AP Chemistry
Unit 3 Topics:
•
Covalent, ionic, & acids: naming and writing formulas*
•
Balancing chemical equations
•
Decomposition, composition reactions, and combustion : predicting products
•
% composition
•
Empirical vs. Molecular formulas: “combustion train” calculation
•
Stoichiometry conversions, limiting reagents, % Yield
* You will not have formal notes for writing the formulas for compounds.
Instead, we will just do practice problems until you get the hang of it. If you
need “formal notes”, you can go to my website and download the 1st year
Accelerated Chemistry Notes for “Naming and Writing Formulas”.
Naming Compounds
• Covalent Molecules– Contains only 2 nonmetals.
− They are formed by sharing e− between atoms.
• General Format:
Prefix (except mono)-name 1st element prefix-name 2nd element ending in –ide
Naming Compounds
• Ionic Compounds – Starts with metallic cation (or NH4+).
−Formed when atoms transfer e− from a metal to a nonmetal.
General Format
Cation Name Anion Name
• On a later quiz, you will have to memorize the cation and anion symbols &
charges on the Ion Sheet. For this unit, I will let you use the Ion Sheet.
Naming Compounds
• Acids– Starts with “H”
Balancing Equations
• You can only change coefficients.
• You can only use whole #’s.
Example:
C3H8 + __O
3
4 2O
5 2  __CO
2 + __H
Decomposition Reactions
• A reaction that breaks apart ______
one ______________
compound into simpler
substances, (usually two elements or an element and a smaller
compound.)
+
General Form: _____
AX  ___
A + ___
X
H2 + _____
O2
Examples:
H2O  _____
O2
KCl + _____
KClO3  _____
Remember that “HONClBrIF” elements are diatomic when alone!!
Don’t try to balance it until you have correctly written the products!!
Categories of Decomposition (and Composition ) Reactions
a) carbonates  metallic oxide + CO2
CaO + _____
CO2
CaCO3  _____
b) chlorates  metallic chloride + O2
NaCl + _____
O2
NaClO3  _____
c) hydroxides  metallic oxide + H2O
MgO + _____
H2O
Mg(OH)2  _____
d) oxy acids  nonmetal oxide + H2O
SO3 + _____
H2O
H2SO4  _____
e) binary compounds  2 elements
Na + _____
Cl2
NaCl  _____
•
•
Notice that in carbonates, hydroxides, and oxy acids, the oxygen
combines with both the metal AND the nonmetal!!
Every time you try to write the formula for a new ionic compound,
charges of the ions and ___________
cross
you must look up the ___________
them if they are different!!
Composition Reactions
This reaction category is sometimes called “Combination” or “Synthesis”
•
•
two __________________,
substances
A reaction of _____
typically a metal and a
one ______________.
compound
nonmetal to form ______
It is the opposite of decomposition. The same categories of
reactions apply. It’s just in reverse!
+
General Form:
Examples:
___
A + ___
X  _____
AX
Al
AlCl3
+ Cl2  _______
2 elements
 binary compound
Pb(OH)2
PbO + H2O  ______
metallic + water 
oxide
hydroxide
•
•
•
Combustion Reactions
A reaction between O2 and a compound containing Carbon,
Hydrogen, (and sometimes Oxygen).
CO2 + ________
H2O
The products are always the same… ________
This reaction is too easy!! Don’t miss it!
General Form:
Examples:
C2H2
CO2 + ____
H2O
CxHy + O2  ____
CO2 + _______
H2O
+ O2  _______
C7H6O +
CO2 + _______
H2O
O2  _______
Percent Composition (by mass)
% Element 
Atoms of Element AW 
FW of Compound
 100
AW stands for the atomic weight of the atom from the periodic table.
FW stands for the formula weight of the compound, i.e.--the “molar mass”.
AW and FW are nicknamed the “W.O.T.C.” (the weight on the chart) or the
“Wizzle on the Chizzle”.
Here’s another formula for % compostion:
Empirical Formulas
Helpful Rhyme: % to mass, mass to mole, divide by small, times ’til whole.
• The molecular formula is a whole # multiple of the empirical formula.
• When given the molecular weight of the compound, you can compare the “empirical mass”
to the “molecular mass” to see what this whole # multiple will be. (We will practice this
calculation later.)
“Combustion Train”
Dehydrite [Mg(ClO4)2•3H2O] and ascarite (NaOH on asbestos)
A weighed quantity of the substance to be analyzed is placed in a combustion train
and heated in a stream of dry O2. All the H in the compound is converted to H2O(g)
which is trapped selectively in a previously weighed absorption tube. All the C is
converted to CO2 (g) and this is absorbed selectively in a second tube. The increase
of mass of each tube tells, respectively, how much H2O and CO2 were produced by
combustion of the sample.
EXAMPLE 1:
A 6.49-mg sample of ascorbic acid (vitamin C) was burned in a
combustion train and 9.74 mg CO2 and 2.64 mg H2O were
formed. Determine the empirical formula of ascorbic acid.
Steps to Follow:
(1)
Calculate the moles of Carbon and Hydrogen in the compound.
(2)
Calculate the moles of Oxygen (if any) in the compound. (Sometimes other elements are
present, such as Nitrogen, but you will be given a way to find these masses as well.)
(3)
The mole ratios are the subscripts of the empirical formula. Use the “Empirical Formula
Rhyme” to complete the rest of the problem.
Ans: C3H4O3
EXAMPLE 2:
The combustion of 62.63 grams of a compound which contains
only C,H and O yields 134.43 grams of CO2 and 13.75 grams of
H2O. What is the empirical formula of the compound? Ans: C2HO
–
If the compound’s molecular weight is 156 g/mol, determine the molecular
formula. Ans: C8H4O4
EXAMPLE 3:
The combustion of 40.10 g of a compound which contains only C,
H, Cl and O yields 58.57 g of CO2 and 14.98 g of H2O. Another
sample of the compound with a mass of 75.00 g is found to contain
22.06 g of Cl. What is the empirical formula of the compound?
–
(Hint: Use the mass data to calculate the % composition of each element in
the compound.) Ans: C4H5ClO2
Stoichiometry Conversion Factors
1 mole = 22.4 L (at STP) = 6.02 x 1023 particles = Molar Mass in grams
• These conversions will take up to 3 steps and no more!
• Always convert to moles of the “given” first!
Limiting Reagent (or Reactant)
• The reactant that runs out first “limits” the amount of product that can
be formed.
• Stoichiometry conversions can be done to determine which substance is
the limiting reagent and how much of the excess is left over.
% Yield
•The amount of product predicted from stoichiometry taking
into account limiting reagents is called the theoretical yield.
− This is the ideal amount of product that should be formed if all
goes as planned in a “perfect world.”
•The percent yield relates the actual yield (amount of material
recovered in the laboratory) to the theoretical yield:
Actual yield
% Yield 
 100
Theoretica l yield