Is the arrangement of elements affiliated with their respective boiling points? I’ve always been told it was my atomic mass/number but my teacher seems to think otherwise.
Sort of. It’s possible to order your periodic table arranged by boiling points.
Like this.
http://www.chemicalelements.com/show/boilingpoint.html
From this site. See the list on the left? You can have your periodic table however you want it–name, atomic number, with anchovies, whatever… 
http://www.chemicalelements.com/
Er…“arranged WITH boiling points”, not “BY boiling points.” I’m assuming that this is what your teacher is talking about. I should point out that the basic, standard, classic arrangement IS by atomic number. In this website’s variations, Hydrogen is always #1, Helium is always #2, and just the extra information appears in the bottom of the little boxes.
No, it’s only by atomic number (it almost matches with atomic weight but there’s two places where that’s off). It doesn’t really have anything to do with boiling point other that the general periodicity that you see with most physical properties, like electronegativity, atomic radius, etc etc.
Boiling point doesn’t have any real linear progression on the periodic table. The heavier elements do boil at higher temperatures, but if you look at tungsten, which is in about the middle, you will note that it has the highest boiling point on the chart.
if that is your chemistry teacher I would be worried
So, molecularly / atomically, what does determine boiling point?
The simplest answer is that the relative energies of the solid and liquid forms determine it- the boiling point is where the two are equal in total energy.
Some solids are really stable (The ideal bonding orientation for carbon is a tetrahedron, and diamond has carbon in tetrahedral orientations- similarly, the ideal orientation of covalently bound atoms to tungsten is at the vertices of an icosahedron, which is where they are in the solid metal.) The other side matters too, as this is a relative measure- more energetically favorable liquid forms will lower the boiling point.
Thanks Tim but should that read “solid and liquid forms” or “liquid and gaseous forms”?
Also, if you have a salt-water mixture (or whatever it’s called) the water will evaporate, and leave the salt as a residue (or whatever it’s called). Is this governed by the same mechanism? I would think so, but I don’t remember much chemistry
I don’t believe this is correct (or maybe I’m misunderstanding you). You have to add energy to liquid water at 100 C to get steam at 100 C. (I’m assuming you meant “liquid and gas forms” above.)
Um, probably what was being implied is that each period has some standard characteristics as regards reactivity, melting and boiling points, etc., and that there is a definite trend within each period (e.g., as Period I’s atomic number increases, the melting point declines, while the exact reverse is true in Period VII).
Surely someone must know the molecular reason why in a salt water solution the water evaporates and not the salt. It’s really bugging me that I can’t figure this out.
It’s the same reason table salt itself doesn’t evaporate or sublimate. It’s vapor pressure just isn’t high enough. Salt crystalizes at relatively high temperatures (compared to water). This means that its preferred state is as a solid (even at relatively low pressure). Liquid is comparatively unstable, and depending on the dew point, there is always a certain amount of water that is soluable in the air. Since air is always moving, evaporations happens by means of disequilibrium processes of molecules on the surface of the water. In their ionic form, no sodium or chloride will end up in the air because the equilibrium far favors them being surrounded by polar water molecules. If you let the salt crystalize it isn’t going to evaporate either. There’s just no way to get that salt in the air.
Ring, it all comes down more or less as said to the difference in free energies between the liquid form and the vapor form. As a general rule, boiling points will be higher for heavier compounds, because heavier compounds tend to have stronger forces between them; they’re bigger, more polarizable, and hence have larger dispersion attractions between them.
However, not all compounds interact predominantly by dispersion forces, and that’s where the complications set in. For starters, some compounds (such as water) form hydrogen bonds, which are stronger than dispersion force; hence, water has a fairly high boiling point. Other compounds have permanent dipoles (also such as water), so the dipole-dipole interaction must be overcome; this also raises the boiling point.
The reason the salt doesn’t evaporate and the water does now becomes somewhat more obvious. In aqueous solution, we’ve really got something like Na[sup]+[/sup] and Cl[sup]-[/sup]. Obviously, these will attract each other more strongly than will water molecules with their comparatively weak dipole-dipole and hydrogen bond forces. So the water boils off first, and what you end up with is NaCl as a solid because the boiling point for ionic crystals is truly ludicrous.
hoping he remembers well from his misguided youth as a chemist
Ring, coming from a physics perspective, you might check out some phase transition diagrams from an intro thermodynamics text. They explain in much better mathematical detail how these mechanisms actually work. A temperature-based boiling point is actually something of a misnomer because pressure is something that is also important in any determination. Of course, chemists always assume STP, but that never really satisfied me:). Phase diagrams are helpful in visualizing such transitions.
Some interesting side notes: One might consider the triple point, which is the temp and pressure at which the gas, liquid, and solid can all coexist. There are also such things as critical points which occur as terminations of boundaries (in, for example, temperature/pressure phase diagrams where you can actually go from one state to another without going through a phase transition!), and eutectics (sp?) which are important in igneous geomorphology as sort of two-dimensional phase transitions (or three- or four- or n-dimensional if you wish).
Whew! Aren’t the states of matter fascinating?
Thank you gentlepersons for those most enlightening answers. I had some strange idea that because the water dissociated the Na and Cl that the Cl being a gas would boil off before the water molecules.
ah…I see the problem. True, elemental Chlorine is a gas. In salt however, the Chlorine exists as a negatively-charged chloride ion which is ionically bound to a positively-charged Sodium ion in a cubic lattice. This is a very stable structure, which is why you have to heat salt to about …well, a very high temperature to break the bonds apart. However, when you dissolve salt in water, the partially-charged water molecules cluster around these ions, in effect shielding their electic charge from each other which causes what is a thermodynamicaly stable structure to fall apart. (or, to use the scientific term, “dissolve”).
When you boil away the water, the shielding effect dissipates, and the ions re-associate, giving you back your salt.
Chlorine is also diatomic in its elemental form in order to fill the valence shell of both constituent atoms with the covalent bond. Ionic chloride is never diatomic because the valence shell is filled with an extra electron (thus the negative charge).
That’s what I like to hear! Except that with “igneous geomorphology”, I picture Shiprock, the Henry Mountains, etc. (I.e., igneous landforms).
But you’re right regarding their importance in igneous geochemistry and petrogenesis–eutectics represent the composition of (for example) a minimum partial melt. Many granites that were derived from a partial melt of crustal rocks, for instance, have an essentially eutectic composition.
This, of course, doesn’t help to answer the OP (which I think has been answered). Sorry.
Speaking of tetrahedral carbon, what would liquid diamond look like?