Name: Honors Unit 3 Holey Moley - WordPress.com · 11/3/2019 · 6. How many legs are found in one mole of elephants? 7. How many carbon atoms are found in one mole of methane molecules?

Honors Unit 3

Holey Moley Topics/ Daily Outline:

# A Day B Day Content: TEXT: CW #: HW #:

1 11/20 11/21 Quiz 1, Moles 10.1, 10.2 1 -- 2 11/22 11/25 Empirical and molecular formulas 10.3 2 1

3 11/26 12/2 Composition of a hydrate 15.2 3 --

4 12/3 12/4 Chemical equations 11.1 4, 5 2 5 12/5 12/6 Types of reactions 11.2 6 3

6 12/9 12/10 Single replacement reactions 11.2, 11.3 7 -- 7 12/11 12/12 Double replacement reactions 11.3 8, 9 4

8 12/13 12/16 Quiz 2, CheMystery 11.3 10 --

9 12/17 12/18 Stoichiometry 12.2 11, 12 5 10 12/19 12/20 Ornament lab -- 13 --

11 1/2 1/3 Molarity 16.2 14 6 12 1/6 1/7 Limiting reactants 12.3 15 7

13 1/8 1/9 Limiting reactants 12.3 16 -- 14 1/10 1/13 Quiz 3, Gravimetric analysis -- 17 8

15 1/14 1/15 Gravimetric analysis -- 17 9

16 1/16 1/17 Gravimetric analysis -- 17 -- 17 1/21 1/22 Review -- -- --

18 1/23 1/24 Unit test -- -- 10 19 1/27 1/28 Dimensional analysis review -- -- --

20 1/29 1/31 Dimensional analysis test -- -- --

Important Due Dates:

• SciResearch: 8 Collect Data, 12/2 (A Day) and 12/3 (B Day)

• SciResearch: 9 Analyze Data and Draw Conclusions, 12/2 (A Day) and 12/3 (B Day)

• SciResearch: 10 Communicate Results, 12/9 (A Day) and 12/10 (B Day)

• Composition of a Hydrate Lab Report, 12/11 (A Day) and 12/12 (B Day)

• Science Fair Competition: after school in the media center on 12/12

• Gravimetric Analysis Scientific Poster, due 1/27 (A Day) and 1/28 (B Day)

For tutorials and additional resources: www.leffellabs.com If you are absent, please use this sheet to determine what you missed and collect the materials from the make-up work bins up front. Get help from a friend, the link above, or the instructor.

Name:

Class:

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Homework:

• HW 1: The Strange Case of Mole Airlines

• HW 2: Chemical Equations I

• HW 3: Chemical Equations II

• HW 4: Chemical Equations III

• HW 5: Stoichiometry Calculations I

• HW 6: Stoichiometry Calculations II

• HW 7: Stoichiometry Calculations III

• HW 8: Limiting Reactants

• HW 9: Review for Unit Test

• HW 10: Dimensional Analysis Review

Quizzes:

• Quiz 1: Dimensional Analysis

• Quiz 2: CW 1 to CW 9

• Quiz 3: CW 10 to CW 16

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CW 1: The Mole

The Elephant and the Methane Molecule

One Elephant has one trunk and four legs One methane molecule, CH4, contains one carbon atom and four hydrogen atoms

1 dozen = 12 items 1 mole = 6.02x1023 items = Avogadro ’s Number

1. How many trunks are found in one dozen elephants?

2. How many legs are found in one dozen elephants?

3. How many carbon atoms are found in one dozen methane (CH4) molecules?

4. How many hydrogen atoms are found in one dozen methane molecules?

5. How many trunks are found in one mole of elephants?

6. How many legs are found in one mole of elephants?

7. How many carbon atoms are found in one mole of methane molecules?

8. How many hydrogen atoms are found in one mole of methane molecules?

9. How is “a mole” similar to “a dozen”?

10. A mole is equal to 6.02x1023 items, which is a very large number. Why would chemists want to use moles as the unit to count atoms in?

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The Mole Road Map

11. Find the molar mass of the following compounds. a. Sulfur Dioxide

b. Lead(II) Nitrate

c. Phosphoric Acid

d. Ammonium Sulfate


12. Convert from 1.56x1030 particles of sodium chloride to grams of sodium chloride (use the mole road map).

13. How many oxygen molecules are in 3.36 L of oxygen gas at STP? (Answer: 9.03x1022 molecules)

14. Find the mass in grams of 2.00x1023 molecules of F2. (Answer:12.6 g)

15. Determine the volume in liters occupied by 14 g of nitrogen gas at STP. (Answer: 11.2 L)

16. Find the mass, in grams, of 1.00x1023 molecules of N2. (Answer: 4.65 g)

1.56x1030

particles NaCl


CW 2: Empirical and Molecular Formulas

Percent Composition

% 𝐶𝑜𝑚𝑝𝑜𝑠𝑖𝑡𝑖𝑜𝑛 𝑜𝑓 𝐸𝑙𝑒𝑚𝑒𝑛𝑡 𝑖 = 𝑀𝑜𝑙𝑎𝑟 𝑀𝑎𝑠𝑠 𝑜𝑓 𝐸𝑙𝑒𝑚𝑒𝑛𝑡 𝑖𝑛 𝐶𝑜𝑚𝑝𝑜𝑢𝑛𝑑

𝑀𝑜𝑙𝑎𝑟 𝑀𝑎𝑠𝑠 𝑜𝑓 𝑡ℎ𝑒 𝐶𝑜𝑚𝑝𝑜𝑢𝑛𝑑× 100%

Table 1: Percent composition (by mass) of some common organic molecules


1. Verify that the % composition given for ethyne in Table 1 is correct. Molar mass of C2H2: Molar mass of carbon in C2H2:

% Composition Carbon:

Molar Mass of hydrogen in C2H2:

% Composition Hydrogen:

2. Fill in the missing molecular formulas and % compositions in Table 1.

3. Can you determine the % composition by mass of H for 2-butene without using the

equation? If so, how?

4. Agree or disagree: compounds with the same percent composition have the same molecular formula. Explain.

5. The molecule 2-hexene has the molecular formula C6H12. Without doing calculations, what is the percent composition of H in this molecule?


Empirical Formula

The empirical formula of a compound describes the relative number of each type of atom in the compound. It is given in terms of the smallest possible whole number ratios (as subscripts). For example, the empirical formula of ethane, C2H6, is CH3. (Note that the subscript "1" is omitted.)

6. Determine the empirical formula of each of the molecules in Table 1.

7. A molecule containing only nitrogen and oxygen contains (by mass) 36.8% N. a. How many grams of N would be found in a 100 g sample of the compound?

b. How many grams of O would be found in the same sample?

c. How many moles of N would be found in a 100 g sample of the compound?

d. How many moles of O would be found in the same sample?

e. What is the empirical formula of the compound?


8. A compound used as a dry-cleaning fluid was analyzed and found to contain 18.00% C, 2.27% H, and 79.73% Cl. Determine the empirical formula of the fluid.

9. A compound was analyzed and found to contain 13.5 g Ca, 10.8 g O, and 0.675 g

H. What is the empirical formula of the compound?


Molecular Formula

The empirical formulas we have calculated in the preceding section express the simplest atomic ratio between the elements in the compound. An empirical formula does not necessarily represent the actual numbers of atoms present in a molecule of a compound; it represents only the ratio between those numbers. The molecular formula of a compound may be the empirical formula, or it may be a multiple of the empirical formula.

10. The empirical formula of a compound is NO2. Its molecular mass is 92 g/mol. What is its molecular formula? a. Determine the molar mass of NO2.

b. Divide the molecular mass by the molar mass. You should get a whole number.

c. Multiply the subscripts in NO2 by the whole number you got in the last step.

11. The empirical formula of a compound is CH2. Its molecular mass is 70 g/mol. What is its molecular formula?

12. A compound is found to be 40.0% carbon, 6.7% hydrogen and 53.5% oxygen. Its molecular mass is 60 g/mol.

a. What is the empirical formula?

b. What is the molecular formula?


CW 3: Composition of Hydrates

A hydrate is an ionic compound (salt) with water molecules loosely bonded to its crystal structure. The water is in a specific ratio to each formula unit of the salt. For example, the formula Na2S·9H2O indicates that there are nine water molecules for every one formula unit of Na2S.

Class Discussion Questions

1. A 5.0 g sample of a hydrate (Cu(NO3)2 · ?H2O) was heated to constant mass, leaving 3.9 g of the anhydrous Cu(NO3)2 behind.

a. What percentage of water was in the hydrate?

b. Determine how many moles of Cu(NO3)2 are present.

c. Determine how many moles of H2O were present before heating.

d. Determine the formula of the hydrate (like finding the empirical formula, divide each of the answers in B and C by the smallest number of moles to get whole numbers).


2. A group of students wanted to determine the formula of a hydrate of MgCO3. They measured the mass of an empty crucible, then added some of the hydrate. They heated the sample until the mass of the sample remained constant.

Mass empty crucible 50.43 g

Mass of crucible with sample 66.10 g

Mass of crucible with sample after heating 58.01 g

a. How many grams of the hydrate did they add to the crucible?

b. How many grams of water did they remove from the hydrate?

c. What is the mass of the anhydrous compound?

d. What is the percent water that was driven off from the sample?

e. What is the formula of the hydrate?

f. Given that the actual formula of the hydrate is (MgCO3 · 5H2O), what is the actual percentage of water?


3. Watch the video found here: http://bit.ly/2A5JSFb a. How can the solid CuSO4 appear to be dry, but contain water molecules?

b. What is meant by an anhydrous salt? How will we drive away the water molecules?

c. Draw a diagram of the heating set up.

d. In the video, the sample was only heated once and massed once. You will heat to constant mass. What does this mean and how does this let you know when you are done heating the hydrate?

e. Why do you need the record the mass of the empty crucible?


Procedure

You will be given an unknown hydrate of copper(II) sulfate. Explain the experimental steps (lab procedure – heating to constant mass) you would need to complete to determine the formula of the hydrate. Then explain how you would use the data you collected to solve for the formula of the hydrate (calculations). Based on your answers to the Class Discussion Questions, write a procedure for how to determine the amount of water in a hydrate of copper(II) sulfate by heating to constant mass. Be sure to include the following:

• Safety precautions

• Heating to constant mass

• How to determine the empirical formula of the hydrate from its mass and the mass of the water removed

• A ready to fill in data table with a place to record all required data and observations


Formal Lab Report

Your final lab report should be typed. The following sections should be completed, in order, as they appear below. COVER SHEET: A cover sheet with nothing else but:

• Title of the lab: a short, descriptive, title that tells the reader what the lab is about.

• Your name

• Your partners’ first and last names

• Due date of the lab report

• Class period PURPOSE: What are we trying to determine/ do in this experiment? PROCEDURE: Include a detailed step by step procedure for setting up the experiment and collecting data over the course of the experiment. Underline any materials you will need once you have written the procedure. DATA: Organize ALL data into a neat data table. This means you will need in depth observations. Include:

• All masses required for calculations

• Observations over the course of the experiment CONCLUSION: Answer the following questions using complete sentences. You will be graded on the quality, completeness, correctness, and style of your writing.

1. Perform the following calculations NEATLY SHOWING ALL WORK: a. Calculate the mass of the hydrate that was added to the crucible (before heating). b. Calculate the mass of the water (H2O) that was driven off. c. Calculate the mass of the anhydrous sample (dry CuSO4) after heating. d. Find the percentage of water (H2O) that was driven off from the hydrate. e. Determine the formula of the hydrate based on your data.

2. The actual formula of the hydrate is CuSO4 · 5H2O. Based on this formula and using molar masses, calculate the actual percentage of water.

3. Calculate the percent error in the percentage of water, using the experimental value you found in Question 1d and the actual value your found in Question 2.

4. Discuss possible sources of error and how they affected your data and calculations.

NOTE: “plugging in numbers wrong” or miscalculations are NOT sources of error, as you can easily go back and fix these. Think about the data you collected and what could have gone wrong with each of those measurements.

% 𝐸𝑟𝑟𝑜𝑟 = |(𝐴𝑐𝑐𝑒𝑝𝑡𝑒𝑑 𝑉𝑎𝑙𝑢𝑒 − 𝐸𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙 𝑉𝑎𝑙𝑢𝑒)

𝐴𝑐𝑐𝑒𝑝𝑡𝑒𝑑 𝑉𝑎𝑙𝑢𝑒| × 100%


Grading Rubric

Cover Page

• Title of the lab

• Your name

• Your partners’ names

• Due date of the lab report

• Class period

/5

Purpose

• Clear statement about what we set out to do with this lab

/5

Procedure

• Repeatable, clear

• Paragraph form

• Materials are underlined in procedure

• Complete, no steps skipped or assumed

/10

Data

• Table is neat, organized, readable, complete

• Reflects student understanding of lab concepts and practices

• Quality/ completeness of observations

/5

Conclusion

• Question 1: 10 points




/25

TOTAL:

/50


CW 4: Balancing Chemical Equations

Atoms are neither created nor destroyed when chemical reactions take place. Therefore, the number of atoms of each element must be identical on the reactant (left) and product (right) sides of a balanced chemical reaction. Such a chemical equation is said to be atom balanced.

Two Balanced Chemical Equations

Two balanced chemical reactions (or chemical equations) are given below: I2(g) + H2(g) → 2HI(g) (1) 2CO(g) + O2(g) → 2CO2(g) (2)

1. Indicate the reactants and products for each reaction in the table below:

Reaction Reactant(s) Product(s)

(1)

(2)

2. What does the arrow represent in a chemical reaction?

3. For reaction (1), how many H atoms and I atoms are represented on: a. The reactant side?

b. The product side?

4. For reaction (2), how many C atoms and O atoms are represented on: a. The reactant side?

b. The product side?

5. Based on your answers to questions 3 and 4, what general statement can be made about the number of atoms of each type on the two sides of a chemical equation?


The Balanced Chemical Reaction

A balanced chemical reaction can be interpreted in two ways.

• First, it can be thought of as describing how many molecules of reactants are consumed to produce a certain number of molecules of products.

• Analogously, it can be thought of as describing how many moles of reactants are consumed to produce the indicated number of moles of products. A mole is equal to 6.02x1023 items.

I2(g) + H2(g) → 2HI(g) (1) 2CO(g) + O2(g) → 2CO2(g) (2)

6. For reaction (2): a. How many CO2 molecules are produced for every O2 molecule consumed?

b. How many CO2 molecules are produced for every CO molecule consumed?

c. How many moles of CO2 are produced when 2 moles of O2 are consumed?

d. How many moles of CO2 are produced when 5 moles of O2 are consumed?

7. How many moles of I2 react to produce 12 moles of HI in reaction (1)?

8. Explain your findings in the previous question with the idea that atoms are neither created nor destroyed when chemical reactions take place (Law of Conservation of Mass).


Balancing Chemical Equations

9. Balance the following chemical equations by inspection or the inventory method. a. ___ SiC + ___ Cl2 → ___ SiCl4 + ___ C

b. ___ KOH + ___ H2SO4 → ___ K2SO4 + ___ H2O

c. ___ N2 + ___ O2 → ___ N2O

d. ___ Li + ___Cl2 → ___LiCl

e. ___ CaSO4 + ___ AlCl3 → ___ Al2(SO4)3 + ___ CaCl2


CW 5: Writing Chemical Equations

State Symbols

Symbol Meaning Look For

s Solid • Solids: Powders, metals

• Precipitates: solids formed when two solutions are mixed l Liquid • Water, molecular compounds that are liquids

g Gas • Bubbles forming

aq Aqueous • Dissolved in water, clear (not cloudy) solutions

Aqueous or Solid?

1. Write the formula for each of the following, then use the solubility guidelines to determine if each compound would be aqueous or solid in solution.

a. Sodium nitrate

b. Silver chloride

c. Barium sulfate

d. Ammonium Carbonate

Solubility Guidelines for Aqueous Solutions Ions that form Soluble Compounds

Exceptions Ions that form Insoluble Compounds

Exceptions

Group 1 ions (Li+, Na+, etc.)

Carbonate (CO32–)

When combined with Group 1 ions or ammonium

Ammonium (NH4+) Chromate (CrO4

2–) When combined with Group 1 ions, Ca2+, Mg2+ or ammonium

Nitrate (NO3–) Phosphate (PO4

3–) When combined with Group 1 ions or ammonium

Acetate (C2H3O2– or

CH3COOH) Sulfide (S2–)


Hydrogen carbonate (HCO3

–) Hydroxide (OH–)

When combined with Group 1 ions, Ca2+, Ba2+, Sr2+ or ammonium

Chlorate (ClO3–)

Perchlorate (ClO4–)

Halides (Cl–, Br–, I–) When combined with Ag+, Pb2+, and Hg2

2+

Sulfates (SO42–)

When combined with Ag+, Ca2+, Sr2+, Ba2+, and Pb2+


Chemical Reactions

Observe the following reactions. Then, use the words to write a balanced chemical equation. Include the use of state symbols on all reagents.

2. Magnesium burns in the presence of oxygen to form magnesium oxide.

3. An electric current is run along a wire submerged in water. Oxygen gas and hydrogen gas are formed.

4. Hydrochloric acid reacts with magnesium metal, forming magnesium chloride and hydrogen gas.

5. Hydrogen gas reacts explosively with oxygen gas to form water vapor.

6. Copper(II) sulfate solution reacts with sodium hydroxide solution to form copper(II) hydroxide and sodium sulfate.


CW 6: Types of Chemical Reactions

Directions

1. Complete your assigned activity. 2. Afterwards, complete the following by discussing in your groups. 3. Return to CW 5 and classify the reactions you wrote.

Combination (Synthesis) General Equation: Reactants:

Example: Products:

Decomposition

General Equation: Reactants:

Example: Products:


Combustion General Equation: Reactants:

Example: Products:

Single Replacement General Equation: Reactants:

Example: Products:

Double Replacement General Equation: Reactants:

Example: Products:


CW 7: Single Replacement: Metal Activity Series

The data below was collected by observing reactions between each metal and each sulfate salt of the metal. It was recorded if a reaction occurred or not.

magnesium

sulfate zinc sulfate iron sulfate copper sulfate

magnesium reacts reacts reacts

zinc no reaction reacts reacts

iron no reaction no reaction reacts

copper no reaction no reaction no reaction

2. Which is the most reactive metal? 3. Which is the least reactive metal? 4. Put the four metals in order of reactivity: ____________________ (Most reactive) ____________________ ____________________ ____________________ (Least reactive)

You have just created an activity series, an empirical tool created from observing several chemical reactions. It can be used to predict the products of single replacement reactions. 5. Compare the activity series you made in question 3 to the more complete activity series on

the next page. Does it match up?


In a single replacement reaction, atoms of one element replace the atoms of a second element in a compound. The activity series of metals helps us to determine if one metal will replace another metal in a compound. A reactive metal will replace any metal listed below it in the activity series.

Activity Series of Metals

Dec

reas

ing

Act

ivit

y

Name Symbol Will replace H from acids and water Will replace H from acids only

Lithium Li Rubidium Rb

Potassium K Calcium Ca

Sodium Na

Magnesium Mg Aluminum Al

Manganese Mn Zinc Zn

Iron Fe

Nickel Ni Tin Sn

Lead Pb (Hydrogen) (H)

Copper Cu Silver Ag

Platinum Pt

Gold Au 6. Magnesium metal can be used to remove tarnish from silver items. Silver tarnish is the

corrosion that occurs when silver metal reacts with substances in the environment, especially those containing sulfur. Why would magnesium remove tarnish from silver?

7. Use the activity series for metals to explain why copper metal is used in plumbing where

the water might contain compounds of many different metals.


8. For each of the following, determine if the lone element is a metal or a halogen. Then determine if a reaction will occur. If no reaction occurs, write NR.

Reactants Lone

Element Element It Replaces

Does reaction occur?

Mg(s) + LiNO3(aq) Mg Li NR: Mg is below Li on activity series, so it cannot

replace Li.

Fe(s) + CuSO4(aq)

Mg(s) + HCl(aq)

Ag(s) + CuCl2(aq)

Activity Series of Metals Virtual Lab

Access the virtual lab from unit 3 page on LEFFELlabs. Make Your Data Table: During the PowerPoint, record if a reaction occurs or not.

A(NO3) B(NO3) C(NO3) D(NO3) E(NO3)

A

B

C

D

E


9. Explain (using evidence from the table you created) which metal is the most reactive. 10. Based on your data table, create an activity series, going from most reactive at the top to

least reactive at the bottom, for the metals in a single displacement reaction. 11. Write the chemical equation for the following reactions (remember – only some are

reactive enough for a single displacement to occur). a. D + BNO3

b. C + DNO3

c. E + ANO3

d. A + ENO3


CW 8: Double Replacement: Eye on the Ions

Observe the demonstration while answering the following. 1. Draw a diagram of each of the following (you must include color!):

a. The “compounds”

b. The two beakers, after the compounds had been added and the solution formed

c. After the solutions were combined

2. Use picture 1b to explain what happens when the compounds are first added to water and form solutions.

3. When the solutions are combined, which of the colored ball(s) represents the precipitate?


4. Aqueous Silver nitrate reacts with aqueous sodium chloride to form silver chloride and sodium nitrate.

a. Write a balanced chemical equation for this reaction.

b. Use your diagrams from above to determine which colored ball represented each of the following ions.

i. Ag+1

ii. Na+1

iii. NO3-1

iv. Cl-1


CW 9: Double Replacement: Net Ionic Equations

Writing Net Ionic Equations

Many important chemical reactions take place in water, or aqueous solution, such as below. AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq) (1) The reactants dissociate (separate) into cations and anions when they dissolve in water. This is shown in the equation (2), which is called a complete ionic equation. Ag+(aq) + NO3

-(aq) + Na+(aq) + Cl-(aq) → AgCl(s) + Na+(aq) + NO3-(aq) (2)

As you can see in equation (2), the nitrate ion and the sodium ion appear unchanged on both sides of the yield arrow. The equation may be simplified by eliminating these spectator ions, which are not directly involved with the reaction. Ag+(aq) + NO3

-(aq) + Na+(aq) + Cl-(aq) → AgCl(s) + Na+(aq) + NO3-(aq) (2)

Once the spectator ions have been removed, the remaining equation is called a net ionic equation, as shown in (3). Ag+(aq) + Cl-(aq) → AgCl(s) (3)

1. For the following, cross out the spectator ions, then write the net ionic equation. a. Ni2+(aq) + 2NO3

-(aq) + 2Na+(aq) + 2OH-(aq) → Ni(OH)2(s) + 2Na+(aq) + 2NO3-(aq)

b. Fe(s) + Cu2+(aq) + 2NO3-(aq) → Fe2+(aq) + 2NO3

-(aq) + Cu(s)

c. Na+(aq) + NO3-(aq) + K+(aq) + Cl-(aq) → K+(aq) + NO3

-(aq) + Na+(aq)+ Cl-(aq)


2. What unusual condition did you encounter in question 1c?

This is an example of a reaction that does not occur – the ions simply float around together in solution.

3. Based on the examples in question 1, which type of chemical reaction will you be

required to write net ionic equations for?

Predicting the Formation of a Precipitate Whether or not a precipitate will form depends upon the solubility of the new compounds that form. By using the solubility rules for ionic compounds, you can predict the formation of a precipitate.



Exceptions


Carbonate (CO32–)












Chlorate (ClO3–)



2+

Sulfates (SO42–)



4. Write the formula for each of the following compounds, then determine if they are soluble (aq) or insoluble (s) using the solubility rules.

a. Calcium carbonate

b. Strontium hydroxide

c. Calcium sulfate

d. Potassium nitrate

e. Iron(II) nitrate

f. Lead(IV) carbonate

5. For each of the following double replacement reactions:

• Predict the products by swapping cations and balancing charges

• Determine if the products are soluble or insoluble

• Balance and include all state symbols

• Write the net ionic equation

• If there is no reaction (see question 2), write NR

a. ___Pb(NO3)2 (aq) + ___Na2SO4 (aq) → Net ionic equation:

b. ___NaClO3 (aq) + ___KCl (aq) → Net ionic equation:


Formation of a Gas or a Liquid

When writing out net ionic equations, it is important to note that liquids and gases do not break apart into ions, just like precipitates (solids). Balance and write net ionic equations for the following:

6. ___Ca(OH)2(aq) + ___H3PO4(aq) → ___H2O(l) + ___Ca3(PO4 )2(s)

7. ___HCl(aq) + ___Na2S(aq) → ___H2S(g) + ___NaCl(aq)

8. ___HCl(aq) + ___NaHCO3(aq) → ___NaCl(aq) + ___H2O(l) + ___CO2(g)


CW 10: CheMystery

Objective

To identity between six unknown aqueous solutions, based on their reactions.

Materials

• Transparency sheet matrix

• Paper towels

• Unknown solutions

• Solubility chart/rules

Procedure

1. Each group will receive six coded dropper bottles, each containing several milliliters of an aqueous solution of a single unknown substance. Each bottle will contain a different solution. The team will have a maximum of thirty minutes to make observations on the individual solutions, and on drops mixed in pairs, to identify the substances. No substances other than the solutions in the bottles may be used.

2. The possible substances are solutions of: HCl, NaOH, Ba(NO3)2, Na2CO3, Cu(NO3)2, and Pb(NO3)2.

3. A transparent plastic sheet matrix will be provided. For each solution, place one drop, mixing the solutions in pairs. The group must carry out all tests on the plastic sheet provided. Testing the solutions by touching or smelling is not allowed.

Conclusion

Identify each solution by letter:

Cu(NO3)2

NaOH

Pb(NO3)2

HCl

Na2CO3

Ba(NO3)2

Note: Each time you try to identify the solutions, you will receive points according to the following scheme:

• First try: 16 % for each correct identification

• Second try: 8 % for each correct identification

• Third try: 4 % for each correct identification


Scaffolding Sheet

1. Determine the charges for each of the ions in the chemical formulas given below. 2. Use the table below to determine the products of the following double replacement

reactions. Use the solubility rules to determine if a solid compound is formed. a. If both products are aqueous, no reaction occurs. b. If H2CO3 is formed, it will decompose to CO2 and H2O, producing bubbles.

• Cu (NO3)2 + Na OH →

• Cu (NO3)2 + Pb (NO3)2 →

• Cu (NO3)2 + H Cl →

• Cu (NO3)2 + Na2 CO3 →

• Cu (NO3)2 + Ba (NO3)2 →

• Pb (NO3)2 + Na OH →

• Pb (NO3)2 + H Cl →

• Pb (NO3)2 + Na2 CO3 →

• Pb (NO3)2 + Ba (NO3)2 →

• Na2 CO3 + Na OH →

• Na2 CO3 + H Cl →

• Na2 CO3 + Ba (NO3)2 →

• Na OH + H Cl →

• Na OH + Ba (NO3)2 →

• H Cl + Ba (NO3)2 →

3. Complete the chart below to assist you in identifying the solutions.

Cu(NO3)2 NaOH Pb(NO3)2 HCl Na2CO3 Ba(NO3)2 Cu(NO3)2

NR -- -- -- -- --

NaOH NR -- -- -- -- Pb(NO3)2 NR -- -- --

HCl NR -- -- Na2CO3

NR --

Ba(NO3)2 NR


CW 11: Mole-Mole Stoichiometry

Chemical Equations

A balanced chemical reaction can be interpreted in two ways.

• First, it can be thought of as describing how many molecules of reactants are consumed to produce a certain number of molecules of products.

• Analogously, it can be thought of as describing how many moles of reactants are consumed to produce the indicated number of moles of products.

1. Balance the chemical equation below. ___H2(g) + ___O2(g) → ___H2O(g) (1)

The Mole Ratio A mole ratio is the ratio between the amounts in moles of any two compounds involved in a chemical reaction. Mole ratios are used as conversion factors between products and reactants in many chemistry problems.

2. Write all possible mole ratios for reaction (1). There should be six total, including inverses.

2 𝑚𝑜𝑙 𝐻2

1 𝑚𝑜𝑙 𝑂2

1 𝑚𝑜𝑙 𝑂2

2 𝑚𝑜𝑙 𝐻2

3. Using you answers to the previous question, fill in the blanks with the correct mole ratio

or the solution to the problem. Use slashes to show units that cancel.

a. 3 𝑚𝑜𝑙 𝐻2

1 ×

= 1.5 𝑚𝑜𝑙 𝑂2

b. 2.3 𝑚𝑜𝑙 𝑂2

1 ×

= 4.6 𝑚𝑜𝑙 𝐻2𝑂

c. 1.7 𝑚𝑜𝑙 𝐻2𝑂

1 ×

= ________ 𝑚𝑜𝑙 𝐻2

d. 8.7 𝑚𝑜𝑙 𝐻2

1 ×

= _________ 𝑚𝑜𝑙 𝐻2𝑂


4. How is the mole ratio used like a conversion factor? What units does it allow you to convert between?

5. Answer the following using this equation: ___H2 + ___O2 → ___H2O a. What is the mole ratio of H2 to H2O?

b. Suppose you had 20 moles of H2, how many moles of H2O could you make?

c. What is the O2 / H2O molar ratio?

d. Suppose you had 20 moles of O2, and enough H2, how many moles of H2O could you make?

6. Use this equation to answer following: ___N2 + ___H2 → ___NH3 a. If you used 1 mole of N2, how many moles of NH3 could be produced?

b. If 10 moles of NH3 were produced, how many moles of N2 would you need?

c. If 3.00 moles of H2 were used, how many moles of NH3 would be made?

d. If 0.600 moles of NH3 were produced, how many moles of H2 are required?


CW 12: Multistep Stoichiometry

Stoichiometry Road Map

____ Li + ____ AlCl3 → ____ Al + ____ LiCl

1. How many grams of Li are needed to produce 10 grams of LiCl? a. Write out a flow map for this problem.

b. Solve.

2. How many grams of Al are consumed if 30.5 grams of LiCl are produced? a. Write out a flow map for this problem.

b. Solve.

Particles Given

Mass Given

Volume (liters) @ STP Given

Particles Wanted

Mass Wanted

Volume (liters) @

STP Wanted

Moles Given

Moles Wanted


3. Write a chemical equation for the reaction of tin(II) oxide with nitrogen trifluoride to produce tin(II) fluoride and dinitrogen trioxide.

4. How many grams of NF3 are required to fully react 36.52 grams of SnO?

5. How many particles of N2O3 are produced from 51.3 grams of SnO?

6. How many moles of SnF2 are produced from 17.89 grams of N2O3?


Theoretical Yield

The theoretical yield is the calculated amount of product that can be produced from a given amount of reactants, if everything in the lab goes perfectly. Lab work is never perfect, resulting in an experimental yield that is less than the calculated theoretical yield.

7. Find the theoretical yield of FeCl3(s) in grams for each of the following.

___Fe(s) + ___Cl2(g) → ___FeCl3(s)

a. What is the theoretical yield in grams of FeCl3 if 3 L of Cl2 gas are used?

b. How many grams FeCl3 can be made from 15.0 grams of iron? Because errors always occur during chemistry labs, the experimental yield of a product is always less than the theoretical (calculated) yield. We express this using percent yield. The highest possible percent yield is 100%, but in practice this will never occur.

𝑃𝑒𝑟𝑐𝑒𝑛𝑡 𝑌𝑖𝑒𝑙𝑑 = 𝐸𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙 𝑌𝑖𝑒𝑙𝑑

𝑇ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 (𝐶𝑎𝑙𝑐𝑢𝑙𝑎𝑡𝑒𝑑) 𝑌𝑖𝑒𝑙𝑑 × 100%

8. Using the theoretical yields from question 6 and the experimental yields given below,

determine the percent yield for each of the problems in question 6. a. Theoretical Yield: (See answer 6a above) g FeCl3 | Experimental Yield: 13.20 g

FeCl3

b. Theoretical Yield: (See answer 6b above) g FeCl3 / Experimental Yield: 40.56 g FeCl3


CW 13: Holiday Ornament Lab

Objective

To create a last minute gift using chemistry.

Materials

• Acidified copper(II) nitrate solution beaker

• Galvanized (zinc-coated) iron, 3” square piece with hole

• Hydrochloric acid solution beaker

• Masking tape

• Pencil

• Scalpel

• Paper towels

• Tongs

• Eraser

• Cotton swab

• Ornament hanger

Safety Precautions

• Hydrochloric acid solution is corrosive to skin and eyes and is moderately toxic by ingestion and inhalation.

• Acidified copper(II) nitrate solution is slightly toxic by ingestion and is a skin, eye, and mucous membrane irritant.

• The edges of the galvanized iron and the scalpel are sharp – be careful to avoid cuts.

• Wear chemical splash goggles and chemical-resistant gloves.

Tips

• Read the whole procedure first, if you mess up, you don’t get another one.

• Keep the designs relatively simple, as simple designs tend to turn out best. Different designs can be drawn on each side of the ornament if desired.

• Note where the hole is located so that they orient the design correctly. If you prepare the ornament upside down or sideways, a new hole can be drilled in the top.

• Acrylic sealer may be applied to the ornaments to slow the tarnishing of the copper. Get an adult to help you apply acrylic sealer evenly to both sides of the ornament. Hang the ornament to allow the sealer to dry completely.


Procedure

1. Completely cover both sides of a piece of galvanized iron with masking tape. Mark where the hole is – this is the top of the ornament. Make sure that the edges of the galvanized iron are also covered.

2. Draw a simple design (such as your initials) on the masking tape with a pencil. Designs may be drawn on both sides of the piece of galvanized iron if desired. The design that is drawn will become the copper-colored part of the ornament.

3. Use a scalpel to gently cut along the pencil marks. Remove the masking tape inside the drawing only so that the design is uncovered.

4. Wash your ornament briefly under a faucet. Dry the ornament as best as possible. Mass the ornament and record.

5. Bring your ornament to the hydrochloric acid beaker in the fume hood. Use tongs to lay the ornament flat on the bottom of the beaker so that it is completely submerged.

6. As soon as the rapid bubbling stops, remove the ornament using the tongs. Rinse the ornament with tap water and dry it with a paper towel.

7. Carefully clean the exposed area of the design by rubbing it with an eraser, dusting off the bits of eraser left behind. Do not remove the masking tape. Mass the ornament and record.

8. Use the copper(II) nitrate solution and a cotton swab to “paint” the exposed metal. 9. Once the entire design area is coated with copper, rinse the ornament with tap water

and dry it with a paper towel. Mass the ornament. 10. Remove the masking tape from both sides of the piece of iron. Attach a hanger to the

hole in the top of the ornament. 11. Dispose of wastes as directed.

Data Table

Step Quantity Value

4 Mass of galvanized iron after design is cut out (with tape)

7 Mass of galvanized iron after erasing (with tape)

9 Mass of galvanized iron after painting (with tape)

Analysis and Conclusion

Complete the Ornament Lab Analysis and Conclusion worksheet.


CW 14: Molarity Calculations

What is Molarity?

The concentration of a solute in an aqueous solution can be expressed in many ways: grams of solute per liter of solution; grams of solute per 1000 grams of water; moles of solute per 1000 grams of water; and so on. One of the most frequently used concentration units is molarity.

𝑀𝑜𝑙𝑎𝑟𝑖𝑡𝑦 𝑜𝑓 𝑆𝑜𝑙𝑢𝑡𝑒 𝑖 = 𝑀𝑜𝑙𝑒𝑠 𝑜𝑓 𝑖

𝑉𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑆𝑜𝑙𝑢𝑡𝑖𝑜𝑛 𝑖𝑛 𝐿𝑖𝑡𝑒𝑟𝑠

The unit for molarity is 𝑚𝑜𝑙𝑒𝑠

𝑙𝑖𝑡𝑒𝑟 and is represented by the symbol M (pronounced “molar”).

1. When one mole of Na2SO4 is dissolved in water:

a. How many moles of sodium ions are found in the solution?

b. How many moles of sulfate ions are found in the solution?

2. Verify that when 10.0 g of sodium sulfate dissolves in water: a. There are 0.0704 moles of sodium sulfate in the water.

b. There are 0.0704 moles of sulfate in the water.

c. There are 0.141 moles of sodium in the water.


3. Given that the solution has a total volume of 250.0 mL, what is the concentration of a. Sodium sulfate?

b. Sulfate ions?

c. Sodium ions?

4. You have 1.0 L of 0.50 M NaCl and 1.0 L of 0.30 M Na2SO4. Which is more concentrated with respect to sodium ions?

𝑀𝑜𝑙𝑎𝑟𝑖𝑡𝑦 𝑜𝑓 𝑆𝑜𝑙𝑢𝑡𝑒 𝑖 = 𝑀𝑜𝑙𝑒𝑠 𝑜𝑓 𝑖

𝑉𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑆𝑜𝑙𝑢𝑡𝑖𝑜𝑛 𝑖𝑛 𝐿𝑖𝑡𝑒𝑟𝑠


Making a 0.5 Molar Solution

5. Your task is to make 50 mL of a 1.0 M solution of Na2CO3. a. What is the equation for molarity? What units should the volume be in?

b. Using the molarity equation, solve for the moles of Na2CO3 required for 50 mL of a 1.0 M solution.

c. Use molar mass to calculate the grams of Na2CO3 needed to make your solution.

d. Use a volumetric flask to make your solution.

6. Consider the measurements you made. What are some possible sources of error that may cause the concentration of your solutions to be “off”?

7. Which of the following has more moles of Na2CO3: 0.15 L of 1.9 M Na2CO3 solution, or 0.25 L of 1.0 M Na2CO3 solution?


CW 15: Limiting Reactant: S’mores

A S’more is constructed with the following ingredients and amounts: 1 teddy graham cracker 1 mini chocolate bar 2 mini marshmallows At the store, these items can only be obtained in full boxes, each of which contains one gross of items. A gross is a specific number of items (like a mole). The boxes of items have the following net weights (the weight of the material inside the box): Box of teddy graham crackers 9.0 pounds Box of mini chocolate bars 36.0 pounds Box of mini marshmallows 3.0 pounds

*Each box contains one gross of items*

1. If you have 100 graham crackers, how many chocolate bars and how many marshmallows do you need to make S'mores with all the graham crackers?

2. If you have 1000 graham crackers, 800 chocolate bars, and 1000 marshmallows: a. How many S'mores can you make?

b. What (if anything) will be left over, and how many of that item will there be? Chemists refer to the reactant which limits the amount of product that can be made from a given collection of original reagents as the limiting reagent or limiting reactant.

3. Identify the limiting reactant for question 2. Explain.


4. Based on the information given, which of the three ingredients (a teddy graham cracker, a mini chocolate bar, or a mini marshmallow):

a. Weighs the most? Explain your reasoning.

b. Weighs the least? Explain your reasoning.

5. If you have 36.0 pounds of graham crackers, 36.0 pounds of chocolate bars, and 36.0 pounds of marshmallows:

a. Which item do you have the most of?

b. Which item do you have the least of?

c. Determine how many gross you have of each ingredient.

6. If you attempt to make S'mores from the materials in question 5: a. How many gross of S'mores can you make?

b. How many gross of each of the two leftover items will you have?

c. How many pounds of each of the leftover items will you have?

d. How many pounds of S'mores will you have?

Box of teddy graham crackers: 9.0 pounds Box of mini chocolate bars: 36.0 pounds Box of mini marshmallows: 3.0 pounds

1 box = 1 gross of items


7. Explain why it is not correct to state that if we start with 36 pounds each of graham crackers, mini chocolate bars, and mini marshmallows, then we should end up with 3 × 36 = 108 pounds of S'mores.

8. Write a balanced chemical equation for the formation of a S’more. Use G as the symbol for teddy graham cracker, Ch for the mini chocolate bar, and M for marshmallow.

9. Using the bag of materials provided to you, complete the following. a. If you had plenty of the other ingredients, how many s’mores could you make

from each ingredient? __________ teddy grahams could make __________ s’mores __________ mini chocolate bars could make __________ s’mores __________ mini marshmallows could make __________ s’mores

b. Given the materials in your bag, what is the maximum number of S’mores you can make?

c. What is the limiting reactant?

d. How much of the excess reactants will be left over?


CW 16: Limiting Reactant: Chemical Quantities

1. Consider the chemical reaction above, in which 2 moles of hydrogen gas (H2) react with

1 mole of oxygen gas (O2) to produce 2 moles of water vapor (H2O). a. How many hydrogen (H2) molecules are in the chemical reagent container?

b. How many oxygen (O2) molecules are in the chemical reagent container?

c. How many water (H2O) molecules are produced from the molecules found in the chemical product container?

d. Is there anything left over in the product container? Why?


2. You react 100.0 g of O2 and 50.0 g of H2 in a container and produce water. 2H2 + O2 → 2H2O

a. How many moles of water could be made from 100.0 g of oxygen and excess hydrogen?

b. How many moles of water could be made from 50.0 g of hydrogen and excess oxygen?

c. Based on your answers to a and b, how many moles of water will be produced from reacting 100.0 g of O2 and 50.0 g of H2? Explain.

d. Which reactant is the limiting reactant in this scenario? Why?

e. What is the theoretical yield of water in this scenario? Report your answer in grams of water.


3. Given the balanced chemical reaction: 2 NO(g) + O2(g) 2 NO2(g) a. Calculate the mass of nitrogen dioxide that can be made from 30.0 grams of NO

and 30.0 grams of O2. (HINT: Do two calculations.)

b. Explain why the amount of NO2 that can be made depends on the limiting reactant.

4. Acetylene gas, HCCH, is commonly used in high temperature torches. a. Write a chemical equation for the reaction of acetylene (HCCH) with hydrogen gas

to form ethane (C2H6).

b. How many grams of ethane can be produced from a mixture of 30.3 grams of HCCH and 4.14 grams of H2?

c. How does the above calculation help determine the limiting reactant?


CW 17: Gravimetric Analysis of PO4–3 Concentration

Objectives

• Predict the formation of a precipitate

• Perform stoichiometric calculations to determine the concentration of a compound in a local water sample

Background Nutrient pollution in the Chesapeake Bay is one of the biggest environmental challenges

facing Maryland and surrounding states. One of the major sources of nutrient pollution comes from phosphates, which are applied as fertilizers in the form of sodium phosphate (Na3PO4). When it rains, excess fertilizer runs off from farmlands and into the Bay. This can lead to algae blooms and pH changes, which are harmful to the Bay ecosystem.

Environmental chemists use stoichiometry to test water samples to determine phosphate levels. They can precipitate out the phosphate from a water sample, then use the mass of precipitate collected to calculate the concentration of phosphate in the original solution. Such tests are known as gravimetric analysis and can determine if the concentration of phosphate is within safe levels and can help scientists target specific areas for remediation or support legal actions such as litigation or lawmaking.

Materials

• Local water sample

• 50 mL graduated cylinder

• 250 mL beaker

• Scale

• Filter paper

• Funnel

• Ring stand with ring

• Wash bottle

• 1 M CaCl2 solution

• 1 M KNO3 solution

• 1 M NaCl solution

• 1 M MgSO4 solution

Class Discussion Questions

1. Given the possible solutions above, determine which one you will need to react to form a precipitate from the sodium phosphate (Na3PO4). The solubility rules/ chart will be helpful in doing this. Write the balanced chemical equation below.


2. Suppose that you perform the reaction from question one using 100 mL of the water sample. You obtain 0.86 grams of the precipitate.

a. Use the grams of the precipitate to determine how many moles of Na3PO4 were in the sample.

b. Given the volume of 100 mL, what was the molarity of Na3PO4 in the sample?

c. What was the molarity of PO4–3 in the sample?

3. Suppose that you failed to add enough of your chosen solution, leaving some of the Na3PO4 unreacted in the water.

a. Would you obtain too much or too little of the precipitate? Explain.

b. Would your calculated molarity be too high or too low? Explain.

c. Which reactant do you want to be the excess reactant? Explain.


4. Determine the amount of your chosen solution to fully react all the phosphate in the water sample. Initial field tests have determined that the concentration of sodium phosphate is around 0.15 M Na3PO4.

a. How many moles of Na3PO4 are in a 25 mL of 0.15 M Na3PO4?

b. How many moles of your chosen compound are needed to fully react with this number of moles?

c. Given that the concentration of your chosen solution is 1.0 M, how many mL of the chosen solution are needed?

d. To ensure that the chosen solution is in excess (and all the Na3PO4 reacts), multiply your answer from the last question by 1.3 to get the volume you should use in the experiment.


5. Watch the video found here: http://bit.ly/1o9bHos a. Explain the difference between clear and colorless.

b. Label the diagram of the filtration set up.

c. Describe how to fold the filter paper so it will fit into the funnel.

d. Why does the demonstrator wet the filter paper with a wash bottle?

e. Compare the composition of the mixture before and after going through the filter.

f. What is an application of this lab technique?

g. What is the purpose of rinsing or “washing” the solid with distilled water?

h. Propose a way to find the mass of the solid collected.


Procedure

Based on your answers to the Class Discussion Questions, write a procedure for how to determine the amount of phosphate ion in a water sample by precipitation. Be sure to include the following:

• Safety precautions • Which reactant will be in excess and why

• How to rinse the glassware to remove as much precipitate as possible

• How to rinse the precipitate and why

• How to determine the mass of the solid collected

• A ready to fill in data table with a place to record all required data and observations

Scientific Poster

In your lab group, you will create a scientific poster to summarize your findings during this experiment. Use a full sized physical poster paper. Choose font sizes large enough to be seen from 3 feet away. All text should be typed or very neatly hand written in dark ink (sharpie).

Grading Rubric

Component Points Introduction

• The purpose of the experiment

• Applications of the technique you used (gravimetric analysis) for measuring and improving the health of the Chesapeake Bay

• The chemistry behind precipitation reactions and gravimetric analysis

10

Methods

• A brief summary of your procedure in paragraph form.

• An original (made by you, not copied) labeled diagram of the filtration set up

10

Data table and observations 10

Original (made by you, not copied) particulate diagram that shows how the reaction occurs

10

Analysis: Show all work neatly.

• Neatly write and balance the chemical equation for your chosen solutions to make the precipitate. Include state symbols (s, l, g, aq).

• Using the mass of precipitate formed, calculate the molarity of the phosphate ion in the water sample.

• Obtain the actual phosphate ion concentration from your instructor and

calculate the percent error.

10

TOTAL: /50

% 𝐸𝑟𝑟𝑜𝑟 = (|𝐸𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙 − 𝐴𝑐𝑡𝑢𝑎𝑙|

𝐴𝑐𝑡𝑢𝑎𝑙) × 100%


Reference Materials


Polyatomic Ions

H3O+ hydronium CrO42– chromate

Hg22+ dimercury(I) Cr2O7

2– dichromate NH4

+ ammonium MnO4– permanganate

C2H3O2–

CH3COO– acetate

NO2– nitrite

NO3– nitrate

C2O42– oxalate O2

2– peroxide

CO32– carbonate OH– hydroxide

HCO3– hydrogen (bi)carbonate CN– cyanide

PO43– Phosphate SCN– thiocyanate

ClO– hypochlorite SO32– sulfite

ClO2– chlorite SO4

2– sulfate

ClO3– chlorate HSO4

– hydrogen sulfate ClO4

– perchlorate S2O32– thiosulfate



Exceptions


Carbonate (CO32–)












Chlorate (ClO3–)



2+

Sulfates (SO42–)



Stoichiometry Map

Documents

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