Naming Ionic Compounds

I recommend doing any problems with a periodic table in hand – get one at

Ionic Bonding

Before we begin – lets review a little – an ionic compound is one where the elements bonded have very different abilities to hold on to electrons. That strength to hold electrons can be called electron affinity or written as a number called Electronegativity

Electronegativity – the effective attraction an atom exhibits for electrons in a bond within a molecule.


The higher the electronegativity (EN), the more an atom likes electrons. So if we take Sodium (Na, EN = 0.9), a metal and Chlorine (Cl, EN = 3.2), a non-metal we get the following:
Sodium gives up its electron to chlorine! This causes the salt sodium chloride to form. Sodium goes from 11 electrons to 10. Chlorine gains one electron and goes from 17 to 18. Notice that both elements now have noble gas configurations – they are stable (the noble gases are Group 18 and contain He, Ne, Ar, Kr, Xe, Rn).

Ionic compounds do not have an actual bond between them – instead,  the ions are strongly attracted because they have opposite charges (kind of like North and South poles on magnets). This attraction causes the ions to form a crystal lattice (a regularly spaced shape):

salt crystal-smallcrystal-955935_960_720-small

To name an ionic compound, you need to find the names of the cation and anion present in the compound. The positive species or cation (think cats have paws or cations are pawsitive!) is written first, followed by the anion or negative species.

Na+ and Cl would be written as sodium chloride.


Notice that the atoms are sodium and chlorine but the ions are called sodium and chloride → whenever an atom has a negative charge we add the suffix -ide!

The ide suffix literally implies a negative ion.

You should first write out the name of the metal, followed by the name of the nonmetal with its respective new ending:

Formula Ions Name
LiBr lithium ion (Li+), bromide ion (Br) lithium bromide
Na2O sodium ion (Na+), oxide ion (O2-) sodium oxide
Na3P sodium ion (Na+), oxide ion (P3-) sodium phosphide
YN yttrium ion (Y3+), nitride ion (N3-) yttrium nitride

Notice how the chemical formula is related to the charges:

lithium bromide (LiBr)


The charges used to make the salt criss-cross. Note how when charges are +1 or -1 we can just write + or -. If a charge is zero there is nothing to write (alternatively we can put a zero to clarify).
This is because charged objects are VERY rare – you rarely get a shock from touching a mineral or object around you. Chemical reactions usually result in a conserved net charge of zero:

sodium oxide (Na2O)


In the case of sodium oxide, two sodiums must balance the charge of one oxygen
2 Na++O2-→Na2O

yttrium nitride (YN)

Notice how the ratio of atoms or ions in yttrium nitride is 1 to 1 – we ALWAYS use the lowest common ratio for salts because salts actually contain billions to trillions of atoms but they are always in a precise, stoichiometric ratio – it would be redundant for us to write Y1000N1000 for a crystal that contains 1000 atoms of each.

Notice that I wrote the above equation backwards – just to illustrate that this is more about charge balance than which way (forwards or backwards) the above reaction is going.

Chemical names are NOT capitalized unless they are at the beginning of a sentence.

Practice Part 1

Try writing the names of the following ionic compounds (solutions at the end of the page)






Try writing the formulas of the following ionic compounds (solutions at the end of the page)

aluminum oxide

lithium fluoride

scandium bromide

scandium nitride

magnesium oxide

Ionic Compounds with Transition Metals

Metals (on the left of the periodic table, hydrogen being an exception) lose electrons and become more stable by approaching noble gas configurations. Non-metals (on the right of the periodic table) gain electrons and add stability by also becoming like noble gases.

There is a limit because at a certain point the charges on the atom become so large it becomes unstable! You can see that the trends for number of electrons lost or gained completely break down in groups 4 to 12.

common charges-small

The charge on an atom is also known as oxidation number or oxidation state. If you think about it, oxidation is related to iron rusting (or oxidizing) and paper burning (or combusting)… and you’d be right! But the tie in for oxidation to chemical reactivity is for a later lesson.

Atoms exchange electrons to increase their stability and become compounds. Stability is correlated with energy and energy is really closely related to temperature. This should give you a hint as to why we see more than one charge for many elements – how many electrons an element loses or gains depends on the temperature and environment!

Transition Metals

Transition metals (groups 3 to 12) have too many electrons to lose to become stable. So they lose as many as they comfortably can. These partially filled/unfilled shells cause transition metals to have a lot of electronic transitions (a transition is just a movement of an electron).

Okay… So what’s important about that? Well when electrons move through the air, they cause sparks. Similarly, transition metals have a variety of colors because their electrons relax or transition to many states and interact with light. The name “transition metals” originally had to do with the metals creating colors in solution:

forms of copper (Small)

Copper (Cu or Cu0) as a metal is brown, as a 1+ ion is green and as a 2+ ion is blue. Copper exists in several oxidation states depending on the temperature and surrounding elements it reacts with (environment).

Because of this we must specify what charge metals like copper have:

copper (I) chloride has the formula CuCl  (from the ions Cu+ and Cl)

copper (II) chloride has the formula CuCl2   (from the ions Cu2+ and Cl)

copper (III) chloride has the formula CuCl3   (from the ions Cu3+ and Cl)

Notice that the charge is always in roman numerals (I, II, III, IV, V, VI, VII, VIII, IX, X). Also notice that I wrote the formulas of copper (I), (II) and (III)… but the periodic table at the beginning of this lesson only had 2+ and 1+ as common states! Again – those numbers are “common” – not an absolute.

IF the compound is a transition metal (except for Zinc), have the charge in brackets

IF there is more than one common oxidation number, have the charge in brackets.

IF the atom has just one common oxidation number but your specific compound has a different oxidation number, you MUST write the charge. ←This doesn’t happen in much in high-school chemistry.

Here are some example ionic compounds where we need to list the charge:

Name Ions Formula
titanium (III) oxide Ti3+, O2- Ti2O3
titanium (IV) oxide Ti4+, O2- TiO2
tantalum oxide Ta3+, O2- Ta2O3
tantalum (VII) oxide Ta7+, O2- Ta2O7

Notice that for titanium (IV) oxide we used the lowest ratio of the atoms – instead of Ti2O4 we wrote TiO2. As a transition metal, tantalum should have its oxidation state written.

Note that you don’t necessarily need to write down the ions like I did in the middle column but it helps A LOT when first learning to name ionic compounds.

Polyatomic Ions

Ionic compounds often contain anions (negative ions) that are more complex than a single atom.

The table below shows the most common polyatomic ions you will encounter.

name formula name formula name formula
acetate (ethanoate) CH3COO chromate CrO42- phosphate PO43-
ammonium NH4+ dichromate Cr2O72- hydrogen phosphate HPO42-
benzoate C6H5COO cyanide CN dihydrogen phosphate H2PO42-
borate BO33- hydroxide OH silicate SiO32-
carbide C22- iodate IO3 sulfate SO42-
carbonate CO32- nitrate NO3 hydrogen sulfate HSO4
hydrogen carbonate HCO3 nitrite NO2 sulfite SO32-
perchlorate ClO4 oxalate OOCCOO2 hydrogen sulfite HSO3
chlorate ClO3 hydrogen oxalate HOOCCOO hydrogen sulfide HS
chlorite ClO2 permanganate MnO4 thiocyanate SCN
hypochlorite OCl or ClO peroxide O22- thiosulfate S2O32-
persulfide S22-

There is a pattern to naming polyatomic ions however even as a graduate level chemist I found it a waste of time – typically we just memorize a few common variations and can easily look up the rest.

Naming ionic compounds with polyatomic ions is just as easy as naming them with a single anion:

name ions formula
sodium chloride Na+, Cl NaCl
sodiu chlorite Na+, OCl NaOCl
sodium hydrogen oxalate Na+, HOOCCOO NaHOOCCOO or HOOCCOONa or NaHC2Oor HC2O4Na (least common)
sodium oxalate Na+, OOCCOO2 Na2OOCCOO or Na2C2O4 (common)
iron  (III) chlorite  Fe3+, OCl Fe(OCl)3
calcium oxalate Ca2+, OOCCOO2  CaOOCCOO or CaC2O4
titanium (IV) oxalate Ti4+, OOCCOO2 Ti(OOCCOO)2 or Ti(C2O4)2

Some big differences to note are that when a formula needs two or more of a polyatomic ion, the entire ion is placed in brackets – which makes sense, its treated like one big, charged piece.

There are a variety of standards in chemical naming as defined by IUPAC (the International Union of Pure and Applied Chemistry) however, depending on the situation or the age (old documents before standardization) a lot of variation exists. We can see that sodium hydrogen oxalate can be written as HOOCCOONa to indicate which oxygen the sodium will spend time around. It can also be written as NaHOOCCOO to follow an IUPAC rule of naming where positive ions are written before negative ions (kind of like how sodium chloride is always Na+ first then Cl or NaCl). If we don’t care about the structure of the ion, we can write NaHC2O4.

oxalate ion with charges shown sodium hydrogen oxalate
oxalate-2 sodium hydrogen oxalate
 The 2- charge is actually split across the two oxygens shown The covalent O-H bond is shown.

Sodium ions and O have an ionic interaction – there is no bond, they just spend time close to each other due to the charge attraction.

Often other chemical formulas like HC2O4Na or C2O4HNa pop up but they should be marked wrong on a test – you’re safest putting positive ions first followed by negative ions.

Practice Part 2

Try writing the formulas of the following ionic compounds (solutions at the end of the page)

zirconium chloride

zirconium oxide

niobium (III) oxide

niobium (IV) oxide

niobium (V) oxide

manganese (II) hydroxide

manganese (II) thiosulfate

manganese (III) thiosulfate

Try writing the names of the following ionic compounds (solutions at the end of the page)









Part 1

NaI sodium iodide

KBr potassium bromide

K2O potassium oxide

Cs3N cesium nitride

CaO calcium oxide

aluminum oxide Al2O3

lithium fluoride LiF

scandium bromide ScBr2

scandium nitride ScN

magnesium oxide MgO

Part 2

zirconium chloride ZrCl4

zirconium oxide ZrO2

niobium (III) oxide Nb2O3

niobium (IV) oxide NbO2

niobium (V) oxide Nb2O5

manganese (II) hydroxide Mn(OH)2

manganese (II) thiosulfate MnS2O3

manganese (III) thiosulfate Mn2(S2O3)3

NaCN sodium cyanide

K2CrO4 potassium chromate

Ca3(BO3)2 calcium borate

OsO2 osmium (IV) oxide

Ca(HOOCCOO)2 calcium hydrogen oxalate

Ru(NO2)3 ruthenium (III) nitrite


  • All of the ionic compounds we have named so far have two parts. This isn’t because all ionic compounds do – its only because we only name the easy ones – the “binary ionic compounds” <– binary meaning composed of two parts.
  • For an overview of ALL the different types of bonds see the article “Forms of Matter” for clarification.
  • All images are my original drawings except for credited photographs.
  • Not all ionic compounds are salts.
  • Electron Affinity is not exactly the same thing as Electronegativity but it is a related property.
  • Salt Crystal photo from
  • Notice that compounds like Osmium have 1 oxidation state given but being in the middle of the transition metal block, it most likely has other states and charge needs to be specified.
  • No copper ions (or any other metal for that matter) float in solution by themselves – there is always a counter ion! For instance Cu+ came from CuCl. This again is because chemical reactions produce a zero charge difference. If they didn’t, salt solutions would be physically attracted to each other and pull each other across the room!


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