Name:



Class Notes: The Mole and Stoichiometry: Using Moles

Objective —Use factor-labeling to convert between moles, masses and number of particles.

Units you already know…

A dozen is ______ objects. A ream of paper is __________ sheets.

An hour is ______ seconds. A decade is _______ years.

New Unit: A mole is ________________________ particles.

Ex: A mole of beans are ________________________ beans.

A mole of pencils are ________________________ pencils.

A mole of _____________ is _____________________ ___________.

The Mole (Avogadro’s Constant)

Where did the mole come from? The mole is a quantity of a substance such that the actual mass in grams has the same value as the average atomic mass listed on the periodic table. It was discovered by Amadeo Avogadro in 1811 and is known as Avogadro’s rule. Avogadro's rule says that a mole of any substance always contains the same number of molecules. A mole of oxygen has as many molecules as a mole of hydrogen, which has as many molecules as a mole of water.

Huh? In other words, the mole is the number of carbon-12 atoms needed to make 12.0 grams of carbon-12. Or the number of oxygen-16 atoms needed to make 16.0 grams of oxygen-16.

In addition, you should remember that the mass of __________ mole has the _________ value as the ______________________ listed in the ______________________. So…

1 mole of any element = average atomic mass of that element in grams

(found on the periodic table)

1 mole of Carbon = _____________ grams of Carbon

1 mole of Nitrogen = _____________ grams of Nitrogen

1 mole of Hydrogen = _____________ grams of Hydrogen

1 mole of Copper = _____________ grams of Copper

Got it? Let’s practice and see

Solve the following problems using factor-labeling….

1) What is the mass of one mole of nitrogen? _______________________

2) What is the mass of five moles of nitrogen? _______________________

3) How many atoms are in one mole of nitrogen? _____________________

4) How many atoms are in five moles of nitrogen? ____________________

5) What is the mass of one mole of iron? ____________________

6) What is the mass of three moles of iron? ____________________

Part II. Molar Mass for Compounds

Like individual atoms, each molecule or ion has a characteristic average mass. For example, we know from the chemical formula H2O that a single molecule of water is composed of exactly _____ hydrogen atoms and _____ oxygen atom. The molar mass of a water molecule is found by _____________ the masses of the three atoms in the molecule.

One mol of Hydrogen = _________ gram

One mol of Oxygen = __________ grams

2 H mol x 1.008 g = 2.016 g

1 H mol

1 O mol x 16.00 g = 16.00 g

1 O mol

2.016 g + 16.00 g = 18.016 g = mass of one mole H2O

1) What is the molar mass of magnesium chloride, MgCl2?

2) What is the molar mass of a phosphate ion, PO4?

3) What is the molar mass of dinitrogen trioxide, N2O3?

Name: ____________________ Block: _____ ________________

Practice – Molar Mass

Instructions: Show your work for the following problems. Circle your final answer with the label.

Part A. Answers should be in grams.

1. What is the mass of one mole of lithium? __________________

2. What is the mass of one mole of calcium? __________________

3. What is the mass of one mole of silver? __________________

4. What is the molar mass of one mole of fluorine, F2? __________________

5. What is the molar mass of one mole of KBr? __________________

6. What is the mass of 4 moles of lithium?

7. What is the mass of 7 moles of strontium?

8. What is the mass of 2 moles of KBr?

9. What is the mass of 3 moles of NaCl?

10. What is the mass of 4 moles of AgO?

11. What is the mass of 3 moles of water?

12. What is the mass of 7 moles of beryllium sulfide (BeS)?

13. What is the mass of 5 moles of hydrochloric acid (HCl)?

Part. B Calculate the moles present in: (Leave answers in decimal form)

1) 2.00 grams of H2O

2) 75.57 grams of KBr

3) 100. grams of KClO4

4) 8.76 grams of NaOH

5) 26.0 gram CaCl

6) 5.08 gram XeF4

7) 32.0 gram O2

8) 10.0 gram V2O5

10.0 gram Al2(SO4)3

Practical Application:

1) If pencil weighs 3.80 grams, how much would a mole of pencils weigh?

2) If a peron eats 200 beans a week, how many weeks would it take for a person to eat a mole of beans?

3) If the cafeteria serves 1,600 lunches a day, how many days would it take the cafeteria to serve a mole of lunches? If the school year is 180 days, how many years would it take to serve?

Additional Problems:

Calculate the mass in grams of each of the following quantities.

9) 0.200 moles of H2S

10) 0.100 moles of KI

11) 1.500 moles of KClO

12) 0.750 moles of NaOH

13) 3.40 x 10-5 moles of Na2CO3

14) 2.55 mole Cu2CrO4

15) 1.95 mole HNO3

16) 2.00 mole HC2H3O2

17) 10.0 mole NaCl

18) 2.20 mole SnCl2

Practice Problems

Complete all work on separate paper.

1) How many moles of compound are there in the following?

a. 4500 g calcium hydroxide, Ca(OH)2

b. 25.0 g sulfuric acid, H2SO4

c. 125 g of sugar, C12H22O11

2) What is the mass of each of the following?

a. 1.00 mol of beryllium chloride, BeCl2

b. 3.50 mol of dinitrogen tetroxide, N2O4

c. 0.625 mol of barium nitrate, Ba(NO3)2

3) What is the mass in grams of 6.25 mol of copper (II) nitrate?

4) How many moles are in 123 g of KCl?

The Mole and Stoichiometry: Molar Math Practice

Objective 8-3—Use dimensional analysis to convert between moles, masses and number of particles.

Molar Math Practice

Complete all work on separate paper and write your final answers on this page.

Be prepared to show your work.

Calculate the moles present in:

19) 2.00 grams of H2O

20) 75.57 grams of KBr

21) 100. grams of KClO4

22) 8.76 grams of NaOH

23) 26.0 gram Ca(ClO4)2

24) 5.08 gram XeF4

25) 10.0 gram Al2(SO4)3

26) 32.0 gram O2

27) 10.0 gram V2O5

28) 2.50 gram CoSO4 . 6H2O

Calculate the mass in grams of each of the following quantities.

29) 0.200 moles of H2S

30) 0.100 moles of KI

31) 1.500 moles of KClO

32) 0.750 moles of NaOH

33) 3.40 x 10-5 moles of Na2CO3

34) 2.55 mole Cu2CrO4

35) 1.95 mole HNO3

36) 2.00 mole HC2H3O2

37) 10.0 mole NaCl

38) 2.20 mole SnCl2

Calculate the number of molecules in each of the following quantities.

39) 3.00 mole H2

40) 3.27 mole O2

41) 0.000300 mole AuCl3

42) 1.55 mole KrF2

43) 0.100 mole NH3

Objective 8-3 Review

Complete all work on separate paper and write your final answers on this page.

Be prepared to show your work.

1) What is the mass of one mole of arsenic? ____________________

2) How many atoms are in one mole of arsenic? ____________________

3) How many moles are contained in 2.005 x 1023 atoms of arsenic? _______________

4) How many moles are in 200.5 grams of arsenic? ________________

5) What is the formula mass of zinc fluoride, ZnF2? _______________

6) What is the formula mass of a sulfate ion, SO42-? _______________

7) What is the molar mass of hydrogen peroxide, H2O2? _______________

8) What is the molar mass of iron (III) hydroxide, Fe(OH)3? _______________

9) How many moles of compound are there in the following?

a. 1.023 g copper (II) nitrate, Cu(NO3)2 _______________

b. 17 g of fructose, C6H12O5 _______________

10) What is the mass of each of the following?

a. 2.30 mol of phosphorus trifluoride, PF3 _______________

b. 0.034 mol of selenium sulfide, SeS _______________

11) What is the mass in grams of 0.52 mol potassium chloride? _______________

12) How many moles are in 123 g of salt? _______________

13) What is the difference between formula mass and molar mass?

14) Compare the size of a mole of salt with a mole of beans—how many particles are in each? What is the mass of each?

15) What is a conversion factor? How are they useful when converting?

Objective 8-3 Review (Part II)

Complete all work on separate paper and write your final answers on this page.

Be prepared to show your work.

Calculate the number of molecules in each of the following quantities.

1) 3.00 mole Zr

2) 3.27 mole Zn

3) 0.000300 mole Au

4) 1.55 mole Kr

5) 0.100 mole Ni

6) 0.00550 mole C

7) 0.100 mole W

8) 0.500 mole Ca

9) 0.300 mole P

10) 1.500 mole K

Calculate the moles present in:

11) 2.00 grams of Ca

12) 75.57 grams of Fe

13) 100. grams of K

14) 8.76 grams of Na

15) 26.0 gram Ca

16) 5.08 gram Xe

17) 10.0 gram Al

18) 32.0 gram P

19) 10.0 gram V

20) 2.50 gram Co

Calculate the mass in grams of each of the following quantities.

21) 0.100 moles of K

22) 1.500 moles of Cr

23) 0.750 moles of Na

24) 3.40 x 10¯5 moles of C

25) 2.55 mole Cu

26) 1.95 mole Ba

27) 2.00 mole Ag

28) 10.0 mole Os

29) 2.20 mole

30) 0.200 moles of Sn

The Mole and Stoichiometry: Stoichiometry

Objective 8-4—Use mole ratios to solve stoichiometry problems.

STOICHIOMETRY PROBLEM-SOLVING

When solving stoichiometry problems, you should always follow the 5-step method below.

1 balance equation

2 identify given

3 identify want

4 set up equation

5 solve the equation

1) __ LiOH + __ HBr ( __ LiBr + __ H2O

If you start with ten grams of lithium hydroxide, how many grams of lithium bromide will be produced?

2) __ C2H4 + __ O2 ( __ CO2 + __ H2O

If you start with 45 grams of ethylene (C2H4), how many grams of carbon dioxide will be produced?

3) __ Mg + __ NaF ( __ MgF2 + __ Na

If you start with 5.5 grams of lithium chloride, how many grams of calcium chloride will be produced?

4) __ HCl + __ Na2SO4 ( __ NaCl + __ H2SO4

If you start with 20 grams of hydrochloric acid, how many grams of sulfuric acid will be produced?

Stoichiometry Problems with a Limiting Reactant

Using your knowledge of stoichiometry and limiting reactants, answer the following questions:

1) Write the balanced equation for the reaction of lead (II) nitrate with sodium iodide to form sodium nitrate and lead (II) iodide:

2) If I start with 25.0 grams of lead (II) nitrate and 15.0 grams of sodium iodide, how many grams of sodium nitrate can be formed? SHOW YOUR WORK.

3) What is the limiting reagent in the reaction described in the above problem?

4) How much of the non-limiting reagent will be left over from the reaction in the above problem?

5) Write the balanced equation for the reaction of hydrogen gas and oxygen gas to form water:

6) If I start with 20 g of hydrogen gas and 20 g of oxygen gas, how many grams of water can be formed? SHOW YOUR WORK.

7) What is the limiting reagent in the reaction described in the above problem?

8) How much of the non-limiting reagent will be left over from the reaction in the above problem?

More Stoichiometry Problems

__ C6H12O6 + __ O2 ( __ CO2 + __ H2O + energy

1. Use the above equation to determine how many grams of CO2 are produced from 3.00 moles of C6H12O6.

__ NO2 + __ O2 ( __ NO3

2. Use the above equation to determine how many grams of nitrogen dioxide are needed to produce 1.55 X 1040 molecules of nitrogen trioxide.

__ CaCO3 ( __ CO2 + __ CaO

3. Use the above equation to determine if 1,212 g calcium carbonate are consumed, how many grams of calcium oxide are produced.

Zinc reacts with excess sulfuric acid to produce zinc sulfate and hydrogen gas

4. In the above reaction, how many atoms of zinc are needed to produce 7 moles of hydrogen gas?

Barium hydroxide plus sulfuric acid produces barium sulfate and water.

5. In the above reaction, what mass of barium sulfate is formed when 75.0 g of barium hydroxide is missed with an excess of sulfuric acid?

__ Fe + __ CuSO4 ( __ FeSO4 + __ Cu

5. In the above reaction, if 48.2 grams of iron react with 5.62 moles cupric sulfate, how many moles of copper can be produced?

6. What is the limiting reagent in the reaction described in the above problem? Explain.

7. How much of the non-limiting reagent will be left over from the reaction in the above problem?

__ Zn + __ H2SO4 ( __ H2 + __ ZnSO4

8. A student reacts 8.32 grams zinc with excess sulfuric acid and produces hydrogen gas and zinc sulfate. How many moles of zinc sulfate will be produced?

The Mole and Stoichiometry: Percent Composition

Objective 8-5—Calculate the percent composition of a certain element in a compound.

Calculating Percent

What is a percent? A percent is a ratio given in terms of parts per hundred (remember: there are 100 cents in a dollar). These ratios are usually expressed in percents and we can use the cross-product method to find our percentage.

For example, let’s say there are five pieces of cake and Doris eats 4 of them. What percentage of the cake did she eat?

To solve this, we set up a proportion and then cross multiply to find the missing value.

4 = x First, multiply the cross: (4)(100) = (5) (x)

5. 100

Then divide to isolate the variable: (4)100 = 5x

5. 5

80 = x

Doris ate 80% of the cake

Now another way to look at the equations we just solved is to see that we are taking the fraction and multiplying it by 100 4 x 100 = 80

5

And it is this equation which is more applicable when looking at percent composition.

Part x 100 = the percent

Whole

Second example: Central High School must hold classes 180 days a year. If there are 365 days in a year, what percentage of the year is spent in school?

180 x 100 = 49.3%

365

Practice Problems (show all work on separate paper):

1) At a basketball game, Sarah shoots 12 times and scores 9 of those times. What is her shooting percentage?

2) Joe just got his nine weeks tests back and wants to find out which test he did best on. He got a 84/95 on his history exam, a 48/50 on his English exam and a 94/120 on his math exam. What were his percentages on each of those tests? Which exam did he perform the best on?

3) There are 50 states in the United States and 11 of them border the Mississippi River. What percentage of states are along the Mississippi?

4) Calculate the percentage of students wearing khakis in this class

5) Calculate the percentage of girls in this class.

Calculating Percent Composition

A chemical formula can tell us the ratio of atoms in a compound. For example, the formula for potassium chlorate tells us that there is 1 potassium atom, 1 chlorine atom, and three oxygen atoms bonded together in one molecule of KClO3.

However, it is also useful for scientists to know what percentage by mass of a particular element in a chemical compound. Suppose we wanted to use potassium chlorate as a source for oxygen—we would want to know how much oxygen is in a sample of KClO3. How can we figure this out? Well we use the basic formula for percents and take the part divided by the whole:

Mass of element in 1 mol of compound x 100 = % element in compound

Molar Mass of Compound

The percentage by mass of each element in a compound is known as the percentage composition of the compound

Example What is the percent composition of potassium chlorate, KClO3?

1) First, find the molar mass:

39.098 + 35.45 + 3(16.00) = 122.55 g/mol

2) Next, find the masses of each element by multiplying the molar mass of the element by the number of atoms of that element in the compound:

K = 1(39.098 g) = 39.098 g

Cl = 1 (35.45 g) = 35.45 g

O = 3(16.00) = 48.00 g

3) Now, take the mass of each element divided by the mass of the compound to find the percent of each element:

K: 39.098 x 100 = 31.9% Cl: 35.45 x 100 = 28.9% O: 48.00 x 100 = 39.2%

122.55 122.55 122.55

4) Check your answer to see if your percents add up to 100%.

31.9% + 28.9% + 39.2% = 100%

Potassium chlorate is 31.9% K, 28.9% Cl and 39.2% O.

Practice Problems (show all work on this paper. If you run out of room you may show some work on separate paper):

1) Find the percent composition of copper (I) sulfide, Cu2S.

2) What is the percent composition of barium nitrate, Ba(NO3)2?

3) Calculate the percent composition of ammonium carbonate, (NH4)2CO3.

4) Determine the percent composition of sodium chloride, NaCl.

5) Find the percent composition of silver nitrate AgNO3.

6) What is the percent composition of sulfurous acid, H2SO3?

7) Calculate the percent composition of beryllium chloride, BeCl2.

8) Determine the percent composition of diarsenic pentoxide, As2O5.

9) Which has a greater percent composition of oxygen—iron (II) oxide (FeO) or iron (III) oxide (Fe2O3)? Explain.

10) Magnesium hydroxide is 54.87% oxygen by mass. How many grams of oxygen are in 175g of the compound? How many mole of oxygen is this?

Practice Problems (Calculate the percentage of EACH element in the following compounds. Show all work on separate paper and write answers on this paper. Be prepared to show your work.):

1) KNO3

2) H2SO4

3) C2H5OH

4) C6H5NH2

5) lithium bromide, LiBr

6) anthracene, C14H10

7) ammonium nitrate, NH4NO3

8) nitrous acid, HNO2

9) silver sulfide, Ag2S

10) iron(II) cyanide, Fe(CN)2

11) lithium acetate LiCH3COO

12) nickel(II) fluoride NiF2

Practice Problems (Calculate the percentage of THE GIVEN element in the following compounds. Show all work on THIS paper and write answers on this paper.):

13) nitrogen in urea, NH2CONH2

14) sulfur in sulfuryl chloride, SO2Cl2

15) thallium in thallium(III) oxide, Tl2O3

16) oxygen in potassium chlorate, KClO3

17) bromine in calcium bromide, CaBr2

18) tin in tin(IV) oxide, SnO2

19) oxygen in chromium permanganate, CrMnO4

20) nitrogen in calcium nitrate, Ca(NO3) 2

Working Backwards

Some Chemistry problems are tough because you cannot always just plug numbers into the formula as you have learned them. The important part of algebra is that it teaches that we can rearrange equations to solve for whatever we’re missing.

Magnesium hydroxide is 54.87% oxygen by mass. How many grams of oxygen are in 175g of the compound? How many mole of oxygen is this?

The equation we learned taught us that:

Mass of element in 1 mol of compound x 100 = % element in compound

Molar Mass of Compound

So here we need to work backwards—instead of knowing the mass of one element and the molar mass, we know the mass of a sample and the percentage.

Mass of oxygen x 100 = 54.87% OR Mass of oxygen = 54.87% x 175 g

175 g 100

Therefore the mass of oxygen is 96.0 g

As we learned in Obj 3, we can use the molar mass to convert to moles:

96.0 g O x 1 mole O = 6.0 moles O

16.00 g O

Let’s try that technique on another similar problem:

Methane is 75% carbon and 25% hydrogen. You have 16.05 g of methane –how many grams of carbon do you have? How many grams of hydrogen? How many moles of carbon do you have? How many moles of hydrogen?

Mass of carbon x 100 = 75% OR Mass of carbon = 75% x 16.05 g

16.05 g 100

Therefore the mass of carbon is 12.01 g = 1 mol C

Mass of hydrogen x 100 = 25% OR Mass of hydrogen = 25% x 16.05 g

16.05 g 100

Therefore the mass of hydrogen is 4.04 g = 4 mol H

How does the relationship between carbon and hydrogen in this problem relate to the chemical formula for methane?

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The Mole and Stoichiometry: Empirical and Molecular Formulas

Objective 8-6— Predict the empirical formula and calculate the molecular formula of a compound using the masses of each element.

Determining Empirical Formula

The methane problem on the last page was easy because our amount of sample was exactly equal to the molar mass of methane (16.05g). However, our lives are not always that convenient. Frequently, we will have a larger amount that is a nice round number (like 100g). In these cases, the moles do not appear in nice round numbers. Therefore we have to divide the moles to find an easy mole ratio.

Example 1: Let’s say we have 100g of the gas diborane. It is 78.1% boron and 21.9% hydrogen. What is the empirical formula for diborane?

First we use the percentage composition formula to find the mass

(Percent( Mass)

Mass boron = 78.1% x 100 = 78.1g Mass of hydrogen = 21.9% x 100 = 21.9g

100 100

Next, convert the masses to moles by using molar mass as a conversion factor

(Mass(Mole)

B: 78.1 g B x 1 mol B = 7.22 mol B H: 21.9 g H x 1 mol H = 21.7 mol H

10.81 g B 1.008g H

Then, we need to figure out the mole ratios. Right now we have 7.22 mol B to 21.7 mol H. How can we simplify this? Just like we do with fractions, divide by the smaller number (in this case 7.22 mol )

7.22 mol B = 1 mol B 21.7 mol H = 3.01 mol H

7.22 mol B 7.22 mol B

Now figure out what the formula will be: 1 boron and 3 hydrogens OR BH3.

Example 2: Analysis tells us that a compound is 32.38% sodium, 22.65% sulfur, and 44.99% oxygen. Find the empirical formula for this compound.

First, find the masses of each element. (Percent ( Mass)

We don’t have a given mass so let’s pretend that we had 100g. It doesn’t much matter because the ratio will be the same no matter what. We like 100g because then the percentage is the same as the mass. (32.38 g Na, 22.65g S, and 44.99 g O)

Next, find the moles of each element (Mass ( Moles)

Na: 32.38 g Na x 1 mol Na = 1.408 mol Na S: 22.65 g S x 1 mol S = 0.7063 mol S

22.99 g Na 32.07g S

O: 44.99 g O x 1 mol O = 2.812 mol O

16.00 g O

Then divide by the smallest (in this case, 0.7063 mol S) S: 0.7063 mol S = 1 mol S

0.7063 mol

Na: 1.408 mol Na = 1.993 mol Na O: 2.812 mol O = 3.981 mol O

0.7063 mol 0.7063 mol

Although the numbers aren’t exact, because we are within .01 of a whole number, we can round to make our answers complete

2 mol Na: 1 mol S: 4 mol O ( Na2SO4 (sodium sulfate)

But what is the numbers aren’t exact and are not within 0.01 of a whole number?

Then we have to multiply both numbers until they are whole.

If the decimal is close to 0.50, multiply by 2

0.25, 0.75 multiply by 4

0.33, 0.66 multiply by 3

Example 3: Analysis of a 10.150 g sample of a compound known to contain only phosphorus and oxygen indicates a phosphorus content of 4.433 g. What is the empirical formula for this compound?

Here we have a given mass so we do not need to worry about the percent composition instead we can start with step 2 (mass( mole)

Total Mass: 10.150 g

Mass P: 4.433 g

Mass O: 10.150 g – 4.433 g = 5.717 g

P: 4.433 g P x 1 mol P = 0.1431 mol P O: 5.717 g O x 1 mol O = 0.3573 mol O

30.97 g P 16.00 g O

Now divide by the smallest (in this case 0.1431 mol)

P: 0.1431 mol P = 1 mol P O: 0.3573 mol O = 2.497 mol O

0.1431 mol 0.1431 mol

2.497 is not close to a whole number but it is close to 0.50 so we will multiply by numbers by 2

1 mol P x 2 = 2 mol P 2.497 mol O x 2 = 4.994 mol O

Since 4.994 is within 0.01 of a whole number we can round it up to 5

The formula is 2 mol P to 5 mol O or P2O5 (Diphosphorus Pentoxide)

Want an easy way to remember these steps? Just learn this poem!

Percent to mass

Mass to mole

Divide by the small

Multiply ‘til whole

Practice Problems ( Show all work on THIS paper and circle your final answers.):

1) A compound is found to contain 63.52% iron and 36.48% sulfur. Find its empirical formula.

2) Find the empirical formula of a compound found to contain 26.56% potassium, 35.41% chromium, and the remainder oxygen.

3) Analysis of a 20.0 g of a compound containing only calcium and bromine indicates that 4.00 g of calcium are present. What is the empirical formula of the compound?

4) Determine the empirical formula of a compound containing 63.50% silver, 8.25% nitrogen and the remainder oxygen.

5) Determine the empirical formula of a compound found to contain 52.11% carbon, 13.14% hydrogen and 34.75% oxygen.

6) A compound is found to contain 36.48% Naa, 25.41% S and 38.11% O. Find its empirical formula.

7) Find the empirical formula of a compound that contains 53.70% iron and 46.30% sulfur.

8) Analysis of a compound indicates that it contains 1.04 g K, 0.70 g Cr, and 0.86 g O. Find its empirical formula.

9) Chemical analysis of citric acid shows that it contains 37.51% C, 4.20% H, and 58.29% O. What is its empirical formula?

10) A 175.0 g sample of a compound contains 56.15 g C, 9.43 g H, 74.81 g O, 13.11 g N, and 21.49 g Na. What is its empirical formula?

Empirical Practice and Molecular Formulas

An empirical formula contains the smallest possible whole numbers that describe the atomic ratio. However, many molecular compounds have the same empirical formula but have unique molecular formulas—which are the actual formulas of a molecular compound.

For example, diborane’s empirical formula is BH3. Any multiple of BH3, such as B2H6, B3H9, B4H12 and so on represents the same ratio of B: H so will have the same empirical formula. However, each of these is a unique substance. How can we distinguish between them? By looking at their molecular formula.

We can find molecular formula by multiplying the empirical formula by a whole number

x (empirical formula) = molecular formula

1 (BH3) = BH3

2 (BH3) = B2H6

3 (BH3) = B3H9

4 (BH3) = B4H12

Formula masses have a similar relationship:

x (empirical formula mass) = molecular formula mass

We rearrange our equation to solve for our variable and get x = molecular formula mass

empirical formula mass

Example 1: Experimentation shows that the formula mass of diborane is 27.67 amu. What is the molecular formula for diborane?

First, follow the steps to find the empirical formula

(see 5-2) BH3

Second, calculate the formula mass of the empirical formula

(see 3-2) 10.81 amu + 3 (1.008 amu) =13.84 amu

Third, divide molecular formula mass by empirical formula mass to solve for x

x = 27.67 amu = 2.00

13.84 amu

Fourth, multiply the empirical formula by x to find the molecular formula

2 (BH3) = B2H6

Which of course makes sense because di- means two

Example 2: The compound studied in Example 3 on 5-2 (between phosphorus and oxygen) has been found to have a molar mass of 283.89 g/mol. What is the compound’s molecular formula?

First, find the empirical formula P2O5

Second, calculate the formula mass of the empirical formula

(see 3-2) 2(30.97 amu) + 5 (16.00 amu) =141.94 amu

Third, divide molecular formula mass by empirical formula mass to solve for x

x = 283.89 amu = 2.00

141.94 amu

Fourth, multiply the empirical formula by x to find the molecular formula

2 (P2O5) = P4O10

Tetraphosphorus Decoxide

Practice Problems ( Show all work on THIS paper and circle your final answers.):

1) Determine the molecular formula of the compound with an empirical formula of CH and a formula mass of 78.110 amu.

2) A sample of a compound with a formula mass of 34.000 amu is found to consist of 0.44 g H and 6.92 g O. Find its molecular formula.

3) If 4.04 g N combine with 11.46 g O to produce a compound with a formula mass of 108.0 amu, what is the molecular formula of this compound?

4) The molar mass of a compound is 92 g/mol. Analysis of a sample of the compound indicates that it contains 0.606 g N and 1.390 g O. Find its molecular formula.

5) What is the molecular formula of the molecule that has an empirical formula of CH2O and a molar mass of 120.12 g/mol?

6) A compound with a formula mass of 42.08 amu is found to be 85.64% carbon and 14.16% hydrogen by mass. Find its molecular formula.

The Mole and Stoichiometry: Empirical and Molecular Formulas ACTIVITY

Objective 8-6— Predict the empirical formula and calculate the molecular formula of a compound

using the masses of each element.

The Strange Case of Mole Airlines Flight 1023

You and your team of medical examiners are called to the scene of a plane crash that has no survivors. The plane shows evidence of a pre-crash explosion. The site of the explosion has a compound with the following analysis: 37.01% C, 2.22% H, 18.5% nitrogen, and 42.27% oxygen. The victims are found and in and around the crash and must be identified by the substance found in their belongings or in their bodies as dental records are not available. One passenger was murdered with the time of death established as 1 hour before the crash.

Your job as the forensic scientist on duty is to:

1. Use the percent compositions to determine the chemical formulas and identities for the compounds found on the victims. (Use a separate sheet of paper and show all work)

2. Use personal data to make a probable identification of each victim.

3. Determine who was murdered and who is the most probable murderer.

4. Determine who caused the plane to explode.

|Victim # | |Analysis of |Compound | |Where found |

|1 |67.31% C |6.96% H |4.62% N |21.10 % O |Blood &Luggage |

|2 |63.15% C |5.30% H | |31.55% O |Briefcase |

| |46.66 % C |4.48% H |31.10% N |17.76 % O |Stomach |

|3 |72.20 % C |7.08 % H |4.68 % N |16.03 % O |Pockets |

|4 |15.87% C |2.22 % H |18.51 % N |63.41 % O |Blood & Pockets |

|5 |75.42 % C |6.63 % H |8.38 % N |9.57 % O |Blood |

| |37.01 % C |2.22 % H |18.50 % N |42.27 % O |Luggage |

|6 |57.14% C |6.16 % H |9.52 % N |27.18 % O |Briefcase |

|7 |80.48% C |7.45 % H |9.39 % N |2.68 % O |Briefcase |

| |81.58 % C |8.90 % H |9.52 % N | |Luggage |

|8 |60.00 % C |4.48 % H | |35.53 % O |Pockets & Briefcase |

| |63.56 % C |6.00 % H |9.27 % N |21.17 % O |Pockets & Briefcase |

| |75.42 % C |6.63 % H |8.38 % N |9.57 % O |Briefcase |

List of substances and their common uses

Acetaminophen C8H9NO2 Tylenol, non-prescription painkiller

Aspartame C14H18N2O5 artificial sweetener

Aspirin C9H8O4 painkiller, over the counter

Codeine C18H21NO3 prescription painkiller, often used in dentistry

Cocaine C17H21NO4 narcotic, illegal substance

Curare C40H44N4O poison

Dimetacrine C40H52N4 antidepressant, prescription

Nitroglycerine C3H5N3O9 explosive; heart medication

Strychnine C21H22N2O2 rat poison

Thriobromine C7H8N4O2 chocolate flavoring used in frosting on pastries

Trinitrotoluene (TNT) C7H5N3O6 TNT, explosive

Vanilla C8H8O3 flavoring, used in baking

List of Passengers and Crew

Amadeo Aviator - a pilot with a heart condition

Curie Drugs - a sales representative of Advil company

Bohr Diet - Chemistry teacher "addicted" to sugar-free drinks

Kelvin Doughnut - baker, specializes in pastries shaped like the famous scientist

Democritus "D" Victory - professional athlete suspended for drug violations

Ruthy Ford-Recycle - electrical engineer, had been severely depressed lately due to a canceled project

Dalton Molar - high school student who recently underwent dental surgery

Norm Einstein - leader of group called the Mighty Order of Lavoisier Enthusiasts, a suspected terrorist organization more commonly know as MOLE. According to Mole-der and Skulley, the groups' goal is to make everyone aware of the power of Chemistry.

1. What compound was found at the crash site?

2. Who is victim #1? What compound(s) did he/she have on them?

3. Who is victim #2? What compound(s) did he/she have on them?

4. Who is victim #3? What compound(s) did he/she have on them?

5. Who is victim #4? What compound(s) did he/she have on them?

6. Who is victim #5? What compound(s) did he/she have on them?

7. Who is victim #6? What compound(s) did he/she have on them?

8. Who is victim #7? What compound(s) did he/she have on them?

9. Who is victim #8? What compound(s) did he/she have on them?

10. Who was the murdered victim?

11. Who was the murderer?

12. Who caused the plane to explode? Why?

The Mole and Stoichiometry: Empirical and Molecular Formulas LAB

Objective 8-6— Predict the empirical formula and calculate the molecular formula of a compound

using the masses of each element.

By knowing the relative weights of the elements in a compound, the empirical formula (or simplest chemical formula) can be deduced. This is a common practice in forensic science. In this experiment the percentage of the carbon in a carbonate salt will be calculated. As the carbonates are decomposed by their reaction with hydrochloric acid, carbon dioxide is released as a gas. The decrease in mass of the salt will correspond to the amount of gas liberated. By knowing the mass of gas, carbon dioxide, released we can calculate the mass of the carbon in the salt sample originally and calculate the percentage of carbon in the salt.

Safety Note: Chemical goggles and laboratory apron should be worn. Use extreme care when obtaining, transporting, dispensing, and discarding hydrochloric acid and its solutions. It is extremely corrosive.

Materials:

• Unknown Carbonate salt

• hydrochloric acid, 6M

• pipet

• spot plate

• balance

Procedure:

1. Place a spot plate on a balance and record the mass to the nearest 0.01 g.

2. Obtain a small sample ( ................
................

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