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Chemistry 1 > Lectures 21-26 > Flashcards

Lectures 21-26 Flashcards

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1
Q

energy

A

quantitative property that provides the ability to ‘do work’ or ‘supply heat’

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2
Q

potential energy (locked-in energy)

A

stored energy

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3
Q

kinetic energy (movement)

A

energy in motion

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4
Q

thermodynamics

A

science of the relationship between heat and other forms of energy

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5
Q

1 J =

A

1 Nm

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6
Q

1 J =

A

1 kg m^2s^-2

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7
Q

calorie (cal)

A

energy required to raise the temperature of 1g of water by 1 degree Celsius (1 cal = 4.184 J)

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8
Q

Calorie (C)

A

used to represent energy content of our food; 1 C = 1000 cal

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9
Q

The Law of Conservation of Energy

A

energy can be converted from one form into another, but it CANNOT be created or destroyed. The total energy of the universe is constant

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10
Q

universe

A

system + surrounding

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11
Q

open system

A

matter can be transferred through opening in the flask; heat can be conducted through flask walls

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12
Q

closed system

A

flask is stoppered so no matter can be transferred; heat can be conducted through the flask walls

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13
Q

isolated system

A

heat transfer is prevented by the vacuum flask; matter cannot be transferred

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14
Q

isothermal change

A

heat is exchanged between system and surroundings, so their temperatures are equal

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15
Q

adiabatic change

A

no heat exchange between system and surroundings, so their temperatures may not be equal

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16
Q

heat (q)

A

energy transferred from one system to another due to a temperature difference

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17
Q

temperature (T)

A

measure of ‘how hot’ something is (kinetic energy) and ability to transfer heat to other systems or surroundings

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18
Q

extensive property

A

a property of matter that changes as the amount of matter changes

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19
Q

is heat an extensive or intensive property?

A

extensive

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20
Q

intensive property

A

property of matter that does not change as the amount of matter changes

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21
Q

is temperature an extensive or intensive property?

A

intensive

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22
Q

work (w)

A

involves energy exchange as a result of motion against an opposing force (F)

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23
Q

work equation

A

w = Fd

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24
Q

when a system does work on the surroundings,

A

it loses energy

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25
Q

when the surroundings do work on the system,

A

it gains energy

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26
Q

volume work equation

A

w = -p(Ext)deltaV where p(ext) is external pressure

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27
Q

endothermic reaction

A

chemical process that absorbs heat

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28
Q

exothermic reaction

A

chemical process that releases heat

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29
Q

activation energy

A

energy required to break the bonds of the reactants

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30
Q

state functions

A

depend only on the amount of substance and its conditions; pathway does not matter

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31
Q

internal energy (U)

A

the sum of all its kinetic and potential energies of all the atoms, ions and molecules within the system

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32
Q

internal energy is an

A

extensive property and a state function

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33
Q

deltaU equation

A

deltaU = U(final) - U(initial)

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34
Q

change in internal energy in a closed system

A

deltaU = q + w

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35
Q

system gains internal energy

A

deltaU > 0

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36
Q

system loses internal energy

A

deltaU < 0

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37
Q

K and Celsius interchangeable for

A

SHC calculations

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38
Q

enthalpy change (deltaH)

A

equal to the heat transferred between the system and surroundings during a process that occurs at constant pressure

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39
Q

heat transferred under constant pressure

A

q(p) (deltaH = q(p))

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40
Q

enthalpy is a

A

state function

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41
Q

deltaH for exothermic

A

< 0

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42
Q

deltaH for endothermic

A

> 0

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43
Q

relationship between internal energy and heat gained/lost

A

deltaH = deltaU + pdeltaV

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44
Q

for systems containing no gases, deltaH

A

approximately equals deltaU

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45
Q

for systems containing gases, deltaH

A

= deltaU + deltan(gas)RT

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46
Q

molar enthalpy change of fusion/latent heat of fusion

A

energy required to melt 1 mol of a pure substance at its melting point

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47
Q

molar enthalpy change of vaporisation/latent heat of vaporisation

A

energy required to vaporise one mole of a pure substance at its boiling point

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48
Q

molar enthalpy change of fusion for water

A

6.00 kJ mol^-1

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49
Q

molar enthalpy change of vaporisation for water

A

40.65 kj mol^-1

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50
Q

heat capacity

A

tells us how much heat energy we need to put into a substance to raise its temperature

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51
Q

specific heat capacity (Cs)

A

heat needed to raise the temperature of 1g of substance by 1 K

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52
Q

specific heat capacity equation

A

Cs = q/mdeltaT

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53
Q

molar heat capacity (Cm)

A

heat required to raise the temperature of 1 mol of substance by 1 K

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54
Q

molar heat capacity equation

A

Cm = q/ndeltaT

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55
Q

for gases, heat capacity depends on whether measurement is carried out at

A

constant pressure (Cp) or constant volume (Cv)

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56
Q

molar heat capacity at constant pressure

A

Cp = q(p)/ndeltaT J K^-1 mol^-1

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57
Q

molar heat capacity at constant volume

A

Cv = q(v)/ndeltaT J K^-1 mol^-1

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58
Q

enthalpy change of a reaction (deltarH)

A

the difference between the sum of enthalpies of the products and the sum of enthalpies of the reactants

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59
Q

solution calorimetry

A

performed under constant pressure

60
Q

bomb calorimetry

A

performed under constant volume

61
Q

equation for heat transferred by a particular amount of substance

A

q = C x m x deltaT

62
Q

as bomb calorimeters maintain constant volume, deltaU =

A

q(v) (no work term)

63
Q

enthalpy change of formation

A

enthalpy change when 1 mol of a substance is formed under

standard conditions, from its constituent elements in their standard states

64
Q

standard enthalpy change of a reaction (deltarH^o)

A

the enthalpy change when all the reactants and products are in their standard states

65
Q

standard conditinos

A
66
Q

standard conditions

A

1 M, 1 bar, 298 K

67
Q

Hess’s Law

A

states if a reaction can be written as the sum of two or more steps, its enthalpy change of reaction is the sum of the enthalpy changes of reaction of each of the steps

68
Q

standard enthalpy change of formation for a substance in its standard state

A

zero

69
Q

standard enthalpy change of formation equation

A

v(i)standard enthalpy change of formation of products - v(i)standard enthalpy change of reactants

70
Q

homolytic dissociation

A

cutting through a bond right down the middle

71
Q

bond-breaking costs energy and is therefore

A

endothermic

72
Q

bond energy (D)

A

the energy required to break 1 mol of bonds

73
Q

calculating enthalpy change using bond energies

A

deltarH = D(bonds broken) - D(bonds formed)

74
Q

summary of ways to calculate enthalpy changes

A

using steps (Hess’s Law), using enthalpy changes of formation equation, using bond energies

75
Q

when temperature increases, the enthalpy of a substance

A

increases

76
Q

calculating deltaH using molar heat capacity at constant pressure

A

deltaH = C(p) x deltaT

77
Q

calculating deltaC(p)

A

v(i)C(p)products - v(i)C(p)reactants

78
Q

Kirchoff Equation

A

standard enthalpy change at T2 = standard enthalpy change at T1 + deltaC(p)(T2 - T1)

79
Q

system absorbs heat

A

q > 0

80
Q

system gives off heat

A

q < 0

81
Q

surroundings does work on system

A

w > 0

82
Q

system does work on surroundings/work done by system

A

w < 0

83
Q

spontaneous process

A

process which, once started, occurs without any external intervention or action

84
Q

some spontaneous processes involve no

A

exchange with heat with surroundings at all

85
Q

some spontaneous reactions occur

A

very slowly

86
Q

entropy

A

degree of disorder in a system, its surroundings and the universe

87
Q

higher the degree of disorder,

A

higher the entropy

88
Q

probability of a disordered state is

A

higher than the probability of an ordered state

89
Q

Second Law of Thermodynamics

A

spontaneous processes are those that increase the total entropy of the universe

90
Q

to calculate SPONTANEOUS change of the universe

A

deltaS(universe) = deltaS(system) + deltaS(surrounding) > 0

91
Q

microstate

A

each possible arrangement of gas molecules

92
Q

macrostate

A

equivalent microstates

93
Q

probability of macrostate

A

add up probabilities of its microstates

94
Q

when the number of atoms is large, we reach the

A

thermodynamic limit

95
Q

as number of atoms increases,

A

distribution of matter with a high degree of order is very unlikely

96
Q

driving force in dispersal of matter

A

maximising system’s entropy

97
Q

maximally dispersed

A

more disordered

98
Q

exception to mixtures always leading to greater disorder (example)

A

LiOH crystal lattice

99
Q

Boltzmann Equation

A

S = kBln(W)

100
Q

Boltzmann constant, kB

A

R/N(A)

101
Q

W

A

number of microstates in a given macrostate

102
Q

more microstates in a given macrostate, the higher the

A

entropy

103
Q

how to calculate W

A

number of possible arrangements for one particle (w) to the power of the number of particles

104
Q

entropy sign of UNIVERSE governs whether reaction is

A

spontaneous or not (even if reaction entropy change is negative)

105
Q

reversible process

A

system changes in such a way that the system and surroundings can be put back in their original states by exactly reversing the process

106
Q

in a reversible system, the entropy of the

A

universe does not change

107
Q

phase changes are

A

reversible

108
Q

Third Law of Thermodynamics

A

there is no disorder (S = 0) in a perfect crystal at 0 K

109
Q

entropy increases with

A

temperature

110
Q

entropy is a

A

state function

111
Q

how to calculate deltaS(sys)

A

q(rev)/T J K^-1

112
Q

standard molar entropy (S^o)

A

defining entropy for 1 mol of substance under standard conditions

113
Q

entropies of substances increase as they change

A

from solid through to liquid and then gas (more freedom of motion)

114
Q

entropies of substances composed of larger molecules have

A

higher S^o than those made up of smaller molecules

115
Q

substances whose molecules have a more complex structure have a

A

higher S^o than substances with simpler molecules

116
Q

within the same phase, the entropy of 1 mol of substance for a temperature T2

A

S^o(T2) = S^o(T1) + C(p)ln(T2/T1)

117
Q

deltarS < 0 from

A

gas to liquid

118
Q

how to calculate deltarS^o

A

v(i)S^o(products) - v(i)S^o(reactants)

119
Q

if deltaS(univ) < 0, the reaction is

A

spontaneous in the opposite direction

120
Q

if deltaS(univ) = 0, the system is at

A

equilibrium

121
Q

deltaS(sys) is the same as

A

deltarS

122
Q

how to calculate deltaS(surr)

A

-deltarH(sys)/T

123
Q

how to calculate delta(vap)S^o

A

delta(vap)H^o/T(boiling point)

124
Q

how to calculate delta(fus)S^o

A

delta(fus)H^o/T(melting point)

125
Q

above freezing point

A

spontaneous melting

126
Q

below freezing point

A

melting not spontaneous

127
Q

Gibb’s free energy change (deltaG) equation

A

deltaH(sys) - TdeltaS(sys)

128
Q

Gibb’s free energy change

A

thermodynamic function that brings together the properties of enthalpy and entropy for a system

129
Q

units for Gibb’s free energy

A

J mol^-1 or kJ mol^-1

130
Q

Gibb’s free energy is a

A

state function

131
Q

Gibb’s free energy change for a reaction (deltarG)

A

G(products) - G(products)

132
Q

if deltarG < 0, reaction is

A

spontaneous

133
Q

if deltarG is > 0, reaction is

A

non-spontaneous

134
Q

if deltarG = 0, reaction is

A

at equilibrium

135
Q

if reaction is exothermic and has negative deltaS, it will be spontaneous

A

at low temperatures

136
Q

if reaction is exothermic and has positive deltaS, it will be spontaneous

A

at all temperatures

137
Q

if a reaction is endothermic and has positive deltaS, it will be spontaneous

A

at high temperatures

138
Q

if a reaction is endothermic and has negative deltaS, it will be non-spontaneous

A

at all temperatures

139
Q

product is favoured,

A

reaction is entropy driven

140
Q

standard Gibbs free energy of formation (deltafG^o)

A

change in free energy when 1 mol of pure substance is formed (in its standard state) from its component elements (in their standard states) at 1 bar pressure

141
Q

standard Gibbs free energy of formation (deltafG^o) equation

A

v(i)deltaG(products) - v(i)deltaG(reactants)

142
Q

elements in their standard states have deltafG^o = 0 and deltafH^o = 0, but their So at 298 K is

A

nonzero

143
Q

Hess’s Law also applies to

A

Gibb’s free energy change

144
Q

free energy work is equivalent to

A

maximum non-expansion work that can be obtained from a change; maximum work after change in volume has been accounted for

145
Q

maximum amount of energy that can be used to do work

A

deltarG^o = w(max)

146
Q

product favoured

A

G < 0

147
Q

reactant favoured

A

G > 0

Chemistry 1 (4 decks)

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