unit 10 notes: reaction...

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UNIT 10 NOTES: REACTION STOICHIOMETRY STUDENT OBJECTIVES: Your fascinating teachers would like you amazing learners to be able to…

1. That you begin to see some real world applications of chemistry in manufacturing industries! 2. That you can “mole road” to make a variety of chemical conversions using:

a. Molar Mass b. Molar Volume c. Density d. Avogadro’s Number

3. That you can apply these concepts to the following common real-life situations: When the reaction is not perfect – it is incomplete, there are losses due to other reactions, or losses due to the steps in the process. In this case, the YIELD is not 100%.

4. In the real world we often have to deal with VERY EXPENSIVE reactants. We want to make sure we maximize the reaction to use every bit of the expensive stuff. The expensive stuff will be the LIMITING REACTANT and the other reactant(s) will be in excess – there will be some left over at the end.

5. Above all, we want to continue to help you develop the critical thinking skills necessary to be successful in college and life!


Leggett PreAP Chemistry Stoichiometry 10-1 (13:23) http://youtu.be/4sJNrDlvVOw http://vimeo.com/35932190

I. WHAT DO THE COEFFICIENTS IN A REACTION TELL US??!?!

Consider the equation: 2 H 2 (g) + O 2 (g) 2 H2 O (l)

The coefficients indicate: “Two molecules of diatomic hydrogen react with one molecule of diatomic oxygen to produce two molecules of water.”

But, since 1 MOLE = (6.022 x 1023) particles, the mole is a “count” as well!

So, the coefficients also indicate: “2 moles of hydrogen gas react with 1 mole of oxygen gas to produce 2 moles of water.” We can use this interpretation to help us make relationships to help us in reaction stoichiometry!!!

II. USING THE LAW OF CONSERVATION OF MASS

Example 10-1. What does the law of conservation of mass state?

Notice in the picture above, you can see how the total mass of hydrogen and the total mass of oxygen total the mass of water produced. We can use this concept to make some basic calculations with reactions!

Example 10-2. Think about the following reaction: 2 NaN3 2 Na + 3 N2. If 500 grams of NaN3 completely decompose to form 323.20 grams of N2, how much Na is produced?

Example 10-3. Something to think about: The Law of Conservation of Mass states that mass is neither created nor destroyed in an ordinary chemical reaction. When an iron nail rusts, it seems to get heavier in mass. Does the iron nail follow the Law of Conservation of Mass?

Look at those AMU

values! Sure look

like molar masses

to me!

http://youtu.be/4sJNrDlvVOw

http://vimeo.com/35932190


III. WHY USE REACTION STOICHIOMETRY?

When you have the amounts of all the substances except one in a reaction, it’s pretty easy to find the mass of the unknown. But what if you have several substances that you do not know the amount of??!?! Then, you must complete some stoichiometry calculations to figure it out! Before we get started, there are some key notes about reaction stoichiometry.

KEY NOTES: Unless stated otherwise, stoichiometry calculations always assume a reaction proceeds 100% to the

product without any side reactions or loss of product. Stoichiometry calculations can go between any substances in a reaction… reactants or products! Stoichiometry helps us to find a ___________________________ yield – a

maximum amount of product that SHOULD be made. (Not the actual yield!)

Steps to a Reaction Stoichiometry Problem: Get your balanced reaction

written Map out a game plan! Calculate!

IV. TYPES OF REACTION STOICHIOMETRY PROBLEMS

A. MOLE-MOLE problems – the most basic type!

The coefficients provide us with the mole ratio needed to convert from one substance to another within a chemical reaction. Think of the mole ratio as a “bridge” between one substance and another!

THE MOLE RATIO IS ALWAYS A RATIO OF THE COEFFICIENTS. Use your dimensional analysis rules to decide what substance should go on top and what should go on bottom of the ratio!

Example 10-4. Use the balanced equation for the synthesis of water (given on the previous page) and

determine how many moles of water would form if you began with 32.38 moles of oxygen gas.

B. MASS-MASS and MASS-PARTICLE problems

A mass of one of the products or reactants is given and you are asked to find the mass of another product or reactant.

Remember, our _______________________________is the “bridge” we use between one substance and another. That means that somewhere in our calculation, we need to use the mole ratio!

Hmm… wonder how much

product I was supposed

to make? Stoichiometry

may come in handy!


It is very helpful on these problems to come up with a game plan, a “map” if you will, to help you solve the problem. We will practice with coming up with a “map” as we work examples!

The “map” will involve the conversions we used last unit. It basically provides a pathway for us to get to the mole ratio bridge – if we can get to moles, then we can set up a mole ratio between substances. Our conversions from last unit are still going to be very important!

SAMPLE MAP: How many grams of H2O could be formed by reacting 50.0 grams of H2 with excess O2?

2 H2 + O2 2 H2O 50.0 g ? g Molar Mass Molar Mass Mole Mole Mole Ratio “Bridge”

SAMPLE SETUP: Here’s what the calculation should look like for the above problem…

OHmol1

OHg02.18

Hmol2

OHmol2

Hg02.2

Hmol1Hgrams0.50

2

2

2

2

2

22 = 446 grams H2O

OKAY, IT’S TIME FOR US TO SOLVE A PROBLEM!!! REMEMBER – GET A BALANCED REACTION WRITTEN, MAP OUT A GAME PLAN, THEN CALCULATE! Leggett PreAP Chemistry Stoichiometry 10-2 (14:35) http://youtu.be/xBsgKNtYw-E http://vimeo.com/35932299

Example 10-5. How many grams of sodium chloride could theoretically be formed in a synthesis reaction in which 30.00 grams of chlorine gas are reacted with excess solid sodium?

Example 10-6. Potassium chlorate decomposes with heat into potassium chloride and oxygen gas. If a

manufacturer must produce 400.0 kilograms of potassium chloride per month to stay in business, how many grams of potassium chlorate must she decompose (assuming 100% yield)?

Excess means we have way

more than what we need,

so it won’t make a

difference in the math.

STEP 1: According

to the chart above,

we calculate from

grams to moles using

molar mass.

STEP 3: According

to the chart above,

we calculate from

grams to moles using

molar mass.

STEP 2: In order to get from one substance to another, we

need the mole ratio. The mole ratio uses the coefficients. In

this case, the ratio is 2 moles of H2 for every 2 moles of H2O.

STEP 1 STEP 2 STEP 3

http://youtu.be/xBsgKNtYw-E



Example 10-7. How many grams of bromine are required to react completely with 45.70 grams of lithium

iodide?

Leggett PreAP Chemistry Stoichiometry 10-3 (14:44) http://youtu.be/MO1AqC4udjo http://vimeo.com/35932968

Example 10-8. How many molecules of hydrogen gas are required to react with 0.334 grams of nitrogen gas in

the synthesis of ammonia (NH3)?

Example 10-9. How many moles of water could potentially be produced by reacting excess sulfuric acid with 120. grams of sodium hydroxide in a neutralization reaction?

http://youtu.be/MO1AqC4udjo



C. MASS-VOLUME OF GAS – an extra piece to the “heart”! BUT ONLY FOR GASES. If you have a substance that is a GAS and you need to go to volume, we can use the molar volume

conversion.

MOLAR VOLUME: 1 mole of a gas = 22.4 Liters (at STP conditions)

STP stands for “Standard Temperature and Pressure”. Standard temperature is 0°C and standard pressure is 1 atm.

If conditions are not at STP, we cannot use this conversion. We will learn what to do in these situations in another unit.

If you have a liquid, you cannot use this conversion.

Example 10-10. How many liters of hydrogen at STP could theoretically be produced by reacting 3.400 grams

of sodium with excess water?

Example 10-11. How many liters of oxygen gas at STP are required to react with 7.98 liters of ethane gas at STP in a combustion reaction?

Example 10-12. If 5.33 X 1023 molecules of hydrogen gas reacts with excess chlorine gas, how many liters of

hydrogen chloride gas could potentially be formed at STP?


Leggett PreAP Chemistry Stoichiometry 10-4 (10) http://youtu.be/AitmQVJoAl4 http://vimeo.com/35933114

D. MASS-VOLUME OF LIQUIDS – what if we can’t use 22.4??!? For liquids, we cannot use the 22.4 L = 1 mole conversion, as we aren’t dealing with a gas. But not to worry – you’ll be given another piece (DENSITY) which you can use as a conversion! Think of the density (x g/mL) as being like an equality ( x grams = 1 mL). See how that looks like a

conversion?

Example 10-13. How many liters of liquid bromine are needed to react completely with 400.0 grams of calcium iodide? The density of liquid bromine is 3.12 g/mL.

Example 10-14. If 46.7 mL of mercury are added to excess silver nitrate, how many grams of mercury (I)

nitrate (Hg2(NO3)2) could theoretically be formed? The density of mercury is 13.6 g/mL.

Example 10-15. How many liters of hydrogen gas at STP are required to react with 40.0 mL of liquid bromine

(density = 3.12 g/mL) in the synthesis of hydrogen bromide?

http://youtu.be/AitmQVJoAl4



Leggett PreAP Chemistry Stoichiometry 10-5 (17:22) http://youtu.be/_weSKP6TDpU http://vimeo.com/36016044

V. PERCENT YIELD

100ltheoretica

alexperiment yield%

Let’s face it… reactions rarely go perfectly to plan. Sometimes, product could splatter out of a container. Or, a reactant could partially evaporate before the reaction proceeds. Maybe an unexpected side reaction occurred, making extra product than you expected. And then sometimes, you just screw up.

Percent Yield allows us to express a ratio of how much product we should have gotten in comparison to how much we actually obtained.

Percent Yield can be above or below 100% - very rarely will you get a 100% yield… unless you are very lucky!

You will use a stoichiometry calculation to find the theoretical yield.

Actual yield will be obtained from experimental values. While many times you will need to find % yield, sometimes the %

yield will be given to you, and you will need to calculate the actual yield. Read your problems carefully!

Types of problems include the following…

Example 10-16. If 32.00 grams of methane at STP are burned in a combustion reaction and 30.00 grams of water are actually produced, what is the % yield (or efficiency) of the reaction?

The “experimental” is how

much you actually made. This

value is determined in lab.

The “theoretical” is the value that is determined from

stoichiometry (a calculation). It is how much you should have

gotten if everything went perfectly with your reaction.

Not always will you actually make

all you were supposed to in a

reaction! The % yield is the %

that you got of what you were

supposed to get. It is sometimes

called the % efficiency.

There is no way we’re getting

100% yield on this lab…

http://youtu.be/_weSKP6TDpU



Example 10-17. The reaction of combustion of octane in an engine is known to be only 74.3% efficient. How many grams of water would you actually expect to form if 90.00 grams of liquid octane are burned?

Example 10-18. A water manufacturer must produce 100 Kg of water per day to stay in business, but the reaction he is using (synthesis of hydrogen and oxygen) is known to be only 86% efficient. How many liters of hydrogen at STP must he plan on reacting (assuming that he has excess oxygen available) in order to actually produce 100. Kg of water per day, considering the 86% yield?

NOTE: remind me to have you complete this in class!

VI. LIMITING REACTANTS (aka Limiting Reagent Problems)

Consider the following… You want to make cookies – as many cookies as possible. But, you don’t want to leave the house. So, you look to see what ingredients you have at home, and compare that to your recipe. Here’s what you find…

Example 10-19. If you can make half-batches, HOW MANY BATCHES OF COOKIES

COULD YOU MAKE?

Basically, you must take each ingredient, and for the amount you have at home, figure out which ingredient will make the smallest amount. This will be the maximum amount of product you can make!

1 BATCH OF COOKIES

RECIPE

4 eggs

2 cups flour

1 ½ cups sugar

2 sticks butter

YOU FIND AT HOME… 6 eggs

5 cups of flour

4 ½ cups of sugar

4 sticks of butter

I make-a the

cookies…


Example 10-20. Which ingredient (reactant) above was the LIMITING REACTANT?

Example 10-21. Which ingredients above were the EXCESS REACTANTS?

Example 10-22. For the butter, how much of it will you end up using to make your cookies?

Example 10-23. How much butter will you have left over?

Example 10-24. For the eggs, how many of them will you end up using to make your cookies?

Example 10-25. How many eggs will you have left over?

Example 10-26. Normally, a batch of cookies should make 48 cookies. However, you ate some dough while you were cooking, and burnt some cookies. You only ended up making only 55 cookies. What was your percent yield?

Leggett PreAP Chemistry Stoichiometry 10-6 (14:23) http://youtu.be/RzBcann7wG0 http://vimeo.com/36016451

While the example above may seem silly, the process is EXACTLY what you will do as you calculate limiting reactant problems!

Your “recipe” will be your chemical reaction. How do I know I have a limiting reactant problem?

o You will be given the starting amounts of BOTH reactants, and asked to find the maximum amount of product you can produce.

How do I solve limiting reactant problems? o TO FIND THE THEORETICAL YIELD: Take both

of your reactants and find how much product they can produce. Whichever product amount is smallest is the maximum amount of product you can make.

o TO FIGURE OUT WHICH IS THE LIMITING REACTANT: Look to see which reactant made the least amount of product. This is your limiting reactant. The other reactant is in excess.

o TO FIGURE OUT THE AMOUNT OF LIMITING REACTANT LEFT OVER: Easy. It will always be NONE. You will use up all of the limiting reactant!

o TO FIGURE OUT THE AMOUNT OF EXCESS REACTANT LEFT OVER: Take the amount of limiting reactant you used, and use stoichiometry to calculate the amount of excess actually used. Then, subtract this value from how much you originally had. (HAD – USED = LEFT OVER)

http://youtu.be/RzBcann7wG0



Example 10-27. 15.00 grams of hydrogen are mixed with 40.00 grams of oxygen in the synthesis of water.

a. What is the theoretical yield (maximum yield) of water in grams? b. What is the limiting reactant? How do you know? c. How many grams of the limiting reactant are left over? d. How many grams of the excess reactant are left over? e. This reaction was done in a laboratory environment and 43.00 g of water was actually collected. What was

the percent yield of this experiment? Reaction:

Start moles

Shift with mmr

Stop (start-shift)


Leggett PreAP Chemistry Stoichiometry 10-7 (14:52) http://youtu.be/wYvfbZEN__Q http://vimeo.com/36016577

Example 10-28. 35.60 grams of zinc are reacted with 100.00 grams of copper (II) sulfate.

a. What is the theoretical yield (maximum yield) of pure copper in grams? b. What is the limiting reactant? What is the excess reactant? c. How many grams of the zinc will be left over? How many grams of the copper (II) sulfate will be left over?

Reaction:

Start moles

Shift with mmr

Stop (start-shift)

http://youtu.be/wYvfbZEN__Q



Example 10-29. 48.0 liters of hydrogen gas at STP are reacted with 15.0 grams of nitrogen gas in a synthesis reaction which produces NH3.

a. What is the theoretical yield (maximum yield) of ammonia in liters, assuming STP conditions?

b. What is the limiting reactant? What is the excess reactant?

c. How many liters of the hydrogen will be left over? How many grams of the nitrogen will be left over?

unit 10 notes: reaction...

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