- London dispersion forces
- Dipole–dipole forces
- Hydrogen bonding
- Ion–dipole forces
- Intermolecular forces and vapor pressure
- Solubility and intermolecular forces
- Surface tension
- Capillary action and why we see a meniscus
- Boiling points of organic compounds
- Boiling point comparison: AP Chemistry multiple choice
- Solubility of organic compounds
- 2015 AP Chemistry free response 2f
- Intermolecular forces
- Intermolecular forces and properties of liquids
Solubility and intermolecular forces
Substances with similar polarities tend to be soluble in one another ("like dissolves like"). Nonpolar substances are generally more soluble in nonpolar solvents, while polar and ionic substances are generally more soluble in polar solvents. Created by Sal Khan.
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- He mentions that the Na+ ion is bigger than the Cl- ions. Isn't it the other way around? As the Na loses its valence electron, the rest of the electrons are attracted to the nucleus with a force that is much stronger because there are more protons than electrons. And as Cl gains an electron, the forces between the electrons increases, causing them to expand and making the ion bigger.(34 votes)
- You got me curious and make a good point so I looked up the values:
Na radius 154 pm; Na+ radius 102 pm
Cl radius 99 pm; Cl- radius 181 pm
Something to think about...(66 votes)
- So if NaCl separates into Na+ and Cl- when it disolves in water, is it still NaCl or just a bunch of separate Na+ and Cl- ions?(18 votes)
- it is still fairly close but each Na+ and Cl- ion is surrounded by water molecules due to ion dipole interactions...hope that helps :)(11 votes)
- At 4.00, Sal says that atoms get smaller as you go from left to right of the periodic table. However, I thought the number of protons in the nucleus increases by one when you go "right" in the periodic table? Confused.
Thanks in advance(9 votes)
- You're correct. The number of protons DOES increase and this increase in positive charge pulls the electrons in the same shell closer to the nucleus, thereby making the overall mass of the atoms smaller. Once a new electron shell is added, the size of the atoms drastically increases again. Cheers!(18 votes)
- Lets say you heat up a can of Coke on the stove. Would this allow the Carbon Dioxide to escape from the water molecules?(5 votes)
- provided the can is open. if it is closed, the pressure would build up in the can and eventually it would explode(13 votes)
- At3:49, Sal says that sodium atoms are larger than chlorine atoms, and then he continues his explanation by saying that as you go across the period, the atomic size gets small and all that... OK that's fine, but, aren't we talking about ions here, a sodium ion would be much smaller than a sodium atom (because it has lost its outer electron, and hence its outer shell) whereas a chloride ion would be slightly larger than its atom, due to greater shielding from inner electrons.... so a chloride ion is larger than a sodium ion!! This is in fact, true:
Na+ has an ionic radius of 0.102 nm
Cl- has an ionic radius of 0.181 nm(6 votes)
- You are correct, as there is some confusion in the video. For atoms the atomic size decreases, in general, from groups 1 to 7, but for ions as you have stated the sodium atom will have lost its' valence electron and thus the outer shell, dropping the size down to the smaller inner full shell. While the chloride ion will have gotten larger than its' original size as an atom due to the greater electron repulsion WITH the same unshielded nuclear charge.(11 votes)
- Just curious at3:59Is Oxygen really that much bigger than Hydrogen?(7 votes)
- This is actually a very difficult question, because it depends on how you define the "size" of an atom. As a point of reference, however, the covalent radius (http://en.wikipedia.org/wiki/Covalent_radius) of oxygen is roughly twice that of hydrogren.(5 votes)
- if NaCl salt could conduct electricity after being dispersed in water than why oftenly do we use H2SO4 as an Electrolyte in electrochemistry reactions.(5 votes)
- What's important is to use a strong electrolyte, so a good conductor of electricity. Sulfuric acid, H2SO4, is one of the best solutes for conducting electricity because it is a strong acid and electrolyte. NaCl salt is strong electrolyte too and can also be used (but H2SO4 is even stronger!)(4 votes)
- Does on increasing temperature increases the solubility in all types of fluids ?(4 votes)
- As a general rule, the solubility of a solute increases with increasing temperature of the solvent, but there are exceptions. Most notably, gases will generally decrease in solubility on increasing the temperature of the solvent, such as O2 and CO2 in water.(6 votes)
- How does the weak partial positive and negative charge from the water molecules able to pull off the strong ionic bonds between the Na and Cl.(3 votes)
- The ionic bonds in NaCl are not very strong at all, but that is irrelevant.
There is more than one water molecule acting on each ion of Na⁺ and Cl⁻. So, although a single molecule of water wouldn't be up to the task, multiple molecules of water acting together can do the job.
Thus, the Na is surrounded by the oxygen side of a number of water molecules and the Cl is surrounded by the hydrogen side of a number of water molecules. As this happens, the net charge between the Cl and Na weaken (being partially removed by each water molecule) until such a time as they are no longer sufficiently attracted toward each other and dissociate.(4 votes)
- What about precipitates? Do they form when the compound doesnt have enough energy to break apart the ionic bonds?(3 votes)
- Almost correct. They form when the solvent doesn't have enough energy to break apart the ionic bonds.(4 votes)
- [Instructor] In this video, we're going to talk about solubility, which is just a way of describing how well certain solutes can dissolve in certain solvents. And just as an example, we could go to our old friend sodium chloride and think about why does it dissolve well in water. Well, to do that, you just have to remind yourself what water is doing when nothing is dissolved in it. So when nothing is dissolved in it, that's an oxygen attached to two hydrogens. This end, we've talked about in other videos, partially negative 'cause the electrons like to spend more time around the oxygen. This is partially positive, partially positive. And you have this hydrogen bond forming, where the partially negative end of one water molecule is attracted to the partially positive end of another water molecule. That's what hydrogen bonds are all about. And we have whole videos on that, actually many videos on that. The reason why sodium chloride dissolves well in water is because sodium chloride, as an ionic compound, it can disassociate into its constituent ions. Into a sodium cation and a chloride anion. And we've seen this before. So that's the chloride anion, this right over here is a sodium cation. The reason why this dissolves well is that the negative charge is able to be drawn to the positive end of the water molecules, so the hydrogen end. And the positive sodium cation is attracted to the negative end of the water molecules. So what's really happening is the attraction between the ions and the water molecules are stronger than the attraction between the ions themselves and the attraction between the water molecules themselves. So the water molecules, don't just bunch up and say, "We want nothing to do with you, sodium and chloride." They say, "Hey, we're kind of attracted to you too. So why don't we mix together?" We can look at things that have less attractive forces, maybe things where the main force is just dispersion forces. If you think about a vat of pentane here, they have those weak forces, kind of attracting them to each other, and then if you think about a vat of hexane here, there's kind of weak forces. But if you were to put some pentane, let's call this the solute here. And if you were to put it into a solvent of hexane, it will dissolve because they are roughly as attracted to each other as they are to themselves. Now, what do you think is going to happen if I try to put, say, some hexane, if I view that as a solute, and I were to put it in water? Well, in that situation, the water is going to be far more attracted to itself than it's going to be attracted to the hexane. Let's say that this is the water here. You're going to have these globs of the hydrocarbon form because the water is more attracted to itself. It's not easy for the hydrocarbon to dissolve. Now, there are many organic molecules that do dissolve well. And that's usually because they have some part of the molecule that has some polarity to it. One example is ethanol, which has an OH group. But ethanol, which has a chain of two carbons, when we talk about alcohol in everyday language, drinking alcohol, that is ethanol. There's many other anols, many other alcohols, but this is ethanol here. And if you were to take alcohol and you were to mix it in water, it does dissolve well. And that's because this oxygen here, it's more electronegative than the things that it is bonded to. And so it still hogs electrons a little bit. And so you still have that partially negative charge. You still have that polarity to the molecule. And so that's able to attract it to neighboring water molecules, which allows it to dissolve. But you could imagine if you had an alcohol that had a much, much longer carbon chain, so you had 10 carbons or 15 carbons, then all of a sudden the relative proportion of how polar it is compared to how large of the molecule it is, it'll make it harder and harder for it to dissolve in a polar solvent like water. And so our mega takeaway here is that like dissolves like.