Heat is a form of energy, sometimes called thermal energy, which can pass spontaneously from an object at a high temperature to an object at a lower temperature. If the two objects are in contact, they will, given sufficient time, both reach the same temperature.
The specific heat can be considered to be the amount of heat required to raise the temperature of one gram of the substance by 1oC (if you make m and Dt in Equation 1 both equal to 1, then q will equal S.H.). Amounts of heat are measured in joules (historically: calories). To raise the temperature of 1 g of water by 1oC, 4.18 joules of heat must be added to the water. The specific heat of water is therefore 4.18 joules/goC. Since 4.18 joules equals 1 calorie, we can also say that the specific heat of water is 1 calorie /goC. Ordinarily heat flow into or out of a substance is determined by the effect that the flow has on a known amount of water. Because water plays such an important role in these measurements, the calorie, which was the unit of heat most commonly used until recently, was actually defined to be equal to the specific heat of water.
The specific heat of a metal can readily be measured in a calorimeter. A weighed amount of metal is heated to some known temperature and is then quickly poured into a calorimeter that contains a measured amount of water at a known temperature. Heat flows from the metal to the water, and the two equilibrate at some temperature between the initial temperatures of the metal and the water. The amount of heat that flows from the metal as it cools is equal to the amount of heat absorbed by the water and the calorimeter. In thermodynamic terms, the heat flow for the metal is equal in magnitude but opposite in direction, and hence in sign, to that for the water. For the heat flow q,
If we now express heat flow in terms of Equation 1 for both the water and the metal M, we get
In this experiment we measure the masses of water and metal and their initial and final temperatures. (Note that Dtmetal <0 and Dt H2O > 0, since At Dt = tfinal - tinitial.) Given the specific heat of water, we can find the positive specific heat of the metal by Equation 3. We will use this procedure to obtain the specific heat of an unknown metal. The specific heat of a metal is related in a simple way to its molar mass. Dulong and Petit discovered many years ago that about 25 joules were required to raise the temperature of one mote of many metals by 1oC. This relation, shown in Equation 4, is known as the Law of Dulong and Petit:
where MM is the molar mass of the metal. Once the specific heat of the metal is known, the approximate molar mass can be calculated by Equation 4. The Law of Dulong and Petit was one of the few rules available to early chemists in their studies of molar masses.
HEAT OF REACTION
When a chemical reaction occurs in water solution, the situation is similar to that which is present when a hot metal sample is put into water. With such a reaction there is an exchange of heat between the reaction mixture and the solvent, water. As in the specific heat experiment, the heat flow for the reaction mixture is equal in magnitude but opposite in sign to that for the water. The heat flow associated with the reaction mixture is also equal to the enthalpy change, AH, for the reaction, so we obtain the equation:
By measuring the mass of the water used as solvent, and by observing the temperature change that the water undergoes, we can find q H2O by Equation 1 and DH by Equation 5. If the temperature of the water goes up, heat has been given off by the reaction mixture, so the reaction is exothermic; q H2O is Positive and DH is negative. If the temperature of the water goes down, the reaction mixture has absorbed heat from the water and the reaction is endothermic. In this case q H2O is negative and DH is positive. Both exothermic and endothermic reactions are observed.
When this reaction occurs, the temperature of the solution becomes much higher than that of the NaOH and water that were used. If we dissolve a known amount of NaOH in a measured amount of water in a calorimeter, and measure the temperature change that occurs, we can use Equation 1 to find qH2O for the reaction and use Equation 5 to obtain DH. Noting that DH is directly proportional to the amount of NaOH used, we can easily calculate DHSolution for either a gram or a mole of NaOH. In the second part of this experiment you will measure DHSolution for an unknown ionic solid.
Chemical reactions often occur when solutions are mixed. A precipitate may form, in a reaction opposite in direction to that in Equation 6. A very common reaction is that of neutralization, which occurs when an acidic solution is mixed with one that is basic. In the last part of this experiment you will measure the heat effect when a solution of HCl, hydrochloric acid, is mixed with one containing NaOH, sodium hydroxide, which is basic. The heat effect is quite large, and is the result of the reaction between hydrogen ions in the HCl solution with hydroxide ions in the NaOH solution;
1. Heat Capacity of the Calorimeter
2. Heat of Neutralization
3. Specific Heat of a Metal
Fill a 400-cml beaker two-thirds full of water and begin heating it to boiling. While the water is heating, weigh your sample of unknown metal (~30 grams) to the nearest 0. 1 g on a top loading balance. Weigh the empty test tube; pour the metal into the test tube. Now weigh the test tube and metal. Put the test tube and metal into the beaker containing the boiling water. Continue heating the metal in the water bath for at least 10 minutes after the water begins to boil to ensure that the metal attains the temperature of the boiling water. The water level in the beaker should be high enough so that the top of the metal is below the water surface. Add water as necessary to maintain the water level.
While the water is boiling, weigh the calorimeter to 0.1 g. Place about 40 cm3 of water in the calorimeter and weigh again. Insert the stirrer and thermistor into the cover and put it on the calorimeter.
Using LabWorks II, measure the temperature of the water in the calorimeter for about 3 -4 minutes before adding the hot metal. Take the test tube out of the beaker of boiling water, remove the stopper, and pour the metal into the water in the calorimeter. Be careful that no water adhering to the outside of the test tube runs into the calorimeter when you are pouring the metal. Replace the calorimeter cover and agitate the water as best you can with the glass stirrer. Continuously record the temperature of the mixture until the temperature again becomes relatively constant. Repeat the experiment, using about 50 cm3 of water in the calorimeter. Be sure to dry your metal before reusing it; this can be done by heating the metal briefly in the test tube in boiling water and then pouring the metal onto a paper towel to drain. You can dry the hot test tube with a little compressed air.
Place about 50 cm3 of distilled water in the calorimeter and weigh the same as in the previous Procedure. In a small beaker weigh out about 4-grams of the solid compound assigned to you. Make the weighing of the beaker and of the beaker plus solid to 0.1 g. Using LabWorks II measure the temperature of the water for about 3 -4 minutes before adding the unknown salt. The temperature should be within a degree or two of room temperature. Add the unknown salt to the calorimeter. Stirring continuously and occasionally swirling the calorimeter, record the temperature until the final temperature remains relatively constant. Check to make sure that all the solid dissolved. A temperature change of at least 5 degrees should be obtained in this experiment. If necessary, repeat the experiment, increasing the amount of solid used.
- Computer-based measurements (must be completed during the laboratory period).
i. Opening LabWorks II .
A view of LabWorks II's Main Menu window should be on your screen.
ii. Calibration Set-up.
iii. Design Set-up.
iv. Acquire Set-up.
v. Analysis Setup View.
The data you have collected will be saved on the 31/2 Floppy (A:) drive or to the desk top for temporary strage. To finish this experiment, you will now analyze this information.
We must accomplish a number of tasks:
Make certain the Spreadsheet and Graph view is on the computer screen. To print your graph, click on the File button found in the menu toolbar.