At 0 degree Celsius, water molecules slow down enough to bond with each other and ice starts to form. By adding more salt or sodium chloride the freezing point of water becomes even lower, why is that? could it be that when water molecules rip apart sodium chloride (fragile ionic bonding) it gains some energy which prevents itself from bonding with other water molecules? Sounds weird because I always thought that the hydrogen bonding is the weakest among the different types of chemical bonding.
How does salt lowers the freezing point of water
ionsphase-transitionphysical-chemistrythermodynamicswater
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The divide is actually not between covalent and ionic, but rather a spectrum between localised and delocalised electrons. The history of all this is actually quite fascinating, and Phil Anderson in his book "More and Different" has a nice chapter on this. Essentially, around the time that people started doing quantum mechanics on molecules seriously, there were two schools of thinking which dominated.
On one side was Mott and more popularly, Hund and Pauli who thought of electrons as primarily attached to atoms and through electromagnetic interactions their motions/orbitals would be deformed and one gets molecules. This is the version usually taught in chemistry classes as with a few rules of thumb it is possible to qualitatively account for a vast range of behaviours.
On the other side was Slater with a dream of a machine which could simply compute the electronic structure by giving it the atoms and electrons. In this picture, the electrons are primarily thought of as delocalised over all the atoms, and through a rigorous procedure of perturbation theory one adds the effect of interactions between electrons and may achieve arbitrarily good precision.
The latter has the problem that the results are not intuitive --- there are no rules of thumb available and one is reduced to simply computing. The problem with the former is that to achieve high accuracy, the "rules of thumb" become exceedingly complex and are not really very easy to use or to compute with --- it lacks the simple regularity of the Slater dream machine. It is telling that essentially the latter has won, and nowadays it is routine to compute the electronic structure of quite large molecules (~1000 atoms) through brute-force (the technique is known as density functional theory, and there are commercial software available to do it).
In finite molecules one can actually show that in principle both approaches will work --- technically we speak of there being an adiabatic connection between the localised and delocalised states. The only practical difference is just how hard it is to carry out the calculations. However, in infinite molecules (e.g. solid crystals) this is not true, and there can be a proper phase transition between the two starting points. In that case, the localised approach corresponds to what is fancily called these days "strongly correlated systems" such as Mott insulators and magnetically ordered materials, and the delocalised approach are essentially metals (technical language: renormalises to be a Fermi liquid).
Nowadays there is a desire (from theoretical condensed matter physicists) to develop the localised approach again, as it may be possible to find some useful rules of thumb regarding magnetic materials, a prominent example of which are the high temperature superconductors.
Adding salt to water makes it freeze at a lower temperature. This fact is being used in two different ways in the two scenarios you mention. Dissolving sodium chloride in water is slighly endothermic, but this effect is small and to the best of my knowledge isn't important in the drink cooling process.
Putting salt on the highway is quite straightforward: we don't want ice to form, so we put salt in the water to prevent that. This doesn't just change the amount of time it takes ice to form, it actually completely prevents ice from forming, unless the temperature gets so low that the water can freeze even with salt in it.
Cooling your drink is a bit more complicated, because in this case the rate at which things happen is important. You don't want your drink to be less than $0^\circ C$ because it would freeze; instead you want to cool it down to a few degrees Celsius nice and quickly.
The rate at which it cools depends on two things: the temperature of its surroundings (the colder the better) and the heat conductivity between it and them. You could try to cool it by putting it in a bowl of ice at $0^\circ C$, but the problem is that the ice is solid and will only touch the bottle at a few points. This results in a poor thermal conductivity, so the drink will only cool slowly.
To get around this, you could try mixing the ice with some water. Now the bottle is touching the liquid over a large surface area, and the liquid itself has a higher thermal conductivity than solid ice due to mixing, so heat will be transferred much more quickly. But the problem is that the water won't be at zero degrees any more, at least not at first (I'm assuming the water comes from a tap, so it's not chilled initially). You have to wait for quite a bit of the ice to melt before the water's temperature will drop. Also, once you put your warm drink into the water it will heat the water up as the drink cools down, so again you have to wait for the ice to melt in order for the water to cool again.
The solution to this is to make the ice melt faster. You can do this by adding salt. This lowers the freezing point, making the water less "happy" about being in the liquid state, so it melts more quickly. This means firstly that the tap water you've added to the ice will cool to close to $0^\circ C$ much more quickly, and secondly that once you've put your drink in there the water will stay cold as the ice continues to melt.
It's also possible that, with the salt added, the water can go to below $0^\circ C$, but this will only happen if the ice is quite a bit colder than $0^\circ C$. This could be the case, but my intuition is that the rate of cooling due to the ice melting faster is more important here than the final temperature. You could easily test this by putting a thermometer in the salty ice water and seeing if it goes much below freezing.
There's also the fact that dissolving the salt is endothermic, as you mentioned. To test whether this is important, you could try adding salt to some chilled water without any ice, and see if the temperature drops a lot. My feeling is that it will only drop by a tiny amount that will be hard to measure with a normal kitchen thermometer, but you can always try the experiment.
Best Answer
Related questions have been asked, but (to my surprise) I couldn't find this specific question being answered before. (Some answers come close, but are not quite there.)
The way that salt affects the temperature of a mix of water and ice involves the concept of entropy. Therefore I will first write about the aspect of entropy that is relevant here.
As introduction I discuss a simpler case of entropy being a relevant factor.
It involves a molecule that is heavier than air, that (on average) moves against gravity. It was demosntrated to a class of students as follows:
The demonstration involved two beakers, stacked, the openings facing each other, initially a sheet of thin cardboard separated the two.
In the bottom beaker a quantity of Nitrogen dioxide gas had been had been added. The brown color of the gas was clearly visible. The top beaker was filled with plain air. Nitrogen dioxide is denser than air.
When the separator was removed we saw the brown color of the Nitrogen dioxide rise to the top. In less than half a minute the combined space was an even brown color.
And then the teacher explained the significance: in the process of filling the entire space the heavier Nitrogen dioxide molecules had displaced lighter molecules. That is: a significant part of the population of Nitrogen dioxide had moved against the pull of gravity. This move against gravity is probability driven.
Statistical mechanics provides the means to treat this process quantitively. You quantify by counting numbers of states. Mixed states outnumber separated states - by far.
What this demonstration illustrates is that in specific circumstances tendency to increase entropy and tendency to go down an energy gradient can be acting in opposite direction, and in this demonstration the entropy increasing tendency dominated.
Dissolution of salt in water
Interestingly, the dissolution of salt in water is by itself already an endothermic process. That is: if you start with water at a particular temperature, and salt at that same temperature, and you allow the salt to dissolve in the water, then the homogeneous solution will be at a (slightly) lower temperature.
So: purely from the point of view of energy level it is unfavorable for salt to dissolve. The point is: the dissolved state is more probable. Without going into more detail, the is-more-probable property is correlated with the fact that the dissolved state is more homogeneous than the not dissolved state.
Ice and brine
The ratio of ice and brine is a dynamic equilibrium. There is a rate of water molecules joining the existing solid ice, and there is a rate of water molecules being knocked off the ice, and joining the brine.
A major factor, of course, affecting those two (opposite) rates, is temperature. At lower temperature the water molecules move slower, reducing the probability of a water molecule being knocked off the ice.
There is a very high energy barrier to incorporating ions of the salt into the ice crystal structure. That is, the crystallization is a process that makes the total system less homogeneous.
As mentioned earlier, being more homogeneous is correlated with being more probable.
Comparing ice-and-water to ice-and-brine: the presence of the salt shifts the probability, moving the equilibrium state of the two (opposite) rates to a lower temperature. As mentioned earlier, that equilibrium is dynamic equilibrium of rate of water molecules joining the ice, versus rate of water molecules being knocked loose from the ice.