3 Periodic Properties 2009

PERIODIC PROPERTIES: A RATIONALE FOR CHEMICAL BEHAVIOUR 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+. (A mole is 6,02x10-23 particles. We will study this somewhere else) 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 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 primarily by:  1-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.  2-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. In the same level, p electrons are slightly farther from the nucleus than the s electrons.  3-The number of electrons between the outer electrons and the nucleus (screening effect). Consider a sodium atom, with the electronic structure 2,8,1. If the outer electron looks in towards the nucleus, it doesn't see the nucleus sharply. Between it and the nucleus

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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. In the same level, p electrons are slightly shielded by the s pair of the same level. 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 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

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Explaining the general trend across periods 2 and 3 The general trend is for ionisation energies to increase across a period. (Factors 2 and 3 keep practically unchanged and factor 1 dominates, however, notice slight drops between groups 2 and 3 and between groups 5 and 6). In period 3, the trend is exactly the same. This time, all the electrons being removed are in the third level (longer distance to nucleus makes it easier) and are screened by the 1s22s22p6 electrons. They all have the same sort of environment, but there is an increasing nuclear charge. 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? In this case factor 1 would make things easier but factors 2 and 3 offset this effect and cause a lower ionisation energy Similar explanations hold as you go down the rest of this group - or, indeed, any other group. 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. 3

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 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. 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 Cl Br I

-328 kJ mol-1 -349 kJ mol-1 -324 kJ mol-1 -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. 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 released as electron affinity. 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 4

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 a first glance 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. 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

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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. 6

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 holding 18 electrons.

QUESTIONS AND PROBLEMS 1- Which three factors influence electron affinity? 2- How and why do electron affinity and ionisation potential vary: a- along a group? b- along a period? 3- Explain why the ionisation energy in group 3 is slightly greater than in group 2. 4- Why it is important to remark that the ionisation energy refers to atoms and ions in the gaseous state? 5- The following table shows the energy required to strip one by one all the 11 electrons of a sodium atom Ionization 1st energy number Enthalpy 49 /kJ mol-1 6

2nd

3rd

4th

5th

6th

7th

8th

9th

10th

11th

456 0

691 0

954 0

1335 0

1661 0

2012 0

2550 0

2890 0

14140 0

15900

a- Plot energy against number of electron stripped off. b- Which electron is more tightly bound to the nucleus? 7

c- Which electrons are closer to the nucleus? d- Which electron is most easily lost? e- Compare all values to the first one (it can be easier for you to divide the values all into 496 (first one) and then compare your results. f- What do these values suggest about the electron distribution in sodium? 6- The following diagram represents the ions in NaCl crystal. Do the darker spheres represent Na+ or Cl-? Explain your answer.

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