‘To heat a planetary surface’ for dummies; Part 1

Happy New Year to everyone! Hope you all had a pleasant celebration.

I will unabashedly start off in 2015 with … another attempt at exposing the chasm that lies between what real physics tells us about the processes of nature (plus what we actually observe in the real world) on the one hand, and what the ‘physics’-like concoctions of the radiative GHE/AGW-establishment proclaim on the other.



The general public understanding (or should we rather call it ‘perception’?) of how the presence of an atmosphere would make the solar-heated planetary surface underneath warmer than if the atmosphere weren’t there, is so riddled with misconceptions and flawed ideas about how the world works, on such a fundamental level, that something needs to be done.

People simply need to understand that the official (and, I’m afraid, ‘authoritative’) rGHE/AGW ‘explanation’ is based altogether on self-invented nonsense physics.

The best way to let people realise this is to explain how things really work and to have this juxtaposed with the standard rGHE postulates advertised by ‘Climate ScienceTM’.


IT’S ALL ABOUT THERMODYNAMICS

The first thing that needs to be recognised by all is that the one field of study that specifically addresses and handles everything concerning energy transfers between different regions or systems and the temperature changes (if any) that result from them, is … that branch of physics called ‘Thermodynamics’.

Any argument trying to introduce for instance quantum theory concepts like ‘photons’ into the mix, or laws and relationships pertaining specifically to ‘pure emitters’ and ‘blackbody radiation’, can safely be dismissed as irrelevant to the issue at hand. Bringing in these concepts changes nothing, illuminates nothing. It only confuses the matter.

Which would be the purpose for bringing them in to begin with.

We’ll get back to this later …

Anyway, it cannot be stressed too much or too often that any problem dealing with energy transfers between regions or systems to change internal energies (see below) and thus absolute temperatures, is a Thermodynamics problem. Period. There is no need, no point and no sense in invoking principles from other preferred branches of physics in order to attempt in some way or fashion to overrule or circumvent the conclusions reached through a strictly thermodynamic analysis of the problem.

To be more specific, when trying to explain why and how the global surface of the Earth is so much warmer on average than the global surface of the Moon, and so much warmer than a hypothetical surface emission temperature reached at pure radiative equilibrium with the Sun, you are operating completely within the realm of Thermodynamics.

I was hoping this would be self-evident. Apparently it’s not.

Further, the thermodynamic approach to any problem such as this is first and foremost one of energy budgeting, simply at all times accounting for the energy coming in to, going out from and held within the systems under study. This is where any proper thermodynamic analysis would start and end. Keeping track of the energy. Temperatures are secondary.


WHAT IS AN ‘ENERGYY TRANSFER’?

The second thing that needs to be understood by everyone is what an ‘energy transfer’ – as defined by thermodynamic principles – really constitutes.

The ‘energy transfer’ is one of the staple concepts of thermodynamics. It comes in two varieties and two varieties only.

An energy transfer is the necessary and automatic result of some imbalance (disequilibrium) between the system under study and its surroundings. The transfer of energy always occurs unidirectionally, either from the system to its surroundings, or from the surroundings to the system.

Energy can be transferred out of or into a thermodynamic system by way of ‘heat’, normally represented by the symbol Q, or by way of ‘work’, normally represented by the symbol W. There is no other way. No other method or mechanism. If it’s not by ‘heat’, it’s by ‘work’. If it’s not by ‘work’, it’s by ‘heat’. If it’s neither of the two, then there simply is no energy transfer going on. And remember, both ‘heat’ and ‘work’ transfers always operate unidirectionally, one way only.

Please reread the above paragraph until you’re blue in the face, or (preferably) until you’re absolutely sure you understand what it says. It is most likely the most essential piece of thermodynamic knowledge you’ll ever learn. So please make sure you do.

What, then, happens to energy that is not transferred as ‘heat’ or ‘work’? Likewise, what happens to energy after it’s been transferred as ‘heat’ or ‘work’? To put it perhaps more succinctly: Where is the energy whenever it’s not being transferred between systems?

It is of course residing inside the separate systems.

Energy in this condition (static (in storage) rather than dynamic (in transit)) is termed ‘internal energy’ within the field of thermodynamics, generally designated by the letter U.


TRANSFER VS. STORAGE OF ENERGY – THE 1ST LAW

There is a very well-known and straightforward relationship between the three possible energy states as defined by thermodynamics: Q, W and U. It is called the 1st Law of Thermodynamics:

ΔU = Q – W

It says: The change in system ‘internal energy’ [ΔU] during the course of some thermodynamic process in which the system is involved, comes directly and solely as a result of the total amount of energy transferred to or from the system as ‘heat’ [Q] minus the total amount of energy lost or gained by the system through ‘work’ done [W]. This is one way of demonstrating the fundamental principle of ‘the conservation of energy’.

Here is what Wikipedia says about ‘internal energy’:

“In thermodynamics, the internal energy is one of the two cardinal state functions of the state variables of a thermodynamic system [the other being ‘entropy’, S]. It refers to energy contained within the system, while excluding the kinetic energy of motion of the system as a whole and the potential energy of the system as a whole due to external force fields. It keeps account of the gains and losses of energy of the system.

The internal energy of a system can be changed by (1) heating the system, or (2) by doing work on it, or (3) by adding or taking away matter. When matter transfer is prevented by impermeable walls containing the system, it is said to be closed. Then the first law of thermodynamics states that the increase in internal energy is equal to the total heat added and work done on the system by the surroundings. If the containing walls pass neither matter nor energy, the system is said to be isolated. Then its internal energy cannot change.

The internal energy of a given state of a system cannot be directly measured. It is determined through some convenient chain of thermodynamic operations and thermodynamic processes by which the given state can be prepared, starting with a reference state which is customarily assigned a reference value for its internal energy.”

(My emphasis.)

From the HyperPhysics site:

“Internal energy is defined as the energy associated with the random, disordered motion of molecules. It is separated in scale from the macroscopic ordered energy associated with moving objects; it refers to the invisible microscopic energy on the atomic and molecular scale. For example, a room temperature glass of water sitting on a table has no apparent energy, either potential or kinetic. But on the microscopic scale it is a seething mass of high speed molecules traveling at hundreds of meters per second. If the water were tossed across the room, this microscopic energy would not necessarily be changed when we superimpose an ordered large scale motion on the water as a whole.

inteng1

(…)

Internal energy involves energy on the microscopic scale. For an ideal monoatomic gas, this is just the translational kinetic energy of the linear motion of the “hard sphere” type atoms, and the behavior of the system is well described by kinetic theory. However, for polyatomic gases there is rotational and vibrational kinetic energy as well. Then in liquids and solids there is potential energy associated with the intermolecular attractive forces. A simplified visualization of the contributions to internal energy can be helpful in understanding phase transitions and other phenomena which involve internal energy.

mceng

(…)

When the sample of water and copper are both heated by 1°C, the addition to the kinetic energy is the same, since that is what temperature measures. But to achieve this increase for water, a much larger proportional energy must be added to the potential energy portion of the internal energy. So the total energy required to increase the temperature of the water is much larger, i.e., its specific heat is much larger.”

intex

[And:]

“Temperature is not directly proportional to internal energy since temperature measures only the kinetic energy part of the internal energy, so two objects with the same temperature do not in general have the same internal energy (see water-metal example [above]). Temperatures are measured in one of the three standard temperature scales (Celsius, Kelvin, and Fahrenheit).

Suppose we are dealing with two equal mass objects at ordinary temperatures and can presume that kinetic temperature gives a reasonable description of their behavior. If the two objects are at the same temperature, then we would say that their average translational kinetic energies are the same. That does not imply that their total internal energies are the same, because the potential energies associated with intermolecular forces can be quite different.

temper1

Even if there are internal kinetic energies other than translational KE, it could be that heat transfer is mainly by collisional transfer. In such cases this picture might help understand that just a portion of the total internal energy of objects is involved in the conditions for thermal equilibrium.”

By the 1st Law of Thermodynamics, a transfer of energy to or from a system (as ‘heat’ or ‘work’) changes the internal energy – and thus normally the temperature – of that system.* This effect (changing system internal energies and thus temperatures) is exclusive to the two varieties of thermodynamic energy transfer. If a thermodynamic process brings about a change in a system’s ‘internal energy’ [U] (and temperature), you have transferred energy to or from it as ‘heat’ [Q] and/or ‘work’ [W]. There is no other way. Transfer heat to it or let its surroundings do work on it, the internal energy increases. And temperature rises. Transfer heat from it or let it do work on its surroundings, the internal energy decreases. And temperature drops.

*The special exception to this rule being phase transitions, where a transfer of heat will NOT start adding to the temperature-relevant part of the internal energy (the microscopic kinetic part as opposed to the microscopic potential part) and thus the temperature, before the phase has changed completely.


THE ‘HEAT TRANSFER’ [Q]

Let’s now single out and turn specifically to the Q term of the 1st Law, since heat transfers’ really are the only kind of energy transfer strictly relevant to the very surface of our planet, the one that is supposed to be so much warmer on average than the surface of the Moon, distinctly by virtue of our atmospheric rGHE.

We have already pointed out that any thermodynamic energy transfer is unidirectional, it always and only moves one way. In the case of energy transferred between systems as ‘heat’, this direction is invariably the same. Spontaneously, in nature, heat moves from hot to cold, never the opposite way. If something were to move all by itself in the opposite direction, it would not be heat. By definition.

So, what is required in a heat transfer? A single heat transfer involves two systems only (or one system and its surroundings). If you want to include multiple systems, then by definition, there will be multiple, separate heat transfers.

For a single heat transfer to occur you would need:

  1. A hot reservoir (a heat/energy source), and
  2. a cold reservoir (a heat/energy sink).

The hot reservoir is the warmer system. It is the system delivering (losing) the energy being transferred as heat. The source. Heating the other system.

The cold reservoir is the cooler system. It is the system receiving (gaining) the energy being transferred as heat. The sink. Cooling the other system.

Now, pay attention! There is no other way that a single heat transfer could function. It could give no other result than this: The hot reservoir loses energy to the cold reservoir. Its internal energy [U] thus decreases and so naturally does its temperature [T]. Conversely, the cold reservoir gains energy from the hot reservoir. Its internal energy thus increases and so naturally does its temperature.

The consequence of this is a fundamental tendency of all natural processes – moving towards increasing entropy [S], ideally ending up maximising it. In the specific case of a heat transfer, this state is reached when the temperatures of the two systems involved have equalled and we have what is called a ‘thermal equilibrium’. In this situation the original hot reservoir cannot lose any more energy to its cold reservoir and the original cold reservoir can likewise not gain any more energy from its hot reservoir. The heat transfer has ceased.

The final equilibrium temperature of the two systems will always – and this is an exceedingly important point! – lie somewhere in between the initial temperatures of the systems. That is to say, it will always be lower than the initial temperature of the hot reservoir and higher than the initial temperature of the cold reservoir. The hot reservoir will ALWAYS cool during a heat transfer process. And the cold reservoir will ALWAYS warm (again barring stages of phase transition).

Why is this such an important point to make?

It is crucial, because knowing this will allow you to recognise immediately how all claims of warming at both ends of a heat transfer process, only more warming at the cooler end and less at the warmer, are fundamentally at odds with basic physics, thus flawed and/or deceptive. Nature doesn’t work like that. Not under any circumstances. You will always have warming at one end only, cooling at the other.


  • Transferring energy as heat TO a system will increase the internal energy of that system and thus (ultimately) its temperature.
  • Transferring energy as heat FROM a system will reduce the internal energy of that system and thus (ultimately) its temperature.
  • A transfer of energy as heat always spontaneously moves from hot to cold only. ‘Heat’ is, by thermodynamic definition, something that in nature always comes IN from … a hotter place.

KEEPING HEAT TRANSFERS APART?

Finally, there is the important point of separating between individual heat transfers. A single system can be involved in two (or more) heat transfers operating simultaneously. This doesn’t mean you’re free to lump them together as if one single process. Quite the contrary. You very much need to keep them apart in your analysis of the situation, so as to enable yourself to discover what’s actually going on as the internal energy and the temperature of, plus the energy input to/output from, the system change.

A relevant example:

300px-Carnot_heat_engine_2.svg

Figure 1.

How many heat transfers do you see? Yes, indeed. Two.

Note how the central system – the so-called ‘working fluid’ or ‘heat engine’ – is involved in both of them at the same time. The ultimate hot reservoir (heat source) [TH] to the left and the ultimate cold reservoir (heat sink) [TC] to the right, are only involved in one each. The central system gains energy from TH. This transfer is a positive one, coming from a hotter place [Qin (QH)]. At the same time it loses energy to TC. This transfer is a negative one, going to a colder place [Qout (QC)]; normally it would be its general surroundings. Notice how the Qout (QC) does not go back towards the hot reservoir (TH), the source of the Qin. Heats do not oppose each other. Energy spontaneously being transferred out of a system as heat will always seek out a cold reservoir, a place cooler than the system itself. It cannot move in the direction of warmer. This is why it is all-essential to keep and tell distinct heat transfers apart. To avoid confusing yourself when interpreting effects and changes.

Take one last look at Figure 1. You might realise something that doesn’t necessarily stand out at once as obvious from the sketch itself, but which, if you think about it, is actually pretty self-evident. To fully apprehend what I’m getting at, mentally replace TH with the Sun, the central system with the Earth’s surface, and TC with the atmosphere.

What you will realise is this: The second heat transfer (QC) is not independent of the first (QH). Actually, the opposite is true. It owes its very existence to that of the first one. QH supplies the energy needed for QC to start working in the first place, and thereafter for it to continue to work. Same with the surface of the Earth. It could never heat the atmosphere if it weren’t first itself heated by the Sun.

Still, the two heat transfers are very much operating distinctly from one another. And should be handled individually. As neatly expressed by the diagram in Figure 1.



Sum-up at this stage:

  • The issue regarding why the mean temperature of the global surface of the Earth is so much higher than the mean temperature of the global surface of the Moon is strictly a THERMODYNAMIC one. Any problem concerning energy transfers (potentially) resulting in temperature changes is. So we conduct a thermodynamic analysis to ‘solve’ it. Not a quantum theory-based one. Or any other kind of analysis, for that matter …
  • According to thermodynamic principles, the ‘internal energy’ [U] of a system corresponds to its ‘temperature’ [T]; not directly, but in the sense that – disregarding the special case of phase transitions – if you change the U of the system, you also change its T, in the same direction (+U ⇒ +T; −U ⇒ −T). You can change the U of a system and not change its T (phase), but you can never change the T of that same system without changing its U.
  • According to the 1st Law of Thermodynamics, you can only change the U (and thereby the T) of a system in two ways: 1) By transferring energy to or from it as ‘heat’ [Q], or 2) by transferring energy to or from it as ‘work’ [W]. 1) is accomplished by heat transfers to/from the system, 2) is accomplished by work being done by/on the system. There is no other way of changing the U and T of a system. The transfer of energy as Q or W. That’s it.
  • A single system can be involved in more than one specific heat transfer process. It is very important when conducting a thermodynamic analysis to keep several different heat transfers separate. If you don’t, you will not be able in the end to say anything meaningful about cause and effect. You will be confused and might very well draw spurious conclusions based more on your own inherent preconceptions about how the world should work than on how it really does work.

The next post in this series will move on to take a closer look (still, of course, from a strictly thermodynamic perspective) at how real-world objects actually warm and reach a steady state temperature.

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5 comments on “‘To heat a planetary surface’ for dummies; Part 1

  1. markstoval says:

    I really enjoyed this post and am looking forward to the next in the series.

  2. Truthseeker says:

    Clarity has been achieved … again!

  3. Nice if you limit your thinking to thermodynamics. Only temperature and pressure, and thinking of energy as something other than an accumulation of power (force acting), work, or power collected over time your Q. It also leaves out all forms of inertia, ever the angular. This would leave out most chemical processes, even the spontaneous processes, which require mass and a changing structure. Your concept of internal energy is what the rest of us call sensible heat of any mass that has a specific heat, a blob with an attitude (temperature)!
    It also excludes the accumulation of power done by the work of by forcing the halves of one blob apart, (gravitational force). You excludes latent heat also requiring both mass and structure, ad its relationship to pressure, one of the more important parts of all thermodynamics.
    Most significantly, you exclude all that does not directly require some change in the properties of mass. That is everything electrical, that can produce work, force times distance or torque times revolutions. Electro-motive, or magneto-motive force (power) actually requires no mass.
    And certainly excludes all electrical relativistic such as the transmission of information or power EMR, which only “sometimes” requires mass, temperature, and sensible heat at one or both ends, and very troublesome in between the ends. Please remember EMR has its own laws that “proceed” any and all laws of thermodynamics. Nothing is more relativistic than EMR. The relativistic has its own set of Lorentz transforms and conservations.

    • okulaer says:

      Sorry, Will. I’m gonna do this my way. Keeping it simple.

      If you want an explanation to include all of those things you mention above, then go ahead, feel free to write it up … start your own blog. If you haven’t already got one. I’m not one to hold you back …

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