Unit 6: Moles and Stoichiometry

Chemistry

Chemistry

Learning Objectives Moles and Stoichiometry

Essential knowledge and skills:

Perform conversions between mass, volume, particles, and moles of a substance.
Perform stoichiometric calculations involving the following relationships:

mole-mole;

mass-mass;

mole-mass;

mass-volume;

mole-volume;

volume-volume;

mole-particle;

mass-particle; and

volume-particle.

Identify the limiting reactant (reagent) in a reaction.
Calculate percent yield of a reaction.
Calculate the % Composition
Calculate the Empirical & Molecular Formula
Calculation of Formula (Molar) Mass

Essential understandings:

Atoms and molecules are too small to count by usual means. A mole is a way of counting any type of particle (atoms, molecules, and formula units).
Avogadro’s number = 6.023 × 1023 particles per mole.
Molar mass of a substance is its average atomic mass in grams from the Periodic Table.

Molar volume = 22.41 L/mol for any gas at standard temperature and pressure (STP).
Stoichiometry involves quantitative relationships. Stoichiometric relationships are based on mole quantities in a balanced equation.
Total grams of reactant(s) = total grams of product(s).
The empirical formula shows the simplest whole-number ratio in which the atoms of the elements are present in the compound. The molecular formula shows the actual number of atoms of each element in one molecule of the substance.

Cheaper by the Mole?

Moles Explained

Think of moles as a "chemist's dozen". Just as 12 eggs is a dozen eggs, 6.023 × 1023eggs is a mole of eggs. 6.023 × 1023molecules of oxygen is a mole of oxygen. The number of grams in a mole is different from substance to substance. If you're like most students, it'sthisthat's confusing you. You can find the number of grams by looking on the periodic table. Just look at the atomic mass of the element.

Picture it this way: a dozen elephants have a different mass than a dozen rabbits- but in each case, you have a dozen animals. Similarly, a mole of oxygen gas has a different mass than a mole of water- but in each case, you have 6.023×1023molecules.

A mole is a standard scientific unit for measuring large quantities of very small entities such asatoms,molecules, or other specified particles. The mole designates an extremely large number of units, 6.02214179 × 1023, which is the number of atoms determined experimentally to be found in 12 grams ofcarbon-12. Carbon-12 was chosen to serve as the reference standard of the mole unit for the International System of Units (SI).

These entities could be atoms, molecules, formula units, ions

1 mole 12C = 6.023 x 1023 carbon atoms

1 mole H2O = 6.023 x 1023 H2O molecules

1 mole NaCl = 6.023 x 1023 formula units

1 mole of electrons = 6.023 x 1023 electrons

1 mole of Na+ ions = 6.023 x 1023 Na+ ions

1 mole of electrons = 6.023 x 1023 electrons

Molar Mass of Compounds

Molar Mass Tutorial

Molar mass is the mass in grams per mol. of a substance. Molar mass is determined by using the average atomic masses of each element. Also called molecular weight or gram formula mass.

Determine the molar mass of nitrogen (III) oxide

N2O3 N 2 x 14.01

O 3 x 16.00

76.01 g/mol

Determine the molar mass of copper (II) sulfate pentahydrate

CuSO4 ●5H2O Cu 1 x 63.55

S 1 x 32.08

O 4 x 16.00

H2O5 x 18.02

249.73 g/mol

Calculation of Molar Mass worksheet

Part 1.

Formula Molar mass (g/mol)Name of compound

1. CH4 ______g/mol______

2. CaCO3______g/mol______

3. SO2______g/mol______

4. NaClO4______g/mol______

5. NaMnO4 ______g/mol______

6. LiCl______g/mol______

7. H2SO3______g/mol______

8. FeSO4______g/mol______

9. CS2______g/mol______

10. K3PO4______g/mol______

Part 2.

Formula Molar mass (g/mol)Name of compound

11. C2H5OH______g/mol ______

12. Cu(OH)2______g/mol ______

13. Ca3(PO4)2______g/mol ______

14. Ba(NO3)2______g/mol ______

15. Al2(SO4)3______g/mol ______

Formula Molar mass (g/mol)Name of compound

16. Zn(HCO3)2______g/mol______

17. NaHCO3______g/mol______

18. C11H22O12______g/mol______

19. Mg(NO2)2______g/mol______

20. Cu(NO3)2______g/mol______

Part 3. Write the formula of each compound. Then calculate its molar mass.

Name / Formula / Molar Mass (g/mol)
Potassium carbonate
Copper (II) sulfate
Carbon tetrachloride
Sodium phosphate
Iron (III) hydroxide

Challenge Problems

Formula Molar mass (g/mol) Name of compound

21. Pb(C2H3O2)2______g/mol ______

22. (NH4)2Cr2O7______g/mol______

23. Cr2(SO4)3______g/mol______

Percent Composition

Percent Mass

The percent composition is the percentage of each element present in a compound by mass. To find the percent composition, you calculate how much of the molar mass is supplied by a particular element, divided by the molar mass, then multiplied by 100 to convert to a percentage.
Here are two examples of this type of calculation using the examples above for calculating molar masses.
The percent composition of NaCl would be as follows:

Notice that although the percentages should add up to 100%, these only add up to 99.99%, probably because of rounding.The percent composition of NH3 would be as follows:

Percent Composition Worksheet

Determine the percentage composition of each of the compounds below

1. CuBr2

Cu: ______

Br: ______

2. NaOH

Na: ______

O: ______

H: ______

3. (NH4)2S

N: ______

H: ______

S: ______

4. N2H2

N: ______

S: ______

5. KMnO4

K: ______

Mn: ______

O: ______

6. (NH4)3PO4

N: ______

H: ______

O: ______

P: ______

Composition of Hydrates

Hydrate Composition

Empricial Formula of Hydrates

A hydrate is an ionic compound with water molecules loosely bonded to its crystal structure. The water is in a specific ratio to each formula unit of the salt. For example the formula CuSO45H2O indicates that there are five water molecules for every one formula unit of CuSO4.

Hydrate – contains water moleculesAnhydrous – water has been removed

What percentage of water is found in CuSO45H2O

What percentage of water is found in Na2S9H2O

A 5.0 g sample of a hydrate of BaCl2 was heated and only 4.3 g of the anhydrous salt remained. What percentage of water was in the hydrate?

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

A 3.0 g sample of Na2CO3H2O was heated to a constant mass. How much anhydrous salt remains?

A 5.0 g sample of Cu(NO3)2xH2O is heated and 3.9 g of the anhydrous salt remains. What is the value of x?

Empirical and Molecular formulae

Empirical formula (also called the simplest formula) is the smallest whole number ratio of atoms in a compound.

Example: C6H12O6 is the molecular formula

CH2O is the empirical formula for this compound

Steps in determining the empirical formula:

Problem: Find empirical formula for a compound that contains:

18.8 % Na 29% Cl 52.2 % O

Step1: Change % to grams using the assumption that there is 100 g of compound

18.8 g Na 29 g Cl 52.2 g O

Step 2: Convert grams to moles by dividing by molar mass

*Note: Use three significant figures for moles

Step 3: Determine whole number ratios by dividing each by the smallest number of moles.

* Note: If one ratio is not a whole number or more than 0.06 from a whole number then you may have to multiply to get a whole number ratio. Example if one of your ratios was 1.5, then multiply all ratios by 2 to make a whole number.

Step 4: Write metal or most metallic element first (These are elements furthest to left on periodic table). Follow this order to the least metallic. Oxygen is always last. If carbon is present, write hydrogen directly after carbon.

NaClO4

* Note: For an ionic compound, the empirical formula is always the chemical formula

Steps for determining the chemical formula (molecular formula) from the empirical formula:

Problem: Find the molecular formula for a compound that has a molecular weight of 180.18 g/mol with an empirical formula of CH2O.

Step 1: Find molar mass of empirical formula.

CH2O has a molar mass of 30.03 g/mol

Step 2: Divide molecular weight of compound by molar mass of the empirical formula.

Step 3: Multiply all subscripts in empirical formula by this number.

1 C = 6

2 H = 12

1 O = 6

Step 4: Write correct molecular or chemical formula

C6H12O6

*Note: You always need the molar mass and empirical formula in order to determine correct chemical or molecular formula.

Empirical Formula Introduction

Write the Empirical Formula (lowest whole number ratio) for each of the following:

a. P4O6 ______

b. C6H9______

c. CH2OHCH2OH ______

d. BrCl2______

e. C6H8O6______

f. C10H22______

g. Cu2C2O4______

h. Hg2F2______

Calculating Empirical and Molecular Formula from percentages

EF and MF Calculations from percentages

Calculate the Empirical Formula for all of the Following (show your work):

a. A compound composed of: 72% iron (Fe) and 27.6% oxygen (O) by mass.

b. A compound composed of: 9.93% carbon (C), 58.6% chlorine (Cl), and 31.4% fluorine (F).(This compound is commonly known as Freon)

c. A compound composed of: 0.556g carbon (C) and 0.0933g hydrogen (H).

Empirical Formula Worksheet 1:

1.Find the empirical formula of a compound that is 48.38% carbon, 8.12% hydrogen, and 43.5% oxygen by mass.

C.I.Pigment Yellow 45 ("sideran yellow") is a pigment used in ceramics, glass, and enamel. When analyzed, a 2.164 grams sample of this substance was found to contain 0.5259 grams of Fe and 0.7345 grams of Cr. The remainder was oxygen. Calculate the empirical formula of this pigment. Answer: Fe2Cr3O12

The composition of nicotine is 74.0% C, 8.7% H, and 17.3% N. The molecular mass of nicotine is 162. What is its molecular formula? Answer: C10H14N2

One of the most deadly poisons, strychnine, has a formula weight of 334 and the composition 75.42% C, 6.63% H, 8.38% N; the rest is oxygen. Calculate the empirical and molecular formulas of strychnine, arranging the atomic symbols in alphabetical order. Answer:C21H22O2N2

Empirical and Molecular formula worksheet 2

1. Find the empirical formula for a compound which contains 0.463 g Tl (#81), 0.0544 g of carbon, 0.00685 g of hydrogen and 0.0725 g oxygen by finding its empirical formula.

2. What is the empirical formula for a compound which contains 67.1% zinc and the rest is oxygen?

3. The characteristic odor of pineapple is due to ethyl butyrate, an organic compound which contains only carbon, hydrogen and oxygen. If a sample of ethyl butyrate is known to contain 0.62069 g of carbon, 0.103448 g of hydrogen and 0.275862 g of oxygen, what is the empirical formula for ethyl butyrate?

4. 300 grams of a compound which contains only carbon, hydrogen and oxygen is analyzed and found to contain the exact same percentage of carbon as it has oxygen. The percentage of hydrogen is known to be 5.98823%. Find the empirical formula of the compound.

5. 200.00 grams of an organic compound is known to contain 83.884 grams of carbon, 10.486 grams of hydrogen, 18.640 grams of oxygen and the rest is nitrogen. What is the empirical formula of the compound?

6. A certain compound contains 4.0 g of calcium and 7.1 g of chlorine. Is relative molecular mass is 111 g/mol. Find its empirical and molecular formulas.

7. A certain compound was found to contain 54.0 g of carbon and 10.5 grams of hydrogen. Its relative molecular mass is 86.0 g/mol. Find the empirical and the molecular formulas.

8. A certain compound was found to contain 26.4 g of carbon, 4.4 grams of hydrogen and 35.2 grams of oxygen. Its relative molecular mass is 60.0 g/mol. Find the empirical and the molecular formula.

9. A certain compound was found to contain 78.2 % Boron and 21.8 % hydrogen. Its relative molecular mass is 27.7 g/mol. Find the empirical and the molecular formula.

15. A certain compound contains 7.3%Carbon, 4.5 % hydrogen, 36.4% oxygen, and 31.8% nitrogen. Its relative molecular mass is 176.0. Find its empirical and molecular formulas.

Empirical and Molecular Formula Worksheet 3

Show ALL your work for credit!

Identify the following as molecular formulas, empirical formulas or both.
Ribose, C5H10O5, a sugar molecule in RNA.
Ethyl butanoate, C6H12O2, a cmpd w/ the odor of pineapple.
Chlorophyll, C55H72MgN4O5, part of photosynthesis.
DEET, C12H17ON, an insect repellent.
Oxalic acid H2C2O4, found in spinach and tea.

Calculate the empirical formula of each compound with the following percent composition.

94.1% O, 5.9% H

79.9% C, 20.1% H

The compound meythl butanoate smells like apples. Its percent composition is 58.8% C, 9.8% H, and 31.4% O. If its gram molecular mass is 102 g/mol, what is its molecular formula?

a. A compound of carbon and hydrogen has the composition of 92.25% carbon and

7.75% hydrogen by mass. What is the empirical formula of this composition?

If the compound has a mass of 52.03 g/mol, what is the molecular formula of the compound?

The Mole Diagram

Mole-Particle Conversions

Moles to Particle Conversions

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

3.01 x 1022 atoms = 5.00 x 10-2 moles

2. How many molecules are there in 4.00 moles of glucose, C6H12O6?

3. How many moles are 1.20 x 1025 atoms of phosphorous?

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

5. How many molecules are in 0.400 moles of N2O5?

Mole-Mass Conversions

Mass to Mole Conversions

How many moles in 28 grams of CO2?

Molar mass of CO2 1 C = 1 x 12.01 g = 12.01 g

2 O = 2 x 16.00 g = 32.00 g

44.01 g/mol.

28 g CO2 = 0.64 moles CO2

What is the mass of 5.0 moles of Fe2O3?

Find the number of moles of argon in 452 g of argon.

Find the grams in 1.26 x 10-4 mol. of HC2H3O2.

Find the mass in 2.60 mol. of lithium bromide.

Mole-Volume Conversions

Moles and Molar Volume

1. Determine the volume, in liters, occupied by 0.030 moles of a gas at STP.

0.030 mol = 0.67 L

2. How many moles of argon atoms are present in 11.2 L of argon gas at STP?

3. What is the volume of 0.05 mol of neon gas at STP?

4. What is the volume of 1.2 moles of water vapor at STP?

Mixed Practice

You have 1.20 x 10-2 moles of Tantalum (Ta). How many grams is this?

You discover that the head of a match contains 1.66 g of Sulfur, S. How many atoms of S does a match contain?

While cleaning a cut, you spill a bottle of Iodine. The label says that the bottle holds 500. grams of I2. How many moles of I2 are there? How many I atoms are present?

Your silver watchband masses out at 326 g. How many moles of Ag do you have?

EXTRA STEP HERE! Can you catch it? While dropping off you recycling, you are overcome by the urge to weigh the tin cans you brought in. You find that the mass of cans in the box you brought massed out at 23.0 kg. How many moles do you have?

Water has a molar mass of 18.02 grams (that’s 18.02 grams per mole…). You drink a 2-liter bottle of water everyday, and you are such a smarty that you know that 1-ml of H2O weighs 1 g. Can you tell me how many moles of water you consume a day?

Your toothpaste probably contains around 62.0 g of fluorine per tube. How many molecules and atoms of fluorine are in one tube of toothpaste?

The head of a golf club might contain 250.0 grams of titanium. How many atoms is this?

10. The shaft of that same golf club probably contains around 35.0 moles of graphite, a natural form of carbon. What is the mass of the shaft of the club? Also how many C atoms are present?

Multiple Variable Conversions

1. How many oxygen molecules are in 3.36 L of oxygen gas at STP?

3.36 L = 9.03 x 1022 molecules

2. Find the mass in grams of 2.00 x 1023 molecules of F2.

3. Determine the volume in liters occupied by 14.0 g of nitrogen gas at STP.

4. Find the mass, in grams, of 1.00 x 1023 molecules of N2.

5. How many particles are there in 1.43 g of a molecular compound with a molecular mass of 233 g/mol?

6. Aspartame is an artificial sweetener that is 160 times sweeter than sucrose (table sugar) when dissolved in water. It is marketed by G.D. Searle as Nutra Sweet. The molecular formula of aspartame is C14H18N2O5 .

a) Calculate the gram-formula-mass of aspartame.

b) How many moles of molecules are in 10 g of aspartame?

c) What is the mass in grams of 1.56 moles of aspartame?

d) How many molecules are in 5 mg of aspartame?

e) How many atoms of nitrogen are in 1.2 grams of aspartame?

Volume-Volume Worksheet

If volumes of reactants or products are provided you can use the ratios of the coefficients from the balanced equation to work out the answer

N2 + 3H22NH3

What volume of Hydrogen is necessary to react with 5.78 L of nitrogen?

What volume of ammonia is produced?

C3H8 + 5O2 3CO2 + 4H2O

If 20.5 L of Oxygen are consumed in the above reaction, how many litres of carbon dioxide are produced?

2H2O  2H2 +O2

If 30.0 ml of hydrogen are produced in the above reaction, how many milliliters of oxygen are produced?

2CO + O2 2CO2

How many litres of carbon dioxide are produced if 75.5 L of CO2 are combusted in oxygen? How many litres of Oxygen are needed?

Mole-Mole Problems

From the balanced equation the coefficients are also used to predict the moles of reactants/products needed or produced in a chemical reaction. Use the ratios given from the balanced equation.

N2 + 3H2 2NH3

How many moles of hydrogen are needed to completely react with two moles of nitrogen?

2KClO3 2KCl + 3O2

How many moles of oxygen are produced by the decomposition of six moles of potassium chlorate?

Zn + 2HCl  ZnCl2 + H2

How many moles of hydrogen are produced from the reaction of three moles of zinc with an excess of hydrochloric acid.

Three step mole conversions

Mass to Mass Conversions

Normally you will not have nice whole number of moles in your problems but the method you use to change the number of moles will be the same. For the next set of problems you cannot just use the coefficients to change the amounts. When converting from the mass of a reactant/product to the mass of another reactant/product you need to do three steps, these are

Mass  Moles

Moles  Moles (using coefficients from the balanced equation)

Moles  Mass

Three steps will be required for anything on the mole diagram. The mole diagram below shows the steps involved to convert between the different quantities you may be given in these types of reactions.

Expanded Mole Diagram

* Use the conversion factor formed from the coefficients of A and B in the balanced equation.

The Mole and Chemical Equations

A student weighed out 2.30 g of magnesium and burned it in air. Magnesium burns in air to form magnesium oxide. The equation for the reaction is:

Mg(s) + O2(g)  MgO(s) (unbalanced)

Calculate the mass of magnesium oxide produced in the reaction.

The reaction for the decomposition of calcium carbonate is:

CaCO3(s)  CaO(s) + CO2(g)

If 100.0 g of calcium carbonate is heated, what mass of calcium oxide will form?

Calcium burns in air according to the equation:

Ca(s) + O2(g)  CaO(s) (unbalanced)

How much calcium is needed when 8.00 g of oxygen is used up?

Iron oxide is converted into iron by carbon monoxide according to the equation:

Fe2O3 + CO  Fe + CO2 (unbalanced)

Calculate the mass of iron, which could be obtained from 1.60 tonnes (1 tonne = 1000 kg) of iron oxide.

Calculate the mass of water that will react completely with 4.00 g of pure calcium metal according to the following equation:

Ca(s) + H2O(l)  Ca(OH)2(s) + H2(g) (unbalanced)

Calculate the mass of ammonia is that is required to produce 182.0 kg of urea, CO(NH2)2, according to the following equation:

CO2(g) + NH3(g)  CO(NH2)2(s)+ H2O (unbalanced)

Mixed Stoichiometry Problems

For the problems involving gases, assume that the reactions are being performed at STP

1)Given the reaction:

4NH3(g) + 5O2(g)  4NO(g) + 6H2O(l)

What is the total number of molecules of water formed when 1.20 L of ammonia reacts with excess oxygen?

2)Ethylene burns in oxygen to form carbon dioxide and water vapor:

C2H4(g) + 3O2(g) 2CO2(g) + 2H2O(g)

How many liters of carbon dioxide can be formed if 1.25 x 1024 molecules of oxygen are consumed in this reaction?

3)Calcium carbonate decomposes at high temperatures to form carbon dioxide and calcium oxide:

CaCO3(s) CO2(g) + CaO(s)