Title: CHEMISTRY I - Lecture
Author: WISD
Last saved by: Jamie Tharp
Company: Weatherford ISD

Chemistry I – Lecture Notes
“Formula Math & The Mole”

In  chemistry, we commonly measure _____________, _____________, _____________ .

Each of these can be related to a single quantity called the  ____________.

The “MOLE” is the SI unit for the “__________ ___ ________________”.

WHY use it?

The quantities above are related to the “MOLE” in the following way:
VOLUME:  __________________ = _________________________________

		(STP = standard temperature and pressure)
    	Standard temperature = ____________
    	Standard pressure = __________________ or ____________________

PARTICLES:  _________________ = __________________

       	for:	element = “atom”
				diatomic element or molecular compound= “molecule”
				ionic compound = “formula unit”
	
MASS:  ________________ = _________________________
	
		for:	element = “gram atomic mass”
				ionic compound = “gram formula mass”
				molecular compound = “gram molecular mass”
						diatomic elements = “gram molecular mass”
	
		You calculate all masses above the same way, so we will refer to them as  the calculation 
of ______________ _______________ (mass of 1 mole of any substance).	

		TO CALCULATE MOLAR MASS:
		  1)  Write the formula for the substance.
		  2)  COUNT the number of each element in the formula.
		  3)  MULTIPLY this number by the mass number of that element.
		  4)  ADD all products to determine molar mass.
		
		REMEMBER:  Round all mass numbers to the whole number first.

Example:  What is the molar mass of carbon tetrachloride?

Example:  What is the molar mass of oxygen?

Example:  What is the molar mass of aluminum carbonate?

All “mole related” quantities from above are equivalent, so they are interchangeable.
 
     ______________ = _______________ = _______________ = ________________________

	Because of this, you can convert from:
		moles  ↔  mass
		moles  ↔  particles
		moles  ↔  volume
		particles  ↔  volume
		mass  ↔  volume
		particles  ↔  mass

	TO CONVERT BETWEEN UNITS OF THE SAME SUBSTANCE:
	  
1)  READ the entire problem.
	  	2)  IDENTIFY “given” and the “unknown”.
	  	3)  FORMULATE a “Conversion Factor”
🡪  states the equivalence between known and unknown.
	  	4)  PLUG information into “box”.
	  	5)  SOLVE, ROUND CORRECTLY, INCLUDE PROPER UNITS!!!
	
EXAMPLE:  What is the mass of 45.8 L of carbon dioxide gas, at STP?

	EXAMPLE:	How many moles are equal to 3.01 x 10^22 atoms of magnesium?

	
 EXAMPLE:	How many molecules are in 38.69 grams of chlorine gas?

	 EXAMPLE:	What is the volume, in L, occupied by 0.600 moles of sulfur dioxide gas, at STP?

Chemistry I – Lecture Notes
“Formula Math & The Mole”

Percent Composition:  Gives the percent, by mass, of the elements in a compound.

	Formula:		  grams element_   _X _100 _= _%
						grams compound

*To Calculate:
	1)  Write chemical formula of compound.
	2)  Count the number of atoms of each element in compound.
	3)  Multiply by the mass number (from periodic table) for that element.
	4)  Add answers up to find grams of compound.
	5)  Use formula above to calculate % for each element.

Example:  What is the percent composition of water?

Example:  What is the percent composition of calcium nitrate?

Example :  What is the percent composition of a compound formed when 27.07 grams of 
calcium reacts with 47.93 grams of chlorine.

**Empirical and Molecular Formulas**

•	Empirical Formula:	Gives the lowest whole number ratio of elements in a 
compound.  (Subscripts ________________  be reduced)

•	Molecular Formula:	A ___________ ____________  _________________ of the substance’s Empirical Formula.

***From the “Mole perspective”, E.F. gives the lowest whole number ratio of “moles” of 
  each element in a compound.

				Ba_3N_2

				Ca_3(PO_4)_2

				C_6H_12O_6
				
				H_2O_2

**Empirical Formula:
	
	**To calculate:
		1)	Change % sign to grams (assume a 100 gram sample).
			**if grams are already given, go to step 2.
2)	Convert from grams to moles.
		3)	Create “mole ratio” by dividing all answers by the smallest answer.

**If all answers are whole numbers, you are finished.  
Write formula for compound.
			
**If all answers ARE NOT whole numbers, see step 4.
		
4)	Multiply all answers from step 3 by “a number” that will give the lowest whole 
number ratio of elements.  Write formula for compound.

EXAMPLE:
	
What is the E.F. of a compound composed of 32.00% Carbon, 42.66% Oxygen, 
18.67% Nitrogen, and 6.67% Hydrogen?

EXAMPLE:	

What is the E.F. of a compound composed of 25.9% Nitrogen and 74.1% Oxygen?

EXAMPLE:

		What is the E.F. of a compound that is composed of 36.6 grams Carbon and 
9.2 grams Hydrogen?

**Molecular Formula:

•	You need to know:
			1)	E.F. mass (you calculate)
			2)	M.F. mass (always given)

				M.F. mass       
•	-----------------------	= Whole number multiple (# of EF’s in the MF)	
E.F. mass 

	**To calculate:
1)	Calculate E.F. mass
		2)	Divide M.F. mass by E.F. mass.
		3)	Multiply E.F. by whole number answer from step 2.  
		4)	Write M.F.

EXAMPLE:
		What is the M.F. of a compound with a mass of 92 grams and an E.F. of NO_2?	

EXAMPLE:	
		What is the M.F. of a compound with a mass of 150 grams and an E.F. of CH_2O?	

EXAMPLE:
What is the M.F. of a compound with a mass of 132 grams if that compound is composed of 54.6% Carbon, 13.6% Hydrogen, and 31.8% Nitrogen?    (hint: calculate E.F. first)