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Honors Chemistry

Unit 8 (Chapter 9)

Stoichiometry

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Cooking with Chemistry Stoichiometry is a way of figuring out how much stuff you're going to make in a chemical reaction, or how much stuff you'll need to make a chemical reaction do what you want. Calculating a stoichiometry problem is very similar to following a recipe. In both, the proportions of the reactants and products, or the ingredients, much remain the same. Before you can begin a stoichiometry problem, you need the recipe—a balanced chemical reaction (BCE). Vocabulary Review Reactant: a substance that is consumed in the reaction. In a chemical equation, the formulas on the

_______________ side of the arrow are the reactants.

Product: a substance that is produced in the reaction. In a chemical equation, the formulas on the

________________ side of the arrow are products.

Molar Mass: the mass in grams of one mole of a substance. This information can be obtained from the

_____________________________ _____________________________.

Stoichiometry Terms Coefficient: the number in front of a formula in a BCE. It indicates the relative number of moles of each reactant and product in the reaction. Write the equation for the reaction of iron with oxygen to produce rust, iron III oxide. What is the coefficient of iron?________; of oxygen ________; of rust ________ Mole Ratio: the ratio used to compare the number of moles of two substances in the same chemical reaction. The numbers that fit into this ratio are the coefficients from the BCE. For the above reaction, write all possible mole ratios.

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Cooking with Chemistry ~ Follow the Boxes

Given the problem you’re trying to solve, figure out what box in the diagram is your starting point and which box is the ending point. Then, use the factor label method to convert stepwise from the starting box through each intermediate step until the designation box is reached.

NEVER SKIP A BOX!!

1. EXAMPLE: mole mole

1. How many moles of oxygen gas are produced when 3.0 moles of water decompose?

BCE: __________________________________________________________________________________

box ______ box _______

2. How many moles of water must decompose to produce 20 moles of hydrogen?

BCE: __________________________________________________________________________________

box ______ box _______

1 2 3 4

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2. EXAMPLE: mole gram

1. How many moles of hydrogen gas are needed to react with 20. grams of nitrogen to produce

ammonia?

BCE: __________________________________________________________________________________

box ______ box _______

2. How many grams of ammonia can be produced from 5.00 moles of hydrogen gas?

BCE: __________________________________________________________________________________

box ______ box _______

3. EXAMPLE: gram gram 1. How many grams of magnesium are required to react with silver nitrate to produce 23.5 grams of

silver?

BCE: __________________________________________________________________________________ box ______ box _______

2. How many grams of silver nitrate are required to react completely with 25 grams of magnesium?

BCE: __________________________________________________________________________________ box ______ box _______

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Reaction Stoichiometry #1

In order to compare reactants and products in a chemical reaction, one must know the “ingredient ratio,” which in a balanced chemical equation is the mole ratio derived from the coefficients. Follow these basic steps to perform a stoichiometric calculation based upon a chemical reaction.

1. Write a balanced chemical equation. 2. If necessary, change the given to the basic chemical unit—the mole. 3. Using the mole ratio, change moles of the given substance to moles of the new substance. 4. Change the new substance into the desired unit.

1. How many moles of hydrogen gas will react with 15.1 g of chlorine gas to produce hydrogen

chloride?

2. Calcium hydride can be used as a portable hydrogen generator. When water is added to calcium hydride, hydrogen gas and a base are formed. How much gas (in grams) is produced from 100. grams of the hydride?

3. How many moles of calcium oxide are produced from the decomposition of 10.0 grams of

calcium carbonate?

4. How many grams of silver can be recovered from the reaction of 13.6 moles of silver nitrate

with copper?

5. How many grams of carbon dioxide are produced from the combustion of 1500 g of propane, C3H8?

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6. Methanol (CH3OH) is made by reacting carbon monoxide and hydrogen in the presence of a catalyst. How much methanol can be produced from 40.0 grams of the poisonous gas?

7. One problem with silver jewelry is that it tarnished when the silver reacts with the sulfur found

in the air. How much tarnish is produced from 0.05 g of sulfur?

8. Ethyl alcohol (C2H5OH) is made by fermenting dextrose (C6H12O6), the sugar found in grapes.

Carbon dioxide is also produced in the fermentation process. How may grams of alcohol can be made from 100 pounds of grapes, which are approximately 30% dextrose by weight?

9. Baking soda (sodium hydrogen carbonate) is often used as an antacid. It reacts with the excess

hydrochloric acid secreted by the stomach. Milk of Magnesia, an aqueous suspension of magnesium hydroxide, is also used as an antacid to neutralize any excess stomach acid. Which is the more effective antacid, 1.00 g of baking soda or 1.00 g of milk of magnesia? (HINT: Write two balanced equations, one for each antacid reacting with hydrochloric acid. Then find which remedy consumes more acid).

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Reaction Stoichiometry #2 Directions: Solve the following problems on a SEPARATE sheet of paper. Use the factor label method and show all your work. 1. The reusable booster rockets of the U.S. space shuttle employ a mixture of aluminum and ammonium

perchlorate for fuel. The unbalanced equation is: Al + NH4ClO4 Al2O3 + AlCl3 + NO + H2O

For each kilogram of aluminum, what mass of NH4ClO4 should be used in the fuel mixture? 2. One of relatively few reactions that takes place directly between two solids at room temperature is

Ba(OH)2 · 8H2O + NH4SCN Ba(SCN)2 + H2O + NH3 a. Balance the reaction. b. Name the first reactant. c. What mass of ammonium thiocyanate must be used if it is to react completely with 6.5 g

of Ba(OH)2 · 8H2O? d. What is the total mass of the three products that are produced in the reaction described in

part C? 3. Elixirs such as Alka-Seltzer use the reaction of sodium bicarbonate with citric acid (H3C6H5O7) in aqueous

solution to produce fizz. a. Write the balanced equation for this double replacement reaction. b. What mass of citric acid should be used for every 100.0 mg of sodium bicarbonate? c. What mass of carbon dioxide is produced from the mixture described in part b? d. If the volume of 1 mole of a gas is approximately 22 L, what volume of gas is produced from the

reaction described in part b?

4. Aspirin (C9H8O4)is synthesized by reacting salicylic acid (C7H6O3) with acetic anhydride (C4H6O3). The unbalanced reaction is as follows: C7H6O3 + C4H6O3 C9H8O4 + H2O

a. What mass of acetic anhydride will react completely with 5.00 x102 g of salicylic acid? b. What mass of aspirin will be produced in the reaction described in part a?

5. Write the balanced chemical reaction for the complete combustion of isooctane (C8H18). Assuming gasoline is 100% octane with a density of 0.692 g/mL, determine the mass of carbon dioxide produced by the combustion of 1.2 x 1010 gallons of gasoline (the approximate annual consumption of gasoline in the United States in 1984).

6. Hydrazine (N2H4) and hydrogen peroxide have been used as rocket propellants. When they react, they

produce nitric acid and water. a. Write the balanced equation for the reaction. (This equation is a little difficult to balance.

Balance the nitrogen first; then place coefficients in front of the hydrogen peroxide and water to balance the hydrogen and oxygen).

b. How many grams of hydrogen peroxide are needed to react with 120 mL of hydrazine (D = 1.01 g/mL) which is 97% pure?

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7. An unknown metal forms a solid oxide with the formula XO3. 3.31 grams of the oxide reacts with hydrogen gas to form the free metal and 1.24 grams of water. What is the unknown metal?

8. A sample of sodium azide, NaN3, a compound used in automobile air bage, thermally decomposes to produce sodium metal and nitrogen gas.

a. Write the equation for the reaction. b. An air bag becomes indlated to a volume of 368 L inside a car. If 1 mole of any gas occupies 22 L

at room conditions, how many moles of nitrogen must be produced to inflate the air bag to 368 L?

c. How many grams of sodium azide must decompose to produce the amount of nitrogen derived in part (b)?

9. One of the many problems encountered on the Apollo 13 mission was a build-up of carbon dioxide in the command module, Odyssey. The built-in system failed, so Mission Control had to devise a plan for CO2’s removal using items on board the spacecraft. To remove carbon dioxide from the air in the spacecraft, scientists react the gas with lithium hydroxide. The products of the reaction are lithium carbonate and water. If a person exhales about one kilogram of carbon dioxide a day, how many grams of lithium hydroxide are required to remove the gas formed in a six-day lunar expedition with three astronauts on board?

10. In emergency situations, oxygen gas can be produced in an oxygen mask by the reaction between potassium superoxide (KO2), carbon dioxide, and water. Besides the oxygen gas, the other product is potassium bicarbonate.

a. Write a balanced equation for the production of oxygen gas. b. If a person wearing such a mask exhales 0.02 g of carbon dioxide every minute, how many

grams of oxygen are produced in one hour?

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The Reactant’s the Limit! (but which one??) In carrying out a reaction, chemists rarely use exact stoichiometric mixtures of

reactants. They usually use an excess of one reactant to ensure that the reaction

goes to completion, and that the more expensive reactant is entirely consumed.

STEPS FOR COMPLETING A LIMITING REAGENT PROBLEM:

1. Convert grams of the two reactants to grams of the SAME product

a. smaller answer = MAXIMUM amount of product that can be produced

b. LR = reactant that produces the smallest amount

c. ER = the other reactant

2. Convert grams of LR grams of ER to determine amount ER you NEED

3. Excess = HAVE (original amount given) – NEED (calculated in #3)

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1. Because of its many uses, ammonia is one of the most highly-produced inorganic chemicals. Eighty percent of the ammonia which is produced commercially by combining hydrogen and nitrogen gases is used for fertilizers. How much ammonia can be produced from a mixture of 1.00 x 106 g of nitrogen and 5.00 x 106 g hydrogen? How much excess reactant is left over?

BALANCED EQUATION: STEP 1:

a. maximum amount of product = _____________________

b. LR = _______________________

c. ER = _______________________

STEP 2: STEP 3:

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2. Sulfuric acid, the most widely manufactured chemical in the world, is produced in a three-step process. The final step of this process is the reaction between sulfur trioxide and water. How much sulfuric acid is produced from the reaction 5.2 x 106 grams of sulfur trioxide with 2.6 x 106 grams of water? How much excess reagent is left over?

BALANCED EQUATION: STEP 1:

d. maximum amount of product = _____________________

e. LR = _______________________

f. ER = _______________________

STEP 2: STEP 3:

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3. In the following equations, determine which reactant is the limiting reagent and which reactant is in excess. How much of the excess reactant is left over? Check to make sure each equation is balanced!

a. KOH + HNO3 --> KNO3 + H2O 16.0 g 12.0 g

b. NaOH + H2SO4 --> Na2SO4 + H2O 10.0 g 10.0 g

4. Welding torches can reach temperature near 2000 0C. The reaction involves the combustion of

acetylene, C2H2. a. Starting with 175 grams of both reactants, what is the limiting reagent? b. How much carbon dioxide can be produced from this reaction? c. How much of the excess reactant is left over?

5. When a certain nonmetal whose formula is X8 burns in air, XO3 forms. Write a balanced equation for

this reaction. If 120.0 g of oxygen gas and 80.0 g X8 are consumed completely, what is element X?

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Name: ___________________________________________________________________ Period _______

Limiting Reagent Practice

1. You drop a 4.00 g hunk of zinc metal into 6.00 g of HCl in solution. How much hydrogen gas do you expect to produce from this reaction? What is the limiting reagent? What is the excess reagent? How much excess do you expect to have left over?

g of hydrogen gas: _____________ LR: _____________ ER: _____________ Left over: _____________ 2. In an extremely exothermic reaction, powdered aluminum metal will react with powdered iodine to

form aluminum iodide. If you react 5.00 g of aluminum with 20.0 g of iodine, how much aluminum iodide to you expect to produce from this reaction? What is the limiting reagent? What is the excess reagent? How much excess do you expect to have left over?

g of aluminum iodide:______________ LR: _____________ ER: ____________ Left over: _____________

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3. Being the pyro that she is, Miss Uhernik loves detonating pumpkins. To accomplish her explosion, Miss Uhernik adds 10.0 g of hydrogen gas and 5.00 g of oxygen gas to a balloon placed inside the gourd. How much water should be produced from this mixture? What is the limiting reagent? What is the excess reagent? How much excess do you expect to have left over?

g of water: ______________ LR: ______________ ER: ______________ Left over: ______________ 4. After overdoing it at wing night, you are experiencing some intestinal distress. To calm your stomach,

you take Maalox to neutralize the excess acid in your tummy. If you take 2.00 g of Maalox (magnesium hydroxide) to neutralize the 4.00 g of excess acid (HCl) in your stomach, how much magnesium chloride do you expect to produce? What is the limiting reagent? What is the excess reagent? How much excess do you expect to have left over?

g of Mg(OH)2: ______________ LR: ______________ ER: ______________ Left over: ______________

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Percent Yield or “How much did you get”

All of the quantities we have been calculating so far represent the maximum yield of product according to the balanced equation. Many reactions, however, fail to give a 100% yield of product. The main reasons for this failure are as follows: 1. Reactions do not always go to completion or may be reversible. 2. Impure reactants and competing side reactions may cause other products to form. 3. Some product may be lost in handling and transferring from one reaction vessel to another. The theoretical yield of a reaction is the calculated amount of product that should be obtained. In order to determine this amount, we must have a balanced chemical equation and do a stoichiometry problem. The actual yield of a reaction is the amount of product that actually forms when the reaction is carried out in the laboratory. An actual yield is an experimental value. The percent yield is the ratio of the actual yield to the theoretical yield multiplied by 100. Both the theoretical and the actual yields must have the same units to obtain a percent. Calculating percent yield measures the efficiency of the reaction in changing the reactants to products. % Yield = actual yield x 100 theoretical yield Example 1: 1.8 g of aluminum is reacted with an excess of copper II sulfate. What is the percent yield of the reaction if 3.74 g of copper is actually produces? Example 2: Carbon disulfide can be made from sulfur dioxide and coke (coke is nothing more than carbon, C). Carbon dioxide gas is also produced. If the percent yield of carbon disulfide is 86%, how much will be produced from 950. g of coke?

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Percent Yield or “How much did you get?”

1. When 4.0 moles of carbon tetrachloride react with an excess of hydrogen fluoride, 3.0 moles of CCl2F2 (freon) is obtained. The other product is hydrogen chloride gas.

State which of the following statements are true about the reaction and make the false statement true.

a. The theoretical yield of CCl2F2 is 3.0 mol.

b. The theoretical yield for hydrogen chloride is 146 g.

c. The percent yield for the reaction is 75%.

d. The theoretical yield cannot be determined unless the exact amount of hydrogen fluoride gas is

known.

e. From just the information given above, it is impossible to calculate how much hydrogen fluoride

is unreacted.

f. For this reaction, as well as for any other reaction, the moles of the reactants must equal the

moles of the product.

g. Half a mole of hydrogen fluoride is consumed for every mole of carbon tetrachloride used?

h. Theoretically, at the end of the reaction, no carbon tetrachloride is left unreacted.

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2. Silicon nitride is a ceramic. It is made by reacting silicon and nitrogen at high temperatures. How much nitrogen must react with sufficient silicon to prepare 2.00 x 102 kg of the ceramic if the reaction yield is 72%?

3. An astronaut excretes (the technical term for pees) about 2500 grams of water a day. If lithium oxide is used in the spaceship to absorb this water, how many kilograms of lithium oxide must be carried for a 30-day space trip for three astronauts? Assume that the reaction yield is 95%.

4. A student calculated the theoretical yield of barium sulfate in a precipitation experiment to be 1.352 g. When she filtered, dried, and weighed her precipitate, however, her yield was only 1.279g. Calculate the student’s percent yield.

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5. Astronauts also exhale an average of 900 g of carbon dioxide a day. This exhaled gas can be removed by its reaction with lithium hydroxide. The unbalanced chemical reaction is

LiOH + CO2 --> Li2CO3 + H2O

a. How many grams of LiOH is needed per astronaut per day? b. Suppose a lithium hydroxide canister contains 1.00 kg of the base. Will the canister take care of

the carbon dioxide exhaled by one astronaut in a day? c. If so, what percentage remains unused?

6. Gastric juice contains about 3.0 g hydrochloric acid per liter. If a person produces about 2.5 L of gastric

juice per day, how many antacid tablets, each containing 400. mg of aluminum hydroxide, are needed to neutralize all the HCl produced all day?

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RXN Stoichiometry Prequiz #2 Name ____________________________ DIRECTIONS: SHOW ALL YOUR WORK for the following problems; report all answers to 3 significant figures. 1. Write the balanced equation for the reaction between aluminum and phosphoric acid.

a. How many grams of hydrogen are produced from a mixture of 15.0 grams of aluminum and 24.0

grams of phosphoric acid?

b. How much excess reagent is left over? 2. Write the balanced equation for the single replacement reaction between sodium and iron III oxide.

a. If 6.50 grams of iron are actually produced from 0.078 moles of iron III oxide, what is the percent

yield of the reaction?

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Name ____________________________________________________ Period ___________________

STOICHIOMETRY - Vocabulary Review Match the correct vocabulary term to each numbered statement. Write the letter of the correct term on the line. Each answer can only be used once.

a. mole i. theoretical yield b. stoichiometry j. limiting reagent c. mass-mass calculation k. mole ratio d. reactants L. actual yield e. excess reagent m. percent yield f. atoms n. molar mass g. coefficient o. anion h. cation

___________ 1. the starting materials in a chemical reaction

___________ 2. a conversion factor derived from the coefficients of a balanced chemical

equation interpreted in terms of moles

___________ 3. the maximum amount of product that could be formed in a reaction

___________ 4. the amount of a substance that contains 6.02 x 1023 representative particles

of that substance

___________ 5. the substance completely used up in a chemical reaction

___________ 6. the ratio of how much product is produced compared to how much is

expected, expressed as a percentage

___________ 7. the calculations of quantities in a chemical reaction

___________ 8. the actual amount of product in a chemical reaction

___________ 9. the substance left over after a reaction takes place

___________10. a stoichiometric computation in which the mass of products is determined

from the given mass of reactants

___________11. mass and _______ are always conserved in a chemical reaction

___________12. the mass in grams of one mole of a substance

___________13. indicates the relative number of moles of each reactant and product in a

chemical reaction

___________14. a negative ion

___________15. a positive ion

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Reaction Stoichiometry Review DIRECTIONS: Show all of your work. Remember to use proper sig figs and units! 1. Natural waters often contain relatively high levels of calcium ion and bicarbonate ion from the leaching

of minerals into the water. When such water is used commercially or in the home, heating of the water leads to the formation of solid calcium carbonate, which forms a deposit (“scale”) on the interior of boilers, pipes, and other plumbing fixtures. The reaction involved the decomposition of calcium bicarbonate into calcium carbonate, carbon dioxide, and water. If a sample of well water contains 2.0 x 103 mg of calcium bicarbonate per milliliter, what mass of calcium carbonate scale would 120. gallons of this water be capable of depositing?

2. When a particular metal X reacts with hydrochloric acid, the resulting products are XCl2 and H2. When 79.1 g of the metal reacts completely, 2.42 g of hydrogen gas results. Identify element X.

3. If 25.0 g of aluminum and 100. g of bromine are reacted, 64.2 g of aluminum bromide product is recovered. What is the percent yield of the reaction? How much of the excess reactant is left over?

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4. The white limestone cliffs of Dover, England, contain a large percentage of calcium carbonate. A sample of limestone weighing 84.4 g is reacted with an excess of hydrochloric acid. The mass of calcium chloride formed is 81.8 g. What is the percent of calcium carbonate in the limestone?

5. A car gets 9.2 km to a liter of gasoline. Assuming that gasoline is 100% octane, C8H18 (D = 0.69 g/mL), how many liters of air will be required to burn the gasoline for a 1250 km trip? Assume complete combustion. Are is 20% oxygen and has a volume of 22 liters per mole.

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Reaction Stoichiometry Pretest 1. The amount of product obtainable under perfect conditions is called the ________________________. 2. When Paul O’Neum weighs his final product and records his result, he has just recorded his

__________________ _________________. 3. In any chemical reaction the quantities that are conserved are ____________________ and

______________________. 4. In a chemical reaction the reactant that is completely consumed is called the ___________________

__________________ . 5. Tess Tube produces 25 lb of Smiley cookies, but according to the ingredients, Tess should be able to

produce 30 pounds. What is her % yield? Questions 6-9 refer to the following equation. Sodium hypochlorite, the active ingredient in bleach, is produced by the following reaction:

2 NaOH + Cl2 -----> NaCl + NaClO + H2O 6. How many moles of sodium hypochlorite (NaClO) can be produced from 10.2 moles of sodium

hydroxide?

7. How many grams of chlorine gas are needed to produce 14.8 moles of sodium hypochlorite? 8. How many grams of base are needed to produce 900. grams of table salt? 9. In the three problems, 6-8, what step is present in every one of them? ____________ _____________ 10. In order for Professor Brinclhof to measure the efficiency of a reaction, he would calculate his

________________ _________________. 11. If 16 grams of hydrogen gas react completely with 128 grams of oxygen gas, how many grams of water

will be produced? _____________. 12. When sodium bicarbonate is mixed with hydrochloric acid, the gas produced is tested with a flaming

splint. What is the result of the test? ________________________________________________________ 13. Name the seven diatomic elements. _________________________________________________________

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14. What is the molar mass of oxygen gas? ___________________________ 15. What is the formula for ammonia? _____________________ Hydrogen peroxide? ________________ 16. The formula for phosphorous acid is _______________. The formula for nitrous acid is ______________. 17. The products of the decomposition of Na3N are ___________ and ___________. 18. Lake Erie is the fourth largest of the Great Lakes, with a surface area of 25 700 km2. How many sig. figs.

are in this measurement? _______ 19. Describe the relationship between moles of reactants and moles of products in a chemical reaction.

___________________________________________________________________________________________

_______________________________________________________________________________________

20. The coefficients in a balanced chemical equation describe the number of _________________ or number

of _____________________ of reactants and products. 21. Hydrofluoric acid cannot be stored in glass bottles because the acid reacts with silica in glass to produce

hexafluorosilicic acid: SiO2 + 6 HF -----> H2SiF6

+ 2 H2O If 30.0 grams of silicon dioxide and 40.0 grams of hydrofluoric acid react,

a. determine how much hexafluorosilicic acid is produced under ideal conditions. b. determine how much excess reactant is left over. c. determine the actual yield of hexafluorosilicic acid. if the reaction’s efficiency is 95.4%.

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Reaction Stoichiometry ACROSS 1. Chemist’s “Bible,” Periodic

________

3. A product of a reaction

between plumbic nitrate

and zinc

5. A product resulting from

the decomposition of a

base

6. A unit used to measure a

quantity conserved in a

chemical reaction

8. Number of moles of H2 in

excess when 2 moles of H2

react with one

mole of O2 to form water

10. A product formed by the

addition of water to a

nonmetallic oxide

12. The step necessary in all

Reaction Stoichiometry

problems

15. Mercury I ion

16. A metallic bond is formed

between

positive ions and a ____ of

electrons

17. Study of chemical reactions

21. Formula for a possible

product of the reaction

between copper and

chlorine gas.

22. SO4-2

: sulfuric acid::

TeO4-2

: _______ acid

25. negative ions

27. Symbol for an inner

transition element in the

lanthanide series

28. Formula for the soluble

ionic salt formed from the

double replacement

reaction between aqueous

solutions of barium

chloride and potassium

sulftate

29. Place this noble gas in a

box (hint) and it may glow

bright orange; also an old

boxy Dodge

30. Least reactive metal

31. “Heart of Chemistry”,

abbreviation

33. Number of moles of

oxygen needed to react

with 184 grams of sodium

to form sodium oxide

34. The ____ reactant in a

reaction regulates the

______, or amount of

product obtained (2 words)

37. SI unit of mass

39. Old name for inner

transition elements: rare

______ elements

41. Acronym for light

amplification by stimulated

emission of radiation

42. Suffix that indicates one

less oxygen than “-ate”

43. Number of moles of Mg

with which nine moles of

Cl2 will react to form MgCl2

44. Nonmetals ______

electrons to form ions

45. Number of grams of

hydrogen gas formed from

the reaction between 80

grams of calcium and

excess hydrochloric acid

46. Solid, liquid, and gas: the

three _____ of matter

47. An alkali metal

48. Chem is ______

49. The symbol for the

feminine version of helium?

DOWN 1. The amount of product

that should be obtained

2. A metallic hydroxide

4. Reactant that is left over

7. Building block of matter

8. Metal used to coat nails

9. Rubidium bromide

11. SnF and NaF in toothpaste

help prevent _________

13. Substance found on the left

side of an equation

14. Poisonous

18. Amount of matter

19. The amount of product

obtained

20. Bitter: base:: _____:acid

23. Excess reactant

24. An example of evidence

that a chemical change has

occurred

26. Particle of a covalent

compound

31. Used to convert grams to

moles

32. Common name for NaOH

35. Nomenclature deals with

this branch of chemistry

36. Mass/Volume =

37. A sharp cutting instrument

made from carbide alloys

38. Each step of a reaction

stoichiometry problem

40. A use for silver bromide

44. A product formed from the

decomposition of

germanium chloride

46. Symbol for the element to

the left of yttrium

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Determination of the mole ratio of reactants and products

DISCUSSION: An important application of the mole concept is expressing mole relationships between substances in a chemical reaction. According to the Law of Conservation of Matter, the total mass of the products in a chemical reaction must be the same as the total mass of the reactants. Also, the number of atoms in an ordinary chemical reaction must be conserved. Because atoms rearrange themselves into different combinations during a reaction, the total number of moles or molecules of products does not have to equal the total number of moles or molecules of reactants. Moles are not conserved. In this experiment, you will react sodium bicarbonate with hydrochloric acid. The products are sodium chloride, carbon dioxide, and water. The experimental determination of the relative masses of sodium bicarbonate and sodium chloride will enable you to determine the relative number of moles of reactants and products. Using this ratio you can write a balanced chemical equation for the reaction. OBJECTIVES:

to verify experimentally mass relationships and mole relationships between reactants and products of a chemical reaction

write balanced chemical equations and label a reaction by its type

MATERIALS: Bunsen burner evaporating dish wire gauze 10 mL graduated cylinder ring stand balance sodium bicarbonate hydrochloric acid, 3 M PROCEDURE: 1. Clean and dry and evaporating dish. 2. Place the evaporating dish on en electronic

balance and measure its mass. Record. _____________

3. Add one level spoonful of sodium bicarbonate to the evaporating dish. Measure the mass of the sodium bicarbonate and evaporating dish. Record. ___________

4. Slowly add 10 mL of hydrochloric acid to the sodium bicarbonate in the evaporating dish.

5. Then carefully add more acid, drop by drop, until the bubbling stops.

6. Using a ring stand apparatus and wire gauze, gently heat the evaporating dish with a small flame (about one inch below the wire gauze) until a dry solid remains.

7. Cool the evaporating dish and contents. Measure their mass and record. __________

8. Reheat the evaporating dish and contents with the gentle flame for five minutes.

9. Cool the evaporating dish and contents. Measure their mass and record. __________

10. Allow the evaporating dish to cool and then clean with water.

CONCLUSION: On a separate sheet of paper, answer the following questions. Be organized, neat, and label all parts. Pay attention to significant figures. 1. Complete the calculations on the following

page. 2. On the back of the calculations page,

construct a data table of all the measurements taken in the lab. Use a ruler and remember units and significant figures.

3. Under the data table, construct a calculations table to include your calculated answers. Use a ruler and remember units and significant figures.

Pop, Pop Fizz, Fizz

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Name______________________________________________________________ Period _________________ 1. Calculate the mass of sodium bicarbonate used in the reaction. 2. Calculate the moles of sodium bicarbonate used in the reaction. 3. Calculate the mass of the sodium chloride produced in the reaction. 4. Calculate the moles of sodium chloride produced in the reaction. 5. Calculate the experimental mole ratio of sodium bicarbonate to sodium chloride. Record the answer as

a decimal value. In other words, divide the denominator into the numerator. (2.35 moles sodium bicarbonate/1.62 moles sodium chloride = 1.45)

6. Write the balanced chemical equation for the reaction. 7. Using the balanced chemical equation, determine the actual mole ratio of sodium bicarbonate to

sodium chloride. 8. Calculate the percentage error for your experimental mole ratio: | actual - experimental | x 100 actual

9. Why is it important to heat the contents of the evaporating dish twice? ____________________________

___________________________________________________________________________________________

10. Explain the importance of using a gentle flame to heat the contents of the evaporating dish. __________

___________________________________________________________________________________________

____________________________________________________________________________________

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Determination of the percent yield of a reaction

DISCUSSION: In this experiment, a known amount of silver nitrate and copper metal will undergo a single replacement reaction. The two products will be separated, collected and weighed. Keep in mind that you will be graded on how well you recover the products, so be sure to follow the procedure carefully. OBJECTIVES: 1. To determine the actual mass of product. 2. To determine the theoretical mass of product. 3. To calculate the % yield of a reaction and %

error. MATERIALS: copper wire silver nitrate distilled water large test tube ring stand iron ring funnel filter paper 400 mL beaker balance PROCEDURE: DAY 1 1. Cut a piece of copper wire approximately 25

cm long. 2. Coil the lower part of the wire, leaving enough

straight wire to stick out of a large test tube. 3. Mass the copper wire to the nearest 0.01 g and

record the mass. __________________________ Set aside.

4. Clean a large test tube and rinse it with distilled water. Stand the clean test tube in a small beaker and place it on the balance; tare the balance.

5. Carefully spoon solid silver nitrate into the test tube until the mass is approximately 2.0 grams. (a little more or a little less if ok!) Record the mass. _________________ (CAUTION: Silver nitrate will react with your skin or clothing, turning it black. If you should get any on yourself, immediately wash it off.)

6. Fill the test tube about half way with distilled water. Swirl until all the silver nitrate has dissolved. (NOTE: If the water becomes cloudy, the test tube was not clean and you must start over.)

7. Put the coiled copper wire into the test tube. Make sure you have as much wire as possible in the solution.

8. Add distilled water until the water level is almost to the top.

9. Write your name on a piece of masking tape and place it near the mouth of the test tube. Place the test tube in the test tube rack designated for your class.

DAY 2 1. Using a pencil, place your name on a piece of

filter paper. Then weigh the paper to the nearest 0.01 g and record. ________________________ Make sure that you use the same balance as you used yesterday.

2. Set up the filtering apparatus; fold the filter paper and place it in the funnel. Moisten the paper with a few drops of distilled water to hold it in place.

3. Remove all the silver from the copper wire. a. Shake the silver crystals off the copper

wire into the test tube. Remove the wire.

b. Use your wash bottle to rinse onto the filter paper any remaining crystals adhering to the copper.

c. A few remaining crystals may need to be scraped from the wire with a stirring rod, then rinsed with more distilled water.

4. Set the copper wire on a clean paper towel and let it dry.

5. Pour the contents of the test tube into the funnel. Collect the blue solution (the filtrate) in the beaker.

6. Use distilled water to rinse any silver that was left behind in the test tube. Pour this rinse water onto the filter paper and let it drain into the beaker. Continue this process until all the silver has been transferred.

7. Wash the filter paper several times with distilled water.

8. Remove the filter paper from the funnel. Place it on a tray on the center table to dry overnight.

9. Pour the filtrate into the sink; clean the beaker. 10. Weigh the copper wire and record.

___________________ Discard the copper wire into the trash can.

DAY 3 1. Determine the mass of the silver and filter

paper. Record. __________________________ 2. Place the silver in the jar located on the central

distribution table and throw the filter paper into the garbage can.

“Au Gee.” It’s Only Silver

34

“Au Gee, It’s Only Silver” Prelab

Earl N. Meyer dissolves 6.00 grams of silver nitrate in distilled water. He then adds 6.63 grams of

pure copper to the test tube. After the reaction is completed, he separates the leftover copper

from the silver and finds that 5.45 grams of copper is left over. After the silver dried overnight,

he weighed the silver and filter paper and found the mass to be 4.53 grams. The filter paper

weighs 1.02 grams.

1. Write the balanced chemical equation for the above reaction.

_________________________________________________________________________

2. What is the excess reactant? _____________________________________

3. What is the limiting reagent? _____________________________________

Data Table

Mass of copper before the reaction

Mass of silver nitrate used

Mass of filter paper

Mass of copper after the reaction

Mass of filter paper and silver

35

“Au Gee, It’s Only Silver” Prelab 1. Calculate the mass of silver that was actually produced in Earl’s experiment. _______________________ 2. Calculate the mass of silver that should be (theoretical yield) produced in the experiment. _______________________ 3. Calculate the percent yield of his reaction. % yield = actual x 100

theoretical

_______________________ 4. Calculate the mass of copper that was actually consumed in his experiment. _______________________ 5. Calculate the mass of copper that should have (theoretical yield) reacted. _______________________ 6. Calculate the percent error in Earl’s amount of copper. % error = |theoretical – actual| x 100

theoretical

_______________________

36

CONCLUSION: On a separate sheet of paper, answer the following questions. Be organized, neat, and label all parts. Pay attention to significant figures. 1. Construct a data table of all measurements

taken in the lab. Use the Prelab as a guide. Remember to use a ruler and units.

2. Calculate the following: a. the mass of silver nitrate actually produced b. the theoretical yield of silver c. the % yield of silver d. the mass of copper actually consumed e. the theoretical mass of copper consumed

f. % error of copper 3. Construct a Calculations table to include the

answers from a - f above. Remember to use a ruler and units.

4. Write a paragraph explaining the main source of error and how they affect your results. Remember, human error is not an excuse for poor results.

37

Recycling old aluminum into alum crystals

INTRODUCTION:

Recycling old aluminum is mandatory in many cities and towns in the United States. Several states require a deposit on canned and bottled soft drinks to promote the recycling effort. Most states offer a certain price per pound on scrap aluminum. In this experiment, we will contribute to the recycling effort by using scrap aluminum as a starting material for the synthesis of alum, an inorganic compound. To produce the alum, the scrap aluminum will first react with potassium hydroxide according to the

following equation: 2Al (s) + 2KOH (aq) + 6H2O ---> 2K[Al(OH)4] (aq) +3H2 (g) Then the aqueous K[Al(OH)4], potassium tetrahydroxoaluminate III, will be added to sulfuric acid to produce the alum:

K[Al(OH)4] (aq) + 2H2SO4 (aq) + 8H2O –> KAl(SO4)2•12 H2O (s)

Alum is the common name for potassium aluminum sulfate dodecahydrate, KAl(SO4)2 • 12 H2O. The ionic hydrate contains 12 water molecules, six surrounding each of the two metallic cations. Alum has a number of useful applications. Most grocery stores have the substance on the shelves as a pickling agent to help pickles retain their crispness. Pool owners keep it on hand to clear up cloudy water in swimming pools. The textile industry uses it as a

bonding agent in dyes, and people use it every time the apply antiperspirant to block the flow of sweat glands.

OBJECTIVES:

1. To produce alum crystals from scrap aluminum. 2. To calculate the theoretical yield and actual yield

of a chemical reaction. MATERIALS: aluminum pie pan 1.4 M potassium hydroxide 9.0 M sulfuric acid scissors 2 250-mL beakers hot plate Buchner funnel Buchner flask filter paper graduated cylinders stirring rod 600 or 800 mL beaker

PROCEDURE:

DAY 1 1. Cut approximately 1 gram of aluminum from an

aluminum pie pan. Weigh it and record the mass. ________________

2. Cut the aluminum into tiny pieces, making sure you do not lose any of it since you have already weighed the limiting reactant of the reaction.

3. Place the aluminum pieces in a clean 250-mL beaker; add

4. 50 mL of 1.4 M KOH. Bubbles of hydrogen gas should begin to form.

5. Heat the beaker gently on a hot plate (800C) to

speed up the reaction. (DO NOT use a Bunsen burner--the hydrogen gas that is generated is flammable!)

6. Heat until all the aluminum has been consumed. (No more bubbles!)

7. When you are finished heating the solution, the volume should be reduced to around 30 mL. If the volume falls below this level while heating, add distilled water to replenish it.

8. Set up a filtering apparatus; place a piece of filter paper in a Buchner funnel. Moisten the filter paper with a few drops of distilled water to hold it in place. Use rubber tubing to connect a Buchner flask to the aspirator. Place a 600 or 800

mL beaker in the sink directly below the aspirator and slowly fill it with water. This will help cushion the force of the water and prevent excessive splashing. Turn on the water full blast to create a vacuum in the flask.

9. Then and only then, filter your hot solution by slowly pouring it onto the filter paper. The filtrate should be clear, with any dark residue left on the filter paper. If it is not clear, filter again.

10. Rinse the beaker twice with 5 mL portions of distilled water. Pour each of the rinses through

the filter residue. 11. Transfer the clear filtrate into a clean 250-mL

beaker. Rinse the filtering flask with 5 mL of distilled water; pour the rinse into the beaker.

12. Discard the filter paper into the trash can. 13. When the filtrate is cool, slowly and carefully,

while stirring, add 20 mL of 9.0 H2SO4 to the solution.

14. As you are adding the sulfuric acid, a white precipitate should form. However, when all the acid has been added, the precipitate should

redissolve. 15. Warm the solution gently on a hot plate to insure

that all the precipitate has redissolved. The volume should be around 50 mL. If it is more than 50 mL keep heating until the volume is reduced (turn the hot plate up all the way to boil off all of the unwanted liquid).

16. STOP!! Write your name on the beaker and allow it to sit overnight in the hood to precipitate the alum crystals.

Junk Bonds

38

DAY 2 1. Obtain your beaker and decant (pour out) the

liquid into the sink. Be careful not to lose any of your crystals.

2. Pour your crystals onto a paper towel and pat dry.

3. Place a small beaker on the balance you used on day 1 and tare. Slowly and carefully transfer the crystals from the paper towel into the beaker. Record the mass of the crystals. ____________________

PRELAB QUESTIONS: 1. Calculate the molar mass of CuSO4 • 5 H2O.

Follow the same procedure when calculating the molar mass of alum.

2. Add these two equations together to eliminate D, a product in the first equation and a reactant in the second.

A + 2B + 3C ---> 2D + E

D + 2F + 3C ---> 4G ∙ 9H2O

Just like in algebra, in order to eliminate D, the second equation must be doubled.

CONCLUSION: Answer the following on a separate sheet of paper. 1. Use a ruler to construct a data table that contains

the measurements taken in the lab.

2. Write the overall balanced equation for the synthesis of alum. This can be obtained by adding together the two equations found in the introduction. Use the prelab as a guide.

3. Using the overall equation determined in #2, calculate the theoretical yield of alum. Use the prelab as a guide when determining the molar mass of alum.

4. Calculate the percent yield of alum. 5. Use a ruler to construct a calculations table.