Compuestos inorganicos

1. NOMENCLATURE
OF INORGANIC
COMPOUNDS
Chapter VI

2. CHEMICAL FORMULAS
 Chemical formulas show the number of atoms present in each
compound.
 AlCl3: One aluminum atom; three chlorine atoms
 Fe2(SO4)3: Two iron atoms; three sulfur atoms; twelve oxygen
atoms
 KCl: One potassium atom; one chlorine atom
 Chemical formulas of ionic compounds represent the minimum
number of atoms necessary to produce a neutral compound
(formula unit):
 Ca2O2 is not a formula unit. The minimum number of atoms
required to form this compound would be CaO.

3. OXIDATION NUMBER
 Oxidation Losing electrons
 Sodium atoms (Na) oxidize to form sodium cations (Na+)
 Fe2+ oxidizes to Fe3+
 Reduction Gaining electrons
 Fluorine atoms (F) reduce to form fluoride anions (F-)
 Ag+ is reduced to Ag when it accepts one electron
 In ionic compounds, the oxidation number is the charge of
each one of the ions present in the formula.
 The oxidation number of aluminum in the aluminum cation
(Al3+) is +3
 The oxidation number of sulfur in the sulfide anion (S2-) is -2

4. OXIDATION NUMBER
 Main group elements usually have a fixed oxidation number:
 Group 1A: +1
 Group 2A: +2
 Group 3A: +3
 Group 4A: +4/-4
 Group 5A: -3
 Group 6A: -2
 Group 7A: -1
 Under special circumstances, some of these elements may have
other oxidation numbers. This is especially true for non-metals.

5. OXIDATION NUMBER
Group
Oxidation number
Frequency
1A
+1
Always
2A
+2
Always
3A
+3
Always
+4 and -4
Most common
+2 and -2
Rare
-3
Most common
+1, +2, +3, +4 or +5
Rare
-2
Most common
-1, +1, +2, +4, +6
Rare
-1
Most common
+1, +3, +5, +7
Rare
4A
5A
6A
7A

6. STOCK NOMENCLATURE
 Elements of the d-block and the f-block usually have a wide variety of oxidation
numbers and all of them are commonly found in nature.
 There are two iron oxides: FeO and Fe2O3
 There are two copper oxides: CuO and Cu2O
 There are three chromium chlorides: CrCl2, CrCl3, CrCl4
 For ions of elements that have multiple oxidation numbers, the stock system of
nomenclature is used to name them.
 In the stock system, the oxidation number of the ion is written in parenthesis using
roman numerals in front of the element’s name: element(oxidation number)
 Fe2+  iron(II)
 Fe3+  iron(III)
 Cu+ --> copper(I)
 Mn7+  manganese(VII)
 Cr6+  chromium(VI)
 For ions of elements with characteristic oxidation numbers (alkali metals, alkaline
earth metals, etc.) there is no need to write their oxidation number in parenthesis
when naming them. (i.e. Na+ is called sodium)

7. STOCK NOMENCLATURE
Elements whose oxidation numbers don’t need to be stated in their stock name
Element(s)
Oxidation number
Group IA
+1
Group IIA
+2
Group IIIB
+3
Al, Ga, In
+3
Zn, Cd
+2
Ag
+1

8. DETERMINING THE OXIDATION
NUMBER OF A VARIABLE ELEMENT
1.
Look for atoms in the formula that possess a constant oxidation
number. (H and alkali metals +1; Alkaline earth metals +2;
earth metals +3; oxygen -2). After this, identify the presence
of elements with a variable oxidation number.
1.
2.
K2Cr2O7 (constant: K and O; variable: Cr)
Write the charge of the known elements at the top of their
positions in the formula unit.
+1
-2
K2Cr2O7

9. DETERMINING THE OXIDATION
NUMBER OF A VARIABLE ELEMENT
3.
Multiply the charge of each of the known elements times the
number of atoms of that element in the formula. Write the
resulting number under the symbol of the element. DON’T
FORGET THE SIGN!!
+1
-2
K2Cr2O7
+2
-14
-2x7

10. DETERMINING THE OXIDATION
NUMBER OF A VARIABLE ELEMENT
4.
Calculate the charge necessary so that the total sum of the
charges is equal to zero (neutral). You may do this mentally or
using algebra
4.
Using algebra: +2+x-14=0
x-12=0
x=+12
5.
Divide the resulting charge by the number of atoms present in
the formula of the variable element. The resulting charge is the
oxidation number of the variable element.
+1 +6 -2
K2Cr2O7
+2+12-14
Divide +12/2= +6

11. WRITING CHEMICAL FORMULAS
 When writing the formulas of ionic compounds, the elements are
written just as we write in English (from left to right)
 Cations always appear to the left of the formula
 Anions always appear to the right of the formula
 Al2O3
 AgBr
 CuCl2
 K2S
 The numbers written as a subscript in front of the symbol of the
element are called stoichiometric numbers. These numbers
represent the number of atoms of that element present in the
formula.

12. FORMULAS WITH POLYATOMIC IONS
 An ion is a charged chemical species. As such, it is not limited to
single atoms.
 Ions may be formed by more than one atom. These are called
polyatomic ions. Another name for them is molecular
ions, because the atoms in a polyatomic ion are bonded together
by covalent bonds.

13. FORMULAS WITH POLYATOMIC IONS
 Polyatomic ions are treated as a whole; as an entity.
 When writing chemical formulas that include polyatomic ions; if
more than one of such ions are needed to ensure neutrality, they
must be written in parenthesis followed by the stoichiometric
number to the right of the parenthesis.
 Fe3+ and SO42-

 Ba2+ and ClO3- 
Fe2(SO4)3
Ba(ClO3)2
 Na+ and NO3- 
NaNO3 (If only one is needed, the
parenthesis is not necessary)
 NH4+ and S2- 
(NH4)2S

14. IUPAC NOMENCLATURE
 The International Union of Pure and Applied Chemistry (IUPAC) is
the organization in charge of standardizing chemical terminology
and nomenclature.
 The IUPAC has set up some rules to name compounds according
to their chemical function and composition.
 Types of chemical compounds:
 Binary: Formed by two different elements (NaI, HF, CaO, Al2S3)
 Ternary: Formed by three different elements (NaOH, BaSO4, HClO3)
 Quaternary: Formed by four different elements (NaHCO3)

15. IUPAC NOMENCLATURE OF
BINARY COMPOUNDS

16. METAL OXIDES
 Metal oxides are formed by a metal cation and the oxide anion
(O2-)
 Nomenclature: “stock nomenclature of the metal ion” +
“oxide”
 Na2O  sodium oxide
 MgO  magnesium oxide
 Fe2O3  iron(III) oxide
 TiO2  titanium(IV) oxide

17. DETERMINING THE FORMULA FROM
THE NAME OF THE COMPOUND
1.
2.
3.
4.
5.
6.
Write the symbol of the elements that must be present in the
formula according to the name of the compound. Cations are
written first, then anions.
From the name of the compound, deduce the charge of each ion.
When metals have no explicit oxidation number in the name of the
compound, it’s because it’s an element with a common or constant
oxidation number. The oxidation number of variable elements is
always stated in the name using stock nomenclature.
Write the charges of each ion above their symbol.
Cross the oxidation numbers of each ion and place them as the
stoichiometric number with NO SIGN, JUST THE NUMBER.
Simplify those numbers until you have the smallest ratio.
Write the formula unit of the compound.

18. EXAMPLE
1.
2.
3.
4.
5.
6.
What is the formula of lead(II) oxide? Elements that must be in
the formula: Pb and O
The name implies this compound is a metal oxide because it
only has oxygen and lead, which is a metal. The charge of the
oxide anion is always -2; the charge of lead is revealed by it’s
stock name (+2).
Pb2+ O2Pb2+ O2-  Pb2O2
As it stands, that formula is not the formula unit. The 2:2 ratio
can be simplified to 1:1.
Formula: PbO

19. NON-METAL OXIDES
 Non-metal oxides arise from the combination of a non-metal and
oxygen.
 These compounds are NOT IONIC, they are molecular
compounds. As a consequence, their formulas are not called
formula units, they are called “molecular formulas”.
 Molecular formulas don’t have to show the minimal relationship
between the elements. (i.e. NO2 and N2O4 are two different
chemical substances and both of them exist).

20. NON-METAL OXIDES
 Non-metal oxides are named the following way:
“prefix+non-metal”+ “prefix+oxide”
 The prefix indicates how many atoms of each element are present in the
formula.
 The prefixes refer to Greek numbering prefixes:
 1 Mono
 2 Bi- or di 3 Tri 4 Tetra 5 Penta 6 Hexa 7 Hepta 8 Octa 9 Nona 10 Deca-

21. NON-METAL OXIDES
 When there is only one atom of the non-metal in the formula, the “mono-”
prefix is omitted. The “mono-” prefix is only used for oxygen.
 In the formula, the non-metal is written first, then the oxygen.
 Examples:
 NO: nitrogen monoxide
 CO2: carbon dioxide
 SO3: sulfur trioxide
 N2O: dinitrogen monoxide
 Cl2O7: dichlorine heptaoxide (also written as dichlorine heptoxide)
 Chemical formulas are very easy to deduce from the name of the
compounds because the number of atoms of each element is explicitly
stated.

22. BINARY SALTS
 In chemistry, salt is a general term used to describe ionic
compounds resulting from a neutralization reaction. A
neutralization reaction is one that occurs between an acid and a
base.
 Binary salts contain a metal cation and a non-metal anion
(other than the oxide anion)
 Examples:
 NaCl
 KI
 CaS
 AlCl3

23. BINARY SALTS
 Nomenclature: “stock system name of the metal cation” +
“root of the name of the non-metal + suffix –ide”
 Examples of roots for non-metals:
 Chlorine  “chlor” + -ide  chloride
 Bromine  “brom” + -ide  bromide
 Sulfur  “sulf” + -ide  sulfide
 Selenium  “selen” + -ide  selenide
 Phoshphorus  “phosph” + -ide  phosphide
 Nitrogen  “nitr” + -ide  nitride

24. BINARY SALTS
 Examples:
 RbCl  rubidium chloride
 SrI2  strontium iodide
 Ag2S  silver sulfide
 CdSe  cadmium selenide
 LiF  lithium fluoride
 AlP  aluminum phosphide
 CuBr  copper(I) bromide

25. BINARY SALTS
 In binary salts, the non-metal always has the characteristic
oxidation number of its group.
 Group 14 anions  -4
 Group 15 anions  -3
 Group 16 anions  -2
 Group 17 anions  -1

26. METAL HYDRIDES
 Metal hydrides are ionic compounds formed by the hydride anion
(H-) and a metal cation.
 They are one of the few types of compounds in which hydrogen
has a charge of -1, instead of the usual +1.
 Ionic metal hydrides are only formed with metals of Group 1 and
2. The rest of the metals form hydrides with a covalent nature.
The nomenclature of such hydrides will not be discussed.
 Their nomenclature is similar to that of salts: “stock system
name of the metal cation” + “hydride”
 NaH  sodium hydride
 CaH2  calcium hydride
 BeH2  beryllium hydride

27. NON-METAL HALIDES
 Halide is a term used to refer to “salts” where the anion is one of
the halogens.
 Non-metal halides, however, are MOLECULAR.
 They contain one non-metal, other than hydrogen, bonded to
halogens.
 The nomenclature of these compounds is similar to that of non-
metal oxides: “numeric prefix+non-metal” + “numeric prefix +
root of the name of the halogen + suffix –ide”
 CCl4  carbon tetrachloride
 PBr3  phosphorus triiodide
 SF4  sulfur tetrafluoride

28. BINARY ACIDS
 Also known as hydracids.
 They are formed by a non-metal and hydrogen.
 They are MOLECULAR compounds when in their pure state.
 Binary acids are usually soluble in water, and when they
dissolve, their covalent bonds suffer electrolytic dissociation (they
separate into ions). The aqueous solution of binary acids is acidic
because of the presence of the hydrogen cation (H+).
 Given that hydracids may exist as molecular compounds or
aqueous solutions; a nomenclature for each case has been
developed.

29. BINARY ACIDS
 When binary acids exist as molecular compounds their
nomenclature is similar to the one of salts: “hydrogen” + “root
of the name of the non-metal + suffix –ide”
 Examples
 HCl  hydrogen chloride
 H2S  hydrogen sulfide
 HI  hydrogen iodide
 H2Te  hydrogen telluride

30. BINARY ACIDS
 When binary acids are present in aqueous solution: “hydro- +
root of the name of the non-metal + suffix –ic” + “acid”
 When binary acids are in aqueous solution, it is indicated in the
formula with the (aq) subscript.
 Examples:
 HCl(aq)  hydrochloric acid
 H2S(aq)  hydrosulfuric acid
 HI(aq)  hydroiodic acid
 H2Te(aq)  hydrotelluric acid

31. IUPAC NOMENCLATURE OF
TERNARY COMPOUNDS

32. METAL HYDROXIDES
 Metal hydroxides contain the polyatomic anion (OH)- called the
hydroxide ion.
 The oxygen and the hydrogen in the hydroxide ion are bonded
through a covalent bond. The oxygen possesses a negative
charge. Metal hydroxides are named as follows: “stock system
name of the metal cation” + “hydroxide”
 Examples:
 NaOH  sodium hydroxide
 Mg(OH)2  magnesium hydroxide
 Mn(OH)2  manganese(II) hydroxide
 Au(OH)3  gold(III) hydroxide

33. OXOACIDS
 Oxoacids are composed of hydrogen, a non-metal, and
oxygen.
 They are MOLECULAR compounds.
 Like other acids, oxoacids dissociate in aqueous solution that are
acidic due to the presence of the hydrogen cation.
 The nomenclature of oxoacids is very particular because each
non-metal may produce more than one oxoacid. This is
determined by the amount of oxygen present in the formula.
 For example, chlorine produces four different oxoacids: HClO,
HClO2, HClO3, and HClO4.

34. OXOACIDS
 Nomenclature: “prefix- + root of the name of the non-metal +
suffix” + “acid”
 For oxoacids, the prefixes and suffixed denote the relative
amount of oxygen present in the formula. These prefixes and
suffixes are not the Greek ones denoting numbers, they are:
Meaning
Prefix
Suffix
Lowest
Hypo-
-ous
Low
-ous
High
-ic
Highest
Per-
-ic

35. OXOACIDS
 Given that the prefixes and suffixes are related to something
relative, rules have been developed to assign the correct affix to
each compound. This is done by determining the oxidation
number of the non-metal
Group IVA
Group VIA
Group VIIA
Nomenclature
+1
+2
+1
Hypo___ous
acid
+3
+4
Group VA
+4
+3
____ous acid
+5
+6
+5
_____ic acid
+7
Per____ic acid

36. OXOACIDS
 The oxidation number of the non-metal can be determined using
the usual method by assuming the compound is ionic.
 Example: Determine the oxidation number of chlorine in HClO3
+1 +5 -2
(+1)(1)
HClO3
(-2)(3)
+1 +5 -6 =0
+5/1
Since chlorine belongs to group VIIA and has an oxidation number of
+5 , this compound is called chloric acid

37. OXYSALTS
 Oxysalts are IONIC compounds derived from the neutralization of
an oxoacid and a base.
 Oxysalts contain a metal cation and an oxyanion, which is the
polyatomic anion derived from the oxoacid.
 HClO3  ClO3
–
2-
 H2SO4  SO4
 The charge of the polyatomic anion can be deduced by
determining the oxidation number of the elements and adding the
total charge (it will be less than zero in this case). An easier way
to do it is by looking at how many hydrogen atoms were removed
from the original oxoacid. The number of hydrogen atoms
removed is the number of negative charges the anion has.

38. OXYSALTS
 The nomenclature of oxysalts is similar to that of binary salts:
“stock system name of the metal cation” + “name of the
oxyanion”
 The name of the oxyanion uses similar affixes than the ones used
for oxoacids and depend of the oxidation number of the nonmetal.
Group IVA
Group VIA
Group VIIA
Nomenclature
+1
+2
+1
“Metal”
Hypo___ite
+3
+4
Group VA
+4
+3
“Metal” ____ite
+5
+6
+5
“Metal”
_____ate
+7
“Metal”
Per____ate

39. OXYSALTS
 Examples:
 NaClO  sodium hypochlorite
 K2CO3  potassium carbonate
 Mn(NO3)2  manganese(II) nitrate
 Al2(SO3)3  aluminum sulfite
 Ba(ClO4)2  barium perchlorate

40. OTHER POLYATOMIC IONS
 Cyanide anion (CN)
-
 Permanganate anion (MnO4)
 Chromate anion (CrO4)
-
2-
 Dichromate anion (Cr2O7)2 Azide anion (N3)
-
 The compounds formed by these ions with metal ions or
hydrogen are named just like any other salt.
 NaN3  sodium azide
 KMnO4  potassium permanganate
 HCN  hydrogen cyanide

41. OTHER POLYATOMIC IONS
+
 The ammonium ion (NH4) is a cation derived from ammonia. It
forms salts with practically all anions. These salts are named like
any other salt.
 (NH4)2S  ammonium sulfide
 (NH4)2CO3  ammonium carbonate
 NH4ClO4  ammonium perchlorate

42. OTHER POLYATOMIC IONS
 Mercury has two oxidation numbers, +1 and +2. When in the +2
state, it exists as a monoatomic cation as usual (Hg2+). However,
when it is in its +1 oxidation state, it exists as a polyatomic cation
(Hg2)2+. In this cation there is a covalent bond between the
mercury atoms, and each one is has an oxidation number of +1.
The name of this cation is mercury(I).
 Hg2Cl2  Mercury(I) chloride
 Hg2O  Mercury (I) oxide