2.7: Nomenclature of Ionic, Covalent, and Acid Compounds

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 Ca^{2+ +} 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 homonuclear 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:	Fe^2^++ 2Cl^-	Fe^3^++ 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.

 

Naming Oxyanions: Oxyanions are polyatomic ions where oxygen is attached to a nonmetal and as was discussed in section 2.6.4.2.1, nonmetals of the same periodic group from homologous oxyanions.  Lets look at the 4 oxyanions of bromine

• perbromate   (BrO₄^-)
• bromate        (BrO₃^-)
• bromite         (BrO₂^-​​​​​​​)
• hypobromite  (BrO^-​​​​​​​)

These all end with -ate or -ite, and the real reasoning deals with the oxidation state of the nonmetal, sort of its hypothetical charge within the covalently bonded polyatomic ion.  We will cover that in a later chapter, but right now we note that the "-ate" ions have more oxygens and the "-ites" have less, with "per____ate" having the most, and "hypo___ite" having the least.

of the same family of the periodic table 

 

Note, members of the same family tend to form similar compounds, so bromine and iodine form similar anions to chlorine (see figure 2.7.6), Selenium and Tellurium form similar anions to Sulfur, and Arsenic forms similar ones to Phosphorous. See video 2.7.3 below.

Note: The second period of the periodic table have nonmetals that behave different than the rest. Nitrogen forms different oxyanions than phosphorous or Arsenic, Oxygen does not form oxyanions, and although I have seen perfluorate, fluorate and fluorite salts on an exam and webpages where they form similar structures to chlorine, I believe the only one that really exists is hypofluorite (FlO^-).

 

 

So why is the second row an exception? We will understand later this semester that as you go down a family or group of the periodic table, the volume of the atoms increases. The reason the second row nonmetals are an exception can best be understood by their small size (how are you going to get 4 oxygens around a small fluorine)? But a search of the web will show you perfluorate, fluorate and fluorite, as if they form the same types of oxyanions as chlorine does, yet there is no perfluorite in a resource like pubchem, which as of 2018 has over 95 million chemical compounds. So on an exam at the freshmen level, I would probably treat them as if they follow the same trend as chlorine, because that is the schema you are learning.

Carboxylate ions: Another class of polyatomic anions are based on the carboxylate functional group of organic chemistry. The compounds phthlate, acetate and oxalate have this functional group. See video 2.7.4 below

Figure\(\PageIndex{4}\): Carboxylate functional group that is the bases for many organic ions.

Figure\(\PageIndex{6}\): Slide showing periodic trends for nomenclature of oxyanions of halogens.

Figure 2.7.4 Shows several types of polyatomic anions, name

Understanding Oxyanions

This video can help you understand oxyanions

Video\(\PageIndex{3}\): 2:20 minute Youtube to help you remember names of oxyanions

Polyatomic Ions based on carboxylate functional group

Organic compounds have functional groups, and many organic anions are based on the carboxylate group. This video helps you use that to remember the formula of polyatomic ions like acetate, phthlate and oxylate.

Video\(\PageIndex{4}\): YouTube to help you remember names of anions containing carboxylate functional groups

 

 

 

 

Exercise \(\PageIndex{4}\)

Name the following compounds:

A) NiCO₃ 

B) Sr(NO₂)₂ 

C) NH₄NO₃

Answer

A) Nickel(II) carbonate. Nickel is a Type II cation, so you need a Roman Numeral. To figure out that charge "x", we need take a look at what we know. There is one carbonate ion at a charge of -2. There is one Nickel ion at unknown charge "x". 1(x) + 1(-2) = 0. To be charge neutral, x=+2. 

B) Strontium nitrite. Strontium is a type I cation and there is no need for a Roman Numeral. NO₂^- is the nitrite ion. 

C) Ammonium nitrate. both are polyatomic ions, so you simply state the name of cation followed by the name of the anion. 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

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:

1. 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)
2. If the ion ends in -ite and
3. 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

 

1. Pettrucci, Ralph H. General Chemistry: Principles and Modern Applications. 9th. Upper Saddle River: Pearson Prentice Hall, 2007
2. Nomenclature of Inorganic Chemistry, Recommendations 1990, Oxford:Blackwell Scientific Publications. (1990)
3. 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..
4. 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⁺ + CO₃^2- --> 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⁺ + C₂O₄^2- --> Beryllium Oxalate

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

c. Cobalt (II) Thiosulfate; CoS₂O₃ --> Co + S₂O₃^2- --> Cobalt must have +2 charge to make a neutral compund --> Co²⁺ + S₂O₃^2- --> 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) 

 

----

 

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

33. 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)H₂SO_{4 Þ}sulfuric 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

 

1. HNO₃ Þ nitric acid
2. H₂S Þ hydrosulfuric acid
3. H₃PO₄ Þ phosphoric acid
4. HBrO₃ Þ bromic acid
5. H₂SO_{4 Þ}sulfuric acid
6. HIO Þ hypoiodic acid
7. H₂SO₃ Þ sulfurous acid
8. HClO₄ Þ perchloric acid
9. HNO₂ Þ nitrous acid
10. H₃PO₃ Þ phosphorus acid
11. H₃AsO₄ Þ arsenic acid
12. HIO₂ Þ iodous acid
13. H₂C₂O ₄ Þ oxalic acid
14. Arsenic acid Þ H₃AsO₄
15. Acetic acid Þ C₂H₄O₂
16. hydrocyanic acid Þ HCN
17. periodic acid Þ HIO₃
18. bromous acid Þ HBrO₂
19. hypoiodous iodic acid Þ HIO
20. hypochlorous acid Þ HClO
21. carbonic acid Þ H₂CO₃
22. phosphoric acid Þ H₃PO₄
23. arsenous acid Þ H₃AsO₃
24. phosporous acid Þ H₃PO4
25. sulfurous acid Þ H₂SO₃
26. 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, [e-mail]. You should contact him if you have any concerns. This material has both original contributions, and content built upon prior contributions of the LibreTexts Community and other resources, including but not limited to:

Modified by Ronia Kattoum (UA of Little Rock)