Download CHM 122 Chapter 8 -Thermochemistry: Chemical Energy

CHM 122
Chapter 8 -Thermochemistry: Chemical Energy
8.1 Energ y and Its Conser vation
A. Definitions
Thermochemistry is a part of Thermodynamics dealing with energy changes associated with physical and
chemical reactions
Why do we care?
- Will a reaction proceed spontaneously?
- If so, to what extent?
However, it won’t tell us:
- How fast the reaction will occur
- The mechanism by which the reaction will occur
Energy is the capacity to do work or to transfer heat
For example if you climb a mountain, you do some work against the force of gravity as you carry yourself and your
equipment up the mountain. You can do this because you have the energy, or capacity to do so, the energy being
supplied by the food that you have eaten. Food energy is chemical energy –energy stored in chemical compounds
and released when the compounds undergo the chemical process of metabolism
- Kinetic Energy: energy associated with mass in motion
- Potential Energy: energy associated with the position of an object relative to other objects (energy that is stored can be converted to kinetic energy)
System: portion of the universe under study
Surroundings: everything else
Open System: can exchange energy and matter through boundary
Closed System: can exchange energy through boundary
Isolated System: can exchange neither with surroundings
We can define the system and surroundings however we want!
Example: You place 10.0g of LiOH in a Styrofoam cup of water. What is the system, and what are the
surroundings?
Energy and units
One of the earliest units of energy to be devised was the unit of calorie, or cal. Since the calorie represents a
small amount of heat energy and we usually work with larger quantities of matter the unit of kilocalorie or kcal is
used instead.
The SI unit is Joules 1 cal = 4.184 J (exactly)
8.2 Internal Energ y and State Functions
Energy (E) can be transferred two different ways:
1. By doing work (w) (applying a force over a distance)
w=Fxd
Work: Can be electrical, mechanical, etc.
2. Transferring heat (q) (results in a change in temperature)
Heat: The amount of thermal energy transferred from one object to another as the result of a temperature difference
between the two
Thermal Energy: The kinetic energy of molecular motion and is measured by finding the temperature of an object
We will symbolize heat energy transferred by the letter q.
HOT OBJECT
q
COLD OBJECT
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Heat transfer occurs by way of collisions between randomly moving particles of matter. The particles with higher
thermal energy are moving more quickly, when they collide with slower moving particles, some of their energy is
transferred to the slower particle as heat energy increasing its speed
Note: w, q, and E all have the same units (Joule), but:
- w & q depend on path (path function)
- E is independent of path (state function)
 A function or property whose value depends only on the present state, or condition, of the system,
not on the path used to arrive at that state
First Law of Thermodynamics
“The total energy of the universe is constant.”
“Energy is neither created nor destroyed in a process, only converted to another form.” -Conservation of Energy
“You can’t win . . . you can only break even.”
The sum of all the kinetic and potential energies of particles making up a substance is called internal
energy, E
The internal energy of systems that
are more complex than an ideal gas
can't be measured directly. But the
internal energy of the system is still
proportional to its temperature. By
monitoring the change in
temperature of a system we can
determine whether the internal
energy of the system increases or
decreases
8.3 Expansion Work
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Work done as the result of a volume
change in the system
Recall, w = F · d and P = F/A
therefore w = F x d = - P∆V
8.4 Energ y and Enthalpy
When there is no work done by the system, the heat given off or absorbed by the reaction would be equal to the
change in the internal energy of the system.
Esys= q (if and only if w = 0)
Esys= qv (at constant volume)
However, if a gas is driven out of the flask during the reaction, the system does work on its surroundings. If the
reaction pulls a gas into the flask, the surroundings do work on the system. The amount of heat given off or
absorbed will no longer equal to the change of the internal energy of the system because some of the heats has been
converted to work.
Esys= q + w
qP = ΔE + PΔV
(at constant pressure)
The amount of heat absorbed or released during a chemical reaction by a system is also call Enthalpy, ΔH.
Under constant pressure, qp = ΔH
qP = ΔE + PΔV = H
Enthalpy is a state function whose value depends only on the current state of the system, not on the path taken to
arrive at that state. Enthalpy is an extensive property because the magnitude of ΔH depends on the quantity of
reactant
ΔH = Hfinal - Hinitial = Hproducts - Hreactants
For convenience, the sign and value of ΔH is always provided in a thermochemical equation.
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2 Al(s) + Fe2O3(s)  2Fe(s) + Al2O3(s)
ΔH = -825 kJ
Write the above thermochemical equation in the reverse reaction
8.6The Ther modynamic Standard State - Enthalpies of Physical and Chemical Change
Enthalpy of Fusion (Hfusion): The
amount of heat necessary to melt a
substance without changing its
temperature
Enthalpy of Vaporization (Hvap): The
amount of heat required to vaporize a
substance without changing its
temperature
Enthalpy of Sublimation (Hsubl): The
amount of heat required to convert a
substance from a solid to a gas without
going through a liquid phase
Thermodynamic Standard State: Most stable form of a substance at 1 atm pressure and at a specified
temperature, usually 25 °C; 1 M concentration for all substances in solution.
Standard enthalpy of reaction is indicated by the symbol ΔHo
Example: Applying Stoichiometry to Heats of Reaction
 A propellant for rockets is obtained by mixing the liquids hydrazine, N2H4, and dinitrogen tetroxide, N2O4.
These compounds react to give gaseous nitrogen, N2 and water vapor, evolving 1049 kJ of heat at constant
pressure when 1 mol N2O4 reacts.

Write the thermochemical equation for this reaction

Write the thermochemical equation for the reverse of the reaction

How much heat evolves when 10.0 g of hydrazine reacts according to the reaction described in (a) ?
Example: Calculate the ΔH for the reaction that occurs when 4.1 g CH4 decomposes according to the following
reaction:
CH4(g) + 2O2(g) -> CO2(g) + 2H2O(g) ΔH = - 802 kJ/mol
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8.6 Calorimetry and Heat Capacity
In general, when heat transfers to an object, the temperature increased. When heat lost, the temperature decreased.
This relationship can be describe in this equation
Heat Capacity (C): The amount of heat required
to raise the temperature of an object or substance a
q = C x ΔT at constant volume
given amount.
q = SH x m x ΔT
at constant pressure
Specific Heat: The amount of heat required to
raise the temperature of 1 g of a substance by 1 °C.
The heat evolved in a chemical reaction can be determined by a process called calorimetry. Calorimeter is a
device that use to measure the heat flow. It is well-insulated so that, ideally, no heat enters or leaves the
calorimeter from the surroundings
At constant pressure
at constant volume
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Example
Iron metal has a specific heat of 0.449 J/goC. How much heat (in kJ) is transferred to a 5.00g piece of iron,
initially at 20.0oC, when it is placed in a pot of boiling water? Assume that the temperature of the water is
100.0oC and that the water remains at this temperature, which is the final temperature of iron
Example You dissolve 5.25 g of NaOH into 100.0 mL of water in a Styrofoam cup. The temperature of the water
increases from 23.3 oC to 40.5 oC. What is the enthalpy of reaction in kJ/mol
Example: A 400. 0 g piece of iron (C (iron) = 0.38 J/g°C) is heated in a flame and dropped into a beaker containing
1000. 0 g of water at 20.0°C. The final temperature is 32.8°C.What was the initial temperature of the iron bar?
8.9 Hess’s Law
Hess’s Law
Energy conservation is the basis of Hess’s Law which states that, if a reaction is the sum of two or more other reactions, the
ΔH overall is the sum of theΔH values of the constituent reactions
ΔH°total = ΔH°1 + ΔH°2 + etc….
For a reaction of
3 H2(g) +
N2(g)  2NH3(g)
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Example: Find ΔHorxn for the following reaction:
3H2(g) + O3(g)  3 H2O(g)
ΔHorxn = ??
Use the following reactions with known ΔH’s
2H2 (g) + O2(g)  H2O(g)
Δ Ho = -483.6 kJ
3O2(g)  2 O3 (g)
Δ Ho =.
+285.4 kJ
Example
 Find ΔHorxn for the following reaction
C(s) + H2O(g)  CO(g) + H2(g)
Use the following reactions with known H’s
C(s) + O2(g)  CO2(g)
2CO(g) + O2(g)  2CO2(g)
2H2 (g) + O2(g)  2H2O (g)
Horxn = ?
ΔHo = -393.5 kJ
Δ Ho = -566.0kJ
Δ Ho = -483.6 kJ
Standard Heats of For mation
The enthalpy change which occurs for a reaction which forms one mole of a compound from its constituent elements in their standard states
at standard conditions is called the
STANDARD MOLAR ENTHALPY OF FORMATION or Enthalpy of Formation for short and is signified by ΔH°f .
Ag(s) + ½ Cl2(g)  AgCl(s)
ΔHrxn = -127.0 kJ
Then
Ho =
ΔH°f
(AgCl(s)) = -1270 kJ/mol
Hof (Products)
-
Hof (Reactants)
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Example: Using standard heats of formation, calculate the standard enthalpy of reaction for the photosynthesis of
glucose (C6H12O6) and O2 from CO2 and liquid H2O.
6CO2(s) + 6H2O(l)  C6H12O6 (s) + 6O2(g)

Calculate the enthalpy change for the following reaction
8Al(s) + 3 Fe3O4(s)  4Al2O3(s) + 9 Fe(s)
Given
ΔH°f (Fe3O4) = -1120.9 kJ/mol
ΔH°f (Al2O3) = -1669.8 kJ/mold
An Introduction to Entropy
Entropy (S): The amount of molecular randomness in a system
Spontaneous Process: A process that, once started, proceeds on its own without a continuous external influence
Spontaneous processes are
•
favored by a decrease in H (negative ΔH).
•
favored by an increase in S (positive ΔS).
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Nonspontaneous processes are
•
favored by an increase in H (positive ΔH).
•
favored by a decrease in S (negative ΔS).
Example
 Predict whether ΔSo is likely to be positive or negative for each of the following reactions
a. H2C=CH2(g) + Br2(g)  BrCH2CH2Br(l)
b. Consider the following figures
Example
 Is the Haber process for the industrial synthesis of ammonia spontaneous or nonspontaneous under standard
conditions at 25.0oC.
3 H2(g) +
N2(g)  2NH3(g)
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