Chemistry of Stoichiometry

Fundamentals of
Chemistry
Module 2 – Chapter 9

Introduction to Stoichiometry

• Equations must always be balanced before calculation of any mass,
moles, or volume of a reactant or product!
• Stoichiometry is the area of chemistry that deals with quantitative

relationships between products and reactants in chemical equations.

• Solving stoichiometry problems always requires the use of:
• A balanced chemical equation (coefficients must be known!)
• Conversion factors in units of moles (mole ratios)

Mole Ratios

•Mole ratio is the conversion
factor between any two
species in a chemical reaction

•The mole ratio will come from
the coefficients of a balanced
chemical equation

Mole Ratios in Practice

•The mole ratio can be used as a
conversion factor to convert
between moles of one substance
and another.
•The desired quantity goes in the
numerator and the known
quantity goes into the
denominator of the mole ratio
•Same method as the solution
map from chapter 2.

Problem Solving for Stoichiometry Problems

Problem Solving for Stoichiometry Problems

Problem Solving for
Stoichiometry

Problems

•Remember that Step 1 is to
always ensure you have a
balanced equation!!!
•You must be in moles to
convert from one substance to
another!

Limiting
Reactants
•In chemical reactions, the
reaction will occur until one of the
reactants runs out

•Think of a burning fire. You need
oxygen, heat and fuel to keep a
fire going. If the fuel (wood) all
burns, the fire goes out. The wood
would be the limiting reactant
because had it not all burned, the
fire would continue to exist.

•In a chemical reaction, the
maximum amount of product
formed depends on the amount of
reactant not in excess, the limiting
reactant

Reaction Yield

• The amount of products formed calculated by stoichiometry are the
maximum yields possible (100%)
• Yields are often lower in practice due to side reactions, loss of

product while isolating/transferring the material, etc.
• The theoretical yield is the maximum possible yield for a reaction,

calculated based on the balanced chemical equation.
• The actual yield is the yield obtained from the reaction
• The percent yield is the ratio of the actual and theoretical yield

Reading
Review

What is stoichiometry?

What unit must you be in to convert from one
substance to another?

What is the limiting reactant?

What is the difference between theoretical
and actual yields?

How do you calculate the percent yield?

  • Slide 1
  • Introduction to Stoichiometry
  • Mole Ratios
  • Mole Ratios in Practice
  • Problem Solving for Stoichiometry Problems
  • Problem Solving for Stoichiometry Problems
  • Problem Solving for Stoichiometry Problems
  • Limiting Reactants
  • Reaction Yield
  • Reading Review

CH1000
Fundament
als of
Chemistry
Module 2 – Chapter 7

The Mole (or mol)

• In chemistry, a mole (mol) is a standard scientific unit for measuring large
quantities of very small entities such as atoms, molecules, or other
specified particles.

• The number represented by 1 mole above is also called Avogadro’s number.

• 1 mol of any element contains the same number of atoms, but can vary
greatly in the overall mass. (Atoms of different elements have different
masses)

Molar Mass

•Molar Mass is the atomic mass
of an element or compound in
grams which contains Avogadro’s
number of particles
• Molar masses are expressed

to 4 significant figures in the
text

•Convert atomic mass units on
the periodic table to grams and
sum the masses of the total
atoms present

Mole Map
** Not found in the textbook,

save for easy access

Molar Mass of Compounds

•Much like an element, molar
mass can be defined for a
compound
•Molar Mass is the mass of one
mole of the formula unit of a
compound
• The molar mass of a

compound is equal to the
sum of the molar masses of
all the atoms in the
molecule

Percent
Composition of
Compounds

Percent composition is the mass percent of each
element in a compound.
• Percent = parts per 100 parts
• Molar mass is the total mass (100%) of the compound

% Composition is independent of sample size

% Composition can be determined by:

• 1. Knowing the compound’s formula
• 2. Using experimental data

Percent Composition from the Compound’s Formula

Percent Composition from Experimental Data

Empirical
and
Molecular
Formula

Empirical Formula
Smallest whole number ratio of
atoms in a compound

Molecular Formula
Actual formula of a compound.
Represents the total number of
atoms in one formula unit of the
compound

Calculating
Empirical
Formulas
•Special Case:
• If fractions are

encountered,
multiply by a
common factor to
provide whole
numbers for each
subscript.

Calculating the Molecular Formula from the Empirical Formula

•If molar mass is known,
the molecular formula can
be calculated from the
empirical formula
•Molecular formula is a
multiple of the empirical
formula.

Reading
Review

What is Avagadro’s
number?

How would you
convert from grams

to atoms of an
element?

What is a mole?

What is the
difference between

empirical and
molecular formulas?

What is the special
case when

calculating empirical
formulas?

  • Slide 1
  • The Mole (or mol)
  • Molar Mass
  • Mole Map
  • Molar Mass of Compounds
  • Percent Composition of Compounds
  • Percent Composition from the Compound’s Formula
  • Percent Composition from Experimental Data
  • Empirical and Molecular Formula
  • Calculating Empirical Formulas
  • Calculating the Molecular Formula from the Empirical Formula
  • Reading Review

CH1000
Fundament
als of
Chemistry
Module 2 – Chapter 8

Chemical Equations

• Chemists use chemical equations to:
• Summarize a chemical reaction by displaying the substances reacting and

forming.
• Indicate specific amounts of materials consumed or produced during the

reaction.
• Reactants: substances consumed during the reaction.
• Products: substances formed during the reaction.

• Atom balance must be maintained in all chemical reactions.
• All atoms from reactants must appear as part of products.

a A + b B c C + d D

The
coefficient
1 is not
written in
a balanced
equation.

Chemical Equations

1. Reactants and products are separated by an arrow.
2. Reactants are on the left side of the arrow, products are on the right.
3. Whole number coefficients are placed in front of substances to

balance the atoms in the equation.
4. The numbers indicate the units of the substance reacted or formed

during the reaction.
5. Information about the reaction (temperature, time) may be placed

above or below the reaction arrow.
6. The physical state is written in brackets after the formula of the

substance. (g) for gas, (l) for liquid, (s) for solid, (aq) for aqueous

a A + b B c C + d D
Reactant

s
Products

Symbol
Summary

Symbol Significance
Produces (points towards products)

(s) Solid (written after substance)
(l) Liquid (written after substance)
(g) Gas (written after substance)
(aq) Substance dissolved in an aqueous

solution
Heat is added (above or below reaction
arrow)

Δ

Law of Conservation of Mass

• The total mass of substances in a chemical reaction must remain
constant.

water hydrogen + oxygen
100.0 g 11.2 g 88.8 g

100.0 g total of productsreactants

In any chemical reaction:
Mass of reactants = Mass of products

Writing and
Balancing
Chemical
Equations

A balanced chemical equations contain the same
number of each kind of atom on both sides of the
equation.

1. Write a word equation for the reaction.

2. Write the correct formula for each substance
(unbalanced):

3. Balance the equation
a) Count the number of each atom on the reactants and

products side and determine what requires
balancing.

b) Balance each element sequentially, using whole
numbers. It is often best to balance metals first.

mercury(II) oxide mercury + oxygenΔ

HgO Hg + O2
Δ

Hg: 1
O: 1

Hg: 1
O: 2

HgO Hg + O2
Δ

Oxygen atoms
need balancing
on the reactants
side.

2 HgO Hg + O2
Δ

Hg: 2
O: 2

Hg: 1
O: 2

Now Hg atoms
need balancing
on the products
side.

Writing and
Balancing
Chemical
Equations

4. Check after adding coefficients that all atoms still
balance. Adjust as needed (a 2 is needed in front of
Hg).

5. Do a final check to make sure all atoms now balance
on both sides of the equation.

2 HgO 2 Hg + O2
Δ

Hg: 2
O: 2

Hg: 2
O: 2

Note: always use the smallest whole
numbers!

4 HgO 4 Hg + 2 O2
Δ

Balanced but incorrect form!

Information in
a Chemical
Equation

© 2014 John Wiley & Sons, Inc. All rights reserved.

Information from a Chemical
Equation

• From the chemical equation below, how many moles of oxygen are
needed to burn 2 molecules of propane (C3H8)?

• a) 5 molecules of oxygen
• b) 6 molecules of oxygen
• c) 10 molecules of oxygen
• d) 15 molecules of oxygen

C3H8 + 5 O2 3 CO2 + 4 H2O

For every 1 molecule of propane,
5 molecules of O2 are needed to fully

react.
Two molecules of propane would then

require
2 x 5 = 10 molecules of oxygen.

Types of
Chemical
Equations

1. Combination reactions
2. Decomposition reactions
3. Single displacement reactions
4. Double displacement reactions
5. Oxidation-reduction (redox) reactions

(Chapter 17)

Reactions are classified into subtypes to aide in
predicting

the products of chemical reactions.

Reactions are classified into five major categories:

Combination Reactions Two reactants combine to give a single product.
A + B AB

Decomposition
Reactions

A single reactant breaks down (decomposes) into
two or more products

AB A + B

Single Displacement
Reactions

One element (A) reacts with a compound (BC) to replace
one element in the compound, giving a new element (B)
and a different compound (AC).

General Types of Single Displacement Reactions

Double Displacement
Reactions

Two compounds exchange partners with one
another to yield two new compounds.

AB + CD AD + CB

General Types of Double Displacement Reactions

Double Displacement
Reactions

Two compounds exchange partners with one
another to yield two new compounds.

AB + CD AD + CB

General Types of Double Displacement Reactions Writing Reaction Equations Practice
1. Write the reaction equation between aqueous

solution of hydroiodic acid and sodium
hydroxide.

2. First convert names to chemical formulas and
determine the type of reaction.

HI (acid)/NaOH(base)

Neutralization Reaction
acid + base salt + water
HI (aq) + NaOH (aq) NaI (aq) + H2O (l)
Salt formula must charge balance (Na+ and I–)

Heat in
Chemical
Reactions

Terminology

Energy transfer and changes accompany any chemical reaction

Heat of reaction: quantity of heat actually produced during a chemical reaction.
Units: kilojoules (kJ) or kilocalories (kcal)

Exothermic reactions: release heat. H2 (g) + Cl2 (g) 2 HCl (g) + 185 kJ
Heat can be treated as a product

Endothermic reactions: absorb heat. N2 (g) + O2 (g) + 181 kJ 2
NO (g)
Heat can be treated as a product

C (s) + O2 (g) CO2 (g) + 393 kJ
1 mol of C reacts with 1 mol of O2 to provide 1 mol of CO2 and 393 kJ
of heat
are released.

Heat in Chemical Reactions
Equations Practice

Heat as an Energy
Transfer

Vehicle in Nature

Graphical
Representations of

Endothermic
Reactions

•Products are at a higher
potential energy than
reactants.
•Activation energy: Amount
of energy needed to initiate a
chemical reaction.

Reaction Coordinate
Diagram

Graphical
Representations of

Exothermic
Reactions

•Products are at a lower
potential energy than
reactants.
•Activation energy: Amount
of energy needed to initiate a
chemical reaction.

Reaction Coordinate
Diagram

Reading
Review

How do you know if a reaction is a
combustion reaction?

What is an endothermic reaction?

What is an exothermic reaction?

What are the four types of chemical
equations?.

How do you know if an equation is
balanced?

  • Slide 1
  • Chemical Equations
  • Chemical Equations
  • Symbol Summary
  • Law of Conservation of Mass
  • Writing and Balancing Chemical Equations
  • Writing and Balancing Chemical Equations
  • Information in a Chemical Equation
  • Information from a Chemical Equation
  • Types of Chemical Equations
  • Combination Reactions
  • Decomposition Reactions
  • Single Displacement Reactions
  • Double Displacement Reactions
  • Double Displacement Reactions
  • Heat in Chemical Reactions Terminology
  • Heat as an Energy Transfer Vehicle in Nature
  • Graphical Representations of Endothermic Reactions
  • Graphical Representations of Exothermic Reactions
  • Reading Review

CH1000
Fundament
als of
Chemistry
Module 2 – Chapter 6

Common and Systematic Names

• Chemical nomenclature is the systematic naming of chemical compounds
• Common names are historical names of compounds which are not based

on systematic rules
• Common names are often used because systematic names are too long

and technical for everyday use
• Chemists prefer systematic names that precisely identify the chemical

composition of compounds.
• Example CaO

• Common name: lime
• Systematic name: calcium oxide

Naming
Flowchart

We will focus on nomenclature of inorganic compounds

Elements and Ions

• The formula for most elements is the symbol of the element off of
the periodic table.
• Diatomic molecules are an exception:

• Two other elements also exist in polyatomic arrangements:

Naming Anions

•Remember from Chapter 5
that any neutral atom that
gains an electron is called
an anion
•When naming anions,
change the element ending
to -ide

Symbols
of the
Elements
•Each element has an
abbreviation called a symbol.

•The first letter of a symbol
must always be capitalized.

•If a second letter is needed, it
should be lowercase.

Predicting Ion
Charge from
Periodic Table

•Metals form cations
•The positive charge is equal
to the group number

Predicting Ion
Charge from
Periodic Table

•Nonmetals form anions
•The negative charge is equal
to 8 – the group number

Writing Formulas from Names of Ionic Compounds

•Ionic compounds contain both a cation and
an anion.

•Ionic compounds must have a net charge of
0

•The sum of charges of the cations and
anions in an ionic compound equal 0

•Rules for writing formulas for ionic
compounds:
• Write the metal ion followed by the

nonmetal ion formula
• Combine the smallest whole numbers

of each ion to provide an overall
charge equal to zero

• Write the compound formula for the
metal and nonmetal, using subscripts
determined from Step 2 for each ion

Naming Ionic
Binary
Compounds
•Binary compounds containing
a metal which forms only one
cation

•By convention, the cation is
written/named first followed
by the anion

•Rules for naming binary ionic
compounds:
• Name the cation
• Write the anion root and

add the –ide suffix

Naming
Compounds
Containing
Metals with

Multiple
Charges

•Rules for Naming Compounds Involving Metals that Could Form
Multiple Charges
• Write the cation name.
• Write the cation charge in Roman numerals in parentheses.
• Write the root of the anion and use the –ide suffix.

•Exception: for metals that only form two cations, a Latin root with
either an –ous or –ic suffix can also be used.

Formula Name Classical Name Formula Name Classical Name

Cu+ Copper(I) cuprous Sn2+ Tin(II) stannous

Cu2+ Copper(II) cupric Sn4+ Tin(IV) stannic

Fe2+ Iron(II) ferrous Pb2+ Lead(II) plumbous

Fe3+ Iron(III) ferric Pb4+ Lead(IV) plumbic

Naming Molecular
Compounds

•Molecular compounds contain two nonmetals
•Rules for naming molecular compounds:
• Write the name for the first element, including the appropriate prefix

(mono is rarely used).
• Write the name for the second element, including the appropriate prefix

and -ide ending (mono is used for the 2nd element).

Prefix Number Prefix Number
mono 1 hexa 6

di 2 hepta 7

tri 3 octa 8
tetra 4 nona 9
penta 5 deca 10

Naming Binary
Acids

Rules for naming Binary Acids:

Write the prefix hydro followed by the root
of the second element and add an –ic suffix Add the word acid

Hydrogen is always written first in an acid formula.

This is the indicator that it is an acid

Certain binary compounds containing hydrogen behave as acids in water and have special names.

HCl(g) is hydrogen chloride HCl(aq) is hydrochloric acid

Naming
Polyatomic Ions

•A polyatomic ion is anion that
contains two or more
elements
•The names, formulas and
charges of common
polyatomic ions should be
learned.
•Rules for naming compounds
containing polyatomic ions
• Name the cation
• Name the anion

Name Formula Charge Name Formula Charge
Acetate C2H3O2- -1 Cyanide CN- -1

Ammonium NH4+ +1 Dichromate Cr2O72- -2
Hydrogen
Carbonate HCO3

– -1 Hydroxide OH- -1

Hydrogen
Sulfate HSO4

– -1 Nitrate NO3- -1

Bromate BrO3- -1 Nitrite NO2- -1
Carbonate CO32- -2 Permanganate MnO4- -1
Chlorate ClO3- -1 Phosphate PO43- -3

Chromate CrO42- -2 Sulfate SO42- -2
Sulfite SO32- -2

Naming
Oxyanions
•Oxyanions are polyatomic ions that contain
oxygen
•Often end in suffix –ate or –ite
•-ate compounds contain more O atoms
than ite compound(s)
•For elements that form multiple ions with
oxygen, prefixes are also needed:
• Per: add one oxygen to the –ate root
• Hypo – subtract one oxygen from the –

ite root

Anion
Formula Anion Name

Anion
Formula Anion Name

ClO- hypochlorite HClO hypochlorous acid

ClO2- chlorite HClO2 chlorous acid

ClO3- chlorate HClO3 chloric acid

ClO4- perchlorate HClO4 perchloric acid

More Complicated
Polyatomics

•Inorganic ions can be formed from more
than 3 elements
•The same method is used as before:
• Identify the ions and name in order
• Cations before anions

Compound Ions Name

NaHCO3 Na+; HCO3-
Sodium hydrogen

carbonate

NaHS Na+; HS- Sodium hydrogen sulfide

MgNH4PO4
Mg2+; NH4+;

PO43-
Magnesium
ammonium
phosphate

NaKSO4 Na+; K+; SO42-
Sodium potassium

sulfate

Naming Acids

• Acids generally begin with hydrogen
• To recognize oxyacids, make sure:
• H is the first element in the formula
• The compound contains a polyatomic ion with oxygen

• The following modifications are made to the name of the
acid:
• -ate ions are changed to –ic acids
• -ite ions are changed to –ous acids
• -ic acids contain one more oxygen than –ous acids

Naming Acids
Flowchart

Reading
Review

What type of ions do metals form?

What type of ions do nonmetals
form?

What is the chemical formula for
potassium sulfide?

Name the compound CrCl3.

Acids often begin with what element?

  • Slide 1
  • Common and Systematic Names
  • Naming Flowchart
  • Elements and Ions
  • Naming Anions
  • Symbols of the Elements
  • Predicting Ion Charge from Periodic Table
  • Predicting Ion Charge from Periodic Table
  • Writing Formulas from Names of Ionic Compounds
  • Naming Ionic Binary Compounds
  • Naming Compounds Containing Metals with Multiple Charges
  • Naming Molecular Compounds
  • Naming Binary Acids
  • Naming Polyatomic Ions
  • Naming Oxyanions
  • More Complicated Polyatomics
  • Naming Acids
  • Naming Acids Flowchart
  • Reading Review
Order a unique copy of this paper
(550 words)

Approximate price: $22