Previously, I've talked about 3 kinds of strong bonding: ionic bonding, metallic bonding, and covalent bonding, and in the covalent bonding post, I wrote "there are other random types like hydrogen bonds, but they're not important for now." Well, hydrogen bonds are now important, and so are other "random types" of bonds like dispersion and dipole-dipole bonds.
I said that I wasn't going to completely adhere to the dot points, but I didn't mean that I wouldn't use them at all. I'm still going to do all the dot-points that are relevant to intermolecular bonds.
Explain polar and non-polar covalent bonds in terms of the electronegativity of the atoms involved in the bond formation
Wow, that sounds so high-brow and sciencey. Not to worry though. Or maybe you should, because my explanation for this is probably going to be a little shoddy? I don't know.
Now, you know how electrons are shared between atoms in a covalent bond? (If not, revisit my post on covalent bonding.) Well, here's the catch: from what I understand, they aren't always shared equally. They're shared equally enough so that both atoms get to have a stable full outer shell configuration, but sometimes those pesky electrons decide that they like one atom more than another. That's due to electronegativity differences: the more electronegative element attracts the electrons more and therefore the electrons spend more time near that atom than the other atom. This results in a polar covalent bond, where one atom is more positively charged than the other. How "positive" or "negative" each end is depends on the difference in electronegativities between the two atoms (the periodic table on my phone- yes, I'm so nerdy that I have a periodic table on my phone- gives information on the electronegativities of atoms of each element). When there's a super big difference in electronegativities, the electrons spend pretty much all their time around the negative end, which results in an ionic bond. (My Chem teacher was saying how it's like one atom is "stealing" electrons from the other, which is like "extreme sharing." Pretty damn good joke, if you ask me.)
Non-polar covalent bonds occur when both atoms on either side of the bond are of the same element because then there's absolutely zero difference between the electronegativities of each atom.
Use the relationship between molecule shape and bond polarity to predict and explain the polarity of a molecule
I'll explain this better later, after I've discussed molecule shape, but basically, if the atoms on one side of the molecule tend to be on the "negative" side of polar covalent bonds and the atoms on the other side tend to be on the "positive" side of polar covalent bonds, then the molecule is probably polar. If there's no general trend like this, then the molecule is probably non-polar. Yeah, I really need a diagram to explain this properly.
Explain the differences between intermolecular and intramolecular forces.
This is easy (unless there's more to it than I'm aware of). Intermolecular forces are forces between molecules. Intramolecular forces are forces between the atoms that make up the molecules. In covalent molecular substances, the intramolecular covalent bonds are strong while the intermolecular Van der Waals forces tend to be weaker.
Describe and explain the origin and relative strength of dispersion forces, dipole-dipole attractions, hydrogen bonds and ion-dipole interactions for molecules of a similar size.
Dipole-dipole attractions are seen only in polar molecules. This is because one side of a polar molecule is more positively charged while the other side is more negatively charged. The more positively-charged side of one molecule can attract the more negatively-charged side of another molecule. Dipole-dipole forces get stronger as the difference in electronegativity between atoms within the molecule increases (because that would cause the ends to be "more positive" and "more negative").
Hydrogen bonds are kind of like dipole-dipole attractions, but with a catch: they only exist when there's an O-H, N-H or F-H bond somewhere in the molecule. Hydrogen bonds are much stronger than dipole-dipole attractions (though still not as strong as ionic, covalent or metallic bonds).
Dispersion forces are seen in all kinds of molecules. In a covalent bond, the electrons do not always remain in the space between the two atoms. Instead, the electrons move around (though with a net movement of zero). Therefore, there are times when there are temporarily more electrons on one side of the bond (and the molecule) than the other. This causes a temporary "positive" dipole and a temporary "negative" dipole. Oppositely-charged temporary dipoles of different molecules can attract each other, much like in dipole-dipole forces, but in the case of dispersion forces, these attractions are not as strong since the dipoles are only temporary. Dispersion forces get stronger as the number of electrons in a molecule increase.
Ion-dipole interactions occur when soluble ionic solids are dissolved in aqueous solution, particularly when said solution is polar. The positive ions of the solid can be attracted by the negative dipole of the molecules in the solution, and vice versa. This causes the ionic solid to dissolve. Complex ions (not really sure how to explain these) are formed by such ion-dipole interactions between ion(s) and whatever other molecules make up the complex ion.
Explain the relationships between physical properties including melting and boiling point, and the types of intermolecular forces present in substances with molecules of similar size.
Melting and boiling involves the breaking of intermolecular bonds between molecules, so stronger bonds result in more heat energy required which in turn results in higher melting and boiling points, and vice versa. Of the types of intermolecular bonds discussed above (except I'm not going to talk about ion-dipole for now since that type of bond doesn't occur between molecules of a covalent molecular substance), hydrogen bonding is the strongest. Dipole-dipole is stronger than dispersion if the molecules are small and don't have that many electrons, but in bigger molecules it might be the other way around. In fact, often it will be the dispersion forces that have a greater effect on melting and boiling points rather than dipole-dipole forces.
Apply an understanding of intermolecular interactions to explain the trends in melting and boiling points of hydrides of groups 15, 16 and 17 accounting for the anomalous behaviour of ammonia, water and HF.
Yay, now I can provide an example for my previous point! Let's use the group 15 hydrides as an example. In group 15 the first four hydrides are ammonia, PH3, AsH3 and SbH3. Ammonia has the highest boiling point, then PH3 has the lowest boiling point. After that, the boiling points gradually increase.
Ammonia has the highest boiling point because it contains hydrogen bonds between molecules due to the N-H bonds within ammonia molecules. The other three do not contain hydrogen bonds so their boiling points are not as high, but as the size of the molecules and their numbers of electrons increases, their dispersion forces also become stronger and their boiling points gradually get higher.
Similar trends are observed in groups 16 and 17.
Explain and describe the interaction between solute and solvent particles in a solution.
When a solute dissolves in a solvent, the bonds between solute particles and the bonds between solvent particles break. Then, new bonds are formed between solute and solvent particles.
Now, of course, breaking bonds requires energy, so in order to make up for that, the amount of energy that is released when bonds are reformed must be close to the amount of energy that was required in the first place. Therefore, dissolving generally only happens when the strength of the bonds within solute and solvent are reasonably similar.
Use the nature of the interactions, including the formation of ion-dipole and hydrogen bonds to explain water's ability to dissolve ionic, polar and non-polar solutes.
Now I have an excuse to provide an example for my last point!
Water contains hydrogen bonds. If you want to dissolve something non-polar in water that only has dispersion forces between molecules (since non-polar covalent molecular substances can't contain dipole-dipole or hydrogen bonds), we have a problem. You see, hydrogen bonds between water molecules would have to be broken (as well as dispersion forces between the non-polar molecules), but when bonds form between the non-polar substance and water, only dispersion forces would technically be formed as the non-polar substance wouldn't be able to bond to the water using dipole-dipole and/or hydrogen bonds. Therefore, a lot of energy would technically be required while very little would be released: a rather selfish reaction, which is why it generally doesn't happen.
The dipole-dipole bonds in polar substances are closer to the strength of hydrogen bonds, so polar substances are generally able to dissolve in water.
Finally, many ionic substances can dissolve as a result of ion-dipole reactions, which I mentioned earlier but I'll say it again. And, because I'm really lazy, I'm just going to copy-paste it from above, so you can skip over this paragraph if you've already read it.
Ion-dipole interactions occur when soluble ionic solids are dissolved in aqueous solution, particularly when said solution is polar. The positive ions of the solid can be attracted by the negative dipole of the molecules in the solution, and vice versa. This causes the ionic solid to dissolve. Complex ions (not really sure how to explain these) are formed by such ion-dipole interactions between ion(s) and whatever other molecules make up the complex ion.
Describe the variation of gas solubility in aqueous solution with temperature
Okay, this is a bit that I'm going to have to revise because I don't really remember. I know that gas solubility tends to decrease with temperature, and I think that is because the added temperature would just cause the gas to evaporate out of the solution. I don't know for sure though, so I guess I'll have to look at my textbook again.
And, of course, this topic seems to be pretty damn elusive in my textbook, so I guess I'll probably have to Google it or search a bit harder. I'll do that later. (Yeah, I am pretty lazy.)
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