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Section 3: Energy Changes in Chemistry
Chemicals in a laboratory can possess potential or kinetic energy, just as a baseball or a person sitting on a bicycle can. Chemical energy is potential energy that is stored in chemical bonds. Different types of bonds store different amounts of energy, which can be released by exchanging high-energy bonds for low-energy bonds. For example, the bonds that hold together a molecule of gasoline are rearranged when the molecule of fuel combusts; the new chemical bonds are formed in the products—CO2 and H2O. Because of the different energies associated with the bonds in the products and the reactants, this reaction will release energy to its surroundings. The engine uses the chemical energy released by the reaction to move the car. (Section 9 will go into more detail on the energies associated with individual bonds.)
As we have seen in our discussions of phase changes, molecules of chemical substances also have kinetic energy. They are in constant, random motion; as a substance heats up, its molecules move faster and faster. The total energy of all the molecules in a substance is called its "thermal energy."
Figure 7-4. Comparing Thermal Energy
Imagine that two flasks of equal size are left on a table, one holding a liter of water and the other holding two liters of water. Once the water has reached room temperature, the water will be the same temperature in the two flasks. But the larger quantity of water will have more thermal energy than the smaller quantity, because two liters of water contain more particles than one liter. Twice as much ice would have to be added to the beaker on the left to cool down the water to the same temperature as the beaker on the right.
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A substance's thermal energy is related to its temperature, which is the average kinetic energy of the particles in the substance. As the substance becomes hotter, its temperature increases and so does its thermal energy. But since thermal energy is the total amount of kinetic energy of all the particles, it depends on the number of the particles in the object, while temperature does not. This can be visualized by imagining two different volumes of water that are the same temperature. (Figure 7-4)
Thermal energy is often incorrectly referred to as heat. Heat is actually the energy that flows from a high temperature object to a lower temperature object when objects are placed in thermal contact. We can say that a certain amount of heat flows from one metal to another, but it is not correct to say that a piece of metal contains a certain amount of heat.
Through the 18th century, scientists thought that heat was a kind of fluid that could flow between substances. Early chemists, including pioneers such as the French nobleman Antoine Lavoisier (1743–1794), called this purported heat fluid "caloric" from the Greek word for "heat."
Figure 7-5. Boring a Cannon
Boring a cannon generated significant amounts of heat. Observing the process, Sir Benjamin Thompson theorized that heat was from the motion, not from a physical substance.
© Wikimedia Commons, Public Domain.
Massachusetts native Sir Benjamin Thompson (1753–1814), later known as Count von Rumford, helped disprove the caloric theory. In a famous experiment, he boiled water with the thermal energy generated by boring a cannon—drilling a hole through a cylinder of solid brass to create the gun barrel of the cannon (the tube through which cannon balls were fired was called a "bore"). (Figure 7-5)
The friction of metal grinding against metal generated so much heat that it warmed a bath of water in which the cannon barrel was submerged. After 2.5 hours, the water boiled, to observers' amazement. Thompson reasoned that if caloric was a physical substance inside the cannon barrel, it should eventually run out. But boring out the cannon seemed to produce an inexhaustible supply. "[A]nything which any insulated body, or system of bodies, can continue to furnish without limitation cannot possibly be a material substance," Thompson concluded. The experiment showed that thermal energy had a "mechanical equivalent." Heat was not a fluid contained in matter. Rather, it was a result of the motion.
Scientists began measuring heat in calories in the 1820s, when French physicist and chemist Nicolas Clément (1779–1842) introduced the unit. He defined it as the amount of heat needed to increase the temperature of one kilogram of water by one degree Celsius. Today, food labels continue to use calories to indicate the energy content of food.1 There is also a specific term for what this physical quantity is often called: the "heat capacity" or the "specific heat" for a substance. It is defined as the amount of heat it takes to raise one gram of a substance by one degree. For water, the heat capacity is 1 calorie/g °C. It turns out that water actually has a very high heat capacity, which is why tanks of water are often used as storage for heat energy collected from solar panels.
Modern chemists use a different unit: the joule, after James Prescott Joule (1818–1889), a British physicist. One joule is equal to 4.18 calories. A joule is a small unit: Driving a small, efficient car one mile burns enough gasoline to release 3 million joules. A nuclear reactor can produce 300 trillion joules in one day. Consequently, large increments of energy are usually measured in megajoules (one million joules), gigajoules (one billion joules), or even larger units.
1Clément's calorie is properly referred to as a "kilocalorie" or "large calorie," and is written with a capital C to distinguish it from a "small calorie"—the amount of heat required to raise the temperature of one gram of water by 1°C.