Theoretical physics: The origins of space and time

Many researchers believe that physics will not be complete until it can explain not just the behaviour of space and time, but where these entities come from. Finding that one huge theory is a daunting challenge. Nature explores some promising lines of attack — as well as some of the emerging ideas about how to test these concepts (see ‘The fabric of reality’). Full Article→

Credit: Zeeya Merali

What happens to a liquid in a cold vacuum? Does it boil or freeze? These animations of liquid nitrogen (LN2) in a vacuum chamber demonstrate the answer: first one, then the other! The top image shows an overview of the process. At standard conditions, liquid nitrogen has a boiling point of 77 Kelvin, about 200 degrees C below room temperature; as a result, LN2 boils at room temperature. As pressure is lowered in the vacuum chamber, LN2’s boiling point also decreases. In response, the boiling becomes more vigorous, as seen in the second row of images. This increased boiling hastens the evaporation of the nitrogen, causing the temperature of the remaining LN2 to drop, the same way sweat evaporating cools our bodies. When the temperature drops low enough, the nitrogen freezes, as seen in the third row of images. This freezing happens so quickly that the nitrogen molecules do not form a crystalline lattice. Instead they are an amorphous solid, like glass. As the residual heat of the metal surface warms the solid nitrogen, the molecules realign into a crystalline lattice, causing the snow-like flakes and transition seen in the last image. Water can also form an amorphous ice if frozen quickly enough. In fact, scientists suspect this to be the most common form of water ice in the interstellar medium. (GIF credit: scientificvisuals; original source: Chef Steps, video; h/t to freshphotons)

You want a physicist to speak at your funeral. You want the physicist to talk to your grieving family about the conservation of energy, so they will understand that your energy has not died. You want the physicist to remind your sobbing mother about the first law of thermodynamics; that no energy gets created in the universe, and none is destroyed. You want your mother to know that all your energy, every vibration, every Btu of heat, every wave of every particle that was her beloved child remains with her in this world. You want the physicist to tell your weeping father that amid energies of the cosmos, you gave as good as you got.

And at one point you’d hope that the physicist would step down from the pulpit and walk to your brokenhearted spouse there in the pew and tell him that all the photons that ever bounced off your face, all the particles whose paths were interrupted by your smile, by the touch of your hair, hundreds of trillions of particles, have raced off like children, their ways forever changed by you. And as your widow rocks in the arms of a loving family, may the physicist let her know that all the photons that bounced from you were gathered in the particle detectors that are her eyes, that those photons created within her constellations of electromagnetically charged neurons whose energy will go on forever.

And the physicist will remind the congregation of how much of all our energy is given off as heat. There may be a few fanning themselves with their programs as he says it. And he will tell them that the warmth that flowed through you in life is still here, still part of all that we are, even as we who mourn continue the heat of our own lives.

And you’ll want the physicist to explain to those who loved you that they need not have faith; indeed, they should not have faith. Let them know that they can measure, that scientists have measured precisely the conservation of energy and found it accurate, verifiable and consistent across space and time. You can hope your family will examine the evidence and satisfy themselves that the science is sound and that they’ll be comforted to know your energy’s still around. According to the law of the conservation of energy, not a bit of you is gone; you’re just less orderly. Amen.

—  Aaron Freeman

The first law of thermodynamics is very famous. It says, “Energy can neither be created nor be destroyed but is converted from one form to another.”

So simple.

But can you explain the second law of thermodynamics? A bit puzzled, aren’t you?

It’s the very reason I call it the monsters of bioenergetics. Let’s convert these monsters into cute little pixies :)

The second law of thermodynamics says, “The entropy of the universe goes on increasing over time.”

What is entropy?

Entropy is the degree of randomness.           

A solid has closely placed molecules. Hence, the randomness in molecules is less. On the on the other hand, in liquids, the distance between the molecules is more. Hence, they have more randomness and more entropy value.

Melting of ice is a good example which illustrates the second law of thermodynamics. When the ice melts, solid gets converted into it’s liquid form. The distance between the molecules increases from solid to liquid and thus, the entropy increases!

Here’s an interesting fact: The human body consumes carbohydrates, breaks it down and stores its energy as ATP, which is a high energy molecule. One would argue that storage of such high energy molecule is against the second law, as entropy of the body is not increasing in this reaction. The entropy increases, but in this case, the entropy of the universe increases because we release carbon dioxide into the surrounding!

Since we are on this topic, let’s address two more terms - Gibbs free energy and enthalpy!

Gibbs free energy

It is the Gibbs free energy which determines whether the reaction will proceed spontaneously to equilibrium without any input from surrounding.

In a reaction, if reactants are unstable (Having more energy) and the products are stable (Having less energy), then the reaction tends to move forward spontaneously without any input from surrounding.

On the other hand, if reactant is more stable than products then for this reaction to happen there has to some input of energy from surrounding.

Hence, if products have less Gibbs free energy than the reactants (i.e. change in Gibbs free energy is negative) then the reaction is spontaneous/exergonic irrespective of whether it is exothermic or endothermic.

If products have more Gibbs free energy than the reactants (i.e. change in Gibbs free energy is positive) then the reaction is non-spontaneous/endergonic.


Enthalpy (H) is a sum of useful energy and non-useful energy. The non-useful part is the Entropy (S) and the useful part is the Gibbs free energy (G).


To summarize all the three terms:

Entropy: Degree of randomness (Non-useful energy)

Gibbs free energy: Energy available to do work (Useful energy)

Enthalpy: Sum of Entropy and Gibbs free energy!

Related post: How to remember the sign and direction of Gibbs free energy change


If you find yourself some place really cold this holiday season, may I suggest stepping outside and having some fun freezing soap bubbles? The crystal growth is quite lovely, as seen in this photograph. If you live in warmer climes, fear not, you can always experiment in your freezer. It would be particularly fun, I think, to see how a half-bubble sitting on a cold plate freezes in comparison to a droplet like this one. (Video credit: Mount Washington Observatory)

You guys ask great questions—I’ve missed replying to them! But I love talking about entropy so prepare yourself.

The Second Law of Thermodynamic is part of a set of three fundamental, beautifully simple physical laws of a thermodynamic system.

  • First Law: Energy is not created or destroyed.
  • Second Law: In an isolated system, entropy must always increase.
  • Third Law: Absolute zero cannot be achieved.

Most people understand the Second Law to mean that “In an isolated system, disorder must always increase”, as in things always tend from order to disorder, but entropy does not automatically equate to disorder. It’s just a measure of disorder.

For those who are unfamiliar with entropy, the basic concept is pretty simple because we see it all around us everyday—objects break, relationships disintegrate, people age. To understand the terms of “high” and “low” entropy means, think of it like this: if a library is highly-organised and has an efficient indexing system, it could be called a low entropy library, whereas if a library has jumbled shelving, a disorganised indexing system and books all over the place, it could be called a high entropy library.

By referring to an ‘isolated system’, the Second Law is referring to a system where no energy is being added to it or taken from it, and entropy is a gauge of the energy in such a system that can’t be used anymore. This unavailable energy hasn’t left the system—it’s just become irretrievably disordered—but even though this will increase, it doesn’t mean that elsewhere in the system other energy can’t become more ordered.

Sure, a system (for example, the universe) contains unavailable energy, but the rest of the universe’s energy needs to be accounted for. It’s still out there, still doing stuff, and can still behave in a bunch of different ways depending on a variety of forces. It’s completely possible for a closed system to create even complex and elaborate order, just as long as there’s a balance and there’s an increase of disorder elsewhere in the system—importantly, you can only create order by increasing disorder too, because creating order involves expending energy, which is inevitably inefficient and so it adds to the unavailable energy.

True to the Second Law of Thermodynamics, the total amount of order in the universe is always decreasing, but that doesn’t mean parts of the universe can’t continue to become ordered too.

It’s like a struggle against the tide on a cosmic-scale. Structures, stars and organisms are created as low-entropy systems, but it’s fruitless because disorder is a byproduct of order and will triumph in the end—and yet, the universe keeps striving for order all the same.



A team from Caltech recently devised a series of demonstrations that illustrate the concepts of oxidation, exothermic reactions, and activation barriers. In the grand finale, sulfuric acid ignites a mixture of sugar and potassium chlorate. That reaction produces enough heat to ignite the thermite reaction (Fe2O3 + Al –> 2 Fe + Al2O3), which in turn generates molten iron and some fireworks, along with a huge quantity of heat.

Credit: Brett McGuire/J. Chem. Ed, 2014 DOI: 10.1021/ed500522c

Watch on

Reader kylewpppd asks:

Have you seen the post of a man in Siberia throwing boiling water off of his balcony? Can you provide a better explanation of what’s going on?

As you can see in the video (and in many similar examples on YouTube), tossing near boiling water into extremely cold air results in an instant snowstorm. Several effects are going on here. The first thing to understand is how heat is transferred between objects or fluids of differing temperatures. The rate at which heat is transferred depends on the temperature difference between the air and the water; the larger that temperature difference is the faster heat is transferred. However, as that temperature difference decreases, so does the rate of heat transfer. So even though hot water will initially lose heat very quickly to its surroundings, water that is initially cold will still reach equilibrium with the cold air faster. Therefore, all things being equal, hot water does not freeze faster than cold water, as one might suspect from the video.

The key to the hot water’s fast-freeze here is not just the large temperature difference, though. It’s the fact that the water is being tossed. When the water leaves the pot, it tends to break up into droplets, which quickly increases the surface area exposed to the cold air, and the rate of heat transfer depends on surface area as well! A smaller droplet will also freeze much more quickly than a larger droplet.

What would happen if room temperature water were used instead of boiling water? In all likelihood, a big cold bunch of water would hit the ground. Why? It turns out that both the viscosity and the surface tension of water decrease with increasing temperature. This means that a pot of hot water will tend to break into smaller droplets when tossed than the cold water would. Smaller droplets means less mass to freeze per droplet and a larger surface area (adding up all the surface area of all the droplets) exposed. Hence, faster freezing!

Broad Summary of Thermodynamics: Sometimes seeing the Bigger Picture helps you Put together the smaller pieces and connect the dots b/w Physics, Chemistry, and Biology. 


Zeroth law (concept of temperature)

0th law of thermodynamics basically says that heat flows from hot objects to cold objects to achieve thermal equilibrium.
Mathematically, if TA = TB, and TB = TC, then TA = TC. Where T is temperature.

First law (ΔE = q + w, conservation of energy)

1st law of thermodynamics is based on the principle of conservation of energy, and it basically says that the change in total internal energy of a system is equal to the contributions from heat and work.
ΔE is the same thing as ΔU, which is the change in internal energy.
Q is the contribution from heat
Q is positive when heat is absorbed into the system (ie. heating it).
Q is negative when heat leaks out of the system (ie. cooling it).
W is the contribution from work.
W is positive when work is done on the system (ie. compression).
W is negative when work is done by the system (ie. expansion).

Equivalence of mechanical, chemical, electrical and thermal energy units

If it’s energy, it’s Joules. It doesn’t matter if it’s potential energy, kinetic energy, or any energy - as long as it’s energy, it has the unit Joules.
Energy is equivalent even if they are in different forms. For example, 1 Joule of mechanical energy can be converted into 1 Joule of electrical energy (ignoring heat loss) - no more, no less.

Second law: concept of entropy

The 2nd law states that the things like to be in a state of higher entropy and disorder.
An isolated system will increase in entropy over time.
An open system can decrease in entropy, but only at the expense of a greater increase in entropy of its surroundings.
The universe as a whole is increasing in entropy.
ΔS ≥ q / T
q is the heat transferred.
T is the temperature in Kelvin.
For reversible processes ΔS = q / T.
For irreversible processes ΔS > q / T.
Real processes that occur in the world are never reversible, so entropy change is always greater than the heat transfer over temperature.
Because of the irreversibility nature of real processes, as long as anything occurs, the entropy of the universe increases.