atomic structure and bonding menu

8/6/2019 Atomic Structure and Bonding Menu

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ATOMIC STRUCTURE AND BONDING MENU

Basic atomic properties . . .

Includes a discussion of orbitals, electronic structures of atoms and ions, ionisation energies, electron affinities, atomic and ionic radii,and the atomic hydrogen emission spectrum.

Bonding . . .

Includes ionic, covalent, co-ordinate (dative covalent) and metallic bonding as well as intermolecular attractions like Van der Waals forces

and hydrogen bonding. Also includes full discussions of electronegativity and shapes of molecules and ions.

Types of structure . . .

Describes and explains how the various types of structure (ionic, giant covalent, metallic, and molecular) affect physical pr operties.

ATOMIC PROPERTIES MENU

Simple background . . .

Revises the simple knowledge you should already have about the structure of atoms from introductory courses (e.g. GCSE).

Atomic orbitals . . .

Explains what atomic orbitals are and discusses their shapes and relative energies. This is essential pre -reading before you go on to anyof the remaining topics in this section.

Electronic structures . . .

How to work out and write the electronic structures for atoms and simple monatomic ions (containing only one atom - e.g. Cl- or Mg2+)using s, p, d notation.

Ionisation energies . . .

Explains what ionisation energies are and how and why they vary around the Periodic Table.

Hydrogen's atomic emission spectrum . . .

An introduction to the atomic hydrogen emission spectrum, and how it can be used to find the ionisation energy of hydrogen.

Electron affinities . . .

Explains what electron affinities are and how and why they vary around the Periodic Table.

Atomic and ionic radii . . .

Looks at the various measures of atomic radius, and explains how and why atomic radii vary around the Periodic Table. Also co nsidershow the radii of positive and negative ions differ from the atoms they come fr om.

A SIMPLE VIEW OF ATOMIC STRUCTURE

The sub-atomic particles

Protons, neutrons and electrons.


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relative mass relative chargeproton 1 +1neutron 1 0electron 1/1836 -1

Beyond A'level: Protons and neutrons don't in fact h ave exactly the same mass - neither of them has a mass of exactly 1 on the carbon-12 scale (the scale on which the relative masses ofatoms are measured). On the carbon-12 scale, a proton has a mass of 1.0073, and a neutron a mass of 1.0087.

The behaviour of protons, neutrons and electrons in electric fields

What happens if a beam of each of these particles is passed between two electrically charged plates - one positive and one negative? Opposites willattract.

Protons are positively charged and so would be deflected on a curving path towards the negative plate.

Electrons are negatively charged and so would be deflected on a curving path towards the positive plate.

Neutrons don't have a charge, and so would continue on in a straight line.

Exactly what happens depends on whether the beams of particles enter the electric field with the various particles having the same speeds or thesame energies

If the particles have the same energy

If beams of the three sorts of particles, all with the same energy, are passed between two electrically charged plates:

y Protons are deflected on a curved path towards the negative plate.

y Electrons are deflected on a curved path towards the positive plate.

The amount of deflection is exactly th e same in the electron beam as the proton beam if the energies are the same - but, of course, it is inthe opposite direction.

y Neutrons continue in a straight line.

If the electric field was strong enough, then the electron and proton beams might curve enough to hit their respective plates.

If the particles have the same speeds

If beams of the three sorts of particles, all with the same speed, are passed between two electrically charged plates:

y Protons are deflected on a curved path towards the negative plate.

y Electrons are deflected on a curved path towards the positive plate.

If the electrons and protons are travelling with the same speed, then the lighter electrons are deflected far more strongly t han the heavierprotons.

y Neutrons continue in a straight line.


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Note: This is potent ially very confusing! Most chemistry sources that ta lk about this give either one or the other of these two diagrams without any comment at a ll - they don't specifically say

that they are using constant energy or constant speed beams. But it matters!

If this is on your syllabus, it is important that you should know which version your examiners are going to expect, and they probably won't tell you in the syllabus. You should look in detail at

past questions, mark schemes and examiner's reports which you can get from your examiners if you are doing a UK-based syllabus. Information about how to do this is on the syllabusespage.

If in doubt, I suggest you use the second (constant speed) version. This actually produces more useful information about both masses and charges than the constant energy version.

The nucleus

The nucleus is at the centre of the atom and contains the protons and neutrons. Protons and neutrons are collectively known a s nucleons.

Virtually all the mass of the atom is concentrated in the nucleus, because the electrons weigh so little.

Working out the numbers of protons and neutrons

No of protons = ATOMIC NUMBER of the atom

The atomic number is also given the more descriptive name of proton number.

No of protons + no of neutrons = MASS NUMBER of the atom

The mass number is also called the nucleon number.

This information can be given simply in the form:

How many protons and neutrons has this atom got?

The atomic number counts the number of protons (9); the mass number counts protons + neutrons (19). If there are 9 protons, t here must be 10neutrons for the total to add up to 19.

The atomic number is tied to the position of the element in the Periodic Table and therefore the number of protons defines wh at sort of element youare talking about. So if an atom has 8 protons (atomic number = 8), it m ust be oxygen. If an atom has 12 protons (atomic number = 12), it must bemagnesium.

Similarly, every chlorine atom (atomic number = 17) has 17 protons; every uranium atom (atomic number = 92) has 92 protons.

Isotopes

The number of neutrons in an atom can vary within small limits. For example, there are three kinds of carbon atom 12C, 13C and 14C. They all have thesame number of protons, but the number of neutrons varies.

protons neutrons mass numbercarbon-12 6 6 12carbon-13 6 7 13carbon-14 6 8 14


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These different atoms of carbon are called isotopes. The fact that they have varying numbers of neutrons makes no difference whatsoever to thechemical reactions of the carbon.

Isotopes are atoms which have the same atomic number but different mass number s. They have the same number of protons but different numbersof neutrons.

The electrons

Working out the number of electrons

Atoms are electrically neutral, and the positiveness of the protons is balanced by the negativeness of the electrons. It follows that in a neutral atom:

no of electrons = no of protons

So, if an oxygen atom (atomic number = 8) has 8 protons, it must also have 8 electrons; if a chlorine atom (atomic number = 1 7) has 17 protons, itmust also have 17 electrons.

The arrangement of the electrons

The electrons are found at considerable distances from the nucleus in a series of levels called energy levels. Each energy le vel can only hold acertain number of electrons. The first level (nearest the nucleus) will only hold 2 electrons, the second holds 8, and the third also seems to be fullwhen it has 8 electrons. At GCSE you stop there because the pattern gets more complicated after that.

These levels can be thought of as getting progressively further from the n ucleus. Electrons will always go into the lowest possible energy level(nearest the nucleus) - provided there is space.

To work out the electronic arrangement of an atom

y Look up the atomic number in the Periodic Table - making sure that you choose the right number if two numbers are given. The atomicnumber will always be the smaller one.

y This tells you the number of protons, and hence the number of electrons.

y Arrange the electrons in levels, always filling up an inner level before you go to an outer one.

e.g. to find the electronic arrangement in chlorine

y

The Periodic Table gives you the atomic number of 17.y Therefore there are 17 protons and 17 electrons.

y The arrangement of the electrons will be 2, 8, 7 (i.e. 2 in the first level, 8 in the second, and 7 in the third).

The electronic arrangements of the first 20 elements

After this the pattern alters as you enter the transition series in the Periodic Table.

Two important generalisations

If you look at the patterns in this table:

y The number of electrons in the outer level is the same as the group number. (Except with helium which has only 2 electrons. The noblegases are also usually called group 0 - not group 8.) This pattern extends throughout the Periodic Table for the main groups (i.e. notincluding the transition elements).

So if you know that barium is in group 2, it has 2 electrons in its outer level; iodine (group 7) has 7 electrons in its oute r level; lead (group4) has 4 electrons in its outer level.


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y Noble gases have full outer levels. This generalisation will need modifying for A'level purposes.

Dots-and-crosses diagrams

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

Note: There are many places where you could still make use of this model of the a tom at A'level . It is, however, a simplif ication and can be misleading. It gi ves the impression that theelectrons are circling the nucleus in orbits like planets around the sun. As you will find when you look at the A'level view of the atom, it is impossible to know exactly how they are actuallymoving.

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

Carbon, for example, would look like this:

Thinking of the arrangement of the electrons in this way makes a useful bridge to the A'level view.

Note: If you have come to this page as a UK GCSE student (or a student on a similar introductor y chemistry course elsewhere) and want some more help, you may be interested in myGCSEChemistry book. This link will take you to a page describing it.

ATOMIC ORBITALS

This page explains what atomic orbitals are in a way that makes them understandable for introductory courses such as UK A lev el and itsequivalents. It explores s and p orbitals in some detail, including their shapes and energies. d orbitals are described only in terms of their energy,and f orbitals only get a passing mention.

What is an atomic orbital?

Orbitals and orbits

When 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 l ooks similar and you mayhave 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. Youcan't do this for electrons.


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The Heisenberg Uncertainty Principle says - loosely - that you can't know with certainty both where an electron is and where it's going next. (What itactually says is that it is impossible to define with absolute precision, at the same time, both the position and the momentu m of an electron.)

That makes it impossible to plot an orbit for an electron a round a nucleus. Is this a big problem? No. If something is impossible, you have to acceptit and find a way around it.

Note: Over the years I have had a steady drip of questions from students in which it is obvious that they still think of electronsas orbiting around a nucleus - which is completely wrong! I

have added a page about why the idea oforbits is wrong to try to avoid having to say the same thing over and over again!

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 justfor clarity.

Suppose you had a single hydrogen atom and at a particular instant plotted the position of the one electron. Soonafterwards, 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 thesecond.

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 tobe found.

In the hydrogen case, the electron can be found anywhere within a spherical space surrounding the nucleus. T he diagram shows a cross-sectionthrough 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 spa ce quite close to thenucleus. 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 electro n 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 ina 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 thenucleus. 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 greatestchance 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 dotsare thicker) nearer the nucleus. ("Electron density" is another way of talking abou t how likely you are to find anelectron 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 ofthis 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 orbitalavailable to electrons is the 1s orbital, but at the second level, as well as a 2s orbital, there are also orbitals called 2porbitals.


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

Taking chemistry further: If you imagine a horizontal plane thr ough the nucleus, with on e lobe of the orbital above the plane and the other beneath i t, there is a zero probability of finding theelectron on that plane. So how does the electron get from one lobe to the other if it can never pass through the plane of thenucleus? At this introductory level 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 ea ch other. These arearbitrarily 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 2p x, 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 mostlikely 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 thirdlevel, there is a set of five d orbitals (with complicated shapes and names) as well as the 3s and 3p orbitals (3p x, 3py, 3pz). At the third level there area 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 thenavailable at all higher energy levels a s well.

For the moment, you need to be aware that there are sets of five d orbitals at levels from the third level upwards, but you p robably won't be expectedto draw them or name them. Apart from a passing reference, you won't come across f orbitals at al l.

Note: Some UK-based syllabuses will eventua lly want you to be able to draw, or at least recognise, the shapes of d orbitals. I a m not including them now because I don't wan t to add confusionto what is already a difficult introductory topic. Check yoursyllabus and past papers to find out what you need to know. If you are a studying a UK-based syllabus and haven't got these, followthis link to find out h ow to get hold of them.

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 roomsaren'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 theelectrons are in some way different.

Taking chemistry further: The need to have all electrons in an atom diff erent comes out of quantum theory. If they live in different orbitals, that's fine - but if they are both in the same orbitalthere has to be some subtle distinction bet ween 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 1s 2. This is read as "one s two" -not as "one s squared".


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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 betwe en 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 porbitals, for example), and helps to minimise the repulsions between electrons and so makes the atom more stable.

The diagram (not to scale) summarises the ene rgies 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 a lways fills with electronsbefore 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, followedby 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 4forbitals don't fill until after the 6s, for example.

ELECTRONIC STRUCTURES

This page explores how you write electronic structures for atoms using s, p, and d notation. It assumes that you know about s imple atomic orbitals -at least as far as the way they are named, and their relative energies. If you want to look at the electronic structures of s imple monatomic ions (suchas Cl-, Ca2+ and Cr3+), you will find a link at the bottom of the page.

Important! If you haven't already read the page on atomic orbitals you should follow this link before you go any further.

The electronic structures of atoms

Relating orbital filling to the Periodic Table

UK syllabuses for 16 - 18 year olds tend to stop at krypton when it comes to writing electronic structures, but it is possible that you could be ask edfor structures for elements up as far as barium. After barium you have to worry about f orbitals as well as s, p and d orbitals - and that's a problemfor chemistry at a higher level. It is important that you look through past exam papers as well as your syllabus so that you can judge how hard thequestions are likely to get.


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This page looks in detail at the elements in the shortened version of the Periodic Table above, and then shows how you could work out thestructures of some bigger atoms.

Important! You must have a copy of yoursyllabus and copies of recent exam papers. If you are studying a UK-based syllabus and haven't got them, follow this link to find out how to get holdof them.

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 lowerenergy than the 2p orbitals. Lithium has an electronic structure of 1s 22s1. Beryllium adds a second electron to this same level - 1s 22s2.

Now the 2p levels start to fill. These levels all have the same energy, and so the electrons go in singly at first.

B 1s22s22px1C 1s22s22px12py1N 1s22s22px12py12pz1Note: The orbitals where something new is happening are shown in bold type. You wouldn't normally write them any different ly from the other orbitals.

The next electrons to go in will have to pair up with those already there.

O 1s22s22px22py12pz1F 1s22s22px22py22pz1Ne 1s22s22px22py22pz2

You can see that it is going to get progressively tedious to write the full electronic structures of atoms as the number of e lectrons increases. Thereare 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 1s 22s22p5, and neon as 1s22s22p6.

This is what is normally done if the electrons are in an inner layer. If the el ectrons are in the bonding level (those on the outside of the atom), theyare sometimes written in shorthand, sometimes in full. Don't worry about this. Be prepared to meet either version, but if you are asked for theelectronic structure of something in an exam, write it out in full showing all the p x, 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 1s 22s22p63s23px23py23pz1.

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 inbonding because they are on the outside of the atom, whereas the 2p electrons are buried deep in the atom and aren't really o f 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 ofneon - in other words: 1s 22s22px22py22pz2 You wouldn't do this with helium because it takes longer to write [He] than it does 1s 2.

On this basis the structure of chlorine would be written [Ne]3s 23px23py23pz1.

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 thesame as in the previous period, except that everything is now happening at the 3-level.

For example:

short versionMg 1s22s22p63s2 [Ne]3s2S 1s22s22p63s23px23py13pz1 [Ne]3s23px23py13pz1Ar 1s22s22p63s23px23py23pz2 [Ne]3s23px23py23pz2


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Note: Check that you can do these. Cover the text and then work out these structur es for yourself. Then do all the rest of this period. When you've finished, check your answers against the

corresponding elements from the previous period. Your answers should be the same except a level further out.

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 seethat the next lowest energy orbital is the 4s - so that fills next.

K 1s22s22p63s23p64s1Ca 1s22s22p63s23p64s2There is strong evidence for this in the similarities in the chemistry of elements like sodium (1s 22s22p63s1) 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 outerelectron 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 ns 1 (where n is a number between 2 and 7). All group 2elements have an outer electronic structure of ns 2. 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 add ed to the atom is in a d orbital. The first series of these contains the elements fromscandium 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.

If you are interested: A transition element is defined as one which has partially fi lled d orbitals either in the el ement or any of its compound s. Zinc (at the right-hand end of the d-block) alwayshas a completely full 3d level (3d10) and so doesn't count as a transition element.

d electrons are almost always described as, for example, d 5 or d8 - and not written as separate orbitals. Remember that there are five d orbitals, andthat the electrons will inhabit them singly as far as possible. Up to 5 electrons will occupy orbitals on their own. After th at they will have to pair up.


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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 1s22s22p63s23p63d14s2Ti 1s22s22p63s23p63d24s2V 1s22s22p63s23p63d34s2Cr 1s22s22p63s23p63d54s1Whoops! Chromium breaks the sequence. In chromium, the electrons in the 3d and 4s orbitals rearrange so that there is one ele ctron in each orbital.It would be convenient if the sequence was tidy - but it's not!

Mn 1s22s22p63s23p63d54s2 (back to being tidy again)Fe 1s22s22p63s23p63d64s2Co 1s22s22p63s23p63d74s2Ni 1s22s22p63s23p63d84s2Cu 1s22s22p63s23p63d104s1 (another awkward one!)Zn 1s22s22p63s23p63d104s2And 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 fromgallium to krypton. Bromine, for example, is 1s 22s22p63s23p63d104s24px24py24pz1.

Useful exercise: Work out the electronic struc tures of all the elements from gall ium to krypton. You can check your answers by comparing them with the elements directl y above them in the

Periodic Table. For example, gallium will have the same sort of arrangement of its outer level electrons as boron or aluminium - except that gallium's outer electrons will be in the 4-level.

Summary

Writing the electronic structure of an element from hydrogen to krypton

y Use the Periodic Table to find the atomic number, and hence number of electrons.

y 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 thatspecially. Fill p and d orbitals singly as far as possible before pairing electr ons up.

y 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

Note: We are deliberately excluding the d-block elements apart from the fir st row that we've already looked at in detail. The pat tern of awkward structures isn't the same in the other r ows. Thisis a problem for degree level.

First work out the number of outer electrons. This is quite likely all you will be asked t o 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 calledgroup 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 fo rgetting the one with H andHe in it).


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Iodineis ingroup 7 andsohas 7 outer electrons. It is inthefifth periodandsoits electrons will bein5s and5p orbitals. Iodinehas theouterstructure5s25px25py25pz1.

What about theinnerelectrons if you needtowork them out as well? The1, 2 and3 levels will all befull, andsowill the 4s, 4p and4d. The4f levelsdon't fill until after anything you will beaskedabout at A'level. Just forget about them! That gives thefull structure:1s22s22p63s23p63d104s24p64d105s25px25py25pz1.

Whenyou'vefinished, count all theelectrons tomakesurethat they cometothesameas theatomic number. Don't forget tom akethis check - it'seasy tomiss anorbital out whenit gets this complicated.

Barium is ingroup 2 andsohas 2 outerelectrons. It is inthesixth period. Barium has theouter structure6s 2.

Including all theinner levels: 1s 22s22p63s23p63d104s24p64d105s25p66s2.

It wouldbeeasy toinclude5d10 as well by mistake, but thedlevel always fills after thenext s level - so5dfills after 6s just as 3dfills after 4s. Aslong as you countedthenumberofelectrons you couldeasily spot this mistakebecauseyou wouldhave10 toomany.

Note: Don't worry toomuch about thesecomplicatedstructures. You need toknow how towork them out inprinciple, but yourexaminers aremuch morelikely toask you for something

simplelikesulphuroriron.

IONISATION ENERGY

This pageexplains what first ionisationenergy is, andthenlooks at theway it varies aroundthePeriodic Table - across periods anddowngroups. Itassumes that you know about simpleatomic orbitals, andcanwriteelectronic structures for simpleatoms. You will finda lin k at thebottom of thepagetoa similardescriptionof successiveionisationenergies (second, thirdandsoon).

Important! If you aren't reasonablehappy about atomic orbitals andelectronic structures you should follow theselinks beforeyou goany further.

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.

Worried about moles? Don't be! For no w, just take it as a measure of a particular amount of a substance. It isn't worth worrying about at the mo ment.

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 v ery low) up to 2370 (whichis 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


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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 atthe pattern from Li to Ne, andthen compare it with the identical pattern from Na to Ar.

These variations in first ionisation energy can all be expl ainedin terms of the structures of the atoms involved.

Factors affecting the size of ionisation energy

Ionisation energy is a measure of the energy neededto pull a particular electron away from the attraction of the nucleus. A high value of ionisationenergy shows a high attraction between the electron andthe nucleus.

The size of that attraction will be governedby:

The charge on the nucleus.

The more protons there are in the nucleus, the more positively chargedthe nucleus is, andthe more strongly ele ctrons are attractedto 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 attractedthan o ne further away.

The number of electrons between the outer electrons andthe 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 usefu l!)

If the outer electron looks in towards the nucleus, it doesn't see the nucleus sharply. Betwee n it andthe nucleus there are the two layers of electronsin the first andsecondlevels. The 11 protons in the sodium's nucleus have their effect cut down by the 10 inner electrons. The outer electrontherefore only feels a net pull of approximately 1+ f rom the centre. This lessening of the pull of the nucleus by inner electrons is known as screeningor shielding.

Warning! Electrons don't, of course, "look in" towards the nucleus - and they don't "see" anything either! But there's no reason why you can't imagine it in these terms if it helps you to

visualise what's happening. Just don't use these terms in an exam! You may get an examiner who is upset by this sort of loose language.

Whethertheelectronisonitsowninanorbitalorpairedwithanotherelectron.

Twoelectronsinthesameorbitalexperienceabitofrepulsionfrom eachother. Thisoffsetstheattractionof thenucleus, sothatpairedelectronsareremovedrathermoreeasily thanyou mightexpect.

Explaining the pattern in the first few elements

Hydrogen has an electronic structure of 1s1. It is a very small atom, andthe single electron is close tothe nucleus andtherefore strongly attracted.There are noelectrons screening it from the nucleus andsothe ionisation energy is high (1310 kJ mol -1).

Helium has a structure 1s 2. The electron is being removedfrom the same orbital as in hydrogen's case. It is close tothe nucleus andunscreened.The value of the ionisation energy (2370 kJ mol -1) is much higher than hydrogen, because the nucleus now has 2 protons attracting the electronsinsteadof 1.

Lithium is 1s22s1. Its outer electron is in the secondenergy level, much more distant from the nucleus. You might argue that that wouldbe of fset bythe additional proton in the nucleus, but the electron doesn't feel the full pull of the nucleus - it is screenedby the 1s2 electrons.


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You can think of the electron as feeling a net 1+ pull from the centre (3 protons offset by the two 1s 2 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 muchgreater 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, andthen 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 lowerthan 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 arescreened 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 a ttraction between thenucleus and the electrons and so increases the ionisation energies. In fact the increasing nuclear charge also drags the oute r electrons in closer tothe nucleus. That increases ionisation energies still more as you go across the period.

Note: Factors affecting atomic radius are covered on a separate page.

In period 3, the trend is exactly the same. This time, all the electrons being removed are in the t hird level and are screened by the 1s 22s22p6electrons. 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)?


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The explanation lies with the structures of boron and aluminium. The outer electron is removed more easily from these atoms than the general trendin their period would suggest.

Be 1s22s2 1st I.E. = 900 kJ mol -1B 1s22s22px1 1st 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 outerelectron is in a 2p orbital rather than a 2s. 2p orbitals have a slightly higher energy than the 2s orbital, an d the electron is, on average, to be foundfurther from the nucleus. This has two effects.

y The increased distance results in a reduced attraction and so a reduced ionisation energy.

y The 2p orbital is screened not only by the 1s 2 electrons but, to some extent, by the 2s 2 electrons as well. That also reduces the pull fromthe 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 1s22s22p63s2 1st I.E. = 736 kJ mol -1Al 1s22s22p63s23px1 1st I.E. = 577 kJ mol -1The 3p electron in aluminium is slightly more distant from the nucleus than the 3s, and partially screened by the 3s 2 electrons as well as the innerelectrons. Both of these factors offset the effect of the extra proton.

Warning! You might possibly come across a text book which describes the drop between gr oup 2 and group 3 by saying that a full s2 orbital is in some way especially stable and that makes

the electron more difficul t to remove. In other words, that the fluctuation is because the group 2 value for ionisation energy is abnormally high. This is quite simply wron g! The reason for thefluctuation is because the group 3 value is lower than you might expect for the reasons we've looked at.

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 isoffsetting it this time?

N 1s22s22px12py12pz1 1st I.E. = 1400 kJ mol -1O 1s22s22px22py12pz1 1st I.E. = 1310 kJ mol -1The 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 2p x2 pair. The repulsion between the two electrons in the sameorbital 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 th e fact that the ionisationenergies in period 3 are all less than those in period 2.

Taking Group 1 as a typical example:


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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 exp ected a much largerionisation energy in sodium, but offsetting the nuclear charge is a greater distance from the nucleus and more screening.

Li 1s22s1 1st I.E. = 519 kJ mol -1Na 1s22s22p63s1 1st I.E. = 494 kJ mol -1Lithium's outer electron is in the second level, and only has the 1s 2 electrons to screen it. The 2s 1 electron feels the pull of 3 protons screened by 2electrons - 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 electronalso 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 extrascreening electrons. The only factor left is the extra distance between the outer electron and the nucleus in sodium's case. That lowers theionisation 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]3d n4s2 (or 4s1 in the cases of chromium and copper). The electron being lost always comesfrom the 4s orbital.

Note: Confusingly, once the orbitals have elec trons in them, the 4s orbital has a higher energy than the 3d - quite the opposite of their order when the atoms are being filled wi th electrons.That means that it is a 4s electron which is lost from the atom when it forms an ion. It also means that the 3d orbitals are slightly closer to the nucleus than the 4s - and so offer some

screening.

You will find this commented on in the page aboutelectronic structures of ions.

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 el ectrons. The3d electrons have some screening effect, and the extra proton and the extra 3d electron more or less cancel each othe r out as far as attraction fromthe centre of the atom is concerned.

The rise at zinc is easy to explain.

Cu [Ar]3d104s1 1st I.E. = 745 kJ mol -1Zn [Ar]3d104s2 1st I.E. = 908 kJ mol -1In 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 theattraction is greater. There will be a degree of repulsion between the paired up electrons in the 4s orbital, but in this cas e it obviously isn't enough tooutweigh the effect of the extra proton.

Note: This is actually very similar to the increase from, say, sodium to magnesium in the third period. In that case, the outer electronic structure is going from 3s1 to 3s2. Despite the pairing-upof the electrons, the ionisation energy increases because of the extra proton in the nucleus. The repulsion between the 3s electrons obviously isn't enough to outweigh this either.

I don't know why the repulsion between the paired electrons matters less for electrons in s orbitals than in p orbitals (I don't even know whether you can make that generalisation!). I suspectthat it has to do with orb ital shape and possibly the greater penetration of s electrons toward s the nucleus, but I haven't been able to find any reference to this anywhere. In fact, I haven't been

able to find anyone who even mentions repulsion in the context of paired s electrons!

If you have any hard information on this, could you contact me via the address on the about this site page.


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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 ionswill 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 solidor in solution. The energy changes in these processes also vary from element to element. Ideally you need to consider the who le picture and not justone 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 ener gy, the faster thereaction 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.

Note: You will find a page discussing this in more detail in the inorganic section of this site dealing with thereactions of Group 2 metals with water.

THE ATOMIC HYDROGEN EMISSION SPECTRUM

This page introduces the atomic hydrogen emission spectrum, showing how it arises from electron movements between energy leve ls within theatom. It also looks at how the spectrum can be used to find the ionisation energy of hydrogen.

What is an emission spectrum?

Observing hydrogen's emission spectrum

A hydrogen discharge tube is a slim tube containing hydrogen gas at low pressure with an electrode at each end. If you put a high voltage acrossthis (say, 5000 volts), the tube lights up with a bright pink glow.

If the light is passed through a prism or diffraction grating, it is split into its various colours. What you would see is a small part of the hydrogenemission spectrum. Most of the spectrum is invisible to the eye because it is either in the infra -red or the ultra-violet.

The photograph shows part of a hydrogen discharge tube on the left, and the three most easily seen lines in the visible part of the spectrum on theright. (Ignore the "smearing" - particularly to the left of the red line. This is c aused by flaws in the way the photograph was taken. See note below.)

Note: This photograph is by courtesy of Dr Rod Nave of the Department of Physics and Astronom y at Georgia State University, Atlant a. The photograph comes from notes about t he hydrogenspectrum in his HyperPhysics pages on the University site. If you are interested in more than an introductory look at the subject, that is a good place togo.


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Ideally the photo would show three clean spectral lines - dark blue, cyan and red. The red smearing which appears to the left of the red line, and other similar smearing (much more difficult tosee) to the left of the other two lines probably comes, according to Dr Nave, from stray reflections in the set-up, or possibly from flaws in the diffraction grating. Ihave chosen to use this

photograph anyway because a) I think it is a stunning image, and b) it is the only one I have ever come across which includesa hydrogen discharge tube and its spectrum in the same image.

Extending hydrogen's emission spectrum into the UV and IR

There is a lot more to the hydrogen spectrum than the three lines you can see with the naked eye. It is possible to detect pa tterns of lines in both theultra-violet and infra-red regions of the spectrum as well.

These fall into a number of "series" of lines named after the person who discovered them. The diagram below shows three of these series, but thereare others in the infra-red to the left of the Paschen series shown in the diagram.

The diagram is quite complicated, so we will look at it a bit at a time. Look first at the Lyman series on the right of the diagram - this is the mostspread out one and easiest to see what is happening.

Note: The frequency scale is marked in PHz - that's petaHertz. You are fam iliar with prefixes like kilo (meaning a thousand or 103 times), and mega (meaning a million or 106 times). Peta means1015 times. So a value like 3 PHz means 3 x 1015 Hz. If you are worried about "Hertz", it just means "cycles per second".

The Lyman series is a series of lines in the ultra-violet. Notice that the lines get closer and closer together as the frequency increases. Eventually,

they get so close together that it becomes impossible to see them as anything other than a continuous spectrum. That's what the shaded bit on theright-hand end of the series suggests.

Then at one particular point, known as the series limit, the series stops.

If you now look at the Balmer series or the Paschen series, you will see that the pattern is just the same, but the series ha ve become more compact.In the Balmer series, notice the position of the three visible lines from the photograph further up the page.

Complicating everything - frequency and wavelength

You will often find the hydrogen spectrum drawn using wavelengths of light rather than frequencies. Unfortunately, because of the mathematicalrelationship between the frequency of light and its wavelength, you get two completely different views of the spectrum if you plot it againstfrequency or against wavelength.

The relationship between frequency and wavelength

The mathematical relationship is:


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Rearranging this gives equations for either wavelength or frequency.

What this means is that there is an inverse relationship between the two - a high frequency means a low wavelength and vice versa.

Note: You will sometimes find frequency given the much more obvious symbol, f.

Drawing the hydrogen spectrum in terms of wavelength

This is what the spectrum looks like if you plot it in terms of wavelength instead of frequency:

. . . and just to remind you what the spectrum in terms of frequency looks like:

Is this confusing? Well, I find it extremely confusing! So what do you do about it?

For the rest of this page I shall only look at the spectrum plotted against frequency, because it is much easier to relate it to what is happening in theatom. Be aware that the spectrum looks different depending on how it is plotted, but, other than that, ignore the wavelength version unless it isobvious that your examiners want it. If you try to learn both versions, you are only going to get them muddled up!

Note: Syllabuses probably won't be very helpful about this. You need to look at past papers and mark schemes.

If you are working towards a UK-based exam and don't have these things, you can find out how to get hold of them by going to the syllabuses page.

Explaining hydrogen's emission spectrum

The Balmer and Rydberg Equations


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Tyingparticularelectronjumpstoindividuallinesinthespectrum

Ifanelectronfallsfromthe3-leveltothe2-level, ithastoloseanamountofenergyexactlythesameastheenergygapbetweenthosetwolevels.Thatenergywhichtheelectronlosescomesoutaslight(where"light"includesUV andIR aswellasvisible).

Eachfrequencyoflightisassociatedwithaparticularenergybytheequation:

Thehigherthefrequency, thehighertheenergyofthelight.

Ifanelectronfallsfromthe3-leveltothe2-level, redlightisseen. Thisistheoriginoftheredlineinthehydrogenspectrum. Bymeasuringthefrequencyoftheredlight,youcanwork outitsenergy. Thatenergymustbeexactlythesameastheenergygapbetweenthe3-levelandthe2-levelinthehydrogenatom.

Thelastequationcanthereforebere-writtenasameasureoftheenergygapbetweentwoelectronlevels.

Thegreatestpossiblefallinenergywillthereforeproducethehighestfrequencylineinthespectrum. Thegreatestfallwillbefromtheinfinity leveltothe1-level. (Thesignificanceoftheinfinitylevelwillbemadeclearlater.)


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The next few diagrams are in two parts - with the energy levels at the top and the spectrum at the bottom.

If an electron fell from the 6 -level, the fall is a little bit less, and so the frequency will be a little bit lower. (Because of the scale of the diagram, it isimpossible to draw in all the jumps involving all the levels between 7 and infinity!)

. . . and as you work your way through the other possible jumps to the 1 -level, you have accounted for the whole of the Lyman series. The spacingsbetween the lines in the spectrum reflect the way the spacings between the energy levels change.


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If you do the same thing for jumps down to the 2-level, you end up with the lines in the Balmer series. These energy gaps are all much smaller thanin the Lyman series, and so the frequencies produced are also much lower.

The Paschen series would be produced by jumps down to the 3-level, but the diagram is going to get very messy if I include those as well - not tomention all the other series with jumps down to the 4-level, the 5-level and so on.

The significance of the numbers in the Rydberg equation

n1 and n2 in the Rydberg equation are simply the energy levels at either end of the jump producing a particular line in the spectrum.


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For example, in the Lyman series, n1 is always 1. Electrons are falling to the 1 -level to produce lines in the Lyman series. For the Balmer series, n 1 isalways 2, because electrons are falling to the 2-level.

n2 is the level being jumped from. We have already mentioned that the red line is produced by electrons falling from the 3-level to the 2-level. In thiscase, then, n2 is equal to 3.

The significance of the infinity level

The infinity level represents the highest possible energy an electron can have as a part of a hydrogen atom. So what happens if the electron exceedsthat energy by even the tiniest bit?

The electron is no longer a part of the atom. The infinity level represents the point at which ionisation of the atom occurs to form a positivelycharged ion.

Using the spectrum to find hydrogen's ionisation energy

When there is no additional energy supplied to it, hydrogen's electron is found at the 1 -level. This is known as its ground state. If you supply enoughenergy to move the electron up to the infinity level, you have ionised the hydrogen.

The ionisation energy per electron is therefore a measure of the distance between the 1-level and the infinity level. If you look back at the last fewdiagrams, you will find that that particular energy jump produces the series limit of the Lyman series.

Note: Up to now we have been talking about the energy released when an electron falls from a higher to a lower level. Obviously ifa certain amount of energy is released when an electronfalls from the inf inity level to the 1-level, that same amount wil l be needed to push the electron from the 1-level up to the infin ity level.

If you can determine the frequency of the Lyman series limit, you can use it to calculate the energy needed to move the elect ron in one atom from the1-level to the point of ionisation. From that, you can calculate the ionisation energy per mole of atoms.

The problem is that the frequency of a series limit is quite difficult to find accurately from a spectrum because the lines a re so close together in thatregion that the spectrum looks continuous.

Finding the frequency of the series limit graphically

Here is a list of the frequencies of the seven most widely spaced lines in the Lyman series, together with the increase in fr equency as you go fromone to the next.

As the lines get closer together, obviously the increase in frequency gets less. At the series limit, the gap between the lines would be literally zero.


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That means that if you were to plot the increases in frequency against the actual frequency, you could extrapolate (contin ue) the curve to the point atwhich the increase becomes zero. That would be the frequency of the series limit.

In fact you can actually plot two graphs from the data in the table above. The frequency difference is related to two frequencies. For example, thefigure of 0.457 is found by taking 2.467 away from 2.924. So which of these two values should you plot the 0.457 against?

It doesn't matter, as long as you are always consistent - in other words, as long as you always plot the difference against either the higher or thelower figure. At the point you are interested in (where the difference becomes zero), the two frequency numbers are the same.

As you will see from the graph below, by plotting both of the possible curves on the same graph, it makes it ea sier to decide exactly how toextrapolate the curves. Because these are curves, they are much more difficult to extrapolate than if they were straight line s.

Both lines point to a series limit at about 3.28 x 10 15 Hz.

Note: Remember that 3.28 PHz is the same as 3.28 x 1015 Hz. You can use the Rydberg equation to calculate the series limit of the Lyman series as a check on this figure: n1 = 1 for the Lyman

series, and n2 = infinity for the series limit. 1/(infinity)2

= zero. That gives a value for the frequency of 3.29 x 1015

Hz - in other words the two values agree to within 0.3%.

So . . . now we can calculate the energy needed to remove a single electron from a hydrogen atom. Remember the equation from higher up the page:

We can work out the energy gap between the ground state and the point at which the electron leaves the atom by substituting the value we've got f orfrequency and looking up the value of Planck's constant from a data book.


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That gives you the ionisation energy for a single atom. To find the normally quoted ionisation energy, we need to multiply this by the number ofatoms in a mole of hydrogen atoms (the Avogadro constant) and then divide by 1000 to convert it into kilojoules.

Note: It would be wrong to quote this to m ore than 3 significant figur es. The value for the frequency obtained from the graph is only to that accuracy.

This compares well with the normally quoted value for hydrogen's ionisation energy of 1312 kJ mol -1.

ELECTRON AFFINITY

This page explains what electron affinity is, and then looks at the factors that affect its size. It assumes that you know ab out simple atomic orbitals,and can write electronic structures for simple atoms.

Important! If you aren't reasonable happy about atomic orbitals and electronic structures you should follow these links before you go any further.

First electron affinity

Ionisation energies are always concerned with the formation of positive ions. Electron affinities are the negative ion equiva lent, and their use isalmost 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 ga seous 1- ions.


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 signshows a release of energy.

The first electron affinities of the group 7 elements

F -328 kJ mol-1Cl -349 kJ mol-1Br -324 kJ mol-1I -295 kJ mol-1

Note: These values are based on the most recent research. If you are using a different data source, you may have slightly different numb ers. That doesn't matter - the pattern will still be the

same.


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Isthereapattern?

Yes- asyou go downthegroup, firstelectronaffinitiesbecomeless(inthesensethatlessenergy isevolved whenthenegativei onsareformed).Fluorinebreaksthatpattern, and will haveto beaccounted forseparately.

Theelectronaffinity isameasureoftheattractionbetweentheincoming electronand thenucleus - thestrongertheattraction, themoreenergy isreleased.

Thefactorswhichaffectthisattractionareexactly thesameasthoserelating to ionisationenergies - nuclearcharge, distanceand screening.

Note: If you haven'tread aboutionisationenergyrecently, itmightbeagood ideato follow thislink beforeyou go on. Thesefactorsarediscussed inmoredetail onthatpagethanthey areon

thisone.

Theincreased nuclearchargeasyou go downthegroupisoffsetby extrascreening electrons. Eachouter electronineffectfeelsapull of 7+ fromthecentreoftheatom, irrespectiveof whichelementyou aretalking about.

Forexample, afluorineatom hasanelectronic structureof 1s22s22px22py22pz1. Ithas9 protonsinthenucleus.

Theincoming electronentersthe2-level, and isscreened from thenucleusby thetwo 1s2electrons. Itthereforefeelsanetattractionfrom thenucleusof 7+ (9 protonslessthe2 screening electrons).

By contrast, chlorinehastheelectronic structure1s22s22p63s23px23py23pz1. Ithas17 protonsinthenucleus.

Butagaintheincoming electronfeelsanetattractionfrom thenucleusof 7+ (17 protonslessthe10 screening electronsin thefirstand secondlevels).

Note: If you wantto be fussy, thereisalso asmall amountofscreening by the2selectronsin fluorineand by the3selectronsinchlorine. Thiswill beapproximately thesameinbo ththesecasesand so doesn'taffecttheargumentinany way (apart from complicating it!).

Theover-riding factoristhereforethe increased distancethattheincoming electronfindsitself from thenucleusasyou go downthegroup. Thegreaterthedistance, thelesstheattractionand so thelessenergy isreleased aselectronaffinity.

Note: Comparing fluorineand chlorine isn'tideal, becausefluorinebreaksthetrend inthegroup. However, comparing chlorineand bromine, say, makesthingsseem moredifficultbecauseofthemorecomplicated electronic structuresinvolved.

Whatwehavesaid so farisperfectly trueand appliesto the fluorine-chlorinecaseasmuchasto anything elseinthegroup, butthere'sanotherfactorwhichoperatesaswell whichwehaven'tconsidered yet - and thatover-ridestheeffectof distanceinthecaseof fluorine.

Why isfluorineoutof line?

Theincoming electronisgoing to becloserto thenucleusinfluorinethaninany otheroftheseelements, so you would expectahighvalueofelectronaffinity.

However, becausefluorineissuchasmall atom, you areputting thenew electroninto aregiono fspacealready crowded withelectronsand thereisasignificantamountofrepulsion. Thisrepulsionlessenstheattractiontheincoming electronfeelsand so lessenstheelect ronaffinity.

A similarreversal oftheexpected trend happensbetweenoxygenand sulphurinGroup6. Thefirstelectronaffinity of oxygen (-142 kJ mol-1) issmallerthanthatofsulphur(-200 kJ mol-1) forexactly thesamereasonthatfluorine'sissmallerthanchlori ne'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 kJmol-1) is less than chlorine's (-349 kJ mol-1). Why?


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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 inboth.

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 theincoming electron is held less strongly as you go down the group and so the negative ion is less likely to form. That explana tion looks reasonableuntil 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 exp lain a trend in terms ofjust one of those steps. Fluorine is much more reactive than chlorine (despite the lower electron affini ty) because the energy released in other stepsin 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.


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!

1st EA = -142 kJ mol-12nd 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 part icularly high becausethe electron is being forced into a small, very electron-dense space.

ATOMIC AND IONIC RADIUS

This page explains the various measures of atomic radius, and then looks at the way it varies around the Periodic Table - across periods and downgroups. It assumes that you understand electronic structures for simple atoms written in s, p, d notation.

Important! If you aren't reasonable happy about electronic structures you should follow this link before you go any further.

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 twotouching atoms, and then halving that distance.


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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 th ey 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 beingmeasured 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 jus t touching. The attractive forces are much less, and the atoms are essentially"unsquashed". This measure of atomic radius is called the van derWaals radius after the weak attractions present in this situation.

Note: If you want to explore these various types of bonding this link will take you to the bonding menu.

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 va n derWaals radii forthose (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 the ir van derWaalsradius - a case where the atom is pretty well "unsquashed". All the other atoms are being measured where their ato mic radius is being lessened bystrong 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 thebond. (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 waybetween 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 nucleusas you go across the period pulls the electrons in more tightly. The amount of screening is constant for all of these elements.

Note: You might possibly wonder why you don't get extra screening from the 2s2 electrons in the cases of the elements from boron to fluorine where the bonding involves the p electrons.

In each of these cases, before bonding happens, the existing s and p orbitals are reorganised (hybridised) into new orbitals of equal energy. When these atoms are bonded, there aren't any 2selectrons as such.

If you don't know about hybridisation, just ignore this comment - you won't need it for UK A level purposes anyway.


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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 pulledcloser 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 screeningdue to the increasing number of 3d electrons.

Note: Confusingly, once the orbitals have elec trons in them, the 4s orbital has a higher energy than the 3d - quite the opposite of their order when the atoms are being filled wi th electrons.That means that i t is the 4s electron s which can be thought o f as being on the ou tside of the atom , and so determine i ts size. It also means that the 3d orbitals are slightly closer to the nucleusthan the 4s - and so offer some screening.

You will find this commented on in the page aboutelectronic structures of ions.

IONIC RADIUS

A warning!

Ionic radii are difficult to measure with any degree of certainty, and vary according to the environment of the ion. For exam ple, it matters what the co-ordination of the ion is (how many oppositely charged ions are touching it), and what those ions are.

There are several different measures of ionic radii in use, and these all differ from each other by varying amounts. It means that if you are going tomake reliable comparisons using ionic radii, they have to come from the same source.

What you have to remember is that there are quite big uncertainties in the use of ionic radii, and that trying to explain thi ngs in fine detail is madedifficult by those uncertainties. What follows will be adequate for UK A level (and its various equiv alents), but detailed explanations are toocomplicated for this level.

Trends in ionic radius in the Periodic Table

Trends in ionic radius down a group

This is the easy bit! As you add extra layers of electrons as you go down a group, the ions are bound to get bigger. The two tables below show thiseffect in Groups 1 and 7.

electronic structureof ion ionic radius (nm)

Li+ 2 0.076Na+ 2, 8 0.102K+ 2, 8, 8 0.138Rb+ 2, 8, 18, 8 0.152


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Cs+ 2, 8, 18, 18, 8 0.167

electronic structureof ion ionic radius (nm)

F- 2, 8 0.133

Cl

-2, 8, 8

0.181

Br- 2, 8, 18, 8 0.196I- 2, 8, 18, 18, 8 0.220

Note: These figures all come from theDatabase of Ionic Radii from Imperial College London. I have converted them from Angstroms to nm (nanometres), which are more often used in thedata tables that you are likely to come across.

If you are interested, 1 Angstrom is 10-10 m; 1 nm = 10-9 m. To convert from Angstroms to nm, you have to divide by 10, so that 1.02 Angstroms becomes 0.102 nm. You may also come acrosstables listing values in pm (picometres) which are 10-12 m. A value in pm will look like, for example, for chlorine, 181 pm rather than 0.181 nm. Don't worryif you find this confusing. Just use

the values you are given in whatever units you are given.

For comparison purposes, all the values relate to 6-co-ordinated ions (the same arr angement as in NaCl, for example). CsC l actually crystallises in an 8:8-co-ordinated structure - so youcouldn't accurately use these values for CsCl. The 8-co-ordinated ionic radius for Cs is 0.174 nm rather than 0.167 for the 6-co-ordinated version.

Trendsinionicradiusacrossaperiod

Let'slook at theradiiof thesimple ionsformedby elementsasyougoacrossPeriod3 of thePeriodicTable - theelementsfrom NatoCl.

Na+ Mg2+ Al3+ P3- S2- Cl-no of protons 11 12 13 15 16 17electronic structure of ion 2,8 2,8 2,8 2,8,8 2,8,8 2,8,8ionic radius (nm) 0.102 0.072 0.054 (0.212) 0.184 0.181

Note: The table misses out sil icon which doesn't form a simple ion . The phosphideion radius is in brackets because it comes from a diff erent data source, and I am not sure whether it is safe

to compare it. The values for the oxide and chloride ions agree in the different source, so it is probably OK. The values areagain for 6-co-ordination, although I can't guarantee that for thephosphide figure.

First of all, notice the big jump in ionic radius as soon as you get into the negative ions. Is this surprising? Not at all - you have just added a wholeextra layer of electrons.

Notice that, within the series of positive ions, and the series of negative ions, that the ionic radii fall as you go across the per iod. We need to look atthe positive and negative ions separately.

The positive ions

In each case, the ions have exactly the same electronic structure - they are said to be isoelectronic. However, the number of protons in the nucleusof the ions is increasing. That will tend to pull the electrons more and more towards the centre of the ion - causing the ionic radii to fall. That is prettyobvious!

The negative ions

Exactly the same thing is happening here, except that you have an extra layer of electrons. What needs commenting on, though is how similar in sizethe sulphide ion and the chloride ion are. The additional proton here is making hardly any difference.


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The difference between the size of similar pairs of ions actually gets even smaller as you go down Groups 6 and 7. For exampl e, the Te2- ion is only0.001 nm bigger than the I - ion.

As far as I am aware there is no simple explanation for this - certainly not one which can be used at this level. This is a good illustration of what Isaid earlier - explaining things involving ionic radii in detail is sometimes very difficult.

Trends in ionic radius for some more isoelectronic ions

This is only really a variation on what we have just been talking about, but fits negative and positive isoelectronic ions in to the same series ofresults. Remember that isoelectronic ions all have exactly the same electron arrangement.

N3- O2- F- Na+ Mg2+ Al3+no of protons 7 8 9 11 12 13electronic structure of ion 2, 8 2, 8 2, 8 2, 8 2, 8 2, 8ionic radius (nm) (0.171) 0.140 0.133 0.102 0.072 0.054

Note: The nitride ion value is in brackets because it came from a different source, and I don't know for certain whether i t relatesto the same 6-co-ordination as the rest of the ions. This

matters. My main source only gave a 4-co-ordinated value for the nitride ion, and that was 0.146 nm.

You might also be curious as to how the neutral neon atom fits into this sequence. It would seem logical that its van derWaals radius would fall neatly between that of the fluoride ion and the

sodium ion. It doesn't! Its radius is 0.154 or 0.160 nm (depending on which source you look the value up in) - bigger than the fluoride ion. I have no idea why that is!

You can see that as the number of protons in the nucleus of the ion increases, the electrons get pulled in more closely to the nucleus. T he radii ofthe isoelectronic ions therefore fall across this series.

The relative sizes of ions and atoms

You probably won't have noticed, but nowhere in what you have read so far has there been any need to talk about the relative sizes of the ions andthe atoms they have come from. Neither (as far as I can tell from the syllabuses) do any of the current UK -based exams for 16 - 18 year olds ask forthis specifically in their syllabuses.

However, it is very common to find statements about the relative sizes of ions and atoms. I am fairly convinced that these st atements are faulty, and Iwould like to attack the problem head-on rather than just ignoring it.

Important!

For 10 years, until I rewrote this ionic radius section in August 2010, I included what is in the box below. You will find th is same information andexplanation in all sorts of books and on any number of websites aimed at this level. At least one non-UK A level syllabus has a statement whichspecifically asks for this.

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 10electrons are being pulled in by the full force of 11 protons.


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Negativeions

Negativeionsarebiggerthantheatomsthey comefrom. Chlorineis2,8,7; Cl -is2,8,8. Although theelectronsarestill all inthe3 -level, theextrarepulsionproduced by theincomingelectroncausestheatom toexpand. Therearestill only 17 protons, butthey arenow havi ngtohold 18electrons.

However, I was challenged by an experienced teacher about the negative ion explanation, and that forced me to think about itcarefully for the first time. I am now convinced that the facts and the explanation relating to negative ions are simply illogical.

As far as I can tell, no UK-based syllabus mentions the relative sizes of atoms and ions (as of August 2010), but you should checkpast papers and mark schemes to see whether questions have sneaked in.

The rest of this page discusses the problems that I can see, and is really aimed at teachers and others, rather than atstudents.

If you are a student, look carefully at your syllabus, and past exam questions and mark schemes, to find out whether you need toknow about this. If you don't need to know about it, stop reading now (unless, of course, you are interested in a bit of controversy!).

If you do need to know it, then you will have to learn what is in the box, even if, as I believe, it is wrong. If you like your chemistry tobe simple, ignore the rest of the page, because you risk getting confused about what you need to know.

If you have expert knowledge of this topic, and can find any flaws in what I am saying, then please contact me via the addresson the about this site page.

Choosingtherightatomicradiustocomparewith

Thisisattheheartoftheproblem.

Thediagramsinthebox above, andsimilaronesthatyouwill findelsewhere, usethemetallicradiusasthemea sureofatomicradiusformetals, andthecovalentradiusfornon-metals. I wanttofocusonthenon-metals, becausethatiswherethemainproblemlies.

Youare, ofcourse, perfectly freetocomparetheradiusofanionwithwhatevermeasureofatomicrad iusyouchoose. Theproblemcomesinrelatingyourchoiceofatomicradiustothe"explanation"ofthedifferences.

Itisperfectly truethatnegativeionshaveradiiwhicharesignificantly biggerthanthecovalentradiusoftheatominque stion. Andtheargument

thengoesthatthereasonforthisisthatif youaddoneormoreextraelectronstotheatom, inter-electronrepulsionscausetheatomtoexpand.Thereforethenegativeionisbiggerthantheatom.

Thisseemstometobecompletely inconsistent. If youaddoneormoreextraelectronstotheatom, youaren'taddingthemto acovalently boundatom. Youcan'tsimply addelectronstoacovalently-boundchlorineatom, forexample- chlorine'sexistingelectronshavereorganisedthemselvesintonewmolecularorbitalswhichbindtheatomstogether.

Inacovalently-boundatom, thereissimply noroomtoaddextraelectrons.

Soif youwanttousetheelectronrepulsionexplanation, theimplicationisthatyou areaddingtheextraelectronstoarawatomwithasimpleuncombinedelectronarrangement.

Inotherwords, if youweretalkingabout, say, chlorine, youareaddinganextraelectrontochlorinewithaconfigurationo f 2,8,7 - nottocovalentlyboundchlorineatomsinwhichthearrangementoftheelectronshasbeenalteredby sharing.

Thatmeansthatthecomparisonthatyououghttobemakingisn'twiththeshortenedcovalentradius, butwiththemuchlarger vanderWaalsradius- theonly availablemeasureoftheradiusofanuncombinedatom.

Sowhathappensif youmakethatcomparison?

Group7

vdW radius (nm) ionic radius of X - (nm)


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F 0.147 0.133Cl 0.175 0.181Br 0.185 0.196I 0.198 0.220

Group6

vdW radius (nm) ionic radius of X2- (nm)O 0.152 0.140S 0.180 0.184Se 0.190 0.198Te 0.206 0.221

Group5

vdW radius (nm) ionic radius of X3- (nm)N 0.155 0.171P 0.180 0.212

As we have already discussed above, measurements of ionic radii are full of uncertainties. That is also true of van der Waals radii. The table usesone particular set of values for comparison purposes. If you use data from different sources, you will find differences in the patterns - includingwhich of the species (ion or atom) is bigger.

These ionic radius values are for 6-co-ordinated ions (with a slight question mark over the nitride and phosphide ion figures). But you may rememberthat I said that ionic radius changes with co-ordination. Nitrogen is a particularly good example of this.

4-co-ordinated nitride ions have a radius of 0.146 nm. In other words if you look at one of the co-ordinations, the nitride ion is bigger than thenitrogen atom; in the other case, it is smaller. Making a general statement that nitride ions are bigger or smaller than nitrogen atoms is impossibl e.

So what is it safe to say about the facts?

For most, but not all, negative ions, the radius of the ion is bigger than that of t he atom, but the difference is nothing like as great as is shown if youincorrectly compare ionic radii with covalent radii. There are also important exceptions.

I can't see how you can make any real generalisations about this, given the uncertainties in the data.

And what is it safe to say about the explanation?

If there are any additional electron-electron repulsions on adding extra electrons, they must be fairly small. This is particularly shown if you considersome pairs of isoelectronic ions.


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You would have thought that if repulsion was an important factor, then the radius of, say a sulphide ion, with two negative c harges would besignificantly larger than a chloride ion with only one. The difference should actu ally be even more marked, because the sulphide electrons are beingheld by only 16 protons rather than the 17 in the chlorine case.

On this repulsion theory, the sulphide ion shouldn't just be a little bit bigger than a chloride ion - it should be a lot bigger. The same effect is shownwith selenide and bromide, and with telluride and iodide ions. In the last case, there is virtually no difference in the size s of the 2- and 1- ions.

So if there is some repulsion playing a part in this, it certainly doesn't look as if it is playing a major part.

What about positive ions?

Whether you choose to use van derWaals radii or metallic radii as a measure of the atomic radius, for metals the ionic radius is smaller than either,so the problem doesn't exist to the same extent. It is true that the ionic radius of a metal is less than its atomic radius ( however vague you are aboutdefining this).

The explanation (at least as long as you only consider positive ions from Groups 1, 2 and 3) in terms of losing a complete la yer of electrons is alsoacceptable.

Conclusion

It seems to me that, for negative ions, it is compl etely illogical to compare ionic radii with covalent radii if you want to use the electron repulsionexplanation.

If you compare the ionic radii of negative ions with the van derWaals radii of the atoms they come from, the uncertainties i n the data make it verydifficult to make any reliable generalisations.

The similarity in sizes of pairs of isoelectronic ions from Groups 6 and 7 calls into question how important repulsion is in any explanation.

Having spent more than a week working on this, and discussing

atomic structure and bonding menu

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