Why aren't all crystal lattices 6:6 or 8:8?

Lets start off with a reminder why NaCl is 6:6 and compare thaty to CaCl2

The NaCl the structure is 6:6 because in that lattice the structure the ions take up is controlled by the size of the smaller Na+ ion and how many Cl- ions you can fit around it, (i .e. Na+ is small so you can only get 6 Cl- around the Na+).

The other number 6 (in 6:6) comes from the fact that to maintain the 1:1 ratio that NaCl must have, you need 6 Na+ around each Cl- too. You could get a lot more Na+ around the Cl- but that would spoil the 1:1 ratio.

(Why must NaCl be 1:1?

Sodium has 1 electron in the outer, which it needs to lose to form a stable ion. Chlorine needs to gain one to from a stable ion. So 1 sodium needs to lose 1 electron to 1 chlorine to be stable. Happy Days.)

Now, lets apply the same logic to CaCl2. Whatever structure CaCl2 takes, it must maintain the 1:2 ratio (for the same reason as above).

The 6:6 structure maintains a 1:1 ratio so it would never work for CaCl2.

You don't need to know the structure CaCl2 takes, all you need to know is, it is not 1:1, so can't be 6:6 (i.e. the same as NaCl.)

Bonding Types - What is that all about?

Answering bonding questions in CH2 is a really big problem. How do you know which bonding type a molecule has? How do you describe it? Here is a quick guide to enable you to at least start to answer these questions.

Lets think through first what type of bonds the molecule has.

If it has ONLY METAL atoms it has ONLY METALLIC bonds. Metallic is fairly straight forward higher charged ions = stronger attraction to electrons = higher BP/MP. Easy.

If it has METAL AND NON-METAL atoms it has ONLY IONIC bonds. Again straightforward you need to know your CsCl and NaCl stuff here and the forces of attraction and repulsion. Not too tricky.

If it has ONLY NON-METAL atoms it has COVALENT bonds IN the molecule and one/two or all three of the INTERMOLECULAR forces BETWEEN the molecules. Now this is the one that trips people up.

There are so many possible molecules here that you can't learn them but you can work out which sort of bonds they have.

Before I start there is a difference between what bonds a molecule has and which ones are important, for example water can do ID-ID but they are not important because the hydrogen bonds are so much stronger that only they matter. So here are the intermolecular bond types in order of importance from least to most.

1. Instantaneous Dipole-Induced Dipole (ID-ID)

If a molecule has...
...only one sort of atom (e.g. Cl2)
...all the same sort of atom on the outside (e.g. CH4)

...then the only sort of intermolecular force it can do is Instantaneous Dipole-Induced Dipole (ID-ID). This happens because all the electrons are swishing around and creating temporarily positive/negative ends of the molecule that then attract other molecules. As molecules/atoms get bigger there are more electrons and therefore stronger ID-ID, therefore, higher BP/MP. All molecules that have covalent in the molecule will have ID-ID between molecules but they are only important when the molecule can't do the other two types of force.

Don't forget then that all covalently bonded simple molecules can do this bond it is just that it only becomes important when it is the only bond that they can form (i.e. in non-polar molecules)


2. Dipole-Dipole (D-D)

If a molecule has a permanently positive and a permanently negative end then that molecule has a dipole, in other words it can attract molecules/atoms/ions of the opposite charge towards it...permanently. It is like the one above but it doesn't change. More electrons don't make a difference now because it is not down to swishing of gangs of electrons. This bond just gets stronger when dipoles get bigger because electronegativity differences become bigger.

On this one just be careful with shapes, e.g. NH3 might look non-polar because it has all the same type of atom on the outside but it is polar. When you look at the shape it is trigonal pyramidal with the N at the top point and all 3 Hs at the other three points.  So there is a negative end (the N) and 3  positive ends (the Hs) so it will do D-D (as it happens it will also hydrogen bond but that is a different story)

3. Hydrogen Bonding

This is the strongest of the three. This is a special case. IN THE MOLECULE, you need to have a very electronegative atom with an active lone pair directly attached to a hydrogen, i.e. in the molecule there must either be a H-F,H-O or H-N bond.

The bond will then be formed between the N, O or F of one molecule and the H of the other.

Your obvious examples of molecules that can Hydrogen bond are water, ammonia and HF but there are lots of other.

If you want a quick summary, here goes

Metal atoms only - Metallic

Metal and Non-Metal atoms - Ionic

Non-Metal atoms - Covalent IN the molecule and then BETWEEN the molecules...

Non-Polar molecules - only ID-ID
Polar molecules - D-D and ID-ID
A molecule that has N-H, O-H or F-H bonds - ID-ID, D-D and hydrogen bonds

(I have emboldened the one that matters in that last statement)

Simple? Not really that hard once you get your head around it.
Worth understanding and working on? Definitely as it is guaranteed to come up in CH2.

Why are bonds different lengths?

(above) Not a bond

It must have crossed your minds at some point why are some bonds longer than others. Why is a hydrogen bond long, a covalent bond short and a double covalent bond shorter still. 

The answer is simple.

Bonds aren't physical things like pieces of string they are just attractions. The stronger the attraction the shorter the bond. So, simply covalent bonds are stronger than hydrogen bonds so they are shorter. 

Isn't It Ionic?

Ionic or Covalent? What is electronegativity?

This is a fairly confusing area of A level Chemistry (Alanis Morissette certainly found it hard) so I thought I would put together a blog post to explain it.

Think of it this way, we need to measure the ability of atoms to draw electrons towards them (i.e. electronegativity) the only way of doing this is on a comparative scale, that is, compared to other atoms. Then all somebody did (the man who did it was called Linus Pauling) was to give this scale values. The reactivity series from GCSE is the same principle except that nobody has ever bothered to give reactivities values so it is still just a list.

Anyhow, if you are going to compare atoms electronegativity (and therefore assign values) you need to get in to a situation where two atoms are both trying to pull electrons towards them and see which one has the bigger pull.

The only place this happens is in a covalent bond. In a covalent bond the electrons are shared between the two atoms but the electrons will spend more time with the atom that has the greater pull and therefore we say that atom is more electronegative (than the other one it is bonded to).

So, by comparing the covalent bonds between every possible pair of atoms you get a hierarchy of pulling ability (electronegativity) which Linus Pauling assigned numbers to. On Pauling's scale, Fluorine is the highest (best at pulling electrons) and is assigned a value of 4.0 and Francium is the lowest (poorest at pulling electrons) at 0.7. All of this measurement is done by comparing pulling ability in covalent bonds.

So the definition has to be something along the lines of...

"Electronegativity is a chemical property that describes the tendency of an atom to attract electrons towards itself in a covalent bond".

So, "How can compounds that form ionic bonds have electronegativities?" I hear you ask.

Electronegativity is a property of an element, if the value is high (e.g. Chlorine, 3.2) the element will form covalent bonds with other atoms of high electronegativity, because they are both pulling really hard on the electrons and so end up sharing them, but will form ionic bonds with atoms of low electronegativity (e.g. Na, 0.9) because the chlorine has so much of a bigger pull on the electrons compared to the sodium that it ends up pulling the electrons towards itself completely and forming a negative ion, whilst the sodium loses its electron and ends up being a positive ion.

So, don't think of electronegativity as happening "in bonds". It is a property that controls what sort of bonds an atom will form.

The only loose end to tie up is how do we know sodium's electronegativity if it never forms covalent bonds (pure sodium forms metallic bonds and sodium compounds have ionic bonds). The answer to this is that even though sodium doesn't naturally form covalent bonds, in a lab it can be made to share electrons (i.e. form a covalent bonds) with other atoms so that its electronegativity can be measured.

My analogy for all this is tug of war competitions, I won't expand on it now but it does work pretty well.

....and in future Alanis, just buy less spoons.