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CHM135 > UNIT 3 > Flashcards

UNIT 3 Flashcards

1
Q

Define thermodynamics

A

the study of energy and its conversion from one form to another

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2
Q

Define thermochemistry

A

a branch of thermodynamics, which highlights how heat is involved in chemical and physical transformations

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3
Q

there are two basic kinds of energy:

A

potential energy (Ep): stored energy
kinetic energy (Ek): the energy of motion

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4
Q

Ek =

A

1/2mv^2

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5
Q

internal energy, E, of a system

A

the sum of all energy (potential and kinetic) of everything in that system

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6
Q

chemical bond energy

A

potential energy stored in molecular bonds

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7
Q

energy of molecular motion

A
  • translational, rotational, and vibrational
  • sometimes also called thermal energy
  • temperature is a common measure of the energy of molecular motion
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8
Q

define energy

A

the capacity to do work or supply heat

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9
Q

change in energy of a system =

A

work done on the system + heat flow in to the system
ΔE = w + q

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10
Q

energy is measured in

A

Joules or kiloJoules

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11
Q

1st law of thermodynamics

A

energy cannot be created or destroyed
ΔE(universe) = 0
the total energy of the universe is conserved

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12
Q

system

A

chosen to include whatever you are focusing on (and things that cannot usefully be separated from that)

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13
Q

surroundings

A

everything not in the system

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14
Q

universe =

A

system + surroundings

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15
Q

how can we determine the energy change of a system of interest and why?

A

by measuring the energy change in the surroundings; if energy leaves the system, it must enter the surroundings (and vice versa)

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16
Q

any change in the energy of a system must be accompanies by

A

an equal magnitude change in the energy of the surroundings, but the signs of these changes must be opposite

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17
Q

ΔE(universe) =

A

ΔE(system) + ΔE(surroundings) = 0

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18
Q

in thermodynamics, signs are defined from

A

the system’s point of view/perspective

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19
Q

ΔE of a reaction

A

= E(products) - E(reactants)

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20
Q

if a reaction releases energy

A
  • system loses energy
  • reaction produces heat or does work
  • reactants higher in energy than the products
    ΔE < 0
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21
Q

state function

A

values that depend on the state of the substance, and not on how that state was reached. For example, density is a state function, because a substance’s density is not affected by how the substance is obtained.

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22
Q

examples of state functions

A

E, V, P

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23
Q

examples of non-state functions

A

q and w

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24
Q

heat (q)

A

heat (q), also called thermal energy, will flow from higher-temperature objects to lower-temperature objects. eventually, the objects will reach the same temperature

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25
Q

work (w)

A

mechanical work is the product of force (F) and distance (d)

w = F x d

the larger the required force, or the longer the distance on an object is moved, the more work is done on it

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26
Q

example of energy transfer as work (w) only (q=0)

A

gas forming reaction in an insulated container that is attached to a piston-cylinder assembly that pushes against something
- the system pushes the piston out, doing work on the surroundings
- system releases (loses) energy
- ΔE = w < 0

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27
Q

expansion work

A

the work done when the volume of a system changes in the presence of an external pressure

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28
Q

expansion work is also often known as

A

pressure volume (PV) work because the amount of work done depends on both P and V

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29
Q

at a constant pressure, P=P(surr), w=

A

-PΔV = -P(V(final) - V(initial))

more work is done when the volume change (ΔV) is larger and/or when pushing against a higher external pressure

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30
Q

why is there a negative sign for expansion work?

A
  • required to fit with the convention that signs reflect the system’s perspective
  • when a system expands against a pressure:
    -> P, which is the constant pressure of the surrounding, is always >0
    -> the system does work. it loses internal energy by doing this work
    -> it must be that w<0, ΔE<0

we calculate: w = -PΔV <0

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31
Q

enthalpy (ΔH)

A

defined as

H = E + PV
another accounting system for keeping track of energy, similar to internal energy, but does not include expansion work

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32
Q

change in enthalpy (ΔH)

A

ΔH = ΔE + PΔV
- amount of heat absorbed or released in a transformation (real or imagined reaction) at a constant pressure.
- this means enthalpy change (ΔH) is often straightforward to determine via experiment

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33
Q

the sign of ΔH shows

A

whether heat is produced or consumed in a reaction.

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34
Q

ΔH<0

A

exothermic
- evolved heat flows out of the system into the surroundings
- heat is a product of the reaction

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35
Q

ΔH>0

A

endothermic
- heat flows into the system from the surroundings
- heat is a reactant

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36
Q

is enthalpy a state function?

A

yes.
ΔH = H(products) - H(reactants) = H(final) - H(initial)

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37
Q

differences and similarities between enthalpy (ΔH) and internal energy change (ΔE)

A
  • differ in whether they account for the expansion (PV) work done on the surroundings or done by the surroundings on the system
  • most reactions involve little (if any) PV work. this is especially true for reactions involving only solids and liquids, or reactions in which the amount of gas does not change

thus, often (but not always) ΔH = ΔE

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38
Q

effect of multiplying a reaction on ΔH

A

increases ΔH by the same factor

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39
Q

effect of reversing a reaction on ΔH

A

changes the sign of ΔH for a reaction

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40
Q

heat capacity

A

amount of heat (q, or ΔH) required to raise the temperature of an object or substance by 1’C (or 1K)

heat capacity = q/ΔT

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41
Q

things with high heat capacity

A
  • require a lot of heat to increase in temperature and give off a lot of heat when they cool down.
  • act as a good sponge for (or, store of) thermal energy
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42
Q

types of heat capacity

A

specific heat capacity (c): heat capacity per mass (1.00g)
c = q/(mass x ΔT)
units of J/gK

molar heat capacity (C or Cm): heat capacity per mole
C = q/(amount (mol) x ΔT)
units of J/molK

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43
Q

effect of energy from heat, q

A

excites both translational motion of molecules and vibrations and rotations within and between molecules

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44
Q

calorimetry

A

the science of measuring the heat exchanged in chemical reactions

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45
Q

in calorimetry, the heat of reaction (qrxn) is measured

A

indirectly, by means of a calorimeter.
- if the reaction produces heat, the temperature of the surroundings increases and vice versa

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46
Q

q(rxn)

A

= - q(calorimeter)
= -(q(vessel) + q(solution) + q(other))

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47
Q

constant-pressure, or ‘coffee-cup’ calorimetry

A

q(system) = - q(calorimeter)
- reaction or physical transformation is done in an insulated container at a constant pressure in a bath (usually water) of known heat capacity
- heat generated thus tells us about the enthalpy of reaction

q = qp = ΔE + PΔV = ΔH

  • heat capacity of the calorimeter (often ~equal to that of the surrounding water because absorption of heat by the vessel is minimal) provides the link between change in temperature and heat gained or lost, from which you can determine the enthalpy change of the system

Heat capacity (calorimeter) = q(cal)/ΔT(cal)

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48
Q

second type of calorimetry

A

constant volume or ‘bomb’ calorimetry:
- measures the heat change at constant volume such that q = qv = ΔE + PΔV = ΔE
- commonly used to measure heat of combustion reactions

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49
Q

Hess’s Law

A

the enthalpy change for a process is equal to the sum of the changes for individual steps (1, 2, …n) of the process
ΔH(overall) = ΔH1 + ΔH2 + …. + ΔHn

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50
Q

standard states

A
  • pure substance: most stable form at 1atm
  • gas: 1atm and ideal behaviour
  • substance in aqueous solution: 1M concentration all at a specified temperature, which is usually 25’C
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51
Q

standard enthalpy of reaction

A

ΔH’(rxn)
- all reactants and products are in standard states

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52
Q

ΔH’f

A

standard heat of formation
- enthalpy change for the formation of 1 mole of substance in its standard state from its constituent elements in their standard states

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53
Q

we can use our database of ΔH’f’s to calculate ΔH (rxn)

A

ΔH’(rxn) = ΔH’f(products) - ΔH’f(reactants)
- ΔH of each reactant or product must be multiplied by its stoichiometric coefficient in the balanced rxn equation

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54
Q

ΔH’(rxn) =

A

bonds broken - bonds made

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55
Q

Spontaneous change

A

one that occurs without a continuous input of energy from outside the system (though activation energy may be required to initiate it)

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56
Q

If a change is spontaneous in one direction

A

it will be non-spontaneous in the reverse direction

57
Q

how may non-spontaneous reactions and processes be driven?

A

with continual input of energy

58
Q

is enthalpy change a predictor of spontaneity?

A

many exothermic processes are spontaneous (eg combustion reactions)

59
Q

examples of spontaneous endothermic transformations

A
  • melting of ice (at higher temps)
  • dissolution (of some solids at some concentrations and temps)

all show an increase in the freedom of motion of particles in the system

60
Q

relationship between freedom of motion and spontaneity

A

an increase in freedom of motion (dispersal of energy) favours spontaneity

61
Q

entropy, S

A

measure of energy dispersal, or freedom of motion, in a system

62
Q

positive value of ΔS indicates
negative value of ΔS indicates

A

increased dispersal of energy
decreased dispersal of energy

63
Q

is entropy a state function?

A

yes
ΔS = S(final) - S(initial)

64
Q

units of entropy

A

J/K

65
Q

S’

A

standard molar entropy: entropy of 1 mole of the pure substance in its standard state

66
Q

entropy trends

A
  • effect of physical state
  • effect of particle numbers
  • effect of molecular complexity
  • effect of temperature
67
Q

effect of physical state

A

S of solids < S of liquids &laquo_space;S of gas
solids - less energy dispersed, lower entropy

68
Q

effect of particle numbers

A

more molecules have higher entropy than fewer molecules

69
Q

effect of molecular complexity

A

entropy increases with chemical complexity and flexibility
- this only holds for substances in the same physical state
- the effect of physical state dominates the effect of molecular complexity

70
Q

effect of temperature

A

as temperature increases, entropy increases
- higher temperature means more freedom of molecular motion

71
Q

draw graph for temperature vs entropy

A

discontinuous jumps at phase changes

72
Q

entropy change upon dissolution

A
  • salt gains entropy as it is dispersed
  • water loses entropy as it is ordered around the ions
  • net entropy change depends on the relative magnitudes of entropy changes in both solute and solvent entropy
73
Q

why does increased entropy favour spontaneous change?

A

high-entropy configurations can be achieved in more ways than low-entropy configurations. they are therefore more likely to occur

74
Q

draw a diagram for spontaneous expansion of a gas and explain it

A
  • arrangement B has higher entropy (is entropically favoured)
  • gas expands when stopcock is opened because there are more ways to achieve the configuration on the right (B) than on the left (A)
75
Q

absolute entropies can be calculated from

A

the number of micro states (W) a system may occupy

76
Q

2nd law of thermodynamics

A

spontaneous reactions proceed in the direction that increases the entropy of the universe (system + surroundings)

ΔS(Universe) = ΔS(sys) + ΔS(surr) > 0

thus, any decrease int he entropy of the system must be offset by a larger increase in the entropy of the surroundings for that process to be spontaneous

77
Q

2nd law has profound implications:

A
  • isolated systems always evolve toward higher entropy states
  • entropy of the universe is always increasing
78
Q

how are temperature, heat flow, and entropy linked?

A
  • if heat flows into a system from the surroundings, the entropy of the system increases. the surroundings lose entropy
  • the amount by which a given amount of heat flow changes entropy depends on temperature. if it is low, the effect on entropy can be enormous
79
Q

3rd law of thermodynamics

A

a perfect crystal has zero entropy at absolute zero

S(sys) = 0 at 0K

has flawless alignment of all its particles. at absolute zero, particles have minimum energy so there is only one micro state

80
Q

contrast between S and H

A
  • entropy scale is anchored to an absolute value
  • enthalpy does not have an absolute 0
81
Q

standard entropy of a reaction ΔS’(rxn)

A

entropy change that occurs when all reactants and products are in their standard states

ΔS’(rxn) = S’(products) - S’(reactants)

remember to include coefficients

82
Q

Gibbs free energy change ΔG

A

evaluates spontaneity as a function of enthalpy and entropy of the system alone

ΔG(sys) = ΔH(sys) - TΔS(sys)

ΔG < 0 process is spontaneous
ΔG = 0 process is at equilibrium
ΔG > 0 process is non-spontaneous

  • lowering free energy is the driving force of chemical reactions
  • negative ΔH(sys) and positive ΔS(sys) favour spontaneity
  • entropic contribution to free energy change (-TΔS) is increasingly important at higher temperatures
83
Q

give another way to define ΔG

A

the maximum useful work that can be done by a system as it undergoes a spontaneous process at constant temperature and pressure

ΔG = w(max)

also the minimum work that must be done on a system to drive the occurrence of a non-spontaneous process

84
Q

is ΔG a state function?

A

yes

85
Q

describe the extensive property of ΔG

A

scales linearly with amount

86
Q

ΔG’f

A

standard free energy of formation of a compound from its constituent elements in their standard states

87
Q

ΔG’f of an element in its standard state is

A

0

88
Q

thermodynamics vs kinetics

A

ΔG tells us whether a reaction will/won’t proceed
the free energy of activation (including Ea) tells us how fast a reaction proceeds

89
Q

ΔG and spontaneity

A

reaction is spontaneous when ΔG(rxn)<0

90
Q

how do we make a non-spontaneous reaction happen?

A

must be driven by coupling the non-spontaneous reaction with a spontaneous reaction of sufficiently favourable ΔG

91
Q

how are free energies and equilibrium position linked?

A

ΔG depends on how much product and reactant is present at that instant (Q0 compared to their equilibrium values (K), the temperature and the gas constant:

ΔG = RTln(Q/K)

92
Q

how does the magnitude of ΔG tell us how far out of equilibrium the mixture is?

A

ΔG < 0 ; Q < K ; ln(Q/K) < 0 - process proceeds (forward spontaneously)

ΔG = 0 ; Q = K ; ln(Q/K) = 0 - process is at equilibrium

ΔG > 0 ; Q > K ; ln(Q/K) > 0 - reverse process proceeds spontaneously

93
Q

if Q and K are very different

A

ΔG has a very large value (negative or positive). The reaction releases or absorbs a large amount of free energy as it proceeds to equilibrium

94
Q

if Q and K are nearly the same,

A

ΔG has a very small value (negative or positive). The reaction releases or absorbs very little free energy as it proceeds to equilibrium.

95
Q

how are thermodynamic Q and K different to Qc and Kc

A

in the thermodynamic Q and K, all substances are referenced to their own standard states, So, Q and K may include mixed states

96
Q

equation linking ΔG, gas constant, temp, and K

A

-RTln(K)

97
Q

another formula for ΔG

A

= ΔG’ + RTln(Q)

98
Q

draw the two free energy hills with ΔG

A
99
Q

features of free energy hills

A
  • slope at any given point tells us the value of ΔG for that mixture
  • as the system approaches equilibrium, ΔG approach 0
  • at equilibrium, the free energy is at a minimum.
100
Q

Define electrochemistry

A

the interconversion of chemical and electrical energy
- spontaneous reactions can produce electricity and electricity can cause non-spontaneous reactions to occur

101
Q

redox reactions involve

A

the movement of electrons from one reagent to another

102
Q

oxidation

A

loss of electrons

103
Q

reduction

A

gain of electrons

104
Q

oxidising agent

A
  • reduced
  • takes electrons from the substance being oxidised
105
Q

reducing agent

A
  • oxidised
  • gives electrons away
106
Q

how to balance redox reactions

A
  • write down the two half-reactions
  • balance the atoms and charges in each half-reaction
  • first balance atoms other than O and H, then O, then H
  • charge is b balanced by adding electrons (e-) to the reactant side of the reduction half-reaction and to the product side in the oxidation half-reaction
  • electrons should cancel out in net reaction. if necessary, multiply one or both half-reactions by an integer so that number of e- gained in reduction = number of electrons lost in oxidation
  • add the balanced half-reactions and include states of matter
107
Q

two types of electrochemical cells

A
  • voltaic cells
  • electrolytic cells
108
Q

voltaic/galvanic cells

A

spontaneous chemical reaction (ΔG<0) generates an electric current
- batteries contain one or more voltaic cells
- voltaic cell does work on the surroundings, converting higher energy reactants in the cell into lower energy products

109
Q

electrolytic cells

A

non-spontaneous reactions (ΔG>0) are driven by electric current
- external power source supplies free energy to run electrolytic cells. the surroundings thus do work on the cell. lower energy reactants are converted to higher energy products in the cell
- used for electroplating, purification of metals, and more

110
Q

electric current is

A

flow of electrons

111
Q

electrons only flow if

A

the driving force (free energy change) is large enough

112
Q

electrodes

A

usually metal strips/wires connected by an electrically conducting wire

113
Q

anode

A

electrode where oxidation takes place

114
Q

cathode

A

electrode where reduction takes place

115
Q

describe salt bridge

A

U-shaped tube that contains a gel permeated with a solution of inert electrolyte (contains positive and negative spectator ions).

116
Q

function of salt bridge

A

The salt bridge keeps half cells electrically neutral because ions flow in and out of the salt bridge, counteracting charge build-up due to electron flow.

117
Q

conventions for notation for a voltaic cells

A
  • anode components are written on the left
  • cathode components are written on the right
  • components of each half cell are written in the same order as in their half reactions
  • single line shows a phase boundary between the components of a half cell
  • double line shows that the half cells are physically separated
118
Q

addition notation for voltaic cell that is more complex

A

any inactive (inert) electrode is specified
a comma is used to show components that are in the same phase

119
Q

why does a voltaic cell work?

A
  • differing abilities of metals to gain electrons gives rise to a voltage drop
  • this is also known as electromotive force (EMF) or cell potential
120
Q

multimeter

A

can measure voltage

121
Q

cell potential (Ecell)

A

the difference in electrical potential between two electrodes

122
Q

units for Ecell

A

measured in Volts (V), where

V = J/C (Colomb is SI unit of charge)

123
Q

difference in Ecell for voltaic and electrolytic cells

A

Voltaic: Ecell>0 (spontaneous e- flow)
electrolytic: Ecell<0 (non-spontaneous e-flow)

124
Q

when Ecell = 0,

A

the redox reaction has reached equilibrium so the cell can do no more work

125
Q

define standard potential (E’cell)

A

cell potential under standard state conditions

126
Q

how to calculate Ecell from Ehalf-cell

A

the cell potential of any electrochemical cell is the sum of the half cell potentials for the oxidation and reduction half cells

Ecell = E ox + E red

127
Q

what does it mean to say that E is an intensive property?

A

Bigger cells with more moles of redox components will last longer, but will have the same standard output voltage

128
Q

which half cell will form the oxidation half cell?

A

the lower Ered (ie more negative/less positive = worse at being reduced)

129
Q

use of standard hydrogen electrode

A
  • potentials are determined experimentally from the difference in potential between two electrodes
  • the reference point is called the standard hydrogen electrode
130
Q

Standard Hydrogen Electrode

A

consists of a platinum electrode in contact with H2 gas (1atm) and aqueous H+ ions (1M)

131
Q

standard hydrogen electrode is assigned (arbitrarily) a value of

A

exactly 0.00 V

132
Q

how is Ecelll related to ΔG?

A

ΔG = -nFE(cell) = -RTln(Keq)

  • ΔG is in J/mol
  • n is in mol and is the no of moles of electrons transferred per mole of the reaction
  • F is faraday’s constant
133
Q

at equilibrium, ΔG = and Ecell =

A

0

134
Q

Nernst Equation

A

E = E’ - RT/nF (lnQ)

E = E’ - 0.0592V/n logQ

E - The cell potential (electromotive force, EMF) under non-standard conditions, measured in volts (V).
The standard cell potential, which is the voltage of the electrochemical cell under standard conditions

135
Q

how does cell potential depend on the relative concentration of reactants and products?

A

when Q<1, lnQ<0 so Ecell>E’cell
when Q=1, lnQ=0 so Ecell = E’ cell
when Q>1, lnQ>1 so Ecell<E’cell

136
Q

two consequences of cell potential depending on the relative concentration of reactants and products:

A
  • as a cell is operated, concentration of reactants will decrease and products will increase. thus, cell potential will decrease over time
  • concentration cells - capture the electrical energy from a concentration difference
137
Q

how is a concentration cell created

A
  • has the same half-reaction in both cell compartments, but with different concentrations of electrolyte
  • there is a potential difference between cells, which drives current flow until both compartments have an equal concentration of ions
  • Ecell>0 as long as the half-cell concentrations are different
  • once the concentrations equalise, Ecell = 0 and current stops flowing
138
Q

write equations for corrosion

A
139
Q

how to protect against corrosion

A

Galvanisation: coating of iron with zinc.
Zinc is more easily oxidised than iron, giving up electrons to it. Iron is more easily reduced than zinc.

CHM135 (23 decks)

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