Its van der Waals radius is 0.154 or 0.160 nm (depending on which source you look the value up in) - bigger than the fluoride ion. You might also be curious as to how the neutral neon atom fits into this sequence. My main source only gave a 4-coordinated value for the nitride ion, and that was 0.146 nm. Note: The nitride ion value is in brackets because it came from a different source, and I don't know for certain whether it relates to the same 6-co-ordination as the rest of the ions. Remember that isoelectronic ions all have exactly the same electron arrangement.
#Periodic table cl series#
This is only really a variation on what we have just been talking about, but fits negative and positive isoelectronic ions into the same series of results. Trends in ionic radius for some more isoelectronic ions This is a good illustration of what I said earlier - explaining things involving ionic radii in detail is sometimes very difficult. For example, the Te 2 - ion is only 0.001 nm bigger than the I - ion.Īs far as I am aware there is no simple explanation for this - certainly not one which can be used at this level. The difference between the size of similar pairs of ions actually gets even smaller as you go down Groups 6 and 7. The additional proton here is making hardly any difference.
What needs commenting on, though is how similar in size the sulphide ion and the chloride ion are.
The negative ions: Exactly the same thing is happening here, except that you have an extra layer of electrons. That will tend to pull the electrons more and more towards the center of the ion - causing the ionic radii to fall. However, the number of protons in the nucleus of the ions is increasing. The positive ions: In each case, the ions have exactly the same electronic structure - they are said to be isoelectronic. We need to look at the positive and negative ions separately. Notice that, within the series of positive ions, and the series of negative ions, that the ionic radii fall as you go across the period. Is this surprising? Not at all - you have just added a whole extra layer of electrons. The values are again for 6-co-ordination, although I can't guarantee that for the phosphide figure.įirst of all, notice the big jump in ionic radius as soon as you get into the negative ions. The values for the oxide and chloride ions agree in the different source, so it is probably OK. The phosphide ion radius is in brackets because it comes from a different data source, and I am not sure whether it is safe to compare it. The table misses out silicon which does not form a simple ion. Let's look at the radii of the simple ions formed by elements as you go across Period 3 of the Periodic Table - the elements from Na to Cl. 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. In the period from sodium to chlorine, the same thing happens. The amount of screening is constant for all of these elements. The increasing number of protons in the nucleus as you go across the period pulls the electrons in more tightly. (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.)įrom lithium to fluorine, those electrons are all in the 2-level, being screened by the 1s 2 electrons. 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. Leaving the noble gases out, atoms get smaller as you go across a period. You aren't comparing like with like if you include the noble gases. All the other atoms are being measured where their atomic radius is being lessened by strong attractions. 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". You have to ignore the noble gas at the end of each period.
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