1. Warm Up
What is a mole?
What is molar mass?
What is Avogadro’s number?

2. The Mole and Chemical Compostion
Chapter 7
The Mole and Chemical Compostion

3. How can chemical composition be determined?
Unit Essential Question:
How can chemical composition be determined?

4. How is the mole used in conversions?
Lesson Essential Question:
How is the mole used in conversions?

5. Section 1: Avogadro’s Number and Molar Conversions
1 mole = x 1023 particles
SI unit for amount of substance.
It’s a counting unit (like a dozen).
Remember that the unit of particles can be: ions, molecules (mcs.), atoms, formula units (f.u.), etc.
Recall that formula units = simplest ratio of ions in an ionic compound.
Section 1 is mostly a review from Chapter 3. You can go back and complete problems from that chapter if you feel that you need to review more.
ionic compounds
covalent compounds

6. Recall your mole map!

7. Converting moles  particles
Same as Chapter 3, but it will involve molecules, formula units, or ions instead of just atoms.
Steps:
1) Need 1mol = x1023 molecules, etc.
2) Use dimensional analysis- turn this into a fraction!
*Be sure to place the correct units on the top and bottom so they cancel!
We mentioned in Chapter 3 that Avogadro’s number could be labeled as just about anything. For this chapter, we need to look at labels for compounds, so molecules, ions, and formula units will be used.

8. Sample Problems 1 & 2: Moles & Particles
Find the number of molecules in 2.5 mol of sulfur dioxide.
A sample contains 3.01 x 1023 molecules of sulfur dioxide. Determine the amount in moles.
1.5 x 1024 molecules SO2
0.500mol SO2

9. Molar Mass Amount of mass (in grams) in 1 mole of a substance.
Use molar masses from the periodic table.
Round to 2 decimal places!
Use units of g/mol.
Example:
C: 12.01g/mol means that 1 mol C = g
Cl: 35.45g/mol means that 1 mol Cl = 35.45g
Use to convert between moles and mass.
Remember to carefully round your numbers from the Periodic Table, it will make your problem incorrect if you have the wrong value.

10. Sample Problems 3 & 4: Moles & Mass
What is the mass of 5.3mol Be?
If you have 27.0g of manganese, how many moles do you have?
48g Be
0.491mol Mn

11. Molar Masses of Compounds
Add together the molar masses of all elements or ions present.
Ex: CH4
C: 12.01g/mol H: 1.01g/mol
12.01g/mol + 4(1.01g/mol) = 16.05g/mol
This means that 1 mole of CH4 has a mass of 16.05g.
You will need to calculate the molar mass of a compound whenever you are converting between mass and moles!

12. Additional Molar Mass Examples:
Element
Ag = g/mol
Diatomic Element/molecule
Br2 = x 2 = g/mol
Molecule (Covalent compound)
H2O = (1.01 x 2) = g/mol
Formula unit (Ionic compound)
Ca(NO3)2 = (2 x 14.01) + (6 x 16.00) = g/mol
Remember you will still have 2 decimal places in ALL of your answers if you follow the sf rules. You must keep/add zeros to make this happen.

13. Sample Problem 5: Mass to Moles with a Compound
Find the number of moles present in 47.5 g of glycerol, C3H8O3.
Hint: you will need to calculate the molar mass of glycerol!
Glycerol’s molar mass: 92.11g/mol
0.516mol C3H8O3

14. Sample Problem 6: Number of Particles to Mass
Remember- you can’t go directly between mass (g) and the number of particles! You must convert to moles first!
Find the mass in grams of 2.44 x 1024 atoms of carbon.
48.7 g C

15. More Practice
How many moles of iron (III) sulfate, Fe2(SO4)3, are there in a 178g sample?
0.445mol

16. How are molar masses on the periodic table determined?
Lesson Essential Question:
How are molar masses on the periodic table determined?

17. Mole Ratios in Chemical Formulas
Ratios can be formed between amounts of elements or ions within a compound.
Look at the subscripts.
Example #1: CaCl2
For every 1mol of CaCl2 there is 1mol of Ca+2 ions and 2mol of Cl- ions.
Example #2: Na2CO3
For every 1mol of Na2CO3, there are 2mol of Na+ ions and 1mol of CO3-2 ions.
Example #3: N2O3
For every 1mol of N2O3 there are 2mol of N atoms and 3mol of O atoms.

18. Practice
If you have one mole of strontium cyanide, Sr(CN)2, how many moles of strontium ions are there? How many moles of cyanide ions are there?
Given the compound P2O5 what is the mole ratio of P atoms to O atoms?

19. Section 2: Relative Atomic Mass and Chemical Formulas
Periodic table masses are averages of all isotopes present.
Recall that we said a weighted average is used- takes into account the amount of each isotope.
Average atomic mass:
(% x atomic mass)+(% x atomic mass)+…
100
Note: % is the percent abundance (how often the element is found as that isotope in nature).
We mentioned in Chapter 3 that the masses on the periodic table are not whole numbers because they took into account all the isotopes that exist. We did not talk about that it actually involves the % of each isotope present – how abundant it is on Earth. The formula would take each % and multiply it times the mass of that isotope and then add all of them together – be careful to use parentheses on your calculator! You must also divided by the total number of isotopes – which should be 100 if you are using percentages.

20. Sample Problem
The mass of a Cu-63 atom is amu, and that of a Cu-65 atom is amu. If the abundance of Cu-63 is 69.17% and the abundance of Cu-65 is 30.83%, what is the average atomic mass of copper?

21. Lesson Essential Questions:
What information can be determined from formulas? How can formulas be determined?

22. Calculating Percent Composition
Tells you the percent each element makes up of the whole compound.
Step 1: Determine the molar mass of the entire compound.
Step 2: Divide each element’s total molar mass by the molar mass of the compound.
Step 3: Multiply by 100 to get percent.
Step 4: Check your answer by adding up the percentages to makes sure they equal 100%.
Percentage problems are like figuring out your grade – divide part by the total. You must remember to use the subscripts to determine the parts and total and to multiply by 100 to get percent. You should get 100% total when you add up the parts – with sig figs and rounding, it might be or so.

23. Percent Composition Cont.
Calculating the percent composition of a compound can be helpful in determining the formula/identity.
Example:
Iron and oxygen form two compounds:
Fe2O3 and FeO
Fe2O3 = 69.9% Fe and 30.1% O
FeO = 77.7% Fe and 22.3% O
Scientists used percentages before they had formulas. They knew certain compounds were a certain amount of each element, but a system for displaying it as a symbol/formula had not been invented yet.

24. Sample Problem #I Sample Problem #2
Calculate the percent composition of copper (I) sulfide.
Calculate the percent composition of isopropyl alcohol, (CH3)2CHOH.
Sample Problem #2

25. Determining Empirical Formulas
The empirical formula shows the simplest ratio of elements/ions in the compound.
Ionic compounds are represented with empirical formulas.
Given percent composition data, you can determine the empirical formula of a compound.
Step 1: Assume 100 g of the sample- put ‘g’ in for ‘%’. Ex: 18.2% O  18.2g
Step 2: Convert grams to moles.
Step 3: Divide each mole value by the smallest mole value. This will tell you the number of each element that appears in the formula.
Remember you must have the lowest ratio possible in whole numbers. If you get fractions, you must multiply ALL of your numbers by the denominator in order to get the whole number ratio. If you don’t get a simple fraction like those mentioned, you probably made a mistake. Check your math.

26. Determining Empirical Formulas Cont.
Step 4: If you get a decimal, multiply ALL numbers by a whole number to turn the decimal into a whole number.
The numbers you will need to multiply by should be relatively small (2, 3, etc.)

27. Sample Problem #1 Sample Problem #2
Chemical analysis of a liquid shows that it is 60.0% C, 13.4% H, and 26.6% O by mass. Calculate the empirical formula of this substance.
A compound is found to contain 38.77% Cl and 61.23% O. What is the empirical formula?
Sample Problem #2

28. Molecular Formulas
Show the actual numbers of elements in the compound- not necessarily the simplest formula.
Often seen for covalent compounds.
They will be a whole number multiple of the empirical formula (can’t be a decimal).
In other words:
n(empirical formula) = molecular formula
where n is a whole number.
Ex: 6(CH2O)  C6H12O6
Molecular and empirical formulas can be the same!
Although for ionic compounds we always look at the simplest ratio, in covalent compounds we need to look at the molecular formula. Ionic compounds exist as crystals with repeating patterns, so we need to know the simplest part of that crystal to write a formula. For molecules (covalent compounds), we need to know the exact number of atoms since many covalent compounds simplify to the same empirical formula. This can be confusing since covalent compounds are very different in their properties. Take a look at Table 3 – formaldehyde is a preservative which causes cancer. Acetic acid is the main ingredient in vinegar. Glucose is a form of sugar. All have the same empirical formula, but very different properties.
These problems CAN NOT come out to fractions. If you don’t get an whole number, check your math.

29. Molecular Formulas Cont.

30. Molecular Formulas Cont.
The molecular formula can be determined from the empirical formula and experimental molar mass of a compound.
Step 1: Determine the molar mass of the given empirical formula.
Step 2: Solve for n by dividing the experimental molar mass by the molar mass of the empirical formula.
*Remember: n(empirical formula) = molecular formula
Step 3: Multiply the subscripts in the empirical formula by n.

31. Sample Problem #1 Sample Problem #2
The empirical formula for a compound is P2O5. Its experimental molar mass is 284g/mol. Determine the molecular formula of the compound.
A brown gas has the empirical formula NO2. Its experimental molar mass is 46g/mol. What is the molecular formula?
Sample Problem #2

32. Hydrates- Honors Only Not in the textbook.
Hydrates – ionic compounds that contain water molecules within the crystal structure.
Example: CuSO4•5H2O
Anhydrous – without the water = CuSO4
Although your textbook does not cover this – it really is an application of the empirical formula problems. Instead of using elements, you will use the anhydrous part and the water to find the lowest ratio.
Silica gel – the little white packets you find in leather and electronics are anyhdrous materials that will absorb any moisture to keep items dry.

33. Determining Hydrate Formulas
Formula can be determined if given: the mass of the hydrate, the anhydrous mass, and the formula of the ionic compound.
Step 1: Determine the mass of water in the hydrate (subtract anhydrous mass).
Step 2: Convert the anhydrous ionic compound mass and water mass to moles.
Step 3: Divide both molar amounts by the smallest number. This gives you the number of water molecules in the hydrate.

34. Sample Problem #1
A 5.82 g sample of Mg(NO3)2· XH2O in an evaporating dish is heated until it is dry.  The mass of the anhydrous sample is 2.63 g Mg(NO3)2.  What is the formula for the hydrate?

35. Determining % Water in a Hydrate
Formula can be determined if given the formula of the hydrate.
Step 1: Calculate the mass of the entire hydrate and the mass of just the water.
Step 2: Divide the mass of the water by the mass of the entire hydrate and multiply by 100 to get a percent.

36. Sample Problem #2
What percentage, by mass, of water is found in the hydrate CuSO4·5H2O?