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4.6 Ionic Compounds

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    Ionic Formulas

    Chemical formulas for ionic compounds are called ionic formulas. A proper ionic formula has a cation and an anion in it; an ionic compound is never formed between two cations only or two anions only. If you've ever tried to force two refrigerator magnets together you'll remember that there was one side of each magnet that clicked together and there was one side of each magnet that pushed against the other. One side of each of our magnets is positive and one side of each of our magnets is negative. When the magnets push away from each other we're trying to force two like charges together. Ions behave the same way on the atomic scale. Two ions with the same charge will push away from each other. Two ions with opposite charges will be attracted to each other as being close to one another allows them to cancel out the additional charge they are carrying.  A Ca2+ cation has twice as much charge as a Cl- anion, so it will be able to attract and cancel out the charge on two Cl- anions.


    The key to writing proper ionic formulas is simple: the total positive charge must balance the total negative charge. Because the charges on the ions are characteristic, sometimes we have to have more than one of a cation or an anion to balance the overall positive and negative charges. It is conventional to use the lowest ratio of ions that are needed to balance the charges.

    For example, consider the ionic compound between Na+ and Cl. Each ion has a single charge, one positive and one negative, so we need only one ion of each to balance the overall charge. When writing the ionic formula, we follow two additional conventions: (1) write the formula for the cation first and the formula for the anion next, but (2) do not write the charges on the ions. Thus, for the compound between Na+ and Cl, we have the ionic formula NaCl (Figure \(\PageIndex{1}\)). The formula Na2Cl2 also has balanced charges, but the convention is to use the lowest ratio of ions, which would be one of each. (Remember from our conventions for writing formulas that we don’t write a 1 subscript if there is only one atom of a particular element present.) For the ionic compound between magnesium cations (Mg2+) and oxide anions (O2−), again we need only one of each ion to balance the charges. By convention, the formula is MgO.


    Figure \(\PageIndex{1}\): NaCl = Table Salt © Thinkstock The ionic compound NaCl is very common.

    For the ionic compound between Mg2+ ions and Cl ions, we now consider the fact that the charges have different magnitudes, 2+ on the magnesium ion and 1− on the chloride ion. To balance the charges with the lowest number of ions possible, we need to have two chloride ions to balance the charge on the one magnesium ion. Rather than write the formula MgClCl, we combine the two chloride ions and write it with a 2 subscript: MgCl2.

    What is the formula MgCl2 telling us? There are two chloride ions in the formula. Although chlorine as an element is a diatomic molecule, Cl2, elemental chlorine is not part of this ionic compound. The chlorine is in the form of a negatively charged ion, not the neutral element. The 2 subscript is in the ionic formula because we need two Cl ions to balance the charge on one Mg2+ ion.

    Example \(\PageIndex{2}\):

    Write the proper ionic formula for each of the two given ions.

    1. Ca2+ and Cl
    2. Al3+ and F
    3. Al3+ and O2−


    1. We need two Cl ions to balance the charge on one Ca2+ ion, so the proper ionic formula is CaCl2.
    2. We need three F ions to balance the charge on the Al3+ ion, so the proper ionic formula is AlF3.
    3. With Al3+ and O2−, note that neither charge is a perfect multiple of the other. This means we have to go to a least common multiple, which in this case will be six. To get a total of 6+, we need two Al3+ ions; to get 6−, we need three O2− ions. Hence the proper ionic formula is Al2O3.

    Exercise \(\PageIndex{2}\)

    Write the proper ionic formulas for each of the two given ions.

    1. Fe2+ and S2−
    2. Fe3+ and S2−


    1. FeS
    2. Fe2S3



    4.6 Ionic Compounds is shared under a CC BY-NC-SA 4.0 license and was authored, remixed, and/or curated by LibreTexts.

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