Changes in Energy

These two are not the same thing! Thermal energy is, well, energy, and matter contains thermal energy but not heat. Heat is thermal energy in transit. This transfer of thermal energy only happens in one direction: from a higher temperature substance to a lower energy substance and NEVER in the opposite direction. This is the 2nd law of thermodynamics.

There are three temperature scales that are currently in use. They all differ by the number of degrees between water freezing and boiling. The one that is most familiar to you is the Fahrenheit scale. This scale has water freezing at 32°F and boiling at 212 °F. The Celsius (same as Centigrade) is a "metric based scale" with water freezing and boiling at 0°C and 100°C respectively. The temperature scale used by many scientists is the Kelvin scale which is a more accurate measure of particle motion which is what temperature really is. Of special interest is the temperature at which all particle motion (kinetic energy) is supposed to stop. The temperature at which KE is the lowest it can be is absolute zero or 0 K. To convert between °C and °F are found below:

°C = 5/9(°F-32) or °F = 9/5°C +32

The difference here is one of types of motion. Even though the picture here says heat (scientists can be sloppy about using the terms heat and thermal energy interchangeably). Thermal energy (internal energy) represents all types of motion (vibrational, rotational, and translation) articles can have as well as attractions and repulsions between the particles (the latter two represent potential energy) whereas temperature is just translational energy or kinetic energy. Something at a lower temperature can, in fact, have more thermal energy than something at a higher temperature...remember our demo!

You will see this symbol (∆) all over this chapter. This symbol means "change". The equation above means that the change in energy is always a combination of work produced and heat (q). Obviously, heat, work and energy are all in the same units: the joule is the unit we will use the most (calories used to be used for heat and still are and horsepower was and is used for work). This topic in physics and chemistry has not only numbers but sign as well. When energy increases, the energy is positive and when it is negative, energy is decreasing. Think graphically (see figures below).

This topic is very "universal" meaning that you think not only about the "system" you are studying but also its "surroundings". If we looking at changes in energy for a metal chair on a sunny deck, that situation does not involve only the chair (if it did, the chair would not get hot) but also its surroundings. Energy is transferred in this example from the surroundings to the system (the chair). This same analogy would apply to a chemical reaction that required an input of energy as shown in the figure. The surroundings are losing energy (negative ∆E; work and heat are negative too) and the system is gaining energy (positive ∆E; work and heat are positive too). Having said all of that, we will focus on the system.

When thermal energy is transferred in and out of a system, work may be involved. Conversely, work done on or by the system can influence the amount of thermal energy in a system. The figure to the right shows the relationships between work and thermal energy.

There are two basic forms of energy: kinetic energy (energy of motion) and potential energy (energy of position). Whenever energy is transformed (e.g., electrical from chemical (fuel)); you are turning potential into kinetic or vice versa. Potential energy is locked up in chemical bonds and can be released when those bonds are broken. When a molecule forms, potential energy drops (think ball on an incline). This is really the only reason molecules do form-the system rests at a lower energy. When chemical reactions occur, energy is put into breaking the bonds (ball has to be pushed back up the hill). Once the bonds break, the ball rolls down the hill again. There is a balance between this constant switching from potential to kinetic and that is what thermochemistry is all about.

Heats of reactions or enthalpies can also be calculated from other kinds of data (besides calorimetry). When reactants react to form products, bonds are broken and then reformed into new arrangements. It takes energy to break bonds (endothermic) and energy is released when bonds form (exothermic). This endothermic and exothermic business occurs in every chemical reaction. The balance between the two dictates whether a chemical reaction is overall exothermic (bond forming step is bigger) or endothermic (bond breaking step is bigger). If a reaction is endothermic, then enthalpy can be thought of as entering the reaction on the reactant side and if a reaction is exothermic, heat can be thought of as exiting the reaction on the product side.

C (s) + O2 (g) → CO2 (g) + 393.5 kJ

H2O (l) + 285.8 kJ → H2 (g) + 1/2 O2 (g)

Which is endothermic above? Exothermic?

Take the enthalpy out of the equations above and it could be written separately from the reaction. If it is, the sign of the enthalpy value tells you whether the reaction is endothermic or exothermic. If ∆Hrxn is negative, the reaction is exothermic. If ∆Hrxn is positive, the reaction is endothermic.

As you will see in our activities, the amount of thermal energy (Q) transferred, aka heat, depends on three things: amount of substance (mass), the substance itself (specific heat), and the temperature of the substance before and after the transfer of heat (∆T). This equation takes all three into account: Q = mc∆T

There are special chemical reactions that essentially define the energy change when making compounds from elements. These equations are formation equations and they can be used to calculate the energy for any chemical reaction using the equation below. For these equations, the energy change is essentially the energy "locked up" in the formation of a compound.

Thermal energy is the least useful type of energy in terms of being able to do work. Every time energy is transferred, the amount of energy available to do work decreases. This dispersal of energy into "non useful" energy is entropy. It can also be defined as amount of "randomness". Entropy (S) can be expressed mathematically as ∆S = ∆Q/T. It takes energy to decrease entropy. Think about how easy it is for your room to become messy but how hard it is to reverse it! Entropy encompasses two out of the three laws of thermodynamics.

The differences between solids, liquids, and gases is that the volume each phase increases in the order of solid<liquid<gas. Water, however, expands just before it freezes. Between 0°C and 4°C, water takes on a slightly different molecular structure that pushes the molecules farther apart. Kind of a good thing too because what would have happened during colder climate conditions if ice (solid water) sank to the bottom of the pond?

Problems (basic) like the ones in the video and using Q equation.

Problems calculating heat for phase changes and more Q equation examples.

Answers to the problems above.

Hess's Law examples with step-by-step instructions.

Ch 6: 13, 15, 17, 27, 29, 31, 43, 45, 51, 53, 55, 57, 59, 69, 71, 73, 75, 79, 83, 85

Chemistry is a subject that must be practiced everyday if possible. Work through the lecture examples stopping the video clips and then restarting to check yourself. Take advantage of the practice outlined above to give you the practice you need to be successful. DO NOT PROCRASTINATE! Check announcements on Blackboard everyday. Keep a printed copy of the most recent course calendar (included in syllabus) next to your work area. Email me with questions!!

Email: melinda.oliver@faulknerstate.edu
Location: Faulkner State Community College, Fairhope Avenue, Fairhope, AL, United States
Phone: (251)405-4504