1. Define and Explain the Following Terms
Chapter 6
Thermochemistry
This chapter develops the concepts of thermochemistry. Upon completion of Chapter 6, your students should be able to:
1. Define and explain the following terms:
· Energy
· Radiant energy
· Thermal energy
· Chemical energy
· Potential energy
· Thermochemistry
· Open system
· Closed system
· Isolated system
· Endothermic
· Exothermic
· Enthalpy (DH)
· Calorimetry
· Heat capacity
· Specific heat
2. Classify common processes as endothermic or exothermic.
3. Use thermochemical equations and stoichiometry to determine amount of heat lost or gained in a chemical reaction.
4. Perform calculations involving specific heat, mass and temperature change.
5. Sketch the main components of a constant-volume bomb calorimeter.
6. Determine heats of reactions given experimental data collected in a calorimetry experiment.
7. Calculate standard enthalpy of reactions given the standard enthalpy of formations for products and reactants.
8. Apply Hess’s law to a multi-step process to determine standard enthalpy of reaction.
9. Describe heat of solution, lattice energy, heat of hydration, heat of dilution, system, surrounding, and internal energy.
10. Classify properties of materials as state functions or non-state functions.
11. Restate the First Law of Thermodynamics.
12. Name the sign conventions for work and heat used in the textbook.
13. Apply heat and work relationships to gas-phase problems.
14. Define H in term of E, P, and V.
15. Calculate change in internal energy (DE) given thermochemical equations.
Section 6.1 The Nature of Energy and Types of Energy
Energy, the ability to do work, takes many forms, including kinetic, potential, radiant, thermal, and chemical energy. Students sometimes confuse energy and temperature. If we place a pot of water on the kitchen stove and turn the dial to high, we observe that the water’s temperature increases (using a thermometer) as the water warms up. This process continues until we reach the boiling point of the water. Once this temperature is reached, the water no longer increases in temperature, but we are still adding energy to the system because the dial on the stove still reads “high”. What this shows is that the temperature of the water and the energy pumped into it are proportional until the boiling point is reached. At that point, the energy is used to change the physical state of the liquid water to steam without changing the temperature of either. Using the concept of conservation of energy, we know that the energy provided by the stove at the boiling point of water must be used to break the forces of attraction that hold the water molecules together in the liquid phase.
Section 6.2 Energy Changes in Chemical Reactions
Heat is defined as the transfer of thermal energy from a hot body to a cold body. It is a process; it is not energy. It is incorrect to refer to heat energy. The study of transfer of energy that occurs during chemical reactions is known as thermochemistry. In thermochemistry, we have the system and the surroundings that make up the universe. We also refer to open, closed, or isolated systems. An open system allows for the transfer of both energy and mass; a closed system allows the transfer of only energy; and an isolated system does not allow the transfer of either energy or mass.
Exothermic processes give off energy. The prefix exo- refers to external, thus exothermic means energy that is given off. The opposite of exothermic is endothermic which means energy is absorbed. In an exothermic reaction, energy can be thought of as one of the reaction products. Energy can be thought of as a reactant in endothermic reactions. Students sometimes think that since energy is added to the system in an endothermic reaction, the temperature of the system should go up. Their logic is that if energy is added to the system, then the temperature should rise. In fact, the energy is removed from the surroundings; thus a cooling effect occurs and is incorporated in the products as energy stored in the chemical bonds. Exothermic reactions seem to be easier for students since, if energy is released, then the temperature of the surroundings should rise.
Section 6.3 Introduction to thermodynamics
It can be difficult for students to understand state functions. Properties of systems that are independent of how they were achieved are known as state functions. Energy, volume, pressure, and temperature are all state functions. Figure 6.4 on page 234 demonstrates state functions.
On the other hand, heat and work are not state functions because they depend upon the path taken to reach the final state. It can also be explained by asking students which one will require more work--going to first floor taking stairs versus taking the elevator.
The first law of thermodynamics states that energy can be converted from one form to another but cannot be created or destroyed. We know it as the law of conservation of energy. This can be expressed as
It is important to understand the sign convention for heat and work. The convention for q is positive for endothermic processes and negative for exothermic reactions. Work is positive for work done on the system by the surroundings and negative for work done by the system on the surroundings. The sign convention for q is the same as for ΔH for exothermic and endothermic reactions. If we think of work being done on the system as a “reactant”, then it follows that the sign on w will be positive as q is positive for endothermic reactions. For systems where work is done by the system, work can be considered as a “product” and thus is similar to exothermic reactions and the sign would be negative. It is important to point out difference in internal energy, heat and work. Work being directly proportional to force is an easy one to explain by using the example of pushing a table versus pushing a book. Relationship of work to distance can be explained by using the example of carrying their text book down the stairs versus carrying their desk.
Example 6.1 on page 237 can be difficult for students to understand since they cannot see the relationship between the units of pressure and volume to joules. This example also draws the conclusion that work is not state function.
If we recall that R = 0.0821 L•atm / mol•K = 8.314 J/mol•K, then we have 1 L•atm = 101.3 J. The terms L•atm may not be easily recognized as units for energy, but indeed they are.
We defined change in enthalpy as
and also
We get
If we substitute w = -PΔV and hold pressure constant we get
ΔH = qp
Since most of the common reactions that we encounter are performed at constant pressure, the heat that we observe, either lost or gained, is a direct measure of ΔHrxn.
Section 6.4 Enthalpy and Chemical Reactions
Since most of the reactions that we do are open to the atmosphere, we are usually interested in constant-pressure processes. The energy that is transferred in a constant-pressure process is called enthalpy and is symbolized by H. H is an extensive property as wells as a state function which means that ΔH, the change in enthalpy, is independent of the path taken in going from state one to state two. ΔH is equal to H for the products minus H for the reactants. If ΔH is negative, the reaction is exothermic (energy is released) and if ΔH is positive, the reaction is endothermic (energy is absorbed).
One method to assist your students in this is to consider the following sets of reactions. Assume the following reaction occurs:
A + B C
Assume further that reactants A and B contain a total of 100 units of energy and product C contains 80 units of energy. Therefore, because the law of conservation of energy must hold, 20 units of energy must be released. If energy is released, then the reaction must be exothermic. For this reaction
The important point is that ΔH<0 or is negative. Therefore, ΔH for exothermic reactions must be negative. The converse is true for endothermic reactions. For example, assume the following reaction occurs:
F + E G
Assume that the reaction is endothermic and that the sum of the energy in F and E is 70 units and the energy of G is 120 units. Thus 50 units of energy must be supplied for this reaction to proceed. The reaction is endothermic.
Here ΔH is >0 which is always true for endothermic reactions.
We could think of the exothermic reaction as
A + B C + energy
where energy is a product for the exothermic reaction and
energy + F + E G
where energy is a reactant for the endothermic reaction.
The melting of ice is
H2O(s) H2O(ℓ)
The heat that is required to do this is 6.01 kJ for one mole and is known as the heat of fusion of water. Therefore, ΔH = +6.01 kJ. The reverse process
H2O(ℓ) H2O(s)
would have ΔH = -6.01 kJ. Therefore, the melting of ice is endothermic because heat is removed from the surroundings while the freezing of water releases heat to the surroundings and is exothermic.
The process of converting liquid water to steam
H2O(ℓ) H2O(g)
has ΔH = +44.0 kJ for one mole and is known as the heat of vaporization. The reverse process
H2O(g) H2O(ℓ)
is exothermic. It is for this reason that steam burns are often so severe. The conversion of the steam to liquid water releases a large amount of energy to the surroundings resulting in a burn to exposed skin.
Section 6.5 Calorimetry
Specific heat and heat capacity are often confused. Specific heat is defined as the amount of heat required to raise the temperature of one gram of material one degree Celsius. It has units of joules per gram degree Celsius. Heat capacity is the amount of energy required to raise a given quantity of material one degree Celsius. It has the units of joules per degree Celsius. Just looking at the units, students will see the specific heat times mass is equal to heat capacity.
Both constant-volume bomb calorimeters and constant-pressure calorimeters are introduces in this section. Most students will not experience bomb calorimetry since it requires fairly sophisticated equipment.
Example 6.8 on page 250 shows that the heat of neutralization for HCℓ with NaOH is -56.2 kJ/mol. This is a good opportunity to review the concept of net ionic equations since in fact -56.2 kJ/mol is the ΔH for the reaction of
H+(aq) + OH—(aq) H2O(ℓ)
Therefore, any neutralization reaction that results in this net ionic equation will always have the same ΔH. This concept may not be obvious to your students. The reverse of this reaction is given in Table 6.3 and has a ΔH equaling +56.2 kJ/mol.
Section 6.6 Standard Enthalpy of Formation and Reaction
The most stable form of oxygen at 25oC is molecular oxygen, O2, and not ozone, O3. Your students will likely accept that fact because of news coverage of destruction of the ozone layer by chlorofluorocarbons. They may have a more difficult time accepting that graphite is more stable than diamond because they are likely aware that graphite will burn but diamonds "last forever". What needs to be explained is that by stability we are referring to thermodynamic stability and not chemical reactivity.
Perhaps the following analogy may assist in explaining the difference between thermodynamic stability and chemical reactivity. With respect to potential energy, a ten-pound box sitting on top of a cliff overlooking a canyon is not as stable as a similar box sitting one inch above the canyon floor. If the box on the top of the cliff is sitting firmly on the ground while the box at the bottom of the canyon is balancing on the head of a pin, certainly the box at the bottom of the canyon is the more reactive of the two boxes. The box with the lowest energy is the more reactive. The analogy to be made is the box on top of the cliff corresponds to diamond (higher energy, less reactive) while the box on the pin corresponds to graphite (lower energy, more reactive).
The change in enthalpy for a reaction, ΔH, can be determined experimentally using calorimetry or indirectly by using Hess’s law. Example 6.5 demonstrates the traditional method of using Hess’s law to determine the for C2H2. Some students become confused or intimidated by all of the chemical formulas. They may be more comfortable if each of the chemicals is assigned a single letter variable and then, after writing the various equations, solve for the desired equation. For example, using the equations labeled a, b, and c in example 6.9, let us define the following:
C (graphite) = A
O2(g) = B
CO2(g) = C
H2(g) = D
H2O(ℓ) = E
C2H2(g) = F
Reaction a in example 6.9 becomes
aa)
Reaction b
bb)
Reaction c
cc)
The desired reaction for the formation of C2H2 from its elements is
dd)