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Holding elements together: Ionic bonds

The views expressed are those of the author and are not necessarily those of Scientific American.


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A while ago, I wrote a couple of posts describing some intra-molecular forces, forces that hold atoms and molecules together. I enjoyed writing them, and people come back to read them quite frequently, so I thought I’d continue and write about a couple more.

The previous posts covered van der Waals forces and hydrogen bonds (and dipoles!) These forces are both fun, and incredibly important in determining the properties of water (without which there would be no life) and petrol (without which there would be no cars) but they aren’t particularly strong forces. Both van der Waals and hydrogen bonds are relatively weak, and I think it’s about time I started covering the three Big Bads of molecular forces: ionic bonds, covalent bonds and metallic bonds.

Starting with ionic bonds. In order to properly cover these, I need just a little bit of background on what atoms are actually like. Now there are many different models for what atoms are ‘like’ but for the purposes of ionic bonds the easiest way to think of them is as a small nucleus with a positive charge surrounded by small whizzing negatively charged electrons. Electrons don’t just all fly around with no direction, they exist in orbitals, each a different distance from the nucleus. To find out how many electrons are in the outermost orbital, furthest away from the nucleus, you can look at the periodic table.

The periodic table. The number of electrons in the outer shell can be found by looking at which column the element is in. This does get a bit complicated for some of the larger elements, but for the first few rows it works fine.

The outermost orbital is the most important because all the other electrons are a) in orbitals with the maximum number of electrons and b) unable to take much part in reactions due to the outermost electrons being in the way. Atoms in general are at their most stable when they have a full outer shell of electrons. As many atoms don’t have a full outer shell of electrons, they try to find ways of getting one.

Take sodium, for example, with the chemical symbol Na (blame the Greeks for that!) Now sodium is in the first column of the periodic table, so it has one electron in the outermost shell. All the other shells are full of electrons. So the quickest way to get a full outer shell is to loose that outermost electron, leaving you with a full shell underneath. Loosing one electron, however, requires energy, which means that unless this can be combined with a process which produces energy, it is not going to happen.

Loosing one electron also means that rather than having an equal number of positive nuclear charges and negative electron charges. the sodium now has one fewer negative charges. Instead of a neutral sodium atom, it is now a positively charged sodium ion.

So how does the sodium get the energy back? One way of producing a lot of energy in the world of molecular forces is to form a bond. Breaking bonds requires energy, but making them produces energy (ask a physicist about this if you’re unsure why, the explanation might not make you any more sure but it’s nice for physicists to have someone to talk to). Luckily, a positively charged ion is very good at forming bonds with a negatively charged ion.

To see how a negative ion forms, take a look at the other side of the periodic table at the element chlorine. Chlorine has seven electrons in its outer shell (a full shell is eight electrons for the smaller elements). In order to fill its outer shell it can take an electron gaining one extra negative charge and forming the negative chloride ion. It doesn’t just take that electron from anywhere, it will actually tear it away from the sodium atom. Leaving one negative and one positive ion; Na+ and Cl-

ionic bond electron transfer

A simplified diagram of the electron transfer, with outer shell electrons shown. Image in the public domain, credit below

The bond formed between these ions is called an ionic bond.

The two ions don’t just form a single bond between them, ideally they want to make as many bonds as possible. Ionic bonds create giant ionic lattices of interchanging ions. These are incredibly strong, making ionic compounds very hard to break, melt, or crush.

Not only do ionic forces help to hold elements together, the formation of these bonds can completely change the properties of the elements they bind. Sodium is a soft grey metal (you can cut it with a knife!). Chlorine is a deadly green gas. The giant ionic lattice sodium chloride is salt, a white grainy edible food.

Sodium metal and chlorine gas in elemental form, and magnified crystals of sodium chloride. Credit for all images below.

Credit for image 2

Credit for images of chlorine, sodium and magnified salt crystals used in image 3

S.E. Gould About the Author: A biochemist with a love of microbiology, the Lab Rat enjoys exploring, reading about and writing about bacteria. Having finally managed to tear herself away from university, she now works for a small company in Cambridge where she turns data into manageable words and awesome graphs. Follow on Twitter @labratting.

The views expressed are those of the author and are not necessarily those of Scientific American.





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  1. 1. Lance Gritton 6:43 pm 02/19/2012

    I tell my students (in a proper Bristol accent) Bond…Ionic Bond, charged not stirred…

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  2. 2. sciencegem 8:51 pm 02/20/2012

    Thanks! Awesome article! Just wondering, I don’t know any physicists to ask, would anyone mind explaining (if they have a moment) why atoms gain energy forming bonds, not lose it? I know it should probably be obvious, but I can’t seem to figure it out.

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  3. 3. S.E. Gould in reply to S.E. Gould 7:13 am 02/22/2012

    Well I stopped worrying about physics when I was about 16, but I’ll give it a go! One way to think about it is the energy state of the atom. Using the example of covalent bonds (where rather than stealing electrons, two atoms will share them) 14 electrons circulating 2 Chlorine nuclei are in a lower energy state than 7 electrons each circling one nucleus each. To get from a higher energy state (Cl + Cl) to a lower energy state (Cl-Cl) releases energy.

    Atoms generally are more stable when bonded, so forming the bond will release energy. Breaking it, on the other hand, requires an energy input because you have to force a molecule from a stable state (Cl-Cl) into a less stable state (Cl + Cl). I wouldn’t write that as an exam answer, but it’s a reasonable way to think about things.

    I’ll do covalent bonds next time, probably some time next month!

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