Academic Chemistry UNIT 5 – CHEMICAL

QUANTITIES & THE MOLE Name: ____________________________ Period: ____

MOLE QUIZ: __11/3/14__ MOLE TEST: __11/13/14____

Chemistry Calendar

Monday Tuesday Wednesday Thursday Friday October 27

28 Notes: Molar Mass & Mole Map Conversion Factors (Hwk: pg. 7)

29 Particles to Moles (calculations) (Hwk: pg. 10) Moles of Chalk Lab

30 Volume <->Mole Conversions: Mixed Calculations (Hwk: pg. 12)

31 Multi-step Conversions (g > mol > L) (Hwk: pg. 14)

November 3 Conversions QUIZ % Composition Problems % Bubble gum lab (Hwk: pg. 18)

4 Empirical Formulas (Hwk: pg. 22)

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Cumulative 6 weeks TEST

6 Empirical & Molecular Formulas (Hwk: pg 22)

7 End 2nd 6 weeks Half day

10 “Crime Scene” Activity (Hwk: Review Pg.23-25)

11 Set up Penny Lab (Hwk: Review Pg.23-25)

12 Test Review Due (p. 21-23) % Comp of a Penny Lab

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Unit 5-Mole TEST

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DLA

Unit 5: MOLE MATH

We use a unit called the _____________________, or _______________, to measure the amount of a

substance. The mole can represent _________________ (__________), _____________________

(___________), or ________________________ (__________).

*You must ALWAYS go through the MOLE!*

• 1 mole = _______________________________ • 1 mole = _______________________________ • 1 mole = _______________________________

Avogadro’s Number = and is also called a

• 1 dozen = _______________ items.

• 3 dozen cookies = ______________ cookies

• .5 dozen doughnuts = _________ doughnuts

• A “dozen” is a counting unit equal to _________ of any object.

• A “Mole” is a counting unit equal to of any object, even really

small ones like , , or .

What is so special about 6.02 X 1023? Why do scientists use that number?

________________________________________________________________________________

________________________________________________________________________________

________________________________________________________________________________

Example Conversion Factors:

o 1 mole = 6.02 x 1023 particles (atoms, molecules, or formula units)

o 1 mole of copper = 6.02 x 1023 of copper

o 1 mole of CuCl2 = formula units of CuCl2

o 1 mole of CO2 = molecules of CO2

The amount of substance containing Avogadro’s number of any kind of chemical unit is called a _____________ of

that substance.

o One mole of K contains atoms.

o One mole of NaOH contains formula units.

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Molar Mass Molar mass is the mass of ______________________________ of a substance.

• Other names for molar mass include:

•

•

•

One ________________ of any element will have a mass in grams corresponding to the value of its

___________________________. For example:

• 1 mol Carbon = ____________ g/mol

• 1 mol Nitrogen = ____________ g/mol

• its atomic mass from the _______________________________.

Diatomic Elements The following elements do not naturally occur as one atom- they are always in a pair:

1.

2.

3.

4.

5.

6.

7.

*Trick to remembering the seven diatomics*

One of any molecule/compound will have a mass in grams corresponding to the value of its molar

mass (the sum of the masses of the elements that compose it). Ex. H2O H: 1.008 X 2 = 2.016

O: 15.999 X 1 = 15.999 ----------------------- 18.015 g/mole

Ex. Ca(NO3)2 Ca: 40.078 x 1 = 40.078

N: 14.007 x 2 = 28.014 O: 15.999 x 6 = 95.994 --------------------- 164.086 g/mole

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Sample Problems #1-2: Calculate the molar mass for the following compounds:

1. Potassium carbonate _________________

# of K atoms: x =

+

# of C atoms: x =

+

# of O atoms: x =

________________

2. Ammonium sulfate ____________________

# of N atoms: x =

+

# of H atoms: x =

+

# of S atoms: x =

+

# of O atoms: x =

________________

PRACTICE: (SHOW ALL WORK & UNITS to receive full credit.)

1. What is the atomic mass of sodium?

2. Calculate the molar mass for Al2(SO4)3.

3. Calculate the molar mass for nitrogen (hint: is it diatomic!?!?!).

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Mole Highway

PRACTICE: Moles & Mass

Draw the Mole Road Map:

NOTES: To convert between units, follow the highway. Notice, there is no shortcut from grams to liters or between any of the three units surrounding the mole. This means you have to convert to ______________ before converting to another unit!

1 MOLE

1 mol =

= 1 mol =

Mole 22.4 L Volume

(gases only)

Molar Mass (from periodic table)

in grams

6.02 X 1023 Representative Particles

(atoms, ions, molecules, formula units)

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Molar mass:

• called gram atomic mass when single element is used. • called gram formula unit when ionic compound is used. • called gram molecular unit when molecular compound or diatomic molecules used.

o Diatomic molecules are atoms that bond with themselves. There are SEVEN of these that you need to remember: Br2 I2 N2 Cl2 H2 O2 F2 (Remember this by the name “BRINClHOF”)

Conversion Factors: 1 mole = 6.02 x 1023 particles (Avogadro’s number) 1 mole = molar mass (Sum of atomic masses in grams) 1 mole = 22.4L of any gas at STP (Standard temperature & pressure) Any gas density at STP = molar mass / molar volume

PRACTICE: (SHOW ALL WORK & UNITS to receive full credit.) Determine the number of moles in each of the quantities below.

1. 25 g sodium 1. 1.

2. 85 g H2SO4 2.

Determine the number of grams in each of the quantities below.

3. 2.5 moles of sodium 3. 3.

4. 0.50 moles of H2SO4 4.

Solve the following mole-mass conversion problems.

5. How many moles are there in 27 g of ethanol (C2H5OH)? 5.

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Homework: Moles & Mass Convert to moles (SHOW ALL WORK & UNITS to receive full credit.)

1. 100 g KMnO4 1.

2. 74 g Potassium Chloride 2. 2.

3. 35 g Copper (II) Sulfate 3. 3. Convert to grams (SHOW ALL WORK & UNITS to receive full credit.)

4. 1.7 moles of KMnO4 4.

5. 0.25 moles of Potassium Chloride 5. 5.

6. 3.2 moles of Copper (II) Sulfate 6. Solve the following (SHOW ALL WORK & UNITS to receive full credit.)

7. How many moles of NaCl are in 16.0 grams of NaCl? 7.

8. How many grams of glucose, C6H12O6, are there in 4.67 moles of glucose? 8.

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NOTES : REPRESENTATIVE PARTICLES MOLES

Representative Particles A representative particle is the __________________________________________________. Types of Particles (atoms / molecules / formula units)

• Monatomic elements = ____________________

• Diatomic elements = _____________________

• Ionic Compounds = ________________________

• Molecular (Covalent) Compounds = _______________________

• Ions = ______________________

• Acids = _____________________

Avogadro’s number, which is , represents the number of “chemical units” in one

mole of any substance. For the monatomic elements, the “chemical unit” is an __________________. 1 mol of

any chemical = particles.

Examples (1 mol = 6. 02 x 1 023 atoms / molecules / formula units)

• 1 mol CaCl2 =

• 1 mol Ca2+ =

• 1 mol HCl (aq) =

• 1 mol P2O5 =

• 1 mol Ca =

• 1 mol Cl2 =

Particles Moles OR Moles Particles Sample Problems

1. How many moles are in 4.50x1025 atoms of manganese?

2. How many atoms are found in 3.27mol of magnesium?

3. Chalk is composed primarily of calcium carbonate. How many particles are in 3.4 moles of calcium carbonate?

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Moles of Chalk Lab

Purpose: To determine how many moles of chalk and how many formula units of chalk are used to write a name. Materials: chalk paper balance Part A:

1. Obtain a piece of chalk, find its mass, and record in your data table. 2. Write your full name on the paper. Write large to fill up the paper and press hard enough to use up some

chalk. 3. Find the mass of your chalk again and record.

Chalk data

Mass (g)

Mass of chalk

Mass of chalk after writing

Show work and include units when answering the following questions:

1. How many grams of chalk did you use? _________________

2. a. What is the formula for chalk? (chemical name for chalk is calcium carbonate) ________________ b. What is the molar mass of calcium carbonate?

3. How many moles of chalk did you use?

4. How many formula units of chalk did you use all together?

5. How many letters are in your name? ___________

6. How many formula units of chalk per letter did you use?

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Particles Moles OR Moles Particles Homework

Now You Try! Remember to show all work and report your answers with the correct number of significant figures and the correct units!

1. How many moles are found in 9.26 x 1026molecules of CO2?

2. How many atoms are found in 12 mol of barium?

3. How many formula units are found in 3.55 mol of NaCl?

4. How many moles are in 4.27 x 1045 molecules carbon tetrachloride?

5. How many moles are present in 2.45 x 1023 molecules of CF4?

6. Iodine is an element required by humans in order to produce thyroid hormones. To help prevent iodine deficiencies in the US, our table salt is enriched with iodine. If you need 0.0219 g per year, how many moles of iodine are required per year? (Remember iodine is a diatomic!)

7. How many particles would that be?

Ha Ha…

Q: Why was the mole of oxygen molecules so excited after he left the party? A: He got Avogadro’s number!!!

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NOTES: VOLUME MOLES

In Chemistry, we use the initials, , to represent standard temperature and standard pressure.

standard temperature = standard pressure =

The molar volume of any gas at STP is L. In other words, one mole of any gas

at STP is equal to .

Conversion Factors: 1 mole = 6.02 x 1023 particles (Avogadro’s number) 1 mole = molar mass (Sum of atomic masses in grams)

1 mole = ______L of any _______ at STP (Standard temperature & pressure) Any gas density at STP = molar mass / molar volume

Moles Liters and Liters Moles Sample Problems (SHOW ALL WORK & UNITS to receive full credit.)

1. The average lung capacity of a male is 6.0 L. The average lunch capacity of a female is 4.7 L. Assume the following: If your TEACHER’s lungs are completely filled with oxygen, determine the number of moles of oxygen gas in the lungs of your chemistry teacher at STP.

2. At STP, how many moles are found in 54 L of neon gas?

3. How many liters are found in 3.02 mol of helium at STP?

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Homework: Mixed Mole Problems Now You Try! Remember to show all work and report your answers with the correct number of significant figures and the correct units! USE YOUR MOLE MAP!!

1. How many grams are in 14.05 mol silver?

2. How many moles are in 895 g of Ba(OH)2?

3. At STP, neon is a gas. How many moles are found in 0.317 L?

4. How many liters are in 44.3 moles of helium gas at STP?

Mixed Mole Solve the following problems. SHOW ALL WORK! Report your answer with the correct units and the correct number of significant figures. 1. Assuming STP, how many moles are in 5.42 x 10-1 L Ne?

2. Convert 1.3 mol of neon into particles.

3. Determine the number of grams in 3.0 mol of calcium nitride.

4. How many moles are in 45 grams carbon tetrachloride?

5. How many moles of magnesium is 3.01 x 1022 atoms of magnesium?

6. How many atoms are in 0.750 moles of zinc?

MA

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RO

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MS

V

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ME

PRO

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MS

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MULTI-STEP CONVERSIONS Remember how to use your Mole Road Map! (You must always go through the ____________!) I Moles! Sample Problems:

1. How many grams are in 45 L of neon gas at STP?

2. How many formula units are in 36.4 g of NaCl?

3. What is the mass of 550 L of helium gas at STP?

4. How many liters are in 5.6 x 1024 atoms of helium gas at STP?

Mole

Particles atoms

formula units molecules

Mass (grams) Liters at STP

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MULTI-STEP CONVERSIONS Homework

Now You Try! Remember to show all work and report your answers with the correct number of significant figures and the correct units!

1. How many liters are in 56 g of helium gas at STP?

2. A 12 oz can of Coke® contains 0.0345 g of caffeine, C8H10N4O2. Determine the number of molecules of caffeine in one can of Coke®.

3. Laughing gas is dinitrogen monoxide. How many grams of dinitrogen monoxide gas are in 45.0 L of laughing gas at STP?

4. Bromine comes from the Greek word meaning “stench.” It has a strong, irritating odor. How many atoms of bromine are found in 92.1 g of bromine?

5. How many atoms are in 65.6 L of Fe?

6. What is the volume of 5.6 g of sulfuric acid?

7. If you burned 6.10x1024 molecules of ethane (C2H6), what mass of ethane did you burn?

8. How many formula units are in 5.1 g of titanium (IV) oxide?

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PERCENT COMPOSITION

• The percent composition of a substance or mixture is the _________________________ of each

_____________________ or component.

• To find % of an element in a compound, use as the whole amount.

Example: Percent Composition of Water (H2O): ____________ oxygen ____________ hydrogen When asked to find the _____________________________, you are calculating the percentage of just _________ element in the substance or mixture. PRACTICE:

1) Find % sodium in Na2O. Molar mass = Na 22.99 x 2 = 45.98 % Na =

O 16.00 x 1 = 16.00 + 61.98 g

2) Find the % composition of a compound that has 5.34g C, 0.42g H, & 47.08g Cl. Mass (total)= 5.34g + 0.42g + 47.08g = 52.84g

%C = 5.34g C x 100% = C 52.84g

%H = 0.42g H x 100% = H 52.84g

%Cl = 47.08g Cl x 100% = Cl 52.84g

3) How many grams of sodium are in 650.0 g of Na2CO3?

Na: X = (part sodium)

C: X = (part carbon)

O: X = (part oxygen)

TOTAL: (Whole) FORMULA:

part x grams(total) = Answer grams Na = whole

4) Calculate the percent composition of C3H7OH (2-propanol).

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Percentage of Sugar in Bubble Gum Lab

Purpose: To determine the percentage of sugar in bubble gum. Materials: balance, bubble gum Note: Chewed gum cannot be placed directly on the balance pan; it must be wrapped in paper and the mass of the paper subtracted out. Also, once you are done chewing the gum, make sure you throw it away in the trash can. If I see chewed gum anywhere else, we will no longer perform food related labs for the rest of the school year. Safety: Do not chew gum that has touched the lab table. Procedure:

1. Place the wrapped pieces of gum of everyone in your group on the balance. Record the combined mass of all the wrappers and all the unchewed gum.

2. Unwrap your piece of gum put the gum in your mouth. Do not throw the wrapper away! 3. Place everyone’s empty gum wrappers on the balance. Record the combined mass of all the wrappers. Do not

throw the wrapper away! 4. Chew your gum for 15 minutes. While you are waiting for 15 minutes to pass, answer questions under

the “While you are waiting….” Section. 5. After 15 minutes, put your wrapper up to your mouth and gently spit your gum into the wrapper. Try not to

get too much saliva on the gum. 6. Place the wrapped, chewed gum of everyone in your group on the balance. Record the combined mass. 7. Throw away the gum and wrappers.

Data:

Brand of gum ______________________________ Mass (g)

Unchewed gum with wrappers

Wrappers only

Unchewed gum only

Chewed gum with wrappers

Chewed gum only

Mass of sugar

*Hint: The mass of the gum should have decreased because the sugar dissolved in your mouth… While you are waiting…..

1. The chemical formula for sucrose (commonly known as table sugar) is C12H22O11. Calculate the percentage of carbon, hydrogen, and oxygen in sucrose. Show work for credit!

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2. Nearly 99.5% of the total composition of your saliva is water. Calculate the percentage of hydrogen and

oxygen in water. Show work for credit!

3. EXTENSIVE - Properties that depend on the amount of matter present. This means that the property changes when you alter the amount of matter present. Ex: mass INTENSIVE - Properties that do not depend on the amount of the matter present. This means that the property does not change when you alter the amount present. Ex: color In the weigh boat, you have a large piece and a small piece of bubble gum. Examine the two pieces and come up with at least 3 more properties of the bubble gum that you can classify as extensive or intensive.

Extensive Intensive • Mass (mass of small piece is less than the

large piece)

• Color (both pieces are pink)

Analysis and Conclusion

1. What is the percentage of sugar in the bubble gum?

2. Convert the mass of dissolved sugar to moles.

3. How many molecules of sugar are in the dissolved sugar?

4. How would the results be affected if there was a lot of saliva on the chewed pieces of gum?

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Homework: Percent Composition

1. Calculate the percent composition of KMnO4.

2. Calculate the percent composition of Mg(NO3)2.

3. Calculate the percent composition of (NH4)3PO4.

4. Milk of magnesia, a common antacid, is a solution of magnesium hydroxide. If a sample of antacid contains 2.1 g of magnesium hydroxide, how much oxygen does the sample contain?

5. How much iron can be recovered from 25.0 g Iron (III) Oxide?

6. Which of the following compounds has the greatest mass percent of oxygen? • Nickel (III) carbonate • Lithium phosphate

7. Find the percent composition of a compound that has 2.62 g Na and 4.04 g Cl.

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NOTES: Empirical Formulas

• Empirical formula: __________________________________________________________________

________________________________________________________________________________

o e.g., CH2O or CH

o be reduced. May or may not exist in this form in the real world.

• Molecular formula: _________________________________________________________________

________________________________________________________________________________

o e.g., C2H4O2 or C6H6

o be reduced. Formula describes a substance as it actually exists. Empirical or Molecular? (Remember that for some compounds, its empirical formula can also be its molecular formula.)

• Na2O ____________________________

• C3H6 ____________________________

• K2SO4 ____________________________

• C6H12O6____________________________

Which pair has the same empirical formula?

• Na2O and Na2O2 • C6H12O6 and CH2O

• C3H6 and C5H12 • C6H6 and C5H5

Calculating an Empirical Formula

1. Determine the _______________ of each ____________________. 2. Convert the mass of each to ____________________. 3. Find the ___________ to _____________ ____________________ of each element by dividing the

number of moles of each element by the ____________________________ number of moles. 4. If the ratio is not a ______________ number, multiply each ratio by a factor to

________________________________________________. 5. Write the formula using the mole ratio as the _______________________________ for the formula.

Find empirical formula of 69.5% O & 30.5% N.

Step 1: Divide % or grams by its atomic mass to get moles of each element. 69.5g O 1 mole O = 4.344 mole O 30.5g N 1 mole N = 2.177 mole N

15.999g O 14.077g N

Step 2: Divide smallest mole number in each element to get ratio of that element. 4.344 mole O = 2.177 mole N =

* This answer becomes the subscript for that element; round to nearest whole number if .8 or higher. ANSWER = NO2

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INDEPENDENT PRACTICE:

1. Analysis of a compound shows that it contains 10.88g of calcium and 19.07g of chlorine. Determine the empirical formula of this compound.

2. One of the most commonly used white pigments in paint is a compound of titanium and oxygen that contains 5.99 g titanium by mass and 4.01 g oxygen by mass. Determine the empirical formula and name for this compound.

3. Used in the production of nylon, adipic acid is an organic compound composed of 49.31% C, 43.79% O, and the rest is hydrogen. Determine the empirical formula of adipic acid.

Calculating Molecular Formulas

• The molecular formula will have the _________________ _________________ as the empirical formula.

• To determine a molecular formula, we will _______________________ the ___________________

____________________ by a whole number factor (WNF).

WNF=

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1. Empirical Formula = P2O5

Molar Mass= 283.88 g/mol

What is the molecular formula of this compound?

2. Nitrogen and oxygen form multiple molecular compounds together. One of these compounds is used to fuel space shuttles and has the empirical formula NO2. If the molar mass of this compound is 92.02 g/ mol, what is the molecular formula?

3. Butane is commonly used in lighters. It is composed of 17.37% carbon and 82.63% hydrogen. It has a molar mass of 58.17 g/mol. What is the molecular formula of butane?

4. Vitamin C is 40.91% C, 4.587% H, and the remaining is oxygen. If the molar mass of Vitamin C is about 180 g/mol, determine the empirical and molecular formula.

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Homework: Empirical and Molecular Formulas

What is the empirical formula (lowest whole number ratio) of the compounds below? 1) 74.9% carbon, 25.1% hydrogen 2) 32.4% sodium, 22.5% sulfur, 45.1% oxygen What is the molecular formula (whole number multiple ratio) of the compounds below? 4) A compound is 64.9% carbon, 13.5% hydrogen and 21.6% oxygen. Its molecular mass is 74 g/mol. What is its

molecular formula? 5) A compound is 54.5% carbon, 9.1% hydrogen, and 36.4% oxygen. Its molecular mass is 88 g/mol. What is its

molecular formula? 6) The empirical formula of a compound is NO2. Its molecular mass is 92 g/mol. What is its molecular formula? 7) The compound methyl butanoate smells like apples. Its percent composition is 58% C, 9.8% H, and 31.4% O. If

its gram molecular mass is 102 g/mol, what is its molecular formula?

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Academic Chemistry Unit 5 Test Review: Due 11/12

Chemical Quantities & The Mole TEST on 11/13

Problems: For the following problems, show all work and report your answers with the correct units and significant figures.

1. What is the molar mass of BaCO3?

2. In one molecule of P2O5, how many phosphorous atoms are present?____________

3. List all seven diatomics:

4. Calculate the percent composition of Ca(OH)2.

5. Calculate the percent composition of diboron hexahydride.

6. How many atoms of carbon are in 3.6 moles of carbon?

7. Determine the number of moles of hydrogen that are in 362.8 g of this element.

8. A chemical reaction produces 13.8 mol of carbon monoxide gas. What volume will the gas occupy at STP?

9. What is the mass of 5.0 X 1025formula units of Fe(NO2)3?

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10. How many liters are there in 5.25 g CO2 at STP?

11. How many grams of Ne are there in 5.5 moles of the gas?

12. Nomenclature Review Name the following compounds

a. N2O ______________________

b. Li2S ______________________

c. FeO ______________________

d. (NH4)2CO3 ____________________

e. Ag2SO4 ______________________

f. Ni(OH)2______________________

Write the formulas for the following compounds.

a. dinitrogen tetroxide _______________ b. copper(I) sulfide _______________ c. calcium nitrate _______________ d. strontium bromide _______________ e. gold(II) phosphate _______________ f. pentaphosphorus decoxide __________

13. The term STP means: ___________________________________ which means:

14. Define empirical formula.

15. Define molecular formula.

16. Find the empirical formula for the following substances (a. and b.) using the information given: a. A compound is composed of 38.67% potassium, 13.85% nitrogen, 47.48% oxygen. What is the

empirical formula?

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b. A compound is made up of 44.33% phosphorus and 57.18% oxygen- what is the empirical formula?

17. Determine the molecular formula for each of these compounds. a. empirical formula = CH2O; the molar mass = 90 g/mol

b. empirical formula = HgCl; molar mass = 472.2 g/mol

18. The action of bacteria on meat and fish produces a poisonous compound called cadaverine. As its name and origin imply, it stinks! It is 58.77% C, 13.81% H, and the remainder is Nitrogen. Its molar mass is 102.2 g/mol. Find its empirical and molecular formula. (Hint: Find the empirical formula first.)

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