1. Wednesday 1.4.17 Do Now Agenda Welcome back! What element am I?
The beginning…of ALL THE MATH!
Homework
(P)BJ procedure
Pages 1-2 of HW packet
Do Now
Look over there 

2. Friday
Agenda
Homework
Do Now
Crate (P)BJ procedures

3. Math
All the math

4. The Molar Mass
Molar Mass – the mass of one mole of any substance, reported in grams (gram atomic mass)

5. The Molar Mass of a Compound or Molecule
Molar mass – sum of the molar masses of the elements in the compound
Example: CH4
The molar mass of CH4=16.05 g/mole
Element
# mol element/mol compound
Molar Mass
Mass contribution C 1
12.01 g/mol
12.01 g H 4
1.01 g/mol
4.04 g

6. For example, the molar mass for the element, Al, is 27. 01 g
For example, the molar mass for the element, Al, is g. This means that 1 mole of Al atoms has a mass of 27.01g.
In other words, 6.02 x 1023 Al atoms has a mass of g.

7. Percent Composition Relative amounts of each element in a compound
Can be used as a conversion factor

8. Practice What is the percent by mass of nitrogen in NH3?
MM of NH3: g/mol
Mass contribution of N: g/mol
(14.01g/mol)÷(17.04g/mol) * 100 = 82.22%

9. The mole We use it to count, just like we use dozen or pair.
If we have:
1 dozen roses= 12 roses
If we have one mole of anything:
We have 6.02x1023 of that thing
Avogadro’s number

10. Empirical Formulas
A chemical formula showing the lowest whole # ratio of elements in a compound
Can be calculated from percent composition
Example: Find the empirical formula for a compound that is composed of 62.1% C, 13.8% H, and 24.1% N.

11. Empirical Formula Contd
To determine formula:
Convert the masses of each element into moles.
Divide each mole value by the smallest mole value to get mole ratio
Write empirical formula using symbols, and use the whole-number ratio as subscript.
Non-whole numbers – multiply by 2 or 3 to make them whole

12. If percents are given, assume that the % = the number of grams.
C 62.1g
H 13.8g
N 24.1 g

13. Molecular Formulas
will be identical to or some multiple of the empirical formula.
Show the actual number of each kind of atom present
Example: C2H6=Molecular Formula
CH3=Empirical Formula
Example: CH4=Molecular Formulas
CH4=Empirical Formula

14. Obtaining Molecular Formula from Empirical Formula + Molar Masses
Steps *Find empirical formula *Calculate molar mass for empirical formula *Divide MM molecular/MM empirical *Multiply this factor by each of subscripts in empirical formula to give molecular formula

15. Dimensional Analysis Three steps: List the knowns and the unknown.
Solve for the unknown.
Select the appropriate unit equality and make it into a conversion factor, with the unknown’s unit in the numerator and the known’s unit in the denominator.
Multiply the known quantity by the conversion factor.
Make sure to eliminate all but the unknown’s unit.
Evaluate - Does the result make sense?
Chapter 10

16. Converting Moles to Particles and vice versa.
We use moles to count particles like: atoms, molecules, formula units, or ions.
One mole of moles (the animal) would have the equivalent mass to 60x the Earth’s oceans.

17. Mole to particles How many atoms are there in 1 mole of sodium?
6.02x1023 atoms
How many formula units are there in 1 mole of KOH?
6.02x1023 formula units
How many molecules are there in one mole of water?
6.02x1023 molecules

18. Particles to mole
Example: How many moles of water molecules are in 1.21 x 1024 water molecules?

19. Mole to Particles
Example: How many Cu atoms are in 2.54 moles of Cu atoms?

20. Mass to mole Calculations
If you are given a certain mass of a substance, you can find out how many moles you have by using the molar mass.
Example: How many moles of carbon atoms are found in g of carbon?
24.02g x (1mole) = moles C
(12.01g)

21. The Molar Volume of a Gas
One mole of any gas at STP occupies a volume of 22.4L.
S.T.P.-Standard Temperature and Pressure. The values are 0 ◦Celsius and 1 atmosphere of pressure.
Example: What is the volume of 2.55 moles of hydrogen

22. The mole is the heart of the matter
molar mass
1.00 mol
6.02x1023particles
1.00 mol
MOLE
Particles
Mass
1.00 mol
molar mass
1.00 mol
6.02x1023particles
1 mol
22.4 L
22.4 L
1 mol
Volume
(gas at STP)

23. Stoichiometry The calculation of quantities in chemical reactions
A balanced chemical equation becomes a recipe that we can use

24. What do equations tell us?
2Na + Cl2  2NaCl
In the above equation, the coefficients tell you:
The relative number of particles
The relative number of moles
If and only if all substances are gaseous, the relative number of

25. Mole to Mole Conversions
How many moles of chlorine gas are needed to react with excess sodium to produce 5.00 moles of sodium chloride?

26. Molarity A unit of concentration Units are mole solute/liter solution
Conversion factor! 

27. Practice
If you dissolve 5 moles of NaCl in enough water to produce 1 L of solution, what is the concentration?

28. Gram to Gram Conversions
How many grams of water can be produced when grams of oxygen react with excess hydrogen?

29. Volume to Volume Conversions
2H2 (g) + O2(g)  2H2O (g)
Example: How many liters of S.T.P. are needed to completely react with 6.0 liters of hydrogen to produce water?
Notice, what is different about this reaction from the one we’ve previously been working with?

30. Percent Yield
Measurement of efficiency of a reaction

31. Practice
A factory worker in a paint company follows a batch slip to prepare a 1500 gallon batch of paint. However, when the batch goes to the filling department only 1489 gallons are filled. What is the % yield?

32. Limiting reactant/reagent
As in the cookie recipe, if you run out of eggs (and can’t get more), but have the rest of the recipe, you can’t make any more cookies.
In chemistry we call the reactant that runs out first, the limiting reactant/reagent.
All other reactants that we have in excess we call excess reagents.

33. Practice
2H2 + O2  2H2O
If 40.00g of hydrogen react with g of oxygen, what mass of water will be produced?