From UCDavis Chemwiki
In previous sections we have indicated that heat is a form of energy and show how the quantity of heat energy absorbed or released by a chemical change can be related to the corresponding chemical equation. We also state the law of conservation of energy, and arguments in other sections have often been based on the idea that energy can neither be created nor destroyed. The law of conservation of energy is the first of three important laws involving energy and matter, which were discovered over a century ago. These laws were originally based on the movement or transfer (dynamics) of heat (thermo), and the law of conservation of energy is therefore referred to as the first law of thermodynamics.
We assign the symbol `DeltaH` and the name enthalpy change to the quantity of heat absorbed by a chemical or physical change under conditions of constant pressure. You may wonder just how heat energy could be absorbed or given off when atoms and molecules change position and structure during a chemical reaction, but we have not yet developed theories of chemical bonding, molecular structure, intermolecular forces, and molecular motion to the point where a satisfactory explanation can be given. We are in a position to investigate what can happen to molecules when matter absorbs or releases heat. One result of this study will be a clearer understanding of enthalpy. At the same time we will begin to appreciate what molecular factors contribute to making a reaction exothermic or endothermic.
The first law of thermodynamics (the law of conservation of energy) states that when heat energy is supplied to a substance, that energy cannot disappear-it must still be present in the atoms or molecules of the substance. Some of the added energy makes the atoms or molecules move faster. This is called translational energy. In the case of molecules, which can rotate and vibrate, some of the added energy increases the rotational and vibrational energies of the molecules. Finally, any atom or molecule will have a certain electronic energy which depends on how close its electron clouds are to positively charged nuclei.
The total of translational, rotational, vibrational, and electronic energies is the internal energy of an atom or molecule. When chemical reactions occur, the internal energy of the products is usually different from that of the reactants, and the difference appears as heat energy in the surroundings. If the reaction is carried out in a closed container (bomb calorimeter, for example), the increase in internal energy of the atoms and molecules is exactly equal to the heat energy absorbed from the surroundings. If the internal energy decreases, the energy of the surroundings must increase; i.e., heat energy is given off.
When a chemical reaction occurs at constant pressure, as in a coffee-cup calorimeter, there is a change in potential energy of the atmosphere (given by `PDeltaV`) as well as a change in heat energy of the surroundings. Because the heat energy absorbed can be measured more easily than `PDeltaV`, it is convenient to define the enthalpy as the internal energy plus the increased potential energy of the atmosphere. Thus the enthalpy increase equals the heat absorbed at constant pressure.
Enthalpy changes for a variety of reactions may be calculated from standard enthalpies of formation. They may also be estimated by summing the bond enthalpies of all bonds broken and subtracting the bond enthalpies of all bonds formed. Because the dissociation enthalpy for the same type of bond varies from one molecule to another, the second method is not as accurate as the first. However, it has the advantage that enthalpy changes for reactions of a particular compound can be estimated even if the compound has not yet been synthesized.
The enthalpy change for a reaction depends on the relative strengths of the bonds broken and formed and on the relative number of bonds broken and formed. A good fuel is a substance which can combine with oxygen from the air, forming more bonds and/or stronger bonds than were originally present. The fossil fuels, coal, petroleum, and natural gas consist mainly of carbon and hydrogen. When they burn in air, strong O—H and C=O bonds are formed in the resulting H2O and CO2 molecules. The supply of fossil fuels is limited, and they constitute a nonrenewable resource. Coal supplies ought to last another century or two, but petroleum and natural-gas supplies will be essentially depleted in half a century or less. During the next few decades it will be possible to gasify or liquefy coal to extend our supply of gaseous and liquid fuels. Conservation of these fuels can also make a major contribution toward continuing their use. Eventually, however, it will be necessary to develop nuclear or solar energy or some unknown source of energy if we are to continue our current energy-intensive way of life.
Energy
Energy is usually defined as the capability for doing work. For example, a billiard ball can collide with a second ball, changing the direction or speed of motion of the latter. In such a process the motion of the first ball would also be altered. We would say that one billiard ball did work on (transferred energy to) the other.
Energy due to motion is called kinetic energy and is represented by Ek. For an object moving in a straight line, the kinetic energy is one-half the product of the mass and the square of the speed:
`Ek =1"/"2m u^2` (1)
Where
`m` = mass of the object
`u` = speed of object
If the two billiard balls mentioned above were studied in outer space, where friction due to their collisions with air molecules or the surface of a pool table would be negligible, careful measurements would reveal that their total kinetic energy would be the same before and after they collided. This is an example of the law of conservation of energy, which states that energy cannot be created or destroyed under the usual conditions of everyday life. Whenever there appears to be a decrease in energy somewhere, there is a corresponding increase somewhere else.
Potential Energy is energy that is stored by rising in height, or by other means. It frequently comes from separating things that attract, like rising birds are being separated from the Earth that attacts them, or by pulling magnets apart, pulling an electrostatically charged balloon from an oppositely charged object to which it has clung, or many other forms (nuclear energy, chemical energy, etc.).
Energy is usually measured in the units of Joules or Kilojoules. Another unit of energy still widely used by chemists is the calorie. The calorie used to be defined as the energy needed to raise the temperature of one gram of water from 14.5°C to 15.5°C but now it is defined as exactly 4.184 J.
Energy
When a chemical reaction occurs, there is usually a change in temperature of the chemicals themselves and of the beaker or flask in which the reaction is carried out. If the temperature increases, the reaction is exothermic—energy is given off as heat when the container and its contents cool back to room temperature. (Heat is energy transferred from one place to another solely because of a difference in temperature.) An endothermic reaction produces a decrease in temperature. In this case heat is absorbed from the surroundings to return the reaction products to room temperature.Thermochemistry, a word derived from the Greek thermé, “heat,” is the measurement and study of energy transferred as heat when chemical reactions take place. It is extremely important in a technological world where a great deal of work is accomplished by transforming and harnessing heat given off during combustion of coal, oil, and natural gas.
The 1st Law of Thermodynamics
The First Law of Thermodynamics states that energy can be converted from one form to another with the interaction of heat, work and internal energy, but it cannot be created nor destroyed, under any circumstances. The equation that supports the First Law of Thermodynamics mathematically is:
`DeltaU=q+ w` here `DeltaU` is the total change in internal energy of a system, `q` is the heat exchanged between a system and its surroundings, and `w` is the work done by or on the system. Work is also equal to the negative external pressure on the system multiplied by the change in volume:
`w=-pDeltaV` where `P` is the external pressure on the system, and `DeltaV` is the change in volume.
The internal energy of a system would decrease if the system gives off heat or does work. Therefore, internal energy of a system increases when the heat increases (this would be done by adding heat into a system). The internal energy would also increase if work were done onto a system. Any work or heat that goes into or out of a system changes the internal energy. However, since energy is never created nor destroyed (thus, the first law of thermodynamics), the change in internal energy always equals zero. If energy is lost by the system, then it is absorbed by the surroundings. If energy is absorbed into a system, then that energy was released by the surroundings:
`DeltaU_"system"=-DeltaU_"surroundings"`
where `DeltaU_"system"` is the total internal energy in a system, and `DeltaU_"surroundings"` is the total energy of the surroundings.
Table 1
| Process | Sign of heat (q) | Sign of Work (w) |
|---|---|---|
| Work done by the system | N/A | - |
| Work done onto the system | N/A | + |
| Heat released from the system- exothermic (absorbed by surroundings) | - | N/A |
Table 2
| The Process | Internal Energy Change | Heat Transfer of Thermal Energy (q) | Work `w=-PDeltaV` | Example |
| `q = 0` Adiabatic | `+` | `0` | `+` | Isolated system in which heat does not enter or leave similar to styrofoam |
| `Deltav=0` Constant Volume | `+` | `+` | `0` | A hard, pressure isolated system like a bomb calorimeter |
| Constant Pressure | `+` o r `-` | enthalpy `DeltaH` | `-PDeltaV` | Most processes occur are constant external pressure |
| `DeltaT = 0` Isothermal | `0` | `+` | `-` | There is no change in temperature like in a temperature bath |
| Example 1 |
|---|
| A gas in a system has constant pressure. The surroundings around the system lose `62J` of heat and does `474J` of work onto the system. What is the internal energy of the system? Solution: To find internal energy, `DeltaU`, we must consider the relationship between the system and the surroundings. Since the First Law of Thermodynamics states that energy is not created nor destroyed we know that anything lost by the surroundings is gained by the system. The surrounding area loses heat and does work onto the system. Therefore, q and w are positive in the equation `DeltaU=q+w` because the system gains heat and gets work done on itself. `DeltaU=(62J) + (474J)` `DeltaU=536J` |
Subpages (1): Example 1