations)- have a higher melting point than solf metallic solids- Heavier are more reactive than lighter ones.
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133
Q

Halogens

A

F₂: gas at RT and 1 atmCl₂: gas at RT and 1 atmBr₂: liquid at RT and 1 atmI₂: solid at RT and 1 atm* all highly reactive unless in diatomic form.-7 valence e-

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134
Q

Noble gases

A
  • 8 valence e– complete octet so very stable and non-reactive.
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135
Q

How do you determine the number of valence electrons in a main group atom?

A
  • look at the group numberif group 1= 1 e- on valence shell (+1)group 2= 2 e- on valence shell (+2)recall: First shell closest to nucleus = ground stateLast shell closest to outside world = Valence e- shell. *Valence e- are involved in chemical bonding
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136
Q

Atoms

A

Define

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137
Q

Protons

A

Define

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138
Q

Electrons

A

Define

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139
Q

Nucleus

A

Define

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140
Q

Neutrons

A

Define

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141
Q

Elements

A

Define

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142
Q

Mass Number

A

(A)

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143
Q

Atomic Number

A

(Z)

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144
Q

Isotopes

A

Define

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145
Q

Atomic Weight

A

Define

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146
Q

Molar Mass

A

(MM or M)

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147
Q

Atomic Mass Units

A

(amu)

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148
Q

Mole

A

Define

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149
Q

Avogadro’s Number

A

Define

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150
Q

Periodic Table

A

Define

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151
Q

Period

A

Define

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152
Q

Groups or Families

A

Define

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153
Q

Metals

A

Define

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154
Q

Nonmetals

A

Define

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155
Q

Metalloids

A

Define

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156
Q

Transition Metals

A

Define

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157
Q

Inert Gases

A

Define

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158
Q

Shields

A

Define

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159
Q

Effective Nuclear Charge

A

(Z eff)

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160
Q

Periodic trends

A

Define

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161
Q

Atomic Radius

A

Define

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162
Q

Ionization Energy

A

Define

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163
Q

Second Ionization Energy

A

Define

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164
Q

Electronegativity

A

Define

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165
Q

Electron affinity

A

Define

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166
Q

Metallic Character

A

Define

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167
Q

Bonds

A

Define

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168
Q

Covalent Bonds

A

Define

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169
Q

Bond Length

A

Define

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170
Q

Bond Energy

A

*Or bond dissociation energy

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171
Q

Compound

A

Define

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172
Q

Empirical Formula

A

Define

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173
Q

Molecules

A

Define

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174
Q

Molecular formula

A

Define

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175
Q

Ionic Compounds

A

Define

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176
Q

Acids

A

Define

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177
Q

Binary Molecular compounds

A

Define

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178
Q

Physical reaction

A

Define

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179
Q

Chemical reaction

A

Define

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180
Q

Runs to completion

A

Define

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181
Q

Fundamental Reaction types

A

1) Combination: A+B —> C2) Decomposition: C —> A + B3) Single Displacement: A + BC –> B + AC4) Double Displacement: AB +CD —> AD + CB

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182
Q

Principle Quantum Number

A

(n)*Note: quantum numbers are the equivalent of a mailing address for an electron.- Description: Energy level and average distance from nucleus- Possible values: n= 1, 2, 3 etc.

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183
Q

Shell

A

Define

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184
Q

Valence electrons

A

Define

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185
Q

Azimuthal Quantum number

A

(l)*Also known as angular momentum quantum numberDescription: Orbital shape (s, p, d, f)Possible values: l=0, to …n-1l=0 is the s-subshelll=1 is the p-subshelll=2 is the d-subshelll= 3 is the f-subshell

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186
Q

Subshell

A

eg. s, p, d and fs= sphericalp= dumbbell shapedd= clover leaf

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187
Q

Magnetic Quantum number

A

m- Description: Orbital orientation (px, py, pz)- Possible values: m= -l to +lfor l=0 (spherical s orbital) there is only one orientation.for l=1 (dumbbell p orbital_ there are three orientations along each of the x, y, and z axes.

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188
Q

Atomic Orbital

A

Space around a nucleus in which e- have the ability to exist.- max 2 e- in each orbital.

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189
Q

Electron Spin quantum Number

A

m (s)- Description: Describes spin of electron- Possible values: m(s) = +½ or -½Either clockwise or counter clockwise.

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190
Q

Pauli exclusion Principle

A

NAME?

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191
Q

Heisenberg Uncertainty Principle

A
  • tells us that the exact position and momentum of such an electron cannot be measure simultaneously. *Ie the more you know about the momentum of an electron, the less you know about it’s exact position.
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192
Q

Aufbau Prinicple

A

*Write out on a piece of paper before writing mcat.1s2s 2p3s 3p 3d4s 4p 4d 4f5s 5p 5d 5f6s 6p 6d7s 7p

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193
Q

Electron Configuration

A

Know 5 points on page 29 of lecture manual.

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194
Q

Ground State

A

Define

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195
Q

Hund’s Rule

A

Define

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196
Q

Planck’s Quantum theory

A

Define

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197
Q

Photoelectric effect

A
  • an important experiment demonstrating the particle nature of light. An electron can be ejected from the surface of a metal by shining light of a certain frequency. Incoming light (photon) must have enough energy to eject the electron.E(photon) = h(Plank
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198
Q

Wavelength vs. Frequency vs. Velocity

A

Wavelength = distance btwn two consecutive crests of troughs. measured in meters.Frequency= the number of crests that pass through a given point per unit of time (Hertz)Velocity= the distance a wave travels through space/ unit time (m/s)Frequency= (velocity) / (wavelength)

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199
Q

Wave particle duality

A

all matter, including light can be thought of as both a wave and as a particle.

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200
Q

Electromagnetic Radiation

A

(EMR) is simple energy. as energy travels forward electric and magnetic fields are created perpendicular to each other and to the direction of travel.

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201
Q

Order of Colours

A

RedOrangeYellowGreenBlueViolet

202
Q

Absorption Line spectrum vs. Emmision Line spectrum

A

ALS= Dark lines on a light backgroundELS= Light lines on a dark back ground.

203
Q

How do you determine the max number of e- an atom can have?

A

ERROR!:parse

204
Q

Mass Percent (of a solute)

A

Mass(solute)/Mass(solution)

205
Q

Archimedes Principle

A

Any object wholly or partially immersed in a fluid is buoyed up by a force equal to the weight of the fluid displaced by the object

206
Q

Buoyant Force

A

Buoyant Force=Density(medium)Volume(object)g

207
Q

Apparent immersed weight

A

Weight-Weight of Displaced Fluid

208
Q

Molar Volume of an Ideal Gas at STP

A

22.4 Liters

209
Q

Empirical Formula

A

Simplest whole number ratio of atoms in a molecule

210
Q

Mass Percent (of an element)

A

Mass percent=[(mass atoms)/(mass compound)]*100%

211
Q

Molecular Formula

A

Requires multiplying the empirical formula by the whole number ratio of the molecular mass:empirical mass; requires knowing the molecular mass of the compound

212
Q

Hygroscopic salt

A

Absorbs water; examples include calcium chloride and magnesium sulfate

213
Q

KOH

A

Potassium Hydroxide; undergoes a combustion reaction with carbon dioxide to form solid potassium bicarbonate

214
Q

Molarity

A

Moles of Solute/Liters of Solution

215
Q

Molality

A

Moles of Solute/Kilograms of Solvent; does not change with temperature so it’s used to calculate the boiling-point elevation and freezing-point depression of solutions containing non-volatile impurities

216
Q

Dilution:Fold vs. Parts

A

Fold (based on the total volume); parts (based on the volume added)

217
Q

Beer’s Law

A

The specific wavelength of light absorbed and the intensity of the absorbance varies with the solute and its concentration respectively

218
Q

Absorbance

A

Absorbance=εcl, where ε refers to the absorption constant (a constant for the solution at λ max), c refers to concentration, and l refers to path length (or the width of the cuvette).

219
Q

Limiting Reagents

A

The reactant with the lowest ratio of actual moles to needed moles

220
Q

Precipitation Reactions (or metathesis or double-displacement reactions)

A

Involves two aqueous salts that react to form spectator ions and a solid salt precipitate

221
Q

Acid-Base Reactions

A

A reaction between an acid (proton donor) and a base (proton acceptor) that results in the formation of a neutral salt and water

222
Q

Composition Reaction

A

The number of reactants exceeds the number of products; entropy decreases and bonds are formed

223
Q

Decomposition Reactions

A

Reactant(s) decompose to form multiple products; entropy increases and bonds are broken

224
Q

Oxidation-ReductionReactions

A

Transfer of electrons from one atom to another ; the oxidation state must change in a redox reaction

225
Q

Oxidation

A

Loss of electrons

226
Q

Reduction

A

Gain of electrons

227
Q

Reducing Agent

A

The atom (or compound) that is losing electrons

228
Q

Oxidizing Agent

A

The atom (or compound) that is gaining electrons

229
Q

Combustion Reactions

A

A special type of oxidation-reduction reactions where the oxidizing agent is the oxidizing agent and the products are oxides (water and CO2)

230
Q

Molarity

A

(M) * units: mols/LM= (moles of solute) / (Volume of solution)

231
Q

Mole Fraction

A

(m) *has no unitsm= (moles of solute) / (kilograms of solvent)

232
Q

Solving for X

A

X= (moles of solute) / (total moles of all solutes and solvent)

233
Q

Parts per million

A

PPM= (mass of solute) / (total mass of solution) x10⁶*Note: PPM is NOT the number of solute molecules per million molecules. It is the mass of the solute per mass of solution times on million.

234
Q

Solution Formation

A

Define

235
Q

Vapor Pressure

A

Define

236
Q

Nonvolatile Solute

A

Define

237
Q

Raoult’s Law

A

Define

238
Q

Volatile Solute

A

Define

239
Q

Solubility

A

Define

240
Q

Precipitation

A

Define

241
Q

Saturated

A

Define

242
Q

Solubility product K(sp)

A

Define

243
Q

Spectator Ions

A

Define

244
Q

The common Ion effect

A

Define

245
Q

Solubility Factors

A

Define

246
Q

Solution

A

Define

247
Q

Solute vs. Solvent

A

Define

248
Q

Colloids

A

Define

249
Q

London Dispersion Forces

A

Define

250
Q

Solvation

A

Define

251
Q

Hydration

A

Define

252
Q

electrolyte

A

Define

253
Q

Arrhenius Acid vs. Base

A

Acid = Is anything that produces hydrogen ions in aqueous solution. [ H+ ]Base = is anything that that produces hydroxide ions in an aqueous solution. [ OH- ]*Only aq solutions.

254
Q

Bronsted and Lowry

A

DefinesAcid: as anything that DONATES a PROTON.Base: as anything that ACCEPTS a PROTON.

255
Q

Lewis

A

More general then Bronsted/ Lowry or Arrhenius.Acid: Anything that ACCEPTS a pair of ELECTRONSBase: anything that DONATES a pair of ELECTRONS.*Includes B+L acids and bases and more because it also accepts molecules that have incomplete octets of electrons eg. BF3 or AlCL3. Also includes simple cations (Smaller the cation the higher the charge, the stronger the acid) Excludes alkali and heavy alkaline earth metals.

256
Q

Acid + Base vs. Conjugate Acid + Base

A

Acid + Base= Refers to the Reactants [A]Conjugate Acid + Base= Refers to the Product [HA][HA] is the conjugate ACID of BASE [A-][A-] is the conjugate BASE of ACID [HA]*The stronger the acid/Base the weaker the conjugate base/Acid, HOWEVER, weak acids may have EITHER weak OR strong conjugate bases.

257
Q

Polyprotic

A

Define

258
Q

Diprotic

A

Define

259
Q

Amphoteric

A

Define

260
Q

How does molecular structure affect Acid strength

A

Define

261
Q

Hydrides

A

Define

262
Q

Autoionization of water

A

Define

263
Q

Acid Dissociation Constant

A

K(a)

264
Q

Finiding the PH

A

Define

265
Q

Titration

A

Define

266
Q

Titration Curve

A

Define

267
Q

Equivalency point

A

or Stoichiometric point

268
Q

Half equivalence point

A

Define

269
Q

Buffered

A

Define

270
Q

Henderson- Hasselbalch Equation

A

pH= pK(a) + log ( [A-] / [HA] )

271
Q

Indicator

A

Define

272
Q

Endpoint

A

Define

273
Q

Polyprotic titrations

A

Define

274
Q

Strong Acids

A
  • Hydroiodic Acid (HI)- Hydrobromic Acid (HBr)- Hydrochloric Acid (HCl)- Nitric Acid (HNO₃)- Perchloric Acid (HClO₄)- Chloric Acid (HClO₃)- Sulfuric Acid (H₂SO₄)
275
Q

Strong Bases

A
  • Sodium Hydroxide (NaOH)- Potassium Hydroxide (KOH)- Amide Ion (NH₂-)- Hydride Ion ( H-)- Calcium Hydroxide Ca(OH)₂- Sodium Oxide (Na₂O)- Calcium Oxide (CaO)
276
Q

______ occurs at the cathode in an electrochemical cell. Electrons flow _____ the cathode.

A

reduction, toward

277
Q

_____ occurs at the anode in an electrochemical cell. Electrons flow ____ the anode.

A

oxidation, from

278
Q

The moles of solute over the kg of solvent

A

molality

279
Q

Normality is the molar _____ per unit volume

A

equivalents

280
Q

A ____ function is one in which the result in dependent only on the initial and final measurements and independent of the path needed to get there

A

state

281
Q

_____ describes the the cagelike assortment of solvent particles around solute particles

A

solvation

282
Q

Hess’s Law

A

The enthalpy of reaction is equal to the difference between the ∆Hformationproducts-∆Hformationreactants

283
Q

Isobaric process

A

One that occurs under constant pressure

284
Q

Isochoric process

A

One that occurs under constant volume and no volume-pressure work is done

285
Q

Constant-volume calorimeters (bomb calorimeters) are used to measure

A

the total heat absorbed or given off by a reaction

286
Q

Adiabatic process

A

One that occurs without the transfer of heat

287
Q

Collision theory of chemical kinetics

A

The rate of reaction is proportional to the number of collisions that occur between reaction molecules per second

288
Q

Disproportionation

A

When a species is oxidized and reduced in the same reaction

289
Q

Partial pressure equation

A

Ppartial = Ptotal * X(mole fraction)

290
Q

Kinetic Molecular Theory of Gases

A

Gases have completely elastic collisions with themselves and the walls of their container. Gas particles have negligible volume, negligible attractive forces and exhibit random motion

291
Q

Ideal gases

A

Monoatomic, no attractive forces or volume

292
Q

Henry’s Law:

A

The partial pressure of a gas above a solution is directly proportional to the partial pressure of the gas dissolved in the solution

293
Q

Graham’s Law:

A

Particles diffusion and effusion rates are inversely proportional to the squareroots of their molecular weights

294
Q

Effusion

A

The passage of gas from high to low pressure through a small opening

295
Q

Charles and Gay-Lussac’s Law

A

At constant pressure, the volume of an ideal gas is directly proportional to its temperature

296
Q

Boyle’s Law

A

At constant temperature, the pressure of an ideal gas is inversely proportional to its volume

297
Q

Molecular orbital

A

A regions where bonding or antibonding orbitals overlap, resulting in a low-energy bonding orbital or high-energy unstable antibonding orbital

298
Q

Avagadros Principle

A

Different gases at the same temperature, pressure and volume, contain the same number of particles

299
Q

Electrochemical reaction

A

One which needs or produces electricity

300
Q

Nernst Equation

A

V=Vo-(0.6/n)logQ

301
Q

VSEPR

A

Shows the 3D geometrical shape of a molecule that is based on the electronic interactions between bonding and non-bonding electrons

302
Q

Isoelectronic

A

When two atoms have the same electron configuration

303
Q

Magnetic Quantum number

A

M(L). Range from -L-L. Determines what orbital within the subshell the electron is likely to reside in.

304
Q

Azimuthal quantum number

A

L. Ranges from 0-(n-1). Determines what subshell the electron is likely to be found in.

305
Q

Aufbau Principle

A

Electrons fill an atom in order of increasing energy level

306
Q

Exceptions to the aufbau principle

A

Cu and Cr. Cu takes an electron from a s orbital and places it in a d-orbital, completing its d-orbital. Cr takes an electron from an s orbital and places it in an d orbital, giving it the maximum number of unpaired electrons.

307
Q

Hunds Rule

A

Electrons fill orbitals such that a maximum number of unpaired electrons results

308
Q

Pauli Exclusion Principle

A

No two electrons in an atom can have the same set of 4 quantum numbers

309
Q

Arrhenius Definition

A

Acid dissociates to form H+ in aqueous solution. Base dissociates to form OH- in aqueous solution.

310
Q

Colligative properties

A

Those that depend only on the number of solute particles and not on the chemical makeup of the particles. (Boiling point elevation, freezing point depression, osmotic pressure, partial pressure)

311
Q

Raoults Law

A

The vapor pressure above a solution is directly proportional to the mole fraction in solution. Ppartial = Ptotal*X

312
Q

Energy of emitted electron =

A

hf- electron binding energy

313
Q

malleability and other metallic characteristics ____ down and to the left on the periodic table

A

increase

314
Q

most likely to have multiple oxidation states

A

transition metals

315
Q

ionization energy

A

: E needed to remove an electron from a gaseous state

316
Q

atomic mass increases ____ and ____ on the periodic table

A

right and down

317
Q

which ions are typically the largest

A

negative ones (anions)

318
Q

density _______ with increasing atomic radius

A

decreases

319
Q

diatomic gases will have _____ densities than monoatomic gases

A

larger

320
Q

T/F: solutions containing ions of transition metals are frequently colored

A

T

321
Q

Units of ideal gas law

A

J/K*mol

322
Q

Kinetic energy of n moles of gas

A

(3/2)nRT

323
Q

When an outside force of other gases causes the molecule to move

A

Brownian forces

324
Q

Variations from ideal gas conditions occur at _____

A

low volumes, low temperatures, high pressures

325
Q

The a and b in the van der Waals equation represent…

A

a = attractive forces between particles, b=volume of molecules themselves

326
Q

Ways of determining the reaction rate

A

Measure the inital reaction rate for a variety of reactant concentrations. Graph the concentration of the reactants as a function of time. Find the mechanism of the reaction.

327
Q

N+L rule

A

determines which electron configuration comes next

328
Q

The rate law constant (k) is affected by ____ and ____

A

temperature, catalyst

329
Q

At equilibrium, the rate of forward reaction is ______ to/than the backwards reaction

A

equal

330
Q

If a collision between two reactants does not form product, it can be assumed that….

A

the positioning was not right and or the energy of collision was not sufficient

331
Q

As temperature increases, the range of reaction kinetic energy at which particles collide _____. The rate of collisions peaks at a higher kinetic energy than at _____ temperatures

A

increases, lower

332
Q

Do liquids and solids enter into equilibrium equations?

A

no

333
Q

Rates _____ increase with increasing temperature in endothermic and exothermic reactions

A

always

334
Q

The emissivity value of blackbodies

A

1

335
Q

Is it more efficient to heat quickly or slowly?

A

quickly

336
Q

The area under the graph of a P vs. V graph is equal to _____

A

the work done on the gas

337
Q

Free adiabatic expansion

A

No change in temperature since no work is done and no heat is lost (ideal gas)

338
Q

Intensive properties

A

Not dependent on the amount of substance

339
Q

Extensive properties

A

Dependent on the amount of substance

340
Q

Percent yield

A

actual/theoretical * 100

341
Q

Rapid cooling of polymers will result in an _____ _____ whle slow cooling will result in _____ _____

A

amorphous solid, crystalline solid

342
Q

Diamond is a _____ _____ solid

A

network covalent

343
Q

Heisenburg uncertainty _____ with increasing mass

A

decreases

344
Q

Enthalpy will differ from the energy of reaction when…

A

The change in entropy is high

345
Q

Heat of formation

A

Amount of heat needed to produce 1 mole of a given product

346
Q

The heat of formation of natural elements is…

A

zero

347
Q

Gases have ____ entropy than phases of the same substance

A

higher

348
Q

Example of 2nd law of thermo

A

Heat always flows from high to low T

349
Q

Bond formation is ____thermic and _____ entropy

A

exo, decreases

350
Q

Any machine is less than ___% efficient because of ____ loss

A

50, heat

351
Q

A substance at absolute zero has this entropy value

A

zero

352
Q

Constant ____ and ____ are necessary in classfying spontaneity of reactions

A

temperature and pressure

353
Q

When the velocity of a fluid increases, the pressure _____

A

decreases

354
Q

When temperature decreases in a fluid, velocity _____, and thus pressure _____

A

increases, decreases

355
Q

When density of a fluid decreases, volume ______ and pressure _____

A

increases, decreases

356
Q

The pressure of a fluid against a flat surface is equal in momentum/∆t*A

A

Define

357
Q

Ideal fluid characteristics

A

Volume flow rate is constant, no viscosity, incompressible

358
Q

The flow of an ideal fluid _____ be determined by pressure changes alone

A

cannot

359
Q

Possible phases of a solution

A

gas, liquid, solid

360
Q

In ideal dilute solutions, solute particles ____ interact, and the mole fraction of the solvent approaches ____

A

never, 1

361
Q

Ideal solutions obey _____ law, where solute and solvent molecules are similar and interact with each other similarily

A

Roults

362
Q

Non crystalline homogenous solution

A

Colloid

363
Q

Colloids do not ____ in solution. Cannot be _____ by centrifugation. May be separated by addition of _____, by raising_____ or by this technique.

A

settle, separated, electrolytes, temperature, dialysis

364
Q

Scattering of light results in the appearance of the lights path

A

Tyndall effect

365
Q

Colloid particles are ____ than thos of the solvent but not _____ enough to precipitate out

A

larger, large

366
Q

Ionic compound naming:

A

-ate, -ite, hypo, per

367
Q

conjugate base of carbonic acid

A

bicarbonate

368
Q

Hydration involves

A

breaking of water-water H-bonds, formation of water-solute bonds

369
Q

Ion number

A

number of water molecules that bind to an ion in an aqueous solution

370
Q

If a compound is composed of two non-metals, it must be a ____ compound

A

molecular

371
Q

A negative heat of solution indicates that the solute-solvent bonds are ____ than the solute-solute bonds.

A

stronger

372
Q

Sublimation occurs when the vapor pressure of the solid is _____ than the partial pressure above it

A

greater

373
Q

Boiling occurs when the vapor pressure of the liquid is greater than the _____ pressure above it

A

total

374
Q

Vapor pressure ____ with temperature

A

increases

375
Q

Atmospheric gas will condense when ____ pressure of the atmospheric gas is at least as great as the ____ pressure of the liquid at that temperature

A

partial, partial

376
Q

Steam refers to

A

water vapor above 100ºC

377
Q

A non volatile solute has ____ vapor pressure

A

zero

378
Q

Ionic compounds with greater charges are typically ____ soluble in water

A

less

379
Q

Deviation from Ksp in real exeriments is usually witness because of ____ ____ and _____ reactions that take ions out of solution

A

ion pairing, hydrolysis

380
Q

Salt crystals nucleate gas bubbles causing them to ____

A

coalesce

381
Q

Amorphous solids have ____ melting points

A

poorly defined

382
Q

Gases nearly always form ____ phase(s)

A

one

383
Q

Coffe cup calorimeters are ____ systems. They are used to measure ____ change.

A

open, enthalpy

384
Q

When an impurity is added to a solid, its melting point ____ and ____

A

decreases, broadens

385
Q

When an impurity is added to a liquid, its boiling point ____

A

increases

386
Q

vant hoff factor for non-electrolytes

A

1

387
Q

Strong acids

A

HCl, HNO3, H2SO4, HBr, HI, HClO4

388
Q

Amphiprotic

A

Capable of gaining or losing a proton

389
Q

Strong acids have ____ conjugate bases

A

weak

390
Q

Strong bases have ____ conjugate bases

A

weak

391
Q

weak conjugate acids can have ____ conjugate bases

A

weak

392
Q

Strong acids are stronger than ____

A

H3O+

393
Q

Strong bases are stronger than OH-

A

Define

394
Q

Strong bases

A

H-, Na2O, N(3-)

395
Q

Percent ionization of an acid is dependent on

A

temperature, identity of acid, concentration of acid

396
Q

In living organisms, small pH changes can great ____ rate chagnes

A

large

397
Q

Atoms with unfilled d orbitals can _____ visible light and move to these orbitals, causing brilliant colors in transition metals

A

absorb

398
Q

HCl reacts with CO3- to form

A

CO2

399
Q

Reaction involving the exchange of bonds between the chemical species. Often called double-displacement.

A

metathesis reaction

400
Q

The strength of oxyacids _____ with the central atoms oxidation state

A

increases

401
Q

Which acid is stronger HClO or HIO

A

HClO

402
Q

Which acid is stronger HCl or HI

A

HI

403
Q

Acidity increases with bond _____ and decreases with bond ____ and and increases with the conjugate base ____

A

polarity, strength, stability

404
Q

A hydride contains ____ and one other element

A

H-

405
Q

A substance with a high boiling point will generally have a ____ vapor pressure

A

low

406
Q

The melting point of water will _____ under low pressure

A

increase

407
Q

The melting point of most liquids will ____ under low pressure

A

decrease

408
Q

Osmolarity

A

Π=iMRT

409
Q

Isobaric

A

constant pressure

410
Q

Pico (p)

A

Define

411
Q

Alkaline Earths

A

Group IIa

412
Q

Ksp of MX

A

Define

413
Q

Homogenous Catalyst

A

same phase

414
Q

Heisenberg uncertainty principle

A

Define

415
Q

Reaction Entropy

A

Define

416
Q

Anode Galvanic

A

negative electrode (oxidation)

417
Q

Arrhenius equation

A

k = Ae^(-Ea/RT)

418
Q

Metals

A

Elements that are characteristically electropositive, malleable, and ductile. These elements tend to be found on the left side of the periodic table, be lustrous, and have relatively low ionization energies and electron affinities

419
Q

Halogens

A

Contains nonmetals, 7 valence electrons in it’s outermost energy level. Very reactive

420
Q

Equilbrium constant is ______ dependent

A

temperature

421
Q

Effusion

A

one gas moves through air

422
Q

Standard delta G =

A

-RT ln(K)

423
Q

Titration

A

Analytical procedure in which a solution of known concentrations is slowly added to a solution of unknown concentration to the point of molar equivalency, thereby providing the concentration of the known solution.

424
Q

Valence Electrons

A

Electrons occupying the outermost electron shell of an atom, participating in chemical bonds. Atoms with the same number of valence electrons tend to have similar properties (families in the Periodic Table).

425
Q

STP

A

Standard Temperature and Pressure. 273 Kelvin (0 Celsius), 1 atmosphere (760 torr, 760 kPA).

426
Q

Lewis definition

A

Acids defined as electron-pair acceptors and bases as electron-pair donors.

427
Q

Half-Cell

A

An electrode immersed in an electrolytic solution that is the site of either oxidation or reduction in a galvanic (voltaic) cell

428
Q

Noble gases

A

Contains nonmetals that are non-reactive. Full outermost energy level except helium which has 2.

429
Q

Which group is most metallic?

A

group IA

430
Q

Best insulator has _____ specific heat

A

highest

431
Q

Speed is dependent on

A

kinetic energy only

432
Q

autophagy

A

A process that describes lysosomes using their hydrolytic enzymes to recycle the cell’s own organic material

433
Q

Phase Diagram

A

A pressure vs. temperature plot showing the conditions under which a substance exists in equilibrium between different phases or in which the substance exists in pure phase.

434
Q

Ion

A

A single or polyatomic particle with an electric charge.

435
Q

Normality

A

Gram equivalent weight of solute per liter of solution, often denoted by N.

436
Q

Disproportionation

A

Redox reaction, in which the same species is both oxidized and reduced.

437
Q

Faraday’s Constant

A

Denoted by F, 9.65x10⁴ coulombs/mol e⁻. Commonly used in the formula It = nF (I = Current, t = time (s), n = mol e⁻).

438
Q

Indicator

A

A chemical species that changes color during dissociation, used to signal the end point of a titration.

439
Q

Redox Half-Reaction

A

Hypothetical equation showing only the species that is oxidized or reduced in a redox reaction and the correct number of electrons transferred between the species in the complete, balanced equation.

440
Q

solvation

A

sol, a chemical process in which solvent molecules and molecules or ions of the solute combine to form a compound

441
Q

Amphoteric

A

having characteristics of both an acid and a base and capable of reacting as either

442
Q

System

A

The part of the universe under consideration that is separated by some real or imaginary boundary from its surroundings

443
Q

atomic radius

A

one-half the distance between the nuclei of two atoms of the same element when the atoms are joined

444
Q

mole

A

the amount of a substance that contains as many particles as there are atoms in exactly 12g of carbon-12

445
Q

bond length

A

the average distance between the nuclei of two bonded atoms

446
Q

Decomposition reaction

A

(chemistry) separation of a substance into two or more substances that may differ from each other and from the original substance C»»A+B

447
Q

Molecule

A

The smallest unit of a substance, composed of two or more atoms joined in covalent bonds, that still retains all the chemical properties of that substance

448
Q

What charge to group 7 elements have?

A

-1

449
Q

Entropy when energy into the system

A

entropy increases

450
Q

Chemical Similarity between Atoms

A

stay in same column

451
Q

Money Metals

A

Au, Pt, Ag, Hg, Cu, Ni

452
Q

Speed of light (c)

A

3.00 x 10^8 m/s

453
Q

Ion Product

A

Product of the molar concentrations of dissociated ions in solution at any point in the reaction other than equilibrium or saturation, where each ion is raised to the power of its stoichiometric coefficient. Denoted IP.

454
Q

Molecular Formula

A

A chemical formula showing the actual number of atoms present in a certain compound.

455
Q

Limiting Reagent

A

Reactant of a chemical equation that, given nonstoichiometric amounts, determines the amount of product that can form; the reactant that runs out first.

456
Q

Molecular Orbital

A

Region in a molecule where atomic orbitals overlap, resulting in either a stable low-energy bonding orbital or an unstable high-energy antibonding orbital.

457
Q

Acid Dissociation Constant

A

An equilibrium expression used to measure weak-acid strength, given by the ratio of the product of the products’ molar concentrations to the product of the reactants’ molar concentrations, with each term raised to the power of its stoichiometric coefficient. Denoted Ka.

458
Q

Percent Yield

A

A ratio, calculated as a percentage, of the actual mass of product yielded to the theoretical yield of product mass.

459
Q

Isothermal System

A

Process in which the system either gains or loses energy to maintain a constant temperature.

460
Q

Isolated System

A

A system that can exchange neither energy nor matter with its surroundings.

461
Q

Exothermic

A

A reaction that proceeds with the net release of energy (heat) into the surroundings

462
Q

Magnetic quantum number

A

specifies the specific orbital in which the electron is most likely to be found., Third quantum number, designated as ml. Describes a particular orbital within a subshell where an electron is very likely to be found. Possible values are integers in the -1 to 1 range, including 0.

463
Q

Chemical Kinetics

A

the area of chemistry that is concerned with reaction rates and reaction mechanisms

464
Q

Rate-Determining Step

A

The slowest step in a reaction mechanism that determines the overall rate of the reaction

465
Q

Effective Nuclear Charge

A

The resulting positive nuclear charge an outer electron senses after accounting for the shielding effect of inner core electrons. Abbreviated Zeff. Increases from left to right and from bottom to top on the periodic table

466
Q

Concentration

A

tells you how much solute is present compared to the amount of solvent

467
Q

Theoretical Yield

A

The expected amount of product yielded in a reaction according to reactants’ stoichiometry

468
Q

Atomic Absorption Spectrum

A

The spectrum of certain absorbed wavelengths of light corresponding to an atom’s spectrum of emitted frequencies of light

469
Q

Diamagnetic

A

An atom or a substance that contains no unpaired electrons and is consequently repelled by a magnet

470
Q

Reduction

A

A reaction in which a species gains electrons

471
Q

Atom

A

The basic building block of all matter in the universe. An atom is made up of three main components: protons, neutrons, and electrons

472
Q

Azeotrope

A

A liquid mixture of two or more substances that has a constant boiling point greater than or less than the boiling points of its constituents. The vapor of this unique mixture has the same composition as the liquid state, making difficult to separate the constituents

473
Q

Isochoric Process

A

A process in which volume remains constant and no net pressure-volume work is done

474
Q

Bronsted-Lowry definition

A

Common definition of acids as proton (H+) donors and bases as proton acceptors

475
Q

What charge do group 1 elements tend to have?

A

1

476
Q

From left to right across the periodic table, metallic characteristics (increase, decrease)

A

decrease

477
Q

Metallic Solid

A

molecules held in place by delocalized bonding

478
Q

State Variables

A

conditions that must be specified to establish the state of the system, pressure, volume, temperature, and amounts of substances

479
Q

Glavanic Reduction potentials

A

higher reduction potential is cathode, lower reduction potential is anode

480
Q

Hund’s Rule

A

orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin

481
Q

Combination Reaction

A

A reaction in which two or more reactants combine into a single product.

482
Q

Pfund Series

A

Set of spectral lines resulting when a hydrogen atom undergoes a transition from energy levels n>5 to n=5.

483
Q

Ideal Gas Law

A

A unification of Boyle’s Charles and Gay-Lussac’s, and Avogadro’s Principle into a formula describing the behavior of ideal gases: PV=nRT. (PressureVolume = molesAvogadro’s Number*Temperature).

484
Q

Reducing Agent

A

A species that is oxidized in the process of reducing another species.

485
Q

Lyman Series

A

Set of spectral lines appearing in the UV region when a hydrogen atom undergoes a transition from energy levels n>1 to n=1.

486
Q

Specific Heat

A

The amount of heat required to raise one gram of a substance by 1 degree Celsius;heat capacity

487
Q

Molality

A

The ratio of the number of moles of solute dissolved in one kilogram of solvent. molality (M = moles solute/kg of solution)

488
Q

Adiabatic Process

A

A process in which no heat is transferred to or from the system by its surroundings

489
Q

Reaction Quotient

A

A ratio of the concentrations of the products to the concentrations of the reactants at any point during the reaction aside from equilibrium, where each reactant and product in the expression is raised to the power of its stoichiometric coefficient. Commonly denoted by Q

490
Q

Solution Equilibrium

A

When a solute is dissolved in a solvent, it will dissociate until reaching an equilibrium point at which the rate of dissociation equals the rate of precipitation of the solute, regardless of any additional solute introduced into the mixture

491
Q

Solubility Product Constant

A

Product of the molar concentrations of dissociated ions in solution at saturation, where each ion is raised to the power of its stoichiometric coefficient. Denoted Ksp.

492
Q

Aqueous Solution

A

a solution in which water is the solvent

493
Q

Hydrogen Bonding

A

involves lone pairs of electron on an electronegative atom of one molecule and a polar bond to hydrogen in another rmolecule. They are confined tomolecules that contain O, N, and F atoms

494
Q

Ideal Gas Postulates

A
  1. molecules are very small compared to the distance between them.2. molecules are constantly moving3. pressure of the gas- collisions of the molecules with container walls4. molecules do not experience intermolecular forces5. KEave proportional to T
495
Q

Osmotic Pressure

A

increase in pressure due to a solvent crossing a membrane into a more concentrated solutionΠV = nRT

496
Q

Assumptions of ideal gases

A

No volume, no repulsive forces, elastic collisions, kinetic energy is proportional to temperature

497
Q

Le Châtlier’s Principle

A

When a system in equilibrium is placed under one of several stressors, it will react in order to regain equilibrium. In other words, act on a system, it will work back towards equilibrium.

498
Q

Vapor Pressure

A

Partial pressure of a vapor when it is in equilibrium with its solid or liquid phase.

499
Q

Single Displacement Reaction

A

Chemical reaction in which an atom or ion of one compound is replaced by another atom or ion.

500
Q

Balmer Series

A

A set of spectral lines that appear in the visible light region when a hydrogen atom undergoes a transition from energy levels n>2 to n=2

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