This document provides information about chemical nomenclature and reactions. It discusses writing chemical formulas, empirical formulas, chemical naming conventions, chemical equations, balancing equations, and types of chemical reactions. Specifically, it covers how to write formulas using oxidation numbers and valence, how to determine empirical and molecular formulas from composition data, rules for naming ionic and molecular compounds, the components and symbols used in chemical equations, and methods for balancing equations. It also defines four main types of chemical reactions: combination, decomposition, single displacement, and double displacement.

1. Chemical Nomenclature and
Reactions
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2. Contents
• Formula Writing
• Empirical Formula
• Chemical Nomenclature
• Chemical Reactions and Equations
• Balancing Chemical Equations
• Types of Chemical Reactions
• Balancing Oxidation Reduction Reactions

3. 6.1 Formula Writing
• Chemical Formula- a representation of the composition of compounds, it
contains:
-The symbol and formula of elements or radicals present.
-The number of atoms of each element present written as numerical
subscripts to the right of each symbols.

4. Things to Know in Writing a Chemical Formula
• Correct Symbols of element and radicals
- VALENCE denoted the number of electrons in the outermost energy
level; it also describes combining power of an atom in a compound.
-RADICAL (polyatomic Ions) are group of elements which act as one or
single atom in chemical reaction and show definite valence.
• Corresponding Valence or Oxidation numbers
-OXIDATION NUMBER refers to the combining capacity of an atom but
specifying charge (whether positive or negative)

5. Writing a Formula Using a Criss-Cross Rule
• Rule 1: Identify the charge of each element. Drop down the charge of each
element to become the subscripts.
• Rule 2: When valence number is 1, subscript is not written.
• Rule 3: When the oxidation number of both elements are numerically equal
but greater than 1, the subscript is not also written.
• Rule 4: All radicals take more than one (the subscript is 2 or more), must be
enclosed with a parenthesis ().
• Rule 5: All subscripts must be reduced to lowest terms (except for molecular
or covalent compounds)

6. 6.2 Empirical Formula
• Empirical formula or the simplest formula of a compound gives the smallest
whole number ratio of atoms that make up the compound.
• Determined from the percent composition of the compound or the
experimental determined mass relationship of the elements that make up the
compounds.

7. To calculate the empirical formula of a compound
• Calculating the empirical formula from the percent composition, one can convert the
percentages to gram.
• Divide these numbers of grams of each element by the respective atomic weight of
the element; gives as the ratios.
• Choose the smallest quotient in step 2. Divide each quotient by the smallest quotient.
• The number obtained becomes the subscript of the element.
• If the whole number is not obtained, multiply by whatever number is necessary to
obtain whole number.
• Write the simplest formula.

8. Molecular Formula
• Formula whose subscript represents the absolute exact numbers of atoms of
each element per molecule of the compound or the absolute number of
moles od each element per mole of the compound.
• Molecular formula may be reducible to a simple formula if all its subscripts
are divisible by the common denominator.
• Some compounds have the same empirical and molecular formula.
• There are many situation where two or more compounds have the same
simplest formula, but differ by their molecular formulas.

9. Structural Formula
• Formula that not only gives via its subscripts the exact number of atoms of
each element per molecule but it displays the way that the atoms are bonded
together and the shape of the molecule is revealed.
• There are compounds that have the same empirical formula and even the
same molecular formula and the only way that they can be distinguished is
through their structural formulas.

10. Interpretation of Formulas
• The number between the letters are subscripts and represent the number of
atoms of the element that they follow in one molecule of the substance.
• H2O has two atoms of Hydrogen and 2 atom of oxygen for every molecule of
water. This can be interpreted as 2 moles of hydrogen and 1 mole of
hydrogen in one mole of H2O. The mole interpretation is the more practical
interpretation because we are not capable of seeing single molecules and
atoms for everyday work.
• C12H12O11 has 12 atoms of carbon, 22 atoms of hydrogen, and 11 atoms of
oxygen in one molecule of table sugar. The mole interpreted would be 12
moles of carbon, 22 moles of Hydrogen, and 11 moles of Oxygen in every
mole of table sugar.

11. To Calculate the Molecular Formula the Molecular
weight must be know.
• Determine the empirical formula
• Add all the relative masses of the atom in an empirical formula
• Divide the molecular mass of the compound by the mass of the empirical
formula to obtain the multiple.
• Multiply the number of atoms of each element in the empirical formula by this
multiply and write the molecular formula of the compound.
• It is also possible to do this with one of the elements in the formula; simply divide
the mass of that element in one mole of a compound by the mass of the element
in the empirical formula. The result should always be a natural number

12. 6.3 Chemical Nomenclature
• In naming the compound, one has to first decide whether
you are looking at an ionic compound or a molecular
compound,
Identifying a Compound as Ionic or Molecular

13. Ionic Compound
• Metals combined to non-metals will produced compound that are ionic. Those
formulas that have one symbols from the metal combined with symbol from
non-metals would make the compound an ionic compound.
• For Example: N𝑎2 SO4 would be ionic because it has the sulfur ion in which is
composed of non-metals combined with sodium which is a metal. Metals
combined with non-metals produce ionic compounds.

14. Molecular Compound
• Non-metals combined with non-metals will produce compounds that are
molecular. If the formula has both symbols from non-metals, then the
compound would be classified as a molecular compound.
• For example: PCl5 would be molecular because Phosphorus and Chlorine
are both non-metals produce molecular compounds.

15. Rules in Naming Compounds
• Compounds with only two elements present are called binary compounds.
• Binary compounds that are formed metals and non-metals are usually ionic in
nature. Ionic compound is a compound made up of a positive and negative
ions joined together by electrostatic force of attraction.
 Binary Ionic Compounds (containing 2 different elements Metal and Non-
Metal) The end in –ide.
• If there is a multi-valent element involved like Iron, copper, lead, tin or
mercury, one will have to determine which valence is involved before the
name can be established.

16. • We can use the Stock System or Classical Methods in naming Binary Ionic
Classical Method in naming Binary Ionic Compounds.
1. Stock System: Roman numeral enclosed in the parenthesis is written
immediately following the name of the metal to indicate the valence of
metal.
2. Classical Method: Name of metal is modified with ending –ic for higher
valence and –ous for lower valence.

17.  Ternary Ionic Compound
• Naming Ternary compounds uses the same procedure as Binary Ionic
Compounds. The only difference is the ending of the name.
• The name ends in-ite for less oxygen, -ate for more oxygen,
 Binary Molecular Compounds (Containing 2 non-metals). Greek prefixes are
used to indicate the number of atoms. The prefix mono is dropped at the start
of the name.
 Binary Molecular Compound (containing Hydrogen) listed as the first
element. They are named without the Greek numerical prefixes.

18. List of Cations According to Charge
• Cations with fixed 1+ charge (Group IA elements of the periodic table,
ammonium ion, and silver ion)
Name Symbol
Ammonium ion NH4
+
Cesium ion Cs+
Hydrogen H+
Lithium Li+
Potassium ion K+
Rubidium ion Rb+
Sliver ion* Ag+
Sodium ion N𝑎+

19. • Cations with fixed 2+ charge and fixed 3+ charge (Group IIA elements of the
periodic table, zinc ion, nickel ion, and cadmium ion)
Name Symbol
Aluminum ion Al3+
Barium ion Ba2+
Baryllium ion Be2+
Cadium ion* Cd2+
Calcium ion C𝑎2+
Magnesium ion Mg2+
Radium ion Ra2+
Strontium ion Sr2+
Zinc ion Zn2+

20. • Cations with variable charge
A. Cations with 2+ and 3+ charge
Name Symbol Name Symbol
Chromium (II) C𝑟2+
Iron (III) Fe3+
Chromium (III) Cr3+
Manganese (II) Mn2+
Cobalt (II) C𝑜2+ Manganese (III) Mn3+
Cobalt (III) C𝑜3+ Nickel (II) Ni2+
Iron (II) Fe2+ Nicken (III) Ni2+

21. B. Cations with 2+ and 4+ charge
Name Symbol Name Symbol
Lead (II) Pb2+
Tin (II) Sn2+
Lead (III) Pb3+
Tin (IV) Sn4+

22. C. Cations with 1+ and 2+ charge
Name Symbol Name Symbol
Copper (I) C𝑢+
Mercury (I)** Hg2
2+
or Hg+
Copper (II) C𝑢2+
Mercury (II) Hg2+

23. D. Others
Name Symbol Name Symbol
Antimony (III) Sb3+
Bismuth (III) Bi3+
Antimony (V) Sb5+ Bismuth (V) Bi5+
Arsenic (III) As3+
Gold (I) Au+
Arsenic (V) As5+ Gold (III) Au3+
*Have other charges but for nomenclature purposes they are considered
to have fixed 2+ charge because it is the common or stable oxidation state. For
this reason, older books consider nickel ion with fixed charge, 2+
**Mercury (I) ions are always bonded in pairs so it is more appropriate to
use Hg2
2+
than Hg+
.

24. List of Anions According to Suffixes
I. Anions with suffix -ide
Name Symbol Name Symbol
Arsenide As2 Nitride N3
Bromide Br Oxide O5
Chloride Cl Peroxide O2
2−
Cynide CN Phosphide P3−
Fluoride F Selenide Se2−
Hydride H Sulfide S2−
Hydroxide OH Telluride Te2−
Iodide I

25. II. Anions with suffic -ite
Name Symbol Name Symbol
Bisulfite or Hydrogen
Sulfite
HSO3 Nitrite NO2
Chlorite ClO2 Phosphite PO3
3−
Hypobromite BrO Sulfite SO3
2−
Hypochlorite ClO

26. III. Anions with suffix -ate
Name Symbol Name Symbol
Acetate CH3 or COO Dichromate Cr2 O7
2−
Bicarbonate or
Hydrogen Carbonate
HCO3 Nitrate NO3
Bisulafate or Hydrogen
Sulfate
HSO4 Oxilate C2 O4
2−
Borate BO3
3−
Perchlorate ClO4
Bromate BrO3 Permanganate MnO4
Carbonate CO3
2− Phosphate PO4
3−
Chlorate ClO3 Sulfate SO4
2−
Chromate CrO4
2−
Tetraborate B4 O7
2−

27. 6.4 Chemical Reactions and Equation
• Is an expression of a chemical process that uses symbols and formulas
instead of words to describe the changes that occur in a chemical reaction.
AgNO3 (aq)+NaCl(aq) AgCl(s) + NaNO3 (aq)
• In this equation, AgNO3 is mixed with NaCl. The equations shows that the
reactants (AgNO3 and NaCl) react through some process ( ) to form the
product (AgCl + NaNO3 ). Since they undergo a chemical process, they are
changed fundamentally.
• A chemical equation illustrates the Law of Conversion of Mass, that matter
cannot be created nor destroyed in a chemical reaction. Therefore, the
equation must be balanced, that is the same number of atoms of each elemet
must appear on both sides of the equation.

28. Parts of a Chemical Equation
• Reactants- The substances that combine in the reaction. Formulas of
reactants must be correctly written on the left side of the equation.
• Products- The substances that are formed by the reaction. Formulas of
products must be correctly written at the right side of equation.

29. Other Terms and Symbols Used in Chemical
Equations
• Essential Symbols
Found between reactants and Products,
means “reacts” to form: yeald(s),
product(s) (This symbols points to
products).
Separates each reactants and each
product. Means “reacts with; combines
with” (if it used used to separate reactant).
It means “and” (if it used to separate each
product).

30. • Optional Symbols
 Physical States indicates the physical state of the substance whose formula it follows
(g)- indicates that the substance is gas
(l) Indicates that the substance is liquid
(s) – indicates that the substance is solid
(aq)- means that the substance is in aqueous (water solution)
Placed after the formula of a product that is
gas. (gas being liberated)
Placed after the formula of a product that is an
insoluble solid-that is precipitate.

31.  Coefficient. Are used in all chemical equations to show the relative amounts
of each substance present. This amount can represent either the relative
number of molecules, or the relative number of moles. If no coefficient is
shown, a one (1) is assumed.
 Conditions. Words or symbols placed above or below the arrow to indicate
conditions used to make the reaction occur, such as a value for temperature.
 Catalyst. Is also indicate above or below the arrow. It is a substance that alter
the speed of the reactions without being consumed in the reaction. In
general, a catalyst is used to slow down reactions. These are sometimes
called negative catalyst or inhibitors.

32. 6.5 Balancing Chemical Equations
• Balancing Chemical Equation is absolutely essential if you want to determine
quantities of reactants or products.
• Balancing equation assures that the Conservation of Matter is obeyed. The
total mass of reactants must equal the total mass of products. A balanced
equation is like a recipe. It tells you the proportional quantities of each
substance involved.
• It is essentially done by the trial and error. There are many different ways and
systems of doing this, but for all methods, it is important to know how to count
the number of atoms in an equation. For example, we will look at the
following term.
2Fe3 O4

33. • Reminder:
1. Never touch subscripts when balancing equations since that will change the
composition and therefore the substance itself.
2. Check to be sure that you have included all sources of a particular element
that you are balancing in a particular side since there may be two or more
compounds that contain the same element on a given side of an equation.
3. Adjust the coefficient of mono atomic elements near the end of the
balancing act since any change in their coefficient will not affect the balance
of other elements/

34. • Now let us try to balance this equation:
Al + Fe3 O4 Al2 O3 + Fe
Steps
1. Count the number of each atom on the reactant and on the product side.
2. Determine a term to balance first. When looking at this problem, it appears that the oxygen will
be most difficult to balance so we’ll try to balance the oxygen first.
• Be sure to notice that the subscript times the coefficient will give the number of that element.
On the reactant side, we have a coefficient of the three (3) multiplied by a subscript of four (4),
giving twelve(12) oxygen atoms. On the product side, we have a coefficient of four (4)
multiplied by the subscript of three (3), giving 12 oxygen atoms. Now the oxygen are balanced.

35. 3. Choose another term to balance. We’ll choose iron, Fe. Since there are nine iron atoms in the
term in which oxygen is balanced we all a nine coefficient in the front of Fe.
4. Balance the last term. In this case, since we had 8 aluminum on the product side we need to
have 8 on the reactant side we add 8 in front of the Al term on the reactant side.

36. 6.6 Types of Chemical Reactions
• Elements and compounds react with each other in numerous ways. Memorizing
every type of reaction would be challenging and also unnecessary, since nearly every
inorganic chemical reactions falls into one or more of the four broad categories.
1. Combination Reactions. Two or more reactants from one product in a
combination reactions.
2. Decomposition Reactions. In decomposition reactions, a compound breaks
down into two or more substances. Decomposition usually results from electrolysis or
heating.
3. Single Displacement Reactions. A single displacement reaction is
characterized by an atom or ion of a single compound replacing an atom of another
element.

37. 4. Double Displacement Reactions. May also be called metathesis reactions. In this type of
reaction, element from two or more compound displace each other to form new compounds.
It may occur when one product is removed from the solution as a gas or precipitate or
when two species combine to form a weak electrolyte that remains undissociated in solution.
A neutralization reaction is specific type of double displacement reaction that occur when
an acid reacts with a base, producing a solution of salt and water.
• Remember that reactions can belong to more than one category. Also, it would be possible to
present more specific categories, such as combustion reaction or precipitation reactions.
Learning the main categories will help you balance equations and predict the types of
compounds formed from a chemical reaction.

38. 6.7 Balancing Oxidation Reduction Reactions
• Redox equations involved the change in the oxidation state of atoms involved
in the reaction. If the oxidation state becomes more negative then electrons
are gained and we say that the atom that is changing is being reduced. If, on
the other hand, the oxidation state of the atom becomes more positive, then
we say that electrons are being lost and the atoms that are changing are being
oxidized.
• Whenever there is a reduction there will also be an oxidation.
• In order to balance a REDOX equation one must balance not only the atoms
of each, element represented, but also the number of electrons being gained
and lost, and the electrical charge. Therefore, such system require a more
systematic approach compared to the trial and error method of balancing by
inspection.

39. Ruling for Assigning Oxidation State
• All elemental substances have a zero oxidation state per atom.
• Group 1 elements are always in the +1 state in a compound.
• Group 2 element are always in the +2 state in a compound.
• Oxygen containing compounds usually have oxygen atoms in the -2 state.
The exception, is the peroxide state where each Oxygen is in the -1 state.
• Hydrogen containing compounds usually have each Hydrogen atom in the +1
state except when combined with elements less electronegative that
Hydrogen in which case the Hydrogen is in the Hydride state and have a -1
Oxidation state.

40. • Neutral Compounds will have the total positive and negative oxidation state
by zero.
• Polyatomic ions will have the sum of the positive and negative oxidation
states add up to the charge indicated on the ion species.
• Monoatomic ions will have an oxidation state equal to the charge in the ion.

41. Solution
1. Assign an oxidation state to each to each atom in the reactants and
product, use the above guidelines in assigning oxidation states.
2. Identify which element underwent a change in the oxidation state and place
the reactant and product on opposite, side of an arrow.
3. Determine whether electrons are being gained (appear on the left) or lost
(appear on the right) and place them on the proper side if the half reaction
(make sure that the atoms undergo a change are balanced by inspection
before determining how many electrons are lost or gained.
4. Balance the number of electrons being gained and lost. To do this we need
to multiply every term in the first half reaction by two.

42. 5. Add up the two half reactions.
6. Balance the change with H+
if the solution is acidic and OH of the solution is
basic.
7. Balance the Hydrogen and Oxygen with neutral H2O so as not to upset the
charge balance.
8. Check by inspection to make sure that all spectator ions have been
balanced.