Unit 2- StoichiometryAP Chemistry

Basic Types of Reactions

SynthesisType General FormatFormation of a binary compound A + B ABMetal oxide + Nonmetal oxide MO + (NM)O SaltsMetal oxide in water MO + H2O Base Nonmetal oxide in water (NM)O + H2O Acid

1. acids (formulas begin with H- except for some organic acids like acetic acid which are often written with the H- at the end, as in CH3COOH)

2. bases (formulas end with –OH except for ammonia NH3 and organic bases which are similar to ammonia and contain hydrogen)

3. metal oxides (binary compounds of a metal and oxygen)4. non-metal oxides (binary compounds of a non-metal and oxygen)5. salts (compounds of metals and nonmetals that are NEITHER bases NOR oxides. Format is

metal,nonmetal,oxygen)

Practice. Don’t forget diatomic elements!

1. Hydrogen gas is burned in air. (Note that AIR means oxygen)

2. Solid magnesium is heated in nitrogen gas.

3. A piece of solid zinc is heated in chlorine gas.

4. A piece of solid sodium is placed in hydrogen gas.

5. Lithium oxide reacts with carbon dioxide.

6. Calcium oxide reacts with sulfur dioxide.

7. Calcium oxide is added to excess water

8. Sulfur trioxide gas is added to excess water

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DecompositionTypes General FormatBinary Compounds AB A + B

Metallic Carbonates MCO3 MO + CO2

Metallic Hydrogen Carbonates MHCO3 MO + H2O + CO2

Metallic Hydroxides MOH MO + H2O

Metallic Chlorates MClO3 MCl + O2

Oxyacids Acid H2O + (NM)O

Salts with Oxygen Salts MO + (NM)O

1. Calcium Hydroxide is decomposed

2. Carbonic Acid is heated

3. Calcium Carbonate is heated

4. Aluminum Chlorate is heated

5. Calcium Sulfate is heated.

6. Sodium oxide decomposes.

7. Sodium hydrogen carbonate decomposes.

8. Potassium nitride is heated.

9. A solution of hydrogen peroxide is heated*** Memorize this one***

10. Sulfurous acid is heated.

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Combustion

General Format:

Alkanes: Organic Alcohols:

1. Ethane is burned in air.

2. Methanol is burned in oxygen

3. Propane is burned in air

4. Butanol is burned in oxygen.

Other “combustion reactions” you may see include sulfides:

General:

1. Carbon disulfide is burned in excess oxygen.

2. Zinc sulfide is heated in an excess of oxygen

Synthesis, Decomposition, and Combustion Practice:

1. ethanol is burned in oxygen

2. aluminum chloride is heated

3. beryllium reacts with oxygen

4. butane combusts

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5. lithium hydrogen carbonate is heated

6. propane is burned in oxygen

Review- Percent Composition

Formula: Mass of part ÷ Mass of whole x 100

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

_________________2. How much silver can be produced from 125 g of AgS?

________________

Review Mole Conversions

Write the conversion factors on each line:

Mole

Mass (g) Particles (atoms, molecules)Volume (L)

Do the following conversions:

1. 0.5 moles of C6H12O6 to grams

2. 35 g of CuSO4 • 5 H2O to moles

3. How many grams are there in 1.5 x 1025 molecules of CO2?

4. A sample of NH3 gas occupies 75.0 liters at STP. How many grams is this?

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5. How many atoms are there in 1.3 x 1022 molecules of NO2?

6. How many atoms of oxygen are in a 5.0 gram sample of O2 at STP?

Review- Empirical and Molecular Formulas

Empirical Formula-

Molecular Formula-

Ex. C6H12O6 CH2O Molecular Empirical

Example 1: A compound with elements C, H, and O is found to have 9.1% hydrogen and 54.5% carbon. What is the empirical formula?

Steps: 9.1 g H x (1 mole/ 1 g) = 9.1 mol / 2.275 = 4 H1. Assume 100 g 54.5 g C x (1 mol/ 12 g) = 4.5417 mol C / 2.275 = 2 C2. Convert to moles 36.4 g O x (1 mol/ 16 g) = 2.275 mol O /2.275 = 1 O 3. Divide by smallest # moles C2H4O (put in alphabetical order)

Example 2: A compound with an empirical formula of C2H4O has a molecular mass of 176 g/mol. What is the molecular formula?

Step 1- complete chart for what you know:Molecular Formula Molecular Mass Empirical Formula Empirical Mass X Factor

176 C2H4OStep 2- Find the empirical mass. Divide the molecular mass by empirical mass to find X Factor

Molecular Formula Molecular Mass Empirical Formula Empirical Mass X Factor176 C2H4O 44 4

Step 3- Distribute the x factor across the empirical: C8H16O4

*Tips:- Carry decimal places out at least 4 digits- You can only round to a whole # if it is within 1 tenth of that #. (Ex. 1.03 1)- When you are not within 1 tenth you will need to multiply all of your answers so they are whole numbers. (Ex. 1.5 x 2 = 3; 1.2 x 5 = 6)

Example: A compound is found to be 64.9 % carbon, 13.5% hydrogen, and 21.6% oxygen. Its molecular mass is 74 g/mol. What is its molecular formula?

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Practice: Percentage Composition and Empirical Formulas*You cannot round to the nearest whole # if your ratio is not 0.1 away**

1. Phencyclidine (“angel dust”) is C17H25N. A sample suspected of being this illicit drug was found to have a percentage composition of 83.71% C, 10.42% H, and 5.61% N. Do these data acceptably match the theoretical data within 5% for phencyclidine?

2. Iso-octane is a standard for gasoline octane rating. A sample weighing 0.5992 g was found to contain 0.5040 g of carbon and 0.09515 g of hydrogen. Its molecular mass is 114. Write the empirical and molecular formulas of isooctane (arranging the atomic symbols in the order CH).

3. Adenosene triphosphate, ATP, is an important substance in all living cells. A sample with a mass of 0.8138 g was analyzed was analyzed and found to contain 0.1927 g of carbon, 0.02590 g of hydrogen, 0.1124 g of nitrogen, and 0.1491 g of phosphorus. The remainder was oxygen. It’s molecular mass is 507. Calculate the empirical and molecular formulas of ATP. (Arrange the symbols in the formulas in their alphabetical order.)

4. One compound of mercury with a molecular mass of 519 contains 77.26% Hg, 9.25% C, and 1.17% H (with the balance being O). Calculate the empirical and molecular formulas (arranging the symbols in the order HgCHO).

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Empirical Formula 1You are given a 0.1843 gram sample of a compound. This compound has 3 elements in it: G, X, and Z. The atomic mass of element G is 56 grams/mole. The atomic mass of element X is 72 grams/mole. It is found that 0.0248 g of element G are in the sample. In this same sample it is found that there are 2.667 x 1020 atoms of element X. Also in this same sample are 0.001329 moles of element Z.

a. What is the empirical formula of this compound?

b. What is the atomic mass of element Z?

Empirical Formula 2 A metal fluoride with the formula XF2 is 48.7% F. Find the identity and molecular mass of X.

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Empirical Formula 3

Two elements “X” and “Y” exist. These two elements form a compound that has the formula XY4. In this compound 1.00g of X is found for every 5.07 g of Y.

It is also noted that the element X will form a compound with oxygen with the formula XO2. In this oxide, 1.00 g of X is found for every 1.14 g of oxygen.

From this information, find the atomic masses of element X and element Y.

Empirical Formula 4

A 12.58 gram sample of ZrBr4 was dissolved, and after several chemical steps, all of the combined bromine was precipitated as AgBr solid. The silver content of the solid AgBr was found to be 13.2 grams. Assume the atomic mass of silver to be 108 and the atomic mass of bromine to be 80.0. What value would be obtained for the atomic mass of Zr?

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Empirical Formula by Combustion Analysis

Combustion reactions:

Example 1:A 0.255 g sample of compound with C, H, and O is burned resulting in 0.561 g CO2 and 0.306 g H2O being collected. What is the empirical formula?

Practice- 2005 AP Exam2. Answer the following questions about a pure compound that contains only carbon, hydrogen, and oxygen.

A 0.7579 g sample of the compound burns in O2(g) to produce 1.9061 g of CO2(g) and 0.3370 g of H2O.

a. Calculate the individual masses of C, H, and O in the 0.7579 g sample.b. Determine the empirical formula

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2006 AP Exam3. Answer the following questions that relate to the analysis of chemical compounds.

A compound containing the elements C, H, N, and O is analyzed. When a 1.2359 g sample is burned in excess of oxygen, 2.241 g of CO2 is formed. The combustion analysis also showed that the sample contained 0.0638 g of H.

a. Determine the mass, in grams, of C in the 1.2359 g sample of the compound.b. When the compound is analyzed for N content only, the mass percent of N is found to be

28.84%. Determine the mass, in grams, of N in the original 1.2359 g sample.c. Determine the mass, in grams, of oxygen in the original 1.2359 g sample of the

compound.d. Determine the empirical formula of the compound.

Review Extended Mole MapMole: Mole Calculations

Fill in the conversion factors on the lines.

g A g B

L A Mole A Mole B L B

Particles A Particles B

Example: Given the formula: N2 + 3 H2 2 NH3, how many grams of NH3 will be produced if 67.2 L of N2 reacts with excess H2? Steps: L N2 moles N2 moles NH3 g NH3 (Ignore “excess”)

67.2 L N2 x 1 mole N2 x 2 mole NH3 x 17 g NH3

22.4 L N2 1 mole N2 1 mole NH3

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Practice:

1. 2KClO3 2KCl + 3O2

How many grams of potassium chloride are produced if 25 g of potassium chlorate decompose?

_______________2. N2 + 3H2 2NH3

What volume of NH3 at STP is produced if 25.0 g of N2 is reacted with an excess of H2?

_______________4. 2AlCl3 2Al + 3Cl2

If 10.0 g of aluminum chloride are decomposed, how many molecules of Cl2 are produced?

_______________5. If 20.0 g of aluminum is produced, how many atoms of chlorine in AlCl3 reacted? Use the equation from problem 4.

_______________

Limiting Reactant

Limiting Reactant:

Excess Reactant:

Example 1:2 Al + 3 Cl2 2 AlCl3

A mixture of 1.5 mol Al and 3 mol Cl2 are allowed to react. Which one is the limiting reactant?

How to tell you have a L.R. problem:

Solve:

Example 2:

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A 2 g strip of zinc is placed in a solution containing 2.5 g silver nitrate causing this reaction to occur:

Zn + 2 AgNO3 2 Ag + Zn(NO3)2

How many grams of silver will form?

Solve:

**ALWAYS use the amount of the LIMITING REACTANT from the problem when solving for the product!!!***

Practice Limiting Reactant

1. N2 + 3 H2 2 NH3

a. How many grams of NH3 can be produced from the reaction of 28 g of N2 and 25 g of H2?

b. How much of the excess reagent is left over?

2. Mg + 2 HCl MgCl2 + H2

a. What volume of hydrogen at STP is produced from the reaction of 50.0 g of Mg and the equivalent of 75 g of HCl?

b. How much of the excess reagent is left over?

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3. 3 AgNO3 + Na3PO4 Ag3PO4 + 3 NaNO3

a. Silver nitrate and sodium phosphate are reacted in equal amounts of 200 g each. How many grams of silver phosphate are produced?

b. How much of the excess reactant is left ?

More PracticeGiven the balanced equation for questions 1-4:

4NH3(g) + 5O2(g) 4NO(g) + 6H2O(g)1. Given six (6) moles of each reactant at the beginning of a reaction, which one will be the limiting reactant

(used up first) when the reaction finishes?

a. How much (if any) of some reactant is left over?b. How many moles of NO are formed?

2. If 35.0 grams of oxygen gas and 19.8 grams of ammonia gas reacted according to the above equation, what would be the limiting reagent? How many grams of the excess reactant is left over?

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3. If 47.2 liters of oxygen gas and 39.5 liters of ammonia gas at STP react, determine:

a. The limiting reagentb. The amount of the excess reactant left over

4. If 3.91 moles of oxygen gas and 125 liters of ammonia gas at STP react, determine the following:

a. The limiting reagentb. The amount of water formed in the reactionc. The amount of the excess reactant left over

Theoretical and Percentage Yield

Theoretical Yield-

Actual Yield-

Percent Yield-

Example 1:

Fe2O3 + 3 CO 2 Fe + 3 CO2

A. You determined the limiting reactant of this experiment to be Fe2O3. If you start with 150 g of Fe2O3, what is the theoretical yield of Fe?

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B. In the lab, you produce 87.9 grams of iron. Calculate the percent yield.

Practice Percent YieldRewrite the formula for percent yield here:

1. A student performs the following experiment: CaCO3(s) CaO(s) + CO2(g)

The student uses 270 g of CaCO3 and obtains 120 g of CaO. Calculate the percent yield.

2. 2Al(s) + Fe2O3(s) Al2O3(s) + 2Fe

A. A student uses 125 g of iron oxide and obtains 80 g of iron. Calculate the percent yield.

B. A student uses 40 g of aluminum and obtains 75 g of iron, what is the percent yield?

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