Thermodynamics Flashcards
1
Q
What is thermodynamics?
A
Thermodynamics describes the macroscopic state of a (microscopically) complex system through a small number of macroscopic variables (e.g. pressure, temperature), state-variables, and through thermodynamic potentials.
2
Q
In summary, what is thermodynamics?
A
Thermodynamics summarises the properties of energy and its transformation from one form to another (deals with the relations between heat and other forms of energy)
3
Q
Can matter and energy be destroyed?
A
No, A from of matter is only converted to another form of matter A form of energy is converted to another form of energy
4
Q
Universe is made of two main components, what are they?
A
System and Surroundings
5
Q
What is a System?
A
What we are interested in (i.e. a block of iron, a beaker of water, an engine, a human body, etc.)
6
Q
What are the Surroundings?
A
The remainder of the Universe outside the system (where we stand to make observations about the system and infer its properties)
7
Q
What is an open system?
A
where matter and energy can be exchanged between system and surroundings (i.e. an open flask)
8
Q
What is a closed (diathermic) system?
A
Only energy can be exchanged (i.e. a sealed bottle)
9
Q
What is an isolated system?
A
neither matter or energy can be exchanged (i.e. a stoppered vacuum flask)
10
Q
The properties of a system depend on what?
A
the prevailing conditions : extensive and intensive properties
11
Q
What are extensive properties?
A
Depending on the quantity of matter in the system (i.e. mass, volume, etc. – 2 Kg of iron occupy twice the volume of 1 Kg of iron)
12
Q
What are intensive properties?
A
Independent of the amount of matter present (i.e. temperature, density, etc. - The density of iron is 8.9 Kg.cm3 regardless of whether we have a 1 Kg block or a 2 Kg block)
13
Q
What is a state function?
A
A state function is one whose value depends only on the state of the substance under consideration; it has the same value for a given state no matter how that state came about.
14
Q
What are the three reactions that are state functions?
A
Internal Energy (U), Enthalpy, H and Entropy, S
15
Q
What is a path function?
A
A path function depends on the path which the system takes in going between two states.
16
Q
What are the two state functions?
A
heat (q) and work (w)
17
Q
What is the 0th Law of Thermodynamics?
A
All parts of system within a thermodynamic equilibrium have the same temperature
18
Q
What is the 1st Law of Thermodynamics?
A
Conservation of energy. The inner energy U of a system can only be changed by heat supply to/extraction from or work done on/performed by the system. (in other words: energy can not be created or destroyed, but only transformed)
19
Q
What is the 2nd Law of Thermodynamics?
A
Direction of state-changes. Reversible processes have zero-change in entropy S. Irreversible (spontaneous) processes have a positive change in entropy. In other words: state-changes will follow the direction of maximum entropy(-change).
20
Q
What is the 3rd Law of Thermodynamics (Nernst Theorem):
A
Approaching zero-temperature (T → 0) , the entropy becomes constant (set to 0). That is, the absolute zero-temperature (T = 0) cannot be reached.
21
Q
What is the basis of all thermometers?
A
If two bodies/systems having different temperature are brought in “contact’’ then the warmer body will get colder and the colder body will get warmer unti an equilibrium temperature has been reached.
22
Q
What is the (thermodynamic) temperature T is measured in?
A
Kelvin
23
Q
How did Lord Kelvin (W. Thomson, 1824-1907) define the temperature scale?
A
Defined the temperature-scale by the triple-point of water.
24
Q
What is the triple-point of water?
A
The triple-point is the point at which the solid, liquid and gaseous state of a system is in equilibrium, at a particular temp and pressure.
25
What is Kelvin?
1K is the 273,16th part of the thermodynamic temperature of the triplepoint of water (0.01°C).
26
What is the celcius scale?
The Celsius Scale uses the melting- and the boiling-point of water (0 and 100 degrees respectively) under normal atmospheric pressure to define temperature
27
How are the celsius and kelvin scales related?
Ø / C = T/K - 273,15
28
What thermometers are based on thermal expansion?
Liquid-glass (mercury, alcohol) thermometer Vapour-pressure thermometer Spring thermometer Bimetal thermomete
29
What thermometers are based on electrical propertise?
Resistance thermometer Thermoelement (thermo-voltage)
30
What thermometers are without mechanical and electrical contact?
Pyrometer (electromagnetic radiation) Acustic thermometer (velocity of sound) Magnetic thermometer (magnntic susceptibility) Glass-fibre thermometer (refractive index)
31
Which of the thermometers is the least accurate?
Mercury
32
What is thermal expansion?
Thermal expansion is the tendency of matter to change in shape, area, and volume in response to a change in temperature. (When a substance is heated, the kinetic energy of its molecules increases)
33
What is Charles law?
Charles' Law is a principle that deals with the effect of heat on the expansion of gases. The volume of a gas increases when heated at constant pressure
34
What occurs when a gas is heated (Charle's Law, Thermal Expansion)
• When a gas is heated, the gas molecules move faster and hit the wall of the container violently. • The volume of gas must increase to keep the pressure constant. So that the gas molecules hit the wall less frequently.
35
What do the volumes of gases depend on and who was it thanks to?
The experiments of J.A.C. Charles and J.L. Gay-Lussac have shown that the volume of gases depends on both pressure and temperature.
36
What is Gay Lussacs law equation?
P1/T1 = P2/T2
37
What is Charle's law equation?
V1/T1 = V2/T2
38
Wha is Gay Lussacs law?
When gases react, the volumes of the reacting gases and any gaseous products bear a simple whole number ratio when volumes are measured at the same temperatures and pressure
39
What is an ideal gas?
The ideal gas is a model-system consisting of non-interacting particles with negligible extension.
40
What is the ideal gas law equation?
PV = nRT pressure - Pascals [x1000 from kPascals] volume - m3 [cm - x 10^-6] n - moles r - give t - +273
41
In what conditions do real gases act as ideal gases?
lower pressure higher temperatures
42
Why do we study gases?
Easiest possible system – molecules in gas are completely disconnected, free to move and occupy the volume of vessel
43
What laws combine for the ideal gas law?
Boyle's Law, Charle's Law, Avogadro's Law
44
What is Boyles Law?
At a constant temperature, as pressure increases, volume decreases
45
What is Avogadro's law?
Equal volumes of gases at the same temperature and pressure contain equal numbers of molecules
46
According to Avogadro's law, what conditions does an ideal gas need?
According to Avogadro’s law an ideal gas always needs the same volume at a given temperature and pressure.
47
What is the universal value of stp?
1 atm (pressure) and 0 degrees C.
48
In STP how much volume does 1 mole of gas occupy?
In STP, 1 mole of gas will occupy a volume of up to 22.4 L.
49
What are deviations from ideality described as?
the COMPRESSION FACTOR, Z (sometimes called the compressibility)
50
For ideal gases, what does Z equal?
For ideal gases Z = 1,
51
Why does deviation of gas from ideal gases arise?
Attractive forces vary with nature of gas At High Pressures repelling forces dominate
52
What does the 1st law of thermodynamics state about the energy in an isolated system?
The quantity of energy in a fully closed system remains constant. (The internal energy of an isolated system is constant)
53
What is work?
is a transfer of energy that utilizes or causes uniform motion of atoms in the surroundings (how we transfer energy) (motion against an opposing force)
54
When is work done?
Work is done when a force moves.
55
What is the equation for work?
W = Force x Distance
56
What is the most common form of work (and often the only one which we have to consider?)
Work done against (opposing motion) the surrounding pressure when volume, V increases.
57
What is energy?
A measure of the capacity of a system to do work
58
What is heat?
It is the means by which energy is transferred from a hotter body to a cooler one in order to equalize their temperatures.
59
What is internal energy?
(energy contained within the system) Total of the kinetic energy and the potential energy
60
What does kinetic energy refer to in a system?
due to the motion of molecules (translational, rotational, vibrational)
61
What does potential energy refer to in a system?
associated with the atoms within molecules or crystals
62
What does the internal energy of a system (U) depend upon?
Translational kinetic energy Molecular rotation Bond vibration Intermolecular attractions Chemical bonds Electrons
63
Speaking of differentials, which are exact and inexact differentials (path, state functions)
State functions are exact differentials (a difference between amounts of things) Path functions are inexact differentials
64
State function differential formula
Change in U = dU = Uf- Ui
65
Path functions differential formula
q = dq is not equal to qf - qi
66
Heat is a signed quantity, if heat is a) absorbed b) given out what are the signs?
+q: heat is absorbed by the system (an endothermic process) -q: heat is given out by the system (an exothermic process)
67
When do we define work as being positive?
We define work as being positive when the system does work on the surroundings (by the system) (energy leaves the system) (when a gas is compressed by an external force the work is positive (+w))
68
When do we define work as being negative?
If work is done on the system (energy added to the system), (a gas expands by pushing against an external force the work is negative (-w) the work is negative.
69
An ideal piston (one with no mass that moves without friction) with an area A, which contains a gas at a pressure pint and where the external pressure is pext. Under conditions where pint\>pext, what does the piston do?
It will move out (by a distance dx) and in doing so does work, dw’ (=-dw) against the external pressure. Because internal pressure is bigger, it is work done BY the gas.
70
The work done by the gas equation
dx = p(ext) dV
71
The work done on the gas equation
-dw' = -p(ext) dV
72
If pext= 0 (in vacuum), what is the value of work?
0, clearly no work is done i.e. no work is done by a gas expanding into a vacuum
73
If p(ext) is constant then we can calculate the work done by the piston when it expands from an initial volume, Vi to a final volume Vf by what equation?
dV = p(ext) (Vf -Vi)
74
What are the three “forms” of energy?
heat (q), work (w) and internal energy (U).
75
How can the first law be expressed mathematically?
If we take a system from state A to state B, then there is a definite change in the internal energy, ΔU = UB - UA Where, UB and UA are the internal energies of the system in the two states. ALSO ΔU = q + w (dU = dq + dw)
76
If we now make the simplification that only PV work is done on the system, what is the equation?
dU = dq - d( pV)
77
What are the three possible expansion scenarios for the internal energy change of a system where only PV work
- A change where the surroundings are a perfect vacuum (p=0) - A change at constant volume (“isochoric”), (dV=0) - A change where the surroundings are at constant pressure (dp=0)
78
What is the equation where there is a change in which the surroundings are a perfect vacuum, where no work is done (p=0):
dU = dq (ΔU = q + w) NO WORK IS DONE BECAUSE PDV = W AND ITS 0
79
What is the equation where there is a change at constant volume (“isochoric”) where again no work is done (dV=0) :
dU = dq (ΔU = q + w)
80
What is the equatiion for a change where the surroundings are at constant pressure (“isobaric”):
dU = dq - pdV (ΔU = q + w)
81
In what conditions would the max amount of work be done?
The external pressure should be as high as possible in order to maximise the work (However if the external pressure exceeds the internal pressure then the gas is compressed)
82
C3H8(g) + 5 O2(g) → 3CO2(g) + 4H2O(l) at 298 K & 1 atm (1 atm = 101325 Pa) What is the work done by the system?
For an ideal gas; pV = nRT (p = p[ext]) 1) V= nRT/p 6 moles of gas: Vi = (6 × 8.314 × 298)/ 101325 = 0.1467 m3 3 moles of gas: Vf = (3 × 8.314 × 298)/ 101325 = 0.0734 m3 work done = -pex x (Vf – Vi) = -101325 (0.0734 – 0.1467) = +7432 J
83
Zn(s) + 2HCl(aq) ZnCl2 (aq) + H2 (g) How much work is done when we throw 50 g of Zinc in HCl in a open beaker?
1 ) pV = n× Ri×T DV =Vf -Vi » Vf = nRT/ p[ex] 2) dw = -p[ex]dV WE KNOW VOLUME = nRT/p[ex] - SUBSTITUTE 3) dw = -p[ex] nRT/pex = -nRT 4) FIND MOLES (50/65.4) = .76 5) -.76 x 8.3145 x ´298K = -1.9kJ
84
When we supply heat to an object its temperature rises and the relationship between the heat supplied, q and the temperature rise, dT is what?
dq = CdT C is the “heat capacity” (J/K)
85
What is the “molar heat capacity”?
The amount of heat required to raise one mole of substance through one degree (units JK-1Mol-1).
86
What is heat capacity?
It is the amount of heat needed to raise the temperature of a certain mass 1 degree Celsius.
87
If the heat capacity of water is 4.19 kJ/(kg K), what does this mean?
That means it takes 4.2 joules of heat energy to raise one gram of water one degree Celsius.
88
How much energy is needed to heat 1.0 kg of water from 0oC to 100oC when the specific heat of water is 4.19 kJ/kg.K?
Q = (4.19 kJ/kg.K) (1.0 kg) ((100 oC) - (0 oC)) = 419 (kJ)
89
What does the heat capactity of a substance depend on?
- the pressure (not in the case of ideal gases), - the volume - the process-path OF THE SYSTEM
90
Two important experimental conditions are to set the volume or the pressure constant, what are they?
a) constant volume - isochore heat-capacity (Cv ,cv ,Cmv) b) constant pressure - isobar heat-capacity (Cp ,cp ,Cmp)
91
What is the definition of the molar heat capacity at constant volume in equation form
𝐶𝑉 = (𝑑𝑈/𝑑𝑇)V
92
How to measure dq – change in heat of a system? The calorimeter directly gives dU
dU = qV U = q+w (dU = dqv + dw) dqV - p[ex]dV =0 (V is constant)
93
Is the heat absorbed during a reaction in a bomb calorimeter a state or path function?
a quantity that depends only on the initial and final states because U, P and V are all state functions.
94
If heat is absorbed by a gas under constant pressure, what will happen to the heat?
-some of the heat will increase the internal energy -some of the heat will appear as work of expansion
95
What is the equation for Enthalpy?
H = U + pV (Internal energy + work system does = enthalpy) dH = q + Vdp (du = q) dH = qp (under constant pressure conditions)
96
If something has a high heat capacity, what does it mean?
Requires a lot of heat to increase temperature
97
Give an example of a substance with a high heat capacity
Water
98
What is the equation for heat change? (q)
heat change = mass × specific heat capacity × temperature change q = m × cg × ΔT
99
How to calculate enthalpy change per mole of a substance
-q/1000 ÷ n /1000 = to get KJ
100
When calculating the enthaply change, what must you just calculate?
Heat change and / 1000 = to get in KJ
101
What is the equation for the work the system does?
w = pV
102
What is the enthalpy change?
Heat absorbed or given out in a reaction occurring at constant pressure. (heat absorbed (taken in) -ENDOTHERMIC: ΔH positive heat evolved (given out) - EXOTHERMIC: ΔH negative)
103
What is a standard state?
when a substance is in its most stable form at 298 K and 1 atmosphere pressure.
104
Elements in their standard states are assigned what value of enthalpy
Zero
105
What is the standard enthalpy change?
ΔHo298: the enthalpy when reactants in their standard states are converted to products in their standard states.
106
What is the standard enthalpy of formation?
ΔHof: use elements in their standard states to form 1 mole of the substance in its standard state.
107
If we know the enthalpy of a substance at one temperature, H(T1) but we want to know its enthalpy at another temperature, H(T2), what equation do we use?
(h2 - h1)p = cp (T2 - T1)p Kirchhoff’s Equation
108
What is a phase?
A phase is a specific state of matter that is uniform throughout in composition and physical state.
109
Why is the term phase more specific than ‘state of matter’?
because a substance may exist in more than one solid form, each one of which is a solid phase
110
Is the process of vaporization of a liquid, such as the conversion of liquid water to water vapor exothermic or endothermic and why?
An endothermic process (ΔH \> 0) because heating is required to bring about the change. (At a molecular level, molecules are being driven apart from each other and this process requires energy.)
111
How does the body use the endothermic character of vaporisation of water to maintain the temperature at about 37°C?
because the evaporation of perspiration requires energy and withdraws it from the skin. (ENDOTHERMS)
112
What is the standard enthalpy of vaporization of a liquid?
The energy that must be supplied as heat at constant pressure per mole of molecules that are vaporized under standard conditions (1 bar) ( ΔVAPH°.)
113
If the standard enthalpy of vaportization of water is 44KJ/Mol
44 kJ of heat is required to vaporize 1 mol H2O(l) at 1 bar and 25°C
114
If the stoichiometric coefficients in the chemical equation are multiplied through by 2, then the thermochemical equation would be written: 2 H2O(l) → 2 H2O(g) If the enthalpy of vaporization of water is 44 (for 1 mole) what is it for 2 moles?
ΔH° = +88 kJ This equation means that 88 kJ of heat is required to vaporize 2 mol H2O(l) at 1 bar and at 298.15 K.
115
What does the difference in values mean? • the value for water is 44 kJ mol−1 • the value for methane, CH4 is only 8 kJ mol−1
Even allowing for the fact that vaporization is taking place at different temperatures, the difference between the enthalpies of vaporization means that water molecules are held together in the bulk liquid much more tightly than methane molecules.
116
What is the standard enthalpy of fusion, ΔfusH°?
The change in molar enthalpy that accompanies fusion under standard conditions (pure solid at 1 bar)
117
The enthalpy of fusion of water is much less than its enthalpy of vaporization, why is this so?
In melting the molecules are merely loosened without separating completely whereas in vaporization the molecules become completely separated from each other.
118
What is the reverse of vaporization
condensation
119
What is the reverse of fusion (melting)
freezing
120
Is condensation and freezing negative or positive
Negative, because heat is released
121
What is Bond Enthalpy?
the energy that is needed to break/form a particular bond in a gaseous compound.
122
The total standard enthalpy change for the atomization (the complete dissociation into atoms) of water: H2O(g) → 2 H(g) + O(g) ΔH°= +927 kJ Why is it not twice the O–H bond enthalpy in H2O, even though two O–H bonds are dissociated?
There are in fact two different dissociation steps: 1) First an O–H bond is broken in an H2O molecule: H2O(g) → HO(g) + H(g) ΔH° = +492 kJ 2) Second, the O–H bond is broken: HO(g) → H(g) + O(g) ΔH°= +428 kJ The sum of the two steps is the atomization of the molecule.
123
How to calculate combustion when the energy released at constant volume as heat by the combustion of the amino acid glycine is -969.6 kJ mol-1 at 298.15 K
ΔU = -969.6 kJ mol-1 From the chemical equation: NH2CH2COOH(s) + 9/4 O2(g) → 2 CO2(g) + 5/2 H2O(l) + 1/2 N2(g) Δν gas = (2 + 1/2) -9/4 = 1/4 Therefore, ΔH = ΔU + 1/4RT = −969.6 kJ mol−1 + ¼(8.3145×10−3 kJ K−1 mol−1) (298.15 K) = −969.6 kJ mol−1 + 0.62 kJ mol−1 = −969.0 kJ mol−1
124
What is hess's law?
This implies that the Enthalpy change for a process is the same no matter what pathway we take in going from the initial to the final state. (becasue it is a state function)
125
If something is Delta, is it a state or path function?
State function
126
What is the standard reaction enthalpy?
Δ rH, is the difference between the standard molar enthalpies of the reactants and the products, with each term weighted by the stoichiometric coefficient, n (nu), in the chemical equation
127
What is a reversible change in thermodynamics?
One that can be reversed by an infinitesimal modification of a variable
128
If the external and internal pressures are equal in a system, what does it mean?
The system is in equilibrium.
129
What occurs to the gas inside the piston if the external pressure is reduced infinitesimally, and how can it be reversed?
The gas inside the piston expands slightly which can be reversed by an infinitesimal increase in the external pressure.
130
If the external pressure is reduced infinitesimally however there is then an infinitesimal increase in the external pressure, what is it known as?
Reversible thermodynamic change.
131
If the external pressure is significantly lower than the internal pressure then an infinitesimal increase in the external pressure does not reverse the expansion, what does this tell us about the system?
The system is not in equilibrium and the process of expansion is irreversible
132
What are the properties of reversible processes?
- Are infinitely slow - Are at equilibrium - Do maximum work
133
What are the properties of irreversible processes?
- Go at finite rate - Are not at equilibrium - Do less than the maximum work
134
What are the properties of reversible processes?
- Are infinitely slow - Are at equilibrium - Do maximum work
135
What is entropy?
It is a measure of how organized or disorganized energy is in a system of atoms or molecules. A change depends on how much energy is transferred as heat which stimulates disordered motion in the surroundings.
136
Does work influence entropy?
Work stimulates a uniform motion of the atoms in the surroundings and hence does not change the entropy of the system)
137
What is entropy in an equation? (entropy change)
dS = dq (rev) /T q(rev) = heat change for a reversible path between the two states of the system
138
If matter and energy are distributed in a disordered way (as in a gas for example) is entropy high or low?
High
139
If matter and energy are distributed in an ordered way (as in acrystal for example) is entropy high or low?
Low
140
What is a spontaneous change?
Don’t need to provide external stimuli – on their own
141
What does the second law of thermodynamics say about the spontaneous process?
For a reaction to be Spontaneous, there must be an increase in entropy (disorder) of universe (system + surroundings)
142
What is the equation for the entropy change of the universe
DS(universe)= DS(sysem) + DS(surroundings)
143
The simplest place to apply the equation is in what kind of system?
an isolated system because here dq = 0 so that the three possibilities are: ΔS \> 0 spontaneous (irreversible) process ΔS = 0 reversible process ΔS \< 0 process not feasible
144
if ΔS \> 0, will the reaction be a)spontaneous b)reversible c)not feasible
a) spontaneous (irrevisible process)
145
if ΔS = 0 , will the reaction be a)spontaneous b)reversible c)not feasible
b) reversible
146
If ΔS \< 0, will the reaction be a)spontaneous b)reversible c)not feasible
c) not feasible
147
When a natural (spontaneous) process occurs in an isolated system, is there an increase or decrease in entropy?
There is an increase in entropy – the entropy increases to a maximum at equilibrium
148
At constant pressure the heat change is equal to what?
to the Enthalpy change and hence (dqrev=dH):
149
What does entropy equal?
dS = dq/t dH/t CpdT/T CpIn T2/T1
150
What is the entropy change for the phase transition is given as and why?
△S = △H/T At the temperature of the phase change the two phases are in equilibrium and therefore any transfer of heat between the system and its surroundings is reversible and hence at constant pressure
151
When is the entropy at 0?
If the substance is perfectly crystalline, with every atom in a well defined location, then there is no spatial disorder. At T = 0, all the motion of the atoms is eliminated and there is no thermal disorder, we can therefore say that at T = 0, the entropy is zero.
152
What does the third law state about entropy?
The Third Law: The entropies of all perfectly crystalline substances are the same at T = 0
153
How do you go about phase transitions and entropy
To add the entropy of transition (ΔHtrans/Ttrans) for each phase transition, between T=0 and the temperature of interest. ST=0, Solid -\> ST=FUS, Solid -\> ST=FUS, Liquid CdT/T + △H/T (everytime you add until you get to end and there is no + △H/T.
154
What is Gibbs Free Energy?
It is a way of combining the entropy and enthalpy of a system into one value. It describes the energy available to do work and thus can be used to determine the spontaneity of a reaction.
155
What is the equation for Gibbs Function (for constant pressure processes)
△G = △H -T△S H= change in enthalpy T = Temperature in kelvin S = Change in Entropy
156
What does Gibbs function tell us?
It tells us nothing about the rate of the reaction or the mechanism it may follow, but it is the criterion for the feasibility of a chemical reaction
157
If △G \< 0, what will the reaction be?
Spontaneous (irreversible) process
158
If △G = 0, what will the reaction be?
At equilibrium (you can always go back = reversible)
159
If △G \> 0, what will the reaction be?
Process is not feasible
160
If △H \> 0 and △S is \> 0, the reaction is possible in what conditions?
At suffiently high temperatures
161
If △H \> 0 and △S is \< 0, when is the reaction possible?
Never
162
If △H \< 0 and △S is \> 0, when is the reaction possible?
Always
163
If △H \< 0 and △S is \< 0, when is the reaction possible?
At sufficiently low T
164
What is △G for melting ice at -10 degrees
1) First we need to calculate our temperature to kelvins: T=-10∘C+273=263K 2)The enthalpy of fusion and entropy of fusion for water have the following values: Δ fus H=6.01 mol/KJ Δ fus S=22.0 mol/KJ 3) ΔG =ΔH−TΔS 6.01 −(263K)(0.022) 6.01 - 5.79 = .22 Thus we can see at -10 degrees, the Gibbs free energy change G is positive for the melting of water. Therefore, we would predict that the reaction is not spontaneous at -10. In fact, we would predict that the reverse reaction should be spontaneous: we would expect to see a puddle of water turn into ice.
165
What happens at sufficiently high temperatures (when T reaches boiling point)
at sufficiently high T (i.e., when T reaches the boiling point) TΔS becomes equal to ΔH ΔHvaporization = Tboiling point x ΔS vaporisation so that ΔG is zero and there is equilibrium – any further heating and the liquid boils. This is why there are definite boiling points for liquids e.g., water boils at 373 K
166
What is the first master equation?
dU = TdS - pdV (as it combines the 1st and 2nd laws)
167
What is the second master equation?
dH = TdS + VdP
168
What is the third master equation?
dG = Vdp - SdT
169
If the temperature is constant dT = 0, what is the third master equation?
dG = Vdp (Because all volumes are positive, the Gibbs energy increases when the pressure increases!)
170
If the pressure is constant VdP=0 what is the third master equation?
dG = -SdT
171
What is the relationship between Gibbs Free energy and temperature (and entropy)
for a given change of temperature, the change in Gibbs energy is proportional to the entropy.
172
What does an increase in temperature result in?
a decrease in G (because entropy is positive)
173
Why does the Gibbs energy falls more steeply with temperature for a gas than for a condensed phase, for a given substance,
because the entropy of the gas phase is greater than that for a condensed phase.
174
Why in a diagram is the slope is less steep for a solid
The entropy of the liquid phase of a substance is greater than that of its solid phase
175
What is a phase diagram?
The phase diagram of a substance is a map showing the conditions of temperature and pressure at which its various phases are thermodynamically most stable.
176
How to read a phase diagram : DETERMINE STATE OF SUBSTANCE
Always have two variables : Pressure on Y axis and Temperature on X axis (can change) Each line : corresponds to an equilibrium existing between two phases Between two lines : in what phase it will behave Point on line – at that combination of temp and pressure, you will have two phases existing at the same time Where all lines meet = triple point, it exists only at one combination of pressure and temperature
177
What is vapor pressure?
The pressure of a vapor that is in equilibrium with its condensed phase. It is an indication of a liquid's evaporation rate. It relates to the tendency of particles to escape from the liquid.
178
A substance with a high vapour pressure at normal temperatures is often referred to as what?
volatile
179
Why does vapour pressure increase with temperature?
Vapor pressure increases with temperature because, as the temperature is raised, more molecules have sufficient energy to leave their neighbours in the liquid and escape…
180
The relation between dT and dp that ensures that in either case the two phases remain in equilibrium is given by what equation?
Clapeyron equation, for the slope of the phase boundary at any temperature. dP/dT = ΔS/ΔV
181
If you have a case where a Solid ⇄ Gas and Liquid ⇄ Gas, (for boundaries with gas), what can you assume and whats the equation? (Clasius, Claperon)
In these cases we can safely assume that the molar volume of gas is much greater than that of either the liquid or solid 1/p dp/dT = ΔH / RT^2
182
Solid ⇄ Liquid, whats the phase equation?
p2-p1 = ΔHm/ΔVIn T2 / T1
183
In an Open vessel, at a certain temperature what is the vapor pressure?
becomes equal to the external pressure (The temperature at which the vapor pressure of a liquid is equal to the external pressure is called the boiling temperature)
184
What is the boiling temperature of a liquid?
The temperature at which the vapor pressure of a liquid is equal to the external pressure
185
What occurs at boiling temperature (when the vapour pressure = external pressure)
At this temperature, the vapor can expand indefinitely. Moreover, because there is no constraint on expansion, bubbles of vapor can form throughout the body of the liquid, the condition known as boiling.
186
When the external pressure is 1 atm, the boiling temperature is called what?
the normal boiling point, Tb.
187
How can we predict the normal boiling temperatures of liquids?
We can predict the normal boiling point of a liquid by noting the temperature on the phase diagram at which its vapor pressure is 1 atm
188
When a liquid is heated in a sealed container, what occurs?
the density of the vapor phase increases the density of the liquid phase decreases There comes a stage at which the two densities are equal and the interface between the two fluids disappears and this occurs at a critical temperature
189
For the stage at which the two densities are equal and the interface between the two fluids disappears, what are the conditions needed?
The container needs to be strong: the critical temperature of water is at 373°C and the vapor pressure is then 218 atm!
190
Proteins and biological membranes can exist in ordered structures stabilized by a variety of molecular interactions, such as hydrogen bonds and hydrophobic interactions. However, when certain conditions are changed, the helical and sheet structures of a polypeptide chain may collapse into a random coil and the hydrocarbon chains in the interior of bilayer membranes may become more or less flexible. How is this explained?
These structural changes may be regarded as phase transitions in which molecular interactions in compact phases are disrupted at characteristic transition temperatures to yield phases in which the atoms can move more randomly
191
What is the phase transitions of biological membranes?
All lipid bilayers undergo a transition from a state of high to low chain mobility at a temperature that depends on the structure of the lipid.
192
At physiological temperature, how does the bilayer exist?
At physiological temperature, the bilayer exists as a liquid crystal, in which some order exists but the chains writhe. There is sufficient energy available at normal temperatures for limited bond rotation to occur and the flexible chains to writhe around. However, the membrane is still highly organized in the sense that the bilayer structure does not come apart
193
At lower temperatures, how does the bilayer exist?
At lower temperatures, the amplitudes of the writhing motion decrease until a specific temperature is reached at which motion is largely frozen. The bilayer is said to exist as a gel.
194
Unfolding of regular secondary protein structure causes: a) Large decrease in the entropy of the protein b) Little increase in the entropy of protein c) No change in the entropy of the protein d) Large increase in the entropy of the protein
d) Large increase in the entropy of the protein, protein is unfolding so it becomes more disordered. Entropy = DISORDER.
195
What does first law of thermodynamics state? a) Energy can neither be destroyed nor created b) Energy cannot be 100 percent efficiently transformed from one type to another c) All living organisms are composed of cells d) Input of heat energy increases the rate of movement of atoms and molecules
a) Energy can neither be destroyed nor created
196
If enthalpy change for a reaction is zero, then ∆G° equals to: a) -T∆S° b) T∆S° c) -∆H° d) lnkeq
a) -T∆S° △G = △H -T△S
197
If ∆G‘° of the reaction A → B is -40kJ/mol under standard conditions then the reaction: a) Will never reach equilibrium b) Will not occur spontaneously c) Will proceed at a rapid rate d) Will proceed from left to right spontaneously
THERMO DYNAMICS NEVER TELL YOU HOW RAPID IT MOVES = d) Will proceed from left to right spontaneously
198
Which of the following statements is false? a) The reaction tends to go in the forward direction if ∆G is large and positive b) The reaction tends to move in the backward direction if ∆G is large and negative c) The system is at equilibrium if ∆G = 0 d) The reaction tends to move in the backward direction if ∆G is large and positive
a)The reaction tends to go in the forward direction if ∆G is large and positive (we move backwards) b)The reaction tends to move in the backward direction if ∆G is large and negative (we move forwards)
199
Which of the following is true for a closed system? a) mass entering = mass leaving b) mass does not enter or leave the system c) mass entering can be more or less than the mass leaving d) none of the mentioned
b) mass does not enter or leave the system
200
A piston cylinder contains air at 600 kPa, 290 K and a volume of 0.01m3 . A constant pressure process gives 54 kJ of work out. Find the final volume of the air. a) 0.05 m3 b) 0.01 m3 c) 0.10 m3 d) 0.15 m3
c) 0.10 m3 ChangeV = W/P = 54/600 = .09m3 (final volume, initial and final) V2 = V1 + ChangeV = .01 + .09 = .1m3
201
Which of the following represents an increase in entropy? A) freezing of water (dec – putting structure) B) boiling of water (inc) C) crystallization of salt from a supersaturated solution (dec) D) the reaction 2NO(g) → N2O2 (g) (dec -\> 2 moles to 1 mole) E) the reaction 2H2 (g) + O2 (g) → 2H2O(g) (dec, 3 moles – 2moles)
B) boiling of water (inc)
202
Calculate the standard entropy change for the following reaction, Cu(s) + ½ O2(g) → CuO(s), given that S˚[Cu(s)] = 33.15 J/K•mol S˚[O2(g)] = 205.14 J/K•mol S˚[CuO(s)] = 42.63 J/K•mol A) 195.66 J/K B) 93.09 J/K C) -45.28 J/K D) -93.09 J/K E) 195.66 J/K
D) -93.09 J/K = STATE FUNCTION. Delta S of products – Delta S of all reactants 42.63 – (205.14/2 + 33.15) = -93.09 (1.2 mole = 205.14 /2)
203
In which of the following reactions do you expect to have a decrease in entropy? A) Fe(s) → Fe(l) (inc) B) Fe(s) + S(s) → FeS(s) (inc) C) 2 Fe(s) + 3/2 O2 (g) → Fe2O3 (s) (dec) D) HF(l) → HF(g) (inc) E) 2 H2O2 (l) → 2 H2O(l) + O2 (g) (inc)
c) It generates more moles of gas
204
Which of the following statements is false? A) The change in entropy in a system depends on the initial and final states of the system and the path taken from one state to the other. (false) B) Any irreversible process results in an overall increase in entropy. (true) C) The total entropy of the universe increases in any spontaneous process. (true) D) Entropy increases with the number of microstates of the system. (how many faces are in the system)
A) The change in entropy in a system depends on the initial and final states of the system and the path taken from one state to the other. (false)
205
Only those processes are possible in nature which would give an entropy ____ for the system and the surroundings together. a) decrease b) increase c) remains same d) none of the mentioned
b) increase
206
When heat is imparted to a system, (give heat) a) the disorderly motion of molecules increases b) the entropy of the system increases c) both of the mentioned d) none of the mentioned
c) both of the mentioned
207
The heat capacity of the substance is defined as the amount of heat required to raise a unit mass of the substance through a unit rise in temperature. a) True b) False
a) True (definition of heat capacity)
208
The enthalpy of a substance(denoted by h), is defined as a) H= U-PV b) H= U+PV c) H= -U+PV d) h= -U-PV
b) H= U+PV
209
In a constant volume process, internal energy change is equal to a) heat transferred b) work done c) zero d) none of the mentioned
a) heat transferred
210
Heat transferred at constant pressure _____ the enthalpy of a system. a) decreases b) increases c) first decreases then increases d) first increases then decreases
b) increases
211
When is a system in equilibrium?
a system is at equilibrium when the chemical potential of each substance has the same value in every phase in which it occurs
212
What is the chemical potential?
a thermodynamic function expressing the ability of an uncharged atom or molecule in a chemical system to perform physical work.
213
Is the molar Gibbs energy of a pure substance a) the same in all the phases at equilibrium b)different in all phases at equilibrium
a) the same in all the phases at equilibrium
214
Are all chemical reactions in equilibrium?
Yes, every chemical reaction can theoretically be in equilibrium. all reactions are reversible, at least to a (perhaps vanishingly small) degree. To say otherwise would violate microscopic reversibility
215
What are the two solutions we can have
Ideal Solutions and Non-Ideal solutions
216
Does the chemical potential of a species a)increase b)decrease with concentration…
a) increase
217
What is Ranoult's law?
Raoult’s law provides a means for defining ideal behaviour. The partial vapor pressure of a substance in a liquid mixture is proportional to its mole fraction in the mixture and its vapor pressure when pure: pA = xApA\* pB = XBpB\* pA \* is the vapor pressure of the pure substance J - a substance in general, A - solvent, B - solute
218
How do we obtain the total vapor pressure
pA + pB = xApA\* + xBpB\*
219
How does the vapor pressure increase?
The vapor pressure of a liquid varies with its temperature. By increasing the temperature
220
What is an ideal solution?
A hypothetical solution of a solute B in a solvent A that obeys Raoult’s law : The partial vapor pressure of a substance in a liquid mixture is proportional to its mole fraction in the mixture and its vapor pressure when pure
221
When is Ranoult's Law most reliable?
When the components of a mixture have similar molecular shapes and are held together in the liquid by similar types and strengths of intermolecular forces
222
If we plot pressure over mole fraction :
Intersection of two axis = 0 , which means 0 moles, no pressure If you have .1 molar fraction of A, then you have .9 fraction of B (if molar fraction = .1, 10%) 0– \>Pa = only a 1-\> Pb = only b Line which joins Pb to PA = total pressure Want to find total pressure from percentage, go up until hits total pressure line
223
When might Ranoult's law be a very good approximation and why?
A good approximation for the solvent if the solution is very diluted - have such a tiny amount that solute molecules would not interact = solution behaves like an ideal solution
224
What is Henry's Law?
The mass of a dissolved gas in a given volume of solvent at equilibrium is proportional to the partial pressure of the gas. pB = a = γXB pB = activity
225
What is an ideal-dilute solution?
is one in which the solvent obeys Raoult’s Law but the solute obeys Henry’s Law!
226
How does Gibbs energy change when two components are mixed?
ΔG(mix) = G(mixed) - G(seperated) Ideal solutions= naRTInxA + nbRTinxB Non-ideal solutions = ΔG (mix)/n total = RT (xAINxA +xBInxB)
227
Is the mixing of ideal solutions spontaneous?
As mole fractions are never greater than 1 the logarithmic terms are negative, and hence ΔGmix \< 0. Therefore the mixing of ideal solutions is spontaneous in all proportions.
228
Let’s look at the isomerism of glucose-6-phosphate (1, G6P) to fructose-6-phosphate (2, F6P), which is an early step in the anaerobic breakdown of glucose: G6P(aq) ⇄ F6P(aq)
Suppose that in a short interval while the reaction is in progress, the amount of G6P changes infinitesimally by dn. - As a result, the contribution of G6P to the total Gibbs energy of the system changes by -µG6Pdn,(µG6P is the chemical potential - the partial molar Gibbs energy - of G6P in the reaction mixture. - In the same time interval, the amount of F6P changes by +dn, so its contribution to the total Gibbs energy changes by +µF6Pdn (where µF6P is the chemical potential of F6P). The change in Gibbs energy of the system is: dG = µF6Pdn - µG6Pdn On dividing through by dn, we obtain the reaction Gibbs energy, ΔrG = µf6P - µG6P
229
binding of O2(g) to the protein hemoglobin, Hb, in blood: Hb(aq) + 4 O2(g) ⇄ Hb(O2)4(aq)
When the amount of Hb changes by dn, from the reaction stoichiometry the change in the amount of O2 is 4dn and the change in the amount of Hb(O2)4 is +dn. The overall change in the Gibbs energy of the mixture is: dG = µHb(O2)4 × dn - µHb× dn - µO2 × 4dn = (µHb(O2)4 - µHb - 4µO2)dn µJ are the chemical potentials of the species in the reaction mixture! The reaction Gibbs energy is:µHb(O2)4 - µHb - 4µO2 - each chemical potential is multiplied by the corresponding stoichiometric coefficient reactants are subtracted from products ΔrG = (cµC + dµD) - (aµA + bµB)
230
How will an ideal solution behave?
It will behave more like a gas – no interactions between molecules
231
What is xA = molar fraction
Total moles of a / a+ B in mixture)
232
The chemical potential of a solution is the same as?
Gibbs Energy
233
What is the ''chemical potential?''
Expresses the ability of an uncharged atom or molecule in a chemical system to perform physical work.
234
Why is the graph of a real mixture pressure vs molar fraction different to the one of ideal mixtures?
Real mixture, has deviations and bumpy curves = There are going to be interactions between molecules
235
What is the activity of a substance?
activity = concentration x gamma factor (number which changes in every solution you have, which tells you how strong the bonds that are formed in the new solution) PB = activity
236
When you mix A + B, what value do you want for delta g and why?
You want delta g as negative as possible, its not stable until it reaches perfect mixture where delta g is minimized as much as it can be – therefore it is at equilibrium (at .5 mole fraction (50/50, because it can go back and forth)
237
What is partial pressure?
the pressure that would be exerted by one of the gases in a mixture if it occupied the same volume on its own.
238
Which statement about these two thermodynamic processes is correct? Process one : metal box and ice @ 0 degrees -\> metal box and ice @ 0 degrees Process two : metal box at 70 degrees and ice at 0 degrees -\> metal box at 40 degrees and ice at 40 degrees A. Both are reversible. B. Both are irreversible. C. Process one is reversible and process two is irreversible. D. Process one is irreversible and process two is reversible. E. Not enough information is given to decide.
C. Process one is reversible and process two is irreversible.
(If you increase temperature, you increase entropy)
(An irreversible process increases the entropy of the universe)
239
An ideal gas is taken around the cycle shown in this p-V diagram, from a to b to c and back to a. Process b c is isothermal. Which of the processes in this cycle could be reversible?
A. a -b
B. b - c
C. c - a
D. two or more of A, B, and C
E. none of A, B, or C
B. b - c
240
An ideal gas is taken around the cycle shown in this p-V diagram, from a to c to b and back to a. Process c b is adiabatic. Which of the processes in this cycle could be reversible?
A. a - c
B. c - b
C. b - a
D. two or more of A, B, and C
E. none of A, B, or C
B. c - b
241
A copper pot at room temperature is filled with room-temperature water. Imagine a process whereby the water spontaneously freezes and the pot becomes hot. Why is such a process impossible?
A. It violates the first law of thermodynamics.
B. It violates the second law of thermodynamics.
C. It violates both the first and second laws of thermodynamics.
D. It violates the law of conservation of energy.
E. none of the above
B. It violates the second law of thermodynamics.
The second law of thermodynamics states that the total entropy of an isolated system can never decrease over time
242
What is an isothermal process?
An isothermal process is a change of a system, in which the temperature remains constant: ΔT = 0
243
An ideal gas is taken around the cycle shown in this p-V diagram, from a to b to c and back to a. Process b-c is isothermal. What can you conclude about the net entropy change of the gas during the cycle?
A. It is positive.
B. It is negative.
C. It is zero.
D. Two of A, B, and C are possible.
E. All three of A, B, and C are possible
C. It is zero.
(isothemal, temp is constant, entropy is 0) - third law of thermodynamics
244
An ideal gas is taken around the cycle shown in this p-V diagram, from a to b to c and back to a. Process b - c is isothermal. What can you conclude about the net entropy change of the gas and its environment during the cycle?
A. It is positive.
B. It is negative.
C. It is zero.
D. Two of A, B, and C are possible.
E. All three of A, B, and C are possible.
A. It is positive.
245
How to caluclate partial pressure?
1) Find total pressure = PV=NRT
2) moles/total moles x total pressure for each one
246
What are intensive properties and give examples
Physical properties that do not depend on the amount of matter eg, melting point, boiling point, density
247
Give examples of extensive properties
Volume, Mass, Energy
248
How to calculate bond enthalpy
Co (g) + H20 -\> Co2 (g) + H2
C-O in carbon monoxide+1077
C-O in carbon dioxide+805
O-H+464
H-H+436
Enthalpy change = Products - Reactants
(805 (2) + 436) - (1077 + 464 (2) )
-41 KJ
249
What does a positive and negative change in bond enthalpy mean?
A positive change in enthalpy is required to break a bond, while a negative change in enthalpy is accompanied by the formation of a bond. In other words, breaking a bond is an endothermic process, while the formation of bonds is exothermic
250
Analyse the table below and answer these questions
1. When can a reaction occur in a) all temperatures and b) not at all (not feasible)
2. When can a reaction occur in a) high temperatures and b) low temperatures
3. What does it mean if gibbs free energy is positive?
1
a) S is positive and H negative
b) S is negative and H positive
2
a) S positive, H positive
b) S negative, H negative
3. Spontaneous!
251
What changes would you expect to observe in the colligative properties of water on addition of a few mg of NaCl?
(a) an increase in boiling point and a decrease in freezing point
(b) a decrease in vapour pressure and a decrease in boiling point.
(c) an increase in boiling point and an increase in vapour pressure
(d) an increase in freezing point and an increase in vapour pressure
(e) none of the above
(a) an increase in boiling point and a decrease in freezing point
vapour pressure decreases
It occurs anytime you add non-volatile solute (e.g., salt) to a solvent (e.g., water).
252
The value of the Gibbs free energy change for a reaction at 35 oC is -27.9 kJ Mol-1 and the value of the change in enthalpy is –36.29 kJ Mol-1 . The value of the change in entropy for the reaction at this temperature is:
(a) -239.7 J Mol-1K-1
(b) -27.2 kJ Mol-1K-1
(c) 239.7 J Mol-1K-1
(d) cannot be calculated from the data supplied.
(e) -27.2 J Mol-1K-1
G = H-TS
G-H = -TS
8. 39/-T = S
8. 39/ -308 = S (KJ)
- .0272KJ x 1000
(e) -27.2 J Mol-1K-1
