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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 down groups. It assumes that you understand electronic structures for simple atoms written in s, p, d notation. | ||
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Important! If you aren't reasonable happy about electronic structures you should follow this link before you go any further. | ||
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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. | ||
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Note: If you want to explore these various types of bonding this link will take you to the bonding menu. | ||
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Trends in atomic radius in the Periodic Table The exact pattern you get depends on which measure of atomic radius you use - but the trends are still valid. The following diagram uses metallic radii for metallic elements, covalent radii for elements that form covalent bonds, and van der Waals radii for those (like the noble gases) which don't form bonds. Trends in atomic radius in Periods 2 and 3
Trends in atomic radius down a group It is fairly obvious that the atoms get bigger as you go down groups. The reason is equally obvious - you are adding extra layers of electrons. Trends in atomic radius across periods You have to ignore the noble gas at the end of each period. Because neon and argon don't form bonds, you can only measure their van der Waals radius - a case where the atom is pretty well "unsquashed". All the other atoms are being measured where their atomic radius is being lessened by strong attractions. You aren't comparing like with like if you include the noble gases.
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. | ||
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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 2s electrons 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 pulled closer to the nucleus by increasing numbers of protons - in each case, screened by the 1- and 2-level electrons. Trends in the transition elements
Although there is a slight contraction at the beginning of the series, the atoms are all much the same size. The size is determined by the 4s electrons. The pull of the increasing number of protons in the nucleus is more or less offset by the extra screening due to the increasing number of 3d electrons. | ||
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Note: Confusingly, once the orbitals have electrons in them, the 4s orbital has a higher energy than the 3d - quite the opposite of their order when the atoms are being filled with electrons. That means that it is the 4s electrons which can be thought of as being on the outside of the atom, and so determine its size. 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 about electronic structures of ions. | ||
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IONIC RADIUS Ions aren't the same size as the atoms they come from. Compare the sizes of sodium and chloride ions with the sizes of sodium and chlorine atoms.
Positive ions Positive ions are smaller than the atoms they come from. Sodium is 2,8,1; Na+ is 2,8. You've lost a whole layer of electrons, and the remaining 10 electrons are being pulled in by the full force of 11 protons. Negative ions Negative ions are bigger than the atoms they come from. Chlorine is 2,8,7; Cl- is 2,8,8. Although the electrons are still all in the 3-level, the extra repulsion produced by the incoming electron causes the atom to expand. There are still only 17 protons, but they are now having to hold 18 electrons.
 © Jim Clark 2000 (modified 2004) |
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