unit 5: moles and chemical reactions notes unit · pdf file3 mini poster project- classifying...

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Unit 5: Moles and Chemical Reactions Notes

Unit Objectives:

Understand the mole concept

Calculate gram formula mass

Differentiate between formula mass and gram formula mass

Convert between grams and moles

Balance a chemical reaction by adjusting only the coefficients

State the Law of Conservation of Mass and Energy and relate it to balanced chemical equations

Create and use models of particles to demonstrate balanced equations

Identify various types of reactions: synthesis, decomposition, single replacement, & double replacement

Solve mole-mole Stoichiometry problems given a balanced reaction

Calculate the empirical formula from percent mass

Differentiate between empirical and molecular formulas

Determine the molecular formula from the empirical formula and molecular mass

Define the following vocabulary:

Empirical formula Percent CompositionSubscriptReactantProduct

Mole Formula mass (FM)Molar MassGram formula mass (GFM) Coefficient Law of conservation of mass and energy Balanced equation Synthesis reaction Decomposition reaction Single-replacement reaction Double-replacement reaction Molecular formula

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The bold, underlined words are important vocabulary words that you should be able to define and use properly in explanations. This is a study guide for what you will be tested on throughout the year. The objectives are divided into categories of “Knowledge” (what you have to know) and “Application” (what you have to be able to do).

I. CHEMICAL FORMULAS, REACTIONS & STOICHIOMETRY

Knowledge Application

1.

o Chemical formulas are used to represent compounds.o The main types of chemical formulas include: empirical,

molecular, and structural.o An empirical formula is the simplest whole-number ratio of atoms

in a compound.o Molecular formulas are chemical formulas that show the actual ratio

of atoms in a molecule of that compound.o Structural formulas can also be used to represent covalent

compounds. These use lines to show covalent bonds between atomsand also show the geometrical arrangement of atoms in thecompound.

o Determine the empiricalformula from a molecularformula

o Draw structural formulas forcovalent (molecular)compounds

2.

o One mole of any substance is equal to 6.02 x 1023 pieces of thatsubstance.

o The formula mass of a compound is equal to the sum of the atomicmasses of its atoms (units are atomic mass units)

o The molar mass (gram-formula mass) of a substance is equal tothe formula mass in grams – hence “gram-formula mass”.

o The mass of one mole of any substance is equal to its molar mass(gram-formula mass).

o Calculate the molar mass (gram-formula mass) of a substance

o Determine the molecularformula, given the empiricalformula and the molar mass

o Determine the number of molesof a substance, given its massand vice versa

3. o The percent composition by mass of each element in a compound

can be calculated mathematically.

o Calculate the percentcomposition of any element in agiven compound

o Calculate the percentcomposition of water in a givenhydrate

4.

o Balanced chemical equations show conservation of matter, energy,and charge.

o The coefficients in a balanced equation can be used to determinemole ratios in the reaction.

o Balance equations, given theformulas for reactants andproducts

o Calculate simple mole-moleratios, given balanced equations

5. o Types of chemical reactions include synthesis, decomposition,

single replacement, and double replacement.

o Identify the different types ofchemical reactions, given theirchemical equations

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Mini Poster Project- Classifying Reaction Types

Background There are many different types of chemical reactions in the world. We have studied five types of them during this unit. You are now going to create a mini poster describing each chemical reaction and differentiating between the types of chemical reactions.

This will be an independent project. During your allotted time you will further research each of the types of reactions as well as find real-world examples and pictures of each reaction type for visual aide on your poster.

Objective In this project you will create a mini poster that provides information about the five types of chemical reactions (synthesis, decomposition, single replacement, double replacement and combustion). This posted should be created on a manila folder as a bi-fold or as a tri-fold.

Materials You will be provided with colored pencils, markers, paper, and folders. If there is anything additional that you would like to use on your poster you must supply it.

Procedure The following information should be included for each reaction type:

• Name of reaction type• Generic equation• Real Life example of reaction• Definition of reaction type

Each bullet point below should also be considered: • Title of Poster• Color• Pictures• Readability• Sizing & Usage of space• Creativity• Originality

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Lesson 1: The Mole & Gram Formula Mass(GFM)

Objective: To define and calculate molar mass. To apply the formula relating mass in grams to moles

Chemistry is a basic science whose central concerns are -

the structure and behavior of atoms (elements)

the composition and properties of compounds

the reactions between substances with their accompanying energy exchange

The laws that unite these phenomena into a comprehensive system.

We have already studied the structure and behavior of atoms (elements) as well as the composition and properties of compounds. We will now be moving on to learn about the reactions that occur between substances. In order to do this we must be able to quantify the reactants and products to see the changes that are occurring.

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What is a Mole? = 6.02 x 1023 things

Symbol:

Like a dozen, only waaaaaaaaaaay bigger

Why we use it: atoms are soooooo tiny that we have LOTS of them in a given sample

1 dozen eggs = 12 eggs 1 mol eggs = 6.02 x 1023 eggs

Both subscripts and coefficients represent MOLES

Ex: in 1 mole of NaCl, there are 1 mol of Na+ ions 1 mol of Cl– ions

in 1 mole of H2O, there are 2 mol of H atoms 1 mol of O atoms

in 1 mole of Na2SO4, there are 2 mol of Na+ ions 1 mol of SO4

2- ions

How many moles of atoms in total? 7

in 1 mole of (NH4)2CO3, there are 2 mol of NH4 + ions

1 mol of CO3 2- ions

How many moles of atoms in total? 14

*Hydrates:

Ex: Na2CO3 7 H2O in 1 mole of this substance, there are

Molar Mass:

a. Example: the molar mass of carbon is 12.011 g/mol

b. molar mass unit is :

c. molar mass is also called :

d.

Example: H2O =

*Honors Chemistry Only:

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1. Fill in the table below. Put an “M” if the substance is molecular/covalent and an “I” if ionic.

Formula

Moles of each atom

Gram

formula

mass

Formula

Moles of each atom

Total

moles of

atoms

a.

Example: HClO3

M

1 mol H atoms

1 mol Cl atoms

3 mol O atoms

84

g/mol d.

CaCl2

b.

NH4C2H3O2

e.

Mg3(PO4)2

c.

Mg(OH)2

f.

CH3CH2CH3

Classwork 5-1:

Chemical Formula Gram Formula Mass

1. MgBr2

2. KCl

3. FeCl2

4. CrF2

5. Al2S3

6. PbO

7. TiI4

8. Mg3P2

9. SnCl2

10. HgCl2

11. H2SO4

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Summary: Gram-formula mass (also called “molar mass” or GFM)

Definition:

Calculate:

Unit:

Ex: Calculate the gram-formula mass Calculate the gram-formula mass of NaCl of Ba(NO3)2

# Atoms x Mass # Atoms x MassNa Ba Cl N

O

We will need to convert from grams to moles and vice versa for this class. The diagram below summarizes these processes:

Converting Grams to Moles:

From Table T, you would use the MOLE FORMULA:

Example: How many moles are in 4.75 g of sodium hydroxide? (NaOH)

Step 1: Calculate the GFM for the compound.

Na = 1 x = = O = 1 x = = H = 1 x = = +

=

Step 2:

Lesson 2: Mole-Mass Conversions

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Converting Moles to Grams:

The same formula can be used to convert moles back to grams

Example: You have a 2.50 mole sample of sulfuric acid (H2SO4). What is the mass of your sample in grams?


H = 2 x 1.0= 2.0S = 1 x 32.0= 32.0O= 4 x 16.0= 64.0

= 98.0 g/mol

Step 2: Plug the given value of moles and the GFM into the “mole calculation” formula and solve for the mass ofthe sample.

grams =

= 245 g of H2SO4

Classwork 5-2:

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Classwork 5-2(con't):

Activity - Moles of Snacks

Directions: Take the necessary measurements, and record them with units and without rounding. Show all your calculations, rounding your answers to three significant digits and labeling units clearly. Remember to round molar mass to the tenth (0.1) place. Show ALL work!

Pre-lab 1. Find and record the molar mass of sodium________________

2. Find and record the molar mass of sodium chloride_____________

Materials Small bag of snack food

Balance

Procedure 1. Using the information on the snack package answer the following questions.SHow ALL work for any conversions that must be done Final answers must be in threesignificant figures. Molar masses must be rounded to the tenth place with units.

What is the mass of one serving of the snack in the snack bag?

How many mg of sodium are in one serving of the snack?

How many g of sodium are in one serving of the snack?

How many moles of sodium are in one serving of the snack?

How many individual atoms of sodium are in one serving of the snack?

2. Using the electronic balance, mass the bag of snack food and record the result__________

3. Open the snack bag and either empty the contents in the designated container in the lab oreat the contents, sharing with the entire lab group.

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Name:_____________________________________________________Date:_________Period:______

4. Using the electronic balance, mass the empty snack bag and record theresult_________________

5. Calculate the mass of the snack food that was in the bag and record the result. Show yourwork below. Circle your final answer.

6. Is the mass you calculated in #5 the same as the mass of one serving that was recorded onthe label of you snack bag. Yes/No (circle one). Why do you believe this happened? ________ __________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________

Extension: Healthy American adults should restrict their sodium intake to no more than 2,400 milligrams per day. This is about 1¼ teaspoons of table salt (sodium chloride [NaCl]).

Answer the following questions.

1. What is the maximum number of moles of sodium recommended in your diet? How manysodium atoms would this be?

SHow ALL work for any conversions that must be done Final answers must be in threesignificant figures. Molar masses must be rounded to the tenth place with units.

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Moles Lab Activity –Aluminum

Remember: Take the necessary measurements and record them with units. Do not round these

numbers. Show all calculations. Molar mass is rounded to the tenth (10th) place. All other calculations

are rounded to three (3) significant figures. Include units in your final answers.

Materials

Empty aluminum can

Balance

Aluminum foil

Procedure

1. Determine the molar mass of aluminum and record it here:__________________

2. Place the aluminum can on the balance and record the mass._____________

3. Answer the following questions:

Does the aluminum can contain more than, less than, or exactly one mole of aluminum? Explain

your answer using complete sentences or calculations.

Calculate how many moles of aluminum are in one aluminum can. Show ALL work!!

Calculate how many atoms of aluminum are in one aluminum can.

Extension

1. How many cans would you need to have one mole (1.00 mole) of aluminum? Show your work using

the factor label method.

2. Aluminum foil is essentially pure aluminum that has been pounded into a thin sheet. How many

grams of aluminum foil would be necessary to have 1.00 mole of aluminum foil? Explain your

answer using complete sentences or calculations. Have your teacher check your answer before

going on to the next step.

Go to the balance and measure out exactly one mole (1.00 mole) of aluminum foil. As a lab group3.

make a “sculpture” out of the aluminum foil, using the entire mole of foil.

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Lesson 3: Types of Formulas

Objective: To compare and contrast Empirical and molecular formulas. To calculate the molecular formula from the empirical formula and molecular mass.

In order to make it easier to describe elements and molecules, chemical formulas are used.

--

The empirical formula is the

***Ionic compounds are already empirical formulas.

The molecular formula is the It is always whole number multiples of empirical formula. The Subscripts can be reduced same as when you reduce a fraction. Remember…subscripts are just ratios of how the atoms are coming together in a compound!

Determining Empirical Formula

Divide subscripts by the greatest common factor

Example: molecular formula = C4H10

Divide by 2 (greatest common factor)Answer: C2H5

Determining MOLECULAR Formula from the Empirical formula and molecular mass

Step 1: Calculate Gram Formula Mass of the EMPIRICAL FORMULA

Step 2: Divide the MASS of the molecular formula by the MASS of the empirical formula.

Step 3: MULTIPLY the subscripts of the empirical formula by the answer in step 2.

-

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Example: The empirical formula of a compound is C2H3, and the molecular mass is 54.0 grams/mole. What is the molecular formula?

Step 1. Determine the GFM of the empirical formula.

CC22HH33 == ((22 CC XX 1122..00 gg//mmooll)) ++ ((33 HH XX 11..00 gg//mmooll)) == 2277..00 gg//mmoollee

Step 2. Divide the molecular mass by the empirical mass. This will give you a whole-number multiple that tells you how many times larger the molecular formula is than the empirical formula.

((5544..00 gg//mmooll)) // ((2277..00 gg//mmooll)) == 22

Step 3. Multiply the whole number by the empirical formula. This will give the molecular formula.

22 XX CC22HH33 == CC44HH66

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1. Identify each of the following as an empirical or molecular formula. If a formula is molecular, write its empiricalformula.

Formula Empirical or

Molecular?

Simplify if

Molecular

Formula Empirical or

Molecular?

Simplify if

Molecular

NaCl N2O4

C2H6 Ra(CN)2

C6H12O6 Ba(NO3)2

Classwork 5-3:

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6. What is the molecular formula of a compound that has a mass of 289g and an empirical formula of NH3?

7. What is the molecular formula of a compound with a mass of 760g and an empirical formula of Cr2O3?

8. What is the molecular formula of a compound that has an empirical formula of NO2 and molecular mass of 92.0 g?

9. A compound has an empirical formula of HCO2 and a molecular mass of 90 grams per mole. What is the molecularformula of this compound?

VitaminChasanempiricalformulaofC3H4O3andamolecularmassof264.0g/mole.Determinethemolecularformula.

Ibuprofen,acommonheadacheremedy,hasanempiricalformulaofC7H9Oandamolarmassof545.0g/mole.Determinethemolecularformula.

OxalicacidhastheempiricalformulaHCO2andamolarmassof90.0g/mole.Determinethemolecularformula.

10.

11.

12.

Methane Bubbles Demo

Prelab Questions 1. What does the word combustion mean?2. What type of gas is methane and what is its chemical formula?3. What do you thing will happen if we added fire/heat to methane gas?

Analysis 1. Why are the methane bubbles able to rise?2. Write the balanced chemical equation for the combustion of methane.

methane reacts with oxygen to produce carbon dioxide, water, and heat)3. What are the products for the complete combustion of all hydrocarbons?4. List three examples of a hydrocarbon other than methane.5. What do you think will happen if you placed some of the methane bubble

solution on your hands and touched a burning candle to it?6. Why is your hand unharmed by the flame?

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Lesson 4 Types of Chemical Reactions

Objective: Identify various types of reactions: synthesis, decomposition, single replacement, & double replacement

1. SynthesisDuring a

General pattern:

Example:

2Mg(s) + O2(g) → 2MgO(s)

2. DecompositionDuring a

General pattern:

Example: 2HgO(s) → 2Hg(ℓ) + O2(g)

3. Single ReplacementDuring a

General Pattern:

Use Table J

Example: Zn(s) + CuSO4(aq) Cu(s) + ZnSO4(aq)

Cu(s) + ZnSO4(aq)

4. Double Replacement

During a

General Pattern:

Example: AgNO3(aq) + NaCl(aq) → NaNO3(aq) + AgCl(s)

A PRECIPITATE

Check Table F to see if there is a PRECIPITATE during a double replacement reaction.

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Classwork 5-4:Practice: Determine if the following reactions are synthesis (S), decomposition (D), single replacement (SR), or double

replacement (DR) reactions.

1. H2 + Cl2 2HCl ________

2. NaOH + HCl NaCl + H2O________

3. 2NH3 N2 + 3H2________

4. Mg + H2SO4 MgSO4 + H2________

5. F2 + 2HBr Br2 + 2HF________

6. Zn(NO3)2 + CaCO3 Ca(NO3)2 + ZnCO3________

7. 4Al + 3O2 2Al2O3________

8. CuSO45H2O CuO + H2SO4 + 4H2O________

9. SiF6 + 6Xe SiXe6 + 3F2________

10. 2 NaClO3 2 NaCl + 3 O2 ________

11. 2 AgNO3 + Ni Ni(NO3)2 + 2 Ag________

12. H2CO3 H2O + CO2 ________

13. BaCO3 BaO + CO2________

14. 4 Cr + 3 O22 Cr2O3________

15. Ca + 2 HCl CaCl2 + H2 ________

16. Ca(C2H3O2)2 + Na2CO3 CaCO3 + 2 NaC2H3O2 ________

17. Cu(OH)2 + 2 HC2H3O2 Cu(C2H3O2)2 + 2 H2O________

18. 8 Cu + S8 8 CuS ________

19. P4 + 5 O2 2 P2O5________

20. 2 K + 2 H2O2 KOH + H2 ________

Single Replacement Lab

Objective: Learn how to use the activity series to predict the products of a single replacement

reaction.

Materials:

6M HCl

Well Plates

Scoops

Metals (magnesium, iron, zinc, copper, aluminum)

Procedure:

1. Label each well for a particular piece of metal on a separate

piece of paper.

2. Put a very small piece of the metal into each well as

assigned in step #1.

3. Place 2 drops of the 6M hydrochloric acid into the first well. Observe the reaction.

4. Repeat step 4 for each of the other metal samples.

5. Rinse off larger pieces of metal before disposing of them in the trash can. Rinse all

wells of the tray. Let dry.

Analysis:

Write a balanced chemical equation for each reaction that occurred. Include the

proper states of matter (i.e. (s) = solid, (l) = liquid, (g) = gas, (aq) = aqueous

solution).

Mg Fe Zn Cu Al

Observations: Note any observations for each metal

MgCuFeAlZn

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Introduction Not all chemicals react when combined. The activity series is a listing of the level of reactivity that elements possess. It will help you determine whether or not a specific single replacement reaction will occur.

Purpose The purpose of this demonstration is to exhibit how the Activity Series can be useful in determining whether or not a single replacement reaction will occur.

Predictions The following “reactants” (chemicals) will be mixed together. I hesitate to call them all reactants because they will not all react with each other. Determine which will react and which will not using the Activity Series.

Will a reaction occur? Did a reaction occur?

# “Reactants” (Chemicals) Yes No

1 CuSO4 + Fe →

2 Al + CuSO4 →

3 NaCl + Al →

4 Mg + CaCl2 →

5 NaCl + Cu →

6 HCl + Al →

Follow-up 1. After the demonstration, indicate in the table above whether or not a reaction took place (in the furthest-right column).

2. For the reactions that took place:a. Finish the chemical equations in the table above by writing what products were formed.b. Balance each of the reactions in the space below. Be neat.

3. For the reactions that did not take place:Finish the chemical equations in the table above by writing ‘NR’ on the product side. This indicates that no reaction willtake place.

Lab - Single Replacement Reactions & Table J

Lesson 5: Balancing Equations

Objective: Balance a chemical reaction by adjusting only the coefficients

Review:

REACTANTS =

PRODUCTS =

COEFFICIENT =

SUBSCRIPT =

*COEFFICIENTS and SUBSCRIPTS tell us how many moles we have for each element

Chemical Symbols (states of matter)

(s) (g) (l) (aq)

l

Example: : 2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)

CATALYST=

Example: 2 H2O2(aq) ---- -> 2 H2O(l) + O2(g)

The catalyst in this example is KI

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A chemical equation is a set of symbols that state the products and reactants in a chemical reaction.

This goes for mole amounts as in balancing equations and mass amounts.

Law of Conservation of Mass

Example: Given the reaction: N2 + 3H2 ---------- 2NH3 What is the total number of grams of H2 that reacts when 14 grams of N2 are completely consumed to produce 20 grams of NH3?

https://youtu.be/UGf60kq_ZDI

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An equation is balanced when:

**NOTE:

Steps for Balancing Equations

Step 1: Find the most complex compound in the equation. Balance the elements found in that compound on the opposite side of the arrow by changing the coefficients for those atoms.

Step 2: Continue balancing until all atoms are balanced (save pure elements for last)

Step 3: Go back and check each atom to see if it is balanced on both sides of the equation.

Step 4: POLYATOMIC IONS may be balanced as a SINGLE UNIT rather than as separate elements as long as they stay intact during the reaction.

Lab -Investigating the Law of Conservation of Mass

Objective: To corroborate the law of conservation of mass through laboratory experimentation.

Background:

When wood burns or water evaporates, some or all of it appears to disappear but does it? A pile of

ash is a lot smaller than the wood it originally started as and a dry empty cup seems a lot lighter if all

of the water is allowed to evaporate. When you pop open a soda can, the fizz is really carbon

dioxide, tiny gas bubbles that rise out of the soda and mix back with the atmosphere. To the naked

eye, it seems like these materials just vanish. However, does matter ever really disappear?

According to the law of conservation of mass, mass is always conserved. In many situations, the end

products appear to have decreased but when all the products are captured and their mass

measured, it was discovered that the mass of the reactants is equal to the mass of the products. In

other words, matter can neither be created nor destroyed. This is called the law of conservation of

mass. All atoms present in the reactants are also present in the products.

According to this principle, matter cannot just appear or disappear, but matter can change its

nature. When atoms chemically combine with or split away from other atoms the chemical and

physical properties of that substance can be completely altered. Take table salt, for example. The

chemical formula for salt is NaCl, or sodium chloride. Sodium by itself is an element an unstable and

poisonous metal. Chlorine by itself is also an element, which at room temperature is a poisonous

gas. However, combine these two atoms together with a chemical bond and they are completely

transformed into a very stable substance that is edible crystalline solid at room temperature.

Materials:

Sodium bicarbonate, NaHCO3 scoopula Funnel

Acetic acid, CH3COOH 250 mL Erlenmeyer flask Balloon

Triple beam balance Graduated cylinder Masking tape

The "Law of Conservation of Mass" states that when matter goes through a physical or

chemical change, the amount of matter stays the same before and after the changes occur.

In other words, matter cannot be created or destroyed.

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Procedure:

1. Observe the physical properties of acetic acid and sodium bicarbonate and record your

observations in data table 1.

2. Put your funnel inside the balloon. Make sure the neck of the funnel is deep inside the

balloon, so you do not get any sodium bicarbonate stuck in the neck of the balloon.

3. Put 4 scoops of sodium bicarbonate in the funnel.

4. Measure out 50 mL of acetic acid. Pour the acetic acid into the Erlenmeyer flask.

5. Stretch the balloon over the mouth of the flask, with the balloon hanging to the side of the

flask. (DO NOT let the reactants mix). Use masking tape to seal the balloon to the neck of

the flask.

6. Record the combined mass of the ENTIRE SETUP in data table 2

7. Lift the balloon so that the sodium bicarbonate now drops into the Erlenmeyer flask.

8. Observe for 5 minutes.

9. Observe the physical properties of the products and record your observations in data table

1.

10. Record the mass of the ENTIRE SETUP again data table 2.

Data:

Table 1:

Substance Physical Properties

reactants Acetic Acid, CH3COOH(aq)

Sodium Bicarbonate, NaHCO3(s)

products NaCH2COOH(aq)

H2O(l) +CO2(g) 29

Table 2:

Total Mass Before Reaction (grams)

Total Mass After Reaction (grams)

Results:

1. What evidence that a chemical reaction took place did you observe when the sodiumbicarbonate and acetic acid were mixed?

2. Calculate the change in mass that you measured between the initial reactants and finalproducts of the reaction.

3. Calculate the percent error between your measured final mass and the final mass predicted bythe law of conservation of mass.

4. Why was it necessary to do this experiment in a sealed container (closed system) in order toaccurately illustrate the principle of conservation of mass?

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LawofConservationofMassWorksheetSince there is a conservation of mass in all balanced chemical reactions, these problems are also quite easy. Consider the following balanced chemical reaction CaCO3 à CaO + CO2. In this reaction the mass of CaCO3 must equal the combined mass of CaO and CO2. If you know the mass of the reactants you automatically know the mass of the products and vice-versa.

Using the reaction above, determine the mass of CaO produced if 200 grams of CaCO3 decomposed and produced 88 grams of CO2.

CaCO3 à CaO + CO2 200g = Xg + 88g à 200g – 88g = 112g CaO produced

Complete the following problems

1) Hydrogenandoxygenreactchemicallytoformwater.Howmuchwaterwouldformif14.8gramsofhydrogenreactedwith34.8gramsofoxygen?(H2+O2àH2O)

2) Whenammoniumnitrate(NH4NO3)explodes,theproductsarenitrogen,oxygen,andwater.When40gramsofammoniumnitrateexplode,14gramsofnitrogenand8gramsofoxygenform.Howmanygramsofwaterform?(NH4NO3àN2+O2+H2O)

3) 40gofcalciumreactswith71gofchlorinetoproduce_____gofcalciumchloride.

4) _____gofpotassiumreactswith16gofoxygentoproduce94gofpotassiumoxide.

5) 14goflithiumreactionwith_____gsulfurtoproduce46goflithiumsulfide.

6) 24gofmagnesiumreactswith38goffluorinetoproduce_____gmagnesiumfluoride.

7) 65.5gcopperreactswith_____goxygentoproduce81gcopper(I)oxide.

8) 88gofstrontiumreactswith160gbrominetoproduce_____gstrontiumbromide.

9) 46gofsodiumreactswith_____gchlorinetoproduce117gsodiumchloride.

10) ____gironreactswith71gchlorinetoproduce129gofiron(II)chloride.

11) 137gofbariumreactswith______giodinetoproduce391gbariumiodide.

12) _____ghydrogenreactswith32gofoxygentoproduce34gofhydrogenperoxide.

13) Whydowebalancechemicalreactions?31

Classwork 5-5:

32

Practice : Balance the following equations in the space provided - Use a pencil

1) ___ NaNO3 + ___ PbO ___ Pb(NO3)2 + ___ Na2O

2) ___ AgI + ___ Fe2(CO3)3 ___ FeI3 + ___ Ag2CO3

3) ___ C2H4O2 + ___ O2 ___ CO2 + ___ H2O

4) ___ ZnSO4 + ___ Li2CO3 ___ ZnCO3 + ___ Li2SO4

5) ___ V2O5 + ___ CaS ___ CaO + ___ V2S5

6) ___ Mn(NO2)2 + ___ BeCl2 ___ Be(NO2)2 + ___ MnCl2

7) ___ AgBr + ___ GaPO4 ___ Ag3PO4 + ___ GaBr3

8) ___ H2SO4 + ___ B(OH)3 __ B2(SO4)3 + ___ H2O

9) ___ Fe2O3 + ___ H2 ___ Fe + ___ H2O

10) ___ Li + ___ N2 ___ Li3N

11) ___ Zn + ___ HCl ___ ZnCl2 + ___ H2

12) ___ NaCl + ___ AgNO3 ___ NaNO3 + ___ AgCl

13) ___ Ca3P2 ___ Ca + ___ P

14) ___ HCl + ___ F2 ___ HF + ___ Cl2

15) ___ (NH3)2CO3 + ___ CaSO4 ___ CaCO3 + ___ (NH3)2SO4

333

Activity - The Balanced Equation is a Recipe Introduction Chemists spend a lot of time and effort making things. Just like a chef spends lots of time and effort making food, a chemist spends lots of time making chemicals. Chefs use recipes to make food; chemists use recipes to make chemicals.

Purpose The purpose of this activity is to relate a recipe for brownies to a balanced chemical equation and use this relationship to determine quantities of ingredients (reactants) and products.

This is my mom’s recipe for one batch of homemade brownies – to be cooked in a 9” x 13” pan. o 2 sticks butter (B) o 4 eggs (E) o 1.67 cups sugar (S) (which

is 1 & 2/3 cups)

o 1.5 cups flour (F) o 1 tsp salt o 2 tsp vanilla

o 4 squares unsweetened chocolate

o ¼ tsp baking powder o 1 ¾ tsp baking soda

So the recipe goes like this:

2B + 4E + 1.67S + 1.5F (+ more) → 1 batch

1. If you had 8 eggs (and everything else you need), how many batches of brownies could you make?

2. Using the 8 eggs, how many sticks of butter would you need to take out of the fridge?

3. If you had 4.5 cups of flour and wanted to make as many brownies as possible, how many eggs would you need to gowith your 4.5 cups of flour?

4. You are really hungry. If you had 3 eggs and wanted to make the most brownies that you could, how many cups of sugarwould you need to go with those eggs?

5. To make 3 batches of brownies, how many cups of flour (and of course all the other necessary ingredients) would youneed? Use the flour to batch ratio to calculate.

6. 2.5 cups of sugar needs how much flour in order to use up all of the sugar?

You see, deciding how much of each ingredient or how many batches you can make is nothing more than manipulating the ratios or proportions that exist between the ingredients. If you don’t maintain those set proportions, it is likely that your brownies will not taste very good.

34

7. Write the balanced equation for the synthesis of ammonia (nitrogen trihydride) from nitrogen gas and hydrogen gas.(Use this balanced equation to solve problems 8 – 11)

8. Given 7 moles of nitrogen gas, how many moles of hydrogen gas do you need to go with it?

9. Given 7 moles of nitrogen gas, how many moles of ammonia can you make? Would your answer be the same if youstarted with 21 moles of hydrogen?

35

Lesson 6 : Calculating Mole Ratios

Objective: Solve mole-mole Stoichiometry problems given a balanced reaction

Stoichiometry (pronounced as: stow – ik – ee – om′ – etree) is the calculation of quantities in chemical equations. The question involved with Stoichiometry is “How much…? For example: based on the balanced chemical equation 3 A2(gs) + 2 B(aq) → 2 A3B(s) How much:

a) of reactant A is needed to completely react with 5.00 g of reactant B?b) solid product will form if 0.25 g of reactant A is used?

One of the important concepts in Stoichiometry is that of mole ratios. Mole ratios are the ratios of the number of moles of each reactant and product to each other.

When we balance chemical equations with the coefficients, we set up the whole number ratios between reactants, between reactants and each product, and between products.

For example, consider the following balanced chemical equation:

N2 (g) + 3 H2 (g) 2 NH3 (g)

Using the coefficients to represent moles of a substance,

Based on the correctly balanced chemical equation, the mole ratios for this reaction are:

ratios between reactants: 1 mole N2 3 moles H2 3 moles H2 1 mole N2

ratios between each reactant and the product:

1 mole N2 2 moles NH3 3 moles H2 2 moles NH3

2 moles NH3 1 mole N2 2 moles NH3 3 moles H2

*****Mole ratios for a given chemical equation are always based on the balanced chemical equation and NEVER change. ******

For example: N2 + 3 H2 2 NH3

1 mole 3 moles 2 moles

2 moles 6 moles 4 moles

3 moles 9 moles 6 moles

0.5 moles 1.5 moles 1.0 moles

Remember: the balanced chemical equation represents the SMALLEST whole number ratios of reactants and products.

Each of the mole ratios from the original balanced equation may be used as a conversion factor to change from moles of one substance to moles of another substance in the same chemical equation. You MUST use the ORIGINAL balanced chemical equation with the correct coefficients as your reference for mole ratios.

36

Using the reaction of nitrogen gas and hydrogen gas to form ammonia gas;

N2 (g) + 3 H2 (g) 2 NH3 (g)

Suppose you had only 0.25 mole N2(g) available. How much hydrogen gas would you need to react completely with the available nitrogen gas? How much ammonia would be formed?

(1) Start with the amount you are given in the problem

(2) multiply the starting amount by the appropriate mole ratio to convert from the unit with which you start to the substance desired

(3) cancel units

(4) do the arithmetic and write the answer with appropriate significant figures, correct unit, and correct chemical formula

0.25 mole N2 x 3 moles H2 = 0.75 moles H2 are needed 1 mole N2

0.25 mole N2 x 2 moles NH3 = 0.50 mole NH3 would be formed 1 mole N2

NOTE: In these calculations, the numbers in the mole ratios do NOT have significant figures. USE THE NUMERICAL VALUES AS PRINTED IN THE QUESTION AS THE BASIS FOR YOUR SIGNIFICANT FIGURES. In this calculation, the given value of 0.25 mole N2(g) has 2 significant figures, therefore the answers must each have 2 sig. figs.

More examples: Consider the reaction between sodium metal and chlorine gas to form solid sodium chloride.

2 Na(s) + Cl2(g) → 2 NaCl(s)

If 25.0 moles of sodium chloride are formed, how many moles of sodium metal were used? How many

moles of chlorine gas were used?

25.0 moles NaCl x 2 moles Na = 25.0 moles Na were used 2 moles NaCl

25.0 moles NaCl x 1 mole Cl2 = 12.5 moles Cl2 were used 2 moles NaCl

37

C3H8 + 5O2 → 3CO2 + 4H2O

1. If 12 moles of C3H8 react completely, how many moles of H2O are formed in the reaction above?

C3H8 + 5O2 → 3CO2 + 4H2O

2. If 20 moles of CO2 are formed, how many moles of O2 reacted in the reaction above?

C3H8 + 5O2 → 3CO2 + 4H2O

3. If 8 moles of O2 react completely, how many moles of H2O are formed in the reaction above?

N2 + 3H2 → 2NH3

4. If 2.5 moles of N2 react completely, how many moles of NH3 are formed in the reaction above?

Classwork 5-6:Determine the molar amount using mole to mole ratios for the following reactions

5.

6.

7.

38

8. Use the balanced reaction below and the relationship that “1 mole of a compound = the gram formula mass of thatcompound” to answer the questions below.

2H2O 2H2 + O2

a. Calculate the gram formula masses for the three substances seen in the reaction abov

b. How many moles are present in 54 grams of H2O? (Remember: one mole of H2O is ALWAYS equal to 18 grams)

c. What is the ratio of H2O to H2 moles according to the balanced reaction above?

d. Using the reaction above (remember, it is just like a recipe!), how many moles of H2 would be produced if 4 moles ofH2O are used?

e. How many grams of H2 are present in 4 moles of H2?

f. What is the ratio of H2 to O2 in the reaction above?

g. If you have 2.5 moles of H2O, how many moles of O2 will be produced?

h. What is the ratio of H2O to O2 in the reaction above?

i. If you produce 0.25 moles of O2, how many moles of H2O did you react?

39

Lesson 7: Percent Composition

Objective: To apply a formula to calculate % composition

A common practice in the chemical laboratory is to determine the percent composition of each element in a compound and use that information to determine the formula that shows the simplest whole-number ratio of elements present in the compound, the empirical formula of the compound.

Percent composition is defined as

Calculating Percent Composition :Formula located on Table T

EXAMPLE: What is the percent composition of Calcium in CaCl2


= 111.0 g/mol

Step 2: Use the formula to find the % composition of each element or “part” in our compound (to the nearesttenth of a %).

%%CCaa iinn CCaaCCll22 ==4400..0088 gg == 3366..1111 %%

111111..00 gg

40.1 g

70.9 g

40

*% Composition of Hydrates

Water, the most common chemical on earth, can be found in the atmosphere as water vapor. Some chemicals, when exposed to water in the atmosphere, will reversibly either adsorb it onto their surface or include it in their structure forming a complex in which water generally bonds with the cation in ionic substances. The water present in the latter case is called water of hydration or water of crystallization. Common examples of minerals that exist as hydrates are gypsum (CaSO

4•2H

2O), Borax (Na

3B

4O

7•10H

2O) and Epsom salts (MgSO

4•7H

2O). Hydrates generally contain water in

specific amounts. The hydrate’s formula are represented using the formula of the anhydrous salt (non-water) component of the complex followed by a dot then the water (H

2O) preceded by a number corresponding to the ratio of

H2O moles per mole of the anhydrous component present. They are typically named by stating the name of the

anhydrous component followed by the Greek prefix specifying the number of moles of water present then the word hydrate (example: MgSO

4•7H

2O: magnesium sulfate heptahydrate).

Hydrate- Ionic solids with water trapped in the crystal lattice. It is written like this:

Ionic Compound’s Formula • n H2O (n) is a whole number

Notice how WATER molecules are BUILT INTO the chemical formula

Anhydrous salt - A crystal with NO WATER trapped inside its lattice

Calculate % Composition of water in a Hydrate Step 1: Calculate GFM of the HYDRATE Step 2: Plug into % composition formula from Table T

% composition by mass = mass water X 100 mass of whole hydrate

Example: What is the percentage by mass of water in sodium carbonate crystals (Na2CO310H2O)?

Step 1- Calculate GFM of Hydrate

=part

=whole

Step 2- Plug values into Formula

% H2O by mass =

****Pay attention because in some problems you will have to calculate the amount of water lost by subtracting the anhydrate mass from the hydrate mass.

*= Honors Chem Only

Water in a Hydrate using Lab data:

RESULTS:DataRemember to record masses to two decimal places:1. Mass of Crucible _____________g2. Mass of Crucible + Hydrate before heating ___________g3. __________ g

0.50

5.00= 10.0%

41

Mass of Crucible + Anydrate after heating

Honors Chem Only

Percent of Sugar in Gum

Purpose: Calculate the percent composition by mass of an ingredient in a commercial food product.

Background: All commercial food producers must list all ingredients found in their food products on the retail packaging for the foods. The Wrigley Company provides the following nutritional facts for Juicy Fruit Gum.

Serving size 1 stick (2.7 g)

Calories 10

Total Fat 0 g

Sodium 0 g

Total carbohydrate 2 g

Sugars 2 g

Protein 0 g

Chewing gum is composed mostly of sugar. The most convenient way to remove sugar from gum is to chew it. During chewing the sugar is dissolved by saliva and is swallowed, leaving an insoluble gum base. This laboratory exercise will establish the percent composition of sugar found in chewing gum.

All percentages are calculated the same way. Percent = Part/whole x 100. For example, if a piece of gum had a mass of 5.0 grams, and the sugar in the gum had a mass of 3.0 grams, the percent of the gum that was sugar would be calculated by: 3.0/5.0 x 100 = 60 %.

Procedure:

Mass of Gum + Wrapper (g)

Mass of Wrapper (g)

Calculated Mass of Gum alone (g)

Mass of Chewed Gum + Wrapper

(g)

Calculated Mass of Chewed Gum

alone (g)

Calculated Mass of Sugar

Removed after chewing(g)

1. Obtain a piece of gum from your teacher. DO NOT OPEN IT YET.2. Mass the gum in the wrapper, record the mass in the data table above.3. Open the gum, begin chewing. Mass the empty wrapper, record in the data table, and save it.4. After chewing until the flavor disappears (or 10 minutes, whichever comes first), place the chewed gum

on your empty wrapper. (While you are chewing the gum, answer the following questions. Youmay work together on the questions, but each group member must write the answer on his/herown paper)

Questions:

At home, I have a jar of coins. If I have 350 coins in the jar, and 120 of them are quarters, whatpercent of the coins are quarters?

42

Determine the percentage of sugar in the gum from the manufacturer’s nutritional statement. *** (Later,you will use this answer to calculate percent error in your analysis section.)

In this lab, you will mass a piece of gum to find it’s total mass. You will then chew the gum to removethe sugar. If the sugar is gone after chewing, how will you be able to find the mass of sugar that was inthe gum?

5. After 10 minutes have passed, re-mass the chewed gum and wrapper together and record the value inthe data table. You may now throw the gum and wrapper away.

6. Complete the calculations that are necessary to fill in your data table.

Analysis of Data:

Calculation for your sample % sugar in your sample Group Average (Show work)

1. Determine the percentage of sugar in your sample. Label and show work for your calculations in thetable above. Record your final answer with the correct number of significant figures in the % sugarcolumn.

2. Determine the average percentage of sugar for your lab group. Show work and record in table above.3. Determine the percent error between your group average (measured) and the manufacturer’s

nutritional statement that you already calculated in the question ***section (accepted).

4. Explain the percent error that you calculated, including a discussion of possible errors. Remember:Human or math error is not an acceptable answer.

43

Lab - Percent of Water in Epsom Salt

Goggles must be worn!

Overview:

Epsom salt (aka magnesium sulfate) is a combination of MgSO4 and H2O. Many ionic

compounds incorporate a fixed number of water molecules into their crystal structures. These are

called hydrates. Heat can be used to dehydrate a hydrated salt causing the H2O molecules to

evaporate and produce an anhydrous salt which often will appear different than its hydrate.

When expressing the formula for a hydrate, it is necessary to notate the fixed number of H2O

molecules. A large dot is placed between the formula and the H2O molecules following the formula

for the ionic compound.

For example: CuSO4 • 5 H2O is the formula for copper sulfate. This formula indicates that for

every 1 mole of CuSO4, 5 moles of H2O are present.

Purpose: 1) Dehydrate Epsom salt using chemistry procedures and determine the percent by mass of water. Extra Credit 2) Use your math skills to determine the mole ratio, or how many moles of H2O

there are compared to moles of MgSO4 (Epsom salt) 3) Compare your lab determined mole ratio

to the real mole ratio on the box of Epsom salt.

Materials: Bunsen burner

BalanceRing stand w/ ring

Pipestem triangle

crucible Tongs

Epsom salt (MgSO4)

Procedure:

1. Put on your goggles. Secure the iron ring on the ring stand acouple of inches above the height of the burner. Place

the pipestem triangle on the ring.

2. Place a clean crucible on the set-up. Light

the burner and heat for a couple of minutes to make certain

container is thoroughly dry. Turn off burner and cool the

container for several minutes until it is comfortable to touch.

Record the mass of the dry container.

3. Add about 5 grams of Epsom salt to the container. Record

the mass of the container and Epsom salt.

4. Place the container back on the set-up and heat gently with hot flame until the water has been

released from the hydrate. This will require about 5 minutes. (see illustration)

5. When no more H2O appears to be coming from the hydrate, turn off the burner and cool for

several minutes until container is comfortable to the touch. Record the mass of the container and

salt.

6. If time allows, reheat the container with salt, cool and find the mass again. If the two final

masses agree, you can be confident that you have indeed released all of the H2O. If not, continue

the heating and cooling as directed in this lab to make sure all the water is released.

7. Put your excess Epsom salt into the trash can and carefully clean out your crucible. All other

lab materials need to be ready for the next class.

Honors Chem Only

44

Data / Observations:

Mass of Epsom Salt before and after heating

1. Mass of a clean, dry, empty container g

2. Mass of container & hydrated salt

(before heating)

g

3a. Mass of container & anhydrous salt

(after heating 1st time)

g

3b. Mass of container & anhydrous salt (2nd time)

*repeat until no change in mass

g

3c. Mass of container & anhydrous salt (3rd time)

*repeat until no change in mass

g

Calculations/ Analysis (must show work):

1. a. Determine mass (g) of H2O lost from your salt (how much did you “cook out?”): Using your

data table, determine the numbers to subtract in order to figure out mass (g) H2O. Use this number to calculate the percent by mass of water in the hydrate-Use the formula on Reference Table T.

b. Determine mass (g) of the anhydrous MgSO4 used in this lab (just the anhydrous salt- not

the molar mass). Using your data table, calculate the mass of the anhydrous MgSO4 used.

2. a. Calculate how many moles of H2O were dehydrated from your salt.

Convert your grams of H2O (found in 1a) to moles of H2O.

b. Calculate how many moles of MgSO4 salt. Convert your grams of

MgSO4 (found in 1b) to moles of MgSO4.

3. a. Determine the mole ratio of H2O molecules to MgSO4 particles. Divide the moles of H2O

(calculated in 2a) by the moles of MgSO4 (calculated in 2b) to determine the ratio of moles of

H2O to moles of salt (moles H2O / moles MgSO4) . Round your answer to the nearest small

whole number.

b. Write the correct formula for your hydrate using the ratios 1MgSO4•_____H2O (useyour answer from #3 to fill in the blank).

Conclusion/Discussion:

1. Copper Nitrate is a hydrate with the following formula: Cu(NO3)2 • 3 H2O. What is the ratio

between moles of copper nitrate and moles of water in this hydrate?

2. What is the percent by mass of water in the hydrate in Cu(NO3)2 • 3 H2O? Show All Work!!!

Extra Credit (must show work):

45

46

1. What is the percentage by mass of carbon in CO2?

2. What is the percent by mass of nitrogen in NH4NO3?

3. What is the percent by mass of oxygen in magnesium oxide(MgO)?

4. What is the percent by mass of water in BaCl22H2O?

5. A 10.40 gram sample of hydrated crystal is heated to a constant mass of 8.72 grams. Thismeans all of the water has been driven out by the heat.

a) Calculate the mass of water that was driven out:

b) Calculate the %mass of water in the hydrate.

Classwork 5-7: Show all calculations - Round percentages to whole number

Question 4 & 5 - Honors Chem Only

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51

Unit Review:

Moles and Chemical Reactions Place a checkmark next to each item that you can do! If a sample problem is given, complete it as evidence.

Place a checkmark next to each item that you can do! If a sample problem is given, complete it as evidence.

_____1. I can still do

everything from Unit 1.









______6. I can calculate the

gram formula mass of a

compound or substance

a. O2

b. Na3PO4

______7. I can calculate the

number of moles of a substance

when given the mass

a. 64g O2

b. 567 g Na3PO4

______8. I can calculate the of

the mass a substance when

given number of moles

a. 7 moles O2

b. 0.6 moles Na3PO4

52

_____9. I can define empirical

formula, molecular formula,

and hydrate.

Definitions:

empirical formula

molecular formula

_____10. Given the empirical

formula and the molar mass, I

can determine the molecular

formula of a compound.

What is the molecular formula of a compound that has the empirical

formula of CH and a molar mass of 78 g/mol.

_____11. I can use particle

diagrams to show conservation

of mass in a chemical equation.

Using the symbols shown below, complete the equation below to

illustrate conservation of mass.

2Al + 3Br2 -----> 2AlBr3

_____12. I can balance a

chemical equation showing

conservation of mass using the

lowest whole number

coefficients.

Balance the following chemical equation using the lowest whole number

coefficients.

_____Al2(SO4)3 + _____Ca(OH)2 -----> _____Al(OH)3 +

_____CaSO4

_____13. Given a partially

balanced equation, I can

predict the missing reactant or

product.

Use the law of conservation of mass to predict the missing product.

2NH4Cl + CaO -----> 2NH3 + ______________ + CaCl2

= Al

= Br

Br

53

_____14. Given a list of

chemical reactions, I can

classify them as being a

synthesis reaction,

decomposition reaction, single

replacement reaction, or

double replacement reaction.

Classify the following reactions as synthesis, decomposition, single replacement, or double replacement.

_____15. Given a balanced

equation, I can state the mole

ratios between any of the

reactants and/or products.

Given the following balanced equation, state the mole ratios between the

requested substances.

C3H8(g) + 5O2(g) -----> 3CO2(g) + 4H2O(l)

The mole ratio between C3H8 and O2 is _______C3H8:_______O2.

The mole ratio between C3H8 and CO2 is _______C3H8:_______CO2.

The mole ratio between C3H8 and H2O is _______C3H8:_______H2O.

The mole ratio between CO2 and O2 is _______CO2:_______O2.

The mole ratio between H2O and CO2 is _______H2O:_______CO2.

_____16. I can define

stoichiometry.

Definition:

stoichiometry

_____17. Given the number of

moles of one of the reactants or

products, I can determine the

number of moles of another

reactant or product that is

needed to completely use up

the given reactant/product.

Using the equation from question #20, determine how many moles of O2

are needed to completely react with 7.0 moles of C3H8.

Using the equation from question #20, determine how many moles of

CO2 are produced when 7.0 moles of C3H8 completely react.

1) How many oxygen atoms are represented in the formula Al2(CO3)3?

a) 3 b) 9 c) 10 d) 6

2) What is the total number of atoms present in 1 mole of Ca3(PO4)2?

a) 8 b) 5 c) 10 d) 13

3) What is the total mass of iron in 1.0 mole of Fe2O3?

a) 72 g b) 112 g c) 56 g d) 160 g

4) What is the gram formula mass of Li2SO4?

a) 206 g b) 55 g c) 110 g d) 54 g

5) What is the percent composition by mass of sulfur in H2SO4? [formula mass = 98]

a) 98% b) 16% c) 65% d) 33%

6) A hydrate is a compound with water molecules incorporated into its crystal structure. In an

experiment to find the percent by mass of water in a hydrated compound, the following data

were recorded:

Mass of test tube + hydrate crystals before heating 25.3 grams

Mass of test tube 21.3 grams

Mass of test tube + anhydrate crystals after heating 22.3 grams

What is the percent by mass of water in the hydrate?

a) 75% b) 50% c) 8.0% d) 95%

7) Which of the following statements explains why mass is lost when a student heats a sample of

BaCl22H2O crystals?

a) water is given off as a gas c) chlorine is given off as a gas

b) the crystals sublime d) the crystals fuse (melt)

8) When the equation H2S + O2 H2O + SO2 is completely balanced using the smallest whole

numbers, the sum of all the coefficients is

a) 9 b) 11 c) 7 d) 5

Unit 5 Practice Test

54

9) What is the percent by mass of water in the hydrate Na2CO3●10H2O [formula mass = 286]?

a) 26.1% b) 62.9% c) 6.89% d) 214.5%

10) How many grams are there in 2.5 moles of C2H5OH?

a) 18 b) 115 c) 46 d) 0.05

11) Which quantity is equivalent to 146 grams of NaCl?

a) 1.0 mole b) 2.5 moles c) 2.0 moles d) 1.5 moles

12) When the equation H2 + Fe3O4 --> Fe + H2O is completely balanced using the smallest

whole numbers, the coefficient of Fe would be

a) 1 b) 2 c) 3 d) 4

13) When the equation NaBr + H3PO4 Na3PO4 + HBr is balanced using the smallest whole

numbers, the sum of the coefficients will be

a) 6 b) 8 c) 5 d) 4

14) Given the equation: 2C2H2 + 5O2 4CO2 + 2H2O

How many moles of oxygen are required to react completely with 1.0 mole of C2H2?

a) 2.5 b) 5.0 c) 10 d) 2.0

15) Given the reaction: 2C2H6 + 7O2 4CO2 + 6H2O

What is the ratio of moles of CO2 produced to moles of C2H6 consumed?

a) 7 to 2 b) 2 to 1 c) 1 to 1 d) 3 to 2

16) Given the following balanced equation: N2 + 3 H2 2 NH3

If you have 8 moles of nitrogen, how many moles of NH3 will you produce?

a) 16 b) 18 c) 24 d) 2

17) Given the reaction: 2Na + 2H2O 2NaOH + H2, what is the total number of

moles of hydrogen produced when 4 moles of sodium react completely?

a) 1 b) 2 c) 3 d) 4

55

18) What type of reaction best describes the following chemical reaction?

Zn + CuSO4 --> ZnSO4 + Cu

a) single replacement c) decomposition

b) double replacement d) synthesis

19) Which chemical equation best represent a decomposition reaction?

a) Cl2 + 2KI 2KCl + I2 c) H2CO3 H2O + CO2

b) 2Al + 3Cl2 2AlCl3 d) KCl + AgNO3 KNO3 + AgCl

20) What is the molecular formula of a compound that has a molecular mass of 54 and an

empirical formula of C2H3?

a) C8H12 b) C6H9 c) C4H6 d) C2H3

21) What is the empirical formula of the compound whose molecular formula is P4O10?

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

22) Which of the following is an empirical formula?

a) H2O2 b) H2O c) C2H2 d) C4H8

23) Which represents both an empirical and molecular formula?

a) P2O5 b) C3H6 c) N2O4 d) C6H12O6

For questions 24 and 25, show all work and express your answer in the appropriate units.

24) Calculate the gram formula mass of ZnSO4.

25) Use your answer from 23 to calculate the percent by mass of zinc in ZnSO4.

26) Balance the following reaction and reduce to the lowest whole number coefficients

_____ H2SO4 + _____ B(OH)3 _____ B2(SO4)3 + _____ H2O

56

27) Li and KNO3 according to the following equation:

Li + KNO3 LiNO3 + X

Write the formula for the missing product X.

Use the chemical equation below to answer questions 28-30.

___ N2 + ___ H2 ___ NH3

27) Balance the equation above using the lowest whole number coefficients.

29) How many moles of H2 are required to produce 6.5 moles of NH3? Show all work and make

sure your answer has the correct number of significant figures and proper units.

30) How many grams of NH3 are produced if 50.0 grams of N2 are consumed? Show all work and

make sure your answer has the correct number of significant figures and proper units.

57

unit 5: moles and chemical reactions notes unit · pdf file3 mini poster project- classifying...

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