Thermodynamics


The origins of thermodynamics lies primarily in the 19th century, the Steam Age. It was James Watt who invented the steam engine. The cost of fuel (i.e. coal) and manpower had to be calculated to determine the efficiency and performance of steam engines. In other words, how much work can you get for your dollar? Efficiency is equal to work output divided by the heat required to do work. The field of thermodynamics is highly mathematical, and you really need to have a strong understanding of calculus to appreciate it -- so, I won't go into that aspect. For our purposes you only really need to know just a few of the following concepts.

A system is any part of the universe that is separated form the rest of the universe by physical or conceptual boundaries. Physical boundaries might include an egg shell, a coffee cup, a balloon, and a calorimeter (a calorimeter is used to measure heat change). Conceptual or imaginary boundaries would include the "sphere" around the Earth, and the Sun's surface.This will become clear in a moment. The surroundings is defined as the "rest of the universe" -- e.g. an ocean.

There are three types of systems. In a closed system, energy can be exchanged with its surroundings, but matter cannot. An example of this would be a sealed, clear plastic bottle of water -- i.e. light energy can pass through, but water cannot get out, until the cap is taken off (which would then be an open system). In an open system both matter and energy can be exchanged with the surroundings. As well as an opened bottle of water, other examples include cells (i.e. all life forms) and the Earth. The Earth is somewhat of a "quasi-closed system" -- i.e. matter can be exchanged with Earth's surroundings in the form of meteorites, dust particles, and gases, while energy comes in from the Sun and leaves the planet in the form of infrared radiation. On the other hand, matter doesn't leave the Earth on a significant scale. An isolated system can exchange neither matter nor energy with its surroundings. A "thermos" might be an example, but it is not completely isolated, since energy will eventually dissipate from it. The Universe itself is, perhaps, the only true isolated system (assuming that there is only one universe, as opposed to a "multiverse" with the possibility of different universes interconnected to one another).

There are three laws in thermodynamics. The First Law of Thermodynamics is the law of conservation of energy. This means that energy is neither created nor destroyed, and can only be transferred or transformed (ultimately, all energy is converted into heat). In other words, the total amount of matter and energy is before and after a reaction is the same. The Second Law of Thermodynamics states that the entropy of a system increases during a spontaneous change. The Third Law of Thermodynamics states that at absolute zero (T=0 K, where K=Kelvin), all thermal motion stops (i.e. the entropy of a perfect crystal is zero). On the Kelvin temperature scale, absolute zero is the lowest (theoretical) temperature possible. The gradations of degrees are equivalent to the degrees in the Centrigrade scale. So, the freezing point of water is 0 degrees Celsius (C) or 273.15 Kelvin (or 273.15 K). The boiling point of water is 100 degrees C or 373.15 K (i.e. just add 100 to 273.15 K). In the physical sciences, the Kelvin scale is used, while in biology the Centigrade or the Kelvin scales are used (which ever one is most appropriate). Incidentally, the average human body temperature is 37 degrees Celsius, or 310.15 K.

There are some other terms you should be aware of:

  • Heat (q) is a form of energy related to mass. For life forms, heat is an "unuseable" energy form -- i.e. it cannot be converted back into some other form by life forms, and so it is dissipated into the surroundings. As heat increases, molecules move faster, and as heat decreases, molecules tend to move more slowly. Temperature (T) is an estimate of heat (or of the average molecular motion), and it is not related to mass. Heat "flows" from a higher temperature to a lower temperature. The total energy of a system is termed internal energy (U). Energy is the capacity to do work (w). An example of work, would be a process that can move an object to different heights in the surroundings (e.g. work equals force times the distance). Heat, work, and internal energy are all measured in Joules (J). One Joule equals one kilogram meter squared per second squared (or 1 J = 1 kg m^2/s^2). And, one calorie is equal to 4.184 Joules (1 cal = 4.184 J).
  • An exothermic (or "exergonic") chemical reaction is a process that releases energy as heat. For example, combustion reactions (i.e. burning fossil fuels) are exothermic reactions. An endothermic reaction is a process that absorbs energy (i.e. it requires some energy input to the system). An example of this would be the vaporization of water.
  • Enthalpy (H) is the total heat content of a system. Entropy (S) is often equated with chaos, disorder, or randomness. Negentropy is the reciprocal of entropy (i.e. 1/S). Free Energy, otherwise known as Gibb's Free Energy (G), is the amount of energy that is available to do work.

The energy of a system can change over time. To refer to a "change in" something the symbol "delta" is used in front of some quantity (I don't actually have a symbol to show you but, it is essentially a triangle -- I will use the name "delta" for my own convenience). For example, delta-H means a change in total heat content, delta-S is a change in Entropy, and delta-G is a change in free energy. Reactions may be characterized as follows:

delta-H

-ve (exo)

-ve (exo)

+ve (endo)

+ve (endo)

 delta-S

+ve

-ve

+ve

-ve

 delta-G

-ve

-ve
+ve
-ve
+ve
+ve

 Process

Spontaneous at all T

Spont. at low T
Non-spont. at high T
Spont. at high T
Non-spont. at low T
Non-spont. at all T
rev-reaction is spont.

A positive value of delta-H indicates that the reaction is endothermic (i.e. energy must be added to the system), while a negative value indicates that the reaction is exothermic (i.e. energy is released from the system). A positive value of delta-S is an increase in entropy (e.g. an increase in randomness, and a decrease in orderliness), and a negative value is a decrease in entropy (e.g. a decrease in randomness, and an increase in orderliness). The value of delta-G indicates the amount of free energy available at certain temperatures. Reactions can go forwards and backwards (which can refer to the speed of the reaction). A spontaneous reaction is one that "goes forward" as long as the conditions shown in the table are met. A non-spontaneous reaction means that the reverse reaction is spontaneous. At equilibrium, both the forward and the reverse reactions are going on at equal rates (i.e. there is no net direction for the reaction).

A few more terms and we'll be done with this section. An isothermal process means that the temperature is kept constant (i.e. delta-T = 0; q = -w). An isobaric process is one where the pressure is kept constant (i.e. neither q nor w equals zero; q = delta-H), and an isochoric process is one that has constant volume (i.e. w = 0; delta-U = q). An adiabatic reaction is one where there is no heat flow (i.e. q = 0; delta-U = w).

Consider an adiabatic process, where dq = 0 (i.e. d is "the partial derivative of"). Delta-S can be greater than or equal to zero. If the process is reversible, delta-S = 0. If the process is irreversible, then delta-S > 0. The biggest conceivable/known system is the Universe itself. Consider the Universe as a whole, in which case, all processes within it are by definition, adiabatic (i.e. no heat flows out from or into the universe). Therefore, delta-S for the Universe is greater than or equal to zero for any process. In other words, delta-S(universe) = 0 for all reversible processes, while delta-S(universe) > 0 for all irreversible processes. The total entropy of the Universe then, is delta-S(universe) = delta-S(system) plus delta-S(surroundings). Every real process increases the entropy (i.e. randomness or disorder) of the Universe. The Universe favors reactions that lead to the lowest energy level, which is a random distribution of particles. Thus, the Universe favors randomness, simplicity, and exothermic reactions (i.e. delta-H is less than zero, and delta-S is greater than zero).

Life prefers a different set of conditions than the universal trend to randomness. Life favors order, complexity, and endothermic reactions (i.e. delta-H is greater than zero, and delta-S is less than zero). If the Universe favors randomness, then how does life even exist? The Universe is essentially patchy with respect to order and disorder. For life, equilibrium means death. Thus, living things are not at equilibrium with their surroundings. A living cell equals order, and therefore has a low entropy (a low entropy, but not a zero entropy). Metabolic processes of living cells decrease the entropy of cells, but there is a larger increase in the entropy of the surroundings that is associated with those processes. Therefore, cells can create order locally, and at the same time increase the disorder of their surroundings. Cellular reactions also generate heat, so overall, cells are exothermic. Cells may then be described as "non-equlibrium, dynamic, steady-state systems." That is, cells never achieve equilibrium until they die; cells are in continual motion (i.e. dynamic, with respect to molecular turn-over); and, cells have no net change in their chemical composition (i.e. steady-state). Thus, life follows all the rules of Thermodynamics. If you don't believe that, then look at the buildings around you. Buildings are orderly things. The Universe tends towards disorder, yet these bulidings still exist. Why? Because energy was used to build and maintain them. Thus, orderly things can and do exist in a Universe that tends towards disorder. In other words, orderly things exist without violating the laws of thermodynamics.

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