What is an atomic orbital?
ATOMIC ORBITALS
What is an atomic orbital?
Orbitals and orbits
When the a planet moves around the sun, you can plot a definite path for it which is called an orbit. A simple view of the atom looks similar and you may have pictured the electrons as orbiting around the nucleus. The truth is different, and electrons in fact inhabit regions of space known as 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.
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
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.
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.
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.
d and f orbitals
In addition to s and p orbitals, there are two other sets of orbitals which become available for electrons to inhabit at higher energy levels. At the third level, there is a set of five d orbitals (with complicated shapes and names) as well as the 3s and 3p orbitals (3px, 3py, 3pz). At the third level there are a total of nine orbitals altogether.
At the fourth level, as well the 4s and 4p and 4d orbitals there are an additional seven f orbitals - 16 orbitals in all. s, p, d and f orbitals are then available at all higher energy levels as well.
For A'level purposes, you have to be aware that there are sets of five d orbitals at levels from the third level upwards, but you will not be expected to draw them or name them. Apart from a passing reference, you won't come across f orbitals at all.
Fitting electrons into orbitals
You can think of an atom as a very bizarre house (like an inverted pyramid!) - with the nucleus living on the ground floor, and then various rooms (orbitals) on the higher floors occupied by the electrons. On the first floor there is only 1 room (the 1s orbital); on the second floor there are 4 rooms (the 2s, 2px, 2py and 2pz orbitals); on the third floor there are 9 rooms (one 3s orbital, three 3p orbitals and five 3d orbitals); and so on. But the rooms aren't very big . . . Each orbital can only hold 2 electrons.
A convenient way of showing the orbitals that the electrons live in is to draw "electrons-in-boxes".
"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.
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.
This filling of orbitals singly where possible is known as Hund's rule. It only applies where the orbitals have exactly the same energies (as with p orbitals, for example), and helps to minimise the repulsions between electrons and so makes the atom more stable.
The diagram (not to scale) summarises the energies of the orbitals up to the 4p level.
Notice that the s orbital always has a slightly lower energy than the p orbitals at the same energy level, so the s orbital always fills with electrons before the corresponding p orbitals.
The real oddity is the position of the 3d orbitals. They are at a slightly higher level than the 4s - and so it is the 4s orbital which will fill first, followed by all the 3d orbitals and then the 4p orbitals. Similar confusion occurs at higher levels, with so much overlap between the energy levels that the 4f orbitals don't fill until after the 6s, for example.
For A'level purposes you simply have to remember that the 4s orbital fills before the 3d orbitals. The same thing happens at the next level as well - the 5s orbital fills before the 4d orbitals. All the other complications are beyond A'level.
Knowing the order of filling is central to understanding how to write electronic structures. Follow the link below to find out how to do this.
ELECTRONIC STRUCTURES
The first period
Hydrogen has its only electron in the 1s orbital - 1s1, and at helium the first level is completely full - ...
This is a preview of the whole essay
For A'level purposes you simply have to remember that the 4s orbital fills before the 3d orbitals. The same thing happens at the next level as well - the 5s orbital fills before the 4d orbitals. All the other complications are beyond A'level.
Knowing the order of filling is central to understanding how to write electronic structures. Follow the link below to find out how to do this.
ELECTRONIC STRUCTURES
The first period
Hydrogen has its only electron in the 1s orbital - 1s1, and at helium the first level is completely full - 1s2.
The second period
Now we need to start filling the second level, and hence start the second period. Lithium's electron goes into the 2s orbital because that has a lower energy than the 2p orbitals. Lithium has an electronic structure of 1s22s1. Beryllium adds a second electron to this same level - 1s22s2.
Now the 2p levels start to fill. These levels all have the same energy, and so the electrons go in singly at first.
B
s22s22px1
C
s22s22px12py1
N
s22s22px12py12pz1
The next electrons to go in will have to pair up with those already there.
O
s22s22px22py12pz1
F
s22s22px22py22pz1
Ne
s22s22px22py22pz2
You can see that it is going to get progressively tedious to write the full electronic structures of atoms as the number of electrons increases. There are two ways around this, and you must be familiar with both.
Shortcut 1: All the various p electrons can be lumped together. For example, fluorine could be written as 1s22s22p5, and neon as 1s22s22p6.
This is what is normally done if the electrons are in an inner layer. If the electrons are in the bonding level (those on the outside of the atom), they are sometimes written in shorthand, sometimes in full. Don't worry about this. Be prepared to meet either version, but if you are asked for the electronic structure of something in an exam, write it out in full showing all the px, py and pz orbitals in the outer level separately.
For example, although we haven't yet met the electronic structure of chlorine, you could write it as 1s22s22p63s23px23py23pz1.
Notice that the 2p electrons are all lumped together whereas the 3p ones are shown in full. The logic is that the 3p electrons will be involved in bonding because they are on the outside of the atom, whereas the 2p electrons are buried deep in the atom and aren't really of any interest.
Shortcut 2: You can lump all the inner electrons together using, for example, the symbol [Ne]. In this context, [Ne] means the electronic structure of neon - in other words: 1s22s22px22py22pz2 You wouldn't do this with helium because it takes longer to write [He] than it does 1s2.
On this basis the structure of chlorine would be written [Ne]3s23px23py23pz1.
The third period
At neon, all the second level orbitals are full, and so after this we have to start the third period with sodium. The pattern of filling is now exactly the same as in the previous period, except that everything is now happening at the 3-level.
For example:
short version
Mg
s22s22p63s2
[Ne]3s2
S
s22s22p63s23px23py13pz1
[Ne]3s23px23py13pz1
Ar
s22s22p63s23px23py23pz2
[Ne]3s23px23py23pz2
The beginning of the fourth period
At this point the 3-level orbitals aren't all full - the 3d levels haven't been used yet. But if you refer back to the energies of the orbitals, you will see that the next lowest energy orbital is the 4s - so that fills next.
K
s22s22p63s23p64s1
Ca
s22s22p63s23p64s2
There is strong evidence for this in the similarities in the chemistry of elements like sodium (1s22s22p63s1) and potassium (1s22s22p63s23p64s1)
The outer electron governs their properties and that electron is in the same sort of orbital in both of the elements. That wouldn't be true if the outer electron in potassium was 3d1.
s- and p-block elements
The elements in group 1 of the Periodic Table all have an outer electronic structure of ns1 (where n is a number between 2 and 7). All group 2 elements have an outer electronic structure of ns2. Elements in groups 1 and 2 are described as s-block elements.
Elements from group 3 across to the noble gases all have their outer electrons in p orbitals. These are then described as p-block elements.
d-block elements
Remember that the 4s orbital has a lower energy than the 3d orbitals and so fills first. Once the 3d orbitals have filled up, the next electrons go into the 4p orbitals as you would expect.
d-block elements are elements in which the last electron to be added to the atom is in a d orbital. The first series of these contains the elements from scandium to zinc, which at GCSE you probably called transition elements or transition metals. The terms "transition element" and "d-block element" don't quite have the same meaning, but it doesn't matter in the present context.
d electrons are almost always described as, for example, d5 or d8 - and not written as separate orbitals. Remember that there are five d orbitals, and that the electrons will inhabit them singly as far as possible. Up to 5 electrons will occupy orbitals on their own. After that they will have to pair up.
d5 means
d8 means
Notice in what follows that all the 3-level orbitals are written together, even though the 3d electrons are added to the atom after the 4s.
Sc
s22s22p63s23p63d14s2
Ti
s22s22p63s23p63d24s2
V
s22s22p63s23p63d34s2
Cr
s22s22p63s23p63d54s1
Whoops! Chromium breaks the sequence. In chromium, the electrons in the 3d and 4s orbitals rearrange so that there is one electron in each orbital. It would be convenient if the sequence was tidy - but it's not!
Mn
s22s22p63s23p63d54s2
(back to being tidy again)
Fe
s22s22p63s23p63d64s2
Co
s22s22p63s23p63d74s2
Ni
s22s22p63s23p63d84s2
Cu
s22s22p63s23p63d104s1
(another awkward one!)
Zn
s22s22p63s23p63d104s2
And at zinc the process of filling the d orbitals is complete.
Filling the rest of period 4
The next orbitals to be used are the 4p, and these fill in exactly the same way as the 2p or 3p. We are back now with the p-block elements from gallium to krypton. Bromine, for example, is 1s22s22p63s23p63d104s24px24py24pz1.
Summary
Writing the electronic structure of an element from hydrogen to krypton
* Use the Periodic Table to find the atomic number, and hence number of electrons.
* Fill up orbitals in the order 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p - until you run out of electrons. The 3d is the awkward one - remember that specially. Fill p and d orbitals singly as far as possible before pairing electrons up.
* Remember that chromium and copper have electronic structures which break the pattern in the first row of the d-block.
Writing the electronic structure of big s- or p-block elements
First work out the number of outer electrons. This is quite likely all you will be asked to do anyway.
The number of outer electrons is the same as the group number. (The noble gases are a bit of a problem here, because they are normally called group 0 rather then group 8. Helium has 2 outer electrons; the rest have 8.) All elements in group 3, for example, have 3 electrons in their outer level. Fit these electrons into s and p orbitals as necessary. Which level orbitals? Count the periods in the Periodic Table (not forgetting the one with H and He in it).
Iodine is in group 7 and so has 7 outer electrons. It is in the fifth period and so its electrons will be in 5s and 5p orbitals. Iodine has the outer structure 5s25px25py25pz1.
What about the inner electrons if you need to work them out as well? The 1, 2 and 3 levels will all be full, and so will the 4s, 4p and 4d. The 4f levels don't fill until after anything you will be asked about at A'level. Just forget about them! That gives the full structure: 1s22s22p63s23p63d104s24p64d105s25px25py25pz1.
When you've finished, count all the electrons to make sure that they come to the same as the atomic number. Don't forget to make this check - it's easy to miss an orbital out when it gets this complicated.
Barium is in group 2 and so has 2 outer electrons. It is in the sixth period. Barium has the outer structure 6s2.
Including all the inner levels: 1s22s22p63s23p63d104s24p64d105s25p66s2.
It would be easy to include 5d10 as well by mistake, but the d level always fills after the next s level - so 5d fills after 6s just as 3d fills after 4s. As long as you counted the number of electrons you could easily spot this mistake because you would have 10 too many.
IONISATION ENERGY
Defining first ionisation energy
Definition
The first ionisation energy is the energy required to remove the most loosely held electron from one mole of gaseous atoms to produce 1 mole of gaseous ions each with a charge of 1+.
This is more easily seen in symbol terms.
It is the energy needed to carry out this change per mole of X.
Things to notice about the equation
The state symbols - (g) - are essential. When you are talking about ionisation energies, everything must be present in the gas state.
Ionisation energies are measured in kJ mol-1 (kilojoules per mole). They vary in size from 381 (which you would consider very low) up to 2370 (which is very high).
All elements have a first ionisation energy - even atoms which don't form positive ions in test tubes. The reason that helium (1st I.E. = 2370 kJ mol-1) doesn't normally form a positive ion is because of the huge amount of energy that would be needed to remove one of its electrons.
Patterns of first ionisation energies in the Periodic Table
The first 20 elements
First ionisation energy shows periodicity. That means that it varies in a repetitive way as you move through the Periodic Table. For example, look at the pattern from Li to Ne, and then compare it with the identical pattern from Na to Ar.
These variations in first ionisation energy can all be explained in terms of the structures of the atoms involved.
Factors affecting the size of ionisation energy
Ionisation energy is a measure of the energy needed to pull a particular electron away from the attraction of the nucleus. A high value of ionisation energy shows a high attraction between the electron and the nucleus.
The size of that attraction will be governed by:
The charge on the nucleus.
The more protons there are in the nucleus, the more positively charged the nucleus is, and the more strongly electrons are attracted to it.
The distance of the electron from the nucleus.
Attraction falls off very rapidly with distance. An electron close to the nucleus will be much more strongly attracted than one further away.
The number of electrons between the outer electrons and the nucleus.
Consider a sodium atom, with the electronic structure 2,8,1. (There's no reason why you can't use this notation if it's useful!)
If the outer electron looks in towards the nucleus, it doesn't see the nucleus sharply. Between it and the nucleus there are the two layers of electrons in the first and second levels. The 11 protons in the sodium's nucleus have their effect cut down by the 10 inner electrons. The outer electron therefore only feels a net pull of approximately 1+ from the centre. This lessening of the pull of the nucleus by inner electrons is known as screening or shielding.
Whether the electron is on its own in an orbital or paired with another electron.
Two electrons in the same orbital experience a bit of repulsion from each other. This offsets the attraction of the nucleus, so that paired electrons are removed rather more easily than you might expect.
Explaining the pattern in the first few elements
Hydrogen has an electronic structure of 1s1. It is a very small atom, and the single electron is close to the nucleus and therefore strongly attracted. There are no electrons screening it from the nucleus and so the ionisation energy is high (1310 kJ mol-1).
Helium has a structure 1s2. The electron is being removed from the same orbital as in hydrogen's case. It is close to the nucleus and unscreened. The value of the ionisation energy (2370 kJ mol-1) is much higher than hydrogen, because the nucleus now has 2 protons attracting the electrons instead of 1.
Lithium is 1s22s1. Its outer electron is in the second energy level, much more distant from the nucleus. You might argue that that would be offset by the additional proton in the nucleus, but the electron doesn't feel the full pull of the nucleus - it is screened by the 1s2 electrons.
You can think of the electron as feeling a net 1+ pull from the centre (3 protons offset by the two 1s2 electrons).
If you compare lithium with hydrogen (instead of with helium), the hydrogen's electron also feels a 1+ pull from the nucleus, but the distance is much greater with lithium. Lithium's first ionisation energy drops to 519 kJ mol-1 whereas hydrogen's is 1310 kJ mol-1.
The patterns in periods 2 and 3
Talking through the next 17 atoms one at a time would take ages. We can do it much more neatly by explaining the main trends in these periods, and then accounting for the exceptions to these trends.
The first thing to realise is that the patterns in the two periods are identical - the difference being that the ionisation energies in period 3 are all lower than those in period 2.
Explaining the general trend across periods 2 and 3
The general trend is for ionisation energies to increase across a period.
In the whole of period 2, the outer electrons are in 2-level orbitals - 2s or 2p. These are all the same sort of distances from the nucleus, and are screened by the same 1s2 electrons.
The major difference is the increasing number of protons in the nucleus as you go from lithium to neon. That causes greater attraction between the nucleus and the electrons and so increases the ionisation energies. In fact the increasing nuclear charge also drags the outer electrons in closer to the nucleus. That increases ionisation energies still more as you go across the period.
In period 3, the trend is exactly the same. This time, all the electrons being removed are in the third level and are screened by the 1s22s22p6 electrons. They all have the same sort of environment, but there is an increasing nuclear charge.
Why the drop between groups 2 and 3 (Be-B and Mg-Al)?
The explanation lies with the structures of boron and aluminium. The outer electron is removed more easily from these atoms than the general trend in their period would suggest.
Be
s22s2
st I.E. = 900 kJ mol-1
B
s22s22px1
st I.E. = 799 kJ mol-1
You might expect the boron value to be more than the beryllium value because of the extra proton. Offsetting that is the fact that boron's outer electron is in a 2p orbital rather than a 2s. 2p orbitals have a slightly higher energy than the 2s orbital, and the electron is, on average, to be found further from the nucleus. This has two effects.
* The increased distance results in a reduced attraction and so a reduced ionisation energy.
* The 2p orbital is screened not only by the 1s2 electrons but, to some extent, by the 2s2 electrons as well. That also reduces the pull from the nucleus and so lowers the ionisation energy.
The explanation for the drop between magnesium and aluminium is the same, except that everything is happening at the 3-level rather than the 2-level.
Mg
s22s22p63s2
st I.E. = 736 kJ mol-1
Al
s22s22p63s23px1
st I.E. = 577 kJ mol-1
The 3p electron in aluminium is slightly more distant from the nucleus than the 3s, and partially screened by the 3s2 electrons as well as the inner electrons. Both of these factors offset the effect of the extra proton.
Why the drop between groups 5 and 6 (N-O and P-S)?
Once again, you might expect the ionisation energy of the group 6 element to be higher than that of group 5 because of the extra proton. What is offsetting it this time?
N
s22s22px12py12pz1
st I.E. = 1400 kJ mol-1
O
s22s22px22py12pz1
st I.E. = 1310 kJ mol-1
The screening is identical (from the 1s2 and, to some extent, from the 2s2 electrons), and the electron is being removed from an identical orbital.
The difference is that in the oxygen case the electron being removed is one of the 2px2 pair. The repulsion between the two electrons in the same orbital means that the electron is easier to remove than it would otherwise be.
The drop in ionisation energy at sulphur is accounted for in the same way.
Trends in ionisation energy down a group
As you go down a group in the Periodic Table ionisation energies generally fall. You have already seen evidence of this in the fact that the ionisation energies in period 3 are all less than those in period 2.
Taking Group 1 as a typical example:
Why is the sodium value less than that of lithium?
There are 11 protons in a sodium atom but only 3 in a lithium atom, so the nuclear charge is much greater. You might have expected a much larger ionisation energy in sodium, but offsetting the nuclear charge is a greater distance from the nucleus and more screening.
Li
s22s1
st I.E. = 519 kJ mol-1
Na
s22s22p63s1
st I.E. = 494 kJ mol-1
Lithium's outer electron is in the second level, and only has the 1s2 electrons to screen it. The 2s1 electron feels the pull of 3 protons screened by 2 electrons - a net pull from the centre of 1+.
The sodium's outer electron is in the third level, and is screened from the 11 protons in the nucleus by a total of 10 inner electrons. The 3s1 electron also feels a net pull of 1+ from the centre of the atom. In other words, the effect of the extra protons is compensated for by the effect of the extra screening electrons. The only factor left is the extra distance between the outer electron and the nucleus in sodium's case. That lowers the ionisation energy.
Similar explanations hold as you go down the rest of this group - or, indeed, any other group.
Trends in ionisation energy in a transition series
Apart from zinc at the end, the other ionisation energies are all much the same.
All of these elements have an electronic structure [Ar]3dn4s2 (or 4s1 in the cases of chromium and copper). The electron being lost always comes from the 4s orbital.
As you go from one atom to the next in the series, the number of protons in the nucleus increases, but so also does the number of 3d electrons. The 3d electrons have some screening effect, and the extra proton and the extra 3d electron more or less cancel each other out as far as attraction from the centre of the atom is concerned.
The rise at zinc is easy to explain.
Cu
[Ar]3d104s1
st I.E. = 745 kJ mol-1
Zn
[Ar]3d104s2
st I.E. = 908 kJ mol-1
In each case, the electron is coming from the same orbital, with identical screening, but the zinc has one extra proton in the nucleus and so the attraction is greater.
Ionisation energies and reactivity
The lower the ionisation energy, the more easily this change happens:
You can explain the increase in reactivity of the Group 1 metals (Li, Na, K, Rb, Cs) as you go down the group in terms of the fall in ionisation energy. Whatever these metals react with, they have to form positive ions in the process, and so the lower the ionisation energy, the more easily those ions will form.
The danger with this approach is that the formation of the positive ion is only one stage in a multi-step process.
For example, you wouldn't be starting with gaseous atoms; nor would you end up with gaseous positive ions - you would end up with ions in a solid or in solution. The energy changes in these processes also vary from element to element. Ideally you need to consider the whole picture and not just one small part of it.
However, the ionisation energies of the elements are going to be major contributing factors towards the activation energy of the reactions. Remember that activation energy is the minimum energy needed before a reaction will take place. The lower the activation energy, the faster the reaction will be - irrespective of what the overall energy changes in the reaction are.
The fall in ionisation energy as you go down a group will lead to lower activation energies and therefore faster reactions
ELECTRON AFFINITY
First electron affinity
Ionisation energies are always concerned with the formation of positive ions. Electron affinities are the negative ion equivalent, and their use is almost always confined to elements in groups 6 and 7 of the Periodic Table.
Defining first electron affinity
The first electron affinity is the energy released when 1 mole of gaseous atoms each acquire an electron to form 1 mole of gaseous 1- ions.
This is more easily seen in symbol terms.
It is the energy released (per mole of X) when this change happens.
First electron affinities have negative values. For example, the first electron affinity of chlorine is -349 kJ mol-1. By convention, the negative sign shows a release of energy.
The first electron affinities of the group 7 elements
F
-328 kJ mol-1
Cl
-349 kJ mol-1
Br
-324 kJ mol-1
I
-295 kJ mol-1
Is there a pattern?
Yes - as you go down the group, first electron affinities become less (in the sense that less energy is evolved when the negative ions are formed). Fluorine breaks that pattern, and will have to be accounted for separately.
The electron affinity is a measure of the attraction between the incoming electron and the nucleus - the stronger the attraction, the more energy is released.
The factors which affect this attraction are exactly the same as those relating to ionisation energies - nuclear charge, distance and screening.
The increased nuclear charge as you go down the group is offset by extra screening electrons. Each outer electron in effect feels a pull of 7+ from the centre of the atom, irrespective of which element you are talking about.
For example, a fluorine atom has an electronic structure of 1s22s22px22py22pz1. It has 9 protons in the nucleus.
The incoming electron enters the 2-level, and is screened from the nucleus by the two 1s2 electrons. It therefore feels a net attraction from the nucleus of 7+ (9 protons less the 2 screening electrons).
By contrast, chlorine has the electronic structure 1s22s22p63s23px23py23pz1. It has 17 protons in the nucleus.
But again the incoming electron feels a net attraction from the nucleus of 7+ (17 protons less the 10 screening electrons in the first and second levels).
The over-riding factor is therefore the increased distance that the incoming electron finds itself from the nucleus as you go down the group. The greater the distance, the less the attraction and so the less energy is released as electron affinity.
Why is fluorine out of line?
The incoming electron is going to be closer to the nucleus in fluorine than in any other of these elements, so you would expect a high value of electron affinity.
However, because fluorine is such a small atom, you are putting the new electron into a region of space already crowded with electrons and there is a significant amount of repulsion. This repulsion lessens the attraction the incoming electron feels and so lessens the electron affinity.
A similar reversal of the expected trend happens between oxygen and sulphur in Group 6. The first electron affinity of oxygen (-142 kJ mol-1) is smaller than that of sulphur (-200 kJ mol-1) for exactly the same reason that fluorine's is smaller than chlorine's.
Comparing Group 6 and Group 7 values
As you might have noticed, the first electron affinity of oxygen (-142 kJ mol-1) is less than that of fluorine (-328 kJ mol-1). Similarly sulphur's (-200 kJ mol-1) is less than chlorine's (-349 kJ mol-1). Why?
It's simply that the Group 6 element has 1 less proton in the nucleus than its next door neighbour in Group 7. The amount of screening is the same in both.
That means that the net pull from the nucleus is less in Group 6 than in Group 7, and so the electron affinities are less.
First electron affinity and reactivity
The reactivity of the elements in group 7 falls as you go down the group - fluorine is the most reactive and iodine the least.
Often in their reactions these elements form their negative ions. At GCSE the impression is sometimes given that the fall in reactivity is because the incoming electron is held less strongly as you go down the group and so the negative ion is less likely to form. That explanation looks reasonable until you include fluorine!
An overall reaction will be made up of lots of different steps all involving energy changes, and you cannot safely try to explain a trend in terms of just one of those steps. Fluorine is much more reactive than chlorine (despite the lower electron affinity) because the energy released in other steps in its reactions more than makes up for the lower amount of energy released as electron affinity.
Second electron affinity
You are only ever likely to meet this with respect to the group 6 elements oxygen and sulphur which both form 2- ions.
Defining second electron affinity
The second electron affinity is the energy required to add an electron to each ion in 1 mole of gaseous 1- ions to produce 1 mole of gaseous 2- ions.
This is more easily seen in symbol terms.
It is the energy needed to carry out this change per mole of X-.
Why is energy needed to do this?
You are forcing an electron into an already negative ion. It's not going to go in willingly!
st EA = -142 kJ mol-1
2nd EA = +844 kJ mol-1
The positive sign shows that you have to put in energy to perform this change. The second electron affinity of oxygen is particularly high because the electron is being forced into a small, very electron-dense space.
ATOMIC RADIUS
Measures of atomic radius
Unlike a ball, an atom doesn't have a fixed radius. The radius of an atom can only be found by measuring the distance between the nuclei of two touching atoms, and then halving that distance.
As you can see from the diagrams, the same atom could be found to have a different radius depending on what was around it.
The left hand diagram shows bonded atoms. The atoms are pulled closely together and so the measured radius is less than if they are just touching. This is what you would get if you had metal atoms in a metallic structure, or atoms covalently bonded to each other. The type of atomic radius being measured here is called the metallic radius or the covalent radius depending on the bonding.
The right hand diagram shows what happens if the atoms are just touching. The attractive forces are much less, and the atoms are essentially "unsquashed". This measure of atomic radius is called the van der Waals radius after the weak attractions present in this situation.
Trends in atomic radius in the Periodic Table
The exact pattern you get depends on which measure of atomic radius you use - but the trends are still valid.
The following diagram uses metallic radii for metallic elements, covalent radii for elements that form covalent bonds, and van der Waals radii for those (like the noble gases) which don't form bonds.
Trends in atomic radius in Periods 2 and 3
Trends in atomic radius down a group
It is fairly obvious that the atoms get bigger as you go down groups. The reason is equally obvious - you are adding extra layers of electrons.
Trends in atomic radius across periods
You have to ignore the noble gas at the end of each period. Because neon and argon don't form bonds, you can only measure their van der Waals radius - a case where the atom is pretty well "unsquashed". All the other atoms are being measured where their atomic radius is being lessened by strong attractions. You aren't comparing like with like if you include the noble gases.
Leaving the noble gases out, atoms get smaller as you go across a period.
If you think about it, the metallic or covalent radius is going to be a measure of the distance from the nucleus to the electrons which make up the bond. (Look back to the left-hand side of the first diagram on this page if you aren't sure, and picture the bonding electrons as being half way between the two nuclei.)
From lithium to fluorine, those electrons are all in the 2-level, being screened by the 1s2 electrons. The increasing number of protons in the nucleus as you go across the period pulls the electrons in more tightly. The amount of screening is constant for all of these elements.
In the period from sodium to chlorine, the same thing happens. The size of the atom is controlled by the 3-level bonding electrons being pulled closer to the nucleus by increasing numbers of protons - in each case, screened by the 1- and 2-level electrons.
Trends in the transition elements
Although there is a slight contraction at the beginning of the series, the atoms are all much the same size.
The size is determined by the 4s electrons. The pull of the increasing number of protons in the nucleus is more or less offset by the extra screening due to the increasing number of 3d electrons.
IONIC RADIUS
Ions aren't the same size as the atoms they come from. Compare the sizes of sodium and chloride ions with the sizes of sodium and chlorine atoms.
Positive ions
Positive ions are smaller than the atoms they come from. Sodium is 2,8,1; Na+ is 2,8. You've lost a whole layer of electrons, and the remaining 10 electrons are being pulled in by the full force of 11 protons.
Negative ions
Negative ions are bigger than the atoms they come from. Chlorine is 2,8,7; Cl- is 2,8,8. Although the electrons are still all in the 3-level, the extra repulsion produced by the incoming electron causes the atom to expand. There are still only 17 protons, but they are now having to hold 18 electrons.
