Download Thermochemistry (Energy Relationships in Chemical Reactions

Thermochemistry (Energy Relationships in Chemical Reactions)
Some forms of energy
Kinetic (thermal) energy:
energy associated with motion
Every chemical reaction obeys two fundamental laws: the law of
conservation of mass and the law of conservation of energy
Potential energy:
energy available by an object by virtue of its position
Energy
The nature of energy and types of energy
Energy is the capacity to do work
Some fundamentals
Chemical energy:
energy stored within the structural units of chemical substances
Force: F = ma (kg m s-2, newton)
Law of conservation of energy:
Energy can neither be created nor destroyed.
However, energy can be converted from one form to another.
Work = Force × Distance (kg m 2 s-2, joule)
Energy has the same units as work
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Explosive agents
cyclotetramethylene-tetranitramine
(HMX)
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Energy changes in chemical reactions
Some fundamentals
Heat: the transfer of energy between two bodies that have different
temperatures. Heat has the same units (Joule) as energy and work.
Year 1955; 9110 m/s
Thermochemistry is the study of heat changes in chemical reactions.
Year 1940; 8754 m/s
Year 1910; 6930 m/s
Some terminologies
Year 1870; 7580 m/s
year of introduction
and detonation rate
Year 1990; 9380 m/s
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Some terminologies
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A reaction is endothermic if it takes in heat. Energy flows from the
surroundings to the system.
A reaction is exothermic if it gives off heat. Energy flows from the
system to the surroundings.
Weaker chemical bonds in products than in reactants.
More potential energy in products
Stronger chemical bonds in products than in reactants.
Less potential energy in products
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Enthalpy
An endothermic reaction
To quantify the heat flow of a process at constant pressure.
An exothermic reaction
At constant pressure, the heat transfer associated with
the reaction is quantified by ∆H.
The sign of ∆H is plus (+) for an endothermic process
The sign of ∆H is minus (-) for an exothermic process
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Thermochemical equations:
Standard enthalpies of formation
Enthalpy change for the formation of 1 mole of a substance from
the stable form of its elements at 1 atm and 250C
1. Stoichiometric coefficients always refer to the number of moles of
each substance.
2. Reverse the direction of a given reaction, reverse the sign of ∆H.
3. If a reaction is multipled by some factor, its ∆H must be multipled by
the same factor.
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Comment on those whose
standard enthalpies of
formation are zero
Standard enthalpy of formation for CaCO3 at 25 oC is -1206.9 kJ/mol
Ca(s) + C(s) + 3/2 O2(g) CaCO3(s)
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Enthalpies of Reactions
the Enthalpy of chemical species is a state function and is therefore
independent of the Path of the reaction!
Hess Law: If a change of state occurs in stages or steps (even if only
hypothetically), the enthalpy change for the overall (net) change is the
sum of the individual enthalpy stages for the individual steps.
{This statement is of course true for any state function, not just enthalpy}
C3H8(g) + 5O2 3 CO2(g) + 4 H2O(l)
∆H = ?
3C(graphite)+4H2(g) --> C3H8(g) ∆H = -103.9 kJ
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Two different ways of calculating the heat of combustion of methane:
Calorimetry – the measurement of heat changes
The corresponding apparatus is called calorimeter.
Another example:
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Exercise: 1.922 g of methanol was burned in a calorimeter. The
temperature of the water increased by 4.2 degree. If the amount of
water was 2000 g and the heat capacity of the calorimeter was 2.02
kJ/degree, what is the molar heat of combustion of methanol?
Please remember that ∆H is the heat flow measured at constant pressure.
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The first law of thermodynamics
Energy cannot be created or destroyed, but only converted to other forms
∆E = q + w
∆E: the change in the internal energy
of the system
q: the heat exchange
w: work done on the system
Internal energy
It is not possible to measure the internal energy of any system. However,
we can measure the changes in the internal energy.
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Heat (q) and work (w)
State functions (狀態函數
狀態函數)
狀態函數
Properties of a substance that don't depend on its History are
called Properties of State or State Functions.
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Enthalpy and the first law
∆E = q + w
∆E = ∆H – p ∆V
At constant pressure, q = ∆H
For a gas, w = -p ∆V (p: pressure; ∆V: the volume change)
The water in the center of the figure can be formed in many different
ways; two possible ways are shown above, that of melting and heating
ice, or cooling boiling water. You can't drink either the frozen or boiling
water, because they have energy content that is either too low or too
high to be comfortable by virtue of their temperature. But the room
temperature glass of water formed from either initial state would be
exactly the same, and just as drinkable. The State of the water does
not depend on where or what its been, just what it is now.
Any property that does not depend on history is a State Function.
Enthalpy, and Energy are State Functions.
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Entropy, Free Energy and Equilibrium
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Spontaneous processes
A spontaneous process is one that does occur under a given set conditions
The three laws of thermodynamics
Whether a process is spontaneous or not has nothing to do with its rate.
First Law: the law of conservation of energy
Second and third laws are covered in this chapter
A spontaneous process can be very fast or so slow that it appears not to
take place as all.
Spontaneous process and entropy
In chemistry, we are interested in whether a particular reaction
will “go”
usually means a favorable equilibrium constant and a conveniently rapid rate.
or will “not go”
either an unfavorable equilibrium constant or a rate too slow to be useful.
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Two examples spontaneous endothermic processes
Four possibility:
Examples of all
four are known.
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Entropy
Entropy – A process is spontaneous if it leads to an increase in
the entropy of the universe.
W=1
Entropy is a measure of the randomness or disorder of a system and
is related to probability.
W = number of microstates
spontaneous
non-spontaneous
S = k ln W
W=4
∆S = Sf - Si
∆S = k ln
Increasing
the
disorder
Wf
Wi
W f > W i then ∆S > 0
Decreasing
the
disorder
W=6
W f < W i then ∆S < 0
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An ordered state has a low probability of occurring and a small entropy,
while a disordered state has a high probability of occurring and a large
entropy.
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Entropy
State functions are properties that are determined by the state of the
system, regardless of how that condition was achieved.
Practices: How does the entropy of a system change for each of the
following processes?
(1)
(2)
(3)
(4)
energy, enthalpy, pressure, volume, temperature , entropy
Condensing water vapor
Forming sucrose crystals from a supersaturated solution
Heating hydrogen gas from 60 degree to 80 degree
Subliming dry ice
The second Law of Thermodynamics
Potential energy of hiker 1 and hiker 2 is the
same even though they took different paths.
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Entropy and the second law
Entropy change in the surroundings
How are the surroundings affected by heating and cooling?
Exothermic reactions (increase or decrease) entropy of surroundings?
How about endothermic reactions ?
How about temperature?
Therefore, we need to consider how the entropies of the system and
the surroundings changes during a process
Entropy change in the system
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Spontaneity and temperature
A process (or a reaction) may be spontaneous at one temperature
and non-spontaneous at a different temperature.
The third law of thermodynamics and absolute entropy
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The tabulated values of the absolute entropy allow us to calculate the
change in entropy for chemical reactions:
S = k ln W
W=1
S=0
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Gibbs Free Energy
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Standard free energies of formation
Change in free energy that accompanies the formation of 1 mole of a
substance from its constituent elements (all materials in their standard
states)
Criterion for
spontaneity is
the Gibbs free
energy change
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Standard free energy at 25oC
Calculate the standard entropy change for
Quantitative prediction: the sign of ∆S ?
Quantitative
prediction ?
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Examples
Applications of ∆ G = ∆ H – ∆S · T
(1)
(2)
(3)
(4)
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Temperature and chemical reactions
CaCO3 (s) ↔ CaO (s) + CO2 (g)
∆G0
=
∆H0 - T ∆S0
=
177.8 - 298 (165/1000)
=
130 kJ
(using data in appendix 2)
∆G0 > 0 under 1 atm and 298K and not spontaneous
To make ∆G0 < 0 ( to make the reaction forward)
∆H0
< T ∆S0
T
> (∆H0/ ∆S0) = 177.8/(160.5/1000) = 1108 K or 835°C
I2
Since both ∆S0 and ∆H0 change with T, this kind of estimation is only approximate
Phase Transition
H2
Stage 1
Stage 2
Stage 3
Phase transition occurs always at equilibrium ∆G=0
Therefore, one can calculate ∆S
∆S = ∆H/T
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Practice:
At what temperature is the following process spontaneous at 1 atm?
Br2(l) Br2(g)
∆H0 = 31.0 kJ/mol∆S0 = 93.0 J/K
What is the normal boiling point of the liquid Br2?
Solution:
∆G0 = ∆H0 – T ∆S0
0 = 31.0 – T (93.0/1000)
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Glucose
Proteins
CO2 + H2O
Amino acids
T = 333 K
Thermodynamics in living systems
The principle of couple reactions
ZnS Zn + S
S + O2 SO2
∆G0 = 198.3 kJ
∆G0 = -300 kJ
Overall reaction
ZnS + O2 Zn + SO2
∆G0 = -108.1 kJ favored
Couple reactions In biological systems
unfavored
favored
Alanine + Glycine
∆G0 = +6.9 kcal
ATP + H 2O + Alanine + Glycine
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∆G0
= -0.4 kcal
Alanylglycine
K<1
ADP + H 3PO4 + Alanylglycine
K>1
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Practice
Free energy and equilibrium
Is the reaction spontaneous under the following conditions?
aA + bB
cC + dD
Q: reaction quotient
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Applications
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