I'm asking for curriculum planning. Summarize each of the Topic Titles, Learning Objectives,  Enduring Understandings, and Essential Knowledge statements defined into a list of simplified learning objectives. 
Example: 
By the end of this section, you will be able to:
Compute the percent composition of a compound
Determine the empirical formula of a compound
Determine the molecular formula of a compound
Also For each objective write a paragraph with info needed to understand the objective:
B. Percentage by Mass and Empirical Formula
1. Calculation of percentage by mass composition. To
determine the percentage by mass composition of an
individual element within a compound, express the mass
of each element as a percentage of the total mass of the
compound.
2. For example(bad example but you get the point):
i. A compound with the formula C;H,40; has a total
mass of
(2)(12.01) + (4X(1.008) + (2)(16.00) = 60.05
ii. Percentages of each element expressed as a function of the total mass are shown as follows:
Carbon 2XA2.00]1 _ 40 009 60.05
Hydrogen [(#)(1.008)] _ 6.714% 60.05
Oxygen (2)(16.00)] = _ 53.29% 60.05
3. All pure compounds have a fixed percentage by mass of
each element, so the purity of any given sample can be
determined by comparing the percentages by mass. A
sample that does not match the percentage by mass of a pure sample is not pure.
4. The empirical formula of a compound is the simplest
wholenumber (integer) ratio of the atoms of each element in that compound.
5. The empirical formula can be calculated from
percentage by mass data.
i. Take the percentage of each element and assume a
sample of 100 g. (This assumption converts percentages to masses.)
ii. Convert the masses (per 100 g) of each element to
the moles (see II. Moles below) of each element by
dividing each element’s mass (per 100 g) by the
correlsipondmg average atomic mass taken from the
periodic table.
iii. Find the ratio of the moles calculated in ii above by
dividing each of the moles by the smallest number
of moles.
The results from iii above will be in a convenient
ratio and gives the empirical formula.
~ It may be that the ratio includes a decimal (fraction) such as
0.500, 0.333, 0.250, and so on. Since empirical formulae are
7 /, the simplest whole-number ratios, you must multiply all of the
numbers in the ratio by 2, 3, or 4 as appropriate in order to
remove the decimal (fraction).
C. Molecular Formula
1. Unlike an empirical formula (which shows the simplest
wholenumber ratio), the molecular formula of a
compound shows the exact whole-number ratio of the
different elements in a compound.
Like the empirical formula, the numbers of each element
are recorded using a subscript to the right of the
element’s symbol. When only 1 atom is present, the
subscript 1 is assumed (understood) and not written.
The following describes the relationship between the
molecular formula and the empirical formula.
The molecular formula will be some simple,
integer-multiple of the empirical formula. For
example, a compound with an empirical formula of
CH,0 will have a molecular formula of either
CH,0, or C,H40,, or C3HgO3 and so on, where the
empirical formula is multiplied by either, 1, 2, or 3, and so on.
To establish the correct multiplier, and therefore
find the molecular formula, it is necessary to know
the molecular mass of the compound.
In the preceding example of an empirical formula
of CH,0, given a molecular mass of 60, divide the
molecular mass of the compound by the mass of
the empirical formula. In this case, the mass of the
empirical formula CH,0 is
12.01 + (2)(1.008) + 16.00 = 30.03
molar mass _ 600 _
mass of empirical formula 30.0
iv. The answer in iii is the multiplier, so (2)(CH,0) =
C,H40,, which is the molecular formula.
v. When two entirely different compounds have the
same percentage by mass, they will have identical
empirical formulas but because they are different
compounds, they will have different molecular
formulas. 

TOPIC 1.1
Moles and
Molar Mass
Required Course Content
ENDURING UNDERSTANDING
SPQ-1
The mole allows different units to be compared.
LEARNING OBJECTIVE ESSENTIAL KNOWLEDGE
SPQ-1.A
Calculate quantities of a
substance or its relative
number of particles using
dimensional analysis and the
mole concept.
SPQ-1.A.1
One cannot count particles directly while
performing laboratory work. Thus, there must
be a connection between the masses of
substances reacting and the actual number of
particles undergoing chemical changes.
SPQ-1.A.2
Avogadro’s number (N = 6.022 × 1023 mol−1
A )
provides the connection between the number
of moles in a pure sample of a substance and
the number of constituent particles (or formula
units) of that substance.
SPQ-1.A.3
Expressing the mass of an individual atom
or molecule in atomic mass units (amu) is
useful because the average mass in amu of
one particle (atom or molecule) or formula
unit of a substance will always be numerically
equal to the molar mass of that substance
in grams. Thus, there is a quantitative
connection between the mass of a substance
and the number of particles that the
substance contains.
EQN: n = m/M