Chemistry Lab Manual - Fairview High School

College Prep

Written, compiled, edited by

Brian Cox and Dan Albritton Thanks to Daniel Knowles, Eddie Taylor, Dylan Muzny, Cammie

Wickham

Chemistry Lab Manual

Table of Contents

Reference Laboratory Safety Guidelines i, ii

Lab Experiments

Pages Expt #1: 5-Solution Unknown 1-2 Data Analysis: How does the Mass of Pennies Change with Age? 3-5 Expt #2: Relationship between Mass of Reactants to Mass of Products 6-13 Expt #3: Distillation: Separation of two liquids 14-17 Expt #4: Observing the Chemistry of Iron(II) and Iron(III) 18-19 Expt #5: An Activity Series of Metals (Single Replacement Rxns) 20-21 Expt #6: Precipitation Reactions (Double Replacement Rxns) 22-23 Expt #7: Recording Measurements with Correct Sig Figures and

Uncertainty

24-26

Expt #8: Mole Concept – How many atoms or molecules are present? 27-30

Expt #9: Empirical Formula of a Hydrate 31-33

Expt #10: Density of Twizzler Sticks 34-35

Expt #11: A Taste of Molarity 36

Expt #12: Preparing Molar Solutions 37- 38

Expt #13: Demos to investigate concept of limiting reactant and optimal

ratio

39-42

Expt #14: Determination of Percent Yield – How efficient is your reaction? 43-46

Expt #15: Determination of Universal Gas Constant 47-51

Expt #16: Emission Spectrum and Quantum Leaps 52-55

Expt #17: Beer’s Law: Determining Molar Concentration by Absorbance

Spectroscopy

56-60

Expt #18: Relationship between Type of Chemical Bonds and Electrical

Conductivity

61-64

Expt #19: Molecular Models (Drawing in 3-D) 65-67

Expt #20: Polar and Non-Polar Molecules in an Electric Field 68-69

Expt #21: Molecular Polarity Phet Simulations 70-72

Expt #22: Polar and Non-Polar Molecules Station Lab 73-74

Expt #23: Acid-Base Titration: Determining Molar Concentration of

Commercial Vinegar

75-78

Expt #24: Experimental Design: Using Specific Heat Capacity and

Density to Identify an Unknown Metal

79-82

Expt #25: Exploring pH PHET simulation 83-86

Expt #26: NOVA Lethal Seas Video 87-88

Expt #27: Investigating pH 89-90

1

5-Solution Unknown Introduction This laboratory exercise is designed to give you an opportunity to use your skills as a scientist to

solve a problem. Working in pairs, each student will receive either 5 numbered micropipets or 5

lettered pipets. Your task is to determine which of the lettered micropipets A – E match (have the

same chemical) the numbered pipettes 1 – 5. In order to successfully complete this task, you will

have to follow the steps of the scientific method discussed in class: make careful and clearly

recorded observations, communicate with a colleague, critically analyze patterns in your data,

develop a hypothesis, and test your hypothesis experimentally.

Materials • Protected sheets labeled with grids of letters or numbers

• 5 micropipets labeled 1-5 or A-E

Safety • Safety glasses must be worn.

• Avoid skin contact with chemicals, rinse exposed areas with water, and clean any

spills.

Procedure Each student will receive 5 micropipets containing unknown solutions. Half the class will receive

“numbered” pipets (unknowns identified by numbers 1 – 5) while the other half of the class will

receive “lettered” pipets (unknowns identified by letters A – E). The numbered unknowns and

the lettered unknowns each contain the same unknowns but in random order. In other words, 1

does not necessarily correspond to A, 2 to B, and so on.

1) RECORD both data tables for your pre-lab. Setup a data table for the letter AND the number

combinations. Be sure to include ALL possible combinations (there are 10 for both).

2) Each student, working alone, should try all different combinations of his or her

unknowns. Place NO MORE THAN 3 drops of a solution on the cover sheet in the

squares provided and add NO MORE THAN 3 drops of the other unknown on top of the

original drops. RECORD results in your data table. Do not touch the pipette tips to the

solutions at any time; this will contaminate the pipet.

UNKNOWN LETTER SOLUTIONS

Initial Observations: All solutions are clear and colorless.

Combination Tested Result

A + B

A + C Etc.

UNKNOWN NUMBER SOLUTIONS

Initial Observations: All solutions are clear and colorless.

Combination Tested Result

1 + 2

1 + 3 Etc.

2

3) After testing all combinations, absorb the solutions from the cover sheet with paper

towels and dispose of these in the trash. If any residue remains on the cover sheet, wipe

it with a wet paper towel. Return pipettes to tray holder and wash your hands with soap

and water.

Hypothesis To generate a hypothesis, find your partner and analyze your two sets of data (number

combinations and letter combinations). Answer the following questions to guide your

thinking in this process (put boxes around your answers and skip lines).

1. Write the equality (equation) for the two numbers that made a brown precipitate and

set this equal to the two letters that made the same brown ppt.

2. Just as in #1, write the equality for the two letters and the two numbers that made a

yellowish white precipitate when added together.

3. Looking at your two equations from #1 and #2, what letter is common to both? What

number is common to both? Therefore, write the equality.

4. Use your information from the three questions above to match letters A, B, and E to

the possible numbers.

5. Write the four pairs of letter and numbers that formed the clear pink solutions (but

don’t write any equations since you don’t know which set of numbers matches which set

of letters).

6. What letter is common in the four pairs? What number is common in the four pairs?

Therefore, write the equality.

7. By process of elimination, what is the last matching number and letter?

8. Write your complete hypothesis; that is, the complete matches of A-E to 1-5.

Hypothesis Testing

An essential part of the scientific method is to test your hypothesis. 1. Make a third data table, exactly as it is shown below:

2. Predict what result you will see from these four test crosses.

3. Actually mix the four test cross combinations to “test” your predictions.

Questions

1. In generating your hypothesis, what color was the most useful to begin the matching

process and why?

2. Why is it important to have test crosses; that is, what purpose do they serve?

TEST CROSS DATA

Test Cross Predicted Result Actual Result

B + 4

A + 2

D + 2

B + 3

3

Data Analysis activity: How does the Mass of Pennies Change with Age?

Introduction: In this activity, we will consider patterns in the masses of pennies in order

to introduce the concept of measurement variation and uncertainty and to practice data

analysis skills we will need for the next lab investigation (comparing the masses of

reactants and products in a chemical reaction).

1A) If we weighed the masses of pennies over many years, would you expect to see

any differences? Predict whether the masses will remain constant, decrease,

increase or show random fluctuations.

1B) Propose an explanation for how pennies might lose mass over time.

1C) Propose an explanation for how pennies might gain mass over time.

2) Data analysis: The mass of 5 brand new pennies in mint condition were measured on

the same balance on the following results were obtained:

A) If the pennies are all brand new (uncirculated), what is the most likely explanation for

differences in the mass?

# of

pennies

Mass of Pennies (g)

2.49

g

2.50

g

2.51

g

4

The graphs below represent the masses of five pennies in circulation for twenty-five

years. Use the graphs below to answer questions # B-D. B) Which graph is consistent with the hypothesis that the mass of circulated pennies does

not change over time?

C) Which graph is consistent with the hypothesis that pennies become scratched or lose

material off the surface over time?

D) Which graph is consistent with the hypothesis that pennies oxidize or become dirty over

time?

5

3) Actual Data: The masses of sets of pennies from each of the past fifty years were

measured and the results were plotted below. Propose an explanation for the actually

observed results and support your explanation with specific references to the data.

7)Penny Mass by year: The mass data for pennies is presented below by year.

Propose an explanation for this data and support your explanation with specific

references to the data.

6

Investigating the relationship between the Mass of Reactants and

the Mass of Products in a Chemical Reaction

Purpose: To investigate the relationship between the mass of reactants and mass of

products in a chemical reaction.

Introduction Early investigations into the nature of matter focused on mass relationships in chemical systems

as a way to learn more about how elements combined to make compounds. In this lab you will

investigate the relationship between the mass of reactants and the mass of products in three

separate chemical reactions. For each reaction you will carefully measure the mass of the

reaction system before and after the reaction. By analyzing the results of your experiments, you

will be able to draw conclusions about the relationship between the mass of reactants and the

mass of products in a chemical reaction.

Materials

• Electronic balance, Alka-Seltzer tablets, and 50 mL beakers

• Long stem beral pipets and modified jumbo beral pipets

• Modified 30 mL plastic bottles with screw caps

• Four solutions: FeCl3, KSCN, CaCl2, Na2CO3

Safety • Safety glasses must be worn at all times.

• Avoid contact with solutions. Wash your hands at the conclusion of the lab.

• BE SURE NOT TO USE MORE THAN THE INDICATED AMOUNT OF ALKA-

SELTZER IN PART 4.

PreLab Read through the experiment carefully.

On a sheet of notebook paper, write the Title, Purpose, Safety, Data (see tables below).

You will have 4 data tables, one for each part of the lab.

Use the following example for Parts 1, 2, and 3 Data Tables.

Part 1: Reaction of Fe+3 and SCN-

Initial observations of FeCl3 solution (color and clarity):

Initial observations of KSCN solution (color and clarity):

Final observations after mixing (color and clarity):

Total mass of reaction system before reaction: ______ g

(FeCl3 soln + pipet, KSCN soln + pipet)

Total mass of reaction system after reaction: ______ g

(FeCl3 soln + pipet, KSCN soln + pipet)

7

Part 2: Reaction of Ca+2 and CO3-2

Initial observations of CaCl2 solution (color and clarity):

Initial observations of Na2CO3 solution (color and clarity):

Final observations after mixing (color and clarity):


(CaCl2 soln + pipet, Na2CO3 soln + pipet)



Part 3: Reaction of Alka Seltzer and Water in a Beaker

Initial observations of alka seltzer solution :

Initial observations of water in beaker (color and clarity):

Final observations after mixing:


(alka seltzer tablet + beaker of water)



Part 4: Reaction of Alka-Seltzer and water inside a plastic bottle

Total mass of system before reaction: ______ g

(Water, tablet, bottle, and lid)

Total mass of system after reaction: ______ g

(Water, tablet, bottle, and lid)

Final mass of system after loosening lid: ______ g

Observations after loosening the lid: __________________________________

8

Procedure for Parts 1 and 2

Part 1: Reaction of Fe+3 and SCN-

STEP 1: Obtain one modified jumbo pipet and one long stem beral pipet (see below).

Long stem beral pipet

Modified Jumbo Beral Pipet

STEP 2: Fill the modified jumbo beral pipet to the black mark with KSCN using a clean

beral pipet. Fill the long-stem beral pipet to the black mark with FeCl3 solution by

pulling solution up from a beaker. Carefully wipe off any stray drops on the outside of

the bulbs with a paper towel.

STEP 3: Zero the balance. Place the two modified bulbs together on the balance.

RECORD the initial mass of the system and the initial appearances of the reactants.

STEP 4: Insert the long tip of the one of the modified beral pipets into the open end of

the other pipet (see diagram below). Check to make sure the connection is tight.

9

STEP 5: Squirt the contents of one bulb into the other. Then replace the system, with

bulbs still connected, on the same balance and RECORD the final mass.

Important: BE CAREFUL NOT TO SPILL ANY SOLUTION OUT OF THE BULBS!

IF YOU SPILL SOLUTION YOU WILL HAVE TO START THIS TRIAL OVER!

STEP 6: RECORD your final observations. What evidence exists to suggest that a

chemical reaction has occurred?

STEP 7: Subtract the final mass from the initial mass and RECORD the difference

below your Before and After mass data.

STEP 8: Clean up by thoroughly rinsing each bulb with tap water and drying the outsides

with paper towels.

STEP 9: RECORD a total of three good (no spills) trials.

Part 2: Reaction of Ca+2 and CO3-2

Repeat the procedure described in part 1 except substitute CaCl2 and Na2CO3 solutions.

Procedure for Parts 3 and 4

Part 3: Reaction of Alka-Seltzer and water in beaker

STEP 1: Place approximately 10 mL of tap water in a 50 mL beaker.

STEP 2: Take a whole Alka-Seltzer tablet and break it into fourths.

STEP 3: Place the beaker with the water and 1/4 of an Alka-Seltzer tablet on the balance

pan and RECORD the initial mass of the system. Also RECORD your initial

observations of the reactants.

STEP 4: Drop the tablet into the water inside the beaker and swirl gently for several

minutes until there is no longer any evidence that a chemical reaction is occurring.

RECORD your observations. What evidence exists to support the idea that a chemical

reaction occurred?

STEP 5: RECORD the final mass of the system on the same balance and calculate the

difference in mass between reactants and products. RECORD this difference.

STEP 6: Thoroughly rinse the beaker with tap water, dry the outside, and set it on a

paper towel upside down.

10

Part 4: Reaction of Alka-Seltzer and water inside a plastic bottle

STEP 1: Obtain a modified 60 mL plastic bottle equipped with a suspended tray (see

diagram) and 1/4 of an Alka-Seltzer tablet. Dry the outside of the bottle/top.

STEP 2: Remove the lid with the attached tray and fill the bottle with water up to the

black mark which is drawn around the outside of the bottle.

STEP 3: Place the 1/4 tablet of Alka-Seltzer on the tray and carefully replace the lid in

such a way that the tablet remains on the tray and is out of contact with the water.

Important: Do not use more than 1/4 of an alka-seltzer tablet!

This is to minimize risk of excessive pressure build-up.

STEP 4: This is a critical step. CHECK TO MAKE SURE THAT THE LID IS

SCREWED ON AS TIGHTLY AS POSSIBLE!

STEP 5: Weigh the system on the balance and RECORD the mass and your initial

observations.

STEP 6: Gently shake the bottle in such a way as to dislodge the Alka-Seltzer from its

tray and send it plummeting into the water. Continue gentle shaking until no further

evidence of a reaction can be observed. RECORD your observations.

STEP 7: RECORD the final mass of the system using the same balance.

STEP 8: Loosen the lid and RECORD any observations. Completely remove the lid and

then replace it again; REMEASURE the mass of the entire setup.

STEP 9: Rinse the inside of the bottle thoroughly with water, dry the outside of the

bottle, and set it upside down on an open paper towel.

************************************************************************

11

Lab Report Requirements

• Required Elements: Title, Purpose, Safety, Your personal data tables recorded in

the lab including initial and final observations, and answers to questions, Class data table

with calculated mass differences, Conclusion

• CONCLUSION MUST BE TYPED (optional to type other elements of report)

• Use the data in the table below entitled “Conservation of Mass Data for Conclusion”

when writing your conclusion (DO NOT use your data).

Mass of Reactants vs Products for Conclusion

Part 1: FeCl3 + KSCN Trial #1 Trial #2 Trial #3

Before: 4.78 g 4.70 g 5.03 g

After: 4.76 g 4.70 g 5.03 g

Difference:

Part 2: CaCl2 + Na2CO3 Trial #1 Trial #2 Trial #3

Before: 4.54 g 5.00 g 5.00 g

After: 4.54 g 5.00 g 4.98 g

Difference:

Part 3: Alka Seltzer + H2O

in beaker

Trial #1 Trial #2 Trial #3

Before: 36.95 g 40.43 g 36.50 g

After: 36.68 g 40.18 g 35.80 g

Difference:

Part 4: Alka Seltzer + H2O

in plastic container

Trial #1 Trial #2 Trial #3

Before: 29.80 g 32.63 g 30.65 g

After rxn with top closed: 29.78 g 32.61 g 30.64 g

Difference:

Before: 29.80 g 32.63 g 30.65 g

After rxn with top open: 29.70 g 32.52 g 30.55 g

Difference:

• Type the conclusion and write in narrative form using complete sentences.

(Don’t simply answer the questions; the question and hints are provided to help

you organize your writing.)

¶1: In one to two sentences, summarize the objective of the experiment and the

general experimental approach.

12

¶2: Using the data from the table, ANALYZE THE DATA. Using the hints

below to guide your thinking, draw a general conclusion or statement

regarding the relationship between the mass of reactants and the mass of

products in a chemical reaction, citing examples from all four Parts of the

experiment to support your conclusion.

HINTS:

For each of the four parts of the experiment, compare the DIFFERENCES

between the before and after.

Comment on which experiments are the data the same within the experimental

uncertainty? (Important: Our balances have an uncertainty of + 0.01 g for each

individual measurement; thus difference values that are within + 0.02 g are the

SAME within the experimental uncertainty. They may be different but our scale

cannot determine that the before and after measurements are different).

Identify expts for which the data are definitely DIFFERENT (i.e. difference is

greater than + 0.02 g)?

Comment on the phases of matter are present in the products. Comment on what

is critically different about the experimental equipment used in the alka seltzer

water experiment in parts 3 vs part 4.

¶3: List a source of potentially significant error in your data.

Hints: Be specific and DO NOT say “human error” or “we could have read

the scale incorrectly.”

FOCUS ON FACTORS THAT CHANGE MASS DIFFERENCES!

Questions:

1) Explain the concept of experimental uncertainty in a measurement. What does it mean to say

the mass of a sample is 1.25 + 0.01 g. What is the range of possible values?

2) John Dalton proposed Atomic Theory in the early 1800, based in part on explaining

the results of experiments comparing mass of reactants to the mass of reactants. Below is

a diagram of the reaction between calcium chloride and sodium carbonate to form

calcium carbonate and sodium chloride: CaCl2 + Na2CO3 → CaCO3 + 2 NaCl

+

+

Compare the number of atoms of each element (Cl, Ca, Na, C, O) on the reactant side and

the product side. How can the concept of atoms explain the relationship between mass of

reactants and mass of products you observed for this reaction?

13

B) In parts 3 and 4, a reaction of alka seltzer in water reaction to form carbon dioxide gas,

CO2 can be represented symbolically as

→ +

where the oval shape represents carbon dioxide gas.

Below is a diagram of a beaker on the left containing products and on the right containing

reactants except for carbon dioxide. Using an oval symbol to represent carbon dioxide

molecules, draw a diagram to represent the location of the carbon dioxide molecules

after the reaction. Why is the mass of the product beaker less than the mass of the

reactant beaker?

Reactants Products

Below is a diagram of the sealed bottle containing the alka-seltzer and water reaction

reactants and products. Using an oval symbol to represent carbon dioxide molecules,

draw a diagram to represent the location of the carbon dioxide molecules after the

reaction. Why is the mass of the product beaker the same as the mass of the

reactant beaker?

14

Distillation: Separation of two liquids

Introduction One of the most important methods for separating a mixture of liquids is distillation, a process

that depends on differences in the boiling points of the two liquids. In distillation, a mixture of

liquids is heated. When a liquid reaches its boiling temperature the vapor passes through a cooled

tube (a condenser), where it changes back into a liquid.

Purpose To separate two volatile liquid by distillation and identify the component liquids by boiling point

and chemical tests.

Prelab: Title, Purpose, Safety, 2 data tables

Safety

1) Don’t forget to add boiling chips and BOIL GENTLY. 2) Never add boiling chips to water that is already heating.

3) Tie hair back and keep papers away from the bunsen burner.

4) If anything catches on fire, turn the burner off.

5) Wear goggles.

6) STOP HEATING BEFORE ALL THE LIQUID IS GONE.

Materials

• 3 test tubes, racks, and masking tape

• 2 ring stands, wire gauze, burner, 2 clamps, 2 rings, and blocks

• thermometer inside a 1-holed stopper

• 250 mL Florence flask, condenser, and rubber tubing

• 250 mL beaker and 25 mL graduated cylinder

• boiling chips, liquid A, and food coloring

Data Table #1: Temperature oC vs. Time (min)\ (Every ½ min until complete.

Typically it takes 15-20 minutes to complete experiment. See example of first 3 data

points below.

Temperature (oC) Time (min)

0.5

1.0

1.5

Data Table #2: Odor and Flammability Tests

Construct a table for distillate #1 and distillate #2 to describe odor test and flammability test

results.

15

Setup 7) Check to see if the clamp/gauze/flask is at least 7 inches above the bunsen burner.

8) Measure 25 mL of the mixture in a graduated cylinder. Pour the liquid into the distilling

flask. Add 3 boiling chips and 1 drop of food coloring.

9) Turn on the cold water in order to fill the condenser.

10) Have your teacher approve your setup, after which you may light the burner.

Data Collection and Recording

• Light the burner and take a temperature reading every 1/2 minute. RECORD your values as

described in step 3 below.

• BOIL THE MIXTURE GENTLY! If you boil too vigorously, you will have to start over.

The key is to keep the food coloring in the distilling flask. The temperature change around

70 degrees has special importance.

IMPORTANT: Switch collecting test tubes when the temperature reaches about 90

degrees. The first test tube will contain the 1st Distillate and the second test tube will contain

the 2nd Distillate. Label the test tubes with these titles.

• While the distilling flask is heating, RECORD two data tables. The first will be a two-

column set of time and temperature readings, time in minutes and temperatures in degrees

Celsius. The second will be a chart with “Mixture”, “1st Distillate”, and “2nd Distillate” down

the left side and “Odor” and “Flammability” across the top line.

• STOP HEATING WHEN THERE IS JUST A LITTLE LIQUID LEFT, the temperature

has stabilized, and you have at least 1.5 inches of 2nd Distillate. Turn off the burner.

16

• Fill a third test tube (1/3 up the tube) with the original mixture. RECORD odors from the

three tubes by waving the vapors towards your nose.

• The second test is a flammability test. Perform this test in the sink. Dip the end of a small

piece of scratch paper or paper towel into each solution. Hold a lighted match to the wet end.

If it burns, quickly put the fire out by running it under some water. Do not throw the paper in

the sink. Please find the trash can for its disposal. Never use the sink as a trash can.

RECORD your results.

• Clean up as described by your instructor.

FINAL LAB REPORT: (In addition to prelab and data tables)

GRAPH Follow these directions carefully! Make a graph (on graph paper and using the entire sheet) with time in minutes on the x-axis (long

side of paper) and temperature in degrees Celsius on the y-axis (short side). Label your axes with

variables and units. You may connect the points dot-to-dot because the graph does not represent

a function.

Discussion and Questions Hint: Read the Lab’s Introduction Section! 1. What is the purpose of a distillation procedure?

2. Fill in the missing word: In order to separate two liquids from each other in a distillation

experiment, the two liquids must have different _______________ points (or temperatures).

3) Use the graph below to answer questions, A - D

3A) Which region I, II, III or IV represents the MIXTURE HEATING?

3B) Which region I, II, III or IV represents the 1ST COMPONENT VAPORIZING?

3C) Which region I, II, III or IV represents the 2ND COMPONENT HEATING?

3D) Which region I, II, III or IV represents the 2ND COMPONENT VAPORIZING?

Testing Distillates after Separation

17

4) Using the symbols to represent liquid distillate #1 and to represent distillate #2, draw

atomic/molecular level pictures of the following. (Hint: How do you represent liquids vs gases in

terms of spacing between particles and organization of particles?)

A) Mixture heating -The two liquids mixed together and heating but neither has boiled off yet.

B) First component is gas phase (has already boiled off) and second component is still liquid (2nd

component heating).

5) In the graph below, what is the boiling point of each liquid distillate?

6) Identification of the mixture: The original mixture contained TWO of the following THREE

clear, colorless liquids:

Acetone (nail polish remover) : flammable, distinct odor, approximate boiling point 51oC

Isopropyl alcohol : flammable, distinct odor, approximate boiling point 75 oC

Water – nonflammable, odorless, approximate boiling point 95 oC.

Based on the results of your experiments, which two liquids were present in the

original mixture? Support your conclusion with three lines of evidence

(flammability test results, odor test and approximate boiling point – use best guess from

graph of your data).

Dist #1

Dist #2

18

Observing the Chemistry of Iron(II) and Iron(III)

Purpose: To determine the charges of iron in different compounds and determine

whether different charges result in different chemical behavior.

Safety: Wear safety goggles. Wash off chemicals if contacted. Wash hands after lab.

Materials: 5 chemicals listed below. Pipets for each chemical and one for water.

Procedure:

Copy the data table below in your lab report.

• Obtain one dropper for each of the chemicals on the data table.

• Slide a single sheet of these instructions into a sheet protector. You will place drops on the

plastic sheet over the rectangle representing the two chemicals.

• Don’t touch the pipet to a drop already on the sheet. This will contaminate the dropper.

K4Fe(CN)6 K3Fe(CN)6 KSCN

• Place 3 drops of Iron(II) ammonium sulfate in each of the three rectangles on the top row of

the data table. RECORD your initial observations (color, clarity) above the chemical

name.

• Place 3 drops of Iron(III) chloride in each of the three rectangles on the bottom row in the

data table. RECORD your initial observations (color, clarity).

• Using the chemicals listed at the top of the columns, RECORD your initial observations

and then place ONE drop of each chemical into the 3 already present.

• Lastly, add 25 drops of tap water to each pile. RECORD your final observations (in the

blanks on the grid of your paper) of the chemicals after the water has been added

CLEAN-UP:

• Place paper towels over the sheet protector to absorb the chemicals; throw these

away. Rinse off the protector and dry its surface. Return pipets to containers.

Potassium iron(II) cyanide

Initial obs: Potassium iron(III) cyanide

Initial obs:

Potassium thiocyanate

Initial obs:

Initial obs:

FeCl3

Initial obs:

Fe(NH4)2(SO4)2

19

LAB REPORT WRITEUP:

• Title, Purpose, Safety

• Completed Data Table

– Initial color observations for each chemical

- Final color observations for each reaction.

• QUESTIONS (Do not have to recopy questions)

Hints for #1 and 2: (What do the Roman numerals in the name mean?)

1) What is the charge on iron in the compound potassium iron(II) cyanide?

2) What is the charge on iron in the compound potassium iron(III) cyanide?

Hints for #3 and 4: Sum of charges must total 0. What are the charges on ammonium and sulfate

ions? How many ions are present? What is their total charge?

3A) How many total ions are present in the compound, FeCl3?

3B) What is the charge on the iron in the compound FeCl3?

4A) How many total ions are present in the compound, Fe(NH4)2(SO4)2?

4B) What is the charge on the iron in the compound, Fe(NH4)2(SO4)2?

IMPORTANCE OF DIFFERENT CHARGES ON THE IRON

5A) Does the charge on the iron matter? If the charge does not matter then the different

iron compounds should react in exactly the same way. Did the two compounds,

Fe(NH4)2(SO4)2 and FeCl3 react in the exact same way or differently when each was

reacted with K4Fe(CN)6, K3Fe(CN)6 and KSCN? Support your answer with specific

examples from your data.

5B) Does the charge on the iron matter? If the charge does not matter then the different

iron compounds should react in exactly the same way. Did the two compounds,

K4Fe(CN)6, K3Fe(CN)6 react in the exact same way or differently when each was

reacted with Fe(NH4)2(SO4)2 and FeCl3? Support your answer with specific

examples from your data.

5C) In conclusion, why is it a requirement that compounds such potassium iron(II)

cyanide, have a different name than compounds such potassium iron(III) cyanide

even though they contain the same type of atoms?

20

An Activity Series of Metals

Introduction: Different substances have different attractions for electrons. The substance

with the stronger attraction will pull electrons away from the substance with the weaker

attraction. When discussing metals, which lose electrons, the metal which has a greater tendency

to lose electrons than another metal is said to be more “active”.

Purpose: In this lab, you will determine the relative attraction that copper, magnesium, and

zinc have for electrons by allowing the elemental (electrically neutral) form of one metal atom to

compete for electrons with the +2 ion of another metal. By examining the results of your

experiments, you should be able to rank the three metals copper, magnesium, and zinc by their

activity, i.e. their tendency to lose electrons.

Materials: • Reaction sheet (with X’s) that matches your data table.

• Reaction sheet inside sheet protector.

• 2 small strips (~ 0.5 cm by 1 cm) each of copper and zinc

• 2 small strips (~ 1 cm long) of magnesium

• 0.5M aqueous solutions of magnesium chloride, zinc sulfate, and copper(II) chloride

Safety: Goggles! Copper(II) chloride is moderately toxic.

Wash hands after completing the lab.

Procedure:

Copy the data table from the NEXT PAGE into your lab report.

1. For each reaction tested be sure to record your initial and final observations next to the

metal or above the solution.

2. Test the following combinations for reactivity:

a) Mg(s) + ZnSO4(aq)

b) Mg(s) + CuCl2(aq)

c) Zn(s) + CuCl2(aq)

d) Zn(s) + MgCl2(aq)

e) Cu(s) + MgCl2(aq)

f) Cu(s) + ZnSO4(aq)

3. Place each of the strips of metal to be tested in the blank space of the covered reaction

sheet.

4. To each piece (strip) of metal, add enough solution of the aqueous ions of one of the

different metals to cover the strip.

5. Record your final observations in the boxes on your data table (abbreviations on next

page).

See back of this sheet for data table.

21

6. Cleanup: If the meal strip is intact, rinse it thoroughly with tap water, pat dry with a

paper towel, and then return to the side bench. Dispose of the metals that reacted and any

remaining solution in the trash can.

Data: You may use the following symbols: X = not tested, + = reacts, and NR = no reaction.

MgCl2 (aq) solution CuCl2 (aq) solution ZnSO4 (aq)

solution

Mg+2 Cu+2 Zn+2

Mg metal (elemental form)

X

Cu metal (elemental form)

X

Zn metal (elemental form)

X

Questions: Answer on your lab report below your data table. 1. Metals in their elemental (neutral) form tend to give up electrons to form positive ions. The

more active the metal, the greater the tendency to give up electrons.

a) Which metal in its elemental form reacted with the most solutions? (2 points)

b) Which metal in its elemental form reacted with the fewest solutions? (2 points)

c) Based on the definition of metal activity and the results from your experiments, rank the

three metals in order from most active to least active: (1 point)

Most active ______ Medium active ______ Least active ______

2. For the six combinations of reactants tested in this experiment, predict the products

and balance the final equation. If a reaction did not occur, simply write NR for the

products. Phases of matter for the products are (s) for the elemental metal and (aq) for

the compound. (1 point for each entirely correct equation.)

a) Mg(s) + ZnSO4 (aq)

b) Mg(s) + CuCl2 (aq)

c) Zn(s) + CuCl2 (aq)

d) Zn(s) + MgCl2 (aq)

e) Cu(s) + MgCl2 (aq)

f) Cu(s) + ZnSO4 (aq)

22

Precipitation Reactions Lab

Introduction When an ionic compound (i.e. a salt) dissolves in water, it dissociates into its constituent ions [for

example: NaCl(s) + H2O(l) Na+(aq) + Cl-

(aq)]. If ions from two or more compounds are

dissolved in water at the same time, oppositely charged ions from different salts may combine to

form an insoluble salt or precipitate. In this lab, you will observe 8 different combinations of

soluble salts that result in the formation of a precipitate. Using the solubility rules and your

observations of the reactions you should be able to write and balance the net ionic equations (the

equations that represent the formation of the solid only!).

Initial Observations Record initial observations of each solution on a separate sheet of paper using the format set

listed below. This is your first Data Table. Note: You may use abbreviations such as

cl = clear, c = colorless, ppt = precipitate.

Solution Initial Observations Solution Initial Observations

CoCl2 NaOH

KI Na3PO4

CuSO4 Na2CO3

Pb(NO3)2 AgNO3

Safety Avoid skin contact with all solutions. If you spill any chemical on your skin, rinse the exposed

area immediately several times with water. Inform the instructor of any spills. NaOH is

corrosive. AgNO3, Pb(NO3)2, CoCl2, and CuCl2 are toxic if ingested. Wear safety glasses.

Procedure

Copy the second Data Table EXACTLY as you find it on the next page.

11) For each reaction, combine 2-3 drops of the two chemicals in the appropriate grid on your

plastic sheet. DO NOT TOUCH PIPET TIP TO CHEMICALS!

• For each reaction, list the final observations (see space below).

• Cleanup by absorbing chemicals on paper towels. Rinse sheet thoroughly with water in the

sink and then dry with a paper towel.

• First, write the two pairs of ions that the aqueous salts represents. Switch “partners”, pairing

the first cation with the second anion and vice versa. Write the products of the double

replacement after the arrow. Balance the molecular equation. Make sure you include the

phases of the products!

• Use the solubility rules chart to identify the precipitate.

23

Example: BaCl2(aq) + K2SO4(aq) Final observations: White PPT

Ions: Ba+2 Cl -1 K+1 SO4-2

Balanced Equation: BaCl2(aq) + K2SO4(aq) 2KCl(aq) + BaSO4(s)

(Solubility rules: BaSO4 is insoluble and KCl is soluble) Precipitate (ppt) = BaSO4

Check to see that your equation gives the appropriate charges for the ions, indicates

the state of matter (aq) or (s), and that the final equation is balanced.

1) CoCl2(aq) + NaOH(aq)

Final observations: ppt=

Ions:

Balanced equation with phases:

2) CoCl2(aq) + Na2CO3(aq)

3) CoCl2(aq) + Na3PO4(aq)

4) CuSO4(aq) + NaOH(aq)

5) CuSO4(aq) + Na2CO3(aq)

6) CuSO4(aq) + Na3PO4(aq)

7) AgNO3(aq) + Na2CO3(aq)

8) Pb(NO3)2(aq) + KI(aq)

24

Recording Measurements with Correct Significant Figures and

Uncertainties

Introduction: Every measurement always has some degree of uncertainty or range of possible

error. All measuring devices including expensive digital electronic devices, have an

experimental uncertainty. The precision of the measurement limits the number of significant

digits (places that are meaningful given the precision of the measuring device) that can be

recorded. The last reported digit is always uncertain. In this lab we will practice reading different

measuring devices and correctly recording measurements.

Purpose: To practice recording measurements using different instruments to the correct

number of significant figures and appropriate uncertainty.

Materials:

• Rulers calibrated with different smallest marks (10 cm increments, 1 cm increments, 0.1

cm increments)

• A strip of laminated paper whose length will be measured by the 3 different rulers

• A 100 mL beaker

• 10 mL and 100 mL graduated cylinders

• A digital electronic scale

• A buret

• A thermometer

Rules for recording significant figures for a nonelectronic instrument. 1) Estimate 1 place to the right of the smallest marked unit.

2) Conservative estimate:

Uncertainty will be + 5 units in the last recorded digit (uncertain).

28 + 5 mL 28.3 + 0.5 mL 28.32 + 0.05 mL

2.8 + 0.5 cm 2.71 + 0.05

cm

25

Data Table:

Important: Circle the uncertain digit in each measurement (last recorded digit)

When measuring liquids, observe the bottom of the curve ("meniscus") at eye level.

Procedure:

1. For each of the measuring tools at your lab table, first determine the smallest marked

unit or increment on the device.

For example, the yellow ruler only has markings 0 and 10; therefore, its smallest marked

unit is 10 cm. You would therefore estimate 1 place to the right of the 10’s place or in

other words to the nearest 1 cm.

2. Measure the blue strip's length with the three colored rulers: yellow, then orange,

then pink.

4. Measure the volume of water in the graduated cylinder.

5. Measure the volume of water in the beaker.

Measured Item Smallest Marked

Unit on Measuring

Device

Measurement (estimate 1

place past smallest

marked unit)

Length of blue paper

(using yellow ruler)

10 cm

(estimate to nearest 1 cm)


(using orange ruler)


(using pink ruler)

Volume of water in 100 mL

graduated cylinder

Volume of water in 100 mL

beaker

Temperature of water (oC)

Mass of stopper on

electronic balance in grams

Initial Volume of water in

buret (mL)

Final volume of water in

buret (mL)

26

6. Measure the temperature of the water in the beaker.

7. Measure the stopper's mass on the electronic balance

8. Measure the initial volume of the water in the buret.

9. Add 10.00 mL to your value and then drain out 10.00 mL of water into a 10 mL

graduated cylinder. Check to see if you drained the correct volume out of the buret by

checking the amount of water drained into the graduated cylinder.

IMPORTANT NOTE: Burets are marked to make it easy to determine the volume of

liquid delivered out of the buret. The numbering starts at zero at the top and numbers

down.

Questions:

1. For each of the measurements in the third column above, circle the digit that is an

estimate.

2. Why does the same length of blue paper have three different measurements in your

data table above?

3. Explain why the measured values for the graduated cylinder and for the beaker have

different numbers of significant figures. (The answer is NOT because there are different

volumes.)

4. A student zeroes the electronic balance and places an object on the pan. The balance

reads 10.00 g. The student writes 10 g on his data table. Is this correct? Why or why

not?

Note: for electronic measuring devices, RECORD ALL DIGITS DISPLAYED.

Since there are no markings for an estimation, the last digit has some uncertainty

associated with it.

27

Mole Concept: Determining the # of Particles

Purpose: To use the mole concept determine the number atoms or molecules present in

samples of different common substances. Introduction In order to determine chemical formulas or chemical recipes we need to determine the

number of particles (atoms or molecules) present. Since atoms and molecules are far too tiny for

us to see with our eyes and too numerous to count even if we could see them, we use the mole

concept to count by weighing and count by grouping. In part 1 of this lab, you will perform a

simulation using plastic blue spheres to represent atoms to explore the concepts of counting by

weighing and counting by grouping. In parts 2,3, and 4, you will count the number of atoms or

molecules present in the samples of 3 different substances: copper, water and sugar.

.

Safety: No safety issues.

Lab Station #1: Modeling the Mole Concept: Counting by weighing and counting by

grouping

In this first simulation, we will use blue plastic spheres to represent atoms. Recall that elements

typically consist of a mixture of isotopes, atoms of the same element with different atomic masses. In our

simulation, the two isotopes of the imaginary element “blueium” are represented by spheres of slightly

different sizes (masses) and shades of blue. “Blueium” consists of 75% dark blue “atoms” and 25% light

blue atoms.

Part 1: Determining the average “atomic” mass element “blueim”.

1) In a small plastic weighing cup, there are 8 “atoms”, 6 larger and darker blue plastic

spheres and 2 smaller and lighter blue plastic spheres. Record the data below and

calculate the average “atomic” mass to the correct number of significant figures.

Mass of cup + 8 “atoms” of “blueium”: ___________ g

Mass of empty cup: ___________ g (Pour “atoms” into a

separate container or your hand to empty cup. Return after weighing.)

Mass of 8 “atoms” of “blueium”: ___________ g ÷ 8 atoms =

Average mass of 1 atom: ___________ g

(Note: 8 is counting number (infinite SF) ; therefore SF of average mass = SF of mass of 8

“atoms” )

You will use the average mass of 1 atom for your calculation in Part 2.

28

Part 2: Counting by Grouping: Determining the “molar mass” of blueium

Chemists count atoms in groups of 6.022 x 1023 (1 mole). In our simulation we will count in

groups of 24 atoms (1 microplate). Instead calculating the mass of 1 mole (6.022 x 1023) atoms

the molar mass, we will calculate the “microplate mass” which is the mass of a set of 24 atoms in

our microplate. Since you calculated the average mass of 1 atom on the previous page, if we

multiply the average mass by 24 (atoms in 1 microplate), we obtain a “microplate mass”.

Calculate the “microplate mass”: 24 atoms/microplate x Average mass of 1 atom(from

previous page) =

Part 3: Determining the number of “microplates” and atoms in the mystery cup:

The mystery cup (a cup with a lid) contains an unknown number of bluim atoms.

(Analogy: Atoms are “invisible” to our eyes because they are too small to be seen). Use

the mass to determine how atoms that are present in the cup.

Mass of cup + ? “atoms” of “blueium”: ___________ g

Mass of empty cup: ___________ g

Mass of ? “atoms” of “blueium”: ___________ g

Calculation of # of microplates: (Hint: Use the “microplate mass” you calculated above

to fill in the grams present 1 microplate in your conversion factor.)

g blueim ( )

Calculation of # of atoms:

( )

Check: Are you right? Open the lid and remove the blue spheres from the mystery box and place blue spheres in

individual wells of 24 microplates. Count the # of microplates filled and total number of blue

spheres present. Does your actual count match your prediction?

“ Microplate mass” of blueium: ___________ g/ 1 microplate

g blueium

1 microplate

= microplates (analogous to # of moles)

microplates

1 microplate

24 atoms

= atoms blueium

(round to nearest whole number of atoms)

29

Directions for Stations 2-4: Rotate in groups through the 3 lab stations. At each station, complete the data table. Record

each value on your lab report to the correct number of significant figures. Using the mole

concept, determine the number of particles present.

Data:

STATION #2 HOW MANY ATOMS OF COPPER (Cu) ARE PRESENT?

Mass of cup + Copper: ________ grams

Mass of cup (empty): ________ grams

Mass of Copper: ________ grams

Calculation (grams → moles → atoms) of # of Copper atoms present:

( )( )

STATION #3 HOW MANY WATER MOLECULES, H2O, ARE

PRESENT IN THE GRADUATED CYLINDER?

Volume of water: ________ mL

A) The density of water is 1.00 g/mL. Use the density equation, (density = mass ÷

volume) to calculate the mass of your water sample. Show work and unit cancellation.

B) Using the mass of water from part A above, calculate the # of water molecules present

in your sample. Show work and unit cancellation.

C) How many H and O atoms are present in 1 water molecule, H2O? H atoms: O

atoms: How many total H and O atoms are present in 3 water molecules, H2O?

H atoms: O atoms:

How many total H and O atoms are present in 10 water molecules, H2O?

H atoms: O atoms:

g Cu

=

30

How many total H and O atoms are present in 100 water molecules?

H atoms: O atoms:

If you have x number of water molecules, write an equation for the number of H atoms

and O atoms present:

H atoms: O atoms:

D) Calculate the number of H atoms present in your sample of water. (Hint: How many

water molecules do have in your sample?)

STATION #4 HOW MANY MOLECULES OF SUCROSE, C12H22O11, ARE

PRESENT?

Mass of cup + Sucrose: ________ grams

Mass of cup (empty): ________ grams

Mass of Sucrose: ________ grams

A) Calculation of # of Sucrose molecules present:

B) Calculate the # of C atoms present in your sample of sucrose:

C) Calculate the # of H atoms present in your sample of sucrose:

31

Empirical Formula of a Hydrate

Purpose:

To determine the empirical formula of a copper sulfate hydrate. Also, to use the concepts

of mole ratio and % composition in the calculations.

Introduction: A compound that contains water molecules as part of its crystal is called a hydrate. One example

of a hydrate is copper sulfate hydrate, CuSO4 ∙ nH2O, where n represents an integer. The ∙ means

that the attachment between the CuSO4 unit and the H2O units is relatively weak. For example,

this attachment is much weaker than the bonds that hold together the hydrogen and oxygen atoms

in water and thus can be broken by heating with a Bunsen burner flame. In this experiment, you

will determine the simplest ratio of CuSO4 units present to H2O units present in the crystal.

Safety: 1. Wear goggles.

• Keep hair and other materials away from the flame.

• The crucible will be VERY HOT. Always handle the crucible with tongs.

If you do burn yourself, immediately run cold water over the area.

• Handle the crucibles with care. They break easily.

• Be careful in Step 10 of the Procedure below!

Materials: Ring stand, ring, triangle, burner, crucible, crucible tongs,

CuSO4·nH2O crystals (blue), and balance

Procedure:

1. Set up the apparatus as shown by your teacher. There should be about 5 inches between the

burner and the crucible.

2. Make sure the pan on your balance is clean.

3. Find the mass of crucible. Record this weight under Data.

4. Remove the crucible from the balance; put in enough hydrate crystals to half fill the crucible.

5. Measure the mass of crucible + copper sulfate hydrate and record the mass.

6. Place the crucible in the clay triangle. Heat gently by moving the flame back and forth under

the crucible for 10 – 15 minutes. Heat until the compound completely changes color (no blue

color visible). DO NOT OVERHEAT!!!

7. Turn off the burner.

8. When the crucible is cool, transfer the crucible with tongs to the balance. DO NOT PUT THE

HOT CRUCIBLE ON THE BALANCE!

9. Record the mass of copper sulfate anhydrous + crucible. The word “anhydrous” means

without water.

10. WHILE STILL WEARING GOGGLES AND LEANING BACK FROM THE

CRUCIBLE, add a few drops of water from the sink to the crucible and observe what happens.

This is cool! ☺ Actually, it’s quite hot; don’t touch the crucible! After adding a few drops of

water, fill the entire crucible with water to cool the hydrate/crucible.

11. Empty the crucible into the trashcan – NOT IN THE SINK. Clean out the crucible with

water and dry it for the next class.

32

Observations: Answer on your lab report!

Initial observations of copper sulfate hydrate prior to heating:

Final observations of copper sulfate anhydrous after heating:

For the DATA and CALCULATIONS sections, recopy what you see

below on your lab report. DO NOT write in the lab manual please.

Data:

Mass of crucible + CuSO4·nH2O __________ g (1)

Mass of crucible __________ g (2)

Mass of CuSO4·nH2O (before heating) __________ g (3)

Mass of CuSO4 (after heating) + crucible __________ g (4)

Mass of crucible (same as above) __________ g (5)

Mass of CuSO4 __________ g (6)

Calculations:

1. Mass of Water: Calculate the mass of water by determining the difference between lines (3)

and (6):

Mass of CuSO4·nH2O – mass of CuSO4 =

2. Moles of Water: Next, convert this mass of water to moles using the molar mass of water.

Show work and keep 3 sig figs for the mole answer.

3. Moles of CuSO4: Calculate the moles of CuSO4 using line (6) and the molar mass of CuSO4.

Show work and keep 3 sig figs for the mole answer.

4. Mole Ratio of H2O to CuSO4: Find the mole ratio by dividing the moles of water and the

moles of CuSO4 by the smaller number of moles. Show work.

5. Final Empirical Formula: If we could completely remove all the water from the CuSO4, our

ratios would be whole numbers. However, since this is difficult to do in the lab, you’ll probably

have a decimal for the number of water moles. It’s OK to round this decimal to the nearest whole

number. Write the Empirical formula by filling the two blanks below using the simplest whole

numbers in a number:

________ CuSO4 · _____ H2O

33

Data Analysis:

Two groups of students performed the CuSO4·nH2O experiment that you just completed. Each

group started with 2.00 g of CuSO4·nH2O and heated off the water. Group 2 heated for a shorter

time and still had some blue left in their crystals after heating. Data for both groups are displayed

below and have simplified to the masses of the CuSO4 and H2O.

Group 1 Group 2

CuSO4 mass = 1.28 g CuSO4 mass = 1.38 g

H2O mass = 0.72 g H2O mass = 0.62 g

1. Calculate the empirical formulas for BOTH Group 1 and Group 2. Follow the same 4 steps

that you used on the previous page.

2. The “accepted” value for the hydrate is CuSO4·5H2O. Why did Group 2 obtain an

incorrect mole ratio? Your answer should use the color change and data above.

Error Analysis:

One way of judging the accuracy of your results is to compare your % by mass of water to the

accepted value.

1. The “accepted” formula for the hydrate is CuSO4·5H2O. Use this formula with the 5 moles of

water to calculate the “accepted” % composition by mass of water (show work):

Recall

{Hint: the mass of water in the numerator is found by 5 x 18.02 grams.}

2. Next, calculate the % by mass of water using YOUR GRAM VALUES in the Data and

Calculations sections. The “part” is grams water from Step 1 in the Calculations and the “whole”

is the grams of CuSO4·nH2O from line (3). (show work)

3. Determine the % error from the “accepted” % composition by mass of water using your

answers from 1. and 2. above. (show work)

% error = your value – accepted value x 100%

accepted value

% mass of water = mass of water x 100%

mass of CuSO4 + mass of water

34

Density of Twizzler Sticks Introduction One of the most important skills for a scientist is data analysis. In this experiment, you will

collect mass and volume data for a yummy candy – Twizzlers and use the data to determine the

density of the candy. In analyzing your data, we will use graphical analysis to determine the

density and to determine the reliability of your data.

Purpose: To use line of best fit to determine the density of two types of candies and analyze the

quality of the data using a correlation coefficient, r2 value.

Safety:

Wear goggles

Do not eat the Twizzlers

Materials

Electronic Balance 25 ml graduated cylinder twizzler

Part 1: Data Collection ***Make sure to record all of your data with the correct number of significant figures and

units!!!

Type of Candy Mass of

Candy

Initial

Volume

Vi

Final

Volume

Vf

Total

Displacement (Vi-

Vf)

1. Obtain a piece of twizzler and a piece of a gummy worm.

2. Weigh the candy on the scale. Record the value

3. Using a 25 mL graduated cylinder, measure the volume, of the twizzler using water

displacement. Record your value. Work fast, as the candies tend to absorb water…

4. Record value on the class google spread sheet

Part 2: Data analysis- You are now going to create two graphs to analyze the class data.

Your first graph will be mass of gummy worms vs volume.

The second, will be mass of twizzler vs volume.

Graphs need to be a scatter plot and include the following:

Title Labeled axis Units Linear Trendline R2 value Workable scale Equation

You may either use google sheets or Microsoft Xcel to prepare graphs.

If you need help making the graphs refer to a partner or one of these videos:

Microsoft Xcel: https://www.youtube.com/watch?v=0VtUQLbfewU

Google Sheets: https://www.youtube.com/watch?v=my7iSXe2kcw

https://www.youtube.com/watch?v=0VtUQLbfewU

https://www.youtube.com/watch?v=my7iSXe2kcw

35

Part 3: Questions

1a) On your graph you plotted mass on the y-axis and volume on the x-axis. What physical

property of the twizzler can be determined from the slope of your graph?

1b) On the class graph, the spreadsheet program calculated the line of best fit for the data.

According to the best fit line equation (expressed in the form of y = mx + b, where m is the

slope), what is the slope of the line?

Big Idea: Why do scientists consider collecting multiple data points more reliable than a single

measurement?

2A) Are all of the class data points exactly on the line of best fit? Are there any data points which

seem to be outliers, i.e. that are far off the line of best fit?

2B) The square of the correlation coefficient abbreviated as the r2 value describes how well your

data points match the line of best fit. The value of r2 describes how well the value of x predicts

the value of y. r2 values range from 0.00 (no predictive value) to 1.00 (perfect predictive value).

Scientists generally like to see an r2 value of 0.95 or higher to be very confident in the reliability

of the data. What is the r2 value for your graph? Is the r2 for our class data 0.95 or above?

2C) Assuming that the slope of the line of best fit is an accurate representation of the density,

does the ratio of mass divided by volume for every single individual point give the same value of

density?

2D) Which value, your calculated value from a single data point, or the calculated value from a

class set of data would generally be considered to be more reliable? Explain. (Hint: How could

you tell from a graph if one data point was suspect?)

3a) Does changing the size (i.e the length) of your twizzler change the mass?

3b) Does changing the size (i.e the length) of your twizzler change the Density (mass/volume)?

Explain.

4) Once a line of best fit has been determined, you can use the equation to predict the outcome of

measurements that are difficult to make. Consider the example below.

A candy shop in Pennsylvania is known for having the Guinness world record for longest

twizzler. It is a whopping 370. m long.

A) Given that 1.00 m of twizzler = 38.8 mL, use dimensional analysis to determine the volume of

a 370. m length to volume of the twizzler in mL.

B) Use density value determined from your graph and the volume of the twizzler calculated in

part A, to determine the mass in grams of 370. m twizzler.

36

A Taste of Molarity

Background Information

Molarity is the moles of a chemical substance (solute) dissolved in a specific volume of

solution in liters:

It represents the chemist’s way of preparing different strengths of a chemical substance. In order

to accomplish this objective, we first pretend that the substance KOOL AID is a chemical

formula. Next, we will assume that the letters represent chemical symbols, each with a fictitious

atomic weight assigned to it: K = 14 ; O = 8 ; L = 10 ; A = 4 ; I = 6 ; D = 12

Activity Directions: Record all calculations and observations on a separate sheet of

paper.

1. Before preparing any solutions, you need to determine the molar mass of Kool Aid.

Calculate the grams/mole of KOOL AID using the above imaginary atomic weights.

2. Next, four groups of students will each make a different molarity of Kool Aid

solution. After each group has made their molar solution, taste observations are recorded.

Groups then rotate and record taste observations of all solutions.

Group A Prepare a 3.00 M solution of Kool Aid.

Calculate the amount of Kool Aid needed to make 1.00 L of a 3.00 M solution.

Show calculations! Record Taste Observations.

Group B Prepare a 1.00 M solution of Kool Aid.



Group C Prepare a 0.500 M solution of Kool Aid.



Group D Prepare a 0.100 M solution of Kool Aid.

Calculate the amount of Kool Aid needed to make 1L of a 0.1M solution. Show

calculations! Record Taste Observations.

Additional (required) Problems: Copy question, write equation (definition) of Molarity, show work, and include answer with sig

figs and units!

1) What mass of Kool Aid would be needed to prepare 500. mL of a 2.5 molar concentration?

2) What volume of solution would you need to prepare a 6.0 M solution if you were using 200.

grams of Kool Aid?

3) If you dissolved 137.65 g of Kool Aid in 250. mL of water, what molarity of solution would

you have?

Molarity = Moles

Volume (in L)

37

Preparing Molar Solutions Introduction: One of the most common tasks that chemists perform is the preparation of a molar solution from

either a solid solute or from a dilution of a concentrated stock solution. In this lab, you make

calculations necessary to prepare solutions and then use the directions you developed to prepare

the solutions in the lab.

Purpose: To prepare aqueous molar solutions from solid solute and by dilution of concentrated

stock solutions.

Safety:

Wear safety glasses. Wash hands if you contact any chemicals.

Prelab: On a sheet of notebook paper, write a Title, Purpose, and Safety ( Lastly, complete all 8 stars

below, calculating the necessary information (watch sig figs!) and writing the description

(“recipe”) of how to make the solution.

Any calculation or recipe with a beside it must be completed before going into

the lab. Hints are provided for # 1-3. Be sure to skip a few lines whenever you see the

asterisks **************** that separate the six stations.

************************************************************************

Station 1 Prepare 500. mL of 0.200 M NaCl from solid NaCl.

Calculation of grams of NaCl : Hint: [(Vol in L) x Molarity = moles; moles x g/mole = g]

Recipe Statement: Dissolve g of NaCl in enough water to prepare mL of

solution.

Picture diagram of steps:

Instructor Initials for Solution Preparation ________

******************************************************************************

Station 2 The reddish-pink solution contains 50.0 g of CoCl2 dissolved in water.

Record the volume of the solution and calculate the molarity.

Volume in mL:

Calculation of Molarity: M = moles of solute

Liters of solution

Instructor Initials for Solution Calculation ________

mL

38

Station 3 The clear blue solution contains a 5.00 M stock solution of CuCl2.

Prepare 100. mL of 1.00 M CuCl2 solution by dilution of the stock.

Calculation of Volume (of 5.00 M CuCl2) stock solution needed:

Recipe Statement: Dilute mL of M CuCl2 solution with enough water to prepare

mL of solution.



******************************************************************************

Station 4 The clear green solution contains a 2.00 M stock solution of NiSO4.

Prepare 100. mL of 0.500 M NiSO4 by dilution of the stock solution.

Calculation of Volume (of 2.00 M NiSO4) stock solution needed:

Recipe Statement: Dilute mL of M NiSO4 solution with enough water to prepare

mL of solution.



******************************************************************************

Station 5 Prepare 250. mL of 0.0500 M sucrose, C12H22O11, solution from solid sucrose.

Calculation of grams of sucrose: Dissolve g of sucrose in enough water to prepare

mL of solution.

Recipe Statement:



******************************************************************************

Station 6 This clear, colorless solution contains 42.0 g of NaHCO3.

Record the volume of the solution and calculate the molarity of the solution.

Volume in mL: Calculation of Molarity:

Instructor Initials for Solution Calculation ________

mL

39

Does increasing the amount of one reactant always increase the

amount of product? Demonstrations to investigate the concept of

limiting reactant and optimal ratio

Part #1: Reaction of copper(II) chloride and sodium oxalate

CuCl2(aq) + Na2C2O4(aq) → CuC2O4(aq) + 2 NaCl(aq)

Demonstration #1: In this experiment different ratios of sodium oxalate and copper

(II) chloride solution are combined.

Experiment # Moles

Na2C2O4

Moles CuCl2 Moles of

product

CuC2O4

Final Observations

1 2.00 10.0 2.00 Blue ppt; liquid

supernatant clear and

colorless

2 4.00 8.00 4.00 Blue ppt; liquid


colorless

3 6.00 6.00 6.00 Blue ppt; liquid


colorless

4 8.00 4.00 4.00 Blue ppt; liquid

supernatant is BLUE

5 10.0 2.00 2.00 Blue ppt; liquid

supernatant is BLUE

Data Analysis: Construct a bar graph with ratio plotted on the x-axis and product height

mm

1:5 2:4 3:3 4:2 5:1

Moles of product CuC2O4

Mole Ratio Cu+2: C2O4-2

40

Questions: 1) Which ratio produced the greatest amount of product?

2) Compare the coefficients of the balanced chemical equation with the optimal ratio from

your data.

Cu(NO3)2 (aq) + Na2C2O4 (aq) → CuC2O4(s) + 2 NaNO3 (aq)

3) Draw a diagram to represent the reactions that took place at each ratio.

Use Cu+2 = Use C2O4-2 =

Draw a circle around combinations that can find a partner ion to react with to form

CuC2O4:

1:5 2:4 3: 3 4:2

4) Why does not simply increasing increase the amount product? Explain

the concept of optimal ratio for a reaction. Use the term limiting reactant and

excess reactant in your answer.

Ratio:

5:1

41

Demonstration #2: In this experiment different amounts of sodium oxalate

are added to a constant amount of copper (II) chloride solution.

Initial observations:

CuCl2 is a clear, blue solution. The blue color is due to the presence of Cu+2 ion dissolved

in water.

Na2C2O4 solution is a clear, colorless liquid.

Summary of Results and Final observations:

Experiment # Moles

Na2C2O4

Moles CuCl2 Moles of

product

CuC2O4

Final Observations

1 3.00 1.00 1.00 Blue ppt; liquid


colorless

2 3.00 2.00 2.00 Blue ppt; liquid


colorless

3 3.00 3.00 3.00 Blue ppt; liquid


colorless

4 3.00 4.00 3.00 Blue ppt; liquid

supernatant is BLUE

5 3.00 5.00 3.00 Blue ppt; liquid

supernatant is BLUE

Construct a Bar Graph of Moles of CuC2O4 moles on the y-axis CuCl2 moles on the

x-axis.

Questions:

1) How can you interpret the graph and the observations of the color of the liquid

supernatant (i.e. why is the supernatant colorless in expts 1-3 but blue in expts 4

and 5?)

42

2) Draw a diagram to represent the reactions that took place at each ratio.

Use Cu+2 = Use C2O4-2 =

Draw a circle around combinations that can find a partner ion to react with to form CuC2O4:

1:3 2:3 3: 3 4:3

3) Why doesn’t simply increasing increase the amount product? Explain the

concept of limiting reactant and excess reactant in your answer.

************************************************************************

Part 2: Potato Gun Demo – Using the reaction of butane and oxygen to

propel a ball across the room at very high velocity! Introduction: Butane, C4H10 is mixed with oxygen, O2, from the air in reaction chamber. Upon

ignition, the gaseous products are very hot and expand rapidly, increasing the pressure. The

pressure pushes the ball out of the barrel of the gun and fires it across the room.

1) Butane gas, C4H10 and oxygen gas, O2, rapidly burn. What type of reaction is this?

2) What two products are produced in this reaction pattern?

3) Write the BALANCED chemical equation for the reaction of butane and oxygen:

C4H10 + O2 →

HOW DOES INCREASING THE AMOUNT OF BUTANE CHANGE THE FORCE OF THE

EXPLOSION?

Background: The total pressure of the reactant gases, butane and oxygen in the reaction chamber

inside the gun before the reaction must always equal the atmospheric pressure. In our experiment

we will represent the pressure of gases in terms of equivalents of seconds of butane sprayed.

Seconds of Butane spray Equivalents of O2 Force of Explosion

2 13

8 7

Question: How can you explain the results of the demonstration?

Ratio:

5:1

43

How Efficient is Your Chemical Reaction? Determination of Percent Yield

Introduction One of the most important tasks of chemists is to synthesize new materials. Making new

chemicals in the lab requires a recipe, which is obtained from the coefficients of the balanced

chemical equation. The balanced chemical equation recipe can be used to calculate the theoretical

yield, which is the maximum amount of product that can be made from the given amounts of

reactants. When reactions are actually conducted in the laboratory however, often the amount of

product we are able to actually isolate (the actual yield), is less than amount predicted by the

theoretical yield calculation. In this experiment, you will carry out a chemical reaction, and by

calculating the theoretical yield and measuring your actual yield, calculate the efficiency of your

reaction. The reaction you will carry out is the reaction of copper(II) chloride and aluminum:

3 CuCl2 (aq) + 2Al (s) → 2 AlCl3 (aq) + 3 Cu (s)

Purpose: To determine limiting reactant in a chemical reaction in order to calculate the

theoretical yield. To measure the actual yield and then calculate the percent yield.

Safety:

• Wear safety glasses.

• CuCl2 is moderately toxic by ingestion and may irritate skin if contacted.

• Be careful touching the outside of the beaker during the reaction - it will get very hot!

Materials: - Aluminum Foil (optional scissors to cut foil to size)

- CuCl2 ∙ 2 H2O

- Two 50 mL beakers

- 10 mL graduated cylinder

- Electronic Balance

- Spatula

- 50 mL beaker

- 125 Erlenmayer Flask

- Funnel

- Filter Paper

- Squirt water bottle (wash bottle)

- Watch glass or glass square

- Drying Oven

- Test tube

- Wooden splint

- Match or ignitor

44

Procedure:

Part 1: Preparation

1) Weigh out a 0.10 g sample of aluminum foil in one 50 mL beaker. Tear the aluminum

foil into small pieces to increase the surface area.

2) In a separate 50 mL beaker, weigh out 1.05 g of CuCl2 ∙ 2 H2O.

3) Using a 10 mL graduated cylinder, measure out 10 mL of 0.10 M HCl.

4) Add the 0.10 M HCl to the beaker containing the CuCl2 crystals. Swirl or stir gently,

until the solid is completely dissolved.

5) Record your initial observations of the aluminum and the copper (II) chloride solution.

Part 2: Reaction BE CAREFUL NOT TO TOUCH THE GLASS – THE REACTION

WILL BE VERY HOT!

6) Gas collection step – once the CuCl2 solution is combined with the Al a gas will be

produced. You will trap the gas in the test tube and then test the flammability of the gas

with a burning splint. Make sure that the test tube, wooden splint and match or ignitor are

standing by – the reaction will be fast.

7) Pour the CuCl2 solution into the beaker container the aluminum. Invert the test and

hold it for approximately 20-30 seconds directly above the surface of the solution.

8) Have one partner light the wooden splint. Once the splint is lit, quickly flip the test

tube and insert the burning splint into the test tube. Record all observations of the splint

test and the main reaction.

Part 3: Isolation of Copper Product

9) DECANTING - once the reaction is complete (Bubbling is no longer visible,

aluminum foil is completely dissolved), pour off the excess the liquid from the reaction

into the empty 50 mL beaker, being careful making sure that no copper product is poured

out.

10) FILTRATION – Fold the filter paper and fit it into the funnel. Place the funnel and

filter paper into the 125 mL Erlenmayer flask. The funnel should sit on the opening of the

flask.

11) Pour the reaction solution through the filter. You will probably have to use a

combination of scraping the copper out with a spatula and a series of very small rinses

from the water bottle to transfer all of the copper from the beaker or spatula onto the filter

paper. Do the best you can to ensure that all of the solid copper from the reaction beaker

is transferred to the filter paper.

12) Record the color of the supernatant (liquid that passed through the filter paper).

13) Spread your filter paper out on a watch glass or glass and place into the drying oven

in the location directed by your instructor.

Next Day:

14) Retrieve your filter paper from the oven.

45

15) Weigh a new piece of filter paper or weighing paper. Scrape the copper onto the new

filter paper and record the final mass. Depose of all filter papers in the trash.

Lab Report:

Title, Purpose, Safety

Data:

Initial Observations: Al CuCl2 solution

Mass of Al (g): Mass of CuCl2 ∙ 2 H2O:

Final Reaction Observations:

Final Mass of Filter and Cu:

Mass of Filter Paper:

____________________________________

Mass of Cu

Calculations and Questions:

Part 1: Determination of Limiting Reactant:

Given: Molar Mass of Al = 26.98 g/mole Molar Mass of CuCl2 ∙ 2 H2O = 170.49

g/mole

3CuCl2 (aq) + 2Al (s) → 2 AlCl3 (aq) + 3 Cu (s)

1A) Calculate the theoretical number of grams of Copper, Cu, that could be

produced by the reaction of 0.10 g of Al.

1B) Calculate the theoretical number of grams of Copper, Cu, that could be

produced by the reaction of 1.05 g of CuCl2 ∙ 2 H2O. (Note: CuCl2 comes as dihydrate

crystal with the 2 waters attached; however converting the grams of CuCl2 ∙ 2 H2O into

moles gives the same number of the moles of CuCl2 present in your reaction.)

1C) Which reactant is limiting, CuCl2 or Al? Support your answer by referring to

the results of your calculation.

1D) Supporting lines of visual evidence: Cu+2 ion is a light blue color, Al+3 is colorless.

In the reaction the Cu+2 replaces the Al to form Al+3.

i) Does the blue color of the solution get darker or lighter during the reaction?

ii) Why doesn’t the blue color completely disappear at the end? (i.e, Why is there

still a small amount of blue color still visible?) How does this observation support your

conclusion as to the reactant present in excess and the limiting reactant?

1E) Based on your answer to part C, what is the predicted theoretical yield for this

reaction?

46

2) What is your actual yield of Cu?

3) Calculate the percent yield for the reaction:

4) Factors that impact percent yield

A) A group left some of the solid copper behind in the reaction flask. Would this cause

the final percent yield to increase or decrease?

B) A group’s final weighing of the filter paper produced an actual yield greater than the

predicted maximum yield of 100%. Assuming that all weighing steps and calculations

were performed correctly, propose an explanation for the apparent extra mass of

product.

C) Impact of side reactions: The products of the main reaction we are investigating are

solid copper and aqueous aluminum chloride, yet a gaseous product (hydrogen) is

observed. The hydrogen is produced by the following side reaction:

2 Al(s) + 6 H2O(l) → 3 H2(g) + 2 Al(OH)3 (aq)

If some of the Al reacts with water instead of CuCl2 how will the percent yield be

impacted?

5) Importance of percent yield – If our goal was to produce 0.50 g of Cu, but our

percent yield was only 50% so we only obtained 0.25 g, how could we change the

initial recipe in order to actually obtain the desired 0.50 g of product?

47

Determination of the value of the Universal Gas Constant

Introduction The ideal gas law, PV = nRT, where P = pressure, V = volume, n = # of moles and T =

Temperature in K is the most important of the gas laws. All of the other gas laws are special cases

of the ideal gas law. In this equation, R represent the universal gas constant. The role of the

constants in equations is to make units much on both sides. For the numbers to match on both

sides of an equation the units must match. To experimentally determine the value of R, we will

vary the value of nT and measure the effect on PV. The equation of a line is y = mx + b. For the

ideal gas law, the value of b = 0, so the equation becomes y = mx, where PV = y, nT = x and R

= m the slope. In other words, if we vary nT and measure the effect on PV, and plot the points the

slope of the line will give the value of R. As a class we will conduct two trials for 5 separate

conditions and then plot the data and determine the best fit slope.

Safety:

The hydrochloric acid is extremely corrosive to eye and skin tissue. Safety glasses must

be worn at all times during the experiment. If any liquid splashes in your eye

immediately rinse your eyes for 15 minutes in the eye wash.

Materials:

1.0, 1.5, 2.0, 2.0 and 2.3 cm lengths of Mg

Eudiometer tube

Buret clamp

Ring- stand

1000 mL beaker

1-hole stopper

Copper wire

Procedure:

48

1) As a class we need two to collect two sets of data for each of the following lengths collect

data using 5 different lengths of Mg: 1.0, 1.5, 2.0 and 2.5 and 3.0 cm.

2) Fill the 500 mL beaker 2/3 with water and place the beaker underneath the eudiometer

tube.

3) The eudiometer contains approximately 10 mL of 6 M HCl solution. Carefully add water

to completely fill the top of the tube. Make sure the tube is completely filled with no air

bubbles. If any water drips over the top it will fill into the beaker.

4) Construct a small cage using copper wire to hold the Mg in place. Secure the loop of Cu

wire by lodging it over the edge of the eudiometer tube.

5) Securely fit the 1 hole-stopper in place at the top of the tube. Place 1 finger over the hole

and flip and invert the stopper underneath the surface of the water.

Data: Mass of 100.00 g strip of Mg: 1.04 g

Length of Mg (cm): _________

Atmospheric pressure in mm Hg: _________

Height of water column (distance between top of meniscus and surface of water) in mm:

Volume of trapped gas (mL): __________ ÷ 1000.0 mL/ 1 L = __________ L

Temperature in oC: _________ + 273 = ________ K

Vapor pressure of water at _______ oC in mm Hg (from CRC handout):

Calculations: Part 1: Calculation of Pressure of trapped gas

1) Pressure difference between atmosphere and trapped gas (what would the

height be if the liquid was made of Hg instead of water). Recall that Hg is 13.5 times

as dense of water; thus if the liquid water were made of Hg instead its height would be

13.5 times less.

Height of water column in mm ÷ 13.5 mm Hg/mm water = __________ mm Hg =

Δh

49

2) Calculation of total pressure of trapped gas:

Since the height of the water in the eudiometer tube which is in direct contact with the

trapped gas is higher than the water level in the beaker which is in direct contact with the

atmospheric pressure, the pressure of the trapped gas is less than the atmospheric

pressure, so the pressure difference should be subtracted instead of added.

PT = total pressure of trapped gas

Δh = height difference between liquid levels measured in mm Hg (from

calculation #1):

Patm = atmospheric pressure

PT = Patm – Δh = - = ____________ mm Hg

3) Calculation of Pressure of hydrogen gas, PH2

The trapped gas is a mixture of mostly hydrogen, H2, and a small amount of water vapor

which evaporated from the liquid water surface. Thus to determine the pressure of the

trapped gas we will need to subtract off the vapor pressure of water.

• PT = PH2O + PH2

• PH2 = PT - PH2O = - = mm Hg

50

Part 2: Calculation of Moles of Hydrogen Gas:

In order to collect volumes of hydrogen gas that will fit in the tube, we need to use very

small masses of Mg – so small that our balance cannot accurately weigh them. In order

to determine the mass of the short length of Mg, the mass of 100.00 cm (1 meter) was

weighed. We can then use a proportion to calculate the mass of short lengths.

1) Calculate grams Mg

1.04 g = ________

Mg Mass = your expt length x 1.04/ 100.0 cm = your length x 0.0104

Mass of Mg = ________ cm x 0.0104 g/cm = __________ g Mg

2) Calculation of moles Mg:

( )

3) Moles of H2:

Balanced Chemical Equation: Mg + 2 HCl → MgCl2 + H2

Stoichiometry between Mg and H2 = 1:1 therefore moles Mg from previous step = moles

H2

moles Mg = moles H2

Part 3: Construction of Graph, PV vs. nT

y = mx → PV = nRT ; m = y/x → R = PV/nT

Units: P in mm Hg , V in liters, T in Kelvin

X value: (n)(T)

y value: ( PH2 )(V )

Class Data Set:

nT (mol x K) PV (mm Hg x L)

100.0 cm

Mg mass

Your expt length

g Mg

24.38 g Mg

Mg

1 mole Mg

= moles Mg

51

1) Plug in your values to class Google spread sheet.

2) Run line of best fit. Calculate slope.

3) Error analysis. Report correlation coefficient value squared (r2).

IMPORTANT: The correlation coefficient r2 is NOT the same as the value of the slope,

the universal gas constant R.

Questions: 1) What are the units for R, the universal gas constant?

2) The accepted value of R is 62.4. Calculate % error.

3) Ideally the correlation coefficient r2 value ought to between .95 and 1 for excellent data.

What is our value? How does this value relate to the amount of scatter in the data?

52

Emission Spectra and Quantum Leaps

Introduction Rutherford’s Gold Foil Scattering experiments published in 1911, established that the atom

consisted of a small, dense positively charged nucleus surrounded by an electron cloud. The

Rutherford model offered no details as to the behavior of the electrons and could not explain why

the negatively charged electrons were not pulled into the positive nucleus. Scientists trying to

design experiments to determine how electrons behaved inside the atom faced a daunting

challenge: how do you “see” an electron that is so tiny that it fits inside an atom and is moving

near the speed of light? The key insight was made by a Danish physicist Niels Bohr who realized

the emission of light from elements that had absorbed energy in the form of electric current, light

or heat could be explained by quantum leaps of electrons. In this investigation, we will look at

patterns of light released from different elements and understand how quantum mechanics can

explain the observations.

Purpose: To observe the spectra of white light and the line emission spectra of several

elements including hydrogen and to understand how your observations can be explained

by quantum mechanics.

Safety: No concerns although a giddy euphoria can be produced because the experiments

are so cool!

Materials:

Special spectral glasses fitted with prism lens that separate light into its individual

wavelengths.

High voltage power source fitted with spectral tubes filled with gas samples of different

elements.

Part 1: Spectrum of White light

Procedure: Put on your spectral glasses, look at a source of white light and sketch a

diagram of your observations in a box (see below).

White Light:

What colors do you see? Do the colors appear as distinct lines or do they appear to

blend? (In reality, although the colors appear to blend, they actual consists of thousands

of individual colors line so close together that your brain interprets the colors as

blended)

Part 2: Emission Spectrum of Element

Procedure: Put on your spectral glasses, look at the light source for each element and

sketch a diagram of your observations in a box (see below). What colors do you see?

Do the colors appear as distinct lines or do they appear to blend? Mark lines in the

appropriate color area

Helium

Neon

Hydrogen

Violet blue green yellow orange

red Violet blue green yellow orange

red

Violet blue green yellow orange

red

53

Questions: 1) How are the patterns of light emitted from the element samples differ from the pattern

of white light?

2) Are the line spectra patterns of the different elements the same or different? What does

it mean to describe the line spectra patterns of an element as “fingerprints”?

3) How do astronomers know that our sun is composed mainly of hydrogen and helium if

no one has ever been to the sun?

4) The fact that only particular individual wavelengths and frequencies of light are

emitted is taken as evidence that the light energies are quantized. What does the term

quantized mean?

5) Provide an example of quantized system and an example of a non-quantized system.

6) Bohr proposed the Bohr Model in 1913 which introduced the concept that electron

energies are quantized. A key feature of this model is the concept of the quantum leap to

explain the emission line spectra.

A) Importance of energy absorbance: No light emission is observed until the atom

absorbs energy in the form of heat, light or electricity. According to the Bohr model,

what part of the atom absorbs the light energy?

B) Which represents a lower energy, more stable state? (Does nature prefer oppositely

charged particles close together or far apart?)

Choose the appropriate word to complete each statement. (From POGIL, edited by Laura

Trout, Flinn Scientific).

C) Electrons and protons (attract/repel) each other.

D) As an electron get closer to the nucleus the (attraction/repulsion) to the nucleus gets

(stronger/weaker).

E) For an electron to move from an energy level close to the nucleus to an energy level

far from the nucleus it would need to (gain/lose) energy.

F) For an electron to move from an energy level far from the nucleus to an energy level

close to the nucleus it would need to (gain/lose) energy.

54

7).Which picture shows a quantum leap in which an electron is absorbing energy? Which

picture shows a quantum leap which releases or emit light energy? Label each picture

below as either ABSORPTION OR EMISSION.

8) Use the energy level diagrams below for questions. The y-axis represents energy.

8A) Consider the energy diagram for the hydrogen atom on the left. Is the energy

distance between all of the energy levels exactly the same? In other words, is the

amount of energy released when electron jumps from level 4 down to level 3, the

same amount of energy that would be released when an electron jumps from level 2

down to level 1?

8B) Are the energy gaps the same in both elements? In other words, would an

electron jumping from level 2 to level 1 in hydrogen, release the same amount of

energy as an electron jumping from level 2 to level 1 in a helium atom?

8C) The energy that an electron has in an energy level is determined by the attraction of

the electron from the positive protons and the repulsion the electrons experience from

other electrons. How would these interactions be different in a hydrogen atom vs a

helium atom? (Hint: How many protons and electrons in each type of atom?)

55

9) In the hydrogen spectra, the deep blue/violet line at 434 nm, the teal line at 486 nm and

the red line at 656 nm corresponds to quantum leaps from a higher energy excited state of

either 3,4 or 5 down to level 2. Which color line is produced by the 3 → 2 transition?

Which color line is produced by the 4 → 2 transition?

Which color line is produced by the 5 → 2 transition? EXPLAIN.

10) Each of the lines of color of the hydrogen emission is the result of a different

quantum leap. A hydrogen atom has one electron. How can all of these different leaps be

happening at once? How can all of these different colors of light be emitted at the same

time?

11) Why can we see only 3 lines of color emitted from the hydrogen spectra? Each

quantum leap from an excited state to a ground state should release energy in the form of

light (e.g. 2 → 1, 3 → 1, 4 → 1, etc.), yet we see only 3 colors of light from the 5 → 2, 4

→ 2, 3 → 2 transitions. Why don’t we see more colors of light?

n = 5

n = 4

n = 3

n = 2

56

Determining the Concentration of a Solution Using

Absorbance Spectroscopy (Beer’s Law) Purpose: To construct a standard curve of absorbance vs. concentration for 5 solutions of

known concentration; use your standard curve to determine the molar concentration of an

unknown solution.

Introduction

In nature, “pure” water does not exist; the water always contains substances dissolved in

it. One important job of analytical chemists is to test our drinking water to make sure that

the concentrations toxic substances (for example heavy metals like lead) are not present

in too high a concentration for us to safely consume. One important technique for

quantitatively measuring the exact concentration of solute present is to measure the light

absorbance of the solution: The higher the light absorbance, the higher the

concentration. This relationship is called Beer’s Law after its discoverer and can be

expressed as:

A = εcl where

A = absorbance,

ε = extinction coefficient (a measure of how strongly a particular substance absorbs light

at a given wavelength wavelength

c = the concentration measured in moles/liter (M)

l = the pathlength (the distance the light has to pass through the test tube holding the

solution.) Note in most experiments, the pathlength is 1 cm, and since the value is 1 it can be ignored in

the equation. Concept of Beer’s Law: The number of photons of light (intensity) of light is entering

and exiting a solution are measured and compared.

Background

2 / 5 photons absorbed 3 / 5 photons absorbed

In the figure at left: 5 photons of enter the test tube; 3 photons are detected as having

passed through; thus 2 photons out of 5 were absorbed. Concentration (solute

particles per volume) is directly to the proportional number of photons of absorbed.

57

When atoms absorb light energy, electrons make transitions or quantum leaps from lower

energy levels to higher energy levels. The energy absorbed in making the quantum leap

is specific to the atom. The energy of the transition is quantized and is specific and

unique for each individual atom or atom that is part of a compound or molecule.

Concept of colors of light absorbed vs. transmitted

There is an important distinction between the color of the light that see because it is

reflected or transmitted due and the color of light absorbed. For example, the absorbance

spectrum copper sulfate absorbs light strongly at 750 nm (red light) but absorbs very little

light in the blue region in the lower 400 nm range. Copper sulfate solutions appear blue

because they reflect the blue wavelengths while absorbing other colors of light such as

red.

Concept electron absorbing

energy within atom

Many possible jumps are present within atoms with many electrons, producing a more

complex absorbance graph (see graph above). Note that although copper sulfate looks blue, it

actually strongly absorbs red.

Red, blue

light

Blue transmitted

Red light

absorbed

Red light

absorbed

Blue light

reflected, not

absorbed

58

Safety

Wear goggles.

Materials

Beaker, 250 mL

Pipettes, 6

Cuvettes (test tubes), 7

Spectrophotometer

Tissues or lens paper

Wash Bottle

Data Table Concentration of Stock Solution

Blank 0 M 1 M 2 M 3 M 4 M 6 M Unknown

Color Comparison (Rank Solutions*) 0

____________

%T (measured from spectrophotometer)

Absorbance (calculated from %T)

* Rank Solutions from lightest blue = 1 to darkest blue = 5, mark the blank as 0

PROCEDURE:

Getting Started: 1. Turn on the spectrophotometer. The spectrophotometer is the instrument used to measure

light absorbance. It will take at least 15 minutes to warm up so look at the clock and record

the time. Do the rest of the steps in Part I while you wait for it to warm up.

2. Turn the wavelength knob on the spectrophotometer to 635 nm.

3. At your work station, you have 6 test tubes labeled with a molar concentration ranging from

0 to 6 M. Compare the colors of each of the solutions. Rank them on a scale of 1-5 with 1

being the lightest and 5 being the darkest and record the information in your data table.

59

Measuring Absorbance

Part 1: Blanking 1. The cuvette marked 0 M will serve as a blank. The blank contains the same solvent

as the samples you wish to measure (in our experiment water). In the blanking

process, you are instructing Spec 20 computer to assign a value of 0 to the

absorbance of anything that is not solute. Wipe the cuvette with a tissue, and then

place the cuvette in the spectrophotometer.

2. Press the 0 Abs/100%T button on the spectrophotometer to set 0 Absorbance.

3. Remove the “blank” (distilled water) cuvette from the spectrophotometer

compartment. REPEAT THE BLANKING PROCESS BEFORE EACH NEW

MEASUREMENT.

Part 2: Sample measurement

4. Wipe the cuvette of the sample to be measured with a tissue, and then place

the cuvette in the spectrophotometer.

5. Read the absorbance value off the screen and record it in your data table.

6. Repeat the above steps until you have recorded the absorbance values for all

of the known concentrations as well as the unknown.

LAB REPORT:

1. Your lab report should have Title, Purpose, Safety, Data Table, Graph and

Questions.

2. Plot a calibration curve using the absorbance values of stock solutions of known

concentration obtained from the spectrophotometer.

• Absorbance (unitless) on the y-axis versus Concentration (units of M) on the

x-axis. Note: The absorbance of the 0 M (the blank) should be 0.00.

• Use google or excel to draw a line of best fit through the data and determine the

r2, value (correlation coefficient) for the graph.

3. Use your graph of Absorbance vs. Concentration to estimate the concentration

of the unknown solution. See example below:

Example: Use the standard curve below to determine the alcohol

concentration of a solution with an absorbance of 0.40.

60

Answer: Draw a line across from the y-value absorbance until it intersects the line of best fit;

drop down from the y-value point to the x-axis to determine the x-value. In this example, the

concentration would be approximately 0.42 %.

QUESTIONS

1. a. If you are wearing a red t-shirt, what color(s) of light is/are it reflecting?

b. What color(s) is/are absorbed?

2. If a compound has the spectrum below what color is it?

3) What is the relationship between concentration and absorbance? (As concentration

increases, does absorbance increase or decrease?).

4) Use the absorbance of your unknown and line of best fit on your graph to determine

the molar concentration of your unknown. Draw lines across and down on your printed

graph to show how you determined the concentration of the unknown. (See example

above).

61

Investigating the Relationship Between the Type of Chemical

Bonds and Electrical Conductivity

Introduction

One of the major themes of our bonding unit is the how different bonding models (ionic,

covalent, metallic) explain the observed properties of different substances. In this lab, we

will investigate the electrical conductivity of different substances. Electric current

represents a flow of charged particles. For a substance to conduct electricity two

conditions must be met:

1) A charged particle, either a + or – ion or an electron must be present.

2) The charged particles must be able to move freely.

In this experiment, you will test the conductivity of different substances using a

conductivity tester. A conductivity tester is an LED light connected to a battery with an

open circuit of two wires. If the wires touch a conductor, the circuit will be complete and

the light will light up. After collecting your data, you should be able explain how

different types of bonding can be identified using electrical conductivity and be able to

explain the results in terms of the bonding models.

Prelab: Title, Purpose, Safety, Data Tables, Identify type of bonding based on formula of

substance.

Safety: Wear safety glasses; wash hands after experiment.

Conductivity Tester

62

Data Table

• Copy Data Table and Classify each substance as Ionic (metal + nonmetal), covalent (2

nonmetals) or Metallic (metal) Note: NA = Not applicable

Substance Type of

Bonding

(Ionic,

Covalent,

Metallic)

Conducts solid? Conducts liquid? Conducts aqueous

solution

(dissolved in

water)?

NaCl

KCl YES

KBr YES

CaCl2 YES

Na2CO3 YES

Al YES NA - insoluble

Zn YES NA – insoluble

Cu YES NA- insoluble

C12H22O11

“Pure” H2O NO NA

Ethanol,

C2H5OH (aq)

NO NO

Iso -Propanol

C3H7OH (aq)

NO NO

Note: Conductivity of molten sugar (C12H22O11) and molten salt, NaCl, will be observed on

video.

Procedure:

1) Make sure electrodes wires are clean, dry and not touching.

2) Place tips of electrodes into material tested and record observations.

3) For solids use a very small pile (about the volume of a pencil tip eraser).

4) To make aqueous solution, add small volume of distilled water.

5) For liquid solutions can test directly, don’t need to add water.

Questions: 1) Patterns in the data:

Bonding Conducts Solid? Conducts Liquid? Conducts aqueous

Covalent

Metallic NA - insoluble

Ionic

63

2) Define electric current: Electric current = _______ of _____________

Two conditions for conductor (both must be true)

i) AND ii)

3) Interpreting Results Covalent Compounds

A) Chose correct answer: Covalent compounds consistent of i) neutral molecules ii) + and

– ions iii) metal crystal with sea of electrons “mobile, delocalized electrons”

B) Liquid water molecules are able to move freely – why then does pure water not conduct

electricity?

C) Pure water does not conduct electricity. Why are tap water, ocean water, river and lake

water generally excellent conductors of electricity?

4) Interpreting Results – Metals

A) Chose correct answer: Metals consistent of i) neutral molecules ii) + and – ions

iii) metal crystal with sea of electrons “mobile, delocalized electrons”

B) What particle is carrying the charge when a metal conducts electricity?

5) Interpreting Results – Ionic Compounds.

A) Choose correct answer: ionic compounds consistent of i) neutral molecules ii) + and

– ions iii) metal crystal with sea of electrons “mobile, delocalized electrons”

64

Use the lattice model of sodium chloride present in SOLID salt to answer

questions B + C.

5B) Fill in the blank: Each ____ charged sodium ion is surrounded by and bonded to

______ _____ charged chloride ions.

5C) Are charged particles present? Can they move? Explain.

Use the representation of molten (liquid) NaCl to help answer question 5D.

5D) Are charged particles present? Can they move? Explain. (What is different than

the solid form?)

Use the representation of an aqueous solution of sodium chloride below to answer

part 5E.

5E) Are charged particles present? Can they move? (Be sure to explain the

interaction between water molecules and sodium and chloride ions).

65

Molecular Models

Introduction: Two-dimensional Lewis structures are a simple and convenient method of showing which atoms

are directly connected within a molecule. However, to accurately understand the properties of a

molecule we also need to know its three-dimensional shape or molecular geometry. The

molecular geometry can be predicted using Valence Shell Electron Pair Repulsion Theory

(VSEPR) by arranging bonding and nonbonding electron pairs as far apart as possible around the

central atom.

Purpose: To use VSEPR theory to predict the geometry of a molecule, predict the

bond angles and draw a 3-D picture of the molecule.

PreLab: 1. Make a table just like you see on the next few pages for the 9 chemicals (molecules or ions).

Be sure to leave enough space for the Lewis Dot Structure and 3-D Shape (Drawing)

between tables.

2. Draw the Lewis Dot Structures for all nine chemicals.

Procedure: 1. Count/record the number of bonding and nonbonding electron pairs around the central atom.

2. Using your VSEPR handout, determine and record the geometry of the central atom. Be sure

to include lone pairs in your decision.

3. Draw a 3-D picture of the molecule using the conventions below:

SOLID LINE represents a bond in the plane of the paper.

DOTTED LINE ---------- represents a bond pointing away from the observer

into the plane of the paper.

SOLID represents a bond coming out of the plane of the paper.

TRIANGLE

represents nonbonding valence electron pairs on the central atom

Example:

66

Chemical Bonding

e- Pairs

(Regions)

Non-

Bonding

e- Pairs

Geometry

of Central Atom

Bond

Angles

Teacher

Approval

HCN

Lewis Dot Structure: 3-D Shape (Drawing):

Chemical Bonding

e- Pairs

(Regions)

Non-

Bonding

e- Pairs

Geometry

of Central Atom

Bond

Angles

Teacher

Approval

BF3


Chemical Bonding

e- Pairs

(Regions)

Non-

Bonding

e- Pairs

Geometry

of Central Atom

Bond

Angles

Teacher

Approval

CF4


Chemical Bonding

e- Pairs

(Regions)

Non-

Bonding

e- Pairs

Geometry

of Central Atom

Bond

Angles

Teacher

Approval

PCl5


Chemical Bonding

e- Pairs

(Regions)

Non-

Bonding

e- Pairs

Geometry

of Central Atom

Bond

Angles

Teacher

Approval

H2O


67

Chemical Bonding

e- Pairs

(Regions)

Non-

Bonding

e- Pairs

Geometry

of Central Atom

Bond

Angles

NH3


Chemical Bonding

e- Pairs

(Regions)

Non-

Bonding

e- Pairs

Geometry

of Central Atom

Bond

Angles

BeCl2


Chemical Bonding

e- Pairs

(Regions)

Non-

Bonding

e- Pairs

Geometry

of Central Atom

Bond

Angles

SF6


Chemical Bonding

e- Pairs

(Regions)

Non-

Bonding

e- Pairs

Geometry

of Central Atom

Bond

Angles

NO2-1


68

Investigating the Behavior of Polar and Nonpolar Molecules

How Does Bonding and Molecular Geometry determine the behavior of a molecule

in an electrical field?

Introduction

One of the main themes of the bonding units is the relationship between the structure of a

molecule and its observed properties. In this activity, we will investigate how the chemical bonds

and shape of a molecule affect its behavior in an electric field. The electrical properties of a

molecule determine how molecules attract other molecules or ions. We will model interactions

between water molecules and ions using magnetic models.

Part #1: Review of electrical charges using Phet simulation of balloons and charges

1) When a balloon is rubbed against the sweater what type of charge does it acquire?

2) Why does the balloon stick to the sweater?

3) What happens to the distribution of charges on the wall when the negatively charged side

of the balloon is held near the wall?

4) What happens when a charged balloon is placed near a neutral balloon?

5) How can the two balloons be made to attract each other?

Part #2: Demo: Investigating the properties of water A wand can be given a negative charge by rubbing it with rabbit fur. (You can also do this

experiment by rubbing a balloon in your hair to give the balloon a charge.) In your group, have

someone hold the charged wand or balloon right next to a thin stream of water from a buret

without actually touching the water. Try moving the charged wand or balloon to different

positions near the water.

1) Describe what happens when the charged wand or charged balloon is held near a thin

stream of water. Sketch a picture of your observations.

2) Since the wand or the balloon is charged negative, what do you think the charge on the

water must be?

3) Predict what would happen if you held a positively (+) charged wand or balloon near the

stream of water.

Part 3: Using Magnet Models to investigate how water molecules attract each

other: In your group, hold models of water molecules containing magnets near each other. The

oxygen atom in H2O is red and the hydrogen atoms are white. Which surfaces attract each

other? What must be true about the charges on these surfaces?

Part 4: Using Magnet Models to Investigate how water molecules are able to

pull + and – ions out a salt crystal. Using your magnetic models of a water molecule

and sodium and chloride ions; determine which water surfaces attract the green negative

chloride ions and which attract the blue positive sodium ions. Sketch a diagram showing 4

water molecules surrounding a + sodium ion, and 4 water molecules surrounding a –

chloride. Make sure you have the correct side (H or O) facing the ion. The hydrogen atoms

are white in the models and the oxygen atoms are red.

69

Part 5: Chem Study Video

1) Which of the following streams of liquid are deflected by a negatively (-) charged wand,

Water, acetone, carbon disulfide, benzene?

2) Since opposite charges attract and like charges repel, a reasonable hypothesis to explain

the observations would be that water and acetone are positive charged liquids.

A) If water and acetone are positive liquids, then how should the streams behave when

a positively charged (+) wand is held near the stream?

B) Describe what actually happens when a positively (+) charged wand is held near a

stream of water.

3) Summary of Chem Study Video Experimental Results. Sketch path the stream of

each liquid takes as it passes past a charged rod. Label molecules that are

attracted to the charged wand as POLAR and molecules that are not attracted to

the charged wand as NONPOLAR.

H2O

Acetone Carbon

Disulfide

Benzene

4) Water is overall electrically neutral. Explain using both a diagram and words, why a

stream of H2O is attracted to both a + charge wand and a – charged wand, even

though H2O is overall electrically neutral. Use to represent a

water molecule.

5) Use arrows to show the directions of the bond dipoles in the H-O bond in water and then an

arrow to show the direction of the overall or net dipole for a water molecule.

+

-

70

Molecule Polarity using PhET Purpose: To explore with computer simulations the concept of molecular polarity and

dipoles.

Experimental Controls

• In the View box, if you click Bond Dipole, an arrow will point toward the partial

negative side of the bond. Note the other side of the arrow had a + on it that is

closer to the partial positive side of the bond.

• In the View box, if you click Partial Charges, the δ+ and δ- will appear.

• In the View box, if you click Bond Characters, a scale that indicates “how

covalent vs. how ionic the bond is” will appear below the electronegativity boxes.

Note that two atoms with identical electronegativites create a bond that is

_____________ and that two atoms with the extreme opposite electronegativites

create a cond that is _______________. The scale that we use in class (ΔEN of 0

= ______________, ΔEN of 0.1 to 1.7 = ___________________, and ΔEN>1.7 =

_____________) is a general scheme to help determine bond character.

• Leave the Surface box clicked as None.

• Turn the electric Field On.

Two Atoms

Start by ensuring that the above boxes are set. Atom A EN should be at the far left (less)

and Atom B EN should be at the far right (more). Bond Character should be Ionic.

1. Why does the bond dipole (black arrow) point towards the + plate?

2. Completely switch the EN values of Atoms A and B. What happens to the molecule?

Why is Atom A closer to the positive plate now?

3. Place the cursor over Atom A and you’ll see two white curly arrows. Rotate the

molecule so that A now faces towards the negative plate. What happens and why?

4. Set the EN difference between the atoms to 2 units; anywhere on the scale is fine.

Rotate the molecule to face 180 degrees opposite of its direction. Why does it rotate back

more slowly than in question #3 above?

Three Atoms

Start by clicking Reset, which should default to less EN for Atoms A and C, a middle EN

value for Atom B.

1. Turn on Bond Dipoles and Partial Charges. Why do the two black arrows,

indicating the bond dipoles, point towards Atom B?

Note the partial charges are smaller for atoms A and C, while the partial charge for B is

larger due to more electron density around it.

71

2. The gold arrow pointing directly up from Atom B represents the Molecular Dipole.

This is different from the Bond Dipoles, hence the color difference. The Molecular

Dipole is essentiall the combination of all the Bond Dipoles. Another term that we use

for Molecular Dipole is “Net Dipole.” Turn on the Electric Field. Why did the

molecule rotate to point the Molecular Dipole toward the + plate?

In molecules with 2 atoms, if there is a Bond Dipole, then the molecule is polar. In molecules

with 3 or more atoms, if there is a Molecular Dipole, then the molecule is polar.

3. Increase Atom B EN all the way to “more.” What happens to the Bond Dipoles and

the Molecular Dipole?

Try rotating the Molecular Dipole to point toward the – plate by rotating Atom B. Just

like in the two atom molecule, the three atom molecule rotates back to point the – side of

the molecule toward the positive plate.

4. Try rotating atom A on top of Atom C. Briefly describe what happens to the Bond

dipoles and the Molecular Dipole.

5. Slowly rotate Atom A to a position 180◦ from Atom C and watch the sizes of the Bond

dipoles and the Molecular Dipole.

Why do the Bond Dipoles stay the same size?

Why does the Molecular dipole drop to zero?

Real Molecules

Start by clicking all the boxes in the View box. Below are a series of questions for each

molecule.

1. HF What is the geometry of any two atom molecule? ________________

Is HF polar? ___________

2. N2 What type of bond does nitrogen have: single, double, or triple? _______

3. O2 What type of bond does oxygen have: single, double, or triple? ________

Like nitrogen, is oxygen polar or nonpolar? _________________

4. H2O Water has two bond dipoles pointing toward the __________ atom. The

H atoms are partial _____ and the O atom is partial _____. The molecule

has a geometry of _______________ (any of the 3 terms are fine). Try

rotating the molecule by clicking and holding an atom or arrow. This will

give you a 3D perspective. Lastly, with the molecular dipole pointing up,

turn on the Electrostatic Potential box. Usually, we’ll leave this off since

it’s hard to see the other arrows and numbers.

72

5. CO2 Does carbon dioxide have 2 bond dipoles? ________

Which way do they point? ______________________

Why does the molecule have NO molecular dipole? __________________

Is CO2 polar or nonpolar? _________________

6. HCN List the three EN values for the atoms: H = _____ C = _____ N = _____

Which direction do the two Bond Dipoles point? ____________________

Is there a Molecular Dipole? ______ Which way does it point? _______

7. NH3 The three bond dipoles all point towards the N atoms. The Molecular

Dipole points “up” through the N atom towards the lone pair that is not

shown. Is ammonia polar? ______

What is the geometry of a 3,1? ___________________________

8. BH3 Even though the trigonal planar BH3 has three bond dipoles, it is nonpolar.

Why?

9. CH2O For this molecule, simply turn on the Electrostatic Potential. What does

the red region represent?

Why is the region around the O red and the region around the Hydrogens

blue?

The next six molecules are all 4,0 with is _________________ geometry. Some are polar

while others are nonpolar. Fill in the table below. Rotating the molecules may help.

Molecule Net or

Molecular

Dipole?

Is the molecule

polar or

nonpolar?

If the molecule is polar, which direction

does the Net of Molecular Dipole point?

CH4

CH3F

CH2F2

CHF3

CF4

CHCl3

73

Polar and Nonpolar Molecules Stations Lab

Purpose: To use molecular models and electronegativity values to determine whether a

particular molecule is polar (net dipole) or nonpolar (no net dipole). EN values: C = 2.5 H = 2.1 N = 3.0 O = 3.5 Br = 2.8 Cl = 3.0 F=4.0 Si=1.8

As=2.0

Station 1: Two of the gases that you breath in air are nitrogen (N2) ~ 80% and oxyen

(O2) ~ 20%. Draw Lewis Dot structures below and indicate whether they are polar or

nonpolar by circling the word.

Nitrogen: Oxygen:

Polar / Nonpolar Polar / Nonpolar

Station 2: Draw the Lewis Dot for NBr3, write the geometry, sketch the 3-D shape.

Indicate the bond dipoles with small arrows and the overall polarity of the molecule with

a NET dipole arrow. Is the molelcule polar or nonpolar?

Lewis Dot: 3-D sketch:

Geometry: ________________________ Polar / Nonpolar

Station 3: CHCl3: The black sphere is C, the red spheres are Cl, and the white sphere is

H. Draw the Lewis Dot for the molecule, write the geometry, sketch the 3-D shape.





Station 4: Benzene, C6H6, is a circular molecule with carbons in the center and

hydrogens on the outside. The alternating double and single bonds give the compound

added stability. Draw bond dipoles for each of the six C – H bonds, indicating partial

charges. Is there a NET dipole moment? ______ Is the molecule polar or nonpolar?

____________

Lewis Dot of Benzene (C6H6):

74

Station 5: The two structures below have identical formulas, C2H2Cl2, but different

structures (chemists call them “isomers”). The structure with both chlorine atoms (green

spheres) on the same side of the double bond is termed “cis” while the structure with the

chlorine atoms diagonal from each other is termed “trans”. Indicate whether each of the

two molecules has a net dipole moment and circle polar or nonpolar.

trans cis

polar / nonpolar polar / nonpolar

Station 6: Draw the Lewis Dot for H2CO, write the geometry, sketch the 3-D shape.





Station 7: Draw the Lewis Dot for SiCl4, write the geometry, sketch the 3-D shape.





Station 8: Draw the polar covalent Lewis Dot for AsF5, write the geometry, sketch the

3-D shape. Indicate the bond dipoles with small arrows and the overall polarity of the

molecule with a NET dipole arrow. Is the molelcule polar or nonpolar?

75

Acid-Base Titration

Introduction It is often critically important in chemistry to know the exact molar concentration of

ingredients in a solution. For example, the concentrations of salts and glucose in the solutions of

intravenous bags administered in hospitals must be known extremely accurately in order to ensure

patient safety. Commercial solutions such as vinegar must be checked to ensure uniform

composition. Titration is a technique for determining the number of moles of one chemical by

reacting it with a known amount of a second chemical. In our experiment, we will determine the

exact concentration of acetic acid in vinegar, by reacting with vinegar with a known

concentration of the strong base, sodium hydroxide.

Purpose: The goal of this experiment is to determine the molar (moles/L) concentration

of acetic acid, CH3COOH, in a solution of commercial vinegar.

Safety: 1) The sodium hydroxide solution is very corrosive to eyes and skin. Safety Goggles

must be worn at all times.

2) If NaOH contacts your skin, immediately rinse affected area with water. If you

get NaOH in your eyes rinse your eyes for 15 minutes in the eye wash and

immediately notify instructor.

3) Handle burets with care; they are delicate and expensive.

4) Wash bottles are to be used ONLY for rinsing the inside of the Erlenmayer flask.

Atomic Level Representation:

Acids = H+ ions in water, clear and colorless Bases = OH- , clear and

colorless HInd = initial protonated (H+ attached) form of indicator, clear and colorless

Ind - = indicator after reaction with base, pink and clear

Initial Add OH- , not at endpoint

At Endpoint Add OH-, just past

endpoint

Animation Link

H+ H+ H+

HInd

H+ + OH- = H2O

H+ Hind H+

H2O H2O

H2O

Hind

H2O H2O

H2O

Ind- (pink)

http://www.mhhe.com/physsci/chemistry/animations/chang_7e_esp/crm3s5_5.swf

76

Prelab Questions: 1) The concentration of acetic acid, CH3COOH, in vinegar is given on the bottle label

as 5.0% mass/volume.

5.0 % m/v = 5.0 g acetic acid

100. mL solution

Convert 5.0 g/100 mL into units of Molarity = moles of acetic acid

Liters of solution

2) Write the balanced equation for the reaction of acetic acid, CH3COOH and

sodium hydroxide, NaOH.

Important hints: This is an acid-base reaction; What is an acid? What is a base?

Which is which? NaOH is an ionic base – don’t forget to separate it into ions

before reacting it.

3A) Acetic acid, sodium hydroxide, sodium acetate and water are all clear and colorless.

What color is the Phenolpthlalein indicator in acidic solution and what color is it in basic

solution?

3B) What is the endpoint of a titration? What change will you observe in the solution

when the endpoint is reached.

3C) Draw 3 atomic level pictures:

Hints: Recall H+ + OH- → H2O, once all H+ has reacted, then

HInd + OH- → Ind- + H2O

Note: Hind = colorless, Ind- = colorless

Picture #1: Label as Initial Solution of Acid : draw 2 H+ ions, and 1 HInd molecule.

Picture #2: Add 2 OH- ions; show how they would react with H+ ions

Picture #3: Add 1 more OH- ion to picture #2 and show how it would react with Hind to

produce a color change. :

The titration equation for this experiment is: MacidVacid = MbaseVbase

Where Macid = Molarity of acid

Mbase = Molarity of base

Vacid = volume of acid (mL) = Volume Final – Volume Initial

Vbase = volume of base (mL) = Volume Final – Volume Initial

77

Reading Buret and Calculating Volume of base, VB delivered into flask:

Example:

Initial Volume = 15.72 mL Final Volume = 26.75 mL

Volume Delivered into flask:

VF: 26.75 mL

– Vi :15.72 mL

VB = 11.03 mL

Materials: 2 burets (one labeled “A” for Acid and another labeled “B” for base. 250 mL Erlenmayer flask

Beaker for waste

Wash bottle

Phenolpthlalein indicator

0.50 M NaOH solution

vinegar (acetic acid: CH3COOH), approximately 10 mL per trial

78

Procedure: Each student is responsible for performing at least 1 titration during the

lab period. Average the best two results for the group based on the the trials with the

lightest pink color at the endpoint. 1) The buret labeled “A” will be filled with acid(vinegar) and the one with “B” with base

NaOH. You will fill the acid buret; your instructor will fill your base buret.

2) Record the molarity of the NaOH from the board in your Data Section.

3) Fill each buret with proper solutions. Make sure that the initial liquid level is below 0.00

mL mark, so that the exact volume can be read.

4) Record the initial reading of BOTH the vinegar and NaOH using the following

procedure:

➢ Position your eyes at the same level as the meniscus (the curved part of the

liquid) and read the BOTTOM of the meniscus to the HUNDREDTHS PLACE.

(Do NOT subtract these values from 50 mL).

5) Add about 50 mL of water from the sink the flask.

6) Let approximately 10 mL of acetic acid flow into the erlenmayer flask. Rinse the tip of

the buret with the water from the wash bottle into the flask. Record the final reading for

the vinegar to the hundredths place. Calculate the exact volume of vinegar that was

drained into the flask.

7) ADD 2 DROPS OF PHENOLPHTHALEIN.

8) Begin to dispense base into the flask with vinegar. In the beginning of the titration, you

can let base flow quickly. As the pink color starts to linger, slow down the flow. BE

VERY CAREFUL AS YOU APPROACH THE END POINT (the point that the solution

remains pink). Add the base DROP BY DROP. Occasionally, wash down the sides of the

flask and rinse the tip of the buret to flush any vinegar back into the solution. The

endpoint is when the LIGHT pink color remains for at least 20 seconds.

9) Record the final volume of the base and calculate the volume of base used.

10) Leave solutions in the buret for the next class. Rinse all other glassware after solutions

have been poured down the sink.

Calculations:

1) Calculate Volume of Acid, VA, and Volume of Base, VB used in each trial.

(Recall: VA or VB = Vf – Vi)

2) Using the equation MA= MBVB, solve for the Molarity of Acid for trial.

VA

3) Average the results for Molarity of Acid for your trials.

4) Calculate the % error using the equation

% error = Accepted value – your experimental value x 100 %

Accepted value

Where accepted value = your calculated answer in M for prelab question #1

79

Experimental Design: Specific Heat Capacity and Density

Introduction In this lab, you will develop your own procedures for determining the heat capacity and

density of metal washers. By comparing the values you find for these properties with

those in the literature you can identify the metal of the washers.

Important Procedural Information • You may work in groups of 4 or 5 students.

• You need to perform at least 2 trials for each experiment (2 for density & 2 for

specific heat).

• Before entering the lab, complete the 10 PreLab questions listed below.

Note: you will need two procedures: one for density and a separate procedure

for the specific heat capacity part of the lab.

PreLab Questions for Specific Heat Capacity

1. What is the equation for heat transfer? Make a table listing all of the equation’s

variables.

2. Which variable(s) are you trying to determine in this experiment? Which variables

from the equation can you measure in the lab?

3. Maximizing the change in temperature ~ instructor will go over this.

4. How are you going to determine the temperature of the metal washers?

5. What can you do to ensure uniform heat transfer? In other words, how can you ensure

that

all washers reach the same temperature, instead of some washers being hotter or colder?

6. What can you do to minimize heat loss to the surroundings?

7. Does heat transfer occur instantly? Consider the following example. Two pieces of

metal with identical composition, mass, and the same initial temperature are to be placed

in a boiling water bath at 92 oC. The first piece of metal is held in the bath for 3 seconds

and then removed. The second piece of metal is placed in the water bath for 10 minutes

and then removed. Are both metal pieces at the same temperature when you remove

them from the water bath?

80

PreLab Questions for Density

8. What is the density equation? What variables will you have to measure?

9. How will you measure these variables?

10. How can you vary the number of washers so as to minimize error in both mass and

volume?

Lab Report Requirements Experimental Design:

Specific Heat and Density

General Information • Your lab report must be TYPED.

• Include your name, your partners’ name(s), your teacher’s name, and the class period.

• If one or more members of your group didn’t contribute equally to the report,

privately talk to your instructor.

• We will “pregrade” the reports in class; this will involve checking others’ reports to

ensure that all sections are included and adequately covered. Reports that don’t

contain all the information listed below or that aren’t “B” quality overall will be

returned for additions and/or corrections.

Specific Information Each lab report must include the sections listed below. Be sure you include the Roman

numeral and title for each section so that the reader can readily distinguish the various

parts of the report. Lastly, separate the two parts of the experiment within each section

(beginning with Section II ~ Materials); that is, keep the specific heat part separate from

the density part.

Abstract

This section contains a brief overview or summary of the experiment’s objectives

AND the experimental approach used to achieve these goals. In addition, state

your results and conclusions; that is, provide your average specific heat

capacity/average density results and identify the metal. Below is a sample

abstract for our Molar Volume of Hydrogen at STP lab; this is an example, not

your actual abstract.

The objective of this experiment is to determine the molar volume of hydrogen under standard conditions of temperature and pressure. The procedure involves displacing a volume of water with hydrogen gas (produced from the reaction of hydrochloric acid and magnesium). The mass of the metal is known, as is the volume of hydrogen produced. Using the combined gas law, we converted our experimental conditions of temperature and pressure to those of STP; using the stoichiometry of the reaction we calculated the number of moles of hydrogen produced. Combining moles with volume gives the molar volume at STP, a value we calculated to be 22.4 L.

81

• Materials

List all equipment and items that you actually used for your experiment.

Organize this list according to the two parts of the procedure; that is, list the

materials for the specific heat part and then the materials for the density

measurement.

Procedure

Describe, in detail, each step you followed in completing the experiment. Try to

write the procedure so that a person not familiar with the lab could follow your

steps and measure the same results. As you did with the Materials section, break

the Procedure section into steps for each part of the lab, the specific heat and the

density.

• Data

Be sure you clearly label all measurements and their appropriate units. For

example, in the Density portion of the experiment, include V(initial) and V(final)

measurements. Table format, as you have seen in our specific heat problems, is a

convenient way to display multiple masses, temperatures, and specific heat

capacities.

Calculations

Label all calculations. Actually show how you solved for a particular variable. It

will not be sufficient to put all the numbers in an equation and then write the

answer. Show how you solved the equation. Include the correct number of

significant figures.

Results Add the following data table to the results of your calculations.

Metal Specific Heat (J/g˚C) Density (g/mL)

Aluminum 0.90 2.7

Iron 0.45 7.9

Lead 0.13 11

Nickel 0.44 8.9

Zinc 0.39 7.1

Your Washer

Questions

Answer the following questions. Please retype each question.

1. What is a calorimeter? Describe your calorimeter AND its function in heat

transfer.

2. In the specific heat capacity experiment, which will produce a large change in

temperature (large ∆T): using a large or small volume of water in the

calorimeter? EXPLAIN YOUR ANSWER.

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3. In the density experiment, which will produce a lower % error in the volume

of washers: using a large number of washers or a small number of washers?

EXPLAIN YOUR ANSWER.

Discussion

Write an analysis of your experimental results. Use complete sentences and

paragraphs. Do not use bullets! Include each of the following points in your

paragraphs.

• In one or two sentences, summarize the objective of the experiment AND the

general experimental approach (for both density and specific heat). • Compare your results from the specific heat capacity experiments to values of

the known metals. Calculate the % error from the closest match to your data,

showing the equation and values you used. • Compare your results from the density experiments to the values of the known

metals. Calculate the % error from the closest match to your data, showing

the equation and values you used.

% error = | closest value – your value | x 100 %

closest value

• Comment as to whether the results from the specific heat capacity

experiments lead you to the same results of the density experiments. Can you

identify the metal in the washers with a high degree of confidence? If there is

a discrepancy, comment as to what you think the identity of the washer

actually is.

• List 2 possible sources of potentially significant error in your specific heat

capacity experiment AND 1 source from the density experiment (3 errors

total).

DO NOT SAY “HUMAN ERROR.”

DO NOT SAY “We may have measured ________ incorrectly or

inaccurately.”

These errors must be detailed and specific and should relate to the steps

in your procedure that may affect your results.

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Exploring pH Phet Simulation

Name: Period: 1) Open Phet pH solution simulation Macro program.

2) Action icons:

➢ Drop menu: Select substance to be tested

➢ pH probe: Make sure cross-hair of green ring is completely submerged in

solution being tested.

➢ Drain: Slide to the left (i.e. pull out) the blue knob which is in the bottom

left of the beaker to drain solution out of the beaker

➢ Dropper: Add additional solution by clicking on red button on floating

dropper positioned above beaker.

➢ Water spout: Slide to the left (i.e. pull out) the blue knob which above the

top right corner beaker to add more water (dilute) your solution.

Carry out the steps described below for four different solutions: “Battery Acid” pH =

1.00, “Vomit” pH = 2.00, “Coffee” pH = 5.00. Record your data in a table and then

answer the questions.

1) Use the dropdown menu to select the solution to be tested.

2) Click and drag the green pH probe and place it at the bottom of the beaker.

Record the value of the pH. It should be the same of the value given on the drop

down menu.

3) The volume of your solution is currently ½ L. Use the dropper icon to add

solution to the 1.00 L mark. Record the value of the pH.

4) Press the reset button. The volume of your solution should now reset to ½ L. Use

the drain icon on the bottom left side of the beaker to drain the solution down to

exactly 0.10 L by sliding the blue knob to the left. Make sure that the cross-hair

center of pH probe is still submerged and record the pH reading.

5) Using the blue knob on the water spout icon, add enough water to change the

volume to exactly 1.00 L. (If you go past the 1.00 L mark you will need to start

over). Record the new pH value.

Substance Tested pH of ½ L pH of 1.00 L pH of 0.10 L pH after

diluting 0.10 L

to 1.00 L with

water Battery acid Vomit Coffee

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1) For each substance, compare the pH values of ½ L, 1.00 L and 0.10 L. Use the

concept of concentration (ratio of moles to volume) to explain the results. If

represent 0.10 moles of H+ ions, how many should be drawn in the 1.0 L

volume? Draw the correct number of in the box.

0.10 L 0.50 L 1.0 L

Explanation of why pH values did NOT change:

2) Compare the pH of 0.10 L of solution before and after diluting with water to 1.00

L. Use the concepts of concentration and dilution to explain why the pH changes.

How many should be drawn in the 1.0 L volume? (Hint: Has the moles

changed? Has the volume changed? Has the concentration changed? Why?)

Moles changed? Volume changed? Concentration changed?

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3) Calculating concentration changes: pH is a measure of acidity that is directly linked to

the concentration of H+ ion, [H+], present.

The molar concentrations of H+ ion present in each solution is given below. Use the

dilution equation and the data given to calculate the new [H+] concentration for each

solution.

Dilution Equation: McVc = MdVd → Solving for Md

Md = McVc/Vd

Battery Acid: Mc = 0.10 or 1.0 x 10-1 M Vc= 0.10 L Vd = 1.0 L Md

= ?

Md =

Vomit: Mc = 0.010 or 1.0 x 10-2 M Vc= 0.10 L Vd = 1.0 L Md = ?

Md =

Coffee: Mc = 1.0 x 10-5 M Vc= 0.10 L Vd = 1.0 L Md = ?

Md =

Enter the final Md values you calculated for each in the table below. Substance Tested Initial pH Initial [H+] Final pH after

diluting 0.10 L

to 1.00 L with

water

Final [H+]

After dilution

Battery acid 1.0 1 x 10-1 2.0 Vomit 2.0 1 x 10-2 3.0 Coffee 5.0 1 x 10-5 6.0

What is the relationship between pH and [H+] ? Fill in the blank: Each pH unit

difference represents a factor of ____ difference in [H+].

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Exploring the Relationship between [H+], [OH-] and pH, 1) Open pH scale phet, custom program. Check the box to show the ratio of H3O

+

(H+) and (OH-)

2) Record data on [H+],[OH-] for the following pH values: 1, 4,7, 10, 14. Classify

each pH as acidic, basic or neutral and calculate the product of [H+] x [OH-]

pH value Acidic, Basic

or Neutral

[H+] [OH-1] [H+][OH-1]

1

4

7

10

14

Questions:

1) Label the diagrams below as representing acidic, basic or neutral solutions and

then fill in the appropriate missing information.

Indicate the relationship between H+ and OH- using <, > or = symbols

H+ OH- H+ OH- H+ OH-

Indicate the pH relative to 7 using <, > or = symbols.

pH 7 pH 7 pH 7

2) Patterns in the numbers: The product of H+ and OH- is always equal to 10 to

what power? [H+][OH-1] =

3) Use your equation from the previous question to calculate the [H+] if the [OH-1]

= 1 x 10-4 M

H+ H+ H+ H+ H+

OH - OH- OH- OH- OH-

OH-

H+ H+ H+ H+ H+

H+ H+ H+ OH- OH-

H+ H+ H+ OH- OH-

OH- OH- OH- OH- OH-

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Importance of pH to the chemistry of water and to aquatic life illustrated by

video clip from 2015 PBS NOVA program Lethal Seas. Name:

Period:

Answer the following questions: 1) What happened to the production of pacific oyster larvae off the coast of the pacific

northwest (US coast) starting in 2007?

2) Describe the shell structure of a normal baby oyster observed under a microscope

compared to the shell structures observed in 2007 from baby oysters that did not develop

properly.

3) What is pH?

4) Complete the following sentences:

The lower the pH, the ____________ the acidity. Tests revealed that the acidity of the

bay water was 6 times __________ than normal sea water.

5) In order for the baby oysters to develop properly it was necessary to add soda ash

(sodium carbonate), a chemical base, which neutralized the excess acid and

____________ the acidity to back normal levels. This procedure allowed the baby oysters

to develop sufficiently to be reintroduced into the bay, however it is very expensive.

6) Water sampling around the world over the last forty years has revealed that the average

ocean acidity has increased approximately five percent each decade. The increase acidity

has been linked with the increase in atmospheric levels of _______ gas.

7) Burning gasoline (primarily C8H18) to power car engines is an example of combustion of

a hydrocarbon. What are the products of the following combustion reaction?

C8H18 + O2 → 8) Carbon dioxide, CO2, is a greenhouse gas. Without greenhouse gases to trap heat in our

atmosphere our planet would be too cold to support life as we know it. However in recent

decades average global temperatures have been rising due release of too much CO2

released into the atmosphere relative to level of CO2 absorbed by plants and water.. (Fill

in the blank) Approximately ______ of the CO2 released into our atmosphere dissolves

in the ____________.

9) Write the chemical equation for the reaction of carbon dioxide and water to form

carbonic acid (H2CO3).

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10) Carbonic acid, H2CO3 dissociates in water to release H+ ions, according to the following

equation:

H+ ion concentration, [H+], determines acidity. The higher the [H+], the __________ the

acidity and the ___________ the pH.

11) The shells of sea creatures such as oysters are primarily made of calcium carbonate.

Calcium carbonate forms by the reaction of calcium ion with carbonate ion. Write the

equation for the reaction of calcium ion and carbonate ion to form calcium carbonate.

12) Ocean water generally contains a high concentration of _____________ ion, but

relatively a much lower concentration of carbonate ion. Increased acidity can further

_____________ the availability of carbonate ion to form shells due to the following reaction.

H+ + CO3-2 → HCO3

-1

If the concentration of carbonate ion is too ______, then sea creatures cannot make

calcium carbonate shells.

13) In biology class last year you discussed the concept of an ecosystem and predator/prey

relationships. Why are biologists concerned that increasing acidity in the ocean might disrupt

ocean ecosystems?

Investigating pH Lab Introduction

pH is a measure of the acidity of an aqueous solution. The acidity of a solution

can have a major impact on the types of reactions that will take place in aqueous

solutions. Since approximately ¾ of both the surface of the planet we live on and our

own bodies are water, understanding pH is critical for understanding both the ecosystems

89

of our planet and our own bodies. In this lab you will carry out a series of computer

(virtual experiments) simulations (details on separate sheet), and wet chemistry

investigations to understand more what we mean by pH and how it is important.

Safety:

Wear safety glasses when handling solutions.

If you accidently touch acid or base solutions, immediately rinse affected area with

water.

Part 1: Importance of pH on the rates of chemical reactions Experiment #1 (as demo): Reaction rate of metals

Materials:

• 3 small pieces of mossy zinc

• Test tube rack

• Test tubes containing 6 M HCl, water, and 6 M NaOH. 3 drops of universal indicator

have been added to each solution.

Color key: Acid = red/pink Neutral = yellow/green Base = dark blue

Procedure: Add a small piece of mossy zinc to each test tube; Record observations.

1) Reaction of Zinc with water, acidic, and basic solutions.

Solutions Initial Color

Observations

Classify color as

Acid, Base,

Neutral

Observations after

Zn added

6 M HCl

Water

6 M NaOH

Result of burning splint test on gas produced:

Experiment #2: Reaction of calcium carbonate with Acid, Water, Base Materials:

• Test tube rack

• spatula

• 3 scoops of calcium carbonate

• 3 test tubes containing: 6 M HCl, water, and 6 M NaOH. 3 drops of universal indicator

solution have been added to each test tube.

Procedure: Add scoop of calcium carbonate; Record observations.

Reaction of calcium carbonate with water, acidic, and basic solutions.

Solutions Initial Color

Observations

Classify color as

Acid, Base,

Neutral

Observations after

CaCO3 added

6 M HCl

90

Water

6 M NaOH

Result of burning splint test on gas produced:

Question: In the Nova Video “Lethal Seas”, explain damage to marine shells in

terms of the effects of acid on the calcium carbonate shells. What observations did

you make in this experiment to support this contention?

Experiment #3: Reaction of water, acidic and neutral solutions with solution

of silver nitrate. Materials:

Plastic sheet grid

• Dropper containing 0.10 M AgNO3

• 3 micropipet droppers: 0.10 M HNO3, water, and 0.10 M NaOH

Procedure:

• Add 3 drops of each solution to the AgNO3 solution to each grid.

• In the appropriate grid add 3 drops of acid, base and water to the AgNO3 solution in each

test tube. Record your observations.

• To the AgNO3 + Base solution ONLY: add 15 drops of acid solution and record new

observations.

Solutions Observations

Acid (3 drops)

Water (3 drops)

Base (3 drops)

Base soln + 15 drops acid

Questions:

Many of Colorado’s mountain towns (e.g. Silverton, Leadville) developed around mining

of heavy metals. Heavy metals dissolved in ground water can render water toxic and

unsafe to drink.

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1) Based on the result of your experiments with silver ion, are heavy metals more

soluble in acidic or basic conditions?

2) Under what conditions, acidic, basic or neutral does silver ion form an insoluble salt

which precipitates out of solution?

Experiment #4: Concept of pH scale: Preparing pH 2 and pH 3 solutions

from a pH 1 solution by serial dilution. Materials:

• 0.10 M HNO3 solution with beral pipette

• Universal Indicator Solution (pH 1-7)

• Distilled water containing universal indicator

• Three 10 mL Graduated cylinders

Procedure:

1) Line up the 3 different graduated cylinder. Add 4 drops of universal indicator

solution to each 10 mL graduated cylinder.

2) Fill the first graduated cylinder to the 10.0 mL line with 0.10 M HNO3

solution.

3) Remove 1.0 mL of 0.10 M HNO3 solution from the first graduated cylinder

and transfer it to the second graduated cylinder.

4) Add water to the 2nd graduated until the final volume reaches 10.0 mL.

5) Remove 1.0 mL of solution from the 2nd graduated cylinder and transfer it to

the 3rd graduated cylinder.

6) Add water to the 3rd .graduated cylinder until the final volume reaches 10.0

mL

92

Indicator Color Key: pH 1 = red, pH 2 = orange, pH 3 = yellow

Record the colors of graduated cylinders 1, 2, and 3 and enter the results in the data table

on the next page.

Graduated cylinder Color pH

1 (0.10 M HNO3)

2 (1:10 dil from cyl 1)

3 (1:10 dil from cyl 2)

Important Concept: HNO3 is a strong acid meaning that it completely dissociates into H+

and NO3- ions when dissolved in water. Thus the molar concentration of HNO3 = molar

concentration of H+ ion.

Example: 0.10 M HNO3 → 0.10 M H+ ion

Color A) In the first 0.10 M HNO3 solution, the [H+] = 0.10 M or 10-1 M. What is the pH

of this solution? (Recall pH = -log10 [H+])

B) First Dilution: What would be the concentration of a HNO3 solution of obtained

by diluting 1.0 mL (Vc) of 0.10 M HNO3stock solution (Mc) into a total

solution volume of 10.0 mL (Vd) ? Fill in your answer in the box above solution

2.

Recall: McVc = MdVd → Md = McVc / Vd

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C) Second Dilution: If we repeated our experiment by taking 1.0 mL (Vc) of the

concentration, you prepared in part B (Solution 2 is new stock for 2nd dilution-

your answer from part B is your new Mc) into a total solution volume of 10.

mL,(Vd) what would the final concentration of the new solution be? Fill in

final answer in the box above solution 3.

Md = McVc / Vd

D) Calculate the pH of solutions two and three from the [H+] values and write the

pH values underneath the pictures of the test tubes above. How do the pH values

you calculated compare to the pH value estimated from the color chart?

E) Why does this experiment work? Specifically what is the relationship between

[H+] and pH?

The [H+] concentration of a solution of pH 2 is (what fraction?) the

[H+] concentration of a solution of pH 1.


[H+] concentration of a solution of pH 2.


[H+] concentration of a solution of pH 1. (Two pH unit difference)

Conclusion: Write a paragraph summarizing what you have learned about pH and its

importance which addresses the following questions.

What is pH?

Support the statement “pH is important because it changes the types of chemical

reactions that can take place in water” with at least three examples from this lab:

Chemistry Lab Manual - Fairview High School

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