The fun thing about ionic bonds is that they really aren’t anything like covalent bonds. There is no sharing of electrons, and there is no pairing up of atoms. There isn’t even a single ionic bond that you can point to, since the bonding happens in all directions, attracting ions from top, bottom, right, left, front or back. Ionic bonding can be best described as ions of opposing charges stacking together to form a lattice structure that maximizes the attractive electrostatic forces between the ions. That’s a mouthful if you aren’t used to chemistry and physics already, so it’s a good idea to break it down into subtopics like ions, forces and crystal lattices.
If you’ve studied much atomic chemistry, you already know that atoms strive to achieve a full outer shell of eight electrons, typically known as an octet. (Smaller elements need fewer – lithium will settle for two, as will hydrogen, but hydrogen is even happy dropping to zero). Essentially, they want a nice symmetrical distribution of electrons to give them a stability similar to that already present in the noble gases – the elements at the far right of the periodic table.
There are two strategies for ion formation that provide the full outer shell they crave: gain extra electrons to complete the partial octet the atom already has, or lose the electrons beyond the last full octet. Nonmetals – elements found toward the right end of the periodic table (beyond the stair step, if yours has it drawn in) already have more than half the electrons they need to reach the octet, so they opt for the first option – gaining electrons. Metals – all the elements left of the stair step – follow the second option. They’re only a couple electrons past their last full outer shell, and can shed those extra electrons more readily. Once an atom has gained or lost electrons, it possesses a net charge and is called an ion (charged particle). Positive ions are “cations,” negative ions are “anions.” There are also more complex ions which are groups of covalently bonded atoms with a net charge, known as polyatomic ions, but they are still ions and behave the same in ionic bonding.
While gaining and losing electrons sounds simple enough, no atom actually wants to give up electrons, so there are large amounts of energy transfer involved. In other words, if you want to make sodium ions and chloride ions directly from the elements, beware! The reaction is extremely violent and dangerous. Beginning lessons in ionic bonding tend to make ion formation sound like a cooperative process, but it’s really more like a mugging. The chlorine atom expends a great deal of energy to rip that electron away from sodium.
Electrostatic forces are a fairly simple physics concept. Most people are familiar with the rules “opposite charges attract” and “like charges repel.” Any cation has a positive charge, and will be attracted to an anion, with its negative charge. Cations push away from other cations, and anions repel anions. The rule that not everyone knows is that both attractive and repulsive forces are stronger at closer distances. What this means is that the best arrangements of ions are the ones where cations are right next to anions, and there is is some spacing between cations and other cations, and spacing between anions and other anions.
Since electrostatic forces act in all directions, a cation will pull anions closer from all directions. Anions also pull cations in from all sides. This means that ions can (and do) arrange themselves into a three-dimensional array where they are immediately surrounded by ions with opposing charges.
Ions can be thought of as spheres, like marbles. Unlike marbles, not all ions have the same size. The bromide ion, for instance, is much larger than the chloride ion. Likewise, sodium ions and silver ions have vastly different sizes. What this means is that when cations and anions attract one another, they will pack together differently, depending on their sizes. While all the varying crystal geometries are beyond the scope of this article, it is useful to be aware that all the different crystalline shapes found in nature result from the sizes of the ions that make them up. In this form, ions are locked in place, so students are often surprised to discover that even though ionic compounds contain metal ions, they can’t conduct electricity.
Notice that there is no particular bond between any one pair of ions. Chemists write the formula NaCl to represent sodium chloride, but there aren’t individual particles made up of one single sodium ion and one single chloride ion. What really exists is a nice lattice where each sodium ion is surrounded by six chloride ions, and every chloride ion is surrounded by six sodium ions. While a discrete bond isn’t a good portrayal of what happens in ionic bonding, chemists still talk about bond length and bond energy for ionic bonding.
In packing together a lattice, the assumption is that each spherical ion sits side to side against its neighbor. Therefore, bond length is defined as the sum of the radii of the two adjoining ions. In sodium chloride, this means that the bond length is the radius of a sodium ion plus the radius of a chloride ion. Measuring the size of crystals is the way scientists were originally able to estimate the size of ions.
Bond energy is the amount of energy it takes to break one bond. While it isn’t very meaningful to talk about a single bond in a crystal lattice, it’s still a value that can be calculated. The lattice energy of a crystal can be measured more easily – that’s the energy to break apart all of the ions that form the lattice. Dividing that value by the number of individual ionic bonds within the crystal gives a value for the bond energy.
Those are the basic concepts of ionic bonding. For a more advanced treatment, you can explore the math, including forces, energy, geometry and more. You can discover the marvelous shapes and symmetry that arise from crystalline structure. You can investigate hydrates, solubility, behavior in solution, the molten state and more.