E - 29
INORGANIC NOMENCLATURE
NGSS: HS: CC Patterns.
In E-27 “Atomic Structure I”, you learned that free atoms have a net charge of zero, meaning that the
number of protons and electrons are equal. You also learned that when an atom loses one or more electrons,
it becomes a positive monatomic ion (i.e. cation). This is a characteristic behavior of metals. When an atom
gains one or more electrons, it becomes a negative monatomic ion (i.e. anion). This is a characteristic of
non-metals.
As mentioned way back in E-14 “Molar Mass”, the Periodic Table is a lot more than just a fancy piece of
“artwork”. The Table is comprised of 18 vertical columns called groups. Each group has its own group
number. Some Periodic Tables simply number the groups 1-18 from left to right. Other Periodic Tables use
Roman numerals and an “A” for the Main Group elements and Roman numerals and a “B” for
Transition Elements. In E-28 “Oxidation Numbers” you learned that the elements in the Main Group on
the Periodic Table make monatomic ions with predictable oxidation numbers. Group IA (Group 1) ions
have a +1 oxidation number, Group IIA (Group 2) ions are +2, Group IIIA (Group 13) are +3, Group
IVA (Group 14) are + or - 4, Group VA (Group 15) are -3, Group VIA (Group 16) are -2, Group VIIA
(Group 17) are -1, and the Group VIIIA (Group 18) Noble Gases do not normally have an oxidation
number since they are already stable without a gain or loss of electrons. Monatomic ions from the Transition
Elements can have more than one oxidation number. You also learned that the prominent “zigzag line” or
“jagged staircase” running diagonally down the Periodic Table separates the metals (i.e. elements which
form positive monatomic ions) on the left from the non-metals (i.e. elements which form negative
monatomic ions) on the right. Finally, you learned that sometimes groups of atoms can become a charged
polyatomic ion which can be either positive or negative.
In E-25 “Electrostatics”, you learned that opposite charges attract. Because of that interaction between
oppositely charged particles, positively charged ions and negatively charged ions attract each other and
form ionic compounds with a net charge of zero. Using your “Oxidation Sheet”, you will be able to name
and write the formulas for ionic compounds by combining one cation with one anion as shown below.
Formula Writing for Ionic Compounds
There are just a few simple steps to follow for writing the formulas for ionic compounds:
1. Write the symbol for the positive ion first (on the left) followed by the symbol for the negative ion (on
the right).
2. Use subscripts to balance the oxidation numbers of the positive and negative ions so that the net ionic
charge is zero.
3. When using subscripts with any polyatomic ion, put parenthesis around that ion first and then add the
appropriate subscript.
Example 1: Write the formula for sodium chloride.
1. The symbol for sodium Na. The symbol for chloride is Cl. So, you would write NaCl.
2. Since the charge of Na is +1 and Cl is -1, the net charge is zero (+1-1=0) so no subscript is needed.
Correct answer: NaCl
Example 2: Write the formula for calcium bromide.
1. The symbol for calcium is Ca and the symbol for bromide is Br. So, you would write CaBr.
2. Since Ca is +2 and Br is -1, the net charge is not zero ((+2) + (-1)  +1) so a subscript will be needed.
If we have two Br ions instead of one, the net charge will be zero ((+2) + 2(-1) = 0), so we need to add a
two subscript with the Br making it Br2.
Correct answer: CaBr2
Example 3: Write the formula for iron (III) oxide.
1. The symbol for iron is Fe and the symbol for oxide is O. So, you would write FeO.
2. Since Iron (III) is +3 and O is -2, the net charge is not zero ((+3) + (-2)  +1) so a subscript will be
needed. If we have two Fe ions and three O ions, the net charge will be zero (2(+3) + 3(-2) = 0), so we
need a two subscript with the Fe making it Fe2 and a three subscript with the O making it O3 .
Correct answer: Fe2O3
Example 4: Write the formula for lithium carbonate.
1. The symbol for lithium is Li and the symbol for carbonate is CO3 . So, you would write LiCO3 .
2. Since Li is +1 and CO3 is -2, the net charge is not zero ((+1) + (-2)  -1) so a subscript will be needed.
If we have two Li ions instead of one, the net charge will be zero (2(+1) + (-2) = 0), so we need to add a
two subscript with the Li making it Li2 .
Correct answer: Li2CO3
Example 5: Write the formula for barium nitrite.
1. The symbol for barium is Ba and the symbol for nitrite is NO2 . So, you would write BaNO2 .
2. Since Ba is +2 and NO2 is -1, the net charge is not zero ((+2) + (-1)  +1) so a subscript will be
needed. If we have two NO2 ions instead of one, the net charge will be zero ((+2) + 2(-1) = 0), so we
need to add a two subscript with the NO2 .
3. Since we are using a subscript with the polyatomic ion NO2, we must make sure to put parenthesis
around the NO2 first and then put the two subscript making it (NO2)2 .
Correct answer: Ba(NO2)2
Naming Ionic Compounds
Naming ionic compounds is very simple! Use the “Oxidation Sheet” to name the positive ion first
followed by the negative ion. Don’t forget to include the Roman numeral for the metal ions of
Transition Elements (ex. Fe, Cu, Hg etc.)! Since transition metals like iron and copper may have more
than one oxidation number, you need to mathematically determine which charge the metal has, and insert the
appropriate Roman numeral.
Note that monatomic metal ions have the same names as metal atoms. However, monatomic non-metal
atoms drop their ending and add -ide when they become ions! Compounds formed by one monatomic
metal ion and one monatomic non-metal ion are called binary ionic compounds. Also, negative polyatomic
ions often end in -ite or -ate (exception: hydroxide). Remember from “Oxidation Numbers” that a
polyatomic ion ending in “-ite” has one less O than its matching “-ate” partner (NO2 is nitrite and NO3
is nitrate / SO3 is sulfite and SO4 is sulfate).
* Other subscripts in the formula do not affect the name of the compound.
Examples: Write the names given the chemical formulas for the ionic compounds.
Li2S
=
lithium sulfide.
SrF2
=
strontium fluoride.
AlBr3
=
aluminum bromide.
Na2SO4 =
sodium sulfate.
KMnO4 =
potassium permanganate.
AgNO3 =
silver (I) nitrate.
Cr(NO3)3 =
chromium (III) nitrate.
CuO
=
copper (II) oxide – since there is one atom each for Cu and O, and the charge of O is -2.
This tells you that the charge of the copper must be +2, so it is copper (II) in this case.
SnCl2
=
tin (II) chloride – since there are 2 chlorides per tin, and since chloride ions have a charge
of -1, then the tin must be tin (II).
Fe2S3
=
iron (III) sulfide,
but
FeS
=
iron (II) sulfide.
PRACTICE PROBLEMS
Write the correct formula
1. lithium fluoride
Name the following
1. KCl
2. sodium bromide
2. LiI
3. magnesium sulfide
3. CaS
4. strontium oxide
4. BaO
5. lithium sulfide
5. K2S
6. sodium oxide
6. Li2O
7. calcium chloride
7. MgF2
8. barium bromide
8. SrI2
9. aluminum chloride
9. AlBr3
10. silver (I) oxide
10. PbI2
11. manganese (II) sulfide
11. NiCl2
12. tin (II) chloride
12. ZnS
13. chromium (III) oxide
14. copper (I) oxide
13. MnO
14. CuS
15. iron (III) chloride
15. FeBr2
16. mercury (I) fluoride
16. HgI2
17. ammonium nitrite
17. NH4OH
18. barium nitrate
18. Pb(NO2)2
19. calcium carbonate
19. Mg3(PO4)2
20. aluminum sulfate
20. SrSO3
21. sodium nitrite
22. potassium phosphate
21. Zn(NO3)2
22. MgCO3
23. aluminum hydroxide
23. Zn(OH)2
24. copper (II) hydroxide
24. Fe(OH)2
25. mercury (II) phosphate
25. Cu3(PO4)2
26. mercury (I) acetate
26. NaHCO3
27. strontium chlorate
27. KMnO4
28. nickel (II) chromate
28. ZnCr2O7
29. manganese (II) bicarbonate
29. Cr(C2H3O2)3
30. sodium permanganate
30. SnCrO4
Hydrated Ionic Compounds
Hydrated ionic compounds are simply ionic compounds with water attached in the crystal structure.
When naming hydrated ionic compounds, follow these steps:
1.
Name the ionic part of the compound (like in Part 1). Leave a space.
2.
Write the Greek prefix pertaining to the number of water molecules included in the compound,
followed by the word “hydrate”.
Example 1:
CuSO4 5 H2O = copper (II) sulfate pentahydrate
Example 2:
CoCl2  7 H2O = cobalt (II) chloride heptahydrate
Practice Problems:
Write the names or formulas of the following hydrates
1. CuCl2 • 2 H2O
1. calcium chloride dihydrate
2. Fe(NO3)3 • 9 H2O
2. aluminum chlorate hexahydrate
3. FeCl3 • 3 H2O
3. barium iodate monohydrate
4. BaCl2 • 2 H2O
4. magnesium sulfate heptahydrate
5. NiCl2• 6 H2O
5. iron (II) chloride tetrahydrate
Greek
prefixes
1 = mono
2 = di
3 = tri
4 = tetra
5 = penta
6 = hexa
7 = hepta
8 = octa
9 = nona
10 = deca
Acids
There are two types of acids, binary and oxyacids. All acids can be easily identified because their formulas
start with hydrogen or H.
Binary Acids contain hydrogen and a monatomic anion, and are usually followed by the
physical state (aq) or aqueous, which means “dissolved in water”. The word “binary” means
“TWO”, so a binary acid contains only 2 elements (not necessarily 2 atoms).
When naming binary acids, follow these steps:
1.
Determine the name of the anion.
2.
Write the prefix “hydro” followed by the root word for the anion, and
then the suffix “ic”. Leave a space, then write the word “acid”.
Example 1:
HCl (aq) = hydrochloric acid
Example 2:
H2S (aq) = hydrosulfuric acid
Example 3:
HF (aq) = hydrofluoric acid
Example 4:
H3P (aq) = hydrophosphoric acid
Common Root
Words:
Br = bromCl = chlorF = fluorI = iodoN = nitrP = phosphOR phosphorS = sulfOR sulfur-
Oxyacids contain hydrogen and a polyatomic ion containing oxygen.
When naming oxyacids, follow these steps:
1.
Determine the name of the anion.
2.
Write the root word for the anion.
3.
If the anion’s name ends with “–ate” and then the suffix “ic”. If the anion’s name ends with
“–ite” and then the suffix “ous”.
4.
Leave a space, then write the word “acid”
Example 1:
H2SO4 = sulfuric acid but
H2SO3 = sulfurous acid
Example 2:
HNO3 = nitric acid
HNO2 = nitrous acid
Example 3:
H2CrO4 = chromic acid but
but
H2CrO3 = chromous acid
Note: There is a pattern to the prefixes and suffixes given to oxyanions!
Anion
Corresponding Acid
ClO31- = chlorate
ClO21- = chlorite
HClO3 = chloric acid
HClO2 = chlorous acid
SO42- = sulfate
SO32- = sulfite
H2SO4 = sulfuric acid
H2SO3 = sulfurous acid
When writing the formulas for acids, follow the same rules for writing the formulas of ionic compounds!
* Hint: Find the “ate” anion FIRST (ex. NO31-, ClO31-, C2H3O21-, SO42-, CO32-, CrO42-, PO43-)
**Remember to balance charges!!!
Practice Problems:
Write the names of the following acids. Write the formulas of the following acids. Remember to balance
charges.
1. HBr
1. acetic acid
2. HI
2. hydrofluoric acid
3. HClO3
3. nitrous acid
4. H2CO3
4. sulfuric acid
5. H3P
5. chloric acid
6. H2SO3
6. phosphoric acid
7. H2CrO4
7. permanganic acid
8. H3PO4
8. carbonous acid
9. H2SO4
9. hydroiodic acid
10. HClO2
10. hydrobromic acid
Binary Molecular Compounds
Greek
prefixes
1 = mono
2 = di
3 = tri
4 = tetra
5 = penta
6 = hexa
7 = hepta
8 = octa
9 = nona
10 = deca
These are compounds that are composed of two nonmetals. Naming these compounds differs
from ionic compounds, because molecules contain no ions. So, we need to tell how many of
each atom is found when we write the name. The Greek prefixes to the left are used as
subscripts to tell how many atoms of each element are present.
When naming molecules, follow these steps:
1. Write the prefix pertaining to the number of atoms of the first nonmetal element, then write
the name of the first element. Leave a space.
NOTE: The prefix “mono-” is typically not included in molecule names
2. Write the prefix pertaining to the number of atoms of the second nonmetal element, then
write the root name of the second nonmetal element, and then the suffix –“ide”.
Example 1:
BH3 =
boron trihydride
Example 2:
P2O3 =
diphosphorous trioxide
Example 3:
N2O5 = dinitrogen pentoxide
Example 4:
IF6 =
iodine hexafluoride
Practice Problems:
Write the names of the following molecules.
1. P2O5
Write the formulas for the following molecules.
1. iodine heptafluoride
2. SO3
2. nitrogen trihydride
3. SO2
3. carbon disulfide
4. CCl4
4. diphosphorous hexoxide
5. CO2
5. dinitrogen monoxide
6. CO
6. dihydrogen monoxide
7. SCl6
7. phosphorous tribromide
8. N2O4
8. silicon dioxide
9. SiO2
9. dichlorine heptoxide
10. N2O3
10. carbon tetrafluoride
In E-44 “Organic Nomenclature”, we will learn how to write the names of other types of molecules –
primarily those that are carbon based.