MODERN ANALYTICAL METHODS

Hello everyone! Nearing the end slowly, and with Breaking Benjamin a comforting sound in my ears, it will fly by.

All molecules absorb infrared radiation. This energy that they absorb causes the bond to oscillate in a certain way. It could either be stretching back and forth or a bending of the bonds of to the side.

This absorption shows up on infrared spectrum, and enables us to identify the presence of functional groups. Cool, eh??

You will be given a table in the exam which gives you absorption data and identifies the responsible functional group. Handy. Learn to understand it and identify what's where.

Learn it, learn it and once again learn it! Not of by heart of course but make sure you know what you see when you do see it.

Infrared spectroscopy is very useful and is now used in breathalysers to measure how much alcohol is in your breath.

Next we have mass spectrometry.

It is dead easy. An electron gun fires at a molecules and breaks it up. Just like glass shattering different fragments are formed, big and small. These show up as peaks on the graph. However, unlike glass, the fragments are cations (positive ions), which get acted on by a huge magnet. The further they get pushed away the smaller their mass was (if a person gets hit by a truck, they go flying but if the truck hits a tree - it'll move but slightly). How far they get pushed enables scientists to determine their mass. 

Now the peak furthest to the right (biggest m/z value) is the molecular ion. This bad boy is pretty lucky as the gun managed to hit one of the inner electrons and hence the molecule didn't break up at all. It is the molecule you were testing for but charged. Simple??

Once you have the mass you play around with numbers and figure out what it is. 

Hey! No sad faces... Not on my watch.

Any questions? Ask, I don't bite... Much.

HALOGENOALKANES

Oh my sweet oranges! I am starting to go crazy with all this revision. You too?? Too bad.

We have already covered many things about halogenoalkanes so here is a short summary.

• Electrophilic addition of hydrogen halides.
• Addition of halogens.

Now you will learn how to name them too. Isn't that epic??

Halogenoalkanes (geez, what a mouthful) contain only single bonds. We name them by first identifying where the halogen is attached. Count on which carbon it is on (smallest number counts). Write the number of the carbon then a relevant prefix. Finish off with the name of the alkane.

Halogen:           Prefix:

F                  flouro-

Cl                 chloro-

Br                 bromo- 

I                  iodo-

NUCLEOPHILE: An electron pair donor. 

NUCLEOPHILIC SUBSTITUTION/HYDROLYSIS of HALOGENOALKANES:

Because it's a substitution, you know that on top of the halogenoalkane you are going to have something else left over. That which has been replaced.

Think about it this way. In electrophilic addition we added the halides. Here we will remove them. See? Opposite.

You might come across questions where it will ask which hydrolysis reaction will take place faster. This rate is affected by two factors:

• Polarity - C-F bond is the most polar (greatest difference in electronegativity) so the C delta+ atom attracts the nucleophile more readily - So faster reaction.
• Bond enthalpy - The C-I bond is the weakest (I don't think you need to know why but I'll include it later in this post - probably at the end somewhere) so it is broken more easily. This means a faster reaction.

Faster as you go up. Then faster as you go down. One point contradicting the other? It's not a mistake, I assure you. The think is that in hydrolysis bond enthalpy is more effective than bond polarity. So you should keep the polarity in mind but focus on the bond enthalpy.

• Industries had a craze when it come to compounds containing chlorine. Such as, chloroethene and tetraflouroethene (these coming from radical substitution), especially CFC (ChloroFlouroCarbons).
• These were used for aerosols, air conditioners. This is because they were relatively unreactive, non-toxic and volatile. But every silver lining has a cloud to it (Sorry, I just had to). CFC's had a huge impact on the ozone layer (about that later - I promised).

There are now biodegradable alternatives to the killer CFC's, such as HCFC's and even CO2 as a blowing agent for expanding polymers.  

Now back to bond enthalpy. The C-I bond is the weakest because the outer electrons of Iodine are very far away, this means nuclear attraction is very weak and the electrons can be easily stolen away. You knew this was coming, didn't you?

ALCOHOLS

Don't get any ideas, this isn't going to be a page filled with different types of alcohols. Well, maybe a little. But not in a way that you think.

First up: PROPERTIES OF ALCOHOLS.

It is, oh, so clear that alcohols can form hydrogen bonds. You know, the electrostatic attraction between the electron deficient Hydrogen and lone pair on the oxygen atom. Course you do. 

Hydrogen bonds are the strongest intermolecular (between molecules) forces, this results in relatively high melting/boiling point.

This also means that alcohols have a relatively low viscosity (so difficult to form a gas).

The hydrogen bonds give alcohols another property. Alcohols are soluble in polar solvents. This is because a hydrogen bond can form between the lone pair on oxygen and the electron deficient hydrogen on a water molecule.

As the chain length increases the solubility decreases. This is because a larger proportion of the molecule is not polar.

Next up: PRODUCTION OF ALCOHOLS.

Alcohol can be produced in two main ways, that you need to know.

Hydration of ethene. Sounds scary? It's not. It is a simple reaction between ethene and steam (recognise it?). A H3PO4 catalyst, 300 degrees and 60 atm are required. This is a reversible reaction (more about this later).

Fermentation. We have a carbohydrate that is acted upon by yeast to form the beloved alcohol - not personally - (wine for example) and carbon dioxide.

C6H12O6 -----> 2C2H5OH + 2CO2

CLASSIFYING ALCOHOLS: This is my method of remembering this.

Primary alcohols: When the OH group is on the end carbon in a chain, and no groups are attached.

Secondary alcohols: OH group is in the middle somewhere but no group is attached to that carbon.

Tertiary alcohol: OH group in the middle with an alkyl group attached to the same carbon.

The textbook phrases it differently and actually talks about alkyl groups attached but this works for me and actually gets me the marks.

Here is what the textbook says, if you think about it hard it does make sense. My way seems simpler.

PRIMARY: OH attached to a carbon with no alkyl groups or bonded to one alkyl group.

They take the hydrocarbon chains as alkyl groups.

SECONDARY: OH group attached to a carbon bonded to two alkyl groups.

TERTIARY: OH group attached to carbon with three alkyl groups.

SO here there is no hydrogen next to the carbon.

COMBUSTION OF ALCOHOLS:

Alcohol + O2 ------> Carbon dioxide + Water

OXIDATION OF ALCOHOLS:

This is a summary of all the oxidations. The R here means the rest of the chain.

When alcohols are oxidised during distillation only the first product is formed. During reflux (allowing the reaction to continue so all reactants can react).

The alcohols are oxidised using Metal dichromate XCr2O7 under the presence of sulphuric acid catalyst. 

We can get rid of the OH group to form alkenes, in a reverse reaction by adding an acid catalyst and heat. This is called dehydration reaction - so many names.

Finally, you need to understand how esters are formed. This is dead easy - if you know the rules. 

Alcohol + Carboxylic acid -----> Ester

During esterification (who comes up with these original names? Hah! The blog doesn't recognise half of the chemistry words) the Hydrogen on the alcohol and the OH group on the carboxylic acid break off and form water. This leaves the COO group for esters.

Esters have a fruity smell. - just a little side comment.

ALKENES

Now on to Alkanes brother ALKENE! No excitement? How about this:

Anybody hugging the screen yet? Geez, you people are strange.

As I'm sure you're all aware, because you're smart you all, is that alkenes contain a double bond. This means that they are unsaturated. A double bond means there are TWO pairs of electrons. 

First pair forms a sigma bond (that is an overlap of s orbitals - remember those bad boys?). each carbon donates one electron to the SIGMA bond.

Second pair of electrons forms a PI bond (this is sideways overlap of p orbitals on each carbon).

The PI bond fixes the carbon atoms in position, at either end of the double bond. This prevents any rotation of the bond.

Now when we think about the number of bonded regions, we have two on each carbon that is the bond C-H and One C-C bond (even though there's two bonds, it's one region of bonding). This means that the angle is 120 degrees, forming a trigonal planar shape.

Oh, and another thing. The PI bond is a very small region of space, this means there will be a high electron density (almost like pressure). This means that PI bond is weaker as electrophiles will attack that region first.

Next up we have reactions of alkenes. There's many so be prepared to take notes.

Addition of HYDROGEN / hydrogenation: 150 degrees Celsius

Addition of HALOGENS / halogenation: 

Addition of Hydrogen halides:

Addition of STEAM: H2PO4 catalyst, high temperature and pressure

Make sure that you know when what fission takes place. Keep in mind that when both elements have the same electronegativity, it will be HOMOLYTIC fission. Different elements always have a difference of electronegativity so it will be a HETEROLYTIC fission.

Electrophile: Electron pair acceptor (positive ion).

There are other things for which alkenes are used for. That is addition polymerisation. It is very simple: monomers of an alkene are added together. During polymerisation the double bond is broken to form a long chain made up of the subunits that were used in the reaction. Ensure you remember that the double bond breaks and draw brackets around that subunit to show it is only a section of the polymer.

Continuing on the topic of polymers is a section of waste. Remember that recycling is important to conserve raw resources, as well as reduce CO2 emission. This needs to be done carefully as some plastics contain Chlorine, which has an effect on ozone layer (about that later). Furthermore, chlorine is toxic and can form HCl, which we know is very corrosive. 

To reduce waste biodegradable and decomposable plastics are developed.

ALKANES

I know the exams are close and you're all sick of it all but we need to finish it. Here's something to cheer you up: 

Alkanes are hydrocarbons, which means that they are a compound made of hydrogen and carbon only. Common sense, eh??

You know from GCSE that crude oil is used as a source of hydrocarbons. During fractional distillation you are able to separate the fuels because each chain has a different boiling point.

Moving on to more difficult part is the effect of branching and chain length on boiling points.

CHAIN LENGTH: As the length increases, there are more electrons and hence more points of contact. This means more Van der Waals's forces form and more energy is required to break them. So as chain length increases the boiling point increases.

BRANCHING: Branched molecules have fewer points of contact (same number of electrons but different shape - isomerism, remember?). they are also further apart as molecules cannot get close together. This means less Van der Waal's forces and less energy required to break them. So as a molecule becomes more branched it's boiling point decreases.

When producing alkanes, you should keep in mind that they are fuels. Therefore they will be combusted. 

In real life complete combustion is not possible, especially since most cars have limited supply of oxygen. This results in incomplete combustion, which produces dangerous gases.

C4H10 (l) + 9/2 O2 (g) ----> 4 CO + 5 H2O

Carbon monoxide is poisonous, it prevents haemoglobin from binding with oxygen in red blood cells, so body tissues become deprived of oxygen. This can lead to death. Simply put.

It is much easier to combust branched alkanes than straight chains of alkanes. This is because oxygen only attacks the outside of the chain, so when the alkane is branched it is easier to attack the molecule and more of the molecule will be combusted.

Because of this companies crack the long chains of alkanes into smaller ones, produce branched alkanes or cyclic hydrocarbons, which combust more efficiently.

Time to step up the game now. It's radical substitution time! Yeah! Am I the only one happy? Nevermind. 

• We use ultraviolet light in a process called INITIATION to break Cl2, Br2, etc. into two radicals by homolytic fission (rings a bell?).

Important to note that radicals are very aggressive (reactive in chemistrian speak), so they will attack anything to form a bond.

Usually in a reaction bonds are first broken, using energy and then reformed releasing energy. In this case the radicals do not require any energy as their bonds are already broken.

• During PROPAGATION the radical attacks an alkane, steals a hydrogen and donates it's extra electron to the left over molecule, to form a HCl (or whatever halogen is there) and that radical alkane, minus one hydrogen.
• That radical alkane now wishes to attack everything, when it chances upon a Cl2 molecule it steals one Cl atoms and donates it's extra electron to the other Cl atom forming a Cl radical. We have formed the radical we started off with a ChloroAlkane (or whatever halogen used).
• This will go on to form many compounds, until no reactants are left. The problem is compounds may form that you do not wish to form.
• Final step is TERMINATION. During this process two radicals react together and share their extra electrons to form a covalent bond (shared pair of electrons, NO??) 

Cl. + Cl. ----> Cl2

CH3. +CH3. ----> C2H6

CH3. + Cl. ----> CH3Cl

Here are termination products when dealing with methane.

Ensure you understand that radical substitution forms a mixture of products.

BASIC CONCEPTS

I know, I know. The title is very vague. But to be honest this section is very vague, it has all the gibberish required for F322 but, seriously it's, well, basic. To start of with: 

Because why not??

First of all we have some definitions, which you may not necessarily need to remember but you do need to understand them.

EMPIRICAL FORMULA: The simplest whole number ratio of atoms of each element present in each element.

As it says in the name, remember it's a ratio. So you need to find the moles of each stuff and divide by the smallest to get the simplest WHOLE number ratio. Simple!

MOLECULAR FORMULA: Actual number of atoms of each element in a molecule.

GENERAL FORMULA: Simplest algebraic formula of a member of a homologous series.

That sounds depressing, doesn't it?? It only means a general formula for any compound of a specific type of molecule. E.g. An alkane, alkene, alcohol. Each of these is a homologous series, meaning each member of that group has similar properties (alkenes have double bonds, alcohols have OH group).

STRUCTURAL FORMULA: The arrangement of atoms in a molecule.

Write it out as it looks. E.g. Propane: CH3CH2CH3

DISPLAYED FORMULA: Relative positioning of atoms and bonds between them.

This one is how it actually looks. So its a drawing - you display drawings don't you?? Structural you write what it looks like, displayed you draw what it looks like.

SKELETAL FORMULA: Simplified organic formula, without showing hydrogen. Leaving just carbon skeleton and associated functional groups.

HOMOLOGOUS SERIES: A series of organic compounds having the same functional group, but with each successive member differing by CH2.

Alright, all this means is that they have the same general formula, same functional group but as you go along the group of that series each member (of that group) has an extra CH2.

FUNCTIONAL GROUP: A group of atoms responsible for the characteristic reactions of a compound.

Easy enough??

Here are the first ten member of the alkanes homologous series:

1. Methane
2. Ethane
3. Propane
4. Butane
5. Pentane
6. Hexane
7. Heptane
8. Octane
9. Nonane
10. Decane

Notice all end with the suffix "ane".
Akenes end with the suffix "ene".

The next part is a pain. A pain in the backside, at least for me it is. You need to be able to identify, draw and name different types of isomerisms.

Okay. So there are two main types of isomerisms:

• Structural isomerism: Same molecular formula but different structural formulae.
• Stereoisomerism: Same structural formulae but with different arrangement in space.

Yeah! Well done me, I freaked you out.

Structural isomerism (as in the name) has different structural formula. Remember structural formula is the one where you write what you see. So this means that the number of C, O, H, etc. is the same but the OH group (for example) could be attached to a different carbon. This is the easy one, just think about where all the functional groups are and voila!

Stereoisomerism is the tricky one. Here we have the structural formula the same. So none of the groups are moved about. This means that there MUST be a double bond present, because this prevents the molecule from twisting in the air (and being a nuisance) and therefore a functional group could be on the same carbon (structure the same) but on the top OR bottom of the double bond! Got it??

Then we have an example of stereoisomerism, which is E/Z isomerism (that bad guy - yes).

As I said a double bond has restricted rotation. Now for E/Z isomerism two different groups have to be attached to each carbon atom on the double bond.

• E isomerism is when the two (same) groups are on opposite sides.
• Z isomerism is when the two (same) groups are on the same side.

To make matters more complicated we have a special case for E/Z isomerism, which is cis-trans isomerism. In cis/trans isomerism there must be two non-hydrogen groups and two hydrogen around the double bond. So kind of like E/Z but this time the two different groups could be a Cl and H on one carbon and Cl and H on another.

• Cis isomerism is when the two other (non hydrogen) groups are on the same side of the double bond.
• Trans isomerism is when the two other (non hydrogen) groups are on the different sides of the double bond. Think of trans as being changing gender... so different.

Next up we have different types of covalent bond fission ( breaking of covalent bonds in different way).

During HOMOLYTIC fission the two products have the same charge, they are two radicals as the shared pair (the covalent bond) is split equally so each atom gets one electron each.

During HETEROLYTIC fission the two products have different charges. A cation (positive ion) and anion (negative ion) are formed. This is because one of the two atoms is greedy (more electronegative) and steals both electrons from the covalent bond. This type of fission happens during electrophilic addition (about that later).

We can use CURLY arrows to show where the pair of electrons (from covalent bond) have moved from and to.

When drawing these diagrams, think....

1. Which is more electronegative?
2. Then you'll know if it's hetero/homolytic fission.
3. Next you know that cations are electrophiles (attracted to electrons/ or electron acceptors) so they will attack any double bonds of a compound and steal on electron to become neutral and another one to form a covalent bond.
4. This leaves that carbon a cation (its extra electron was stolen - the one from double bond).
5. This means the left over anion will then attack the cation carbon and donate it's electrons, one to make carbon neutral so the anion becomes neutral too, and one to form a covalent bond.

Ta dah!

Worst for last. You really should keep in mind percentage yield and atom economy, and for crying out loud! Do understand them too, memorising isn't everything.

In real life situations, it is not always possible to get the calculated amount of product in a reaction:

• Reversible reactions may not go to completion
• Some product may be lost when it is removed from the reaction mixture
• Some of the reactants may react in an unexpected way

The yield of a reaction is the actual mass of product obtained. The percentage yield can be calculated:

Percentage Yield =    Actual mass of product obtained
                   -------------------------------------   * 100
                    Maximum theoretical mass of product

So in short percentage yield shows you how efficient that reaction is, how close you got to what you wanted to get.

Atom economy is very similar, and that what gives me a headache. It looks at how much of all the products is what you really want. Ugh! I'll explain in a bit.

Atom economy = Molecular mass of the desired product
               -------------------------------------  *100
                  Total molecular mass of products

Now... You could have a high percentage yield (as you get nearly the same mass of - I don't know! butanol, for example - as you expected to get) but the atom economy is low as there are many waste products produced along with butanol (e.g. carbon dioxide and water).

To clear up. % yield looks at the what you wanted to get and actually do. And atom economy creates a ratio of what you actually got over all products - so how efficient the reaction is. You do not want a low atom economy as it wastes raw resources.

Phew! That's it. Hope you got that, it took me a LOOOOONG white to get my head round it.

F322

YEAH! I have managed to start with the right thing this time!. Anyway, Tuesday is the exam and I got a bit carried away by my Very important things that I have to do. So lets get cracking with F322. Here is the list of topics I will cover:

• Basic concepts
• Alkanes
• Alkenes
• Alcohols
• Halogenoalkanes
• Modern analytical techniques
• Enthalpy changes
• Rates and equilibrium
• Chemistry of the air
• Green chemistry

Don't get put off by the titles of these sections, they just make my life much easier, it enables me to put all the work into chunks, not tiny chapters. Well, what are we waiting for?? 

Lets GO!!

F321

I think it is a bit too late for that but I thought I'd mention the different topics that need to be known for the F321 module. (I know, sorry, should have done it ages ago but I sort of forgot). As a little remedy here is the list of topics for F321 and a very snazzy picture.

• Atoms
• Moles and equations
• Acids
• Redox
• Ionisation energy
• Bonding and structure
• Periodicity
• Group 2
• Group 7

Summary

Hello everyone.

I hope you have all had a wonderful exam, though the questions were sooo dull. Now I wanted just to say that since the F321 exam is over and done with I will now comcentrate on F322.

GROUP 7

To finish off with F321, comes the lovely group 7. After doing F321 for a year I have realised how easy it becomes (you know - with further knowledge of the F322 unit it really does help).

Keep in mind that F321 is your chance to pull your grades to the highest notch to ready yourself for year 13. So remember... Memorise those things I have mentioned. They really do help.

First of we start with a pretty picture to easy you into the toxic topic that group 7 really is. Uhhh, Quite literally. For they are toxic so remember about the fumes cupboard.

Yeah, that's what I thought. To your notes, now.

There many things they could ask you about group 7, all very straight forward if you know how to play the game.

You need to be able to explain the trend in boiling points of group 7. You know the drill, don't you? What are the intermolecular forces here??? Well there's no hydrogen bonding or any permanent dipole-dipole interactions. So that's out. What's left? That's right Van der waals's forces. That old chestnut.

As you go down group 7 there are more electrons, which means that the Van der Waals's forces will be much stronger. When melting or boiling something you break the intermolecular forces first and then you break any intramolecular forces if you have to (like covalent, ionic bonding - these have much higher boiling points as they are usually a giant lattice). Stronger forces mean more energy so higher boiling point as you go down the group.

This can be seen in the physical states of the elements. Flourine and chlorine are gases. Bromine is a liquid, whilst iodine is a solid (all at room temperature).

Up next: Identifying group 7. If you wish to test a substance to see if it contained any halide ions (a halogen ion bonded to a positive ion) you should add some SILVER NITRATE AgNO3 to it. 

If Chloride present: White precipitate forms.
Ag+(aq) +Cl-(aq) ---> AgCl(s)

If Bromide present: Cream precipitate forms.
Ag+(aq) +Br-(aq) ---> AgBr(s)

If iodine present: Yellow precipitate forms.
Ag+(aq) +I-(aq) ---> AgI(s)

Now, what if you have two of these substances in at once? Well obviously one of them is more reactive so will displace the less reactive one and you'll only see one precipitate form. Now different halide precipitates have different solubility's in aqueous ammonia - this will enable you to confirm the halide present.

Chloride: Soluble in dilute ammonia.
Bromide: Soluble in concentrated ammonia.
Iodide: Insoluble in concentrated ammonia.

Another trend that you may be asked to look at is the reactivity of group 7. It is pretty much the opposite of group 1 elements.

As you go down group 7 the reactivity decreases, this is because:

• The atomic radius increases down the group due to the extra shell that is added each time. . This means the nuclear attraction on the outer electrons is less so it is harder to attract an electron (to reach a full outer shell).
• The electron shielding increases as there is an extra shell of electrons each time you go down. This means the nuclear attraction on the outer electrons is weaker. So harder to attract electrons.
• The nuclear charge increases as there are more protons as you go down the group, which mean the nuclear attraction on outer electrons is stronger. However the other two factors outweighs this and overall the attraction on outer electrons is weaker so more difficult to attract electrons so reactivity goes down.

We can show the decrease in reactivity by a REDOX reaction. How you ask?? Well, A halide is already present in our solution and a halogen is introduced. You will get a displacement reaction - the more reactive element taking over and becoming the halide. Halogens form solutions of different colours, so any change in colour will indicate a displacement reaction. An organic solvent is added, usually CYCLOHEXANE. This is very useful, plus it gives very pretty colours.

Cl2 : Pale green in water and in cyclohexane.
Br2: Orange in water and in cyclohexane.
I2: Brown in water and purple in cyclohexane.

This leads on nicely to the DISPROPORTIONATION reaction. It is very simply put: a reaction is which the same element is oxidised and reduced.

We have two reactions that need to be understood and preferable memorised.

First off, is Chlorine in water:

Chlorine added to water will kill bacteria and make the water safe to drink. Chlorine will react with water to produce two acidic products: Hydrochloric acid and Chloric acid.

Cl2 (aq) + H2O (l) ---> HCl (aq) + HClO (aq)
0                        -1                 Chlorine reduced.
0                                   +1      Chlorine oxidised.

Second, is Chlorine in aqeuous Sodium hydroxide.

Remember that Chlorine is only slightly coluble in water and has a mild bleaching action. The house stuff that we use is formed when dilute NaOH and CL2 react together.This is another disproportination reaction.

Cl2 (aq) + 2NaOH (aq) ---> NaCl (aq) + NaClO (aq) + H2O (l)
0                            -1               Chlorine is reduced.
0                                        +1   Chlorine is oxidised.

You may also want to keep in mind that it is the ClO- ion which is responsible for all that bleaching magic.

To finish off with, you are all fully aware that Chlorine is used in water treatment. Now there are good point and bad point to that.

It kills bacteria and enables water to be drinkable. However, chlorine gas is toxic and there are also some risks of chlorinated carbons to form, which are responsible for the destruction of the ozone layer. Which is bad - trust me

GROUP 2

Here we go for the hundredth time. Group 2 elements, you must be sick to death of them by now. I certainly am. But you wan that A or even A* next year lets get to work!

First of all we need to have a look at the trend of group 2. You may have noticed but as you go down group 2 the reactivity increases. This may seem familiar to you by now but its all about three factors:

• Electron shielding: As you go down group 2 there is a new shell added each time. Now group 2 metals want to LOSE the 2 electrons as it is the easiest way to get a full outer shell. As a new shell is added there is more electron shielding so the attraction on the outer electrons is weaker and the 2 outer electrons can be lost much easier.
• Atomic radius: Each element down the group has an extra shell which means the 2 outer electrons are further away. This means that there is a weaker attraction on the outer electrons and they can be lost much easier.
• Nuclear charge: As you go down group 2 the nuclear charge increases so the attraction on the outer electrons increases. HOWEVER, (before you ask) the other two factors out weight this increased attraction.

Overall it is much easier to lose the 2 electrons as you go down the group so the reactivity increases.

There are several reaction that you need to understand, they are fairly basic so don't worry too much.

First, we have a group 2 metal and oxygen:

This is a REDOX reaction (more about these on another post). When a group two element is reacted with oxygen a metal oxide forms.

2Ca (s) + O2 (g) ---> 2CaO (s)

The calcium here is oxidised: goes from 0 to +2

The oxygen is reduced: goes from 0 to -2

Second, we have a group 2 metal and water:

In this reaction when the metal reacts with water a metal hydroxide forms and hydrogen gas.

Ca (s) + 2H2O (l) ---> Ca(OH)2 (aq) + H2 (g)

You would be able to see in the reaction that the metal dissolves (the calcium goes from solid to aqueous. Also gas is produced so there would be fizzing. Furthermore Ca(OH)2 is an alkali (a soluble base) therefore if you added an indicator it would show that you have an alkali present (the pH would be around 10-12 depending on the strength of the alkali).

In this case:

Calcium is oxidised: Goes from 0 to +2

Hydrogen is reduced: Goes from +1 to 0

As you go further down the group each metal will react more violently with water.

Another fact to mention (which you need to know) is that metal hydroxides are used to neutralise acidic soil, however too much will turn the soil too alkali (which is bad).

Third, we have a group 2 oxide and water:

Here a metal oxide is reacted with water to from a solution of metal hydroxide.

CaO (s) + H2O (l) ---> Ca(OH)2 (aq)

Group 2 hydroxides dissolve in water to form alkaline solutions. Now, as you go down the group the solubility increases and the resulting solution in also more alkaline.

There is also a more ugly side to Group 2 that you need to remember: Thermal decomposition.

THERMAL DECOMPOSITION: The breaking up of a chemical substance with heat into at least two chemical substances.

Group 2 carbonates can be decomposed (broken down) by heat to form a metal oxide and carbon dioxide gas.

CaCO3 (s) ---> CaO (s) + CO2 (g)

The carbonates become more difficult to decompose as you go DOWN the group.

Here is a summary of all the reactions I have mentioned above, as well as reactions with acids. Memorise these as some will sneakily pop up in the exam and ruin your awesome chemistry exam day.

PERIODICITY

Welcome back (or hello if you're new) and be warned this section is very simple but very WORDY! If you get one of these in the exam, which you're bound to get, you have gained some free marks (but only if you memorise the rules). Have fun.

First of all: 



Now we can begin.

As you obviously have noticed, cos you're awesome all of you, the periodic table is arranged in a specific order. Well, this is no rocket science but you do need to understand that all the elements are not arranged in an increasing mass order but in an increasing atomic number order. This means that each element along has one more proton and one more electron.

Also as you move along the periods the chemicals there have a repeating chemical and physical property trend. Such as lithium and sodium react similarly when added to water, they also look similar. This is because of the number of electrons in the outer shell, which is very important to note.

Now this repeating pattern is called PERIODICITY.

Like I have mentioned before all the properties of an element depend on its electron structure, especially the arrangement of the outer electrons.

During one of my previous posts I explained how the ionisation energy is affected by these three factors:

• Nuclear charge
• Atomic radii
• Electron shielding

Make sure you have these memorised, but also that you UNDERSTAND them. Once you do, you will fly through Chemistry (at least this unit). Just to ensure you remember I'll go through it again.

NUCLEAR CHARGE: As you go along the period each element along has an extra proton, which increases the attraction on the outer electrons. An extra electron is also added, however it is added to the same shell as the previous one (until you get to the next period then its the next shell) therefore there is no extra shielding. Because of this the ionisation energy INCREASES ACROSS a period.

On top of this as the nuclear charge increases across a period this draws the electrons further in (as there isn't any extra electron shielding - as electrons are added to the same shell). This therefore increases the attraction on the electrons and ionisation increases ACROSS the period.

ELECTRON SHIELDING: A you go DOWN a group there is a new shell added each time. This means there is more repulsion between the electrons and therefore there is a weaker attraction of the nucleus on the outer electron (the electrons act like a shield - shielding the outer electron from the powerful attraction of the nucleus: hence the name being electron shielding). Therefore the ionisation energy DECREASES DOWN a group.

ATOMIC RADII: This is the easiest part. It is linked to the nuclear charge (read above) and extra shells. As you go DOWN a group an extra shell is added so the radius increases. This means that the outer electron experiences a weaker attraction from the nucleus and hence the ionisation energy DECREASES DOWN the group.

The next trend which you need to understand is boiling points. Like I have discussed in previous posts, boiling points are all about how easy it is to break bonds - to make the molecule free to fly off wherever it wants. Also remember when you bring something to its boiling point you do not break the INTRAmolecular forces (the ones the bond the ATOMS together - like covalent bonds as this requires HUGE amounts of energy) but the INTERmolecular forces (the attraction between neighbouring molecules - much lower energy required). Make sure you spot the difference!

I would love to get you a lovely, snazzy table but the blogger doesn't want to cooperate so lets improvise.

Ok. As you go across period 2 we have: 

Li, Be:

• These are giant metallic lattices, which have strong metallic bonding between the POSITIVE IONS AND NEGATIVE DELOCALISED ELECTRONS.
• They have very high melting points, which increase as you go across a period. This is because the nuclear charge increases, so delocalised electrons are more dense (more of them - if you have a +1 metal ion then you need 1 electron to have no charge overall, if you have +2 charge on the ion you need an extra free electron to cancel out the charge). The higher nuclear attraction also draws the electrons closer in and therefore the metallic bonding is stronger.
• Stronger bonding - harder to break the bonds - more energy required - higher boiling/melting point

B, C:

• These are giant covalent lattices, which contain strong covalent bonds BETWEEN ATOMS so a huge amount of energy is required to break them, so high melting/boiling point.
• You don't see molten diamonds often do you? They're made of carbon alone.

N2, O2, F2, Ne:

• These are simple molecular structures (not giant lattices). They only have weak forces between MOLECULES, which are Van der Waals's forces. Weak forces mean low energy required to break them so a low boiling/melting point.

Period 3 is very similar:

Na, Mg, Al: Giant metallic lattice.

Si: Giant covalent lattice.

P4, S8, Cl2, Ar: Simple molecular structure.

BONDING AND STRUCTURE

After that last post I feel like I need to do something easier or you'll all freak out on me. This one is less understand and more just memorising. It's just going through the notions of reading question and seeing how many electron pairs its talking about and also what kind, LONE or BONDED? I'll go through that in a minute. 

In a way you could compare this chapter to Assassin's Creed, you read the question and as soon as the answer flies past you, you attack it and stab it to the page so it stays there, or maybe that's just me.

Coming up: Some definitions.

IONIC BONDING: The electrostatic attraction between oppositely charged ions. 

You may need to construct a "dot and cross" diagram to show this. Simply figure out which atom will lose an electron and which one will gain that electron (or more). The one which lost the electron has a charge of +1 the one which gained the electron will have a charge of -1. Now you all know that opposite charges attract so these ions also attract to form a compound.

Now draw the positive ion with an empty shell (or a lower shell that is full), brackets round it and a charge of +1 and next to it draw a diagram of the -1 ion with its own electrons as crosses and the extra electrons as circles (or visa versa). Don't forget the brackets!

It's quite wordy to explain it but it is very simple in practise, quite literally free marks in the exam.

Example: A reaction between Magnesium and chlorine. The magnesium loses two electrons to have a full outer shell and chlorine needs 1 electron each to have a full outer shell (so two chlorine atoms required). The magnesium donates the 1 electron to each chlorine and they become ions. Chlorine with -1 charge each and Magnesium with +2 charge. 



You are required to know the formulae of the following ions:

NO3: -1
CO3: -2
SO4: -2
NH4: +1

COVALENT BOND: Shared pair of electrons.

Unlike in ionic bonding, during covalent bonding the electrons are shared, not given away. This forms a bond equal in strength to ionic bonding. However covalent bonding occurs between non metals, whilst ionic bonding occurs between metals and non metal (except hydrogen and a non metal).

It is also important to note that a covalent bond is directional (acting only between the two atoms involved in the bond) whilst ionic bonding attracts in all directions (often forming a lattice - Like a net).



Here is a good guide on how the bonds are formed. The individual atoms have only one role is life to form a full outer shell, so that they may be stable and lead a happy atomic life.

Sometimes a DATIVE COVALENT bond forms. This happens when a molecule, once bonded, a lone pair of electrons and it chances to meet a +1 ion so the two react together to form positive ion with no more electrons to pair up.

LONE PAIR: Outer shell pair of electrons which are not involved in the chemical bonding.

DATIVE COVALENT BOND: A shared pair of electrons which has been donated by ONE of the bonding atoms.

Now we get onto the tough memorisation part: SHAPES! <---- Ugh the horror. Make sure you remember the examples AND the bonding angles AND the names.

The shape of the molecule is determined by repulsion between electron pairs surrounding the central atom.Lone pairs repel more than bonded pairs.

VERY IMPORTANT.

The type of bonding very much depends on something called the ELECTRONEGATIVITY.
 ELECTRONEGATIVITY: It is the attraction of an atom on the bonded pair of electrons.

In short? The more electronegative an atom is the more it wants those electrons. So when you get two atoms which have a high difference in electronegativity, this results in a permanent dipole to form (one of the atoms becomes slightly positive and the other slightly negative). The resulting dipole results in a polar bond to form. 

When the difference in electronegativity is very high then ionic bonding occurs where the atoms do not become slightly charged but have a full charge on them.

When there is no difference is electronegativity the reacting atoms form a covalent bond and share the electrons equally, they're not greedy.

However sometimes in a diatomic molecule (E.g. F-F or Cl-Cl) the electrons may be distributed unevenly. This results in an instantaneous dipole. This dipole will induce a dipole in the neighbouring molecule. The attraction between the instantaneous dipoles is Van der Waals's forces. The greater the number of electrons the greater the Van der Waal's forces.

This attraction is temporary and the electrons will move to another random place and create another instantaneous dipole in another direction.

There is another type of bonding that you need to know. It's called hydrogen bonding and yes it involves hydrogen.

HYDROGEN BONDING: A strong dipole-dipole attraction between an electron deficient atom and a lone pair on another molecule.

This type of bonding is strong (weaker than covalent but stronger than Van der Waals) and can only occur between molecules containing N-H and O-H.

This type of bonding give water very special properties:

• Ice is less dense tan water. Usually solid is more dense than liquid. However, in water (when freezing) more hydrogen form which hold then molecules apart. When ice melts again, the bonds break allowing the molecules to come closer together.
• Water has a relatively high melting/boiling point. This is because on top of Van der Waals there are also hydrogen bonds that need to be broken. Extra bonds - Extra energy.
• High surface tension and viscosity (thickness) of water. This all due to hydrogen bonds.

Final type of bonding is Metallic bonding.

METALLIC BONDING: The attraction of positive ions to the delocalised electrons.

Nothing to add there, very simple. All in the name: you have a metal that becomes +1 ion there will be one delocalised electron per each positive ion.

All the bonding is very important so make sure you memorise it as it needs to be applied to some higher end questions. Here is a quick summary:

• Ionic bonding
• Covalent bonding
• Dative covalent
• Electronegativity
• Permanent dipoles - Polar bonds
• Instantaneous dipoles - Van der Waals
• Hydrogen bonding
• Metallic bonding

There are many different structures (Eiffel tower and the pyramids being one of them):

Giant ionic lattices: These include ionic bonding which is very strong - E.g. NaCl (the common salt).

• High melting and boiling point - Strong forces means lots of energy.
• When solid the ions are fixed but when molten or in a solution the ions are free to move.
• Ionic lattices dissolve polar solvents, such as water. The lattice is broken down by the water molecules which surround each ion.

Giant covalent lattices: They have very strong covalent bonds - E.g. Carbon structures such as diamond and graphite.

• High melting and boiling points - Strong covalent bonding.
• Non conductor of electricity as there are no free charged particles that can move.
• They are insoluble in both polar and non polar solvents, as the covalent bonds are too strong to be broken by solvents.

Giant metallic lattices: Quite obvious, they have metallic bonding which is very strong.

• High melting/boiling points - Strong metallic forces.
• Good conductors as delocalised electrons can move.

Simple molecular lattices: Include hydrogen bonds which are moderately strong - E.g. Water.

• Relatively high boiling point - Moderately strong hydrogen bonds.

Simple molecules: Only Van der Waal's forces attract the molecules together, fairly weak - E.g. I2.

• Low melting and boiling points - Weak Van der Waals's forces.

DIAMOND:

These shiny crystals are nothing more than carbon - literally. The carbon atoms are joined together by strong covalent bonds, which means a very high boiling/melting point. 

They are rubbish at conducting electricity as there are no delocalised electrons and all outer electrons are used up in bonding. (I mean if you're running and juggling you can't be playing the piano at the same time, can you??)

On top of this they are very hard as the tetrahedral shape allows the applied force to be spread throughout the whole structure. (Another great example: when you're at a concert and you throw your self into the crowd - their hands will keep you from falling and they won't really feel your weight as it is spread out onto many people, easy eh??)

GRAPHITE:

Believe me or not this is diamond's sister (though not as shiny). It is also made of carbon alone. However graphite is made of strong covalently bonded layers, which is why you can draw with a pencil (the layers slide off).

It is a good conductor as there are delocalised electrons between the layers, which can move.

Unlike diamond, graphite is soft. This is because there are only weak Van der Waals's forces between the layers which allow them to slide off easily, even though the bonding within each layer is strong.