The first law of thermodynamics is basically a bookkeeping law. Its the law of conservation of energy - the change in internal energy of a closed system (ΔU) will be equal to the heat added to the system (Q) minus the work done by the system (W). In other words:
ΔU = Q - W
The internal energy is the sum total of all the energy of the molecules of the system. A system is any object or set of objects that we wish to consider. A closed system is one in which no mass enters or leaves the system (but energy may by exchanged from the environment). Thus the equation makes some intuitive sense - if heat is added to an object, you'd expect the energy of the molecules to increase. Similarly, if work is done by the object, that energy has to come from somewhere, and that somewhere must be the internal energy of the object - i.e. the object would cool down somewhat.
But now consider humans. We do work, don't we? Work in the physicist's sense, that is. We walk, run, eat and smile. We grow - a process that also subtly requires work. Now, if the first law of thermodynamics were true, the work that we did would result in the lowering of our body temperature. Yet everyone knows that our body temperatures are remarkably constant.
Perhaps we simple take in heat from our environment to compensate for its loss during our work? This would balance the equation nicely, but it simply isn't true. Our body temperatures are almost always higher than the environments (37ºC / 98ºF is usually much higher than the temperature the weatherman gives). Therefore on most days, we compound our thermodynamic problem by losing heat to the environment too! Even on hot days, when the environment is actually hotter than us, the body is unable to utilise this energy usefully anyway.
So what's going on? How do we resolve the paradox?
The answer is that humans are simply not an example of a closed system. We clearly can exchange mass with our environment - we can take in food and liquids (and excrete urine and faeces). We are what a physicist would call and open system - we may exchange both energy and mass with our environments. Since the food itself could briefly be regarded as a system, it has internal energy of its own - it has a temperature, etc. too, but it also has wonderful chemical potential energy in the bonds between its atoms.
And so when we eat, we are actually adding some internal energy from the environment (i.e. food) to our own system's internal energy - something that we couldn't do if we were a closed system. And this increased internal energy eventually goes into work, allowing us to keep living (which is always nice) and heat, keeping our body temperature constant.
So in an abstract and bird's-eye-view way, we know why we die if we don't eat or drink - our body simply runs out of (internal) energy to do work with. Also, the first law of thermodynamics explains why people with very low energy reserves (e.g. sick, underfed newborns and starving people) often drop their body temperatures. This is a very grave sign to a doctor, since it indicates that they don't even have enough energy stores to generate heat, and are probably very close to death.
The above also applies to evolution in its broadest sense. To a physicist (but not to a biologist!), evolution is very similar to growth in the sense that, over great lengths of time, simple molecules have become complicated organisms. Clearly, this process also requires energy, and so clearly thermodynamics applies. It's easy to see that, at least for the first law, evolution definitely does not violate thermodynamics - at all stages, the organisms could exchange matter with the environment, and thus avoid any thermodynamic complications. In the next post, I'll tackle the second law.
Reference: "Physics" (5th Edn.) - Giancoli
But now consider humans. We do work, don't we? Work in the physicist's sense, that is. We walk, run, eat and smile. We grow - a process that also subtly requires work. Now, if the first law of thermodynamics were true, the work that we did would result in the lowering of our body temperature. Yet everyone knows that our body temperatures are remarkably constant.
Perhaps we simple take in heat from our environment to compensate for its loss during our work? This would balance the equation nicely, but it simply isn't true. Our body temperatures are almost always higher than the environments (37ºC / 98ºF is usually much higher than the temperature the weatherman gives). Therefore on most days, we compound our thermodynamic problem by losing heat to the environment too! Even on hot days, when the environment is actually hotter than us, the body is unable to utilise this energy usefully anyway.
So what's going on? How do we resolve the paradox?
The answer is that humans are simply not an example of a closed system. We clearly can exchange mass with our environment - we can take in food and liquids (and excrete urine and faeces). We are what a physicist would call and open system - we may exchange both energy and mass with our environments. Since the food itself could briefly be regarded as a system, it has internal energy of its own - it has a temperature, etc. too, but it also has wonderful chemical potential energy in the bonds between its atoms.
And so when we eat, we are actually adding some internal energy from the environment (i.e. food) to our own system's internal energy - something that we couldn't do if we were a closed system. And this increased internal energy eventually goes into work, allowing us to keep living (which is always nice) and heat, keeping our body temperature constant.
So in an abstract and bird's-eye-view way, we know why we die if we don't eat or drink - our body simply runs out of (internal) energy to do work with. Also, the first law of thermodynamics explains why people with very low energy reserves (e.g. sick, underfed newborns and starving people) often drop their body temperatures. This is a very grave sign to a doctor, since it indicates that they don't even have enough energy stores to generate heat, and are probably very close to death.
The above also applies to evolution in its broadest sense. To a physicist (but not to a biologist!), evolution is very similar to growth in the sense that, over great lengths of time, simple molecules have become complicated organisms. Clearly, this process also requires energy, and so clearly thermodynamics applies. It's easy to see that, at least for the first law, evolution definitely does not violate thermodynamics - at all stages, the organisms could exchange matter with the environment, and thus avoid any thermodynamic complications. In the next post, I'll tackle the second law.
Reference: "Physics" (5th Edn.) - Giancoli
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