Stoichiometry

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Stoichiometry Solution Stoich

• Chapter 3 and 4.2/4.6

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Chemical Reactions/Equations

• Reactants are listed on the left
• Products on the right
• Atoms are neither created nor destroyed
• Shows what state of matter the compounds are in
  Solid (s), Liquid (l), Gas (g), or dissolved in
  water (aqueous solution) (aq)

3
Chemical Equations

• Hydrogen burns and reacts with oxygen in the air
  to form water. write the chemical reaction
• H2 (g) O2 (g) ? H2O (l)
• Whats wrong with the above equation?
• Due to the Law of Conservation of Mass, it
  balances out to
• 2H2 (g) O2 (g)? 2H2O (l)

4
Balancing Equations

• Determine what reaction is occurring (sometimes
  it helps to write it in word form)
• Write the unbalanced (skeleton) eqn.
• Balance the equation by inspection, by adding
  coefficients (usually works best going from left
  to right)
• Include phase information

5
Balancing Equations

• Al2(SO4)3 Ca(OH)2 ? Al(OH)3 CaSO4
• H3PO4 NaOH ? Na3PO4 HOH

6
Types of Reactions

• Knowing the basic types of chemical reactions,
  helps you to predict the products of many
  reactions
• Combination Reaction
• 2Mg (s) O2 (g) ? 2MgO (s)
• Decomposition Reaction
• CaCO3 (s) ? CaO (s) CO2 (g)

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• Combustion Reactions
• Rapid reactions that produce a flame
• Most involve O2 (from air) as a reactant
• When Hydrocarbons (CxHy) react with O2 they
  produce CO2 and H2O.
• C3H8(g) 5O2 (g) ? 3CO2 (g) 4H2O(g)

8

• When an air bag deploys, sodium azide (NaN3)
  decomposes, rapidly releasing nitrogen gas and
  sodium.
• What type of RXN took place?
• Write a balanced equation

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Balancing Combustion RXNs

• first start with those elements that occur in the
  fewest chemical formulas.
• Try a few
• CH4 O2 ? CO2 H2O
• C2H5OH O2 ? CO2 H2O

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How much sand?
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3 Ways to Measure Matter

• By COUNT 1 million grains of sand
• By MASS 1,000 grams of sand
• By VOLUME 100 liters of sand

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ATOMIC MASS

• What is the atomic mass of Hydrogen?
• 1.01 a.m.u.
• What is the atomic mass of Oxygen?
• 15.999 ? 16.0 a.m.u.

13
The Mass of a Compound


SO3
1 S atom
3 O atoms
You can calculate the mass of a molecule by
adding the atomic masses of the atoms making up
the molecule. (32.1 16 16 16 80.1 amu)
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Formula Weights

• What is the atomic mass of Water (H2O)?
• 2H 2 (1.0) 2.0 a.m.u.
• 1O 1(16.0) 16.0 a.m.u.
• 18.0 a.m.u
• What is the atomic mass of Ca(NO3)2?
• 1 Ca 1(40.1) 40.1 a.m.u
• 2 N 2 (14.0) 28.0 a.m.u.
• 6 O 6 (16.0) 96 a.m.u.
• 164.1 a.m.u.

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Percentage Composition

• The percentage by mass contributed by each
  element in the substance
• Calculating any percentage is just like
  calculating your grade in a class (the part,
  divided by the whole, multiplied by 100)
• Calculate the percentage by mass of each element
  in Ca(NO3)2

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• Calculate the percentage composition of Oxygen in
  glucose (C6H12O6)

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Problem Solving Hints

1. Analyze the Problem
2. Develop a plan for solving the problem
3. Solve the problem
4. Check the solution

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The Mole
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What is a MOLE?

• A mole is a quantity equal to Avogadros Number
  (6.02 x 1023)
• 6.02 x 1023 particles (atoms or molecules)
  depending on what you are looking at.
• A mole of anything contains the same number of
  things as a mole of anything else.
• one mole is set by defining one mole of carbon 12
  atoms to have a mass of exactly 12 grams.

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Molar Mass

• The mass in grams of one mole of a compound
• It is numerical equivalent to what its atomic
  weight was in a.m.u.s
• Molar Mass of Water (H2O)?
• 18 grams/mole
• Molar Mass of Calcium Nitrate?
• 164.1 grams/mole

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The Mole Road Map
Multiply by 22.4
divide

Grams
Moles
Liters
multiply
Divide by 22.4
Molecules
Use 6.02 x 1023
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Practice Problems

• How many moles are in 5g of Copper?
• How many grams in 3.2 moles of Oxygen?
• How many atoms in .350 moles of Sodium?
• How many moles in 7.2 x 1024 atoms of gold?
• How many glucose molecules in 5.23 grams of
  glucose, C6H12O6?
• How many atoms of Oxygen?

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Determining the Empirical Formula of a Compound

1. Determine the percentage of each element in your
   compound
2. Treat as grams, and convert grams of each to
   moles of each element
3. Find Smallest whole number ratio (divide the
   larger number by the smaller one)
4. If ratio is not all whole numbers, multiply each
   by an integer to make all whole number ratio

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Example 1

• Mercury forms a compound with chlorine that is
  73.9 mercury and 26.1 chlorine.
• Convert to 73.9 g Hg and 26.1 g Cl
• Convert to moles of each element
• Find Smallest whole number ratio (divide the
  larger number by the smaller one)

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Example 2

• Ascorbic acid (Vitamin C) contains 40.92 C,
  4.58 H, and 54.5 O by mass. What is the
  empirical formula of ascorbic acid?
• CHO 3 (11.331) 343
• C3H4O3

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Determining the molecular Formula

• Find the empirical formula mass.
• Divide the known molecular mass by the empirical
  formula mass, deriving a whole number, n.
• Multiply the empirical formula by n to derive the
  molecular formula.

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Example 1

• The empirical formula of ascorbic acid is C3H4O3
• The empirical formula mass is 88.0 amu.
• The experimentally determined molecular weight is
  176 amu.
• Therefore the molecule has twice the mass (176/88
  2.00) and must have twice as many of each atom
  for a molecular formula of C6H8O6

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Example 2

• Mesitylene, has an empirical formula of C3H4.
  The experimentally determined molecular weight of
  this substance is 121 amu. What is the molecular
  formula of mesitylene?

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Combustion Analysis to determine empirical
formulas

• Used to figure out how much Carbon and Hydrogen
  were in the original sample of the Hydrocarbon

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Example 1

• Isopropyl alcohol (rubbing alcohol) is composed
  of C, H, and O. Combustion of .255g of isopropyl
  alcohol produces 0.561 g CO2 and 0.306g H2O.
  Determine the empirical formula of isopropyl
  alcohol.
• Calculate the number of grams of C present in the
  CO2 and grams of H present in H2O using the mole
  concept and dimensional analysis.

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1. Calculate the mass of O in the final sample by
   subtracting the mass of C and H in the sample
   from the total sample mass Mass of O mass of
   sample (mass of C mass of H)
2. Then calculate the number of moles of C, H, and O
   in the sample
3. Find the lowest whole number ratio

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Example 2

• Caproic acid, which is responsible for the foul
  odor of dirty socks, is composed of C, H, and O
  atoms. Combustion of a 0.225g sample of this
  compound produces 0.512g CO2 and 0.209 g H2O.
  What is the empirical formula of caproic acid?
• It has a molar mass of 116 g/mol. What is its
  molecular formula?

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Stoichiometric CalculationsQuantitative info
from balanced EQNs

1. Balance the chemical equation
2. Convert grams of reactant or product to moles.
3. Compare moles of the known to moles of the
   desired substance (use a ratio derived from the
   coefficients in the balanced equation.)
4. Convert from moles back to grams if required.

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Chemical Calculations

• Mole to Mole
• Mole to Gram
• Gram to Gram

m
m
m
m
g
m
m
g
g
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m
m
g
g
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Mole-Mole Calculations

• How many moles of ammonia (NH3) are produced when
  0.60 moles of nitrogen reacts with hydrogen?
• Step 1 Write Chemical EQN
• N2 H2 ? NH3
• Step 2 Balance EQN
• N2 3H2 ? 2NH3

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• Step 3 Find known unknown, then calculate

Known 0.60 mol N2
Unknown ? mol NH3
Mole Ratio 1 mol N2 2 mol NH3
0.60 mol N2 x 2 mol NH3 1.2 mol NH3 1
mol N2
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Gram-Gram Calculations

• Calculate the number of grams of NH3 produced by
  the reaction of 5.40 g of hydrogen with an excess
  of nitrogen.

m
m
2
N2 3H2 2NH3
m
m
5.40 g
? grams
1
g
g
g
3
g
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Step 1 Change grams to moles.
5.40 g H2 x 1 mol H2 2.70 mol H2
2.0 g H2
Step 2 Get a mole ratio from the equation.
3 mol H2
2.7 mol H2
then
2 mol NH3
? mol NH3
Step 3 Solve for the unknown number of moles.
(2.7 x 2) / 3 1.8 moles NH3
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Step 4 Change moles of the unknown to grams.
1.8 mole NH3 x 17.0 g NH3 30.6 g NH3
1 mol NH3
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Examples

• How many grams of water are produced in the
  oxidation of 1.00 g of glucose, C6H12O6?
• Propane, C3H8, is a common fuel used for cooking
  and home heating. What mass of O2 is consumed in
  the combustion of 2.75 moles of propane?

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Examples

• K2PtCl4(aq) NH3(aq) Pt(NH3)2Cl2 (s)
  KCl(aq)
• what mass of Pt(NH3)2Cl2 can be produced from 65
  g of K2PtCl4 ?
• How much KCl will be produced?
• How much from 65 grams of NH3?

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Making Chocolate Chip Cookies

• Ingredients in Kitchen (I have a BIG kitchen)
• 40 lbs of butter
• 2 lbs of salt
• 1 gallon of vanilla extract
• 80 lbs of chocolate chips
• 200 lbs of flour
• 150 lbs of sugar
• 10 lbs baking soda
• 2 eggs

Whats going to determine how many cookies I
can make?
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Limiting Reactants

• limits or determines the amount of product that
  can be formed in a reaction.
• The reactant that isnt used up is called the
  excess reagent
• To determine, book says use ratio method and
  I.C.E. chart, Ill show you a different method
  both work.

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• To determine the limiting reagent requires that
  you do two stoichiometry problems.
• Figure out how much product each reactant makes.
• The one that makes the least is the limiting
  reagent

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Using an I.C.E. chart

• Convert both reactants into moles suppose we had
  10 moles of H2 and 7 moles of O2.

2H2 O2 ? 2H2O
Initial Quantities
Change (reaction)
Expected Quantities
10 mol
7 mol
0 mol
-5 mol
10 mol
-10 mol
0 mol
2 mol
10 mol
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Example

• Ammonia is produced by the following
  reaction N2 H2 NH3
• What mass of ammonia can be produced from a
  mixture of 100. g N2 and 500. g H2 ?
• How much unreacted material remains?

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Success of Reaction

• The amount of stuff you make is the yield.
• The theoretical yield is the amount you would
  make if everything went perfect.
• The actual yield is what you make in the lab.

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Percent Yield

• yield Actual x 100
  Theoretical
• yield what you got x 100 what
  you could have got

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Examples

• Aluminum burns in bromine producing aluminum
  bromide. In a laboratory 6.0 g of aluminum reacts
  with excess bromine. 50.3 g of aluminum bromide
  are produced. What are the three types of yield.
  (actual, theoretical, percent)

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Precipitation Reactions
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• Occur when certain pairs of oppositely charged
  ions attract each other so strongly that they
  form an insoluble ionic solid.
• Can determine if a precipitate will form by
  following certain guidlines

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Solubility Guidelines for Ionic Compounds

• Previous picture Only 1.2 x10-3 mol of PbI2
  dissolves in a liter of water.
• We will consider it insoluble if the solubility
  is less than 0.01 mol/L
• Refer to Table 4.1 (and Ksp values on pg 1045) in
  text for solubility guidelines for common ionic
  compounds in water
• Dateless Dudes
• Ammonium, Nitrate, Acetate, and Alkali Metals

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• Ksp Solubility-Product Constants
• pg. 1045 in text book
• The smaller the number, the more likely it will
  precipitate
• If there is no insoluble product, the reaction
  does not occur.
• Exchange (Metathesis) Reactions are another name
  for double replacement reactions ( and - ions
  switch partners)

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Simple Solubility Rules

1. Most nitrate slats are soluble
2. Most salts containing the alkali metal ions and
   ammonium are soluble
3. Most chloride, bromide, and iodide salts are
   soluble. (exceptions are salts containing the
   ions Ag, Pb2, and Hg2)

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1. Most sulfate (SO4-2) salts are soluble. (notable
   exceptions are BaSO4, PbSO4, HgSO4, and CaSO4)
2. Most OH- salts are only slightly soluble. The
   important soluble ones are NaOH and KOH. Ba, Sr
   and Ca are marginally soluble.
3. Most S-2, CO3-2, CrO4-2 and PO4-3 are slightly
   soluble.

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Ionic and Net Ionic Equations

• Write out the molecular equation
• Pb(NO3)2(aq) 2 KI(aq) ? PbI2 (s) 2KNO3(aq)
• Write out the complete ionic equation and
  Identify and cancel out spectator ions
• Write out the Net Ionic equation

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Example

• An aqueous solution of silver nitrate reacts with
  an aqueous solution of potassium chloride. A
  precipitate is produced.
• Determine the precipitate using Ksp.
• Write the Molecular, Ionic, and Net Ionic
  equations for the above reaction

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Solution Stoichiometry and Chemical Analysis

• Molarity moles of solute / Liters of solution.
• M1V1 M2V2
• We can take the molarity and volume of a solution
  to find out the moles of a solution. And then use
  dimensional analysis to determine the moles or
  grams of another reactant or product.

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Example 1

• How many grams of Ca(OH)2 are needed to
  neutralize 25.0 mL of 0.100 M HNO3?
• Use molarity and volume of HNO3 to convert to
  moles of HNO3
• Balance equation
• HNO3 Ca(OH)2 ? H2O Ca(NO3)2
• Convert moles of HNO3 to moles of Ca(OH)2 and
  then to grams of Ca(OH)2

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Example 2

• How many grams of NaOH are needed to neutralize
  20.0 mL of 0.150M H2SO4 solution?
• NaOH H2SO4 ? H2O Na2SO4