Ch 2 - Notes
2-0 Intro – Stoichiometry – means?
A) Two types of analysis –
1. qualitative…what does it contain, and
2. quantitative…how much of everything does it contain
B) Stoichiometry – composition stoichiometry (this chapter) and reaction stoichiometry (ch 3)
2-1 Atoms and Molecules
A) Aristotle v Democritus
1. Early scientists – philosophers/thinkers – NOT experimenters
2. ~450B.C. Demjocritus (along with teacher Leucippus) – “atomos” = all matter
composed of tiny indivisible particles
B) 1803 - Dalton’s atomic theory
1. Each element composed of extremely small particles called atoms which are
indivisible and indestructible.
2. All atoms of a given element are identical (specifically in their masses).
3. Atoms of any given element are different than atoms of any other element (specifically
in their masses).
4. A given compound always has the same relative numbers (whole number ratios) and
kinds of atoms.
Note: not all this is still considered true! (explain)
C) Atoms are composed of: Subatomic particles
Electron
Proton
Neutron
e (e-)
p (p+)
n (no)
-1
+1
0
Relative mass
(amu)
0.000549
1.007
1.009
(~0)
(~1)
(~1)
Actual mass (g)
9.11x10-28
1.673x10-24
1.675x10-24
Symbol
Relative charge
APChem, 2/15/2016
2-2 Chemical Formulas (see also nomenclature notes) – see figure 2-6 (note: what do you think of
the caption about the space-filling model?)
A) Molecules can be represented by:
1. Chemical formulas: ex. H2O:
 shows: _________________________________________
2. Structural formulas: Ex:
 shows: ___________________________________________
3. Ball and stick model: Ex:
 shows: ___________________________________________
4. Space filling model: Ex:
 shows: ___________________________________________
B) See Nomenclature notes – Molecular compounds
2-3 Ions and Ionic Compounds (see nomenclature notes – Ionic compounds)
2-4 Names and Formulas of Some Ionic Compounds (see nomenclature notes – Ionic Compounds)
2-5 Atomic Weights
A) Def: relative mass
B) 1 amu = 1/12 mass of an atom of carbon-12
C) aka: atomic mass, molar mass,
D) the mass of one mole of a substance – on the periodic table !!!!
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2-6 The Mole (cool figures in this section!!!)
A) Def.: the number of particles contained in exactly 12 g of carbon-12 atoms
B) We will use
6.022x1023 particles/mol
C) Some of the conversions available using the mole
2-7 Formula Weights, Molecular Weights, and Moles
A) Formula weight vs molecular weight (note: round to 2 decimals unless told otherwise)
1. Formula implies ionic: ex. __________________________________
2. Molecular implies covalent: _________________________________
B) Example: what is the molar mass of copper(II) sulfate pentahydrate?
C) Mole Map (ref: http://chemistry.mrs.smith.tripod.com/id12.html )
D) Ex.
1. Convert 3.2 x 10-4 moles of CaCl2 to g
2. Convert 0.020 g of Ca(NO3)2 to total atoms
3. What is the mass of 1 chlorine atom?
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4. How many atoms of hydrogen in 13.2 mL water (assume a density of 0.9982 g/mL)?
2-8 Percent Composition and Formulas of Compounds
A) be able to calculate the % composition from the formula
B) Ex.
1. ? % composition by mass of CO2?
2. ? % water by mass of copper (II) sulfate pentahydrate
2-9 Derivation of Formulas from Elemental Composition (chem team – an excellent reference!)
A) Empirical formulas: simplest whole number ratios of atoms in a formula (“Simplest
formula”) For example: the formula for benzene is C6H6, while CH is the empirical
formula
B) To determine:
A Simple Rhyme for a Simple Formula
by Joel S. Thompson
Journal of Chemical Education
Vol. 65, No. 8; August 1988, p. 704
Percent to mass
Mass to mole
Divide by small
Multiply 'til whole
C) Example:
1. Given % composition: A compound is synthesized and found to contain 9.93%
carbon, 58.64% chlorine and 31.43% fluorine by mass. Determine its empirical
formula.
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2. Given mass: A 15.000 g sample, containing only carbon, hydrogen and oxygen, was
determined to contain 10.320 g carbon and 0.742 g hydrogen. What is its simplest
formula?
3. From combustion data: A 0.2500g sample of a hydrocarbon containing only carbon,
hydrogen, and oxygen was combusted and produced 0.3664 g carbon dioxide and
0.1500 g water. What is the empirical formula of this hydrocarbon?
p. 5
2-10 Determination of Molecular Formulas
A) These are the formulas written as the molecule actually occurs. To determine them you
need to know the molar mass (or be able to calculate it)
B) Example: Many compounds have the empirical formula of CH2O. Here are the molecular
weights of three. Determine the molecular formulas for each.
 30.0 g/mol
 60.0 g/mol
 180.0 g/mol
2-11 Some Other Interpretations of Chemical Formulas
A) Example problem: What mass of oxygen is contained in 8.425 g of H2SO4?
2-12 Purity of Samples
A) Generally on the labels on reagent bottles, a sample purity is given. For example, 99%
HCl. This works like any % problem.
B) Example. A reagent bottle reads 99.44% NaCl. How many grams of the sample are
required to actually have 14.22g NaCl?
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