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.

Q - When is a triangle not a triangle?

A - When it's a pyramid
There are two molecular shapes that could be described as triangular (or trigonal). One is a...errr...a triangle and it is referred to as trigonal planar. The other is effectively a triangular based pyramid, called equally obviously, trigonal pyramidal. The reasons behind the differences in the two are important.

Trigonal planar is flat (i.e planar) so imagine a triangle with three atoms at each corner and one in the middle. Each of the corner atoms are 120degrees apart (i.e. 360/3 = 120). You get trigonal planar when you have three bonded pairs and no lone pairs as that is the furthest the electron pairs can get away from each other. an example is BF3

Trigonal pyramidal is like...erm... a pyramid but a pyramid with a trigonal base (hence the name). The bond angles are 107degrees (no way of working that out you just have to remember). You get trigonal pyramidal when you have 3 bonded pairs and 1 lone pair. The lone pair is more repulsive and pushes the bonded pairs down and  away from being trigonal planar (which is what you would get if the lone pair wasn't there). An example is NH3

P.S. The purple lump on the trigonal pyramidal isn't a comedy hat it is a lone pair of electrons.

Oxidising Agents and Reducing Agents

OK, you have two chemicals, one is being oxidisied (losing electrons) and the other is being reduced (gaining electrons). Both must be happening at the same time. Otherwise where are the electrons going? If you want an analogy, there is no thief that has no victim and no victim of theft without a thief.

So, the species being reduced is always being an oxidising agent (it is causing the other thing to be oxidised) and the species being oxidised is always being a reducing agent (it is causing the other thing to be reduced), No exceptions. 

Does it matter whose fault it is? Does it matter that one species really wanted to gain an electron while the other was fairly ambivalent whether it lost one or not? Conversely, does it matter that one species really wanted to give one away and the other species that happened to be passing was just a willing recipient?

Quite frankly, no. It doesn't matter whether it really wants to or no, the species that loses the electron, causes the other to gain so is a reducing agent and the species that gains the electron caused the other one to lose an electron so is an oxidising agent.

There is no blame game where it comes to redox they are all agents.

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. 

Why do you get that 180 bond angle between water molecules?

Right this is difficult to explain without a board pen and some pointing but here we go.  

The 180 degrees I am referring to here is for the H that is covalently bonded to an O on one side and Hydrogen bonded to an O on the other (see picture). 

This hydrogen has a covalent bond on one side (i.e. a bonded pair) and a hydrogen bond on the other. Now, technically the hydrogen bond is not a bonded pair of electrons but the oxygen that the hydrogen is hydrogen bonded to is a big ball of negativity (imagine a fat dementor) that behaves much like having a pair of electrons on that side too. So effectively that hydrogen has 2 bonded pairs, and what shape do molecules with 2 bonded pairs take...straight line. 

So there we go, that is your answer. The actual answer in the exam goes along the lines of...

"...because it has two bonded pairs of electrons and to take the position of minimal repulsion the electron pairs  get as far away from each other as possible which is a straight line."

What's HOT and What's NOT in the world of Chemistry (or Spot the Difference)

Whenever they change specification (like they did 4 years ago) they always take some stuff out and put other new stuff in. 

There is no good reason why, except to confuse pupils and teachers . So, those of you doing the old specification papers will be very confused from time to time when they ask a question that looks like it is written in greek because you don't recognise the stuff at all or you didn't think you needed to know in that level of detail.
So, to help you through, here is a list of all the changes that come to mind. I may have missed something, so please point in out if there is something else that you think is different.

AS
New In!
Hydrogen Emmission Spectrum
Smart Materials
Carbon Nanotubes
Nanotechnology
Green Chemistry

Gone!
Electron Density Distribution

Equilibrium Constants (now in A2)

A2 

New In!

Hydrogen Fuel Cell

Group 3

Chromatography

Gone!

Group 1 and 2

Hydrogen Emission Spectrum (now in AS)

Testing, Testing, 1, 2, 3... (CH4/CH2)

One of the hardest aspects of organic chemistry is remembering all the wet tests that you need to carry out to differentiate between different functional groups that you are likely to come across. 

Not only do you have to remember the compounds that give a positive (and the ones that don't), you need to  remember the observations, the chemicals you need to add to get this thing to work and some seemingly random facts that someone deemed important. 

To help you through this, here is a list of all the tests you are likely to come across and an attempt at the relevant detail. I have also indicated whether these are relevant for CH2 or CH4

Tollen's Reagent (CH4), Fehling's Reagent (CH4), Acidified Dichromate (CH4 and CH2)
These three are collectively known (by me at least) as the oxidation tests, each depends on a molecule being oxidised. So molecules that are easily oxidised give a postive, namely aldehydes, primary alcohols and secondary alcohols. 

The only thing that changes is the observation.

So for Tollen's Reagent you see a silver mirror (or more likely if your test tubes could do with a wash - a grey ppt)

For Fehling's Reagent you see a orangey brown ppt

For Acidified Dichromate you see the orange solution go green

Other facts worth remembering are...

• Tollen's reagent is fairly rubbish as an oxiding agent so it is not strong enough to work with primary alcohols (such as ethanol)
• Fehling's is almost the same Benedicts solution (remember from biology, the test for reducing sugars)

Iodoform Test (CH4)

Iodoform (or triiodomethane or CHI3) is the name of the test but in this case it is not the name of the reagent you add. In fact, Iodoform is the yellow stuff you see at the end that tells you the test has worked. 

The reagent you add is aqueous alkaline iodine (or iodine, mixed with aqueous sodium hydroxide) it also works with Potassium Iodide and Sodium Chlorate (I), You generally don't need to warm it but if its being a bit slow a bit of heat helps.

The observation is a yellow ppt (of Iodoform) that smells of ...erm...Iodoform which is basically an antiseptic smell, which makes sense as iodoform used to be used as antiseptic.

People often come unstuck in what gives a positive. Positives are given by methyl ketones or methyl hydroxyls. That is -OH or C=O groups being on the carbon next to a CH3. For example -  ethanal, ethanol, propan-1-ol, Propanone  all give positives but methanol, propanal and pentan-3-one all give negatives. Think this one through, draw the molecules and try and work it out, if you are still stuck have a look here. 

2,4 DNPH (CH4)

Nice and easy one this one. Positives are given by aldehydes and ketones and it looks like a bright orange ppt. The only other factoid that they want you to know here is that this is an example of an addition-elimination reaction. Don't bother learning the mechanism just remember that fact.

Lucas Test (CH4)

This is technically not in the syllabus so they can't ask a question where this is the only answer and it won't be in the mark scheme as a possible answer but it does work and could get you out of jail if you don't know the actual answer but if you can don't use this as your answer of choice as it won't be in the mark scheme and a half asleep examiner may well mark it wrong if they aren't concentrating.

Anyway (briefly), what is it? Reagent is Conc Hydrochloric Acid and Anhydrous Zinc Chloride. Tertiary Alcohols give a white ppt v quickly, secondary alcohols give a white ppt v slowly. Primary Alcohols don't give a white ppt but neither does anything else so that fact doesn't help you much.

Sodium Carbonate Test (CH2 and CH4)

Simple idea this, acids bubble with metal carbonates. So if you add sodium carbonate solution to any acid (including carboxylic acids) you will get bubbles that turn lime water milky. 

Tests for Phenol (CH4)

Two of these, either... 

•  add Iron (III) Chloride and the solution goes purple

            or

• add bromine water and it decolourises and you get a white ppt

Enough said, lets move on.

Test for Alkenes (CH2 and CH4)

This is an easy one you should all know from GCSE. Alkenes decolourise bromine water. In other words it goes from being a browny-orange colour to being colourless (NOT CLEAR!!!)

Test for Amines (CH4)

This is a spin off from all that di-azo stuff that you do. When you add nitrous acid to an amine it forms a di-azo compound that is really unstable so it immediately breaks down to give off bubbles of nitrogen gas

Test for Amides (CH4)

Again this is a consequence of a conversion reaction. When you sodium hydroxide to an amide and heat it a bit it gives off ammonia. You will know because it stinks but if your nose isn't working then test it with damp red litmus paper and it will go blue. The only possible confusion here is with ammonium salts which also give off ammonia with sodium hydroxide but that will happen in the cold, so it is usually possible to tell the difference.

Test for Halogenoalkanes (CH4 and CH2)

Halogenoalkanes are known for being unreactive so the only way to test for them is to chop a bit off and test that. So, add some sodium hydroxide, this will then release the halide ion. 

The halide ion can then be tested for by adding a mixture of silver nitrate and nitric acid (the acid is just there to neutralise any excess sodium hydroxide from earlier).

The observation will depend on the halide ion - so white ppt means chloride (so it was a chloroalkane) cream ppt means bromide (so it was a bromooalkane) and yellow ppt means iodide (so it was an iodoalkane).

There we have it all the test in one place. If I have missed anything please leave a comment and I will add it. 

Y12 if you are getting stressed by the 4 you have to learn for CH2 spare a thought for Y13 who had to learn all of these others for CH4, as will all of you next year...  

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.

Crystal Coordination Number

Background

Ionic compounds are made up of + and - ions (Na+ and Cl-) in a regular arrangement that is arranged so that + isn't next to a -. You would've thought that there is one or at best two or three ways of doing this, actually there are loads (visit this link to see 14 of them http://en.wikipedia.org/wiki/Crystal_structure#Lattice_systems of them) depending on the ratio of + to - ions and the relative sizes of + and - ions. Luckily for you, you only need to recognise and explain about two of them.

The two structures - NaCl and CsCl 

Think of a theoretical situation where all the ions are exactly the same size like in a box of bowling balls. When you have a box of bowling balls (or any identically sized spheres) the natural arrangement they will settle in will have each bowling ball being touched by 12 others (this is the maximum number of spheres that can touch a central sphere when they are all the same size, trust me its true!). 

Now to really exaggerate the situation imagine a similar scenario where you have a mixed bag of bowling balls and table tennis balls in a 1:1 ratio and for some weird reason when they arrange themselves naturally the bowling balls can only be touched by table tennis balls (and vice versa). Now there is a problem, because bowling balls are bigger than table tennis balls you can get maybe 50 around a bowling ball but in the reverse situation you can only get maybe 3 bowling balls to touch a table tennis. So in this theoretical box of bowling balls and table tennis balls there is a problem if you are to maintain this 1:1 ratio we can't have 50 of 1 and 3 of the other. So, the structure will have to take the smaller number of 3 as the number of each that touch each other. 

So now we have a pattern, when spheres are identically sized we can fit 12 around the opposite one, when they are different sizes this number decreases (down to possibly 3 for bowling balls and table tennis balls). This is called crystal coordination number. i.e. bowling balls have a crystal coordination number of 12. Mixed bowling balls and table tennis balls have a crystal coordination number of 3

Right to ions. 

Cs and Cl are v similar sizes (think bowling balls) so the crystal coordination number will be high (i.e. 8). Na and Cl are different sizes (think mixed bowling balls and table tennis balls) so the crystal coordination number must be lower (i.e. 6). 

This means that in the NaCl crystal there are 6 Nas next to every Cl and 6 Cls next to every Na.

The total crystal coordination number of the crystal is 6,6.  The crystal coordination number of sodium in the crystal is 6 and the crystal coordination of chloride in the crystal is 6.

What does NaCl and CsCl look like?

I am not going to try and draw this here so look at your notes or google them!

What are they going to ask you in the exam?

Firstly, there are the obvious questions like, what do they look like? what is the crystal coordination number? but what about the more descriptive questions that will involve a bit more thought.

1. Why do they have different crystal coordination numbers?

Cs is bigger than Na, so you can fit more Cl ions around a Cs than an Na. Hence CsCl is 8,8 and NaCl is 6,6

2. Why can't MgCl2 have a 6,6 coordination number?

In MgCl2 there are twice as many Cl ions as Mg ions and the crystal coordination number must reflect this, so the crystal coordination might be 4,8 or 6,12 or any structure that involves there being twice as many of one ion compared to another.