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1 The Mole Honors Chemistry Ms. Poliner Ways of measuring matter: MASS grams, pounds VOLUME liters, gallons COUNTING numbers Dozen= 12 Pair= 2 Gross= 144 Ream= 500 Mole= 6.02 x 10 23 R.P. By using dimensional analysis, we can calculate things like: 1. What is the mass of 90 average-sized apples if one dozen apples weight 2.0kg? xg= 90 apples x (1 dozen/ 12 apples) x (2.0kg/1 dozen) = 15 g 2. What is the mass of .50 bushel of apples? xg= .5 bushel x (1 dozen/ .2 bushel) x (2.0 kg/1 dozen) =50 g

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The Mole Honors Chemistry

Ms. Poliner

Ways of measuring matter: MASS grams, pounds

VOLUME liters, gallons

COUNTING numbers

Dozen= 12

Pair= 2

Gross= 144

Ream= 500

Mole= 6.02 x 1023 R.P.

By using dimensional analysis, we can calculate things like:

1. What is the mass of 90 average-sized apples if one dozen apples weight 2.0kg?

xg= 90 apples x (1 dozen/ 12 apples) x (2.0kg/1 dozen)

= 15 g

2. What is the mass of .50 bushel of apples?

xg= .5 bushel x (1 dozen/ .2 bushel) x (2.0 kg/1 dozen)

=50 g

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3. Assume that a variety of apples has 8 seeds in each apple. How many apple seeds are in 14 kg of apples?

x seeds = 14kg x (1 dozen/ 2.0 kg) x (12 apples/ 1 dozen) x (8 seeds/1 apple)

=672 seeds

THE MOLE! ( What is it?)

Avogadro’s Number: 6.02 x 1023 representative particles

Representative particle: Compound Molecules

Element Atom

Ex. How many moles of magnesium is 1.25 x 1023 atoms of magnesium?

xmol = 1.25 x 1023 atoms x (1 mol/ 6.02 x 1023)

=.208 mol

Practice Problems:

1. How many atoms are in 2.12 mol of propane (C3H8)?

x atoms = 2.12 mol x (6.02 x 1023 atoms / 1 mol)

= 1.28 x 1024 atoms

2. How many atoms are there in 1.14 mol SO3?

x atoms= 1.14 mol x (6.02 x 1023 atoms/ 1 mol)

= 6.9 x 1023 atoms

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3. How many moles are there in 4.65 x 1024 molecules of NO2?

x mol = 4.65 x 1024 molecules x (1 mol/ 6.02 x 1023 molecules)

= 7.72 mol

Molecular Mass- the mass of __1 mol___ of a compound.

= +

SO3 S 3 O atoms

Atomic Mass: _84 g/mol__ = __32 g/mol_ + __3 (16 g/mol)_

Therefore, the molecular mass is the __sum__ of atomic masses of the atoms.

Relationship between atomic mass, molecular mass, and mole:

Atomic Mass- is the mass of 1 mol of an element.

Molecular Mass- is the mass of 1 mol of a compound.

Practice Problems:

1.

Compound Formula Molecular Mass Lithium Sulfide LiS 39 g/mol

Iron Chloride FeCl2 127 g/mol

Calcium Hydroxide Ca(OH)2 74 g/mol

Ammonium Carbonate (NH4)2CO3 96 g/mol

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Oxygen Gas O2 32 g/mol

Water H2O 18 g/mol

Copper II Oxide CuO 79.5 g/mol

2. Describe the relationship between Avogadro’s number and one mole of a substance.

1 mole= 6.02 x 1023 atoms/molecules

3. How many moles is each of the following?

a. 1.50 x 1023 molecules NH3

x mol = 1.50 x 1023 molecules x (1mol/ 6.02 x 1023molecules)

= .25 mol

b. 1 billion molecules O2

x mol = 1 x 109 molecules x (1mol/ 6.02 x 1023molecules)

= 1.67 x 10-15 mol

c. 6.02 x 1022 molecules Br2

x mol = 6.02 x 1022 molecules x (1mol/ 6.02 x 1023molecules)

= .1 mol

4. How many oxygen atoms are in a representative particle of each substance?

a. Ammonium nitrate

3 atoms

b. Acetylsalichylic acid (C8H8O4)

4 atoms

c. Ozone, (O3)

3 atoms

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5. Name the compound and find its molecular mass:

Name Molecular Mass

a. K2O Potassium Oxide 55 g/mol

b. CaSO4 Calcium Sulfate 136 g/mol

c. CuI2 Copper Iodide 190.5 g/mol

d. (NH4)3PO4 Ammonium Phosphate 149 g/mol

e. Al2(SO4)3 Aluminum Sulfate 342 g/mol

f. CuSO4•5H2O Copper Sulfate Pentahydrate 249.5 g/mol

g. Zn(C2H3O2)2•2H2O Zinc Acetate Dihydrate 219.4 g/mol

h. Mg3(PO4)2 Magnesium Phosphate 262.9 g/mol

Another term that can be used is ___molar mass_ which is the mass of one of either an element, a molecular compound, or an ionic compound.

We can now convert between grams and moles!!!

Ex. How many grams are in 9.45 mol of dinitrogen trioxide?

x g = 9.45 mol x (76 g/ 1 mol)

= 718.2 g

Practice Problems:

1. Find the mass, in grams, of each.

a. 3.32 mol K 129.48 g

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b. 4.52 x 10-3 mol C20H42 1.27 g

c. 0.0112 mol K2CO3 1.55 g

2. Find the number of moles in each quantity.

a. 3.70 x 10-1 g B .034 mol

b. 27.4 g TiO2 .34 mol

c. 847 g (NH4)2CO3 8.82 mol

Practice Multiple Choice:

1. Calculate the number of atoms in 4.00 x 10-5 g sodium.

A. 3.24 x 1023 atoms

B. 5.54 x 1020 atoms

C. 3.92 x 1019 atoms

D. 1.05 x 1018 atoms

2. How many molecules of ammonia are present in 34.0 g of ammonia (NH3)?

A. 34

B. 1.20 x 1024

C. 6.02 x 1023

D. 1.77 x 1022

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3. What is the mass of 1.0 x 1012 molecules of O2?

A. 6.0 x 1011 g

B. 1.9 x 1013 g

C. 5.3 x 10-11 g

D. 2.7 x 10-11 g

4. What is the molecular mass of H2CrO4?

A. 118 g/mol

B. 52 g/mol

C. 2 g/mol

D. 64 g/mol

5. What is the mass of 5.0 x 1021 molecules of water?

A. 8.3 x 10-3 g

B. 0.15 g

C. 6.02 x 1023 g

D. 5.0 g

6. How many moles of hydrogen are present in 2.0 grams of hydrogen gas?

A. 1.0 mol

B. 2.0 mol

C. 0.5 mol

D. 4.0 mol

Volume of a Mole of Gas:

At ____standard temperature and pressure__ (STP) , temperature is 0°C and pressure is 101.3kPa (1atm), 1 mol of gas occupies a volume of ___22.4_L. Therefore at STP, 22.4L of gas has _6.02 x 1023__ representative particles of that gas.

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Ex. What is the volume in liters of 0.60 mol SO2 gas at STP?

x L = .60 mol x (22.4 L / 1 mol)

= 13.44 L

Practice Problems:

1. What is the volume at STP of the following gases?

a. 3.20 x 10-3 mol CO2 .072 L

b. .960 mol CH4 21.5 L

c. 3.70 mol N2 82.9 L

2. Assuming STP, how many moles are in these volumes?

a. 67.2 L SO2 2 mol

b. 0.880 L He .039 mol

c. 1.00 x 103 L C2H6 44.64 mol

3. Use your knowledge of dimensional analysis to determine the molar mass of a gaseous compound containing carbon and oxygen with a density of 1.964 g/L at STP.

x g/mol = 1.964 g/L x 22.4 L/mol

= 44 g/mol

4. Using your knowledge of mole calculations and unit conversions, determine how many atoms are in 1 liter of gasoline. Assume that the molecular formula of gasoline is C6H14that the density of gasoline is approximately 0.85 grams/mL.

x atoms = 1 L x (1000 mL / 1L) x (.85 g/ 1 mL) x (1 mol/ 86g) x (6.02x1023 cule/1 mol)x (20 atoms/1 cule) = 1.19 x 1026 atoms

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Refer to Notes for Mole Map

Conversion to/from Mass:

Use Molar Mass to convert to and from moles and mass.

Conversion to/from Volume:

Use 22.4 L/mol to convert between volume and moles

Conversion to/from Particles:

Use Avogadro’s number to convert between moles and representative particles.

Types of Particles: Atoms-Element Molecules-Covalent Compound

Formula Units- Ionic Compound

Practice Problems:

1. Note whether the following is an atom, a molecule, or a formula unit:

a. Na Atom

b. CO2 Molecule

c. CH4 Molecule

d. KNO3 Ionic Compound

2. Find the mass in grams of each quantity:

a. 0.720 mol Be 6.48 g

b. 2.40 mol N2 67.2 g

c. 0.160 mol H2O2 5.44 g

d. 5.08 mol Ca(NO3)2 833.12 g

3a. Calculate the number of molecules in 60.0g NO2 7.85 x 1023 molecules

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b. Calculate the volume, in liters, of 3.24 x 1022 molecules Cl2 at STP

1.21 L

c. Calculate the mass, in grams, of 18.0L CH4 at STP

12.86 g

4. Would three balloons, each containing the same number of molecules of a different gas at STP, have the same mass or the same volume? Explain.

They have the same volume but different masses. To find mass, you would end up multiplying the number of moles by the molecular mass which varies with different molecules. To convert between moles and volume, you would just multiply by 22.4 L/mol since you are at STP. Therefore, same volume, different masses.

5. Find the number of moles in each quantity:

a. 5.00 g hydrogen molecules

2.5 mol

b. 0.000264g Li2HPO4

2.4 x 10-6 mol

c. 187g Al

6.9 mol

d. 333g SnF2

2.13 mol

Percent Composition: Represents the relative amounts of each element in a compound

or is the _percent by mass__ of each element in a compound.

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This can be calculated by:

% Composition = Mass of Element/ Mass of Entire Compound

Ex. An 8.20-g piece of magnesium combines completely with 5.40 g of oxygen to form a compound. What is the percent composition of this compound?

% Composition O = 5.40 g / (5.4 g + 8.2 g) x 100= 39.7 % O

% Composition Mg = 100% - 39.7% = 60.3% Mg

Steps for Calculating Percent Composition:

Calculate the percent composition of propane (C3H8)

Step 1: Determine the atomic mass of the element.

Atomic Mass C- 36 g/mol ; H- 8 g/mol

Step 2: Determine the molecular mass of the compound.

Molecular Mass- 44 g/mol

Step 3: Plug it in! % Composition = atomic mass/ molecular mass

81.8 % is C; 18.2% is H

Step 4: Make sure to answer the question.

Ex. How much iron can be recovered from 25.0g of Fe2O3?

% Composition Fe- 112 g/mol / 160 g/mol = 70%; % O= 30%

Mass that is only iron= .70 x 25.0 g = 17.5 g

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Practice Problems:

1. Calculate the percent composition of these compounds

a. Ethane (C2H6)

% C= 80%

% H= 20%

b. Sodium bisulfate (NaHSO4)

% Na- 19.2%

%H - .83%

%S- 26.7%

%O- 53.3%

c. Ammonium chloride

%N- 66.4%

%H- 26.2%

%Cl- 7.5%

2. Calculate the percent nitrogen in these common fertilizers.

a. CO (NH2)2

23.3%

b. NH3

82.4%

c. NH4NO3

35%

3. Calculate the mass of carbon in 82.0 g of propane (C3H8)

70.3 g

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4. Calculate the mass of nitrogen in 125 g of each of the following fertilizers:

a. CO(NH2)2

29.2 g

b. NH3

102.9 g

c. NH4NO3

43.75 g

5. How many grams of oxygen can be produced from the decomposition of 100g of KClO3?

39.2g

6. How much silver can be produced from 125g Ag2S?

108.9 g

Calculating Percent H2O in hydrates:

What is a hydrate?

Is a compound that has water bound to its atoms.

First Way to Determine Percent Composition:

Ex. CuSO4•5H2O

% water =molecular mass of water/ molecular mass of hydrate; 36%

Practice Problems:

1. Find the percent water in BaCl2•2H2O.

14.7% water

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2. Find the percent water in CoCl2•6H2O.

45.4% water

Second Way to Determine Percent Composition:

Ex. 5.40 grams of BaCl2 hydrate was heated in a crucible. After heating for 10 minutes, 4.64g of anhydrous BaCl2 remains. What is the % water in the hydrate?

% Composition = mass water/mass of hydrate; 14%

Practice Problem:

1. A 2.5g sample of a hydrate of Ca(NO3)2 was heated and only 1.7g of the anhydrous salt remained. What percentage of water was in the hydrate?

32%

2. A 3.0g sample of Na2CO3•H2O is heated to constant mass. How much anhydrous salt remains?

2.56 g

Practice Multiple Choice:

1. The percent composition of aluminum in aluminum (III) hydroxide is:

A. 50%

B. 25%

C. 14%

D. None of these answers is correct.

2. Hydrates are defined as:

A. Compounds with water molecules attached to them.

B. Compounds that have had their water molecules removed.

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C. Compounds that have been heated to high temperatures.

D. None of these answers is correct.

3. A compound contains 6.0 g of carbon and 1.0 g of hydrogen. The percent composition of the compound is:

A. 17% hydrogen and 83% carbon.

B. 86% hydrogen and 14% carbon.

C. 14% hydrogen and 86% carbon.

D. 83% hydrogen and 17% carbon.

Calculating Empirical Formulas:

The empirical formula gives the __lowest___ whole-number ratio of the atoms of the elements of a compound. This gives us important information about the kinds and number of atoms in molecules. This MAY or MAY NOT BE the same as the

__molecular mass__.

Ex. What is the empirical formula of a compound that is 25.9% nitrogen and 74.1% oxygen?

1. Convert to grams 25.9 g N and 74.1 g O 2. Convert to moles 25.9 g x (1mol/14 g) = 1.85 mol; 74.1g x (1mol/16g)= 4.6

mol 3. Plug it in: N1.85O4.6 4. Divide by the smallest number: NO2.5 5. Make sure you have no decimals!!!! Here we will multiply by 2: N2O5

Practice Problems:

1. Calculate the empirical formula in each compound.

a. 94.1% O, 5.9% H

HO

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b. 79.8% C, 20.2% H

CH3

c. 67.6% Hg, 10.8%S, 21.6% O

HgSO4

d. 27.59% C, 1.15% H, 16.09% N, 55.17% O

C2HNO3

2. 1,6-diaminodexane is used to make nylon. What is the empirical formula of this compound if it is 62.1% C, 13.8% H, and 24.1% N?

C3H8N

Calculating Molecular Formulas:

The molecular formula of a compound is either the same as its experimentally determined

__empirical formula__, or it is a simple whole-number multiple of it.

Steps to Calculate Molecular Formula:

1. Calculate the empirical formula mass (the molar mass of the empirical formula). 2. Calculate the molecular formula mass. 3. Divide the molecular formula mass by the empirical formula mass

____number of empirical formula units___ and is the multiplier to convert the empirical formula to the molecular formula.

Ex. Calculate the molecular formula of the compound whose molar mass is 60.0g and empirical formula is CH4N.

1. Calculate the empirical formula mass- 30 g 2. Find the molecular formula mass- 60 g 3. Divide the molecular formula mass by the empirical formula mass- 2 4. Multiply by the multiplier you calculated in #3. C3H12N3

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Practice Problems:

1. Find the molecular formula of each compound given its empirical formula and molar mass:

a. Ethylene glycol (CH3O), used in antifreeze, molar mass=62g/mol

C2H6O2

b. p-dichlorobenzene (C3H2Cl), mothballs, molar mass= 147g/mol

C6H4Cl2

2. Which pair of molecules has the same empirical formula?

a. C2H4O2, C6H12O6 or b. NaCrO4, Na2Cr2O7

a

3. Determine the empirical formula of a molecule with:

a. 40.5% Zinc, 19.9% Sulfur, 39.6% Oxygen.

ZnSO4

b. 28.2% Potassium, 25.6% Chlorine, 46.2% Oxygen.

KClO

c. 79.9% Copper, 20.1% Oxygen

CO

Practice Multiple Choice:

1. If the empirical formula of a compound is CH2, what is a possible molecular formula for the compound?

A. CH2

B. C2H6

C. C4H8

D. More than one could be a molecular formula for CH2.

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2. If the empirical formula of a compound is P2O3, what could be a possible molar mass of the compound?

A. 55 g/mol

B. 165 g/mol

C. 275 g/mol

D. None of these could be a molar mass for P2O3.

3. A compound is 40.0% carbon, 53.3% oxygen, and 6.66% hydrogen. What is its empirical formula?

A. C4O5H7

B. CO2H3

C. COH2

D. None of these is the correct empirical formula.

4. If the compound in problem 3 has a molecular weight equal to 60 g/mol, what is the molecular formula?

A. C3OH6

B. CO2H16

C. C2O2H4

D. None of these is the correct molecular formula.

Finding the Formula of a Hydrate:

Ex. 5 grams of BaCl2 hydrate was heated. 4.26 grams remain. Find the formula of the hydrate.

First determine the mass of water 5g-4.26g = .74 g. Next calculate the percent composition: %H2O= 14.8%; %BaCl2 = 85.2%. Convert to grams. Then to moles:

.82 mol H2O and .41 mol BaCl2 BaCl2•H2O

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Practice Problems:

1. A 5 gram sample of a hydrate of BaCl2 was heated and only 4.3 g of the anhydrous salt remained. What is the formula for this hydrate?

Eliminate

2. A hydrate is found to have the following percent composition: 48.8% MgSO4 and 51.2% H2O. What is the formula for this hydrate?

MgSO4•7H2O

3. 2.5 grams of CuSO4 hydrate was heated. After heating, the mass of the anhydrous crustal was 1.59g. Find the formula of the hydrate.

CuSO5•5H2O

Review: 1. How many molecules is each of the following?

a. 15.5g SiO2

1.56 x 1023 molecules

b. 14.4 mol F2

8.67 x 10E23 molecules

c. 0.780 mol Ca(CN)2

4.69 E23 molecules

d. 7.00 mol H2O2

4.214 E23 molecules

e. 5.60 mol NaOH

3.37E24 molecules

f. 3.21 x 10-2 mol Ni 1.93E22 molecules

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2. Calculate the volume of each of the following gases at STP.

a. 7.6 mol Ar

170.24 L

b. 0.44 mol C2H6

9.9 L

c. 1.20 mol O2

26.9 L

3. Find each of the following quantities:

a. the volume, in liters, of 835g SO3 at STP

233.8 L

b. the mass, in grams, of a molecule of aspirin (C9H8O4) 3x 10-22 g

4. Calculate the percent composition of each compound.

a. H2S 5.44 % H; 94.1% S

b. (NH4)2C2O4

22.6% N; 6.5% H; 19.4% C; 51.6% O

c. Mg(OH)2

41.7% Mg; 54.8% O; 3.4% H

d. Na3PO4

42.1% Na; 18.9% P; 39% O

5. You find that 7.36g of a compound has decomposed to give 6.93g of oxygen. The only other element in the compound is hydrogen. If the molar mass of the compound is 34.0 g/mol what is the molecular formula? Empirical Formula is – HO; Molecular Formula is H2O2

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6. Classify each formula as an empirical or molecular formula:

a. S2Cl2 Molecular

b. C6H10O4 Molecular

c. Na2SO3 Empirical

d. C5H10O5 Molecular

e. C17H19NO3 Empirical

f. (NH4)2CO3 Empirical

7. Determine the molecular formula for each compound:

a. 94.1% O and 5.9% H; molar mass=34g

H2O2

b. 40.0% C, 6.6% H, and 53.4% O; molar mass=120g

C2H4O2

8. Calculate the mass in grams for the following amounts:

a. 2.5moles of H2

5 g

b. 0.50 moles Cl2

35.5 g

c. 1.25 moles of Ar

50 g

d. 2.5 moles of Na2CO3

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9. For each compound listed below: 1. Write the formula; 2. Find the molecular mass; c. Convert 0.25 mol into grams; d. Find the percent composition of the first element

a. Ammonium nitrate

a. NH4NO3

b. 80 g/mol

c. 20 g

d. 17.5% N

b. Aluminum Nitrate

a. Al(NO3)3

b. 213 g/mol

c. 17.25 g

d. 32.4% Al

c. Barium Hydroxide

a. Ba(OH)2

b. 171 g/mol

c. 42.75 g

d. 80.1% Ba

d. Potassium phosphate

a. K3PO4

b. 212 g/mol

c. 53 g

d. 55.1% K

10. Calculate the empirical formula of the compound that contains 1.0 g S for each 1.5 g O. S2O3

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11. Calculate the empirical formula of the compound containing 75.0% C and 25.0% H. CH4 12. Calculate the empirical formula of the compound containing 81.8% C and 18.2% H. C3H8 Multiple Choice Practice: 1. What is the empirical formula of the compound whose molecular formula is P4O10?

a. PO b. PO2 c. P2O5 d. P8O20

2. What is the percent by mass of oxygen in magnesium oxide, MgO?

a. 20% b. 40% c. 50% d. 60%

3. A compound is 86% carbon and 14% hydrogen by mass. What is the empirical formula for this compound?

a. CH b. CH2 c. CH3 d. CH4

4. What is the mass in grams of 3.0 x 1023 molecules of CO2?

a. 22 g b. 44 g c. 66 g d. 88 g

5. What is the percent by mass of water in the hydrate Na2CO3 · 10H2O (formula mass = 286)?

a. 6.89% b. 14.5% c. 26.1% d. 62.9%

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6. How many moles of CO2 are present in 220 mg? a. 5 moles b. 0.005 mole c. 5000 moles d. 10 moles 7, What is the percent water in hydrated calcium chloride (CaCl2·2H2O)?

a. 66.67% b. 32.47% c. 24.51% d. 12.26%

8. One mole of (NH4)2HPO4 contains ? moles of hydrogen atoms. a. 1 b. 5 c. 6 d. 9