The nature of atom-bound electrons, as we have seen, makes the ionic forms of some atoms much more energetically stable—and thus more common—than the neutral form. For example, we saw that inert gas atoms don't tend to form ions because they have a full octet of electrons in their outer shell. Halides (F, Cl, Br, I, At), on the other hand, have one more electron than the next smaller inert gas (e.g. Na compared to Ne), and tend to form +1 ions by losing that "extra" electron.
The periodic table below shows the most common forms of most of the elements. Where two ions are listed, both are observed but the top form is dominant. Notice that while most of the ions follow the trends we'd expect from electron configuration, there are exceptions. These can seem disturbing in a first run through chemistry, but each can be rationalized in terms of an interplay of competing factors that you will learn about as your study of chemistry becomes more sophisticated.
In the atoms section I mentioned that the typical way an atom with more than an octet of electrons loses one (or more) is by coming into contact with one that has the opposite "need." For example, when sodium metal comes into contact with fluorine gas, the fluorine readily takes an electron from sodium, and both are left in a more energetically favorable state. What results is an ion pair with a neutral overall charge, Na+F-, or just NaF.
Nature tends to neutralize charges, so ions of opposite charge tend to attract to make ionic compounds. Some examples are NaCl (table salt), KCl, CsF, and RbBr. Each is composed of a positive ion and a negative ion.
We try not to refer to ionic compounds as molecules (which we'll get to in a while), because the really never exist in that form. The abbreviation NaCl, for example, is best thought of as the formula unit of sodium chloride (we'll talk about how to name ionic compounds later, too). It's really just the ratio in which the constituent atoms are found in the pure material.
The figure below shows the formula unit of sodium chloride. The relative sizes of the atoms are accurate, although we know they don't have hard edges. On the right is a picture of how sodium chloride formula units can stack together in a cubic configuration to form the well known salt crystals you spill on the table. The cubic crystal of NaCl is charge-neutral: For every positive ion, there is a negative ion.
Notice that some of our ions are multiply charged. For example, the magnesium ion carries a +2 charge. Nature still tends to neutralize this charge, except that now we either need two -1 charged ions or one with a -2 charge. We can form many ionic compounds that satisfy this need. MgO consists of Mg2+ and O2-, and MgCl2 is Mg2+ and two Cl- ions.
In MgCl2, our chemical abbreviation for "magnesium chloride", the subscript 2 means that there are two chlorine ions in the compound. This is a general feature of subscripts in chemical formulae. A subscript in a chemical formula always counts the number of the atom that precedes it.
For each of the ionic compounds below, roll over the box to see how many of each kind of ion makes it up.
You probably noticed that we have naming conventions for ions. For atomic ionic compounds, it's pretty easy, just substitute the suffix "ide" for the last syllable of the anion and let the cation name stay the way it is. Generally, the subscripts don't play a role in naming, although, because chemistry is an old field, there are some exceptions. Here are some examples:
Now to introduce another important class of ions, the molecular ions. So far in our study of chemistry, we haven't talked about molecules, and we really won't here. For now, let's just say that atoms can bind together in another way, quite different from the ionic bonding we've seen so far, and leave it at that. There are certain molecules that are much more stable after giving up one or more electrons or taking on one or more extra electrons - molecular ions.
You can download a chart of frequently-encountered ions – atomic and molecular. It will be worth your time to memorize the highlighted ions, as you will encounter them frequently. Knowing them and their charges will make life easier as you work through chemistry. With any luck, with some frequent use of the other molecular ions, you'll come to memorize those names, too.
Clicking on the hyperlinked (blue) ions will take you to the Wikipedia page so you can learn more if you'd like.
Ionic compounds are very important in chemistry. Here are some common ones formed by combining atomic and molecular ions. For historical reasons, some are known by a more common name, like the acids below. We will figure out the naming of acids later.
|H2SO4||Hydrogen sulfate||Sulfuric acid||2H+ + SO42-|
|Na2SO4||Sodium sulfate||2Na+ + SO42-|
|H3COOH||Hydrogen acetate||Acetic acid||H+ + H3COO-|
|NaNO3||Sodium nitrate||Na+ + NO3-|
|KClO4||Potassium (VII) chlorate||Potassium perchlorate||K+ + ClO4-|
|Li3PO4||Lithium phosphate||3Li+ + PO43-|
|CuCr2O7||Copper dichromate||Cu2+ + Cr2O72-|
|NH4OH||Ammonium hydroxide||NH4+ + OH-|
|HNO3||Hydrogen nitrate||Nitric acid||H+ + NO3-|
A. Using the names given, write the formula unit (chemical formula) for the ionic compound.
|3.||iron (III) bromide|
|4.||iron (III) sulfate|
|5.||lead (II) phosphate|
|8.||cobalt (II) phosphate|
B. Write the name each of the following ionic compounds.
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