Stoichiometry

1
Stoichiometry

• Chapter 2-6 2-12
• Chapter 3-1 3-8

2
Key concepts

• Know what a mole is, and how to use it.
• Understand the term molar mass and its
  relationship to formula weight.
• Interconvert between number of particles, moles,
  and mass.
• Understand the term percent composition and know
  how to calculate the percent composition of an
  element in a formula.
• Use percentage composition to determine the
  empirical formula of a substance.
• Understand how to formulate a chemical equation.
• Understand how to balance chemical equations by
  inspection.
• Know how to use moles and a chemical formula to
  determine quantitative information about chemical
  reactions.
• Understand the terms theoretical yield and
  limiting reactant.
• Introduction to solutions in chemical reactions

3
The mole

• THE MOLE IS A COUNTING UNIT!!!!
• You can have a mole of anything.
• 1 dozen 12 things
• 1 gross 144 things (a dozen dozen)
• 1 mol 6.0221367 ? 1023 things

4
Molar Mass

• 1 mole number of atoms in 0.012 kg of
  carbon-12.
• 1 amu 1/12 mass of a carbon-12 atom.
• As a result, the atomic weight of an element (in
  amu) is __________________ to the mass of one
  mole of that element (in g/mol). This is defined
  as the molar mass of the element.
• Review Is molar mass an intensive or extensive
  property?

5
Formula weight/molecular weight

• We may use the atomic weight (in amu) of elements
  in a molecular formula (or formula unit) to
  determine the molecular weight (or formula
  weight) of the molecule (unit).
• Examples Fe2(CO3)3, C11H22O11

6
Converting from mass to moles

• Just as the atomic weight is numerically equal to
  the molar mass of an atom, the molecular weight
  (formula weight) is numerically equal to
  __________________.
• Now we have a powerful conversion between the
  mass of a substance and the moles (or the number
  of particles we have). Why is it so important to
  have this kind of conversion? (or is it
  important at all?)

7
Converting mass ?? moles

• We will need to determine the molar mass of the
  substance(s) in question.
• Examples
• How many moles of CO2 are in 0.56 g?
• Determine molar mass
• Convert from mass to moles
• What is the mass of 0.45 mol of lithium
  carbonate?

8
composition of a compound

• (we discussed this before, but now we have a
  better way to work it.)
• Use the chemical formula and molar masses to
  determine the relative mass of each element in
  the compound.
• What is the C in glucose?

9
Empirical formula

• The relative number of each kind of atom in a
  substance
• To determine empirical formulas, you must count
  atoms. Therefore, you must use moles, a counting
  unit, not grams.

10
Determining empirical formulas

• assume 100 g of sample (when is given). If
  actual masses are known they may also be used.
• convert from grams of element to moles of element
• Ratio all other elements against the element with
  the smallest number of moles (not mass).
• Example a compound contains 52.9 Al and 47.1
  O.

11
More examples

• A certain compound is 73.9 Hg and 26.1 Cl by
  mass. What is the empirical formula?
• (BLB) A sample of a substance is analyzed and
  found to contain 5.28 g Sn and 3.37 g F. What is
  the empirical formula?
• (BLB) A sample contains 11.66 Fe and 5.01 g
  oxygen. What is the empirical formula?

12

• For ionic compounds, the empirical formula is
  generally the formula unit.
• For molecular compounds, molecular formula is not
  always the empirical formula, but will always be
  some whole number multiple of the subscripts of
  the atom.
• The relative ratio between the atoms must be
  preserved (if you multiply one subscript by a
  number, you must multiply all subscripts by the
  same number)
• Law of constant composition.
• What would you need to know in order to find the
  molecular formula using the empirical formula?

13
Combustion analysis of compounds

• Done with hydrocarbons and other organic
  compounds.
• We use deductive reasoning to determine the
  composition from the data.
• Example A 1.000 g sample of an alcohol is burned
  and produces 1.913 g CO2 and 1.174 g H2O. What
  is the empirical formula?

14
Even more fun

• (BLB) Nicotine contains C, H, and N. A 5.250 mg
  sample of nicotine was combusted, and 14.242 mg
  CO2 and 4.083 mg H2O were produced. What is the
  empirical formula for nicotine?
• The molar mass of nicotine is 1605 g/mol. What
  is the molecular formula?

15
Formula of a hydrate

• A hydrate is a complex where an amount of water
  is contained inside an ionic salt crystal in a
  specific ratio.
• We may use the empirical formula calculation
  process in determining the formula of a hydrate
  as well.

16
Chemical equationa visual representation of a
chemical reaction

• There are reactants and products in a chemical
  reaction.
• Reactants
• Products

17
Conservation of matter

• Matter is not created or destroyed in a chemical
  reaction.
• The number of each type of atom in a chemical
  reaction must be the same for reactants as it is
  for products.
• Sandwich example write a reaction for a
  sandwich made of two slices of bread and one
  slice or cheese.

18
Balancing equations by inspection

• Start with the element found in the fewest
  chemical formulas on either side of the reaction.
• Fill out remaining elements, changing
  coefficients as needed, until the number of each
  type of atom is same on both sides of the
  equation.
• When balancing equations, we change the
  coefficients. we do not change the subscripts.
  Changing the subscripts means we have changed the
  molecule involved in the reaction (law of
  constant composition).
• More balancing examples

19
Using the chemical equation to determine
quantitative information

• A balanced chemical equation gives us the ratio
  of reactants and products to each other, whether
  in molecules or moles, but NOT the mass ratio.
• example combustion of methane
• if we know how much we have of one reactant (or
  product), we may determine how much of all other
  reactants (and products) were used (or produced)
  in the chemical reaction.

20
In general, comparing substance A and substance B

• Given mass of A, find moles of A using molar mass
  of A
• use ratio of A and B to find moles of B.
• Use molar mass of B to find mass of B.
• The ratio (called the stoichiometric ratio or
  reaction ratio) is the key to solving these types
  of problems.
• Lets do some examples.

21
Limiting reactants.

• limiting reactantthe reactant that limits how
  much product you can get. The limiting reactant
  is entirely used up in the reaction.
• Back to the sandwich example if I have 6 slices
  of bread and 4 slices of cheese, which is the
  limiting reactant?
• Note it doesnt matter if I start with the
  bread or the cheese, Ill still be able to
  determine the answer if I use the ratio properly.

22

• Limiting reactants are always determined by
  counting how many molecules are needed or, in
  other words, you must use moles, not mass.
• Lets try some more examples.

23

• (WGD) How many grams of NH3 can be prepared from
  77.3 g N2 and 14.2 g H2?
• N2 H2 ? NH3
• (WGD) 12.6 g of AgNO3 and 8.4 g of BaCl2 are
  dissolved in water. A solid, AgCl, forms. How
  much of this solid, in grams, can form?
• AgNO3 BaCl2 ? AgCl Ba(NO3)2

24
Yield

• The yield is the amount produced in a reaction.
• Theoretical yield
• Actual yield
• Percent yield

25
Yields

• (WGD) Ethylene glycol (C2H6O2) is formed
  according to the following rxn
• C2H4Cl2 Na2CO3 H2O ? C2H6O2 2 NaCl CO2
• (this is balanced)
• When 27.4 g C2H4Cl2 is used in the rxn, 10.3 g
  ethylene glycol is formed.
• What is the theoretical yield?
• What is the actual yield?
• What is the percent yield?

26
Sequential reactions

• The same concepts we have been learning in the
  analysis of single reactions can be applied to
  multiple reactions. Lets look at an example

27
Sequential reactions

• (WDP) What mass KClO3 is needed to provide enough
  O2 to react completely with 66.3 g methane?
• KClO3 ? KCl O2
• CH4 O2 ? CO2 H2O
• Break problem into steps plot your course.

28
Solutions

• A solution is a ____________________.
• A Solution contains a solute and a solvent.
• Solute
• Solvent
• We will discuss ways of expressing the
  concentration of a solute in solution. Well
  talk about the solvation process (process of
  dissolving in solution) later in the course.

29
by mass
This is mass of the solution, not mass of the
solvent.
30

• In very dilute solutions, solute is an awkward
  unit to use. Often, parts-per-million (ppm) or
  parts-per-billion (ppb) are used instead.

31
Molarity

• Chemists will often use molarity because it is
  more directly related to the moles of solute (and
  becomes more useful in analyzing chemical
  reactions).

Not volume of solvent.
32
Solution dilution..

• Suppose we change the volume of a solution, but
  we leave the amount of solute the same.

33
Solutions in chemical reactions

• Just as we dealt with the mass of reactants and
  products in a chemical reaction by converting to
  moles, we may also use the molarity of a solution
  in analyzing chemical reactivity.
• The key is to always compare moles to moles.