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Chapter 3
Formulas, Equations,
and Moles
3.1 Balancing Equation
Balancing Equations
Chemical Reaction
Reactants  Products
 Steps:

 Write
the unbalance equation
 Find an appropriate coefficient and place in front of
the formula unit to balance the equation
 Reduce the coefficients to the smallest ratio
 Check your answer
Balancing Equation

Propane, C3H8 is a colorless gas often used as a heating
and cooking fuel in campers and rural homes. Write a
balance equation for the combustion reaction of
propane with oxygen to yield carbon dioxide and water.
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Step 1: Write out the unbalanced equation
Step 2: Balance C atoms
Step 3: Balance C and H
Step 4: Balance C, H and O
Reduce coefficients to the smallest ratio
Check answer
Examples
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When solid potassium reacts with liquid water, the
products are hydrogen gas and potassium hydroxide,
the later remains dissolved in the water. From these
information, write the balanced equation for the
reaction
Write the balanced equation for the following reaction
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Solid carbon reacts with gaseous oxygen to form gaseous
carbon dioxide

Glass is sometimes decorated by etching patterns on its
surface. Etching occurs when hydrofluoric acid (in a aqueous
solution of HF) reacts with the silicon dioxide in the glass to
form gaseous silicon tetrafluoride and the liquid water.
Chemical symbols on Different Levels
Chemical symbols represent both a microscopic and a macroscopic level
Microscopic level – chemical symbols represent the behavior of
individual atoms and molecules
E.g
Atoms and molecules are much too small to be seen  use
microscopic behavior to describe
2 H2 + O2  2 H2O
Two molecules of hydrogen react with one molecule of oxygen to yield
two molecule of water.
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Help to understand how reaction occur
Chemical symbols on Different Levels
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Macroscopic level –
formula and equations
represent the large-scale
behavior of atoms and
molecules that give rise to
observable properties


Deal with macroscopic
behavior in the laboratory
E.g
Weighing
amount of reactants, place
them in a flask and observe
visible changes.
3.3 Avogadro Number and The Mole

Described the number of objects present in a
dozen
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Based on C-12
1 mol of anything = 6.02 x 1023 units of that
substance (atoms, molecules, particles, dollars
etc..)

Avogadro’s number or nA
Avogadro’s number
Moles

or Formula unit
Mass

Molecular mass (molecular weight) – sum of the atomic

masses of all atoms in a molecule (covalent molecule)
Formula mass – sum of atomic masses all atoms in a formula
unit of any substances (ionic salts)

Units = amu
Formula Mass =
Number of atoms
mass
of 1st element in x Atomic
of
Chemical compound 1st element
+
Number of atoms
x Atomic mass
of 2nd element in
of
2nd
element
Chemical compound
Molecular Weight
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What is the molecular weight of CH4?
Calculate molecular weight of sulfur dioxide, a
gas produced when sulfur containing fuels are
burned.
Calculate the formula weight of the substance
below (record to two decimal places)

Ba3(PO4)2
Molar Mass

Like density, it’s a ratio of two numbers.
 Mass per one mole of substance
 Unit

=> g/mol
Replaced with molecular weight because molar
mass described the concept more accurately
E.g
Molecular weight
18.02 amu
Molar Mass
18.02 g/mol
Molar Mass

Calculate the molar mass for each of the
substance below (record to two decimal places)
C2H4
 NH3
 C2H5Cl
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3.4 Stoichiometry Chemical Arithmetic

Recall
Molar mass is the mass per one mole of a substance
 1 mol of substance = # g
 Use as a conversion factor
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E.g
Moles
Grams
if the molar mass of NaCl = 58.44 g/mol


Grams
Moles
or
Stoichiometry

Calcium carbonate (also called as calcite,
CaCO3), is the principal mineral found in
limestone, marble, chalk, pearls and the shells of
marine animals such as clams
Calculate the molar mass of calcium carbonate
 A certain sample of calcium carbonate contains 4.86
mol. What is the mass in grams of this sample?
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Stoichiometry
How many moles of sucrose are in a table spoon of sugar
that contains 2.85 g? The molar mass of sucrose is
342.0 g/mol
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Calculate the mass of 4.85 mol of acetic acid,
HC2H3O2. Vinegar is a dilute solution of acetic acid
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How many water molecules are in a 10.0g sample of
water?
Mass Calculation
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Steps for Calculating the Masses of Reactants and
Product in Chemical Reactions
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Step 1: Balance the equation for the reaction
Step 2: convert the masses of reactants or product to moles
Step 3: use the balanced equation to set up the appropriate
mole ratio(s)
Step 4: Use the mole ratio(s) to calculate the number of
moles of the desired reactant or product
Step 5: Convert from moles back to masses (of the desired
reactant or product)
Example
Aqueous solution of sodium hypochlorite
(NaOCl), best known as household bleach, are
prepared by reaction of sodium hydroxide with
chlorine:
…..NaOH(aq) + ….Cl2(g)  …NaOCl(aq) +
NaCl(aq) + ….H2O(l)
How many gram of NaOH are needed to react
with 2.5g Cl2?
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Limiting Reactant
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Limiting Reactant (limiting reagent): is
the reactant that is completely consumed
in a chemical reaction and limits the
amount of product.
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Excess Reactant: Any of the other
reactants still present after determination
of the limiting reactant.
Steps in Determination Limiting
Reagent
Check to be sure you have a balanced equation
 aA + bB  c C
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i.
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ii. Convert the amount of reagent one that was given
into the number moles product that you could form if that
reagent was completely consumed.
 gA  moles A  moles C
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iii. Convert the amount of reagent two that was given
into the number moles of product that you could form if that
reagent was completely consumed.
 gB  moles B  moles C
Steps in Determination Limiting
Reagent
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iv. The reactant that produced the LEAST amount
of product in step 2 or 3 will the limiting reagent.
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v. Convert the LEAST moles into the number
grams product and that will be your theoretical
yield.
 Theoretical yield is the amount of product that
can be made in a chemical reaction based on the
amount of limiting reagent
Excess Reactant
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To determine how much of reactants are left
over from the reaction:
Larger moles of product – smaller moles of product
 Convert this moles of product to the g of the excess
reactant
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Example
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If we have 42.5 g Mg and 33.8 g O2, what
is the limiting reactant?
2 Mg(s) + O2(g)  2 MgO(s)
Example

Ammonia, NH3, can be synthesis by the
following reaction:
2 NO(g) + 5 H2(g)  2 NH3(g) + 2H2O(g)
Starting with 86.8 g NO and 25.6 g H2, find
the theoretical yield of ammonia in grams
 Calculate the number grams of excess reactant
that are unused from the reaction
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3.5 Percent Yield

Recall
 Theoretical
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yield is the amount of product
that can be made in a chemical reaction
based on the amount of limiting reagent
actual yield
Percent yield = --------------------- x 100
theoretical yield
Examples
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Ethyl alcohol is prepared industrially by the
reaction of ethylene, C2H4 with water. What is
the percent yield of the reaction if 4.6 g of
ethylene gives 4.7g of ethyl alcohol?
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C2H4(g) + H2O(l)  C2H6(l)
Examples

Titanium (IV) oxide is a white compound used
as a coloring pigment. Solid titanium (IV) oxide
can be prepared by reacting gaseous titanium
(IV) chloride with oxygen gas. A second
product of this reaction is chlorine gas.
Suppose 6.71 x 103 g of titanium (IV) chloride is
reacted with 2.45 x103 g of oxygen. Calculate the
mass of titanium (IV) oxide that can form
 If the percent yield of TiO2 is 75%, what mass is
actually formed?
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3.11 Percent Composition and
Empirical Formula

Formula of a compound represents the relative
numbers of various types of atoms present.
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E.g CO2, H2O etc..
Empirical formula: simplest formula
Molecular formula: the actual formula of a
compound
Examples
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In each case below, the molecular formula for a
compound is given. Determine the empirical formula
for each compound
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C6H6. This is the molecular for benzene; a liquid commonly
used in industry as a starting material for many important
products
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H2O2. This is hydrogen peroxide, a substance commonly
diluted with water and used as a disinfectant
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CCl4. This is carbon tetrachloride, an organic solvent
Percent Composition
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Percent composition – the mass present of
each element present in a compound
Atomic mass of an element
% composition
=……………………………….. X 100
Molar mass of the compound
Percent Composition

To calculate the percent composition
(percentage composition) of a compound
Calculate the molar mass of the compound
 Calculate the total mass of each element present in
the formula of the compound
 Calculate the percent composition (percentage
composition):
% by weight (mass) of element = (total mass of
element present ÷ molecular mass) x 100
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Examples
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Glucose or blood sugar, has the molecular
formula C6H12O6. What is the empirical
formula, and what is the percent composition?
Determination of an empirical
formula
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Step 1: Obtain the mass of each element present (in
grams)
Step 2: Determine the number of moles of each type
of atom present
Step 3: Divide the number of mole of each element by
the smallest number of moles to convert the smallest
number to 1. If all of the numbers so obtained are
integers, these are the subscripts in the empirical
formula. If one or more of these number are not
integers, go on to step 4
Determination of an empirical
formula
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Step 4: Multiply the numbers you derived in
step 3 by the smallest integer that will convert all
of them to whole numbers. This set of whole
numbers represents the subscripts in the
empirical formula
Example
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In a lab experiment it was observed that 0.6884 g of
lead combines with 0.2356 g of chlorine to form a
binary compound. Calculate the empirical formula of
this compound
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When a 2.00 g sample of iron metal is heated in air, it
reacts with oxygen to achieve a final mass of 2.573 g.
Determine the empirical formula for this iron oxide
Examples
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What is the empirical formula for sodium
thiosulfate with the percent composition of
30.36% O, 29.08% Na and 40.56%S
Example
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Cisplatin, the common name for a platinum
compound that is used to treat cancerous
tumors, has the composition (mass percent)
65.02% platinum, 9.34 % nitrogen, 2.02%
hydrogen, and 23.63% chlorine. Calculate the
empirical formula for csiplatin
Molecular Formula
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Molecular formula – gives the actual numbers
of atoms in a molecule
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Molecular formula = n x empirical formula
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may be the same as the empirical formula
Molecular Formula
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a multiple of the empirical formula (or # of
empirical unit) is defined:
Emperical mass
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Multiple or n
= --------------------------
( # of empirical unit)
Molecular mass
Examples
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Ribose, a sugar present in the cells of all living
organism, has a molecular mass of 150 amu and
the empirical formula CH2O. What is the
molecular formula?
Examples
A compound used as an additive for gasoline to
help prevent engine knock shows the following
percent composition
71.65 % Cl
24.27 % C
4.07 % H
The molar mass is known to be 98.96 g/mol.
Determine the empirical formula and the
molecular formula for this compound
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Combustion Analysis
CxHyOz (O can be replaced with any other
element)
 gCO2  moles CO2  moles C
 gH2O  moles H2O  moles H
 g of O or unknown element = grams of sample
– (g of H + g of C)
 Follow the steps in determination of an
empirical formula
Combustion Analysis

Caproic acid, the substance responsible for the
aroma of dirty gym socks and running shoes,
contains carbon, hydrogen and oxygen. On
combustion analysis, a 0.450 g sample of caproic
acid gives 0.418 g of H2O and 1.023 g of CO2.
what is the empirical formula of caproic acid?
If the molecular mass of caproic acid is 116.2
amu, what is the molecular formula?
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Determine the percent composition of caproic acid?
Examples
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Coniine, a toxic substance isolated from poison
hemlock, contains only carbon, hydrogen and
nitrogen. Combustion analysis of a 5.024 mg
sample yields 13.90 mg of CO2 and 6.048 mg of
H2O. What is the empirical formula of coniine?
EXAM 1
(Chapter 1-3, except 3.7-3.10)