In our studies thus far we have written simple equations, predicted the products of simple reactions, balanced equations and performed quantitative calculations using balanced chemical equations. There is one important characteristic of a chemical reaction which we have not yet included in a chemical reaction. The heat which is involved in the chemical reaction.

We have seen several reactions which we could classify as violent. Recall the reaction between aluminum and bromine. Although the reaction was slow to start, once it began it proceeded rapidly. We saw bromine vapor bellowing out of the beaker and pieces of aluminum glowed red hot as they danced about on the surface of the liquid bromine. The reaction of potassium in water was another violent reaction.

In class we looked at two new reactions to set the stage for this discussion.

The first reaction is between aluminum and iron (III) oxide (Fe2O3). The reaction will be initiated by adding a few drops of concentrated sulfuric acid to a mixture of sucrose and potassium chlorate. This 'side' reaction will produce enough heat to start the reaction between aluminum and iron (III) oxide.

The reaction is;

Al(s) + Fe2O3(s) ---> Al2O3(s) + Fe(l)

The reaction produces so much heat the product iron changes phase from solid to liquid. The heat produced is enough to change the temperature from 25 C to at least 1537 C (the melting point of iron).

The second reaction is between barium hydroxide and ammonium chloride. Both of these reactants are white crystalline solids, as shown here. The two white solids are mixed together in a beaker which has been placed on a piece of wood (2 x 4) upon which several drops of water have been placed. Before the reaction begins the two solids are at room temperature as is the beaker, 2 x 4 and the water. Show how the beaker does not adhere to the wood before mixing the reactants.

The first observation is the two solids form a slurry, as a liquid is formed. This is very interesting, but not surprising when you are told water is a product of the reaction. The second observation occurs when the beaker is picked up. The block adheres to the bottom of the beaker. In fact the beaker and the block of wood are frozen together. The temperature has fallen when the two reactants are mixed to the extent the water between the bottom of the beaker and the block of wood has frozen.

The reaction is;

Ba(OH)2.8H2O(s) + 2NH4SCN(s) ----> Ba(SCN)2(aq) + 2NH3(g) + 10H2O(l)

As the two white solids were mixed afew students in the front of the class observed how the temperature of the reaction changed. The temperature feel is got cooler. We also notice the change which occurred in the flask. The two solids changed to a liquid. If we could smell the reaction vessel it would smell of ammonia.

So what have we discovered in these two examples? These two reactions were chosen because the particular change I wanted you to note was very evident. In the case of the reaction between aluminum and iron(III) oxide the reaction produced heat, in the form of a flame. Heat was released when the reaction began. In the case of the barium hydroxide and ammonium thiocyanate instead of getting hotter the reaction mixture became cooler. In both of these reactions energy was being transferred in the form of heat. The study of the energy changes when a reaction occurs is called thermochemistry. Thermochemistry is part of thermodynamics,the study of heat, energy, and work and and their transformations.

When discussing heat transfer it is a convenience to identify where the heat flows from and where it goes to. We will use the terms system and surroundings to help focus on how heat flows. System is that portion of the universe we single out for study and which is bounded by some boundary as defined by a container or by the sample itself. The surrounding is everything else in the universe. If in a chemical change or reaction, energy is released and the temperature of the system will get warmer than the surroundings, than heat will flow from the system to the surroundings. This was the case for the reaction between aluminum and iron(III) oxide.

In the second example the reaction occurred in the beaker. That was the system. Had we touched the reaction container (the system) is would have felt cool. Heat would be removed from our hand, which is at a higher temperature than the system. Heat must flow from the surrounding into the system.

A chemical reaction which releases heat is called an exothermic reaction. Heat flows from the system to the surroundings. An endothermic reaction absorbs heat from the surroundings. To begin our study of thermodynamics (and we will encounter thermodynamics many different times during this year) we are going to begin by covering thermochemistry. As an introduction to thermochemistry I need to define several important terms so that we might better understand the two reactions we observed earlier. Those terms are energy, temperature, heat and work.

Because energy is not tangible, as are material objects, it become difficult to define it completely. The scientific definition of energy is the capacity to do work or to transfer heat. In the first reaction we saw, between aluminum and iron(III) oxide energy was given off in the form of heat. No work was done by the reaction, with the exception of pushing back the atmosphere. So no useful work. There appears to have been energy present in the reactants which is liberated when the products are formed, as heat when the reactants were combined.

In the second reaction energy was absorbed when the reaction occurred. This was evident when I touched the container. It felt cold to the touch. Heat was transferred from my hand to the flask. Again no useful work was done.

In either case we could have utilized the heat transfer to do some work. When we discuss energy two forms of energy come to mind potential energy and kinetic energy. Potential energy (U = mgh: g = 9.8 m.sec-2) is energy stored in an object by virtue of its position. A book held above my head has more potential energy than a book held at my side near my waist. The amount of potential energy an object has depends on the mass of the object and its height above the earth's surface. If I drop the book the potential energy is converted to kinetic energy. At the atomic level two nitrogen atoms have a higher potential energy when they are far apart then when they are bonded together, just as the book has a higher potential when separated from the surface of the earth. To separate the two nitrogen atoms in a mole of N2 molecules requires 960 kJ. This amount of energy would raise a 100 g ball to a height of 600 miles.

Kinetic energy is energy of motion. The magnitude of the kinetic energy of an object depends on its mass and velocity. Ek = 1/2 mv2 The heavier and faster an object is going the more energy it has and the more work it can do on whatever.

If we substitute the SI units for mass and velocity into the kinetic energy equation we have the correct units for energy. The SI unit for energy is the joule (J) (1 kg.m2.s-2 ). An object with a mass of 1 kg traveling at a velocity of 1 m.sec-1 has a kinetic energy of 1 J. One particular textbook author has indicated that dropping a six pack of soda pop cans on your foot, the kinetic energy at the moment of impact is between 4 and 10 joules. The kinetic energy of one gas molecule at 25 degree C is about 6 x 10-21 J. A mole of this gas has about 4 kJ of kinetic energy. If that energy could be harnessed it could be used to raise 350 copies of our text about 1 meter.

A 100 watt light bulb produces 100 J of energy for every second it operates. (When the City of Stillwater sends a bill for the energy used at your house it includes an electrical bill for energy used in the form of electricity and the units used are kilowatt-hours. A single kilowatt-hour is equivalent to 1000 watts-3600 seconds or 3.6 x 106 watt-sec or 3.6 x 106 joules. A BTU is the amount of energy required to raise the temperature of one pound of water one degree Fahrenheit.) A calorie is another non-SI energy unit. It is defined as the heat required to raise the temperature of 1 g of water, 1 degreeC. 1 cal = 4.184 J.

If we had had the proper equipment we could have measured the temperature of both reactions. Temperature is a measure of the degree of hotness or coldness of an object. If I indicate the temperature of a sample of water is 95 C we know the sample is very hot. If another sample of water has a temperature of 1C, we know the sample will feel cool when we touch it.

Heat is energy that is transferred as a result of a temperature difference. Heat always flows from a warmer object to a cooler object. Heat causes a change in temperature. So when we 'heat' an object it get hotter. Lighting a bunsen burner and placing it beneath a beaker filled with water causes the temperature of the water to increase. Heat flows from the warmer flame of the bunsen burner to the cooler water. We can measure this temperature change using a thermometer. If we have two beakers, a 100 mL beaker and a 25 mL beaker each filled with water at the same initial temperature, and an equal amount of heat is added to each beaker we will find the water in the larger beaker to have a lower temperature compared to the temperature of the water in the small beaker. The temperature change depends on the amount of heat added to the water and it depends on the amount of matter present. When the bunsen burner is removed, the source of heat, the water in the beaker will cool, as heat flows from the warmer water in the beaker to the cooler air of the room, and return to room temperature. I have described heat as flowing or being transferred and this may be misleading. Heat is not matter, it is not contained in matter. Heat is a way to exchange energy.

As we saw in the example of heating the two samples of water, the addition of equal amounts of heat resulted in a different change in temperature because of the different masses of water. It is interesting that there is a relationship to the amount of energy added to a specified amount of water and the subsequent temperature change. This relationship is called specific heat. Adding energy, in the form of heat, to a sample of water will cause a change in temperature of the water. It is interesting that if 4.184 J of energy is added to 1 gram of water at 4 degrees C the temperature of the sample changes by one degree Celsius. This quantity of heat is called the specific heat of water. This is the amount heat required to change the temperature of 1 gram of water by 1 degree Celsius. The definition of specific heat is;

The specific heat of any substance can be determined using the equation. All that must be measured is how much heat is required to effect a change in temperature for a given amount of the substance.

Some additional specific heats can be found in Table 5.2 on page 159 of your textbook.

Water has a very high specific heat, that is, a large amount of heat is required to change the temperature of water. The greater the specific heat of a substance the more heat is required to effect a change in temperature of 1 gram of the substance.

Heat capacity is an alternative way to express the capacity of a substance, or an object, to absorb heat. But heat capacity has different units compared to specific heat. Heat capacity is the amount of heat required to raise the temperature of an object by 1 degree C. The difference between specific heat and heat capacity is that specific heat is heat required per gram, while heat capacity is heat required for an object whose mass is constant and not expected to changed. This difference will be important when we discuss calorimetry.

Once the specific heat or heat capacity of a substance is known it is possible to determine the amount of heat absorbed or lost in a process. The equation can be rearranged to solve for heat. We will use the symbol, q, to represent heat. Rearranging the equation, heat lose or gain = Specific heat . mass of substance . the change in temperature

This equation tells us the quantity of heat but we also need to know direction. If the sign of q is positive heat flows into the system. If q is negative heat flows out of the system.

Lets look at some sample problems;

Problem #1;

Problem #2;

Problem #3;

Go to Friday, October 10, 1997 Lecture