Download Molecular and Empirical Formulas

The Mole and Chemical Equations
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Mole and Chemical Equations
Concepts to Master
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Be able to covert the moles of a
substance to atoms (and vice
versa).
Be able to convert grams of a
substance to moles (and vice
versa).
Be able to balance equations.
Be able to convert the grams of
one substance to the grams of
another substance using a
balanced equation.
Be able to convert the moles of
one substance to the moles of
another substance using a
balanced equation.
Be able to determine the empirical
formula from gram analysis data.
Be able to determine the empirical
formula from % composition.
Be able to determine the
molecular formula from an
empirical formula.
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Be able to predict the products of
a chemical reaction using Table J
(for single replacement) and Table
F (for double replacement).
How do you know that a reaction
has come to completion?
Be able to use Table F to
determine solubility.
Is a combustion reaction
endothermic or exothermic?
Which reactions release energy?
Which reaction absorb energy?
When energy is absorbed, is it a
product or reactant?
What is the conservation of mass
law and why is it important to this
chapter?
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Labs
Vocab
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Addition reaction
Analysis reaction
Aqueous
Atomic Mass
Coefficient
Combination reaction
Combustion
Composition reaction
Decomposition
Double replacement
Empirical Formula
Endothermic
Exothermic
Formula mass
Formula units
Gas
Liquid
Metathesis
Molar Mass
Mole
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Molecular Formula
Phases of matter
Precipitate
Products
Reactants
Single replacement
Solid
Stoichiometry
Structural formula
Synthesis
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Celebrating the Mole
Burning Mg
Balance Beads
Mole and Mass
Relationships
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THE MOLE!!!
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A mole is the number of protons or neutrons needed to make up one gram of
any element.
• 1 mole = 602 000 000 000 000 000 000 000 = 6.02x1023 particles
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It doesn't matter if it's atoms, donuts or basketballs, a mole is the same
number of things. (Like a dozen of anything is always 12.)
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Atomic weight of an element is the average of the masses of it’s isotopes.
– For hydrogen, this value is 1.00794 atomic mass units, for chlorine its 35.453 amu
– But since a mole of amu’s is a gram, we can use these exact same numbers for
atomic weights in “grams per mole.” Hydrogen weighs 1.00794 grams per mole of
atoms; chlorine weighs 35.453 grams per mole of atoms, etc.
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No one ever deals with just one or two atoms. We deal with enough to see,
which means trillions and trillions of them – measurable numbers or fractions
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of moles.
Molar Mass = Atomic Mass found
on periodic table
Formula Composition Moles
g/mol
H2O
Total of 3 atoms in
each molecule.
1 oxygen and 2
hydrogen
1 mole of
1 (16.00) = 16.00 g/mol
oxygen for every 2 (1.01) = 2.02 g/mol
2 moles of
18.02 g/mol
hydrogen
CO2
Total of 3 atoms in
each molecule.
2 oxygen and 1
carbon
2 moles of
1 (12.00) = 12.00 g/mol
oxygen for every 2 (16.00) = 32.00 g/mol
1 mole of
44.00 g/mol
carbon
O2
Total of 2 atoms in
each molecule.
2 oxygen
2 moles of
oxygen
2 (16.00) = 32.00 g/mol
C6H12O6
Total of 24 atoms in
each molecule.
6 oxygen, 12
hydrogen, and 6
carbons
6 moles of
oxygen, 12
moles of
hydrogen, and 6
moles of carbon
6 (16.00) = 96.00 g/mol
12 (1.01) = 12.12 g/mol
6 (12.00) = 72.00 g/mol
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180.12 g/mol
Mass to Mole Conversions
• How many moles are present in 152 g of
H2SO4?
• How many moles are present in 13.3 kg of
KNO3?
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Mole to Mass Conversions
• How many grams are present in 252.3 moles of
CO2?
• How many grams are present in 1.3 moles of
H2O?
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Mole to Particle Conversions
• Use Avagadro’s number in your dimensional analysis
• How many molecules are present in 8.2 moles of
ammonia?
• How many molecules are present in 10 moles of water?
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Particle to Mole Conversions
• Use Avagadro’s number in your dimensional analysis
• How many moles are present in 2.359 x 1036 molecules
of carbonic acid?
• How many moles are present in 250 x 1023 atoms of
neon?
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Molecular and Empirical
Formulas
• An empirical
formula for a
compound shows
the simplest
whole number
ratio for the
elements in the
compounds.
• A molecular
formula shows
the molecule as
it actually exists
In nature.
CMPD
Empirical Molecular
Water
H2O
H2 O
Hydrogen HO
peroxide
glucose
CH2O
H2O2
Methane
CH4
CH4
Ethane
CH3
C2H6
Octane
C4H9
C8H18
C6H12O6
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Determining an Empirical Formula from Gram Analysis
A compound was analyzed and found to contain 13.5 g Ca, 10.8 g O, and
0.675 g H. What is the empirical formula of the compound?
1. Start with the number of grams of each element, given in the problem.
2. Convert the mass of each element to moles using the molar mass from the
periodic table.
H
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Determining an Empirical Formula from Gram Analysis
3. Divide each mole value by the smallest number of moles
calculated. Round to the nearest whole number.
4. This is the mole ratio of the elements and is represented
by subscripts in the empirical formula.
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Steps for Determining an Empirical Formula
1. Start with the number of grams of each element, given in the
problem.
1. If percentages are given, assume that the total mass is 100 grams so
that the mass of each element = the percent given.
2. Convert the mass of each element to moles using the molar mass
from the periodic table.
3. Divide each mole value by the smallest number of moles calculated.
4. Round to the nearest whole number. This is the mole ratio of the
elements and is represented by subscripts in the empirical formula.
5. If the number is too far to round (x.1 ~ x.9), then multiply each solution
by the same factor to get the lowest whole number multiple.
1. If one solution is 1.5, then multiply each solution in the problem by 2 to
get 3.
2. If one solution is 1.25, then multiply each solution in the problem by 4 to
get 5.
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Determining an Empirical Formula from % Composition
NutraSweet is 57.14% C, 6.16% H, 9.52% N, and 27.18% O. Calculate the empirical
formula of NutraSweet and find the molecular formula. (The molar mass of
NutraSweet is 294.30 g/mol)
1.
Start with the number of grams of each element, given in the problem. If percentages
are given, assume that the total mass is 100 grams so that
the mass of each element = the percent given.
2.
Convert the mass of each element to moles using the molar mass from the periodic
table.
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Determining an Empirical Formula
from % Composition
3. Divide each mole value by the smallest number of moles calculated. Round to the
nearest whole number.
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Determining an Empirical Formula
from % Composition
4. If the number is too far to round (x.1 ~ x.9), then multiply each solution by the same
factor to get the lowest whole number multiple.
5. This is the mole ratio of the elements and is represented by subscripts in the empirical
formula.
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Determining Empirical Formula
• Vitamin C (ascorbic acid)
contains 40.92 % C, 4.58 % H,
and 54.50 % O, by mass. The
experimentally determined
molecular mass is 176 amu.
What is the empirical and
chemical formula for ascorbic
acid?
• In 100 grams of ascorbic acid
we would have:
– 40.92 grams C
– 4.58 grams H
– 54.50 grams O
• This would give us how many
moles of each element?
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Determining Empirical Formula
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Determine the simplest whole number ratio by dividing by the
smallest molar amount (3.406 moles in this case - see Oxygen):
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Determining Empirical Formula
• Multiply the relative amounts of each atom
by '3', to get whole numbers for each
atom.
C = (1.0)*3 = 3
H = (1.333)*3 = 4
O = (1.0)*3 = 3
C3H4O3 = Empirical Formula
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Determining Molecular Formula from
Empirical Formula
• We are told that the experimentally determined molecular mass is
176 amu or g/mol. What is the molecular mass of our empirical
formula?
• (3*12.011) + (4*1.008) + (3*15.999) = 88.062 g/mol
• The molecular mass from our empirical formula is signficantly
lower than the experimentally determined value. What is the
ratio between the two values?
• (176 amu/88.062 amu) = 2.0
• Multiplied the subscripts in the empirical formula by '2', to get
the molecular formula.
• 2* C3H4O3 = C6H8O6
• C6H8O6 = Molecular Formula
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Determining Molecular Formula from
Empirical Formula
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Need to know the molar mass of the molecule
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Calculate the mass of the empirical formula
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Divide the mass of empirical formula by molar
mass of molecule.
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Multiply all the atoms in the empirical formula
(subscripts) by this ratio to find the molecular
formula.
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Formulas (in summary)
• Formulas are used to express
compounds/molecules and
elements
• Types of Formulas
– Empirical
• simplest whole number ratio of
atoms of elements within a molecule
CAFFEINE
– Molecular
• actual number of atoms within the
molecule as found in nature
– Structural
• A formula in which atoms are shown
in relation to each other, and bonds
are shown as lines or dashes
– Formula Units
• Only ionic compounds
• Number of ions within the
compound
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Types of Formulas
Formula
Type?
# of Formula
Units
CaSO4
C6H12O6
CH2O
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Chemical Reactions
– Chemical reactions occur when bonds between
the outermost parts of atoms are formed or
broken
– Chemical reactions involve changes in matter,
the making of new materials with new
properties, and energy changes.
– Symbols represent elements.
– formulas describe compounds.
– chemical equations describe a chemical
reaction.
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Chemical Equations
• Equations explain chemical reactions.
– Reactants are the substances that exist before the
chemical reaction takes place (before the arrow).
– Products are the new substances that exist after the
chemical reaction takes place (after the arrow).
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Symbols used in Chemical Equations
textbook pg 206
Symbol
Explanation
+
Used to separate two reactants or two products
→
Means “yields” or “produces”, separates reactants from products
↔
Used in place of → for reactions that can go either forwards or backwards
(s)
Appears after the formula, the compound or molecule is in the SOLID
state.
(l)
Appears after the formula, the compound or molecule is in the LIQUID
state.
(g)
Appears after the formula, the compound or molecule is in the
GASEOUS state.
(aq)
Appears after the formula, the compound or molecule is dissolved in
water.
∆
Heat is required or supplied to the reaction
Heat
Pt
Platinum is used as a catalyst for the reaction
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Law of Conservation of Mass
• Mass cannot be created or destroyed. It can only change
form.
• Mass of reactants must equal the mass of the products.
• The number of atoms of each element must be the same
on BOTH sides of the arrow.
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Balancing Chemical Reactions
• Balanced Equation – the number of atoms of
each element on the reactant side of the arrow
(left) equal the number of atoms of each element
on the product side of the arrow (right) .
• To do this you can change the COEFFICIENTS
ONLY! These are the numbers in front of the
chemical formulas (Bold).
• NEVER CHANGE THE SUBSCRIPTS!
4 Al (s) + 3 O2 (g) ---> 2 Al2O3 (s)
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Steps to Balancing Equations
There are four basic steps to balancing a chemical equation.
1. Write the correct formula for the reactants and the products. DO
NOT TRY TO BALANCE IT YET! You must write the correct
formulas first. And most importantly, once you write them correctly
DO NOT CHANGE THE FORMULAS!
2. Find the number of atoms for each element on the left side.
Compare those against the number of the atoms of the same
element on the right side.
3. Determine where to place coefficients in front of formulas so that
the left side has the same number of atoms as the right side for
EACH element in order to balance the equation.
4. Check your answer to see if:
– The numbers of atoms on both sides of the equation are now
balanced.
– The coefficients are in the lowest possible whole number
ratios. (reduced)
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Helpful Hints for balancing equations
• Take one element at a time, working left to right
except for H and O. Save H for next to last, and O
until last.
• IF everything balances except for O, and there is
no way to balance O with a whole number, double
all the coefficients and try again. (Because O is
diatomic as an element)
• (Shortcut) Polyatomic ions that appear on both
sides of the equation should be balanced as
independent units
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Balance
5. ______HgO → ______ Hg + ______ O2
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Law of Conservation of Mass
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Which of the following reactions best illustrates the Law of Conservation of Mass?
a) H2O2 → H2O + O2
b) Na + CuS → Na2S + 2 Cu
c) K + AgCl → KCl + Ag
d) NaOH + 2 HCl → NaCl + H2O
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When sodium chloride reacts with calcium oxide to form sodium oxide plus calcium
oxide, which of the following equations best illustrates the Law of Conservation of
Mass?
a) NaCl + CaO → Na2O + CaCl2
b) NaCl + CaO → 2 Na2O + CaCl2
c) 2 NaCl + CaO → Na2O + CaCl2
d) 3 NaCl + 2 CaO → Na2O + 3 CaCl2
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In the following reaction: 2NaN3 decomposes to form 2Na + 3N2. If 500 grams of
NaN3 decomposes to form 323.20 grams of N2. How much Na is produced?
a) 100 grams
b) 176.80 grams
c) 323.20 grams
d) 500 grams
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Types of Chemical Reactions
Composition, Addition, Synthesis or Direct Combination
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2  1
single product is formed
– General Equation:
A + X  AX
–
element + element  compound
–
compound + element  bigger compound
–
compound + compound  bigger compound
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Examples:
Fe + S  FeS
2 KCl + 3 O2  2 KClO3
SO3 + H2O  H2SO4
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Burning Ribbon
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List the reactants
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Write the balanced
chemical equation
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Explain why this is
considered a synthesis
reaction.
Observations
Explain the products in
terms of your
observations
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Types of Chemical Reactions
Decomposition, Analysis
–1  2
Single reactant is used up
– 1 substance splits up into 2 or more substances.
– General Equation:
AX  A + X
– compound  element + element
– big compound  compound + element
– big compound  compound + compound
– Examples
2 H2O  2 H2 + O2
NaClO2  NaCl + O2
Ca(OH)2  CaO + H2O
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Elephant Tooth Paste
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List the reactants
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Write the balanced
chemical equation
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Explain why this is
considered a
decomposition reaction.
Observations
Explain the products in
terms of your
observations
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Good Bye Gummie Bear
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List the reactants
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Write the balanced
chemical equation
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Explain why this is
considered a
decomposition reaction.
Observations
Explain the products in
terms of your
observations
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Types of Chemical Reactions
Combustion, Burning, Rapid Oxidation
Fuel + O2
 CO2 +
H 2O + E
• Fuel is a hydrocarbon
• Hydrocarbons are molecules that contain
only carbon and hydrogen
• 2C8H18 + 25O2  16CO2 + 18H2O
• Lots of energy is released
• oxygen is always a reactant
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Methane Bubbles
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List the reactants
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Write the balanced
chemical equation
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Explain why this is
considered a
combustion reaction.
Observations
Explain the products in
terms of your
observations
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Types of Chemical Reactions
Double Replacement, Ionic, Exchange of Ions,
Metathesis
–
2  2
– compound + compound  different
compound + different compound
– General Equation
– B+Y- + A+X-  B+X- + A+Y–
–
Compounds change partners
Examples
Cu+2S-2 + 2Na+1Cl-1  Na2+1S-2 + Cu+2Cl2-1
CuS + 2NaCl → Na2S + CuCl2
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Types of Chemical Reactions
Single Replacement
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2  2
element + compound  different compound + different
element
General Equation:
Bo + A+X-  B+X- + Ao
positives replace positives
the more active metal replaces the less active metal
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General Equation:
Yo + A+X-  A+Y- + Xo
negatives replace negatives
the more active nonmetal replaces the less active nonmetal
Or
Examples:
Zn + CuS  ZnS + Cu (metals)
F2 + 2NaCl  2NaF + Cl2
(nonmetals)
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Who’s Replacing Who?
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List the reactants
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Write the balanced
chemical equation
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Is this reaction a double
or single replacement
reaction?
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Explain how you know.
Observations
Explain the products in
terms of your
observations
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• Go back to balanced eqns!
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Activity Series – Table J
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Each element from the list displaces any of the
elements below it from a compound.
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The larger the space between the elements in the
table, the more vigorous the reaction.
All metals above hydrogen displace hydrogen
from HCl or H2SO4.
Metals above Mg vigorously displace hydrogen
from water.
Elements near the top of the table are never
found free in nature.
Elements near the bottom of the table are
often found free in nature.
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Complete the following equations using Table J to predict
the products. Write out the equation and be sure it’s
balanced.
• magnesium + copper sulfate yields ?
• calcium + zinc nitrate yields?
• copper + aluminum nitrate produces?
• tin + copper nitrate produces?
• iron + tin sulfate yields?
• aluminum + silver nitrate produces?
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• Write the correct formulas for each reactant and product.
• Balance the equation
• Identify the type of reaction
– Liquid water decomposes to hydrogen and oxygen gases.
– Solid sodium reacts with fluorine gas to produce solid sodium
fluoride.
– Aqueous Iron (III) iodide and chlorine gas yields aqueous iron
(III) chloride and iodine gas.
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• Write the correct formulas for each reactant and product.
• Balance the equation
• Identify the type of reaction
– Aqueous calcium bromide reacts with aqueous sodium sulfide to
produce the aqueous forms of calcium sulfide and sodium
bromide.
– Aqueous calcium chloride combines with aqueous silver nitrate
to yield solid silver chloride and aqueous calcium nitrate.
– Solid magnesium reacts in the presence of a nickel catalyst with
chlorine gas to produce solid magnesium chloride.
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Predicting When a Solid is Produced from the
Reaction between Two Aqueous Substances
• Solubility Guidelines – Table F
• Double Replacement Reactions
• Examples:
–
Ag2SO4(aq) +
–
AlCl3(aq) +
–
Na2CO3(aq) +
–
NaCl(aq)
Ca(NO3)2 (aq) 
(NH4)3PO4 (aq)
ZnCl2 (aq)
+ Pb3(PO4)2(s)
→
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How do you know when a rxn goes to
completion?
• The phrase go to completion means that the reaction
continues until all reactants are used up and the reverse
reaction can NOT happen.
• Conditions under which this occurs:
– When a gas is formed. (Zn + 2 HCl → ZnCl2 + H2 ↑) The
hydrogen gas bubbles out of the test tube so the reverse
reaction cannot occur.
– When a precipitate is formed. (AgNO3 + HCl → AgCl ↓ + HNO3)
The ions that make up silver chloride are no longer present in
the reaction so the reverse reaction cannot occur.
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Stoichiometry and Chemical
Equations
• Stoichiometry is the quantitative
relationship between reactants and
products in a chemical reaction.
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Coefficients = Moles
Stoichiometry means…
4 Al(s) + 3 O2(g) ---> 2 Al2O3(s)
This equation means
4 Al atoms + 3 O2 molecules
---produces--->
2 molecules of Al2O3
or
4 moles of Al + 3 moles of O2
---produces--->
2 moles of Al2O3
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Stoichiometry does NOT mean…
4 grams of Al + 3 grams of O2
---produces--->
2 grams of Al2O3
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Conversions Review
• How many moles are present in 73g of HCl?
(mass to mole)
• How many grams are present in 2.6 moles of
NaCl? (mole to mass)
• How many molecules are present in 5 moles of
glucose? (mole to molecule)
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Mole to Moles Conversions
• Using the coefficients from a balanced chemical
equation in your dimensional analysis.
• How many moles of H2 are produced when 7
moles of Na are allowed to react according to
the following equation?
2 Na + 2 H2O → 2 NaOH + H2
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Mass to Mass Conversions
• Use dimensional analysis to do the following:
– Convert grams to moles (use formula mass)
– Convert moles to moles (use coefficients in equation)
– Convert moles to grams (use formula mass)
• How many grams of carbon dioxide are
produced upon combustion of 44 grams of CH4?
CH4 + 2O2 → CO2 + 2 H2O
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Stoichiometry Problems
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