![]() ![]() A second way to view the internal energy of a system is in terms of its macroscopic characteristics, which are very similar to atomic and molecular average values. ![]() Because it is impossible to keep track of all individual atoms and molecules, we must deal with averages and distributions. Thus internal energy is the sum of atomic and molecular mechanical energy. Recall that kinetic plus potential energy is called mechanical energy. The internal energy U of a system is the sum of the kinetic and potential energies of its atoms and molecules. The first is the atomic and molecular view, which examines the system on the atomic and molecular scale. We can think about the internal energy of a system in two different but consistent ways. Internal energy is a form of energy completely different from either heat or work. However, both can change the internal energy U of a system. Heat transfer and work are both energy in transit-neither is stored as such in a system. Once the temperature increase has occurred, it is impossible to tell whether it was caused by heat transfer or by doing work. Heat transfer into a system, such as when the Sun warms the air in a bicycle tire, can increase its temperature, and so can work done on the system, as when the bicyclist pumps air into the tire. Nevertheless, heat and work can produce identical results.For example, both can cause a temperature increase. Work, a quite organized process, involves a macroscopic force exerted through a distance. Heat transfer, a less organized process, is driven by temperature differences. Heat transfer ( Q) and doing work ( W) are the two everyday means of bringing energy into or taking energy out of a system. The first law gives the relationship between heat transfer, work done, and the change in internal energy of a system. The first law of thermodynamics is actually the law of conservation of energy stated in a form most useful in thermodynamics. Making Connections: Law of Thermodynamics and Law of Conservation of Energy (See Figure 2.) We will now examine Q, W, and Δ U further. Heat engines are a good example of this-heat transfer into them takes place so that they can do work. Note also that if more heat transfer into the system occurs than work done, the difference is stored as internal energy. So positive Q adds energy to the system and positive W takes energy from the system. We use the following sign conventions: if Q is positive, then there is a net heat transfer into the system if W is positive, then there is net work done by the system. W is the net work done by the system-that is, W is the sum of all work done on or by the system. Q is the net heat transferred into the system-that is, Q is the sum of all heat transfer into and out of the system. Here Δ U is the change in internal energy U of the system. In equation form, the first law of thermodynamics is Δ U = Q − W. The first law of thermodynamics states that the change in internal energy of a system equals the net heat transfer into the system minus the net work done by the system. The first law of thermodynamics applies the conservation of energy principle to systems where heat transfer and doing work are the methods of transferring energy into and out of the system. ![]() If we are interested in how heat transfer is converted into doing work, then the conservation of energy principle is important. As the entire system gets hotter, work is done-from the evaporation of the water to the whistling of the kettle. The water in the kettle is turning to water vapor because heat is being transferred from the stove to the kettle. This boiling tea kettle represents energy in motion. The hot gasses (in the form of steam) have to release energy into the environment in the form of heat to cool to the point that they can form liquid water, meaning that the formation of H 2O is exothermic.Figure 1. This makes sense - H 2 and O 2 are gasses, while H 2O, the product, is a liquid. Since the sign is negative, we know that our reaction is exothermic. In our example, our final answer is -13608 J.Beware strongly exothermic reactions - these can sometimes signify a large release of energy, which, if rapid enough, can cause an explosion. The larger the number itself is, the more exo- or endo- thermic the reaction is. ![]() On the other hand, if the sign is negative, the reaction is exothermic. If the sign of your final answer for ∆H is positive, the reaction is endothermic. One of the most common reasons that ∆H is calculated for various reactions is to determine whether the reaction is exothermic (loses energy and gives off heat) or endothermic (gains energy and absorbs heat). Determine whether your reaction gains or loses energy. ![]()
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