Download Empirical Formula

Chapter 3
Than, K. Sugar Found In Space: A Sign of Life?
Organic molecules found in gas swaddling a young
star. National Geographic News. PUBLISHED August
30, 2012
Molecules, Compounds, and
Chemistry Equations
Hydrogen, Oxygen, and Water
• Hydrogen gas (____) is explosive
• Oxygen gas (_____) is not explosive but must
be present for combustion (fire) to occur
• When these two gases come together to
react, they form water
– A molecule now with completely different
properties!!!
Properties
• The properties of compounds / molecules are
generally different from those of the elements
used to compose them
• What “holds” these atoms together??
Chemical Bonds
• Compounds are composed of atoms held
together by chemical bonds
– Chemical bonds are the result of interactions
between charged particles
• Electrons (-) and Protons (+) !!
• Two MAIN types of bonds (more in the
“bonding” chapters  ):
1. Ionic
2. Covalent
Chemical Bonds
• Two MAIN types of bonds
We have two whole chapters on chemical bonds (Chapter 9 AND chapter 10)
1. Ionic (a metal and nonmetal) – transferring of
electrons – or stealing of electrons
2. Covalent (two nonmetals) – sharing electrons to
obtain an octet and make everyone happy 
Image obtained from:
http://kitkatyj.deviantart.com/art/The-four-chemicalbonds-395098750 on October 17, 2016
Ionic Bonds
• Metals tend to lose electrons and form
_____________
• While nonmetals tend to gain electrons and form
______________
• These oppositely charged ions are then attracted
to one another by electrostatic forces
• The results is an ionic compound
– In the solid phase make up a lattice structure – a 3-D
array of alternating cations and anions
Covalent Bonds
• When two nonmetals bonds, neither actually
give up or take on electrons, instead they
share
facts-i-just-madeup.tumblr.com/post/83778517260/sharing-iscaring-the-new-care-bears-series
Covalent Bonds
• When two nonmetals bonds, neither actually
give up or take on electrons, instead they
share
• The shared electrons interact with the nuclei
of both atoms, lowering their potential energy
through electrostatic interactions with the
nuclei
– Think of it like the (-) electron is holding together
the two (+) nuclei
Chemical Formulas
• Chemical formulas can be divided into three
different types:
1. Empirical – gives the relative number of atoms of
each element in a compound (like a ratio)
2. Molecular – gives the actual number of atoms of
each element in compounds (not simplified)
3. Structural – shows lines between the elements
to attempt to illustrate how they are organized or
connected (bonded) together
Empirical Formulas
• Empirical Formulas are just “simplified” down
molecular formulas
Practice:
write out the empirical formula from each
of the following molecular formulas
A. C4H8
B. B2H6
C. CCl4
Empirical Formulas
• Empirical Formulas are just “simplified” down
molecular formulas
Practice:
write out the empirical formula from each
of the following molecular formulas
A. C4H8
A. CH2
B. B2H6
B. BH3
C. CCl4
C. CCl4
Empirical Formulas
• Empirical Formulas can also be calculated
• To do this, you must have some information
given to you
• You must know a Percent Composition
– So…. What is a percent comp?
Percent Composition
(Mass Percent Composition)
• A Percent Composition is the percentage by mass
of each element in a compound.
• Just like your grade on a test
• First, calculate the molar mass of the ENTIRE
compound (total points possible).
• Then, divide that by the mass of the part you are
trying to find the percent for (your score).
Practice with Percent Composition
(Mass Percent Composition)
• Calculate the percent composition of iron in
iron(III)oxide.
• Calculate the percent of Nitrate in magnesium
nitrate.
Empirical Formulas
• The Empirical Formula is a chemical formula
that shows the composition of a compound in
terms of the relative numbers and kinds of
atoms in the SIMPLEST RATIO.
• How do we find this simplest ratio?
– You just did this using the easier way… what if
we are not given a molecular formula but instead
a percent composition of the compound?
Empirical Formulas
• To calculate the empirical formula, you must
first know the percent composition.
• Using the percents, assume a 100g sample.
– How will this help?
– What units will this leave your percents in?
Empirical Formulas
• Now that we are in grams of each component
(from the percent compositions), we need to
convert these to moles.
• Last step, divide each number by the LOWEST
number to determine what the lowest ratio
values will be.
Lets Practice!!
• Chemical analysis of a liquid shows that it is
60.0% C, 13.4% H, and 26.6% O by mass.
Calculate the empirical formula of this
substance.
More Practice!!
• A dead alkaline battery is found to contain a compound of Mn and
O. Its analysis gives 69.6% Mn and 30.4% O.
• A compound is 38.77% Cl and 61.23% O.
• A liquid compound is 18.0% C, 2.26% H, and 79.7% Cl.
Molecular Formulas
• Molecular Formulas are multiples of Empirical
Formulas.
• A molecular formula is a chemical formula
that shows the number and kinds of atoms in
a molecule, but not the arrangement of the
atoms.
Molecular Formulas
• The molar mass of a compound is equal to the
molar mass of the empirical formula times a
whole number, n.
• Fill in the following table:
Empirical
Formula
Molar mass of
compound
Molar mass of
the empirical
formula
MM/EFM
Molecular
Formula
Practice with Molecular Formulas
• The empirical formula for a compound is P2O5.
Its experimental molar mass is 284 g/mol.
Determine the molecular formula of the
compound.
Practice with Molecular Formulas
• The empirical formula for a compound is P2O5.
Its experimental molar mass is 284 g/mol.
Determine the molecular formula of the
compound.
Empirical
Formula
Molar mass of
compound
Molar mass of
the empirical
formula
MM/EFM
Molecular
Formula
More Practice
• A compound has an experimental molar mass
of 78 g/mol. Its empirical formula is CH.
What is its molecular formula?
• A brown gas has the empirical formula NO2.
Its experimental molar mass is 46 g/mol.
What is its molecular formula?
Structural Formulas
• Yep that is it….
• Mostly these are
“organic” but can
represent more than
just these
“hydrocarbons”
Atomic-Level View
Atomic-Level View
• Atomic Elements are those that exist in nature
with single atoms as their basic units
– Examples: Gold, Helium, Iron, etc.
• Molecular Elements do not normally exist in
nature with single atoms, but instead as
molecules or groups of atoms
– Examples: O3, O2, S8, all your other ___________
Compounds
• Molecular Compounds (molecules) are
usually composed of two or more covalently
bonded nonmetals
– Examples: H2O, CO2
• Ionic Compounds are composed of cations
(metal) and anions (nonmetals)
– Example: Na+ & Cl-  NaCl
Polyatomic Ions
• Look at the back of your Periodic Tables.
• You know this list as “Polyatomic Ions”
– Poly = many
– Atomic = atoms
– Ion = with a charge
• Polyatomic Ions – an ion composed of two or
more atoms with an overall net charge. Forms
ionic (electrostatic) attractions / bonds
Practice
• Classify each of the following substances as an
atomic element, molecular element, molecular
compounds, or ionic compound
A.
B.
C.
D.
E.
Xenon
NiCl2
Bromine
NO2
NaNO2
More Practice
• Classify each of the following substances as an
atomic element, molecular element, molecular
compounds, or ionic compound
A.
B.
C.
D.
E.
Fluorine
N2O
Silver
K2O
F2O 3
REVIEW: Ionic Compounds
• Nomenclature Rules for “types A and B”
• Ionic Compounds CAN have polyatomics!!
– Name the same way
• Type A (make the charges cancel and change
the ending)
• Type B (again, make the charge cancel and
change the ending… use Roman numeral!)
REVIEW: Naming Molecular
Compounds
• This is Type C
• Use prefixes
• DO NOT use or even look at the charges
• Will NOT use polyatomics!!!!!
Acids (not review!)
• Acids are molecular compounds that release
hydrogen ions (H+) when dissolved in water
• The ‘acidic’ hydrogen is usually written out
front (ex. HCl, HBr, HI, H2SO4, HNO3, HClO4)
HCl  (when dissolved)  H+(aq) + Cl-(aq)
Aqueous(aq) means ‘dissolved in water’
Acids
• Acids are characterized by their sour taste and
their ability to dissolve many metals
• They are ___________ on the pH scale
• pH actually is a measure of the H+ ion
concentration
Naming Acids
• Binary Acids are composed of hydrogen and a
nonmetal
Example: HCl(aq) = HYDRO + CHLOR + IC ACID
Try to Name: HF(aq), HI(aq), H2S(aq)
Naming Acids (oxyacids)
• Oxyacids contain hydrogen and an oxyanion
(an anion containing a nonmetal and oxygen –
like nitrate or nitrite, sulfate or sulfite, etc.
Naming Acids (oxyacids)
Try naming:
HNO3(aq)
H2SO3(aq)
HNO2(aq)
Write out:
Acetic Acid
Benzoic Acid
Perchlorous Acid
Conversion Factors from Chemical
Formulas
• Think about CCl2F2
• For every one mole of CCl2F2 , there are two
moles of Cl
2 𝑚𝑜𝑙 𝐶𝑙
1 𝑚𝑜𝑙𝑒 CCl2F2
Let’s Practice  What would be the mass of Cl in
25.0 grams of CCl2F2?
Practice
• What mass of hydrogen (in grams) is
contained in 1.00 gallons of water? (the
density of water is 1.00 g/mL)
3.785x103 grams H20
4.23x102 grams H
More Practice
• Determine the mass of oxygen in a 7.2 gram
sample of aluminum sulfate
More Practice
• Butane (C4H10) is used as a liquid fuel in lighters. How
many grams of carbon are presented within a lighter
containing 7.25 mL of butane? The density of liquid
butane is 0.601 g/mL)
Chemical Reactions
Part 1:
Describing Chemical
Reactions
Chemical Reactions (rxns)
• Chemical Reaction (rxn) – a process in which
one or more substances are converted into
one or more different substance
– “bonds are broken, reformed, and gives you
something new”
http://www.harpercollege.edu/tmps/chm/100/dgodambe/thedisk/chemrxn/signs3.
htm
Some Vocab
• Chemical Equations are used to “represent”
or describe a chemical reaction (sometimes
called “word equations”
• Reactants –
• Products –
Reactants MUST Come Together
• You cannot hit a ball without touching it.
• Reactions use the same concept
• The reactants must come together and interact in
order to react with one another
• What happens with you:
• Increase the concentration of the reactants?
• Crush up the reactants?
• Increase the temperature / pressure?
Word Equations Tell Us A Lot
Symbol
(s), (ℓ), (g)
Meaning
Substance in the solid, liquid, and gaseous state.
(aq)
Substance in aqueous solution (dissolved in water).

“produces” or “yields,” indicating result or reaction
Reversible Rxn in which products can reform into reactants; final
result is a mixture of products and reactants
Δ heat
 or 
Pd

Reactants are heated; temperature is not specified
Name or chemical formula of a catalyst, added to speed a reaction.
Chemical Reactions
Part 2:
Balancing Chemical
Reactions
Balancing Equations
• The ________________ states that we cannot
create or destroy matter / energy…
• So, what we start with… we must end with!!
– EXACTLY!!
• Coefficients are used to ‘multiply’ substance
quantities to make the number of moles the
same on both sides of the ‘arrow’
Coefficients
• H2O represents water (2 H, 1 O)
• 2 H20 represents two moles of water (4 H, 2 O)
• H202 represents hydrogen peroxide (2 H, 2 O)
• A small change of the subscripts within the
molecule will change the compound completely
AND alter the entire reaction.
More on Coefficients
• Determine the number of each atom in the following
examples:
• _1_Fe2O3 + _3_H2 
• _3_ Ca(NO3)2 +
• _1_ C6H12O6 + _6_ O2  _6_ H20 + _6_ CO2
Rules to consider when balancing
1. Identify reactants and products
• If no equation is provided, identify the
reactants and products and write an
unbalanced equation for the reaction. (You
may find it helpful to write a word equation
first.)
• If not all chemicals are described in the
problem, try to predict the missing
chemicals based on the type of reaction.
Rules to consider when balancing
2. Count atoms
• Count the number of atoms of each element in
the reactants and in the products, and record the
results in a table.
• Identify elements that appear in only one reactant
and in only one product, and balance the atoms of
those elements first. Delay the balancing of
atoms (often hydrogen and oxygen) that appear in
more that one reactant or product.
• If a polyatomic ion appears on both sides of the
equation, treat it as a single unit in you counts.
Rules to consider when balancing
3. Insert coefficients
• Balance atoms one element at a time by inserting
coefficients.
• Count the atoms of each element frequently as you
try different coefficients. Watch for elements whose
atoms become unbalanced as a result of your work.
• Try the odd-even technique (double the odd
number) if you see an even number of a particular
atom on one side of an equation and an odd
number of that atom on the other side.
Rules to consider when balancing
4. Verify your results
• Double-check to be sure that the
numbers of atoms of each element
are equal on both sides of the
equation. Make sure your own work
did not mess any “previously
balanced” atoms.
Practice:
• Balance the following example:
• Iron (III) oxide and hydrogen react to yield iron
and water.
Practice:
• Balance the following examples:
• ___P4 + ___O2  ___P2O5
• ___C3H8 + ___O2  ___CO2 + ___H2O
• Silicon reacts with carbon dioxide to form silicon
carbide, SiC, and silicon dioxide
NEVER change subscripts to balance a
reaction!!!
Chemical Reactions
Part 3:
Classifying Chemical
Reactions
Reaction Types
• There is an unlimited number of reactions that
may occur in the real world.
• To make learning them easier, we classify
them according to what is taking place.
– Five main types
• Makes predicting products MUCH easier.
1. Combustion Reactions
• Often used to generate energy.
• Combustion Reaction – the oxidation of an
organic compound, in which heat is released.
• One reactant is a hydrocarbon and the other
must be ______________.
• Products are always carbon dioxide and water.
Combustion Reactions
Examples:
___C3H8 + ___O2  ___H2O + ___CO2
___C2H5OH + ___
 ___
Write down the combustion of CH4
+ ___
2. Synthesis Reaction (aka:
combination reactions)
• The word synthesis means “to put together”
• Synthesis Reaction – a reaction in which two
or more substances combine to form a new
compound.
• The reactants in many of these reactions are
two elements or two small compounds.
Synthesis Reactions
C + O2  CO2
2C + O2  2CO
CaO(s) + H2O(l)  Ca(OH)2(s)
Predict the products of the following synthesis rxns:
K + Cl2 
Mg + O2 
3. Decomposition Reaction
• Decomposition reactions are the opposite of
synthesis reactions – they have only one
reactant.
• In a decomposition reaction, a single
compound breaks down, often with the input
of energy, into two or more elements or
simpler compounds.
Decomposition Reactions
Water breaks down into its simpler components
when introduced to electricity:
Electricity
2H2O(l) ---------> 2H2(g) + O2(g)
Compounds made up of three or more elements
usually do not decompose into those elements.
Heat
CaCO3(s) -------> CaO(s) + CO2(g)
Practice So Far:
• Predict the product(s) and write a balanced
equation for each of the following reactions:
– The reaction of butane, C4H10, with oxygen.
– The reaction of water and calcium oxide
– The reaction of lithium with oxygen
– The decomposition of carbonic acid.
4. Single Displacement (replacement)
Reactions
• Single displacement (replacement) reaction– A single element reacts with a compound and
displaces another element of the same charge
from the compound.
• Examples:
2Al(s) + 3CuCl2(aq)  2AlCl3(aq) + 3Cu(s)
• Predict the following products:
CuSO4(aq) + Zn(s) 
Activity Series
• Used to make predictions about displacement
reactions.
• In the activity series, elements are arranged in
order of activity with the most active on top.
• In general, an element can displace those
listed below it but not those above it.
What would
you predict will
happen when
silver is put
into a copper
(II) nitrate
soln.?
5. Double Displacement (replacement)
Reactions
• Double displacement (replacement) reaction– Two compounds react with one another and swap
cations / anions.
• Examples:
HCl(aq) + NaOH(aq)  HOH(l) + NaCl(s)
• Predict the following products:
Cu(NO3)2(aq) + Al2(SO4)3(s) 
Practice Classifying Reactions
• Balance each of the equations below, and
indicate the type of reaction for each equation
___Cl2(g) + ___NaBr(aq)  ___NaCl(aq) + ___Br2(l)
___CaO(s) + ___H2O  ___Ca(OH)2(aq)
___Ca(ClO3)2(s)  ___CaCl2(s) + ___O2(g)
Practice Classifying Reactions
• Balance each of the equations below, and
indicate the type of reaction for each equation
___C8H18(l) + ___O2(g)  ___CO2(g) + ___H2O(l)
__Zn(s) + __CuBr2(aq)  __ZnBr2(aq) + __Cu(s)
__AgNO3(aq) + __K2SO4(aq)  __Ag2SO4 + __KNO3(aq)
Practice With Activity Series
• Predict whether a reaction
would occur when the
materials indicated are brought
together. For each reaction
that would occur, complete
and balance the equation.
___Ag(s) + ___H2O(l)
___Mg(s) + ___Cu(NO3)2(aq)
___Al(s) + ___MgNO3(aq)
Predicting Products
• Predict the products, write a balanced
equation, and identify the type of reaction for
each of the following reactions.
___H2SO4(aq) + ___KOH(aq)
___ C3H7OH + ___ O2 
___ Zn + ___ CuSO4 
Predicting Products
• Predict the products, write a balanced
equation, and identify the type of reaction for
each of the following reactions.
___ BaCl2 + ___ Na2SO4 
___ Zn + ___ F2 
___ C5H10 + ___ O2 
Test Coming up!!
• Continue Practicing!!!
Read on your own 
• Read through Section 3.11: Organic
Compounds (page 117-121) – QUIZ!!
• You need to know this section as well as
naming this alkanes, alkenes, and alkynes
– Know the skeletal notation, the vocabulary, and
the functional groups from page 120. Be able to
‘circle’ possible functional groups seen in a large
drawn out organic molecule
Total Ionic Reactions
• You can write out a chemical reaction
illustrating EXACTLY what is taking place “in
the pot”
• Total Ionic Equations show us ALL the ions
interacting in solution
– Assign all phases – break apart “aqueous”
– Some of these ions play an important role in the
reaction, others just “spectate”
Practice
• Write out the total ionic equations for the
following reaction:
Sodium hydroxide reacts with sulfuric acid
Spectator Ions
• Spectators ions do not influence the reaction
– They remain unchanged from the reactant side to
the product side.
• You may CROSS OUT and forget about
spectator ions – they do not do anything
– Think about your example reaction – cross out the
spectator ions
Net Ionic Equations
• After you cross out your spectator ions, you are
left with a Net Ionic Equation, typically resulting
in a “precipitate”
• Net ionic equations are equations that show only
the soluble, strong electrolytes reacting (these
are represented as ions) and omit the spectator
ions, which go through the reaction unchanged.
• What are you left with from the NaOH + H2SO4 ?
So….
• How do you know what is solid, liquid, or
gas??
– Use your best judgement !!!
• How about aqueous???
– There is a way!!
Solubility Rules
Practice
Silver nitrate reacts with potassium chloride
Magnesium nitrate reacts with sodium
carbonate
strontium bromide + potassium sulfate
manganese(II)chloride + ammonium carbonate
chromium(III)nitrate + iron(II)sulfate
End of Unit 3 