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• Each element has been assigned a unique one-, two-, or
three-letter symbol for its identification
• The first letter of a symbol is always capitalized
• Only recently discovered elements that don’t yet have
permanent names are given three-letter symbols
• Uncombined elements are written as monatomic (without
a subscript)
– Examples: Fe, C, Al
• Diatomic elements – 2 of the same atoms (HOFBrINCl’s)
– Examples: H2, O2
• The atomic mass of an element is given in
atomic mass units (amu)
• Found on the periodic table
• The mass of 1 mole of an element is equal
to its atomic mass (periodic table) in
grams
• The sum of the atomic masses of its
atoms
Example:
CO2 =
Mass Examples
• Find the atomic or formula mass of the
following substances:
(To the nearest tenth)
1.
2.
3.
4.
5.
6.
C
Br
Cl2
H2O
Ca3PO4
Ca(OH)2
• The mass of 1 mole of the substance
• Measured in grams
Example:
1.0 mole of CO2 = 44g
• Use the gram formula mass to convert
between moles and mass
Examples:
1. 1.5 moles of CO2 = _______ g
2. 24 g of C =
_______ moles
3. 3.0 moles of Br = _______ g
4. 5.0 g of Al =
_______ moles
• If you have an ideal gas at standard
temperature and pressure (STP) – 0oC and
1atm
1 mole = 22.4 L
Mole – Volume Examples
1. 1.5 moles of CO2 (g) = _______ L
2. 38 L of Ne(g) =
_______ moles
3. 3.0 moles of O2(g) =
_______ L
4. 11 L of N2(g) =
_______ moles
• 1 dozen = 12
• 1 mole = 6.02 x 1023 (Avagadro’s Number)
• How big is a mole?
– 602,000,000,000,000,000,000,000
– A mole of marbles would spread over the surface of the
earth, and produce a layer about 50 miles thick
– A mole of sand, spread over the United States, would
produce a layer 3 inches deep
– A mole of dollars could not be spent at the rate of a
billion dollars a day over a trillion years
• He stated that equal volumes of all
gases at the same temperature
and pressure contain the same
number of molecules (Acogadro’s
Principle)
• Later his work led to the realization Amedeo Avogadro
that a molecular mass in grams
(mole) of any substance contains
the same number of molecules
(6.02 x 1023)
1 mole of atoms = 6.02 x 1023 atoms
1 mole of particles = 6.02 x 1023 particles
1 mole of molecules = 6.02 x 1023 molecules
1 mole of compounds = 6.02 x 1023 compounds
Mole – Number Examples
1. 4.00 moles of NaCl = _______ molecules
2. 0.50 mole of MgBr2 = _______ molecules
3. 3.01 x 1023 NaCl molecules = _______ moles
4. 6.02 x 1021 Zn atoms = _______ moles
Empirical Formulas
• Represents the simplest integer ratio in
which the atoms combine to form a
compound (the reduced form)
• All ionic formulas are written as empirical
formulas
Molecular Formulas
• The actual numbers of the atoms in a molecule
• May be a multiple of the empirical formula
Examples:
• What are the empirical formulas of the following
molecules?
a. H2O2
b. C6H12O6
c. CCl4
Percent Composition
• Composition in terms of the percentage of
each component present
Example: H2SO4
– Step 1: Calculate the formula mass of H2SO4
– Step 2: Find the percent mass of each element
Percent Mass Example
• Calculate the percent composition of
oxygen in CO2
Water of Hydration
• Percentage of water in the crystal
– Hydrate – a compound that contains water
– Anyhdrous – hydrate without the water
Water of Hydration Example
• Find the water of hydration in
CuSO45H20
• Step 1 - Find the formula mass
• Step 2 - Percentage of water
Water of Hydration Example
• Calculate the water of hydration in
Na2CO310H2O
Determining Mass Ratios from
Formulas
• Example: Find the mass ratio for a
compound with the empirical formula CH2.
Determining the Molecular Formula from
the mass and the empirical formula
A compound has a molecular mass of 180amu
and an empirical formula of CH2O. What is its
molecular formula?
• Step 1: Determine the molecular mass of the
empirical formula
• Step 2: Divide the molecular mass of the
compound by the mass of the empirical formula.
• Step 3: Multiply the subscripts in the empirical
formula by your answer to step 2
Determining the empirical formula
from percent composition
What is the empirical formula of a compound that
consists of 58.80% barium, 13.75% sulfur, and
27.45% oxygen by mass?
• Step 1: Assume that the mass of the sample is 100g
• Step 2: Convert the masses into moles
• Step 3: Find the smallest whole numbers ratio (divide
each number from step 2 by the smallest one)
Chemical Reactions
• Reactant – Substance that enters into a
reaction, written to the left of the arrow,
starting material
• Product – substance that is produced by
the reaction, written to the right of the
arrow, end material
• Example: HCl + NaOH  NaCl + H2O
– Reactants:
– Products:
• Reaction that requires energy (heat) in
order to occur – heat enters
• Heat is absorbed
• Heat is a reactant
• Surroundings will feel cold because heat
has been absorbed from the surroundings
• Example: H2O(s) + heat  H2O(l)
• Reaction that produce energy (heat) when
they occur – heat exits
• Heat is released, given off
• Heat is a product
• Surroundings will feel hot because heat
was released to the surroundings
• Example: H2O(l)  H2O(s) + heat
 Synthesis
 Decomposition
 Single Replacement
 Double
Replacement
• Direct combination
• Substances combine to form a new
compound
• Produces 1 product
Examples:
A + B  AB
2H2(g) + O2(g)  2H2O(l)
• Break down of a compound into simpler
parts
• Starts with 1 reactant
Examples:
AB  A + B
2H2O(l)  2H2(g) + O2(g)
• One substances switches spots with
another
• Element + compound makes a new
element plus a new compound
Examples:
A + BC  B + AC
Cu(s) + 2AgNO3(aq)  2Ag(s) + Cu(NO3)2
*copper replaces silver
• Exchange of ions
• Everything gets a new "partner"
• Compound + compound makes new
compound + new compound
Examples:
AB + CD  AD + CB
AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq)
Balancing Equations
• Coefficient - a number, placed before formulas
to indicate the ratios of moles involved in a
reaction
• Equations must be balanced in accordance with
the Law of Conservation of Mass
• The mass of both sides of the arrow must be
equal
• You must have an equal number of each type of
atom on both sides of the equation
Example: 2H2 + O2  2H2O
Balancing Examples
1.
____ Na + ____ H2O  ____ NaOH + ____ H2
2.
____ CaO + ____ H2O  ____ Ca(OH)2
3.
____ Al + ____ O2  ____ Al2O3
4.
___ PbCl2 + ___ Al2(SO4)3  __ PbSO4 + __ AlCl3
5.
____ Na + ____ O2  ____ Na2O
Determining the Missing Mass in
Equations
If 103.0g of potassium chlorate (KClO3) are
decomposed to form 62.7g of potassium
chloride (KCl) and oxygen gas (O2) according
to the equation 2KClO3  2KCl + 3O2, how
many grams of oxygen are formed?
(Hint: Remember the mass of the products must
equal the mass of the reactants)
How many grams of Fe are needed to react with
8.0g of O2 to produce 28.9g of Fe3O4 according
to the equation 3Fe + 2O2  Fe3O4?
Equation Problems
• Using the balanced equations and mole
conversions you can solve for variety of
problems
• Remember that the coefficients used
represent mole ratios
Given the following reaction answer the
questions below:
2C2H6 + 7O2  4CO2 + 6H2O
1. How many moles of CO2 are produced
when 2.0 moles of C2H6 reacts?
2. How many moles of H2O are produced
when 4.0 moles of C2H6 reacts?
3. How many moles of H2O are produced
when 5.0 moles of C2H6 combusts?
Given the following reaction:
____ H2 + ____ Cl2  _____ HCl
1. The production of 37g of HCl would
require how many moles of H2?
2. If 20.L of H2 completely reacts how many
grams of HCl would be produced?
Use the balanced equation to answer the
questions below:
____ C2H4 + ____ O2  ____ CO2 + ____ H2O
1. How many liters of O2 are used to
produce 1.0 mole of H2O?
2. How many liters of CO2 are produced
when 9.00 liters of O2 is consumed?
3. How many liters of C2H4 are needed to
produce 10.0 liters of CO2?