Structure and Bonding in Chemistry

Ionic Bonds


Elements in the first few columns of the periodic table have a few more electrons than predicted by the octet rule: they therefore lose the electrons in the outermost shells fairly easily. For example, the alkali metals (group I), such as sodium (Na) or potassium (K), which have, respectively, 11 and 19 electrons, easily lose one electron to form monopositive ions, Na+ and K+. These ions have 10 and 18 electrons, respectively, so they are quite stable according to the octet rule.

Elements in the last few columns of the periodic table have one, two or three fewer electrons than predicted by the octet rule: they therefore gain electrons fairly easily. For example, the halogens (group VII), such as fluorine (F) or chlorine (Cl), which have, respectively, 9 and 17 electrons, easily gain one electron to form mononegative ions, F- or Cl-. These ions have 10 and 18 electrons, respectively.

Likewise, elements in group II form doubly positive ions such as Mg++ or Ca++, and elements in group VI form doubly negative ions such as O-- or S--. All these ions obey the octet rule and so are fairly stable.


Now, imagine what will happen when one sodium atom meets one chlorine atom: the sodium atom will lose one electron to give Na+, and the chlorine atom will gain that electron to give Cl-. This can be represented schematically in the following way:

The resulting ions, which have Opposite charges, will be attracted to one another, and will draw closer, until they "touch". This happens when the inner shell of electrons on the sodium ion (shown in blue) starts to overlap with the outer shell of electrons on the chloride anion (shown in green). This pair of ions looks something like this:

Clicking on the picture above (and all subsequent pictures!) will bring up a 3D model of the Na+/Cl- pair of ions, in a new window.

It is possible to determine where the valence electrons are situated in this pair of ions. They are almost entirely situated on the chlorine atom, as expected: the sodium atom has lost its only valence (3s) electron, whereas chlorine has gained an electron and has the 3s23p6 valence configuration. The blue transparent surface on this picture encloses the region of space where the valence electrons spend most of the time:


NaCl, or sodium chloride, is however more complicated than this! This is because charge-charge interaction occurs in all directions. Once an Na+ cation has attracted a Cl- anion in one direction, it can attract another in a different direction. So two pairs of ions such as above can come together to form a species with four ions in total, all placed so as to interact favourably with ions of opposite charge:


Here, too, all the valence electrons sit on the chlorine atoms:


And this need not stop here... The next step is to get 8 ions, 4 each of sodium and chlorine:


The stable form of sodium chloride involves a very large number of NaCl units arranged in a lattice (or regular arrangement) millions of atoms across. Because the lattice is rigid, this means that one gets a solid: the ions do not move much one with respect to another. Also, because atoms are so small, even a small crystal of salt will have billions of sodium chloride units in it! The ions are arranged so that each positive (sodium) ion is close to many negative (chloride) ions, as shown on the following picture:

Can you count how many ions each sodium is next to? And how many ions each chlorine is next to? These pairs of ions in close contact are shown with lines joining them. These lines illustrate the strong ionic bonding between ions of opposite charge which are next to each other. However, you should remember that these close contacts are not the same as covalent bonds - there is no pair of electrons shared between the two atoms which are connected by the two lines. Also, there is some ionic bonding between ions which are further away from each other - ions of opposite charge always attract each other, however far they are from each other. Nevertheless, the force holding them together is largest when they are close together. The lines connecting ions in this lattice (and others below) are there to make it easier to detect the pairs of ions in close contact with each other.


Remember - atoms are very small! The distance between a sodium ion and its nearest chloride ion neighbours is about 3 ten-millionths of a millimeter. Imagine a cubic grain of salt with edges which are 3 tenths of a millimeter long. That means there will be a line of about a million ions along each edge. And the grain will contain one billion billion ions in total. If each ion was replaced by a ping-pong ball (roughly 3 centimeters in diameter), each edge would be one hundred million times longer. Instead of being 3 tenths of a millimeter wide, this "grain" would be roughly thirty kilometers (or twenty miles) wide!! Enough to cover most of London...


All ionic compounds adopt a similar three-dimensional structure in which the ions are close to many ions of the opposite charge. There are however several ways of doing this. Caesium chloride (CsCl), for example, adopts a different structure to that of NaCl, as shown on the following picture:

Can you count how many ions each caesium ion (pink) is in close contact with? And how many ions each chloride ion (green) is close to?


As another example, let us consider a salt with a divalent (doubly positive) ion, for example calcium fluorite, CaF2. This adopts the structure shown below (the calcium atoms are shown as large grey spheres, the fluorine atoms are smaller and orange):

Can you count how many fluoride ions each calcium is in close contact with? And how many caesium ions each fluoride ion is close to?


Experienced chemists can often predict the structure that a given ionic species will adopt, based on the nature of the ions involved. This means that it is often possible to design ionic compounds having certain well-defined and desirable properties. As an example, chemists have been able to make high-temperature superconductors, such as the complicated ionic compound, YBa2Cu3O4. This solid conducts electricity with no resistance at all at low temperature (below ca. -100 degrees centigrade). Previous superconductors only had this property at much lower temperatures. The lack of resistance makes superconductors very useful in a number of technological applications - e.g. in designing high-speed trains that levitate above the track!

The repeating structure of this solid is shown below (oxygen is large and red, barium large and yellow-ish, yttrium small and pink, and copper small and blue). Notice how many oxygen ions surround each barium and yttrium ion.


Ionic Bonds - Conclusions


Ionic bonds form between elements which readily lose electrons and others which readily gain electrons. Because the interaction between charges as given by Coulomb's law is the same in all directions, ionic compounds do not form molecules. Instead, periodic lattices with billions of ions form, in which each ion is surrounded by many ions of opposite charge. Therefore, ionic compounds are almost always solids at room temperature. By careful consideration of the properties of each ion, it is possible to design ionic solids with certain well-defined and desirable properties, like superconductors.


Click Here to return to the main structure and bonding webpage.

Click Here to return to the previous page (Basic Principles).

Click Here to go on to the next page (Covalent Bonding).


This page and all its contents belong to and were written by Jeremy Harvey