Chapter 5 - Thermochemistry
Chapter 8 Section 8 – Covalent Bond Enthalpy
Vocabulary:
Energy –
Work – Energy used to move an object over some distance. w = ___________
w is work, F is the force, d is the distance over which the force is exerted.
Heat – Energy transferred between substances due to a ________________________________.
(From the higher temperature substance to the lower temperature substance.)
Potential Energy – Energy an object possesses by virtue of its ______________________________.
PE = _______________________________
Kinetic Energy – Energy an object possesses by virtue of its________________________________.
KE=
(true except at 0K)
Transferal of Energy Example: As a ball drops –
a. The potential energy of the ball of clay is increases when it is moved from the ground to
the top of the wall.
b. As the ball falls, its potential energy is converted to kinetic energy.
c. When it hits the ground, its kinetic energy falls to zero (it’s not moving); some of the
energy does work on the ball, the rest is dissipated as heat.
Units of Energy –
The SI unit of energy is the Joule
An older, non-SI unit still in widespread use: The calorie (cal).
1 cal =
1 food Calorie =
System – The ___________________________________
(Here the hydrogen & oxygen molecules)
Surroundings – _________________________________
(Here the cylinder & piston)
First Law of Thermodynamics a.k.a. Law of Conservation of Energy
Energy is neither ___________________________________
The total energy of the universe is a constant; if a system loses energy, it must be gained by
the surroundings, and vice versa.
Internal Energy – The sum of all kinetic and potential energies of all components of the
system. (We call it E)
Usually we don’t measure E, we measure _______________________________________.
By definition, the change in internal energy, E, is the final energy of the system minus the
initial energy of the system:
ΔE =
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•
If ΔE is positive (E > 0, Efinal > Einitial)
•
The system _________________ energy from the surroundings.
•
The energy of the system _____________________
•
The energy of the surroundings _____________________
If E is negative (E < 0, Efinal < Einitial )
•
The system _________________ energy from the surroundings.
•
The energy of the system _____________________
•
The energy of the surroundings _____________________
When energy is exchanged between the system and the surroundings, it is exchanged as either
__________________________________________.
E =
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State functions – Depend only on the _____________________ of the system, not on the
_____________ by which the system arrived at that state.
ΔE depends only on Einitial and Efinal and is therefore a state function.
However, q and w are not state functions.
• Whether the battery is shorted out or is discharged by running the fan, its E is
the same, but q and w are different in the two cases.
Generally…
Quantities we look at a change in (E) are state functions.
Quantities we don’t look at a change in (q & w) are not state functions.
P-V Work
When a chemical reaction occurs, commonly the only work done is a change in volume of a gas
pushing on the surroundings (or being pushed on by the surroundings).
We can measure the work done by the gas if the reaction is done in a vessel that has been
fitted with a piston.
w = ________________
Example #1
Hydrogen and oxygen gas are in a cylinder with a movable piston. When ignited, the
system loses 1150 J of heat to the surroundings. The reaction also causes the piston to
rise (from the expanding gases) and 480 J of work is done as the piston is pushed against
the atmosphere. What is the change in internal energy of the system?
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Enthalpy
If a process takes place at constant pressure (as the majority of processes we study do) and
the only work done is this pressure-volume work, we can account for heat flow during the
process by measuring the enthalpy of the system.
Enthalpy is the internal energy plus the product of pressure and volume:
H=
When the system changes at constant pressure, the change in enthalpy, H, is
H = (E + PV)
• This can be written
H = E + PV
Since E = q + w and w = −PV, we can substitute these into the enthalpy expression:
H = E + PV
H =
H =
So, at constant pressure the change in enthalpy is ________________________________!
•
We use H to describe chemical reactions because most reactions happen at constant pressure,
and at constant pressure H= q and q is (relatively) easy to find!
Often for chemical reactions, PΔV is very small and ΔE ≈ ΔH
Endothermic – ΔH is _______________ (heat transferred to the system from the surroundings)
Phase changes from solid to liquid to gas are always endothermic
Exothermic – ΔH is ________________ (heat transferred from the system to the surroundings)
Phase changes from gas to liquid to solid are always exothermic
Thermochemical Equations
Must specify ________________________________ of products and reactants
 Enthalpy depends on state!
Must specify ______________________________
Example:
CaO(s) + CO2(g)  CaCO3(s)
H298 = -178KJ
H, is called the enthalpy of reaction, or the heat of reaction (HRXN)
(Temp at which the reaction occurs is sometimes included)
HOW DO WE FIGURE OUT H TO INCLUDE IN OUR EQUATION???
Method #1 to find H: Calorimetry
When H =q, you can calculate q like you did in your first year class when given the proper
information!
Specific heat of water (and most aqueous sol’ns):
4.184J/gºC or 4.184J/g K
Constant Pressure Calorimetry:
By carrying out a reaction in aqueous solution in a simple calorimeter, we can indirectly
measure the heat change for the system by measuring the heat change for the water in
the calorimeter.
Constant Volume Calorimetry:
Reactions can be carried out in a sealed “bomb” calorimeter to measure the heat absorbed by
the water.
• Because the volume in the bomb calorimeter is constant, what is measured is really the
change in internal energy, E, not H.
• For most reactions, the difference is very small.
Temperature changes are converted to heat energy by using heat capacities.
• Heat Capacity (C) - The amount of energy required to raise the temperature of a
substance by 1 K (J/K)
• Molar Heat Capacity (Cm) – The amount of energy required to raise the temperature of 1
mole of a substance by 1 K (J/molK)
• Specific Heat Capacity (or simply Specific Heat) (Cs) - The amount of energy required to
raise the temperature of 1 g of a substance by 1 K. (J/gK)
• Calorimeter Constant (cc) – The heat capacity of a calorimeter(J/K)
(*temperatures may also be in °C)
Example #2
The specific heat of ammonia is 4.381 J/gK. Calculate the heat required to raise the
temperature of 1.50g of ammonia from 213.0K to 218.0K.
Example #3
2.15 g of methane is combusted (burned) to form CO2 and H2O. All of the heat evolved is used
to raise the temperature of 100.0 g of water by 17.8 K. Calculate the heat (q) of
combustion per mol of methane.
Example #4
50.00 mL of 0.750 M HCl solution and 50.00 mL of 0.875 M KOH solution are combined in a
calorimeter. The temperature of the calorimeter and both solutions goes from 25.0 oC to
29.6 oC. Assume that the specific heat of each solution is the same as water, 4.184 J/g oC
and that the density of each solution is 1.00 g/mL. The heat capacity of the calorimeter
is 125 J/oC. Calculate the heat of reaction per mole of water formed.
Some more about Enthalpy…
1. Enthalpy is an extensive property (it depends on amount).
2. H for a reaction in the forward direction is equal in size, but opposite in sign, to H for
the reverse reaction.
3. H for a reaction depends on the state of the products and the state of the reactants.
Enthalpy is an extensive property so changing the stoichiometry changes the value of ΔH
CaO(s) + CO2(g)  CaCO3(s)
H = -178kJ
2CaO(s) + 2CO2(g)  2CaCO3(s)
H = -356kJ
½ CaO(s) + ½ CO2(g)  ½ CaCO3(s) H = -89.0kJ
Example #5
How much heat is released when 4.50g of methane gas is burned in a constant pressure
system?
CH4(g) + O2(g)  CO2(g) + 2H2O(g) ΔH = -802 kJ
Method #2 to find H: Hess’s Law
Hess’s law states that “If a reaction is carried out in a series of
steps, H for the overall reaction will be equal to the sum of the
enthalpy changes for the individual steps.”
Example #6
Given:
C(s) + O2(g)  CO2(g)
ΔH = -393.5 kJ
CO(g) + ½ O2(g)  CO2(g) ΔH = -283.0 kJ
Find ΔH for the following:
C(s) + ½ O2(g)  CO(g)
Example #7
Given:
C2H2(g) + 5/2 O2(g)  2CO2(g) + H2O(l)
C(s) + O2(g)  CO2(g)
H2(g) + ½ O2(g)  H2O(l)
Find: 2 C(s) + H2(g)  C2H2(g)
ΔH = - 1299.6 kJ
Δ H = - 393.5 kJ
Δ H = - 285.9 kJ
Enthalpy (Heat) of Formation
An enthalpy of formation, Hf, is defined as the enthalpy change for the reaction in which a
compound is made from its constituent elements in their elemental forms.
H2(g) + ½ O2(g)  H2O(g) ΔHf = -242.3 kJ/mol
Hf depends on physical state, pressure and temperature
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Standard enthalpies of formation, H f , are measured under standard conditions (25°C and
1.00 atm pressure).
The standard enthalpy of formation is the heat of formation of one mole of the compound in
its standard state formed from its constituent elements in their standard states.
The standard state of formation of any element in its most stable form is by convention
taken to be zero
• values are tabulated in Appendix C of your textbook (page 1059)
Method #3 to find ΔH
The change in enthalpy, H, is the enthalpy of the products minus the enthalpy of the
reactants:
n and m are the stoichiometric coefficients of the chemical equation
Example #8
(a) Calculate the standard enthalpy change for the combustion of 1 mole of benzene, C6H6(l) to
CO2(g) and H2O(l) (b) If the ΔH°rxn for the combustion of 1 mole of propane is – 2220 kJ,
compare the quantity of heat produced by combustion of 1.00 g propane C3H8(g) to that
produced by 1.00 g of benzene
Example #9
The standard enthaply change for the reaction
CaCO3(s)  CaO(s) + CO2(g)
is 178.1 kJ
Given ΔH°f CaO(s) as -635.5 kJ and Δ H°f CO2(g) as -393.5 kJ
calculate Δ H°f CaCO3(s) (without looking it up in the table)
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Method #4 to find ΔH
The energy required to break a bond between two atoms is called the bond enthalpy.
We use the letter D to represent bond enthalpies. (See page 316 in textbook)
 Ex: D(H-Br) = 366 KJ/mol
Δ Hrxn =
Example #10
Calculate the enthalpy of the reaction when methane gas and chlorine gas react to form methyl
chloride (CH3Cl) and hydrogen chloride.
CH4(g) + Cl2(g)  CH3Cl(g) + HCl(g)
Chapter 5 Book HW
5(3, 15, 27, 40, 42, 44, 50, 56, 64, 66, 70, 74)
Chapter 8 Book HW
8(69, 73)