STRUCTURAL ISOMERISM

This page explains what structural isomerism is, and looks at some of the various ways that structural isomers can arise.

What is structural isomerism?

What are isomers?

Isomers are molecules that have the same molecular formula, but have a different arrangement of the atoms in space. That excludes any different arrangements which are simply due to the molecule rotating as a whole, or rotating about particular bonds.

For example, both of the following are the same molecule. They are not isomers. Both are butane.

There are also endless other possible ways that this molecule could twist itself. There is completely free rotation around all the carbon-carbon single bonds.


Note: Isomerism is much easier to understand if you have actually got some models to play with. If your school or college hasn't given you the opportunity to play around with molecular models in the early stages of your organic chemistry course, you might consider getting hold of a cheap set. The models made by molymod are both cheap and easy to use. An introductory organic set is more than adequate. Find them at www.molymod.com.

Alternatively , get hold of some coloured Plasticene and some used matches and make your own.

If you had a model of a molecule in front of you, you would have to take it to pieces and rebuild it if you wanted to make an isomer of that molecule. If you can make an apparently different molecule just by rotating single bonds, it's not different - it's still the same molecule.


Note: It's really important that you understand this. If you aren't sure, then you must get hold of (or make) some models.

What are structural isomers?

In structural isomerism, the atoms are arranged in a completely different order. This is easier to see with specific examples.

What follows looks at some of the ways that structural isomers can arise. The names of the various forms of structural isomerism probably don't matter all that much, but you must be aware of the different possibilities when you come to draw isomers.

Types of structural isomerism

Chain isomerism

These isomers arise because of the possibility of branching in carbon chains. For example, there are two isomers of butane, C4H10. In one of them, the carbon atoms lie in a "straight chain" whereas in the other the chain is branched.



Note: Although the chain is drawn as straight, in reality it's anything but straight. If you aren't happy about the ways of drawing organic molecules, follow this link.

Use the BACK button on your browser to return to this page.

Be careful not to draw "false" isomers which are just twisted versions of the original molecule. For example, this structure is just the straight chain version of butane rotated about the central carbon-carbon bond.

You could easily see this with a model. This is the example we've already used at the top of this page.

Pentane, C5H12, has three chain isomers. If you think you can find any others, they are simply twisted versions of the ones below. If in doubt make some models.

Position isomerism

In position isomerism, the basic carbon skeleton remains unchanged, but important groups are moved around on that skeleton.

For example, there are two structural isomers with the molecular formula C3H7Br. In one of them the bromine atom is on the end of the chain, whereas in the other it's attached in the middle.

If you made a model, there is no way that you could twist one molecule to turn it into the other one. You would have to break the bromine off the end and re-attach it in the middle. At the same time, you would have to move a hydrogen from the middle to the end.

Another similar example occurs in alcohols such as C4H9OH

These are the only two possibilities provided you keep to a four carbon chain, but there is no reason why you should do that. You can easily have a mixture of chain isomerism and position isomerism - you aren't restricted to one or the other.

So two other isomers of butanol are:



Note: It's essential if you are asked to draw isomers in an exam not to restrict yourself to chain isomers or position isomers. You must be aware of all the possibilities.

You can also get position isomers on benzene rings. Consider the molecular formula C7H8Cl. There are four different isomers you could make depending on the position of the chlorine atom. In one case it is attached to the side-group carbon atom, and then there are three other possible positions it could have around the ring - next to the CH3 group, next-but-one to the CH3 group, or opposite the CH3 group.

Functional group isomerism

In this variety of structural isomerism, the isomers contain different functional groups - that is, they belong to different families of compounds (different homologous series).

For example, a molecular formula C3H6O could be either propanal (an aldehyde) or propanone (a ketone).

There are other possibilities as well for this same molecular formula - for example, you could have a carbon-carbon double bond (an alkene) and an -OH group (an alcohol) in the same molecule.

Another common example is illustrated by the molecular formula C3H6O2. Amongst the several structural isomers of this are propanoic acid (a carboxylic acid) and methyl ethanoate (an ester).



Note: To repeat the warning given earlier: If you are asked to draw the structural isomers from a given molecular formula, don't forget to think about all the possibilities. Can you branch the carbon chain? Can you move a group around on that chain? Is it possible to make more than one type of compound?

Be careful though! If you are asked to draw the structures of esters with the molecular formula C3H6O2, you aren't going to get a lot of credit for drawing propanoic acid, even if it is a valid isomer.

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DRAWING ORGANIC MOLECULES

This page explains the various ways that organic molecules can be represented on paper or on screen - including molecular formulae, and various forms of structural formulae.

Molecular formulae

A molecular formula simply counts the numbers of each sort of atom present in the molecule, but tells you nothing about the way they are joined together.

For example, the molecular formula of butane is C4H10, and the molecular formula of ethanol is C2H6O.

Molecular formulae are very rarely used in organic chemistry, because they don't give any useful information about the bonding in the molecule. About the only place where you might come across them is in equations for the combustion of simple hydrocarbons, for example:

In cases like this, the bonding in the organic molecule isn't important.

Structural formulae

A structural formula shows how the various atoms are bonded. There are various ways of drawing this and you will need to be familiar with all of them.

Displayed formulae

A displayed formula shows all the bonds in the molecule as individual lines. You need to remember that each line represents a pair of shared electrons.

For example, this is a model of methane together with its displayed formula:

Notice that the way the methane is drawn bears no resemblance to the actual shape of the molecule. Methane isn't flat with 90° bond angles. This mismatch between what you draw and what the molecule actually looks like can lead to problems if you aren't careful.

For example, consider the simple molecule with the molecular formula CH2Cl2. You might think that there were two different ways of arranging these atoms if you drew a displayed formula.

The chlorines could be opposite each other or at right angles to each other. But these two structures are actually exactly the same. Look at how they appear as models.

One structure is in reality a simple rotation of the other one.


Note: This is all much easier to understand if you have actually got some models to play with. If your school or college hasn't given you the opportunity to play around with molecular models in the early stages of your organic chemistry course, you might consider getting hold of a cheap set. The models made by molymod are both cheap and easy to use. An introductory organic set is more than adequate. Find them at www.molymod.com.

Alternatively , get hold of some coloured Plasticene and some used matches and make your own. It's cheaper, but distinctly messier!

Consider a slightly more complicated molecule, C2H5Cl. The displayed formula could be written as either of these:

But, again these are exactly the same. Look at the models.

The commonest way to draw structural formulae

For anything other than the most simple molecules, drawing a fully displayed formula is a bit of a bother - especially all the carbon-hydrogen bonds. You can simplify the formula by writing, for example, CH3 or CH2 instead of showing all these bonds.

So for example, ethanoic acid would be shown in a fully displayed form and a simplified form as:

You could even condense it further to CH3COOH, and would probably do this if you had to write a simple chemical equation involving ethanoic acid. You do, however, lose something by condensing the acid group in this way, because you can't immediately see how the bonding works.

You still have to be careful in drawing structures in this way. Remember from above that these two structures both represent the same molecule:

The next three structures all represent butane.

All of these are just versions of four carbon atoms joined up in a line. The only difference is that there has been some rotation about some of the carbon-carbon bonds. You can see this in a couple of models.

Not one of the structural formulae accurately represents the shape of butane. The convention is that we draw it with all the carbon atoms in a straight line - as in the first of the structures above.

This is even more important when you start to have branched chains of carbon atoms. The following structures again all represent the same molecule - 2-methylbutane.

The two structures on the left are fairly obviously the same - all we've done is flip the molecule over. The other one isn't so obvious until you look at the structure in detail. There are four carbons joined up in a row, with a CH3 group attached to the next-to-end one. That's exactly the same as the other two structures. If you had a model, the only difference between these three diagrams is that you have rotated some of the bonds and turned the model around a bit.

To overcome this possible confusion, the convention is that you always look for the longest possible chain of carbon atoms, and then draw it horizontally. Anything else is simply hung off that chain.

It doesn't matter in the least whether you draw any side groups pointing up or down. All of the following represent exactly the same molecule.

If you made a model of one of them, you could turn it into any other one simply by rotating one or more of the carbon-carbon bonds.

How to draw structural formulae in 3-dimensions

There are occasions when it is important to be able to show the precise 3-D arrangement in parts of some molecules. To do this, the bonds are shown using conventional symbols:

For example, you might want to show the 3-D arrangement of the groups around the carbon which has the -OH group in butan-2-ol.

Butan-2-ol has the structural formula:

Using conventional bond notation, you could draw it as, for example:

The only difference between these is a slight rotation of the bond between the centre two carbon atoms. This is shown in the two models below. Look carefully at them - particularly at what has happened to the lone hydrogen atom. In the left-hand model, it is tucked behind the carbon atom. In the right-hand model, it is in the same plane. The change is very slight.

It doesn't matter in the least which of the two arrangements you draw. You could easily invent other ones as well. Choose one of them and get into the habit of drawing 3-dimensional structures that way. My own habit (used elsewhere on this site) is to draw two bonds going back into the paper and one coming out - as in the left-hand diagram above.

Notice that no attempt was made to show the whole molecule in 3-dimensions in the structural formula diagrams. The CH2CH3 group was left in a simple form. Keep diagrams simple - trying to show too much detail makes the whole thing amazingly difficult to understand!

Skeletal formulae

In a skeletal formula, all the hydrogen atoms are removed from carbon chains, leaving just a carbon skeleton with functional groups attached to it.

For example, we've just been talking about butan-2-ol. The normal structural formula and the skeletal formula look like this:

In a skeletal diagram of this sort

*

there is a carbon atom at each junction between bonds in a chain and at the end of each bond (unless there is something else there already - like the -OH group in the example);
*

there are enough hydrogen atoms attached to each carbon to make the total number of bonds on that carbon up to 4.

Beware! Diagrams of this sort take practice to interpret correctly - and may well not be acceptable to your examiners (see below).

There are, however, some very common cases where they are frequently used. These cases involve rings of carbon atoms which are surprisingly awkward to draw tidily in a normal structural formula.

Cyclohexane, C6H12, is a ring of carbon atoms each with two hydrogens attached. This is what it looks like in both a structural formula and a skeletal formula.

And this is cyclohexene, which is similar but contains a double bond:

But the commonest of all is the benzene ring, C6H6, which has a special symbol of its own.



Note: Explaining exactly what this structure means needs more space than is available here. It is explained in full in two pages on the structure of benzene elsewhere in this site. It would probably be better not to follow this link unless you are actively interested in benzene chemistry at the moment - it will lead you off into quite deep water!

Deciding which sort of formula to use

There's no easy, all-embracing answer to this problem. It depends more than anything else on experience - a feeling that a particular way of writing a formula is best for the situation you are dealing with.

Don't worry about this - as you do more and more organic chemistry, you will probably find it will come naturally. You'll get so used to writing formulae in reaction mechanisms, or for the structures for isomers, or in simple chemical equations, that you won't even think about it.

There are, however, a few guidelines that you should follow.

What does your syllabus say?

Different examiners will have different preferences. Check first with your syllabus. If you've down-loaded a copy of your syllabus from your examiners' web site, it is easy to check what they say they want. Use the "find" function on your Adobe Acrobat Reader to search the organic section(s) of the syllabus for the word "formula".

You should also check recent exam papers and (particulary) mark schemes to find out what sort of formula the examiners really prefer in given situations. You could also look at any support material published by your examiners.


Note: If you are working to a UK-based syllabus and haven't got a copy of that syllabus and recent exam papers, follow this link to find out how to get them.

What if you still aren't sure?

Draw the most detailed formula that you can fit into the space available. If in doubt, draw a fully displayed formula. You would never lose marks for giving too much detail.

Apart from the most trivial cases (for example, burning hydrocarbons), never use a molecular formula. Always show the detail around the important part(s) of a molecule. For example, the important part of an ethene molecule is the carbon-carbon double bond - so write (at the very least) CH2=CH2 and not C2H4.

Where a particular way of drawing a structure is important, this will always be pointed out where it arises elsewhere on this site.

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ELECTRONIC STRUCTURE AND ATOMIC ORBITALS

A simple view

In any introductory chemistry course you will have come across the electronic structures of hydrogen and carbon drawn as:



Note: There are many places where you could still make use of this model of the atom at A' level. It is, however, a simplification and can be misleading. It gives the impression that the electrons are circling the nucleus in orbits like planets around the sun. As you will see in a moment, it is impossible to know exactly how they are actually moving.

The circles show energy levels - representing increasing distances from the nucleus. You could straighten the circles out and draw the electronic structure as a simple energy diagram.

Atomic orbitals

Orbits and orbitals sound similar, but they have quite different meanings. It is essential that you understand the difference between them.

The impossibility of drawing orbits for electrons

To plot a path for something you need to know exactly where the object is and be able to work out exactly where it's going to be an instant later. You can't do this for electrons.


Note: In order to plot a plane's course, it is no use knowing its exact location in mid-Atlantic if you don't know its direction or speed. Equally it's no use knowing that it is travelling at 500 mph due west if you have no idea whether it is near Iceland or the Azores at that particular moment.

The Heisenberg Uncertainty Principle (not required at A'level) says - loosely - that you can't know with certainty both where an electron is and where it's going next. That makes it impossible to plot an orbit for an electron around a nucleus. Is this a big problem? No. If something is impossible, you have to accept it and find a way around it.

Hydrogen's electron - the 1s orbital


Note: In this diagram (and the orbital diagrams that follow), the nucleus is shown very much larger than it really is. This is just for clarity.

Suppose you had a single hydrogen atom and at a particular instant plotted the position of the one electron. Soon afterwards, you do the same thing, and find that it is in a new position. You have no idea how it got from the first place to the second.

You keep on doing this over and over again, and gradually build up a sort of 3D map of the places that the electron is likely to be found.

In the hydrogen case, the electron can be found anywhere within a spherical space surrounding the nucleus. The diagram shows a cross-section through this spherical space.

95% of the time (or any other percentage you choose), the electron will be found within a fairly easily defined region of space quite close to the nucleus. Such a region of space is called an orbital. You can think of an orbital as being the region of space in which the electron lives.


Note: If you wanted to be absolutely 100% sure of where the electron is, you would have to draw an orbital the size of the Universe!

What is the electron doing in the orbital? We don't know, we can't know, and so we just ignore the problem! All you can say is that if an electron is in a particular orbital it will have a particular definable energy.

Each orbital has a name.

The orbital occupied by the hydrogen electron is called a 1s orbital. The "1" represents the fact that the orbital is in the energy level closest to the nucleus. The "s" tells you about the shape of the orbital. s orbitals are spherically symmetric around the nucleus - in each case, like a hollow ball made of rather chunky material with the nucleus at its centre.

The orbital on the left is a 2s orbital. This is similar to a 1s orbital except that the region where there is the greatest chance of finding the electron is further from the nucleus - this is an orbital at the second energy level.

If you look carefully, you will notice that there is another region of slightly higher electron density (where the dots are thicker) nearer the nucleus. ("Electron density" is another way of talking about how likely you are to find an electron at a particular place.)

2s (and 3s, 4s, etc) electrons spend some of their time closer to the nucleus than you might expect. The effect of this is to slightly reduce the energy of electrons in s orbitals. The nearer the nucleus the electrons get, the lower their energy.

3s, 4s (etc) orbitals get progressively further from the nucleus.

p orbitals

Not all electrons inhabit s orbitals (in fact, very few electrons live in s orbitals). At the first energy level, the only orbital available to electrons is the 1s orbital, but at the second level, as well as a 2s orbital, there are also orbitals called 2p orbitals.

A p orbital is rather like 2 identical balloons tied together at the nucleus. The diagram on the right is a cross-section through that 3-dimensional region of space. Once again, the orbital shows where there is a 95% chance of finding a particular electron.


Beyond A'level: If you imagine a horizontal plane through the nucleus, with one lobe of the orbital above the plane and the other beneath it, there is a zero probability of finding the electron on that plane. So how does the electron get from one lobe to the other if it can never pass through the plane of the nucleus? For A'level chemistry you just have to accept that it does! If you want to find out more, read about the wave nature of electrons.

Unlike an s orbital, a p orbital points in a particular direction - the one drawn points up and down the page.

At any one energy level it is possible to have three absolutely equivalent p orbitals pointing mutually at right angles to each other. These are arbitrarily given the symbols px, py and pz. This is simply for convenience - what you might think of as the x, y or z direction changes constantly as the atom tumbles in space.

The p orbitals at the second energy level are called 2px, 2py and 2pz. There are similar orbitals at subsequent levels - 3px, 3py, 3pz, 4px, 4py, 4pz and so on.

All levels except for the first level have p orbitals. At the higher levels the lobes get more elongated, with the most likely place to find the electron more distant from the nucleus.

Fitting electrons into orbitals

Because for the moment we are only interested in the electronic structures of hydrogen and carbon, we don't need to concern ourselves with what happens beyond the second energy level.

Remember:

At the first level there is only one orbital - the 1s orbital.

At the second level there are four orbitals - the 2s, 2px, 2py and 2pz orbitals.

Each orbital can hold either 1 or 2 electrons, but no more.

"Electrons-in-boxes"

Orbitals can be represented as boxes with the electrons in them shown as arrows. Often an up-arrow and a down-arrow are used to show that the electrons are in some way different.


Beyond A'level: The need to have all electrons in an atom different comes out of quantum theory. If they live in different orbitals, that's fine - but if they are both in the same orbital there has to be some subtle distinction between them. Quantum theory allocates them a property known as "spin" - which is what the arrows are intended to suggest.

A 1s orbital holding 2 electrons would be drawn as shown on the right, but it can be written even more quickly as 1s2. This is read as "one s two" - not as "one s squared".

You mustn't confuse the two numbers in this notation:

The order of filling orbitals

Electrons fill low energy orbitals (closer to the nucleus) before they fill higher energy ones. Where there is a choice between orbitals of equal energy, they fill the orbitals singly as far as possible.

The diagram (not to scale) summarises the energies of the various orbitals in the first and second levels.

Notice that the 2s orbital has a slightly lower energy than the 2p orbitals. That means that the 2s orbital will fill with electrons before the 2p orbitals. All the 2p orbitals have exactly the same energy.

The electronic structure of hydrogen

Hydrogen only has one electron and that will go into the orbital with the lowest energy - the 1s orbital.

Hydrogen has an electronic structure of 1s1. We have already described this orbital earlier.

The electronic structure of carbon

Carbon has six electrons. Two of them will be found in the 1s orbital close to the nucleus. The next two will go into the 2s orbital. The remaining ones will be in two separate 2p orbitals. This is because the p orbitals all have the same energy and the electrons prefer to be on their own if that's the case.



Note: People sometimes wonder why the electrons choose to go into the 2px and 2py orbitals rather than the 2pz. They don't! All of the 2p orbitals are exactly equivalent, and the names we give them are entirely arbitrary. It just looks tidier if we call the orbitals the electrons occupy the 2px and 2py.

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