# Naming monatomic ions and ionic compounds

Learn how to name monatomic ions and ionic compounds containing monatomic ions, predict charges for monatomic ions, and understand formulas.
Close up view of colorless sodium chloride crystals, which have the overall shape of a cube.
Sodium chloride is an ionic compound made up of sodium ions and chloride ions in a crystal lattice. Image credit: Wikipedia Commons, public domain
Atoms are electrically neutral because the number of protons, which carry a 1+ charge, in the nucleus of an atom is equal to the number of electrons, which carry a 1- charge, in the atom. The result is that the total positive charge of the protons cancels out the total negative charge of the electrons so that the net charge of the atom is zero. Most atoms, however, can either gain or lose electrons; when they do so, the number of electrons becomes different from the number of protons in the nucleus. The resulting charged species is called an ion.

## Cations and anions

When a neutral atom loses one or more electrons, the total number of electrons decreases while the number of protons in the nucleus remains the same. The result is that the atom becomes a cation—an ion with a net positive charge.
The opposite process can also occur. When a neutral atom gains one or more electrons, the number of electrons increases while the number of protons in the nucleus remains the same. The result is that the atom becomes an anion—an ion with a net negative charge. We can illustrate this by examining some very simple cations and anions, those formed when a single hydrogen atom loses or gains an electron.
Note: Hydrogen is actually somewhat unusual in that it readily forms both cations and anions. Most elements much prefer to form only one or the other. In terms of its electron configuration, can you explain why hydrogen can form both cations and anions? Feel free to post in the comments at the end of the article!
A hydrogen cation, a hydrogen atom, and a hydrogen anion.
$~~~~~~~~~\text{H}^+~~~~~~~~~$$~~~~~~~~~~~~~~~~~~~~~~~~\text{H}~~~~~~~~~~~~~~~~~~~~~~~~$$~~~~~~~~~~~~~~~~~~\text{H}^-~~~~~~~~~~~~~~~~~~~$
Classificationcationneutral atomanion
No. of protons$1$$1$$1$
No. of electrons$0$$1$$2$
Net charge$1$$+$$0$$1$$-$
If a neutral hydrogen atom ( $\text{H}$, center) loses an electron, it becomes a hydrogen cation ( $\text{H}^+$, left). Conversely, if the neutral $\text{H}$ atom gains an electron, it becomes a hydrogen anion ( $\text{H}^-$, right), also known as a hydride ion. Image credit: adapted from Boundless Learning, CC BY-SA 4.0.
In the center column, we have a diagram of a single, neutral hydrogen atom. It contains one proton and one electron; thus, its net charge is zero. If hydrogen loses its electron, it forms the cation $\text{H}^+$ (left column). The $\text{H}^+$ cation has a net charge of 1+ from the one proton in the nucleus since there are no electrons to cancel out the positive charge. If neutral hydrogen gains an electron, it forms the anion $\text{H}^-$ (right column). The $\text{H}^-$ anion has a net charge of 1- because it has one extra electron compared to the total number of protons.
Concept check: A certain ion has 20 protons and 18 electrons. What kind of element is this ion, and what is its net charge?
Because the ion has 20 protons, the element in question must be calcium, $\text{Ca}$, which has an atomic number of 20. The 20 protons contribute a total charge of 20+, and 18 electrons contribute a total charge of 18-:
Thus, the net charge on the ion will be 2+. We can write the symbol for the calcium 2+ cation as follows: $_{20}\text{Ca}^{2+}$ or $\text{Ca}^{2+}$.

## Predicting charges on monatomic cations and anions

Did you know that you can use the periodic table to predict the charges certain elements will have when they ionize? This is a very convenient and powerful tool, so it's worth examining in some detail. The following figure summarizes the common charges for the elements in the eight main groups, or families, on the periodic table. Remember that periodic groups refer to columns on the periodic table, whereas rows are known as periods. Keep in mind that these charges only apply when these elements are found in ionic compounds since covalent compounds don't contain ions.
Ionic compounds are made up of ions, which are held together by ionic bonds between ions of opposite charge.
In covalent compounds, however, electrons are shared in covalent bonds, so there are no true ions with full charges on them. Elements in covalent compounds, however, can be assigned oxidation states or oxidation numbers, which are similar to an apparent charge. We can consider oxidation numbers to be what the charge on the element would be if the electrons in the covalent bond were completely transferred to the more electronegative atom. If you would like to learn more about oxidation states, you can watch this video on oxidation states and how they are used.
That said, it is important to realize that the distinction between ionic bonds and covalent bonds isn't always clear. Instead of thinking about all compounds as falling into one of two categories, it is more accurate to think of a continuum between ionic and covalent bonds, and compounds. Many compounds contain bonds with some fraction of covalent character as well as some ionic character.
For Group 14 elements, forming cations with a 4+ charge is much more common than forming anions with a 4- charge. Carbon, however, can form both types of ions, so we include both charges here.
As a general rule of thumb, the main group elements will usually gain or lose electrons in order to get a full octet of valence electrons. By figuring out how many electrons an element is likely to lose or gain to reach a full octet, we can predict the charge on the ion. This requires first knowing how many valence electrons are in the neutral atom.
Tip: The number of valence electrons in the neutral atom is equal to the number in the $\blueD{1}$s place in the new IUPAC group number.

### Elements that form cations

For groups 1, 2, 13, and 14, the elements have one to four valence electrons as neutral atoms, and they will usually give away these valence electrons to become ions—carbon is sometimes an exception to this trend since it can also gain four electrons to form the $\text C^{4-}$ anion. Since the resulting ion has fewer electrons than protons, the net charge on the ion is positive. The magnitude of the charge is equal to the number of electrons lost, which is equal to the number of valence electrons in the neutral atom.
For example, what if we wanted to predict the charge on an aluminum ion? Aluminum is in group 13, or IIIA. Since the group number, $1 \blueD3$, has the number $\blueD 3$ in the $1$s place, we would predict the charge to be $\blueD 3$$\blueD +$ to give $\text{Al}^{3+}$. We can also think about a neutral aluminum atom losing its three valence electrons to become $\text{Al}^{3+}$, which has a full octet.
Yup! When we talk about atoms losing or gaining electrons in order to get to a full octet, we can also think of this process as the atom losing or gaining electrons to get to the same electron configuration as the nearest noble gas.
For our current example, the $\text{Al}^{3+}$ ion has the same electron configuration as the noble gas neon, $\text {Ne}$. They both have the following electron configuration:
$\text{1s}^2 \text{2s}^2 \text{2p}^6$

### Elements that form anions

For groups 15 through 17, the charge is usually negative because these elements are more likely to gain than lose electrons. The charge on the ion is therefore equal in magnitude to the number of electrons gained to reach a full octet of eight valence electrons. Mathematically, we can calculate the magnitude of the charge by subtracting the number of valence electrons in the neutral atom from eight. We can also use the periodic table to count how many columns to the right we need to go to reach the noble gases, group 18, where each adjacent column counts as one electron that needs to be gained to reach the full octet.
If we use these guidelines to predict the charge on a sulfur ion, which is in group 16, we predict that the magnitude of the charge is $8-6=2$ since sulfur has six valence electrons. We can also find the number of valence electrons by checking sulfur's group number, group 16, which has a $6$ in the $1$s place. That means that a neutral sulfur atom will need to gain two electrons to reach a full octet of eight electrons. Therefore, we predict that the most common charge on a sulfur ion will be 2-.
Concept check: What ionic compound would you predict to form in a reaction between potassium metal and liquid bromine?
Potassium metal is a group 1 alkali metal and therefore usually forms cations with a 1+ charge, namely, $\text{K}^+$.
Bromine, as a group 17 halogen, usually forms anions with a 1- charge, namely, $\text{Br}^-$.
Because the charges on these ions are equal and opposite, they will combine in a 1:1 ratio so that their charges exactly cancel. Therefore, we predict that the ionic compound formed will be $\text{KBr}$, potassium bromide.

## Naming cations

Now that we know that many common elements take on predictable charges, let's consider how to name the ions. We'll first look at the alkali metals—the elements in group 1 on the periodic table. From the figure above, we can see that the alkali metals tend to form cations with a 1+ charge. Thus, these cations include $\text{H}^+$, $\text{Li}^+$, $\text{Na}^+$, $\text{K}^+$, and so on. Naming these types of cations requires no special rule. For instance, we can refer to a hydrogen cation, $\text{H}^+$, simply by calling it "$\text{H}$-plus" or a "hydrogen ion". Similarly, a sodium cation, $\text{Na}^+$, can be called "$\text{Na}$-plus", "sodium plus", or most commonly, a "sodium ion". Note that it is unnecessary to say "a one plus sodium ion", because it is understood that a sodium ion usually has a 1+ charge.
The same logic also applies to all other elements that typically form cations of one particular charge. For instance, the alkaline earth metals, group 2, form cations with a charge of 2+: $\text{Be}^{2+}$, $\text{Mg}^{2+}$, $\text{Ca}^{2+}$, etc. While we often refer to an ion such as $\text{Mg}^{2+}$ as "magnesium two-plus", we could also simply say "magnesium ion", since it's understood what the charge on a magnesium ion is.
Note: The discussion in this section is mainly for naming cations by themselves, and the naming convention will be slightly different when the cation is part of an ionic compound. The naming of ionic compounds will be discussed separately below!

## Elements that form multiple types of cations

So far, we have considered elements that typically form cations of one particular charge. For example, the alkali metals and the alkaline earth metals usually form 1+ ions and 2+ ions, respectively. Most transition metals, however, can form cations of various charges. That is why the d-block of the periodic table figure above has been labeled "variable charges". Iron, for instance, is often found as both the $\text{Fe}^{2+}$ and $\text{Fe}^{3+}$ cations, and sometimes other charges as well. Thus, iron is polyvalent, which literally means "many valued"—it is able to form cations of different charges.
On left, glass vial containing chromium (II) chloride, a grey-green powder and, on right, glass vial containing chromium (III) chloride, a bright purple powder.
Chromium commonly forms compounds as $\text{Cr}^{2+}$ and $\text{Cr}^{3+}$. Chromium (II) chloride, on the left, is a grey-green solid with very different properties and reactivity compared to chromium (III) chloride, the bright purple solid on the right. Thus, it is important to specify which one you are referring to! Image credit: Maria Sanford
For metals that are polyvalent, we need to specify the magnitude of the charge on the ion. For instance, we have to call $\text{Fe}^{2+}$ "iron two-plus" or "iron two" because simply referring to it as "iron ion" will not give enough information to specify the type of cation. Most of the transition metals—those metals in the center d-block of the periodic table—are polyvalent. Since they can form cations with different charges, those charges must be specified when naming the ions and when naming compounds containing those ions.
In ionic compounds, the magnitude of the charge for a transition metal cation is usually included using Roman numerals in parentheses after the name of the metal, such as for chromium (II) chloride which contains $\text{Cr}^{2+}$ (see picture on the right). Naming ionic compounds containing transition metal cations will be discussed in more detail in a separate section below.

## Naming monatomic anions

Most often, when we name monatomic anions, we add the suffix -ide to the end of the element's name. Because we can predict the charge on simple cations and anions based on an element's group number, it is not necessary to specify the magnitude of the charge on an anion most of the time. The following table shows how the suffix applies to naming anions of various elements:
Element name$\rightarrow$Ion nameIon formula
HydrogenHydride$\text{H}^-$
Chlorine Chloride$\text{Cl}^-$
Bromine Bromide$\text{Br}^-$
Iodine Iodide$\text{I}^-$
Oxygen Oxide$\text{O}^{2-}$
Sulfur Sulfide$\text{S}^{2-}$
Nitrogen Nitride$\text{N}^{3-}$
Phosphorus Phosphide$\text{P}^{3-}$
Carbon Carbide$\text{C}^{4-}$

## Formulas and naming of basic ionic compounds

Now that we've seen the naming conventions for cations and anions, we can discuss how they apply to naming simple ionic compounds made up of monatomic ions. The following guidelines can be used for naming ionic compounds:
• Always name the cation before the anion; in the chemical formula, the cation will always appear first as well.
• When naming the cation within an ionic compound, we don't include the word ion or the charge unless it is a polyvalent cation. That means we only have to name the element that the ion came from (see Example 2 below).
• Any ionic compound will have a net charge of zero. Another way of saying this is that cations and anions must always combine in such a way so that their charges cancel.
• The number of cations and anions in the formula should be written as the lowest possible integer value. For example, the formula for sodium chloride is $\text{NaCl}$, not $\text{Na}_2 \text{Cl}_2$ or some other multiple of $\text{NaCl}$, even though the charges would still add up to zero.
Let's look at a few more examples.

## Example 1: Finding the chemical formula from the name

What is the chemical formula of potassium chloride?
Remember that potassium is a group 1 element that forms a 1+ ion. Chloride, by definition, is an anion that has formed from an atom of chlorine. Since chlorine is in group 17, it will form a 1- anion. Because their charges are equal and opposite, there will be one $\text{ K}^+$ ion for every one $\text{ Cl}^-$ anion, and the chemical formula will be $\text{KCl}$. Remember that subscripts are not used when there is only one atom/ion of a particular type.

## Example 2: Finding the name from the chemical formula

What is the name of the ionic compound $\text{Mg}_3\text{P}_2$?
Magnesium, $\text{Mg}$, is a group 2 element that will form 2+ cations. Because it usually forms cations of only one type, we don't need to specify its charge. We can simply refer to the cation in the ionic compound as magnesium. Phosphorus, $\text{P}$, is a group 15 element and therefore forms 3- anions. Because it is an anion, we add the suffix -ide to its name to get phosphide as the name of the ion. Therefore, the name for the compound is magnesium phosphide.

## Try it: Names and formulas of ionic compounds

Problem 1
What is the chemical formula for calcium bromide?
Choose 1 answer:
Choose 1 answer:
Calcium forms $\text{Ca}^{2+}$ cations, and bromide refers to a $\text{Br}^-$ anion.
In order for the charges on the ions to cancel, we need one $\text{Ca}^{2+}$ cation to bond with every two $\text{ Br}^{-}$ anions. This is because the total charge contributed by calcium will be:
$1\times(2$$+)=2$$+$
The total charge contributed by bromide will be:
$2\times(1$$-)=2$$-$
The $2$$+$ and the $2$$-$ exactly cancel.
Therefore, $\text{CaBr}_2$ is the correct answer.
Problem 2
What is the name of the compound $\text{SrF}_2$?
Choose 1 answer:
Choose 1 answer:
$\text{F}$ is the anion, so it will have the suffix -ide. $\text{F}$ is the chemical symbol for fluorine; here, fluorine has gained an electron to form the anion fluoride, $\text F^-$.
$\text{Sr}$ is the chemical symbol for strontium, a group 2 element. Based on strontium's position on the periodic table, we would predict it to form $\text{Sr}^{2+}$ cations.
When naming ionic compounds, the name of the element that contributed the cation comes first, followed by the name of the anion—which should end in -ide.
Therefore, strontium fluoride is the correct answer.
Problem 3
What are the constituent ions present in the compound $\text{Al}_2\text{S}_3$?
Choose 1 answer:
Choose 1 answer:
Aluminum is a group 13 metal that usually form cations. Sulfur is a group 16 nonmetal that usually forms anions.
Aluminum forms $\text{Al}^{3+}$ cations, and sulfur forms $\text{S}^{2-}$ anions. From the chemical formula for $\text{Al}_2\text{S}_3$, we can see that there are two $\text{Al}^{3+}$ ions for every three $\text{S}^{2-}$ ions.
Therefore, the answer is two $\text{Al}^{3+}$ ions and three $\text{S}^{2-}$ ions.

## Naming ionic compounds with polyvalent cations

Recall from our earlier discussion that if an element can form more than one type of cation, we have to specify the charge on that cation. The magnitude of the charge for a transition metal cation is usually indicated using Roman numerals in parentheses after the name of the metal—this is also called the systematic name of the ion. The following table lists some of the most common ions for polyvalent metals. The systematic name is included for all ions. For some ions, the common or trivial name is also given. The common or trivial names are somewhat old fashioned nowadays, but they're still used in some places, so they're helpful to know. Notice that the ions of lesser charge take the suffix -ous in the common name; ions of higher charge take the -ic suffix. For example, ferrous chloride ($\text{FeCl}_2$) is the name of $\text{Fe}^{2+}$, while ferric chloride ($\text{FeCl}_3$) is understood to contain $\text{Fe}^{3+}$.
ElementCommon ions formedSystematic nameCommon (trivial) name
Chromium$\text{Cr}^{2+}$chromium (II)chromous
$\text{Cr}^{3+}$chromium (III)chromic
Cobalt$\text{Co}^{2+}$cobalt (II)
$\text{Co}^{3+}$cobalt (III)
Copper$\text{Cu}^{+}$copper (I)cuprous
$\text{Cu}^{2+}$copper (II)cupric
Iron$\text{Fe}^{2+}$iron (II)ferrous
$\text{Fe}^{3+}$iron (III)ferric
Lead$\text{Pb}^{2+}$lead (II)
$\text{Pb}^{4+}$lead (IV)
Tin$\text{Sn}^{2+}$tin (II)stannous
$\text{Sn}^{4+}$tin (IV)stannic
Using this table as a reference, let's look at how to name ionic compounds containing polyvalent metals.

## Example 3: Naming compounds containing polyvalent cations

What is the name of the compound $\text{PbCl}_4$?
When naming ionic compounds that contain transition metals, we first need to determine the charge on the transition metal cation. We can deduce this charge by first calculating the charge contributed by the anion, whose charge we already know for sure.
We recognize that $\text{Cl}$ is a group 17 halogen, so it forms the chloride anion $\text{Cl}^-$. We can see from the chemical formula $\text{PbCl}_4$ that there are four chloride ions in the compound. The total negative charge contributed by the four chloride ions is calculated below:
$\text{Total charge from anions}=4\times(1$$-$$)=4$$-$
In order for the compound to be electrically neutral, the lead cation must be $\text{Pb}^{4+}$. This is because the 4+ charge on this ion will exactly cancel the net 4- charge contributed by the four chloride ions.
Therefore, the name of $\text{PbCl}_4$ is lead (IV) chloride.

## Try it: Ionic compounds containing polyvalent cations

What is the name of the compound $\text{Co}_2\text{S}_3$?
Choose 1 answer:
Choose 1 answer:
We will want to start by determining the charge for the ion that is not a transition metal and therefore has a known charge. In this case, it is sulfur, $\text S$. If you look on the periodic table, $\text{S}$ is a group 16 element that commonly forms anions with a 2- charge.
The anion in the compound is $\text{S}^{2-}$, which is called sulfide. From the chemical formula, we can see that there are three sulfide ions in this compound. Therefore, these three sulfide ions contribute a net charge of 6-:
$3\times (2$$-$$)=6$$-$
In order for the compound to be electrically neutral, the net charge on the two $\text{Co}$ cations must be equal to 6+. Thus, cobalt must be cobalt (III), or $\text{Co}^{3+}$, since the total charge from the cobalt ions will then cancel out the charge from the anions:
$2\times (3$$+)=6$$+$
Therefore, cobalt (III) sulfide is the correct answer.

## Conclusion

Cations are positively charged ions formed when neutral atoms lose electrons; anions are negatively charged ions formed when neutral atoms gain electrons. It is possible to predict the charges of common monatomic ions by looking at the group numbers on the periodic table. However, many of the transition metals are polyvalent, which means they can form cations of multiple charges. When naming these cations or compounds containing these cations, it is necessary to specify their charge.
Cations and anions combine to form ionic compounds. Ionic compounds are named with the cation first and the anion last. The same convention is used when writing their chemical formulas. Ionic compounds must be electrically neutral. Therefore, the cations and anions must combine in such a way that the net charge contributed by the total number of cations exactly cancels the net charge contributed by the total number of anions.

## Attributions

1. Modified from "Ions", Boundless Learning, CC BY-SA 4.0.
The modified article is licensed under a CC-BY-NC-SA 4.0 license.

## Additional References

Zumdahl, S. S., and S. A. Zumdahl. "Atomic Structure and Periodicity." In Chemistry, 290-94. 6th ed. Bston, MA: Houghton Mifflin Company, 2003.