Re-typed from The Ultimate Chemical Equations Handbook by Hague and Smith
SIMPLE INORGANIC FORMULAS AND NOMENCLATURE
Binary Molecules
1. A binary molecule is formed when two nonmetals or metalloids combine. Electrons are shared so the
bonding involved is known as _________________________ bonding.
2. Sometimes these compounds have generic or common names (water) and they also have systemic
names (dihydrogen monoxide). The common names must be memorized. The systemic name is more
complicated but it has the advantage that the formula of the compound can be deduced from the name.
3. Simple binary compounds consist of only a few atoms. Systemic naming of these compounds follow
the rules:

The elements, except for H, are written in order of increasing group number.

Prefixes are used to designate the number of each element present in the molecule.
4. The prefixes are:
1
6
2
7
3
8
4
9
5
10
5. Mono is __________ used in front of the first element. If there is only one atom, the mono is
assumed.
***** Name the following binary molecules:
(a) CO2
(e) O2F
(b) N2O3
(f) SBr6
(c) P4O10
(g) XeF6
(d) N2O
(h) P2S3
Ionic Compounds
1. Ionic compounds are formed between metals and nonmetals. The electrostatic force of attraction
between the positive ion (cation) formed by the metal and the negative ion (anion) formed by the
nonmetal is what holds the compound together.
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2. The cation is named first and written first when writing a formula while the anion is both named and
written second.
3. Cations can have one or more charges. These charges are known as oxidation numbers, or valences.
4. The transition metals and representative elements in Groups 13, 14, 15, and 16 have multiple oxidation
numbers.
5. Roman numerals, enclosed in parentheses, are used after the name of the cation to designate the
oxidation state of the cation – only if the cation has more than one positive oxidation state from which to
choose.
6. This system of nomenclature is known as the Stock system.
7. Here are some simple rules that should help in the determination of the oxidation numbers of cations
from the formulas of their compounds.
(a) The oxidation number of any element in its free state (uncombined with other elements) is zero. Fe
in a bar of iron is zero. O2 and N2 in the atmosphere are zero.
(b) The oxidation number of the alkali metals in a compound is always 1+.
(c) The oxidation number of the alkaline earth metals in a compound is always 2+.
(d) Fluorine in a compound is always assigned an oxidation number of 1-.
(e) The oxidation number of oxygen is almost always 2- in a compound. Exceptions to this rule would
be peroxides, O22- where the oxidation number of each oxygen is 1-, and superoxides, O2- where the
oxidation number of each oxygen is ½-.
(f) In covalent compounds (with nonmetals) hydrogen is assigned an oxidation number of 1+ (examples
are HCl, H2O, NH3, CH4).
(g) In metallic halides the halogen (F, Cl, Br, I, and At) always has an oxidation number equal to 1-.
(h) Sulfide, selenide, telluride, and polonide are always 2- in binary salts.
(i) Nitrides, phosphides, and arsenides are always 3- in binary salts.
(j) All other oxidation numbers are assigned so that the sum of the oxidation numbers of each element
equals the net charge on the molecule or polyatomic ion. In neutral compounds, the sum of the positive
and negative charges must equal zero.
***** Determine the oxidation number of the underlined element:
(a) KMnO4
(c) K2W4O13
(b) NaClO4
(d) Au2O3
Re-typed from The Ultimate Chemical Equations Handbook by Hague and Smith
7. Free elements, no matter how complex the molecule, have an oxidation number (valence or charge)
equal to zero.
8. The following are diatomic or polyatomic elements in nature which must be committed to memory.
These elements exist as neutral molecules in nature:
(a)
(g)
(b)
(h)
(c)
(i)
(d)
(j)
(e)
(k)
(f)
(l)
Charges and the Periodic Table
1. The periodic table can be used to help determine charges on many ions. Cations come from metals
that lose electrons (_________________________) in order to become _________________________
with a noble gas. Anions come from nonmetals that gain electrons (_________________________) in
order to become ________________________ with a noble gas.
*****
Group 1
Group 17
Group 2
Group 16
Group 3
Group 15
Group 14
2. Transition metals, representative metals with ____ and _____ sublevels, and the inner transition metals
typically have more than one oxidation state.
3. Electrons for these metallic elements are lost in the following order: ____________________. Such
elements are _______ isoelectronic with a noble gas when the outermost (valence) electrons are lost.
4. Inner transition elements are also known as the lanthanides, the actinides, the rare earth elements, and
the transuranium elements. These elements are rare and many exist for short periods of time.
5. Both inner transition elements and transition elements are known for their variable oxidation numbers.
The most common oxidation number for transition elements is _______.
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6. The _____ sublevel in transition elements is responsible for the various oxidation numbers that result.
Incomplete d sublevels are also responsible for the many colorful transition compounds that are known to
exist. Complete d sublevels in cations of silver and zinc result in white compounds.
SUMMARY OF CATIONS WITH VARIABLE OXIDATION NUMBERS
– STOCK SYSTEM
1+, 2+
1+, 3+
2+, 3+
2+, 4+
3+, 4+
3+, 5+
2+, 3+, 4+
copper (I), Cu+; copper (II), Cu2+
mercury (I), Hg22+; mercury (II), Hg2+
gold (I), Au+; gold (III), Au3+
indium (I), In+; indium (III), In3+
thallium (I), Tl+; thallium (III), Tl3+
chromium (II), Cr2+; chromium (III), Cr3+
cobalt (II), Co2+; cobalt (III), Co3+
iron (II), Fe2+; iron (III), Fe3+
lead (II), Pb2+; lead (IV), Pb4+
platinum (II), Pt2+; platinum (IV), Pt4+
tin (II), Sn2+; tin (IV), Sn4+
zirconium (II), Zr2+; zirconium(IV), Zr4+
cerium (III), Ce3+; cerium (IV), Ce4+
antimony (III),Sb3+; antimony (V), Sb5+
arsenic (III), As3+; arsenic (V), As5+
bismuth(III), Bi3+; bismuth (V), Bi5+
phosphorus (III), P3+; phosphorus (V), P5+
iridium (II), Ir2+; iridium (III), Ir3+; iridium (IV), Ir4+
titanium (II), Ti2+; titanium (III), Ti3+; titanium (IV), Ti4+
manganese (II), Mn2+; manganese (III), Mn3+; manganese (IV), Mn4+
2+, 4+, 5+
tungsten (II), W2+; tungsten (IV), W4+; tungsten (V), W5+
3+, 4+, 5+
uranium (III), U3+; uranium (IV), U4+; uranium (V), U5+
2+, 3+, 4+, 5+
vanadium (II), V2+; vanadium (III), V3+; vanadium (IV), V4+;
vanadium (V), V5+
Re-typed from The Ultimate Chemical Equations Handbook by Hague and Smith
Polyatomic Ions
1. Polyatomic ions are a group of atoms that behave as a single ion.
2. The bonding within a polyatomic ion is covalent, but because there is either an excess or a shortage of
electrons compared to the number of protons present, an ion results.
3. This short list of polyatomic ions must be MEMORIZED.
NH4+
ammonium
OH-
hydroxide
NO2-
nitrite
PO43-
phosphate
NO3-
nitrate
ClO-
hypochlorite
SO32-
sulfite
ClO2-
chlorite
SO42-
sulfate
ClO3-
chlorate
CO32-
carbonate
ClO4-
perchlorate
HCO3-
bicarbonate
CrO42-
chromate
C2H3O2-
acetate
Cr2O72-
dichromate
CN-
cyanide
SCN-
thiocyanate
You also need to know the common Group ions such as (but not limited to) the ions in Groups 1 &
2, the halogens, O, S, N, P, and “the triangle”: Ag, Zn, Cd, Al, Ga, and In.
OVER
Re-typed from The Ultimate Chemical Equations Handbook by Hague and Smith
Ternary Nomenclature: Acids and salts Containing Halogens and/or Oxygen
1. The halogens, with their variable oxidation numbers, allow for a great variety of compounds.
2. A good way to learn ternary nomenclature is to start with a certain ____________________
polyatomic ion. This is polyatomic ion ending with the suffix __________.
Number of Oxygen Atoms
(compared to home base)
Polyatomic Ion Name
Plus one oxygen atom
ClO4-
perchlorate ion
Home base
ClO3-
chlorate ion
Minus one oxygen atom
ClO2-
chlorite ion
Minus two oxygen atoms
ClO-
hypochlorite ion
No oxygen atoms
Cl-
chloride ion
3. Water solutions of binary hydrides form acids. The root derived from the hydride is given the prefix
_______________ and the suffix __________ and the name ends with the word __________.
4. The binary hydride HCl is known as _________________________ (hydrogen monochloride) gas, but
when aqueous it is known as ______________________________.
5. Many common acids contain only oxygen, hydrogen, and a nonmetallic ion or polyatomic ion. These
acids are called _________________________.
6. If the name of the polyatomic ion ends in __________, the suffix ________ is substituted followed by
the word acid.
7. If the name of the polyatomic ion ends in __________, the suffix _______ is substituted followed by
the word acid.
8. The mnemonic aid is:
***** Name the following compounds:
(a) HIO4 (aq)
(c) H3PO4 (aq)
(b) NaBrO4
(d) HBr (aq)
OVER
***** Write the formulas for the following compounds:
(a) calcium hypochlorite
(c) cyanic acid
(b) hydrotelluric acid
(d) chlorous acid
BALANCING EQUATIONS HINTS
SYNTHESIS & DECOMPOSITION
Chemists write balanced equations to illustrate what is happening during a chemical reaction. Bonds are
broken, atoms are rearranged, and new bonds are formed. Every chemical reaction supports the Law of
conservation of Matter. This means that in every reaction, the number of atoms of each type of element
contained within the reactants must be the same as the number of atoms of each type of element contained
within the products.
Balancing equations is a process which assures that equations are written properly to support the Law of
Conservation of Matter; however, balancing cannot be done until each reactant and product formula is
written correctly. It is important to properly write the seven elements that are diatomic in their elemental
form and also use subscripts and parentheses appropriately when considering the oxidation number of
ions. All compounds must be made neutral before beginning to balance the atoms.
Balancing is accomplished by adding coefficients that multiply the number of atoms represented by the
formula. For example, a coefficient of 2 in front of oxygen (e.g., 2 O2) means that 4 oxygen atoms are
represented. Unlike algebra, in chemistry a coefficient does not need to be outside parentheses or
brackets to b distributed. A coefficient applies to the complete substance; however, it no longer applies
when a plus sign (+) or arrow (→) is encountered. For example:
3 (NH4)2CO3 shows 6 nitrogen, 24 hydrogen, 3 carbon, and 9 oxygen atoms.
3 MgCl2 + NaBr indicates 3 magnesium, 6 chlorine, 1 sodium, and 1 bromine atom.
PREREQUISITE KNOWLEDGE
Before you do anything, you must know and understand the following areas of nomenclature and formula
writing:
1. Ionic compounds
2. Covalent compounds
3. Acids and bases
4. Complex ions (coordination chemistry)
5. Organic nomenclature
Re-typed from The Ultimate Chemical Equations Handbook by Hague and Smith
TIPS FOR BALANCING EQUATIONS
1. Ensure each molecular formula is written correctly and each compound is neutral.
2. Mentally count or tally how many of each type of atom is present on each side of the equation.
3. Begin by balancing elements that are only found in one substance on each side.
4. Balance oxygen and hydrogen LAST – they usually balance out at the end or perhaps only the number
of water molecules needs to be adjusted.
5. If there is an odd number of an element on one side and an even number on the other, the odd number
will need to be evened out – so use a coefficient of 2 for that substance.
6. If there are polyatomic ions that remain together as a unit during the reaction, count the polyatomic ion
as a unit.
7. When tallying, be sure to adjust the count for each and every element that an added coefficient affects.
8. Combustion reactions that don’t seem to balance will often come out better if a coefficient of 2 is used
for the hydrocarbon.
SYNTHESIS REACTIONS
Synthesis reactions occur when two or more reactants combine to form a single product. There are
several types of synthesis reactions.
1.
A metal combines with a nonmetal to form a binary salt.
Example: A piece of lithium metal is dropped into a container of nitrogen gas.
6 Li + N2 → 2 Li3N
2. Metallic oxides and water form bases (metallic hydroxides)
Example: Solid magnesium oxide is added to water.
MgO + 2 HOH → Mg(OH)2
3. Nonmetallic oxides and water form acids. The nonmetal retains its oxidation number.
Example: Dinitrogen pentoxide is bubbled into water.
N2O5 + H2O → 2 HNO3
4. Metallic oxides and nonmetallic oxides form salts.
Example: solid calcium oxide is added to sulfur trioxide.
CaO + SO3 → CaSO4
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DECOMPOSITION REACTIONS
Decomposition reactions occur when a single reactant is broken down into two or more products.
1.
Metallic carbonates decompose into metallic oxides and carbon dioxide.
Example: A sample of magnesium carbonate is heated.
MgCO3 → MgO + CO2
2. Metallic chlorates decompose into metallic chlorides and oxygen.
Example: A sample of magnesium chlorate is heated.
Mg(ClO3)2 → MgCl2 + 3 O2
3. Ammonium carbonate decomposes into ammonia, water and carbon dioxide.
Example: A sample of ammonium carbonate is heated.
(NH4)2CO3 → 2 NH3 + H2O + CO2
4. Sulfurous acid decomposes into sulfur dioxide and water.
Example: A sample of sulfurous acid is heated.
H2SO3 → H2O + SO2
5. Carbonic acid decomposes into carbon dioxide and water.
Example: A sample of carbonic acid is heated.
H2CO3 → H2O + CO2
6. A binary compound may break down into two elements.
Example: Molten sodium chloride is electrolyzed.
2 NaCl → 2 Na + Cl2
7. Hydrogen peroxide decomposes into water and oxygen.
Example: 2 H2O2 → 2 H2O + O2
8. Ammonium hydroxide decomposes into ammonia and water.
Example: NH4OH → NH3 + H2O
Re-typed from The Ultimate Chemical Equations Handbook by Hague and Smith
NOMENCLATURE PRACTICE
1. Name each of the following compounds.
(a) CaF2
(h) PF5
(b) HC2H3O2 (aq)
(i) (NH4)2SO3
(c) NaH
(j) Au2(C2O4)3
(d) HIO (aq)
(k) KSCN
(e) Hg2Cl2
(l) Pb(C2H3O2)4
(f) BaO2
(m) LiH
(g) Mn(HCO3)3
(n) H2CrO4 (aq)
2. Write formulas for the following substances.
(a) vanadium (V) oxide
(g) tin (IV) chromate
(b) zinc hydroxide
(h) nitrous acid
(c) silver chromate
(i) magnesium hydrogen phosphate
(d) tungsten (V) thiosulfate
(j) calcium carbide
(e) cerium (III) nitrate
(k) mercury (II) acetate
(f) hydroiodic acid
(l) perbromic acid
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BALANCING SYNTHESIS AND DECOMPOSITION EQUATIONS
Predict and balance the following synthesis and decomposition reactions. Use abbreviations to indicate
the phase of reactants and products where possible [(aq) (s) (l) (g)]
1. A sample of calcium carbonate is heated.
2. Sulfur dioxide gas is bubbled through water.
3. Solid potassium oxide is added to a container of carbon dioxide gas.
4. Liquid hydrogen peroxide is warmed.
5. Solid lithium oxide is added to water.
6. Molten aluminum chloride is electrolyzed.
Re-typed from The Ultimate Chemical Equations Handbook by Hague and Smith
7. A pea-sized piece of sodium is added to a container of iodine vapor.
8. A sample of carbonic acid is heated.
9. A sample of potassium chlorate is heated.
10. Solid magnesium oxide is added to sulfur trioxide gas.
OVER