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CHM 1020--Chemistry for Liberal Studies--Spring 1999

Chemistry 1020—Lecture 11—Notes

One more bond energy problem. Lets take a look at "Your Turn—4.7"

C2H5OH + 3 O2 --> 2 CO2 + 3 H2O

From Lewis structures: (You draw them this time)


5 C-H bonds = 5 (411 kJ) = 2055 kJ

1 C-C bond = 346 kJ = 346 kJ

1 C-O bond = 358 kJ = 358 kJ

1 O-H bond = 459 kJ = 459 kJ

3 O=O bonds = 3 (494 kj) = 1482 kJ

Total 4700 kJ


4 C=O bonds = 4 (799 kJ) = 3196 kJ

6 H-O bonds = 6 (459 kJ) = 2754 kJ

Total 5950 kJ

Energy change = 4700 kJ – 5950 kJ = -1250 kJ

Again, this is per mole. We can use this value to calculate many other relationships:

What about kJ produced per gram of ethanol?

1250 kJ/mole ethanol x 1 mole ethanol/46.0 g ethanol = 27.2 kJ/g ethanol

What about kJ produced per mole of CO2 produced?

1250 kJ/mole ethanol x 1 mole ethanol/2 mole CO2 = 625 kJ/mole CO2.

What about moles CO2 produced per kJ of energy generated?

1 mole CO2/625 kJ = 0.0016 mole/kJ (or 1.6 mole CO2/J). This would be one way of comparing the effect on the greenhouse gas production for a given amount of energy, for example.

When we speak of energy produced or energy consumed in a process, we use a formalism that refers to the change in the energy state or condition of the material undergoing change. In effect, we are talking about the difference in the energy content of the products and that of the reactants:

Energyproducts – Energy reactants = D E (or delta E if your browser can’t use Greek symbols). For an exothermic process, Eproducts<Ereactants, so that D E is negative.

We say the energy is lost by the system and passed to the surroundings.

system that portion of the universe we are studying. (usually, for example, the components of a chemical reaction confined in a test tube or other reaction site).

surroundings the rest of the universe.

We study the exchange of heat, work, and matter between a system and its surrounding. When no matter is exchanged (what we call a closed system), then one way of stating the first law of thermodynamics is to relate the internal energy change (E) in a process with the exchange of heat (q) and work (w):

Efinal-Einitial = D E = q + w

Where q and w are positive if the heat or work is done on the system (hence internal energy increases) or negative if done by the system on the surroundings (hence internal energy decreases).

If the process does no work, we measure D E by measuring heat produced or taken up. When a process involves a change in volume, there is some work of expansion or contraction done, and in this case the heat produced would be slightly different from the change in internal energy. We call that change enthalpy, using the letter H (or D H). There is a fine distinction between D E and D H, but for your purposes you can think of D H as heat produced or consumed in a process (a process carried out at constant pressure, where the only possible work is work of expansion or contraction).

So the energy changes we have been calculating from bond energies are really D H values.

Now the first law of thermodynamics can tell us how much energy is produced or consumed. But it doesn’t tell us whether the process will occur or not.

Our experience suggests that most processes occur which release energy (reaction of hydrogen and oxygen)—you may need a spark to get it started, and reactions which consume energy will not take place by themselves. Water will not spontaneously decompose into hydrogen and oxygen. You must do work on water—run an electric current through it—in order to get it to decompose.

But there are exceptions. When you dissolve some salts in water, the solution gets cooler instead of warmer. Ice will melt at temperatures above 0oC, but the melting is an endothermic process. Therefore there is something besides energy change that determines whether a reaction will occur.

We are familiar with the concept that some events are spontaneous and some are non-spontaneous.

  • Water runs downhill—not uphill.
  • Heat travels from cold to hot, not from hot to cold.
  • A gas will expand into a vacuum, it will not contract to form a vacuum.
  • You can mix alcohol and water. They will not spontaneously separate.

How many others can you think of? Have you ever seen a movie reel run backwards (for example, a cue ball breaking the set of pool balls. Confetti dropped on a New York parade reassembling itself into the paper it was shredded into).

These are all manifestations of the notion that time has a directionality (called the arrow of time). And they are manifestations of what we call the second law of thermodynamics.

There are many ways of stating the second law. They all fall under the guise of saying that things will spontaneously become more disorganized. It takes extra work to organize something—energy invested in the organization itself. This organization is a property of substances that we call entropy, and one statement of the second law is that the entropy of the universe increases in any spontaneous process.

A consequence of the second law puts a restriction on our interchange of energy. Work can be changed completely into heat, but heat cannot be changed completely into work.

Heat can only flow from a high temperature to a lower one. To convert it into work, such as in the steam engine, you put heat in at a high temperature, and take it out at a lower temperature, so only part of the energy is used to produce work. That puts a limit on the maximum efficiency of any process involving heat to work transformation;

Efficiency = Thigh-Tlow/Thigh

This limitation applies to the steam engine, the automobile engine, and power plants running off of the heat energy produced in combusting fuels. (Hence automobile engines, for example, are more efficient the higher the temperature at which they operate).

So in producing electrical energy in a power plant, for example, the diagram in Figure 4.14 indicates the energy conversion steps:

Chemical energy--> Heat energy--> Mechanical Energy--> Electrical Energy

Each conversion involves some loss due to inefficiency, but the big loss comes when heat is an intermediate form.

It is the second law of thermodynamics which makes impossible the perpetual motion machine (for example, let your automobile engine run a generator that makes electricity that is used to run the engine).

So for any fuel we use, we don’t get all the energy use of it, especially if heat is an intermediate.

Lets take a look at the pro’s and con’s of some of our more common fuels.

From what you have read in the chapter, can you list advantages and disadvantages of each of the following:



Automobile Engines

Power Plants








Natural Gas











(Use this as a take-home exercise to fill in as many things in this table as you can).

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