2.7: Nomenclature of Ionic, Covalent, and Acid Compounds (2024)

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learning objectives

Recognize ionic compound, molecular compounds, and acids.
Give the names and formulas for ionic compounds
Give the names and formulas for molecular compound s
Give the name and formula for acidic compound

Introduction

Nomenclature, the naming of chemical compounds is of critical importance to the practice of chemistry, as a chemical can not only have many different names, but different chemicals can have the same name! So how can you do science if you do not know what you are talking about??? In this section we will look at nomenclature of simple chemical compounds. The rules we use depends on the type of compound we are attempting to name.

Overview of Nomenclature

The first question we ask is if the compound is ionic or covalent? That is, does it have ionic bonds, or covalent bonds? If it is covalent, which is typically between 2 or more nonmetals, we need to ask, is it a simple molecule, or is it an acid. If it is a simple molecule we use Greek prefixes to identify the number of atoms of each type of element in the molecule. If it is an acid, we base it's name on the ionic compound it would form if hydrogen could be a cation. Note, hydrogen can not lose its only electron as then it would be a subatomic particle and the charge density would be too high, so it forms a covalent bond.

If the compound is ionic, we use the principle of charge neutrality to name the compound.

Figure\(\PageIndex{1}\):Outline of strategy for naming molecules (watch the Youtube immediately below for a lesson on how to use this outline)

Video\(\PageIndex{1}\): 2:17 min Youtube describing the logic of the flow diagram for nomenclature.

We will start with ionic compounds, then covalent and then acids.

Ionic Compounds

Ionic compounds have a [+] cation and a [-] anion. We use the Principle of Charge Neutrality, that is, for an ionic compound to be stable its chemical formula MUST BE NEUTRAL. So you do not need to state the number of cations and anions, you only need to state what they are (and what their charge is). You then figure the formula based on the lowest whole number ratio of cations to anions that produces a neutral formula.

We will start with binary ionic, which are between monatomic anions and cations. We note that there are two types of metals, those that have only 1 charge (Type 1), and those that can have more than one stable charge. Finally, polyatomic ions often form which are covalently bonded atoms where the total number of protons is not equal to the total number of electrons.

Binary Ionic

Binary ionic compounds are between a metal and nonmetal. This does not mean there are two atoms, but two types of atoms, so Al₂S₃ is a binary ionic compound. The rules are simple, name the cation first and the anion second, giving the anion the -ide suffix.

Cation (metal) first
Anion (nonmetal) second with -ide suffix

Example 1:

Na⁺ + Cl^- = NaCl Ca²^{+ +}2Br^- = CaBr₂

Sodium + Chlorine = Sodium Chloride Calcium + Bromine = Calcium Bromide

Two Types of Monatomic Ions

As we have seen in the previous section, there are two types of monatomic ions, those of elements that form only one charge state, and those that can form multiple charged states. The anions are all of the first type, and gain electrons until they have the same number as the nearest noble gas. But metals form cations by losing electrons, and some metals form only one stable cation, while others can form many. For simplicity, we will call metals that form only one (invariant) charge state to be Type I and those that form variable charge states to be type 2. It is probably easiest to identify the Type 1 and consider others to be Type 2.

Type 1 Cations of Invariant Charge (Oxidation State)

Remember metals lose electrons to form cations and nonmetals gain electrons to form anions. Figure 2.7.2 lists the ions (cation and anion) that have invariant oxidation states. All the anions are of this type, gaining the number of electrons required to fill the valence shell and become isoelectronic with the nearest noble gas (have the same electron configuration). Many metals are of this type. Alkali metals (Group 1A) lose one electron to becomes isoelectronic to a noble gas. Alkaline earths (Group 2A) lose two electrons and many of the compounds , silver (group 1B) only forms [+1] cations. Alkaline earths lose two electrons (becoming isoelectronic to a noble gas) and many of the group IIIA only form cations [+3] charge.

There are some transition metals that form only one stable ion. Silver (Group 1B) forms a [+1] cation like the 1A alkali metals. Zinc and Cadmium (group 2B) form [+2] cations like the group 1B alkaline earth metals. The lighter Group 3A metals (Aluminum, Galium and Indium), along with Scandium and Yttrium lose 3 electrons to form [+3] cations. All the remaining transition metals form multiple charged ions (iron for example, forms Fe⁺² and Fe⁺³ and thus has multiple charge states).

Figure\(\PageIndex{2}\) The elements in the above table form only one stable charged state. Note, we often refer to the charge of a monatomic ion as its oxidation state.

Video\(\PageIndex{2}\): 3:10 min YouTube on nomenclature for binary ionic compounds (with type I metal cations).

Using the Principle of Charge Neutrality (section 2.6.6) and knowing the charge of the ions allows you to determine the formula.

Exercise \(\PageIndex{1}\)

What is the formula for Barium Chloride?

Answer: Barium is an alkaline earth and always corms a cation of charge of [+2], while chlorine is a halogen and always form the chloride ion of [-1]. For barium chloride to be neutral you would need two chlorides for every barium, and so the formula is BaCl₂.

Exercise \(\PageIndex{2}\)

Name the following compound: Al₂O_3.

Answer: Aluminum oxide. Al is a Type I monatomic ion, no Roman Numeral is needed since the charge does not change. Change the oxygen ending to the -ide ending for the anion.

Type II: Cations with Variable Charge (Oxidation State)

Most transition metals form multiple stable cations with different charges, and so you have to identify the charge state, which is done by writing the charge in Roman numerals and placing it in parenthesis after the name of the metal. So Fe⁺² is iron(II) and Fe⁺³ is iron(III). Some of the metals form very common ions which have latin names that are in common use, and you need to be familiar with those in the following table. Note, "-ic" stands for the higher oxidation state (charge), and "-ous" stands for the lower oxidation state.

Figure\(\PageIndex{3}\): Example of some compounds that have multiple oxidation states.

Note mercury(1) is not a monoatomic cation, but is really a hom*onuclear diatomic ion of two mercury atoms bound to each other, both having lost one electron. So Hg₂Cl₂ it the lowest whole number ratio of cation to anion.

Example 2

Ions:	Fe²⁺+ 2Cl^-	Fe³⁺+ 3Cl^-
Compound:	FeCl₂	FeCl₃
Nomenclature	Iron (II) Chloride	Iron (III) Chloride

-ous ending is used for the lower oxidation state
-ic ending is used for the higher oxidation state

Example 3

Compound	Cu₂O	CuO	FeCl₂	FeCl₃
Charge	Charge of copper is +1	Charge of copper is +2	Charge of iron is +2	Charge of iron is +3
Nomenclature	Cuprous Oxide	Cupric Oxide	FerrousChloride	FerricChloride

Some of the more common chemicals use the -ous/-ic nomenclature, but the use of Roman Numerals to designate the charge is acceptable.

Exercise \(\PageIndex{3}\)

Name the following compound: CoBr₃.

Answer: Cobalt(III) bromide. We know that Br has a -1 charge and the are three bromide ions. There is one cobalt ion. To counterbalance the -3 charge, the charge of the Co ions must be +3. So 1(x) + 3(-1) = 0. The "x" must be +3.

Polyatomic Ions

Polyatomic ions are covalent units (molecules) where the total number of protons \(\neq\) to the total number of electrons and they were introduced in section 2.6.4.2. They end in [-ide], [-ate] and [-ite]. Many of these are oxyanions with oxygen being bonded to a nonmetal and others are carboxylate ions. It should be noted that ionic comounds like ammonium acetate are composed entirely of nonmetals, but they form a crystal structure of cations and anions.

Figure \(\PageIndex{4}\): List of Polyatomic ions freshmen chemistry students are required to know.

Molecular Compounds: Nonmetals and Nonmetals

Compounds that consist of a nonmetal covalently bonded to a nonmetal are commonly known as Molecular Compounds, where the element with the positive oxidation state is written first. In many cases, nonmetals form more than one binary compound, so prefixes are used to distinguish them.

# of Atoms	1	2	3	4	5	6	7	8	9	10
Prefixes	Mono-	Di-	Tri-	Tetra-	Penta-	Hexa-	Hepta-	Octa-	Nona-	Deca-

Example 4

CO = carbon monoxide BCl₃ = borontrichloride

CO₂ = carbon dioxide N₂O₅ =dinitrogen pentoxide

The prefix mono- is not used for the first element. If there is not a prefix before the first element, it is assumed that there is only one atom of that element.

Exercise \(\PageIndex{5}\)

Laughing gas is commonly named as nitrous oxide, N₂O. What is the systematic name for this compound?

Answer: Dinitrogen monoxide. First, we must establish this a molecular compound before we use prefixes. Remember, prefixes are NOT used for ionic compounds.

Exercise \(\PageIndex{6}\)

Name the following compound: NO₃

Answer: Nitrogen trioxide. BE AWARE: This is not the the nitrate ion. If it were, it would a negative charge. It a molecular compound and named as such.

Binary Acids

Although HF can be named hydrogen fluoride, it is given a different name for emphasis that it is an acid. An acid is a substance that dissociates into hydrogen ions (H⁺) and anions in water. A quick way to identify acids is to see if there is an H (denoting hydrogen) in front of the molecular formula of the compound. To name acids, the prefix hydro- is placed in front of the nonmetal modified to end with –ic. The state of acids is aqueous (aq) because acids are found in water.

Some common binary acids include:

HF (g) = hydrogen fluoride -> HF (aq) = hydrofluoric acid

HBr (g) = hydrogen bromide -> HBr (aq) = hydrobromic acid

HCl (g) = hydrogen chloride -> HCl (aq) = hydrochloric acid

H₂S (g) = hydrogen sulfide -> H₂S (aq) = hydrosulfuricacid

It is very important to include (aq) after the acids because the same compounds can be written in gas phase with hydrogen named first followed by the anion ending with –ide.

Example 5

hypo____ite ____ite ____ate per____ate

ClO^- ClO₂^- ClO₃^- ClO₄^-

hypochlorite chlorite chlorate perchlorate

---------------->

As indicated by the arrow, moving to the right, the following trends occur:

Increasing number of oxygen atoms

Increasing oxidation state of the nonmetal

(Usage of this example can be seen from the set of compounds containing Cl and O)

This occurs because the number of oxygen atoms are increasing from hypochlorite to perchlorate, yet the overall charge of the polyatomic ion is still -1. To correctly specify how many oxygen atoms are in the ion, prefixes and suffixes are again used.

Table: Common Polyatomic ion the topic of acids and polyatomic ions, there is nomenclature of aqueous acids. Such acids include sulfuric acid (H₂SO₄) or carbonic acid (H₂CO₃). To name them, follow these quick, simple rules:

If the ion ends in -ate and is added with an acid, the acid name will have an -ic ending. Examples: nitrate ion (NO₃^-) + H⁺(denoting formation of acid) = nitric acid (HNO₃)
If the ion ends in -ite and
is added with an acid, then the acid name will have an -ous ending. Example: nitite ion (NO₂^-) + H⁺ (denoting formation of acid) = nitrous acid (HNO₂)

Naming Acid Salts

Video 2.7.e: (1'28") Acid Salts 1: Youtube for naming compounds that can form from carboxylate (acid, salt and acid salt)

Video 2.7.f: (3'19") Acid Salts 2: YouTube naming the compounds that can form from phosphate (acid, salt, and acid salts)

References

Pettrucci, Ralph H. General Chemistry: Principles and Modern Applications. 9th. Upper Saddle River: Pearson Prentice Hall, 2007
Nomenclature of Inorganic Chemistry, Recommendations 1990, Oxford:Blackwell Scientific Publications. (1990)
International Union of Pure and Applied Chemistry (2005). Nomenclature of Inorganic Chemistry (IUPAC Recommendations 2005). Cambridge (UK): RSC–IUPAC. ISBN 0-85404-438-8. Electronic version..
Biochemical Nomenclature and Related Documents, London:Portland Press, 1992.

Problems

1. What is the correct formula for Calcium Carbonate?

a. Ca⁺ + CO₂^-

b. CaCO₂^-

c. CaCO₃

d. 2CaCO₃

2. What is the correct name for FeO?

a. Iron oxide

b. Iron dioxide

c. Iron(III) oxide

d. Iron(II) oxide

3. What is the correct name for Al(NO₃)₃?

a. Aluminum nitrate

b. Aluminum(III) nitrate

c. Aluminum nitrite

d. Aluminum nitrogen trioxide

4. What is the correct formula of phosphorus trichloride?

a. P₂Cl₂

b. PCl₃

c. PCl₄

d. P₄Cl₂

5. What is the correct formula of lithium perchlorate?

a. Li₂ClO₄

b. LiClO₂

c. LiClO

d. None of these

6. Write the correct name for these compounds.

a. BeC₂O₄:

b. NH₄MnO₄:

c. CoS₂O₃:

7. What is W(HSO₄)₅?

8. How do you write diphosphorus trioxide?

9. What is H₃P?

10. By adding oxygens to the molecule in number 9, we now have H₃PO₄? What is the name of this molecule?

Answer

1.C; Calcium + Carbonate --> Ca²⁺ + CO₃²^- --> CaCO₃

2.D; FeO --> Fe + O^2- --> Iron must have a charge of +2 to make a neutral compound --> Fe²⁺ + O^2- --> Iron(II) Oxide

3.A; Al(NO₃)₃ --> Al³⁺ + (NO₃^-)₃ --> Aluminum nitrate

4.B; Phosphorus trichloride --> P + 3Cl --> PCl₃

5.D, LiClO₄; Lithium perchlorate --> Li⁺ + ClO₄^- --> LiClO₄

6. a. Beryllium Oxalate; BeC₂O₄ --> Be²⁺ + C₂O₄²^- --> Beryllium Oxalate

b. Ammonium Permanganate; NH₄MnO₄ --> NH₄⁺ + MnO₄^- --> Ammonium Permanganate

c. Cobalt (II) Thiosulfate; CoS₂O₃ --> Co + S₂O₃²^- --> Cobalt must have +2 charge to make a neutral compund --> Co²⁺ + S₂O₃²^- --> Cobalt(II) Thiosulfate

7. Tungsten (V) hydrogen sulfate

8. P₂O₃

9. Hydrophosphoric Acid

10. Phosphoric Acid

Contributors

Pui Yan Ho (UCD), Alex Moskaluk (UCD), Emily Nguyen (UCD)
Ronia Kattoum (UA of Little Rock)

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Practice sheets

1402 Nomenclature Problem set 1 (Ionic and Covalent Compounds)

1402 Nomenclature Problem set 2 (Acids)

Key to Practice Set 1:

1)aluminum nitride AlN

2)aluminum carbonate Al₂(CO₃)₃

3)aluminum bicarbonate Al(HCO₃)₃

4)nitrogen dioxide NO₂

5)zinc sulfate ZnSO₄

6)zinc sulfide ZnS

7)zinc phosphate Zn₃(PO₄)₂

8)carbon tetrachloride CCl₄

9)sodium hydroxide NaOH

10)barium hydroxide Ba(OH)₂

11)titanium (II) cyanide Ti(CN)₂

12)cadmium iodide CdI₂

13)cadmium acetate Cd(CH₃CO₂)₂

14)cadmium bromide CdBr₂

15)cadmium nitrate Cd(NO₃)₂

16)cadmium sulfide CdS

17)phosphorus pentachloride PCl₅

18)iron (II) sulfide FeS

19)iron (II) sulfate FeSO₄

20)iron (III) sulfate Fe₂(SO₄)₃

21) iron (III) nitrate Fe(NO₃)₃

22)magnesium nitride Mg₃N₂

23)magnesium nitrate Mg(NO₃)₂

24)diboron hexafluoride B₂F₆

25)tin (IV) fluoride SnF₄

26)tin (II) fluoride SnF₂

27)tin (IV) sulfide SnS₂

28)tin (II) sulfide SnS

29)nitrogen trichloride NCl₃

30)manganese (II) sulfide MnS

31)manganese (II) sulfate MnSO₄

32)manganese (II) phosphate Mn₃(PO₄)₂

manganese (II) nitrate Mn(NO₃)₂

34)vanadium (V) oxide V₂O₅

35)molybdenum (VI) sulfide MoS₃

36)silver phosphide Ag₃P

37)mercury (II) oxide HgO

38)chromium (III) bromide CrBr₃

39)chromium (III) nitrate Cr(NO₃)₃

40)sulfur dichloride SCl₂

41)magnesium sulfide MgS

42)magnesium sulfate MgSO₄

43)cobalt (III) carbonate Co₂(CO₃)₃

44)cobalt (III) phosphide CoP

Write formulas from the following names:

1)Ni(NO₃)₃ nickel (III) nitrate

2)NiCl₃ nickel (III) chloride

3)Mn(OH)₂ manganese (II) hydroxide

4)Hg(CN)₂ mercury (II) cyanide

5)CuS copper (II) sulfide

6)PCl₃ phosphorus trichloride

7)Sr(NO₃)₂ strontium nitrate

8)FeSO₄ iron (II) sulfate

9)Bi(OH)₃ bismuth (III) hydroxide

10)AgHCO₃ silver bicarbonate

11)Co₂(CO₃)₃ cobalt (III) carbonate

12)SF₂ sulfur difluoride

13)Zn(HCO₃)₂ zinc bicarbonate

14)NCl ammonium chloride

15)(NH₄)₂SO₄ ammonium sulfate

16)NiPO₄ nickel (III) phosphate

17)Co(CN)₂ cobalt (II) cyanide

18)NCl₃ nitrogen trichloride

19)Hg₂S mercury (I) sulfide

20)Hg₃(PO₄)₂ mercury (II) phosphate

21) Mn(NO₃)₂ manganese (II) nitrate

22)Fe₂(SO₄)₃ iron (III) sulfate

23)Sn(CN)₂ tin (II) cyanide

24)ZnSO₄ zinc sulfate

25) LiNO₃ lithium nitrate

26) LiCN lithium cyanide

27) Ba(CN)₂ barium cyanide

28) Al(CN)₃ aluminum cyanide

29) CuCN copper (I) cyanide

30) Cu(CN)₂ copper (II) cyanide

31) CuSO₄ copper (II) sulfate

32) Cu(NO₂)₂ copper (II) nitrite

33) Cu(NO₃)₂ copper (II) nitrate

34) CuN copper (III) nitride

35) CuBr copper (I) bromide

36) CuBr₂ copper (II) bromide

37) Cu₂SO₄ copper(I) sulfate

38) Cu₂(SO₄)₃ copper (III) sulfate

39) Cu₃PO₄copper (I) phosphate

40) Cu₃(PO₄)₂ copper (II) phosphate

41) PI₅ phosphorous pentaiodide

42) PbS lead (II) sulfide

43) PbS₂ lead (IV) sulfide

44) SnCl₄ tin (IV) chloride

45) SnBr₂ tin (II) bromide

46) SnSO₄ tin (II) sulfate

47) Sn₃(PO₄)₂ tin (II) phophate

48) SF₆ sulfur hexafluoride

49) P₄H₁₀ tetraphosphorus decahydride

50) CuS copper (II) sulfide

51)HNO₃ nitric acid

52)H₂S hydrosulfuric acid

53)H₃PO₄ phosphoric acid

54)HBrO₃ bromic acid

55)H₂SO₄_Þsulfuric acid

56)HIO hypoiodic acid

57)H₂SO₃ sulfurous acid

58)HClO₄ perchloric acid

59)HNO₂ nitrous acid

60) H₃PO₃ phosphorus acid

61) H₃AsO₄ arsenic acid

62) HIO₂ iodous acid

63) H₂C₂O ₄ oxalic acid

64) Arsenic acid H₃AsO₄

65) Acetic acid C₂H₄O₂

66) hydrocyanic acid HCN

67) periodic acid HIO₄

68) bromous acid HBrO₂

69) hypoiodous iodic acid HIO

70) hypochlorous acid HClO

71) carbonic acid H₂CO₃

72) phosphoric acid H₃PO₄

73) arsenous acid H₃AsO₃

74) phosporous acid H₃PO₃

75)sulfurous acid H₂SO₃

76) hydrosulfuric acid H₂S

1402 Nomenclature Problem set 1 Key

1402 Nomenclature Problem set 2 Key

Key to Practice Set 2

HNO₃Þ nitric acid
H₂S Þ hydrosulfuric acid
H₃PO₄ Þ phosphoric acid
HBrO₃Þ bromic acid
H₂SO₄_Þsulfuric acid
HIO Þ hypoiodic acid
H₂SO₃Þ sulfurous acid
HClO₄ Þ perchloric acid
HNO₂Þ nitrous acid
H₃PO₃Þ phosphorus acid
H₃AsO₄Þ arsenic acid
HIO₂Þ iodous acid
H₂C₂O ₄Þ oxalic acid
Arsenic acid Þ H₃AsO₄
Acetic acid Þ C₂H₄O₂
hydrocyanic acid Þ HCN
periodic acid Þ HIO₃
bromous acid Þ HBrO₂
hypoiodous iodic acid Þ HIO
hypochlorous acid Þ HClO
carbonic acid Þ H₂CO₃
phosphoric acid Þ H₃PO₄
arsenous acid Þ H₃AsO₃
phosporous acid Þ H₃PO4
sulfurous acid Þ H₂SO₃
hydrosulfuric acid Þ H₂S

Contributors

Robert E. Belford (University of Arkansas Little Rock; Department of Chemistry). The breadth, depth and veracity of this work is the responsibility of Robert E. Belford, rebelford@ualr.edu. You should contact him if you have any concerns. This material has bothoriginal contributions, and contentbuilt upon prior contributions of the LibreTexts Community and other resources,including but not limited to:

Modified by Ronia Kattoum (UA of Little Rock)