Atomic Mass

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Atomic Mass

• Atoms are so small, it is difficult to discuss
  how much they weigh in grams.
• Use atomic mass units.
• an atomic mass unit (amu) is one twelth the mass
  of a carbon-12 atom.
• This gives us a basis for comparison.
• The decimal numbers on the table are atomic
  masses in amu.

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They are not whole numbers

• Because they are based on averages of atoms and
  of isotopes.
• can figure out the average atomic mass from the
  mass of the isotopes and their relative
  abundance.
• add up the percent as decimals times the masses
  of the isotopes.

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Examples

• There are two isotopes of carbon 12C with a mass
  of 12.00000 amu(98.892), and 13C with a mass of
  13.00335 amu (1.108).
• There are two isotopes of nitrogen , one with an
  atomic mass of 14.0031 amu and one with a mass of
  15.0001 amu. What is the percent abundance of
  each?

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The Mole

• The mole is a number.
• A very large number, but still, just a number.
• 6.022 x 1023 of anything is a mole
• A large dozen.
• The number of atoms in exactly 12 grams of
  carbon-12.

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The Mole

• Makes the numbers on the table the mass of the
  average atom.

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Representative particles

• The smallest pieces of a substance.
• For a molecular compound it is a molecule.
• For an ionic compound it is a formula unit.
• For an element it is an atom.

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More Stoichiometry
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Molar mass

• Mass of 1 mole of a substance.
• Often called molecular weight.
• To determine the molar mass of an element, look
  on the table.
• To determine the molar mass of a compound, add up
  the molar masses of the elements that make it up.

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Find the molar mass of

• CH4
• Mg3P2
• Ca(NO3)3
• Al2(Cr2O7)3
• CaSO4 2H2O

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Examples

• How much would 2.34 moles of carbon weigh?
• How many moles of magnesium in 24.31 g of Mg?
• How many atoms of lithium in 1.00 g of Li?
• How much would 3.45 x 1022 atoms of U weigh?

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Percent Composition

• Percent of each element a compound is composed
  of.
• Find the mass of each element, divide by the
  total mass, multiply by a 100.
• Easiest if you use a mole of the compound.
• Find the percent composition of CH4
• Al2(Cr2O7)3
• CaSO4 2H2O

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Working backwards

• From percent composition, you can determine the
  empirical formula.
• Empirical Formula the lowest ratio of atoms in a
  molecule.
• Based on mole ratios.
• A sample is 59.53 C, 5.38H, 10.68N, and
  24.40O what is its empirical formula.

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CO2 is absorbed
Sample is burned completely to form CO2 and H2O
H2O is absorbed
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• A 0.2000 gram sample of a compound (vitamin C)
  composed of only C, H, and O is burned completely
  with excess O2 . 0.2998 g of CO2 and 0.0819 g of
  H2O are produced. What is the empirical formula?

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More Stoichiometry
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Empirical To Molecular Formulas

• Empirical is lowest ratio.
• Molecular is actual molecule.
• Need Molar mass.
• Ratio of empirical to molar mass will tell you
  the molecular formula.
• Must be a whole number because...

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Example

• A compound is made of only sulfur and oxygen. It
  is 69.6 S by mass. Its molar mass is 184 g/mol.
  What is its formula?

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Chemical Equations

• Are sentences.
• Describe what happens in a chemical reaction.
• Reactants Products
• Equations should be balanced.
• Have the same number of each kind of atoms on
  both sides because ...

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Meaning

• A balanced equation can be used to describe a
  reaction in molecules and atoms.
• Not grams.
• Chemical reactions happen molecules at a time
• or dozens of molecules at a time
• or moles of molecules.

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Stoichiometry

• Given an amount of either starting material or
  product, determining the other quantities.
• use conversion factors from
• molar mass (g - mole)
• balanced equation (mole - mole)
• keep track.

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Examples

• How many moles is 4.56 g of CO2 ?
• How many grams is 9.87 moles of H2O?
• How many molecules in 6.8 g of CH4?
• 49 molecules of C6H12O6 weighs how much?

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Examples

• One way of producing O2(g) involves the
  decomposition of potassium chlorate into
  potassium chloride and oxygen gas. A 25.5 g
  sample of Potassium chlorate is decomposed. How
  many moles of O2(g) are produced?
• How many grams of potassium chloride?
• How many grams of oxygen?

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Examples

• A piece of aluminum foil 5.11 in x 3.23 in x
  0.0381 in is dissolved in excess HCl(aq). How
  many grams of H2(g) are produced?
• How many grams of each reactant are needed to
  produce 15 grams of iron form the following
  reaction? Fe2O3(s) Al(s) Fe(s) Al2O3(s)

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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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Gases and the Mole
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Gases

• Many of the chemicals we deal with are gases.
• They are difficult to weigh.
• Need to know how many moles of gas we have.
• Two things effect the volume of a gas
• Temperature and pressure
• Compare at the same temp. and pressure.

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Standard Temperature and Pressure

• 0ºC and 1 atm pressure
• abbreviated STP
• At STP 1 mole of gas occupies 22.4 L
• Called the molar volume
• Avagadros Hypothesis - at the same temperature
  and pressure equal volumes of gas have the same
  number of particles.

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Examples

• What is the volume of 4.59 mole of CO2 gas at
  STP?
• How many moles is 5.67 L of O2 at STP?
• What is the volume of 8.8g of CH4 gas at STP?

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Density of a gas

• D m /V
• for a gas the units will be g / L
• We can determine the density of any gas at STP if
  we know its formula.
• To find the density we need the mass and the
  volume.
• If you assume you have 1 mole than the mass is
  the molar mass (PT)
• At STP the volume is 22.4 L.

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Examples

• Find the density of CO2 at STP.
• Find the density of CH4 at STP.

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The other way

• Given the density, we can find the molar mass of
  the gas.
• Again, pretend you have a mole at STP, so V
  22.4 L.
• m D x V
• m is the mass of 1 mole, since you have 22.4 L of
  the stuff.
• What is the molar mass of a gas with a density of
  1.964 g/L?
• 2.86 g/L?

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Stoichiometry

• Greek for measuring elements
• The calculations of quantities in chemical
  reactions based on a balanced equation.
• We can interpret balanced chemical equations
  several ways.

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Look at it differently

• 2H2 O2 2H2O
• 2 dozen molecules of hydrogen and 1 dozen
  molecules of oxygen form 2 dozen molecules of
  water.
• 2 x (6.02 x 1023) molecules of hydrogen and 1 x
  (6.02 x 1023) molecules of oxygen form 2 x (6.02
  x 1023) molecules of water.
• 2 moles of hydrogen and 1 mole of oxygen form 2
  moles of water.

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Mole to mole conversions

• 2 Al2O3 4Al 3O2
• every time we use 2 moles of Al2O3 we make 3
  moles of O2

2 moles Al2O3
3 mole O2
or
3 mole O2
2 moles Al2O3
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Mole to Mole conversions

• How many moles of O2 are produced when 3.34 moles
  of Al2O3 decompose?
• 2 Al2O3 4Al 3O2

3.34 moles Al2O3
3 mole O2

5.01 moles O2
2 moles Al2O3
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Your Turn

• 2C2H2 5 O2 4CO2 2 H2O
• If 3.84 moles of C2H2 are burned, how many moles
  of O2 are needed?
• How many moles of C2H2 are needed to produce
  8.95 mole of H2O?
• If 2.47 moles of C2H2 are burned, how many moles
  of CO2 are formed?

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Periodic Table
Periodic Table
Balanced Equation

• Decide where to start based on the units you are
  given and stop based on what unit you are asked
  for

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For example...

• If 10.1 g of Fe are added to a solution of Copper
  (II) Sulfate, how much solid copper would form?
• Fe CuSO4 Fe2(SO4)3 Cu
• 2Fe 3CuSO4 Fe2(SO4)3 3Cu

1 mol Fe
63.55 g Cu
10.1 g Fe
3 mol Cu
55.85 g Fe
2 mol Fe
1 mol Cu
17.3 g Cu

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More Examples

• To make silicon for computer chips they use this
  reaction
• SiCl4 2Mg 2MgCl2 Si
• How many moles of Mg are needed to make 9.3 g of
  Si?
• 3.74 mol of Mg would make how many moles of Si?
• How many grams of MgCl2 are produced along with
  9.3 g of silicon?

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For Example

• The U. S. Space Shuttle boosters use this
  reaction
• 3 Al(s) 3 NH4ClO4 Al2O3 AlCl3 3 NO
  6H2O
• How much Al must be used to react with 652 g of
  NH4ClO4 ?
• How much water is produced?
• How much AlCl3?

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Gases and Reactions
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We can also change

• Liters of a gas to moles
• At STP
• 0ºC and 1 atmosphere pressure
• At STP 22.4 L of a gas 1 mole
• If 6.45 moles of water are decomposed, how many
  liters of oxygen will be produced at STP?

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For Example

• If 6.45 grams of water are decomposed, how many
  liters of oxygen will be produced at STP?
• H2O H2 O2
• 2H2O 2H2 O2

1 mol H2O
1 mol O2
22.4 L O2
6.45 g H2O
1 mol O2
18.02 g H2O
2 mol H2O
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Your Turn

• How many liters of CO2 at STP will be produced
  from the complete combustion of 23.2 g C4H10 ?
• What volume of oxygen will be required?

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Yield

• How much you get from an chemical reaction

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Limiting Reagent

• If you are given one dozen loaves of bread, a
  gallon of mustard and three pieces of salami, how
  many salami sandwiches can you make?
• The limiting reagent is the reactant you run out
  of first.
• The excess reagent is the one you have left over.
• The limiting reagent determines how much product
  you can make

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Limiting Reagent

• Reactant that determines the amount of product
  formed.
• The one you run out of first.
• Makes the least product.
• Book shows you a ratio method.
• It works.
• So does mine

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Limiting reagent

• 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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How do you find out?

• Do two stoichiometry problems.
• The one that makes the least product is the
  limiting reagent.
• For example
• Copper reacts with sulfur to form copper ( I )
  sulfide. If 10.6 g of copper reacts with 3.83 g S
  how much product will be formed?

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• If 10.6 g of copper reacts with 3.83 g S. How
  many grams of product will be formed?
• 2Cu S Cu2S

Cu is Limiting Reagent
1 mol Cu
1 mol Cu2S
159.16 g Cu2S
10.6 g Cu
63.55g Cu
2 mol Cu
1 mol Cu2S
13.3 g Cu2S
13.3 g Cu2S
1 mol S
1 mol Cu2S
159.16 g Cu2S
3.83 g S
32.06g S
1 mol S
1 mol Cu2S
19.0 g Cu2S
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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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How much excess reagent?

• Use the limiting reagent to find out how much
  excess reagent you used
• Subtract that from the amount of excess you
  started with

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Excess Reagent

• The reactant you dont run out of.
• 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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Your turn

• Mg(s) 2 HCl(g) MgCl2(s) H2(g)
• If 10.1 mol of magnesium and 4.87 mol of HCl gas
  are reacted, how many moles of gas will be
  produced?
• How much excess reagent remains?

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Your Turn II

• If 10.3 g of aluminum are reacted with 51.7 g of
  CuSO4 how much copper will be produced?
• How much excess reagent will remain?

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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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Yield

• The amount of product made in a chemical
  reaction.
• There are three types
• Actual yield- what you get in the lab when the
  chemicals are mixed
• Theoretical yield- what the balanced equation
  tells you you should make.
• Percent yield Actual x 100
  Theoretical

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Example

• 6.78 g of copper is produced when 3.92 g of Al
  are reacted with excess copper (II) sulfate.
• 2Al 3 CuSO4 Al2(SO4)3 3Cu
• What is the actual yield?
• What is the theoretical yield?
• What is the percent yield?
• If you had started with 9.73 g of Al, how much
  copper would you expect?

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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.

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Examples

• Years of experience have proven that the percent
  yield for the following reaction is
  74.3 Hg Br2 HgBr2 If 10.0 g
  of Hg and 9.00 g of Br2 are reacted, how much
  HgBr2 will be produced?
• If the reaction did go to completion, how much
  excess reagent would be left?

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Examples

• Commercial brass is an alloy of Cu and Zn. It
  reacts with HCl by the following reaction Zn(s)
  2HCl(aq) ZnCl2 (aq) H2(g) Cu does not react.
  When 0.5065 g of brass is reacted with excess
  HCl, 0.0985 g of ZnCl2 are eventually isolated.
  What is the composition of the brass?