chem 100 lab manual harrisburg area community college...last updated 8/8/2013 . chem 100 lab manual....

Last updated 8/8/2013

CHEM 100 Lab Manual Harrisburg Area Community College

2013/2014

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Table of Contents

Equipment Illustrations ............................................................................................................................... 5

Introduction: Measurements and Recording Data .................................................................................. 7

Common Lab Equations .............................................................................................................................. 9

Measurements: Density of a Saline Solution .......................................................................................... 11 Pre-lab Questions .................................................................................................................................... 15 Data: ........................................................................................................................................................ 17

Separation of Mixtures .............................................................................................................................. 21 Pre-lab Questions .................................................................................................................................... 25 Data .......................................................................................................................................................... 27

Nomenclature.............................................................................................................................................. 31 Pre-lab Questions .................................................................................................................................... 33

Stoichiometry: Empirical Formula of a Hydrate .................................................................................... 41 Pre-lab Questions .................................................................................................................................... 43 Data .......................................................................................................................................................... 45

Balancing Chemical Equations ................................................................................................................. 49 Balancing Chemical Reactions ................................................................................................................ 51

Conductivity ............................................................................................................................................... 55 Pre-lab Questions .................................................................................................................................... 57 Data .......................................................................................................................................................... 59

Chemical Reactions: Classification and Prediction of Products .......................................................... 65 Pre-lab Questions .................................................................................................................................... 69 Data .......................................................................................................................................................... 71

Double Displacement Reactions ................................................................................................................ 73 Pre-lab Questions: ................................................................................................................................... 77 Data .......................................................................................................................................................... 81

Lewis Structures ......................................................................................................................................... 83 Pre-lab Questions .................................................................................................................................... 89 Lewis Structures ....................................................................................................................................... 91

Gas Laws ..................................................................................................................................................... 97 Pre-Lab Questions ................................................................................................................................. 101 Data ........................................................................................................................................................ 103

Acid-Base Titrations ................................................................................................................................ 107 Pre-lab Questions: ................................................................................................................................. 111 Data ........................................................................................................................................................ 113

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Spectroscopy: Determination of Concentration Using Beer’s Law .................................................... 115 Pre-lab Questions .................................................................................................................................. 117 Data Table .............................................................................................................................................. 119

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Equipment Illustrations

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Introduction: Measurements and Recording Data

Taking good measurements is one of the most important aspects of science. This is closely

followed in importance by being able to indicate to others how “good” the results are. Generally, the final results of a lab can be no better than the data or measurement that the results are based on, so the ability to get “good” results depends on how “good” the original data is.

How “good” are data and results?

The “goodness” (henceforth known as validity) of data relies primarily on two factors:

• the measuring device • how well the experimenter used the device

For example, if I wanted to determine the mass of a paperclip, I would use a digital balance

instead of holding the paperclip in my hand, because the balance is supposedly better. But be careful; the assumption that I should get a better value from the balance implies a whole bunch of other factors (such as calibrating the balance and using the balance correctly.

Two general terms used to discuss the “goodness” of data are accuracy and precision. Accuracy

describes how closely a measured value matches a known value, while precision describes how reproducible or how closely grouped a series of measurements are. “Good” data should be both accurate and precise.

How can someone look at a value and tell how precise it is?

Scientists record all significant digits of their measurements. This allows them to communicate

the precision of data and results without having to specifically identify all devices used and the exact procedures. The reader merely looks at the number to see how well it was measured.

An illustration of how this works follows: Two students are asked to count the amount of money

in a bag that’s in another room. Carol answers around $2, while Joan says $1.92. Who do you think measured the best? If you’re like me, you answered Joan. You didn’t see them count, but more decimal places usually implies better measuring.

The more significant digits—also referred to as significant figures—in a value, the more precisely is was measured. So, when you record a value in a lab, BE CAREFUL. You aren’t just writing down a number; you’re also telling someone how well you measured.

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CORRECTLY RECORDING EXPERIMENTAL DATA How many decimal places should be recorded for data?

• For digital devices, like a digital balance: o Record ALL numbers in the digital display, and record the error limit often printed on the

device. • For non-digital devices like a volumetric pipet or a ruler:

o First look on the device to see if there is a +/- error range on it. For example, if you are measuring volume with a 10 mL volumetric pipet, the pipet has “+/- 0.02 mL” marked on it. That means that the pipet is designed to measure to within +/- 0.02 mL of the actual value (assuming you measured properly), so any data you measure from this device should be recorded to the hundredths place (i.e. 10.00 mL, not 10.0 or 10).

o If there is no error range on the device you will have to determine the least count of the device. The least count is the smallest division on the measuring device. For example, if

the smallest division on a ruler is 1 cm, then the least count of that ruler is 1 cm or 0.01 m.

Once you identify the least count, the general rule is to record the value of the measurement to one decimal place smaller than that of the least count. For example in the ruler above, a measurement would be recorded to the 1/10th cm (12.3 cm).

How many significant figures should be recorded at the end of a calculation? Review your lecture notes for all of the guidelines. Here is a simple reminder:

• For multiplication and division, the answer will have the least number of total significant figures. • For addition and subtraction, the answer will have the least number of decimal places. • Only round the final answer.

Always follow the rules of significant figures when recording values and performing calculations.

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Common Lab Equations Average is the approximate middle value in a series of similar measurements.

Average = Sum of all data valuesNumber of data points

Range is the difference between the highest and lowest valued measurement. Range = | Highest value - Lowest value | Percent Error is the percent an experimental value differs from a known or “true” value.

% error = | Known value – Experimental value |

Known value × 100 Percent Difference is used when a known value is not given.

% difference = | Experimental value #1 – Experimental value #2 |Average of #1 and #2 × 100

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Measurements: Density of a Saline Solution Objectives: In this lab, students will:

• Determine the density of a saline solution by measuring its mass and volume • Compare the effectiveness of using two different devices for measuring volume: a graduated

cylinder and a volumetric pipet. Skills: Upon completion of this lab, students should have learned:

• To measure and record volume of liquids using graduated cylinders and volumetric pipets • To measure and record mass using a digital balance • To record numerical data and calculated values to the appropriate number of significant figures

Textbook References: (Tro, Introductory Chemistry, 4th Ed.) Read these sections PRIOR to lab.

• Measurements and significant figures Section 2.3 • Significant Figures in Calculations Section 2.4 • Converting from One Unit to Another Section 2.6 • Density Section 2.9 • Classifying Matter Section 3.4 • Physical and Chemical Properties Section 3.5

Lab Manual References: Read these sections PRIOR to lab.

• Introduction: Measurements and Recording Data • Common Lab Equations

Introduction:

Measuring data is what differentiates science from other courses. In this lab, you will use measurements to explore the concept of density. The density of an object is defined as the ratio of its mass to its volume.

MassDensityVolume

= or mdV

=

Elements or compounds can often be identified by determining their density. For example, aluminum and titanium look very similar, but the density of aluminum is 2.7 g/mL and that of titanium is 4.5 g/mL. If you measured the mass and volume of an unknown silvery metal and calculated the density to be around 4.5 g/mL, the metal is probably titanium.

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You will use a volumetric pipet to measure out the solution. A suction bulb is used to withdraw air from the pipet while drawing up the liquid to be measured. Always use the suction bulb and always hold the pipet with your dominant hand and hold the pipet bulb in your other hand . Never pipet by mouth.

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Procedure: Determine the density of a saline solution

Part A. Obtain 150-200 mL of saline solution in a 250 mL beaker. Record the known density.

Part B. Graduated Cylinder

1. Measure and record the mass of a dry 150 mL beaker on the digital balance. Record the mass to the correct number of decimal places.

2. Pour a volume of saline solution between 15 mL and 45 mL into a clean and dry 50 mL graduated cylinder. Record the volume to the correct number of decimal places in your data table.

3. Pour the saline from the graduated cylinder into the empty beaker. Measure and record the mass of the beaker with the saline in it. Don’t empty the beaker.

4. Repeat step B2 and B3, two more times, using a different volume between 15 mL and 45 mL each time.

Part C. Volumetric Pipet - Practice using the volumetric pipet with deionized water until you are

comfortable with it.

1. Measure and record the mass of a dry 50 mL beaker on the digital balance. Record the mass to the correct number of decimal places.

2. Suction 10 mL of saline into a volumetric pipet (review instructions on p. 12). Record the volume to the correct number of decimal places in your data table.

3. Release the saline from the volumetric pipet into the empty beaker. Measure and record the mass of the beaker with the saline in it. Don’t empty the beaker.

4. Repeat step B2 and B3 two more times. Use the correct number of decimal places for the volumetric pipet.

5. Clean up.

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Name: _________________________________________ Date due: ___________________________

Measurements: Density of a Saline Solution

Pre-lab Questions: Read the relevant textbook sections and the entire lab before answering pre-lab questions. 1. If a measured quantity is written correctly, which digits are certain? Which are uncertain?

2. A new penny has a mass of 2.49 g and a volume of 0.349 cm3. Is the penny pure copper? (dcopper = 8.96 g/cm3) Show your work!

3. Define accuracy and precision. 4. To the correct number of significant figures, what is the volume, in milliliters, of the liquid in the graduated cylinder?

5. Refer to the procedure and explain how you can make small adjustments to the volume of liquid in a graduated cylinder.

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6. The mass of a piece of copper (known mass = 4.750 g) is measured three times by two different students. The students’ results follow: Gloria Max 4.68 g 4.69 g 4.86 g 4.71 g 4.75 g 4.66 g

a) Calculate the average mass of the copper determined by each student. Show your work

Gloria___________ Max____________

b) Calculate the range of each student’s measurements. Show your work

Gloria___________ Max____________

c) Calculate the percent error for each student’s mass measurement. (Use the calculated average mass as

the experimental value.) Show your work

Gloria___________ Max____________

d) Which student’s measurements were more precise? ___________

e) Which student’s measurements were more accurate? ___________

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Name: _________________________________________ Date lab performed: __________________

Partner(s) name:_________________________________ Date due: ___________________________

Measurements: Density of a Saline Solution

Data: Record all data with correct units and significant figures.

A. Known density of saline solution _________

B. Graduated Cylinder

1. Volume of saline added each trial _________

2. Mass of empty beaker _________

3. Mass of beaker with first addition of saline _________

4. Mass of beaker with second addition of saline _________

5. Mass of beaker with third addition of saline _________

C. Volumetric Pipet

1. Volume of saline added each trial _________

2. Mass of empty beaker _________

3. Mass of beaker with first addition of saline _________

4. Mass of beaker with second addition of saline _________

5. Mass of beaker with third addition of saline _________

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Calculations: • Show neatly labeled and organized work for each of the following calculations below. • Include correct units and significant figures in all calculations.

Using the Graduated Cylinder Data from Part B, calculate the:

1. Mass (g) and density (g/mL) of the first addition of saline

2. Mass and density of the second addition of saline

3. Mass and density of the third addition of saline

4. Average experimental density of the three samples of saline

5. Range of the densities of the three samples of saline

6. Percent error between the known density (from Data A) and the average experimental density (from Calc 4).

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Using the Volumetric Pipet Data from Part C, calculate the:

7. Mass (g) and density (g/mL) of the first addition of saline

8. Mass and density of the second addition of saline

9. Mass and density of the third addition of saline

10. Average experimental density of the three samples of saline

11. Range of the densities of the three samples of saline

12. Percent error between the known density (from Data A) and the average experimental density (from Calc 10).

Results Table: Part B Part C Cylinder Pipet

1. Density of first volume of saline _________ _________

2. Density of second volume of saline _________ _________

3. Density of third volume of saline _________ _________

4. Average density _________ _________

5. Density range _________ _________

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6. % error _________ _________

Conclusion Questions:

1. Which term - range or % error - best describes the precision of data? __________________

2. Which term - range or % error - best describes the accuracy of data? _________________

3. Were your results more precise using the graduated cylinder or volumetric pipet? Explain using your values.

4. Were your results more accurate using the graduated cylinder or volumetric pipet? Explain using your values.

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Separation of Mixtures

Objectives: In this lab, students will: • Separate the components of a mixture using filtration and evaporation. • Verify the Conservation of Mass Law.

Skills: Upon completion of this lab, students will have learned:

• To use a digital balance to measure mass • To prepare solutions • To separate a mixture using vacuum filtration • To separate a mixture using evaporation • To identify common lab chemicals as mixtures or pure substances • To identify chemical and physical changes


• Classifying Matter Section 3.4 • Physical and Chemical Properties Section 3.5 • How Matter Changes Section 3.6 • Conservation of Mass Law Section 3.7

Introduction:

Mixtures are classified as either homogeneous or heterogeneous. When a mixture is homogeneous, techniques such as evaporation and crystallization are used to separate the mixture into its components. When a mixture is heterogeneous, techniques such as filtration and decantation are used to separate the mixture into its components. This experiment will study the reaction between calcium chloride and sodium carbonate, which is shown below. The product mixture will be separated into pure sodium chloride and pure calcium carbonate via filtration and evaporation. CaCl2 (aq) + Na2CO3 (aq) 2 NaCl (aq) + CaCO3 (s)

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Procedure: Part A. Preparation of solutions

1. Measure out 0.9–1.1 g of solid anhydrous calcium chloride, CaCl2, using a digital balance and weighing paper. Examine the solid CaCl2 and record your observations.

2. Transfer the CaCl2 into a large test tube, labeled “A”.

3. Use a graduated cylinder to measure approximately 10 mL of deionized water, and add the water to the solid. Stir the solution until all of the solid CaCl2 dissolves. Record your observations.

4. Measure out 0.9−1.1 g of solid anhydrous sodium carbonate, Na2CO3, using a digital balance and weighing paper. Examine the solid Na2CO3 and record your observations.

5. Transfer the Na2CO3 into a large test tube, labeled “B”.

6. Use a graduated cylinder to measure approximately 10 mL of deionized water, and add the water to the solid. Stir the solution until all of the solid Na2CO3 dissolves. Record your observations.

Part B. Mixing solutions

1. Mix the contents of test tube A and B together in a clean 50 mL beaker. Rinse each test tube with a small amount of deionized water (~ 1 mL) and add the rinse water to the beaker. Swirl the contents of the beaker, and record your observations.

Part C. Vacuum filtration (Buchner filtration)

1. Assemble the vacuum filtration apparatus as indicated by your instructor.

2. Write your initials on a filter paper circle in pencil.

3. Record the mass of the filter paper and a watch glass together.

4. Place the filter paper in the Buchner funnel (initials down) and wet the paper with a small amount of deionized water (~1 mL).

5. Turn on the vacuum and pour the beaker’s contents into the Buchner funnel. Record your observations.

6. Use a small amount of deionized water (no more than 5 mL) to rinse the beaker. With the rubber end of the stirring rod, transfer the rinse water and the remaining precipitate (solid) from the beaker onto the filter paper.

7. Continue to pull a vacuum on the solid for 5 minutes.

8. Remove the filter paper, and place it on the watch glass (measured in step 8). Place the watch glass in the oven.

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9. Record the mass of the watch glass with dried material before the end of the lab period.

Part D. Evaporation of water

1. Record the mass of a clean, dry evaporating dish.

2. Pour all of the liquid from the receiving flask into the evaporating dish. Rinse the flask with a small amount of deionized water (less than 5 mL) and add the rinse to the dish.

3. Place the evaporating dish on the clay triangle on a ring stand (as indicated by your instructor) and evaporate the liquid using a Bunsen burner. Record your observations as the liquid evaporates. Reduce the flame near the end of the heating (when most of the water has evaporated) to prevent the solid from “popping” out of the dish. Continue to heat for approximately 5 minutes after you see no more steam.

4. Turn off the gas and allow the dish to cool.

5. Record the mass of the dish and remaining solid.

6. Dispose of solids and filter papers in the appropriate container.

7. Clean up.

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Name: __________________________________________Date due: ___________________________


Pre-lab Questions: Read the relevant textbook sections and the entire lab before answering pre-lab questions.

1. Define the following terms in your own words: Pure substance—

Mixture—

Heterogeneous mixture—

Homogeneous mixture—

Physical change--

Chemical change—

2. State the Law of the Conservation of Mass.

3. Using the conservation of mass law, calculate how much NaCl, table salt, forms when 27.4 g Na reacts with 42.3 g Cl2.

4. What kind of change occurs when NaCl is formed from Na and Cl2? Hint: Examine the images of elemental sodium and elemental chlorine shown in section 5.1 in your textbook.

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5. If NaCl was placed in a beaker of water and dissolved, what kind of change would occur?

6. Write the chemical equation for the reaction you will perform in this lab.

7. Refer to the procedure and indicate the technique you will use to isolate NaCl from the mixture.

8. Refer to the procedure and indicate the technique you will use to isolate CaCO3 from the mixture.

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Name: __________________________________________Date lab performed: __________________

Partner(s) name: _________________________________Date : ______________________________



A. Mass of weighing paper and solid CaCl2 ___________

Mass of weighing paper ___________

Mass of solid CaCl2 ___________

Mass of weighing paper and solid Na2CO3 ___________

Mass of weighing paper ___________

Mass of solid Na2CO3 ___________

B. Observations:

CaCl2 solid: CaCl2 solution: Na2CO3 solid: Na2CO3 solution: After mixing solutions A and B: Vacuum filtration (see procedure, part C, step 5):

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C. Mass of filter paper and watch glass ___________

Mass of dried filter paper, watch glass, and solid ___________

Observations:

D. Mass of evaporating dish ___________

Mass of evaporating dish and solid ___________ Observations:

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Calculations: • Show your calculations. • Include correct units and significant figures in all calculations.

Calculate:

i. The total mass of dissolved solids (in Part A).

ii. The mass of solid from filtration.

iii. The mass of solid from evaporation.

iv. The total mass of recovered solids.

v. The % error between the total mass of dissolved solids (reactants) and the total mass of recovered solids (products). Use the total mass of the dissolved solids as the “known” value in the calculation,

Results Table: total mass of dissolved solids _________ total mass of recovered solids _________ % error _________

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1. How well do your results support the Conservation of Mass law (very well, OK, not very well)? Support your answer using your result values. Evaluate your technique and discuss where procedural errors may have occurred.

2. Identify the following as an element, a compound, a homogenous mixture, or a heterogeneous mixture.

a. Solid sodium carbonate _________________________

b. Sodium carbonate solution _________________________

c. The result of Part B _________________________

d. The liquid from Part C _________________________

3. Identify the following as a physical or chemical change.

a. The dissolving of CaCl2 in water _________________________

b. The formation of the white precipitate in Part B ______________________

c. The separation of solid from solution in Part C _______________________

d. The evaporation of the liquid in Part D _________________________

4. The solid collected on the filter paper is either NaCl or CaCO3. Explain which one it is.

5. Identify the liquid that is evaporated in Part D. Explain your answer.

6. Which compound remains after evaporation? Explain your answer.

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Nomenclature

Objectives: In this lab, students will learn: • To identify a compound as ionic, molecular, or an acid from either its name or its formula • To identify cations, anions, and charges • To identify compounds as binary or polyatomic • To correctly name compounds when given the formula • To write the correct chemical formula when given the name

Textbook References: (Tro, Introductory Chemistry, 4th Ed.) Read these sections PRIOR to lab,

and bring your textbook with you to lab. • Chemical Formulas Section 5.3 • A Molecular View of Elements Section 5.4 • Writing Formulas for Ionic Compounds Section 5.5 • Nomenclature: Naming Compounds Section 5.6 • Naming Ionic Compounds Section 5.7 • Naming Molecular Compounds Section 5.8 • Naming Acids Section 5.9 • Nomenclature Summary Section 5.10

Introduction:

Names and Formulas are used to identify chemicals. Each different chemical has a unique formula

and a unique IUPAC name. The IUPAC naming system was developed so that everyone names the

compound the same way to avoid confusion. The IUPAC names will be used in class, but many

compounds also have a common name that you already know.

Most of the chemical compounds used in CHEM 100 can be roughly separated into three main

categories – acids, ionic compounds, and molecular compounds.

Acids are easy to identify by their name or formula. The name of an acid always includes the term

“acid” and the formula of an acid usually begins with the chemical symbol for hydrogen, H. There

are two main types of acids – binary acids and oxyacids. Binary compounds contain only two

different elements. The first example below is a binary acid and the other two are oxyacids.

IUPAC name Formula Common name or use

hydrochloric acid HCl muriatic acid; used in concrete work and welding acetic acid HC2H3O2 vinegar sulfuric acid H2SO4 battery acid

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Ionic compounds often include a metal in the formula and name. There are two types of ionic

compounds: binary ionic and polyatomic ionic. The last two examples have polyatomic ions—ions

with more than two different elements. The names of ionic compounds with polyatomic ions in

them will often end in “ate” or “ite”. A table of the most common polyatomic ions is found in Tro

(Table 5.6, p. 141).

Notes:

• The last example has no metal in it, but is still an ionic compound because it contains ions.

• The second example below includes a Roman numeral after the name of the cation, iron. A

Roman numeral is used to identify the ionic charge for metals that can have different charges.

(These metals are referred to as Type II metals.) Iron(III) is the name of Fe3+.

IUPAC name Formula Common name or use sodium chloride NaCl table salt iron(III) oxide Fe2O3 rust potassium nitrite KNO2 used to cure meats ammonium nitrate NH4NO3 used in instant cold packs

Molecular compounds include only nonmetals in the formula, and they aren’t acids. You will only

be responsible for naming binary molecular compounds like the first three examples below.

Molecular compounds that contain more than two different elements, like the last example, are

organic chemicals and have an entirely different naming scheme which is beyond the scope of the

CHEM 100 course.

IUPAC name Formula Common name or use dinitrogen monoxide N2O nitrous oxide or “laughing gas” carbon dioxide CO2 one of the green house gases ammonia NH3 this compound is so common that the

common name has been adopted as the IUPAC name, just like water

ethyl alcohol C2H6O in beer, wine, and distilled spirit

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Name: __________________________________________Date due: ___________________________

Nomenclature

Pre-lab Questions: Read the relevant textbook sections and the entire lab before answering pre-lab questions.

1. Define the following terms in your own words: a. element

b. compound

c. metal

d. non-metal

e. main group metal

f. transition metal

g. acid

h. ion

i. cation

j. anion

k. polyatomic ion 2. Explain how to name ionic compounds containing type II metals. Give two examples.

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Partner(s) name: _________________________________Date due: ___________________________

Nomenclature

A. Look at each formula in the list below and identify each • As Acid (A), Ionic (I), or Molecular (M) in the first column • As binary (B) or polyatomic (P) in the second column • As Type II (T) if it contains a type II metal

Formula Acid, Ionic, Binary or Type II or Molecular Polyatomic ion metal

KI ________ ________ ________

HCl ________ ________ ________

BaSO4 ________ ________ ________

SO2 ________ ________ ________

Fe(NO3)2 ________ ________ ________

HgI ________ ________ ________

H2SO4 ________ ________ ________

B. Write names for the ions found in each ionic compound and indicate the charge on each individual ion (not the total charge). Include roman numerals with the cation if necessary.

Name of Cation Charge Name of Anion Charge

ScCl3

PbS2

K2SO4

Al(C2H3O2)3

Fe2O3

Mn(HCO3)2

CsMnO4

(NH4)2SO3

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C. Naming compounds when given the formula

1. Name each ionic compound. (Remember to look for type II metals.)

Hg2O CaF2

Al2(SO4)3 Au2CO3

2. Name each molecular compound.

CCl4 SF6

N2O5 S2O3

3. Name each aqueous acid.

HNO2(aq) HC2H3O2(aq)

H2SO4(aq) HCl (aq)

4. Name each compound.

CrCl3 Li3N

K2O SrS

CO (NH4)2SO3

Ni(NO2)3 HI (aq)

HF (aq) H2CO3 (aq)

H3PO4 (aq) HNO3(aq)

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D. Write chemical formulas when given the name.

1. Write the formula for each ionic compound. (Remember to look for type II metals.)

gallium oxide copper(II) bromide

aluminum nitrite silver phosphate

2. Write the formula for each molecular compound.

sulfur dioxide phosphorus trichloride

iodine monobromide nitrogen monoxide

3. Write the formula for each aqueous acid.

sulfuric acid hydrobromic acid

perchloric acid sulfurous acid

4. Write the chemical formula.

water lithium sulfate

ammonia ammonium hydrogen sulfate

silver carbonate sodium bicarbonate

ammonium nitrite oxygen difluoride

iron(II) nitrate potassium phosphate

calcium nitride barium sulfide

sodium phosphide lead(IV) oxide

zinc acetate strontium hydroxide

cobalt(III) oxide aluminum cyanide

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Stoichiometry: Empirical Formula of a Hydrate

Objectives: In this lab, students will: • Determine the empirical formula of a hydrate

Techniques: Upon completion of this lab, students will have learned:

• To determine mass percent composition • To determine empirical formula


• Converting between grams and moles Section 6.4 • Mass Percent Composition Section 6.6 • Calculating Empirical Formulas Section 6.8 • Classifying Chemical Reactions Section 7.10

Introduction:

Gravimetric analysis involves comparing the mass of a compound before and after a chemical reaction. Gravimetric analysis is often used in experiments to determine the stoichiometry of a chemical reaction.

Decomposition reactions involve one chemical breaking down into two or more simpler chemicals. Some decomposition reactions occur spontaneously at room temperature while others require higher temperatures.

Hydrates are a common class of compounds that have water molecules loosely bound within the crystal structure of a solid. The number of bound water molecules depends on the hydrate being examined. The loosely bound water molecules can be driven off through heating. The waterless form of the compound is called the anhydrous form of the compound.

In this experiment, you will determine the stoichiometric ratio of water molecules in copper(II) sulfate hydrate. When copper(II) sulfate hydrate is heated, it gives off water vapor according to the decomposition reaction below.

CuSO4

• nH2O (s) → CuSO4 (s) + nH2O (g) blue white

The value for n, which represents the mole ratio of water molecules to CuSO4 in copper (II) sulfate hydrate will be calculated from the experimental data.

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Procedure:

Safety Precautions:

• The Bunsen burner flame is extremely hot and is nearly invisible. The metal stand, clamp

and the clay triangle look the same whether hot or cold. Use extreme caution.

Procedure: 1. Assemble the ring stand, ring clamp, clay triangle and Bunsen burner as demonstrated by

your instructor.

2. Obtain a clean crucible and record its mass.

3. Add approximately 5 g of copper(II) sulfate hydrate into the crucible. Record the mass of the

crucible and copper (II) sulfate hydrate. Describe the appearance of copper(II) sulfate

hydrate.

4. Heat the crucible GENTLY over the Bunsen burner for 15 to 20 minutes, or until the sample

has completely turned white. Avoid strong heat, because it may trigger the decomposition of

anhydrous copper sulfate: CuSO4 (s, white) CuO(s, brown) + SO3(g)

5. Remove the crucible from the heat, and allow it to cool completely.

6. Record the mass of the crucible and anhydrous copper(II) sulfate.

7. Repeat heating (5 minutes) and cooling cycles until the mass of the crucible and anhydrous

copper(II) sulfate changes by less than 0.05 g between two consecutive heatings. Record the

appearance of the anhydrous copper(II) sulfate.

8. Place the anhydrous copper sulfate in the waste container, and remove any remaining solid

from the crucible by washing with water.

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Name: __________________________________________Date due: ___________________________ Stoichiometry: Empirical Formula of Copper(II) Sulfate Hydrate Pre-lab Questions: Read the relevant textbook sections and the entire lab before answering pre-lab questions.

1. Define the terms hydrate and anhydrous in an inorganic chemistry context, and provide an example of each.

2. It is very difficult to predict the number of water molecules in a hydrate compound. Go online to a legitimate source (such as a chemical supplier or a college) and get the actual formula for copper(II) sulfate • n hydrate. This will tell you n, which is the ratio of water molecules to copper(II) sulfate units. Reference your source.

3. A 2.241 g sample of nickel combines with oxygen to produce 2.852 g of a metal oxide.

a. Calculate the number of moles of nickel in the metal oxide.

b. Calculate the number of grams of oxygen in the metal oxide.

c. Calculate the number of moles of oxygen in the metal oxide.

d. What is the empirical formula of the metal oxide?

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Stoichiometry: Empirical Formula of Copper(II) Sulfate • n Hydrate


1. Mass of crucible _________

2. Mass of crucible and copper(II) sulfate • n hydrate _________

3. Mass of crucible and anhydrous copper(II) sulfate (1st heating) _________

4. Mass of crucible and anhydrous copper(II) sulfate (2nd heating) _________

5. Mass of crucible and anhydrous copper(II) sulfate(3rd heating) _________

(if needed)

Observations:

Initial appearance of the copper(II) sulfate • n hydrate:

Observed changes during heating:

Final appearance of the anhydrous copper(II) sulfate:

46

Calculations: • Show calculations and express answers with correct units and significant figures. • Record the results in the results table.

1. Mass of copper(II) sulfate • n hydrate ______________g

2. Mass of anhydrous copper(II) sulfate _______________g

3. Moles of anhydrous copper(II) sulfate _______________mol

Molar mass of copper(II) sulfate _______________g/mol

4. Mass of water in the copper(II) sulfate • n hydrate _______________g

5. Moles of water in the copper(II) sulfate • n hydrate _______________mol

Molar mass of water _______________g/mol

6. Calculate the ratio of moles of water over moles of anhydrous copper(II) sulfate. This number equals n. n = _________________

7. Write the experimental formula for copper(II) sulfate • n hydrate in the form CuSO4 ∙ nH2O. Round n to nearest whole number.

___________________

8. Calculate the number of moles of copper(II) sulfate pentahydrate used in this experiment.

___________________mol

47

9. Compile your data in the table below:

Compound Formula Molar Mass (g/mol)

Experimental mass (g) Moles

Copper(II) sulfate pentahydrate

Copper(II) sulfate (anhydrous)

Water


1. How does your experimental empirical formula of copper(II) sulfate • n hydrate compare to your predicted formula from the pre-lab? Analyze your procedure and identify where your experimental error could have occurred.

2. Using your experimental data, calculate the percent of water in your copper(II) sulfate • n hydrate, by mass. Show your work.

48

3. If you used twice as much copper(II) sulfate • n hydrate at the beginning of the lab, would you expect the empirical formula of the product to be the same or different? Explain.

4. If not all of the water was removed from the copper(II) sulfate pentahydrate after heating, would you expect the empirical formula of the product to be the same or different? Explain.

49

Balancing Chemical Equations

Objectives: In this lab, students will: • Balance equations • Classify chemical equations


• The Chemical Equation Section 7.3 • How to Write Balanced Chemical Equations Section 7.4

Balancing Equations: To balance equations, it is necessary to only CHANGE WHOLE NUMBER COEFFICIENTS IN

FRONT OF formulas for elements and compounds! In the reaction, ALL FORMULAS FOR

REACTANTS AND PRODUCTS ARE INCLUDED, and ALL FORMULAS FOR REACTANTS AND

PRODUCTS ARE CORRECT. Subscripts in the chemical formula are NEVER changed.

Easy Rules for Balancing Equations:

1. Balance everything but hydrogen and oxygen.

2. Polyatomic ions can be treated as one unit if they appear in both a reactant and a product.

3. Balance hydrogen.

4. Balance oxygen.

5. Redo rules 1, 2, and 3 again, if necessary.

51



Balancing Chemical Reactions A. BALANCING EQUATIONS: Balance each of the chemical equations. Coefficients of 1

shouldn’t be written. 1. ___C6H14 + ___O2 → ___CO2 + ___H2O 2. ___C9H20 + ___ O2 → ___CO2 + ___H2O 3. ___Zn + ___HCl → ___ZnCl2 + ___H2 4. ___P4 + ___Cl2 → ___PCl3 5. ___NaHCO3 → ___Na2CO3 + ___CO2 + ___H2O 6. ___HNO3 → ___NO2 + ___H2O + ___O2 7. ___Fe + ___O2 → ___Fe2O3 8. ___CaC2 + ___H2O → ___Ca(OH)2 + ___C2H2 9. ___Mg3N2 + ___H2O → ___NH3 + ___Mg(OH)2 10. ___CaCO3 + ___HCl → ___CaCl2 + ___CO2 + ___H2O 11. ___Zn + ___H3PO4 → ___Zn3(PO4)2 + ___H2 12. ___AgNO3 + ___CaCl2 → ___AgCl + ___Ca(NO3)2

52

13. ___Al2O3 + ___H2SO4 → ___Al2(SO4)3 + ___H2O 14. ___Fe + ___Br2 → ___FeBr3 15. ___Al(OH)3 + ___H2SO4 → ___Al2(SO4)3 + ___H2O 16. ___C2H2 + ___O2 → ___CO2 + ___H2O

17. ___Li2O + ___H2O → ___LiOH 18. ___NH3 + ___H2CO3 → ___(NH4)2CO3 19. ___C2H5OH + ___O2 → ___CO2 + ___H2O 20. ___H3PO4 + ___Ca(OH)2 → ___Ca3(PO4)2 + ___H2O 21. ___HBr + ___K2SO3 → ___H2O + ___SO2 + ___KBr 22. ___Na + ___H2O → ___NaOH + ___H2 23. ___Al + ___Fe2O3 → ___Al2O3 + ___Fe 24. ___KClO3 → ___KCl + ___O2 25. ___(NH4)2SO4 + ___BaCl2 → ___NH4Cl + ___BaSO4 26. ___CH3OH + ___O2 → ___CO2 + ___H2O

53

B. Write a balanced chemical equation for each reaction described below (include states):

1. Solid magnesium oxide is produce by heating solid magnesium metal in the presence of oxygen gas.

2. Solid calcium reacts with nitric acid to form aqueous calcium nitrate and hydrogen gas.

3. Aqueous hydrochloric acid reacts with solid manganese(IV) oxide to produce aqueous

manganese(II) chloride, liquid water, and chlorine gas.

55

Conductivity

Objectives: • Students will study conductivity as they compare:

o Ionic and covalent compounds

o Solutions of varying concentrations

o Weak acids/bases and strong acids/bases

Techniques: Upon completion of this lab, students will have learned:

• To predict conductivity of a compound based on chemical formula

• To classify compounds as non-electrolytes, weak electrolytes, or strong electrolytes based on

formula


• A Molecular View of Elements and Compound Section 5.4

• Aqueous Solutions and Solubility Section 7.5

• Strong and Weak Acids and Bases Section 14.7

Introduction:

You might remember from Section 4.6 of Tro that metals are good conductors of electricity, while

non-metals are poor conductors. A quick electrical conductivity test can identify an element as

either a metal or non-metal. What about compounds? In this lab, you will explore the relationship

between conductivity, chemical formula, and physical state. You will examine the differences in

conductivity for ionic versus molecular compounds, and how acid strength affects conductivity of

acid solutions.

56

Procedure:

Safety Precautions: • Normal precautions need to be taken with acid and base solutions • Dispose of all solutions in the appropriate waste container as instructed

•

Part A: Conductivity Meter Operation 1. Clean and dry a well plate. 2. Place approximately 1 mL of solutions SE, WE and NE into 3 separate wells. (SE stands for

strong electrolyte, WE for weak electrolyte, NE for non-electrolyte.) 3. Immerse the electrodes of the meter into the SE solution. Record your observations. 4. Rinse and dry the meter electrodes. Repeat steps 3 and 4 for the WE and NE solutions.

Part B: Conductivity of pure water, solids and solutions

1. Place ~ 1mL (20 drops) of pure water (distilled or deionized) and ~0.1 g NaCl into separate clean, dry wells. Measure the conductance of each substance (SE, WE, or NE).

2. Mix the water and the solid NaCl in one well. Stir until the solid dissolves. Measure and record the conductance of this saline solution.

3. Dilute the saline solution by taking 1 drop of it and placing it in a clean well. Use a clean dropper to add 9 drops of pure water. Stir the mixture, and then measure the conductivity.

4. Make an even more dilute saline solution by diluting 1 drop of the saline solution from step #3 with 9 drops of pure water. Stir the mixture, and then measure the conductivity.

5. Repeat steps 1-4 using table sugar (C12H22O11), potassium nitrate (KNO3), and calcium acetate [Ca(C2H3O2)2] solids in place of the NaCl.

Part C: Conductivity of liquids

1. Measure the conductivity of approximately 1 mL of ethanol (C2H5OH). 2. Measure the conductivity of approximately 1 mL of tap water.

Part D: Conductivity of acid solutions

1. Measure the conductivity of approximately 1 mL of 0.1 M HC2H3O2, 0.1 M H2SO4, and 0.1 M HCl in 3 separate clean, dry wells.

57

Name: __________________________________________Date due: ___________________________

Conductivity Pre-lab Questions: Read the relevant textbook sections and the entire lab before answering pre-lab questions.

1. Describe the major difference between an ionic and a molecular compound.

2. Describe the difference between a strong acid and a weak acid.

3. Identify the following as strong (S) or weak (W) acids.

a. _______ 0.1 M HCl (aq)

b. _______ 0.1 M H2SO4 (aq)

c. _______ 0.1 M HC2H3O2 (aq)

4. Identify the following as ionic (I) or molecular (M) compounds.

a. _______ H2O

b. _______ NaCl

c. _______ C12H22O11

d. _______ Ca(C2H3O2)2

e. _______ C2H5OH

59



Conductivity

Data: Part A: Conductivity Meter Operation.

Solution Meter Observation

SE

WE

NE

How the meter indicates conductivity. Part B: Conductivity of pure water, solids and solutions.

Substance Conductivity Substance Conductivity Pure water, H2O (l) Pure water

NaCl (s) C12H22O11 (s)

NaCl (aq) C12H22O11 (aq)

Diluted NaCl (aq) Diluted C12H22O11 (aq)

Most Diluted NaCl (aq) Most Diluted C12H22O11 (aq)

Substance Conductivity Substance Conductivity

Pure water Pure water

KNO3 (s) Ca(C2H3O2)2 (s)

KNO3 (aq) Ca(C2H3O2)2 (aq)

Diluted KNO3 (aq) Diluted Ca(C2H3O2)2 (aq)

Most Diluted KNO3 (aq) Most Diluted Ca(C2H3O2)2 (aq)

60

Part C: Conductivity of liquids.

Substance Conductivity

Ethanol, C2H6O (l)

Tap water

Part D: Conductivity of acid solutions.

Substance Conductivity

0.1 M HC2H3O2 (aq)

0.1 M H2SO4 (aq)

0.1 M HCl (aq)

Conclusion Questions: 1) Draw a picture of what happens to the formula units in the following aqueous solutions. Do not draw the water molecules; just the particles present after dissolving the compounds. Assume you have 4 formula units dissolved in the appropriate beaker. See an example on page 215 of text. You may use simple characters or geometric shapes to represent each ion. For example a circle may represent a sodium ion and a square the chlorate ion for the compound sodium chlorate. Label one of both shapes in your drawing. The line on each beaker represents the surface of the solution.

KNO3 (aq) NaCl (aq) 2) Does each the first solutions prepared with these ionic compounds in part B conduct an electrical current? 3) What common feature is present in both solutions?

61

4) You tested a solution CaC2H3O2 for electrical conductivity. What might explain the electrical conductance of CaC2H3O2? 5) Based on your experimental results in part B, describe how dilution of the solutions affected the conductivity. 6) Hypothesize why tap water is a better conductor of electricity than pure water. 7) Hypothesize why solutions of NaCl conduct electricity, but solid NaCl does not. 8) In aqueous solutions containing molecular compounds the molecules of the dissolved compound are separated and distributed throughout the water. For example a solution of sugar consists of individual sugar molecules distributed through the water. Draw a picture of what happens to the molecules in the following aqueous solution. Do not draw the water molecules; just the particles present after dissolving the compounds. Assume you have 4 molecules dissolved in the solution.

C2H6O (aq)

62

9) Draw a picture of what happens to the formula units in the following aqueous solutions. Do not draw the water molecules; just the particles present after dissolving the compounds. Assume you have 4 acid molecules dissolved in each beaker. See examples on page 499-500 of text.

HC2H3O2 (aq) HCl (aq) 10) Is HCl a strong or weak acid? 11) Is HC2H3O2 a strong or weak acid? 12) What is the difference between your drawings of the two acids? 13) Is HCl (aq) a strong, weak or nonelectrolyte? 13) Is HC2H3O2 (aq) a strong, weak or nonelectrolyte? 15) Explain how acid strength affects electrical conductivity.

63

16) What is common among all the solutions that conducted an electrical current? 17) Suppose that each compound below is placed in deionized water. Indicate whether each mixture is a strong, weak or nonelectrolyte. MgI2 Al(NO3)3

HC2H3O2 Cs3PO4

HClO4 (NH4)2SO4

NaOH CH2O

65

Chemical Reactions: Classification and Prediction of Products

Objectives: In this lab, students will:

• Classify reactions • Predict products of reactions

Skills: Upon completion of this lab, students should have learned to

• Perform and observe chemical reactions • Identify chemical changes


• Precipitation Reactions Section 7.6 • Acid-Base and Gas Reactions Section 7.8 • Oxidation-Reduction Reaction Section 7.9 • Classifying Chemical Reaction Section 7.10

Introduction:

Chemical reactions can be classified in many ways. One way is to classify the reaction by what is formed during the reaction (precipitation, acid-base, gas evolution, or oxidation-reduction reactions). Another way, and the one used in this lab, is to classify reactions by what the atoms or groups of atoms do. In this classification method, reactions are classified as double displacement, single displacement, synthesis, or decomposition reactions. Another common type of reaction that will be examined is the combustion reaction. Reactions occur in these general forms:

Type of Reaction Generic Equation

Double Displacement AB + CD AD + CB

Single Displacement A + BC AC + B

Synthesis (or combination) A + B AB

Decomposition AB A + B

Combustion Hydrocarbon + O2 CO2 + H2O

The key to classifying a reaction is to look closely at the reactants (and the products, if you know them). In the table above, whenever a letter is written in the equation by itself, it usually means that material is an element (like the “A” on the reactant side of the single displacement reaction). Decomposition reactions have exceptions to this. If the letters are written together (as in “AB” on the reactant side of the double displacement reaction), it means that material is a compound.

66

Procedure:

Safety Precautions: • Before you start any reaction Carefully read the whole procedure for that reaction, taking note of all safety concerns like

hot glass, a flame, or acids and bases. If you are not sure of the procedure, ask the instructor BEFORE running the reaction.

Gather all of the reactants and equipment at your work station. Do NOT perform the reaction at the side bench.

Return chemicals to the side bench when you are finished with them. Move slowly, and clean-up after each reaction. Dispose of all chemicals in the proper waste container.

Special Notes:

• Write observations before, during, and after completion of each reaction.

• Unless instructed otherwise,

o All reactions are performed in large test tubes.

o Add chemicals in the written order.

o Use graduated pipets to dispense solutions.

o All masses used are approximate and you can use any amount within ± 10% of the

stated value. Do not waste time trying to get exact masses.

o When heating chemicals in a test tube, clamp the test tube to a stand. Angle the test

tube to ensure it is not aimed at anyone. Use the Bunsen burner with a moderately

sized blue flame. (Only #7 will be heated for this lab.)

• Observe each reaction for a color change, precipitation, gas evolution, and/or temperature

change.

• In reactions that produce gases, you will test to see if the gas produced is H2, O2, or CO2.

These tests are performed as the reaction is fully underway and the test tube has filled with

the gaseous product. Don’t wait until the reaction stops.

o H2 test - A wooden splint is lit in a Bunsen burner flame and first placed near the

top of the test tube, and then slowly inserted into the test tube. If the gas “pops”

hydrogen gas is present.

o O2 and CO2 test - A wooden splint is lit and allowed to burn for a few seconds and

then blown out. The red ember of the wooden splint is placed inside the test tube.

If the ember glows very brightly or the flame is rekindled, oxygen gas is present. If

the ember quickly goes out, carbon dioxide gas is present.

67

Reactions: Perform the following seven reactions. Record your observations on the data sheet.

Complete all of the reactions and your observations before you try to classify, write, and/or balance the

chemical equations.

Reaction 1: Mix 2 mL of 0.1 M calcium chloride with 2 mL of 0.1 M sodium phosphate. Reaction 2: Add 2 mL of 3 M hydrochloric acid to a small piece of zinc. Test for H2. Reaction 3: Mix 2 mL of vinegar (HC2H3O2) with a very small amount of baking soda (NaHCO3).

Test for O2 and CO2. Reaction 4: Place 1 g of ammonium chloride and 2 g of strontium hydroxide in a large test tube. Stir

the solid contents with a stirring rod for 3-5 minutes. Gently waft your hand across the top of the tube, and note the odor. Try to identify the common household chemical that has this smell.

Reaction 5: Mix 2 mL of 3 M sulfuric acid with 4 mL of 3 M sodium hydroxide. Reaction 6: Mix 2 mL of 0.1 M lead(II) nitrate with 2 mL of 0.1 M potassium iodide. Reaction 7: Heat 0.5 g of copper(II) hydroxide.

69

Name: _________________________________________ Date due: ___________________________

Chemical Reactions: Classification and Prediction of Products Pre-lab Questions: Read the relevant textbook sections and the entire lab before answering pre-lab

questions.Classify each reaction as either a double displacement (DD), single displacement (SD), synthesis (SYN), decomposition (DEC), or combustion (COM) reaction.

• Predict and write the correct formulas for the products of each reaction. • Balance the equation. • Include the phase of each of the products.

1. ________ KCl(aq) + Pb(NO3)2(aq) 2. _______ C3H8O(l) + O2(g) 3. _______ AgNO3(aq) + Mg(s) 4. ________ Ca(s) + N2(g) 5. ________ H2CO3(aq)

71



Chemical Reactions: Classification and Prediction of Products

Data: Record your observations and any test results. Reaction #: 1. 2. 3. 4. 5. 6. 7.

72

Results: • Classify each reaction as either a double displacement (DD), single displacement (SD),

synthesis (SYN), decomposition (DEC), or combustion (COM) reaction. • Predict and write the correct formulas for the products of each reaction. • Write and balance the chemical equation. • Identify the physical state of each chemical as (aq), (s), (l), or (g).

1. Classification _______ Balanced Chemical Equation: 2. Classification _______ Balanced Chemical Equation: 3. Classification _______ Balanced Chemical Equation: 4. Classification _______ Balanced Chemical Equation:

5. Classification _______ Balanced Chemical Equation: 6. Classification _______ Balanced Chemical Equation:

7. Classification _______ Balanced Chemical Equation: _____Cu(OH)2 (s) __________( ) + _____ CuO ( )

73

Double Displacement Reactions

Objectives: In this lab, students will • Identify six unknown solutions by mixing and observing their reactions

Skills: Upon completion of this lab, students should have learned to

• Write molecular equations Textbook References: (Tro, Introductory Chemistry, 4th Ed.) to be read PRIOR to lab.

• Evidence of Chemical Reactions Section 7.2 • Aqueous Solutions and Solubility Section 7.5 • Precipitation Reactions Section 7.6 • Writing Chemical Equations Section 7.7 • Acid-Base and Gas Reactions Section 7.8

Introduction:

Double Displacement reactions are very common and occur in one of three main ways:

• precipitation reactions

• acid-base reactions

• gas evolution reactions

In this lab, you will be given 6 “unknown” solutions. You will mix different combinations of two

solutions together and observe the results. Using these results and the predictions from the pre-

laboratory exercises, you will use logic to identify the 6 unknown solutions.

To help in your predictions of reaction outcomes, you will use your knowledge of precipitation and

solubility to help identify reactions. It will also help to know that:

• Acid-base reactions often produce heat

• Reactions between an acid and a compound containing CO32- ions will produce CO2 gas

74

Procedure: You will be working with 6 unknown solutions. Each solution is one of the following:

1.5 M H2SO4 1.0 M K3PO4 0.1 M Mg(NO3)2 1.0 M Na2CO3 3.0 M NaOH 0.1 M SrCl2

Safety Precautions and Special Notes:

• Since you do not initially know the identity of the solutions, treat every solution as potentially

hazardous.

• Write observations of the reactions.

• Unless instructed otherwise:

o All reactions are performed in large test tubes.

o Use automatic dispensers to dispense solutions.

1. Place 2 mL (one pump) of solution A in a small test tube, then add 2 mL of solution B. Observe

the reaction and record your results in the data table. Dispose of your waste in the appropriate

waste container.

2. Repeat this process with 2 mL quantities in a small clean test tube for each of the following

combinations of solutions:

• AB, AC, AD, AE, AF

• BC, BD, BE, BF

• CD,CE,CF

• DE,DF

• EF

75

SO

LU

BILI

TIE

S O

F IO

NIC

CO

MPO

UN

DS:

A

PRO

XIM

AT

E #

OF

GR

AM

S O

F SO

LUT

E P

ER 1

00 G

RA

MS

OF

SOL

UTI

ON

C2H

3O2−

ss

60

42

26

20

s 7 31

75

40

70

1 32

27

25

An

arbi

trary

stan

dard

for s

olub

ility

is th

at a

com

poun

d is

cal

led

solu

ble

if gr

eate

r tha

n 1

gram

dis

solv

es in

100

gra

ms o

f sol

utio

n.

vs =

ver

y so

lubl

e, s

= so

lubl

e, ss

= sl

ight

ly so

lubl

e, i

= in

solu

ble,

d =

dec

ompo

ses

Br−

s 43

51

59

3 54

56

s 0.8 62

50

40

8 x

10-6

48

50

82

CO

32−

50

2 x

10-3

6 x

10-3

i i i

1 x

10-4

1.3

0.07

52

3 x

10-3

22

1 x

10-3

2 x

10-2

Cl−

31

27

26

43

50

35

42

70

1 45

35

25

2 x

10-4

26.4

35

79

CrO

42−

25

4 x

10-4

14 i i

7 x

10-6

50

42

39

4 x

10-3

47

0.12

i

OH

−

1 x

10-4

47

4 0.16

i

3 x

10-4

3 x

10-4

1 x

10-5

0.02

11.3

2 x

10-3

53

d 52

1

4 x

10-4

IO3−

2 0.02

0.3

0.1 1 0.1

0.04

2 x

10-3

45

8 7.5

4 x

10-3

8 0.03

0.9

I−

s,d

63

68

68

s 65

1.1

0.07

62

58

59

3 x

10-6

64

64

83

NO

3−

42

66

8 56

64

50

55

46

35

43

41

27

70

47

42

56

PO43−

i 26 i

2 x

10-3

i i i i

1 x

10-5

0.03

0.02

47

6 x

10-4

11 i i

SO42−

27

43

2 x

10-4

0.2 9 26

17

ss

4 x

10-3

26

28

10

0.8 20

0.01

30

S2-

d vs

d 0.02

i

4 x

10-4

2 x

10-4

3 x

10-1

7

9 x

10-5

vs

d s

7 x

10-1

3

16

s,d

10-8

Al3+

NH

4+

Ba2+

Ca2+

Ce3+

Co2+

Cu2+

Fe3+

Pb2+

Li+

Mg2+

K+

Ag+

Na+

Sr2+

Zn2+

77

Name: __________________________________________Date due: ___________________________


Pre-lab Questions: Read the relevant textbook sections and the entire lab before answering pre-lab questions. 1. For the three solutions 0.1 M Ba(OH)2, 1 M Na2CO3, and 2 M HCl,

a. Complete and balance each double displacement reaction (include states). . b. Determine whether each reaction will produce a precipitate (ppt), a gas (gas), heat (heat), or if there

will be no observable reaction (nr), and write the correct abbreviation in the prediction matrix below.

1. ____Ba(OH)2(aq)+ ____Na2CO3(aq)

2. ____Ba(OH)2(aq)+ ____HCl(aq)

3. ____HCl(aq) + ____Na2CO3(aq)

Prediction Matrix: Na2CO3 HCl

1 2 Ba(OH)2

3 Na2CO3

2. Assume that each of the three solutions, 0.1 M Ba(OH)2, 1 M Na2CO3, and 2 M HCl, is placed in a separate

bottle labeled either A, B, or C. Based on the results observed for the reactions, determine which solution is in which bottle.

B C heat gas A

ppt B

A ____________________ B ____________________ C ____________________

78

3. Predict the products and write the balanced chemical equation for each reaction. If the reactants and products are all aqueous, no reaction took place, so write nr. Label the physical states (g), (s), (aq), or (l) of each product. (Use the solubility table.)

_____ H2SO4(aq) + _____ K3PO4(aq)

_____ Na2CO3(aq) + _____ K3PO4(aq)

_____ NaOH (aq) + _____ K3PO4 (aq)

_____ Mg(NO3)2 (aq) + _____ K3PO4(aq)

_____ SrCl2(aq) + _____ K3PO4(aq)

_____ Na2CO3(aq) + _____ H2SO4(aq)

_____ NaOH(aq) + _____ H2SO4(aq)

_____ Mg(NO3)2(aq) + _____ H2SO4(aq)

_____ SrCl2(aq) + _____ H2SO4 (aq)

_____ NaOH (aq) + _____ Na2CO3 (aq)

79

_____ Mg(NO3)2(aq) + _____ Na2CO3(aq)

_____ SrCl2(aq) + _____ Na2CO3(aq)

_____ Mg(NO3)2(aq) + _____ NaOH(aq)

_____ SrCl2(aq) + _____ NaOH(aq)

_____ SrCl2(aq) + _____ Mg(NO3)2(aq)

4. For the six solutions you will use in the procedure of this lab, write whether each reaction will produce a precipitate (ppt), a gas (gas), heat (heat), or if there will be no observable reaction (nr) based on the reactions in question 3. You should predict 7 reactions will form precipitates, 1 reaction will produce heat, 1 reaction will produce gas, and 6 have no observable reaction. Copy this matrix to the data section.

Prediction Matrix:

H2SO4 Na2CO3 NaOH Mg(NO3)2 SrCl2

nr K3PO4

H2SO4

Na2CO3

NaOH

Mg(NO3)2

81




Data: Record your observations and any test results. • Use the abbreviations precipitate (ppt), a gas (gas), heat (heat), no observable reaction (nr) to record

your observations in the matrix below. • Hint: 7 reactions should form precipitates, 1 reaction should produce heat, 1 reaction should produce

gas, and 6 should have no observable reaction.

Observation Matrix: B C D E F

ppt nr heat nr gas A

ppt ppt nr ppt B

nr ppt nr C

ppt nr D

ppt E

Prediction Matrix: (copied from the pre-lab page)

H2SO4 Na2CO3 NaOH Mg(NO3)2 SrCl2

nr K3PO4

H2SO4

Na2CO3

NaOH

Mg(NO3)2

82

Results: Use the data in the Observation Matrix and the Prediction Matrix to determine the identity of solutions A through F.

Solution:

A ____________________________

B ____________________________

C ____________________________

D ____________________________

E ____________________________

F ____________________________

83

Lewis Structures

Objectives: In this activity, students will: • Draw Lewis dot structures with ionic bonds, covalent bonds, or both ionic and covalent bonds

Skills: Upon completion of this activity, students will have learned

• To identify the number of valence electrons in main group elements • To draw structures of molecules, compounds, and ions


• Electron Configurations and the Periodic Table Section 9.7 • Covalent Lewis Structures: Electrons Shared Section 10.4 • Writing Lewis Structures for Covalent Compounds Section 10.5 • Resonance Section 10.6

Introduction:

Chemical compounds contain ionic bonds, covalent bonds, or both. Ionic materials, like sodium chloride

and calcium carbonate contain ions and are often easily recognized by the fact that they contain a metal and

a non-metal in their formula. Covalent compounds, like water and ethyl alcohol contain only non-metal

atoms in their formula. (Covalent compounds are also known as molecular compounds.) There are

exceptions to these simple rules for identifying compounds, but we will not focus on them here.

Valence electrons are the electrons that participate in bonding. Ionic bonds result from the gain or loss of

electrons, while covalent bonds are the sharing of electrons by two atoms. The reason for the covalent

sharing of electrons is so that the atoms can fill their valence shells. Useful chemical information can be

gained by looking at how the electrons are shared in covalent compounds. Drawing a Lewis structure is one

of the most common ways to illustrate this sharing.

84

Drawing Lewis Structures Here are some simple guidelines for drawing correct Lewis structures.

Step 1: Identify Bonding Types

• Write down the ions present. (Look for metals.) They will have an ionic bond between

them. If the ion is polyatomic, it will be made up of covalent bonds.

• If no ions are present, all bonds will be covalent. (All atoms will be nonmetals.)

Step 2: Valence Electrons

• Use the periodic table to count the total number of valence electrons for all atoms in the

molecule or ion. (Work on ions separately.)

• Add one additional electron for each negative charge of an anion or subtract one for each

positive charge of a cation.

Step 3: Connect Atoms

• Connect atoms with lines to represent bonds between atoms.

• Hydrogen is always terminal.

• The least electronegative atom is usually the central atom (unless otherwise noted).

• Symmetrical structures are preferred.

Step 4: Assign Electrons to the Terminal Atoms and Fill Their Valence Shells

• Place lone pairs (non-bonding electrons) of electrons around each terminal atom to

complete each atom’s octet (except for hydrogen).

Step 5: Recount and Adjust

• If you have extra electrons, place around central atom as pairs.

• If you don’t have enough electrons to complete an octet around the central atom, make

double and triple bonds by sharing electron pairs from the terminal atoms.

• Always keep in mind the octet rule.

Step 6: Check

• All atoms have an octet (hydrogen has a duet).

• All valence electrons were used.

Step 7: Put Ions Together

• Put brackets around the ion, and write the charge on the outside of the brackets.

• Place the ions next to each other so that each positive ion is next to a negative ion.

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Example #1, Water, H2O

1. Identify Bonding Types. All of the atoms are nonmetals, so all bonds will be covalent. 2. Valence Electrons.

Atoms Valence Electrons

H 1 H 1 O 6

Total 8

3. Connect Atoms. Draw simple and symmetrical structures. The possible arrangements of 3 atoms and 2

bonds for H2O are:

H H O H O H O H H

Hydrogen is always a terminal atom. Also, H-O-H is the most symmetrical, so use H-O-H.

4. Assign Electrons to the Terminal Atoms and Fill Their Valence Shells. Of the 8 valence electrons

we started with, 4 electrons were used in making the 2 bonds. That leaves 4 electrons to fill the valence

shells. Each hydrogen has 2 electrons in its valence shell (1 bond = 2 electrons) and is filled.

5. Recount and Adjust. The 4 remaining electrons are placed around the oxygen atom as 2 pairs of non-

bonding electrons (also known as lone pairs).

H O H....

6. Check. Two bonds and two lone pairs were used; the total electrons used were eight. Hydrogen has a

duet and oxygen has an octet.

7. Put Ions Together. No ions are present.

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Example #2, Formaldehyde, CH2O 1. Identify Bonding Types. All of the atoms are nonmetals, so all bonds will be covalent. 2. Valence Electrons.


C 4 O 6 H 1 H 1

Total 12

3. Connect Atoms. Draw simple and symmetrical structures. Possible arrangements are

H C O H H C H

O

H O H

C

Hydrogens are terminal atoms in all three structures. The last two structures are more symmetrical. Generally, the atom that has the lowest electronegativity is the central atom. Carbon has a lower electronegativity than O, so C is the better central atom. That means the middle structure is “better”.

H C H

O

4. Assign Electrons to the Terminal Atoms and Fill Their Valence Shells. Of the 12 valence electrons

we started with, 6 electrons were used in making the 3 bonds. That leaves 6 electrons to fill the valence shells of the terminal atoms. The hydrogens are filled, so the remaining 6 electrons are placed around the oxygen atom.

H C H

O

5. Recount and Adjust. In the above drawing, the hydrogens and oxygens are filled; the hydrogens have

duets and the oxygen has an octet. However, the carbon does not have an octet. One of the lone pairs of oxygen is shared covalently to give both oxygen and carbon octets.

H C H

O

6. Check. All valence shells are filled and the correct number of electrons has been used.

7. Put Ions Together. No ions are present.

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Example #3, Sodium Cyanide, NaCN 1. Identify Bonding Types. Sodium is a metal cation, and cyanide is a nonmetal anion. An ionic bond

exists between sodium and cyanide. Cyanide is polyatomic and made up of nonmetals, so there will be covalent bonds within cyanide. Work on the ions separately (as Na+ and CN

-) and then put them together at the end.

2. Valence Electrons. Na+


Na 1 Positive ion charge -1

Total 0

3. Connect Atoms. Sodium is not covalently bonded to any other atom, and has zero valence electrons around it, so don’t draw any dots.

Na There aren’t any covalent bonds, so skip those steps and just put brackets around sodium along with its

charge.

Na+

The sodium cation is finished, so now work on the cyanide anion. 2. Valence Electrons. CN

-

Atoms Valence Electrons C 4 N 5

Negative ion charge 1 totals 10

3. Connect Atoms.

C N N C

The structure is linear, so either one of the above can be used.

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4. Assign Electrons to the Terminal Atoms and Fill Their Valence Shells. Of the 10 valence electrons we started with, 2 electrons were used in making the single bond. That leaves 8 electrons to fill the valence shells. Lone pairs are placed on the carbon and nitrogen until the 8 electrons are used up.

5. Recount and Adjust. In the above drawing, neither the carbon nor the nitrogen is filled. Each atom

shares a lone pair so that each will have an octet. C N

6 Check. All valence shells are filled and the correct number of electrons has been used. Brackets are placed around anions and the overall charge is written outside of the brackets.

....N C[ ]-

7. Put Ions Together. Sodium cyanide has the following structure.

Na+

....N C[ ]-

C N

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Name: __________________________________________Date due: ___________________________

Lewis Structures Pre-lab Questions: Read the relevant textbook sections and the entire lab before answering pre-lab

questions. 1. Define the following terms in your own words.

a. valence shell

b. valence electron

c. covalent bond

d. octet rule

e. resonance structure 2. Determine the number of valence electrons in an atom of each of these elements.

a. Determine the number of electrons needed to fill the valence shell.

Atom Valence Electrons # electrons needed

C

O

H

Cl

N

S

90

Draw Lewis structures for:

CH4

NH4+

SO42-

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Name: _______________________________________Date lab performed: __________________


Lewis Structures

A. Draw Lewis Structures of Molecules and Polyatomic Ions. Draw the Lewis structure for each molecule or ion below. Follow the same process used in the examples.

1. F2

2. O2

3. IO2−

4. CH4

5. CO2

6. NH4+

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7. SO32−

8. C2H6

9. C2H4

10. ClO4−

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11. Cl2CO (carbon is the central atom)

12. CH3OH (O-H bond)

13. NO2+

14. N2

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B. Resonance Structures

Resonance structures occur when more than one correct Lewis structure can be drawn for a molecule or ion. One resonance structure differs from another only by the placement of a double (or triple) bond. The skeleton structure of atoms does not change. Ozone, O3, has been done as an example. The double arrow indicates that the two structures are resonance structures.

O O O OOO

Draw the indicated number of resonance structures for these formulas.

1. SO2 (2 structures)

2. NO2− (2 structures)

3. SeO2 (2 structures)

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C. Ionic Compounds

Draw the Lewis structure for each ionic compound below. Follow the same process used in the examples.

1. CaO

2. MgSO4

3. Ca(OH)2

97

Gas Laws

Objectives: In this experiment, students will: • Determine the atomic mass of zinc through the combination of gas law concepts and

stoichiometry • Calculate the molar mass of natural gas using the Ideal Gas Law

Skills: Upon completion of this lab, students will have learned to:

• Apply Dalton’s law of partial pressures to a gaseous mixture • Calculate the number of moles of a gas generated using the Ideal Gas Law • Use stoichiometry to convert moles of a product to moles of a reactant in a chemical reaction • Calculate the atomic mass of an element from an experimental mass and number of moles


• Dalton’s law of partial pressures Section 11.9 • Ideal Gas law Section 11.8 • Stoichiometry Section 8.1, Section 11.10

Introduction: When zinc reacts with hydrochloric acid, the following single displacement reaction occurs:

Zn (s) + 2 HCl (aq) → ZnCl2 (aq) + H2 (g) In Part A of this experiment, a known mass of zinc is reacted with an excess of hydrochloric acid in order to totally consume the zinc. It is possible to calculate the number of moles of hydrogen gas produced in this reaction by rearranging the Ideal Gas Law.

𝑃𝑉𝑅𝑇 = 𝑛

Where n = number of moles of hydrogen gas produced, P = pressure due to hydrogen gas, V = volume of hydrogen gas produced, R = ideal gas constant = 0.08206 𝐿 𝑎𝑡𝑚

𝐾 𝑚𝑜𝑙, and

T = temperature of hydrogen gas produced.

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Because the hydrogen gas is collected over water, it also contains some water vapor. The total pressure equals the pressure from hydrogen plus the pressure from water. The partial pressure due to hydrogen can be calculated by rearranging Dalton’s law of partial pressures as follows:

PH2 = PT - PH2O

Where PH2 = partial pressure due to hydrogen gas PT = total pressure of gaseous mixture (atmospheric pressure) PH2O = partial pressure due to water vapor

PH2 should be converted to atm using 1 atm = 760 mm Hg or 1 atm = 29.92 in Hg.)

Water Vapor Pressure at Various Temperatures

Temperature ˚ C

Vapor Pressure mm Hg

Temperature ˚ C


Temperature ˚ C


10 9.2 17 14.5 24 22.4 11 9.8 18 15.5 25 23.8 12 10.5 19 16.5 26 25.2 13 11.2 20 17.5 27 26.7 14 12.0 21 18.6 28 28.3 15 12.8 22 19.8 29 30.0 16 13.6 23 21.1 30 31.8

The balanced equation above tells us that for every mole of zinc consumed, one mole of hydrogen is produced. Therefore, the number of moles of hydrogen produced is equal to the number of moles of zinc consumed. Knowing the mass of zinc (in grams), the atomic mass of zinc can be calculated as follows:

Molar Mass = 𝑚𝑎𝑠𝑠 (𝑔)

𝑚𝑜𝑙𝑒𝑠 (𝑚𝑜𝑙)

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Procedure: Safety Precautions: Hydrogen and air mixtures are extremely explosive. Keep all open flames away from flasks containing hydrogen. Determining the atomic mass of zinc. You will complete TWO trials of this procedure

1. Set up the apparatus as shown below. Fill the 500 mL Florence flask to the neck with water as shown in the diagram, and put 25 mL of 6 M hydrochloric acid in the 250 mL Erlenmeyer flask.

Erlenmeyer Flask Florence Flask Collection Beaker (400 mL or larger)

2. Remove both the stopper from the Erlenmeyer flask and the clamp from the rubber tubing between the Florence flask and the beaker. Using a rubber bulb, apply compressed air to the tube inserted through the stopper. When water begins to flow through the tubing, remove the rubber bulb and clamp the tubing through which water is flowing. Remove the stopper gently from the Florence flask, careful to keep the long glass tube under the water. Pour the water collected in the beaker back into the Florence flask, and replace the stopper.

3. Obtain 0.900 g - 1.000 g of zinc. Remove the stopper from the Erlenmeyer flask, and transfer the zinc into the hydrochloric acid in the flask. Immediately replace the stopper in the Erlenmeyer flask and remove the clamp from the rubber tubing.

4. After the zinc has been completely consumed by the HCl, allow the gas in the flasks to cool to room temperature (a few minutes). Raise the flask or the beaker until the levels of water in the flask and the beaker line-up and then clamp the rubber delivery tubing before gently removing the beaker. Use a 500 mL graduated cylinder to measure the volume of water in the beaker. Use a glass thermometer to measure the temperature of water in the beaker.

. 5. Dispose of the contents of the Erlenmeyer flask in the appropriate waste container and repeat the experiments.

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Name: __________________________________________Date due: ___________________________

Gas Laws Pre-Lab Questions: Read the relevant textbook sections and the entire lab before answering pre-lab questions. 1. Perform each of the following pressure conversions:

a. 1.020 atm to mm Hg

b. 30.65 in Hg to mm Hg 2. 472 mL of H2 gas was collected over water when 1.256 g Zn reacted with excess HCl. The atmospheric

pressure during the experiment was 754 mm Hg and the temperature was 26 oC.

a) Write the balanced chemical equation for this reaction.

b) What is the water vapor pressure at 26 oC? ____________mm Hg

c) What is the partial pressure (in atmospheres) of dry hydrogen gas in the mixture? (Show work.)

d) Calculate the number of moles of H2 produced by this reaction using the ideal gas law. (Show work.)

e) Use the data from the experiment to calculate the experimental molar mass of zinc in g/mol.

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f) What is the molar mass of zinc from the Periodic Table (known value)? __________ g/mol

g) Calculate the percent error for the experimental molar mass of Zn.

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Gas Laws Data: Report all measurements in the correct number of significant figures and units. For all responses requiring calculations, the mathematical setup must be shown. Atomic Mass of Zinc

1. Mass of zinc used ____________ g ____________ g

2. Volume of water collected ____________L ____________L

3. Temperature of water collected ____________ °C ____________ °C

____________ K ____________ K

2. Atmospheric pressure ____________ in Hg ____________ in Hg

____________ mm Hg ____________ mm Hg

3. Partial pressure of water vapor at temperature recorded above (See water vapor pressure table in the introduction.) ____________ mm Hg ____________ mm Hg

4. Partial pressure of dry hydrogen gas (show work)

____________ mm Hg ____________ mm Hg

____________ atm ____________ atm

7. Moles of hydrogen gas collected ____________mol ____________mol

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8. Write the balanced chemical equation. ____________________________________________

9. Moles of zinc consumed (Use the mole ratio of hydrogen gas to the moles of zinc from the chemical equation.) ____________ moles Zn ____________ moles Zn

9. Calculated experimental molar mass of

zinc (MUST show work) ____________ ____________

11. Known value for the molar mass of

zinc ____________ ____________

12. Percent error (MUST show work)

____________ ____________ Conclusion Question:

After starting with 0.343 g Al and excess HCl, the volume of H2 gas collected over water was 472 mL. The atmospheric pressure was 754 mmHg and the temperature was 21oC. 1. Write the balanced chemical equation for this reaction. 2. Calculate the pressure in atm of dry hydrogen gas in the mixture. (Show work.)

105

3. How many moles of H2 were produced by this reaction? (Show work.) 4. Calculate the experimental molar mass of Al. (Show work.)

5. What is the percent error for the experimental molar mass?

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Acid-Base Titrations

Objectives: In this lab, students will: • Determine the concentration of an unknown acid solution.

Skills: Upon completion of this lab, students will have learned to:

• Use volumetric pipets and burets • Perform titrations between strong acids and strong bases


• Molarity Section 13.6 • Solution Stoichiometry Section 13.8 • Acid-Base Titration Section 14.6

Introduction:

Titration is a method used to determine the concentration of an acid solution or a base solution. Water hardness of your city’s public water supply is probably determined using a titration. Often when performing an acid-base titration, a base solution of known concentration is slowly added to the acid solution of unknown concentration. A strong base reacts with a strong acid to form water according to the general reaction below.

HA(aq) + MOH(aq) H2O(aq) + MA(aq)

This reaction is often called a neutralization reaction because the strong base neutralizes the strong acid. The chemist knows when all of the acid has been neutralized by either taking the pH of the solution during the titration or by using an indicator. The indicator changes color as the pH of the solution becomes neutral. By titrating—adding base—until the indicator changes color, the chemist knows when the equivalence point is reached. The equivalence point is when the moles of OH- added are exactly enough to neutralize the moles of H+ present in the acid solution.

A common indicator to use is phenolphthalein because it changes from colorless in acidic solution to pink when the solution becomes basic.

Glassware:

Titrations are performed using a buret. A buret is a long, narrow, calibrated tube which is designed to deliver (TD) a quantity of a liquid. The stopcock at the bottom controls the flow of the liquid. Due to the narrow top, a funnel is used to fill a buret.

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Buret calibrations increase in value from the top to the bottom, the reverse of the graduated cylinder.

Graduated Cylinder Buret

When reading the buret, you can hold a colored index card or something else behind the buret to help you see the meniscus better. Read the bottom of the meniscus. Be sure your eye is at the level of meniscus, not above or below. The readings are recorded to ±0.02mL.

You will use a volumetric pipet to measure out the unknown acid. Be sure to always use two hands (one on the pipet bulb and your dominant hand on the pipet itself).

109

Procedure: The instructor will demonstrate the set-up and use of the buret to perform a titration.

Part A. Determination of acid concentration using titration.

1. Obtain approximately 50 mL of the HCl solution of unknown concentration in a small beaker.

Record the ID#.

2. Use a volumetric pipet to transfer exactly 10.00 mL of the HCl solution into a clean Erlenmeyer

flask. Add 2 drops of phenolphthalein indicator to the HCl solution.

3. Obtain approximately 100 mL of the sodium hydroxide solution. Record its concentration.

4. Clean and set-up a buret according to the directions given by your instructor. Place a funnel on the

top of the buret, and carefully pour 45-50 mL of NaOH into the buret. Remove the funnel.

5. Place a large waste beaker underneath the buret tip, and open the stopcock to let ~ 5 mL of NaOH

run down to fill the buret tip. Make sure air bubbles aren’t trapped around the stopcock.

6. Record the initial buret reading of the NaOH solution to the nearest 0.02 mL. See tips for reading

the buret correctly in the introduction.

7. Perform a “rough” titration to get an idea of what color change to expect at the endpoint and also to

get a rough idea of the volume of NaOH needed to get to the equivalence point by adding

approximately 2 mL aliquots of NaOH to the HCl solution. Swirl the flask to stir. Continue to add

NaOH in 2 mL increments until the phenolphthalein indicator remains pink for more than 20

seconds. Record the final buret reading of NaOH in the buret. Dispose of the contents of the flask in

the sink with lots of water.

8. Measure and transfer another 10.00 mL of HCl to a clean flask. Check to see if you need to refill the

buret. Record the initial buret reading of NaOH and titrate until you have added about 2 mL less

than the total amount added in the rough titration. Now add NaOH drop by drop until the HCl

solution just turns light pink and remains for at least 30 seconds while you swirl the flask.

(Proper technique here is essential. You want one drop of added base to cause the acid solution to

turn from colorless to light pink. If you add too much base and “overshoot” the equivalence point

you will need to start over.)

9. Repeat the titration until you have gotten 5 “good” titrations without overshooting the equivalence

point. (Ask your instructor to help you define “good”.)

10. Mix all acid and base solutions together and pour down the sink with lots of water.

11. Clean up.

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Name: __________________________________________Date due: ___________________________

Acid-Base Titrations Pre-lab Questions: Read the relevant textbook sections and the entire lab before answering pre-lab questions.

1. Define “equivalence point” in your own words, and describe how you will know when it has been reached in this lab.

2. Write the balanced chemical equation for the reaction of a sodium hydroxide solution with:

a. hydrochloric acid

b. sulfuric acid

3. What is the level of the liquid in the buret? Use the correct number of significant figures to reflect the precision of the instrument?

____________mL

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4. 17.65 mL of a 0.110 M sodium hydroxide solution is needed to titrate 25.00 mL of a hydrochloric

acid solution to the equivalence point.

a. Write the balanced equation for the neutralization reaction.

b. How many moles of sodium hydroxide are used for the titration?

c. How many moles of hydrochloric acid reacted with sodium hydroxide?

d. What is the molar concentration (Molarity) of the hydrochloric acid solution?

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Acid-Base Titrations

Data:

A. Titrations

Unknown acid ID # _____________

Concentration of NaOH _____________

Record your data to the correct significant figures.

Rough Trial 1 Trial 2 Trial 3 Trial 4 Trial 5

Initial buret reading (mL)

Final buret reading (mL)

Volume of NaOH added

(mL)

Calculated acid concentration

(M)

Calculations:

1. Calculate the concentration of acid for each trial. Show one sample calculation below and remember to use the correct number of significant figures and units.

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2. Calculate the average concentration of the acid. (Use your three closest values to calculate your average.)

Average acid molar concentration ________________mol/L

3. Calculate the % error. Your instructor will provide the known acid concentration. Known acid molar concentration __________________mol/L

Percent error = __________ % Conclusion questions:

1. The equivalence point of a titration is overshot! How will this error affect the calculated concentration of acid? (Will it be too high or too low?) Explain.

2. Describe at least one modification you would make to improve the accuracy of your experimental data.

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Spectroscopy: Determination of Concentration Using Beer’s Law

Objectives: In this lab students will

• Determine the concentration of copper(II) sulfate unknown Skills: Upon completion of this lab, students should have learned to:

• Prepare solutions of known concentration by dilution from a stock solution. • Use a spectrophotometer to measure the absorbance solutions • Draw a Beer’s Law graph (calibration graph) of absorbance versus concentration

Textbook References: (Tro, Introductory Chemistry, 3rd Ed.) to be read PRIOR to lab.

• Specifying Solution Concentration: Molarity Section 13.6 p. 457 • Solution Dilution Section 13.7 p. 461

Introduction: Some chemical solutions appear colored because certain wavelengths of light are absorbed by the solute. A red solution looks red because it absorbs light at wavelengths other than red and then only the red wavelengths pass through to your eye. You can easily guess that the darker the red color of the solution, the higher the concentration of light absorbing material in the solution. This idea of light absorbance being proportional to concentration is known as Beer’s Law, and written in equation form is

A = εbc

A = absorbance ε = is the absorption constant (different for each chemical) b = the path length the light travels through the solution c = the concentration

In this lab we are only interested in the relationship between absorbance and concentration (A and c), and ε and b are constants during our experiment so we can say that the Beer’s Law becomes

A = kc

(Where k is a constant . . . the product of ε x b = k). And now it is obvious that absorbance, A, is directly proportional to concentration, c. Because absorbance and concentration are directly proportional, a graph of absorbance versus concentration should be a straight line. The data for the graph is collected by measuring the absorbance of several solutions with known concentrations. This type of graph is known as a Beer’s law Plot (or calibration plot). We can use the Beer’s Law plot to determine the concentration of an unknown solution. One of the many practical uses of this process is to determine the amount of a contaminant (like nitrate ion, sulfate or lead ion, to name a few) in a drinking water sample.

Absorbance will be measured using a spectrophotometer.

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A detector within the spectrophotometer measures the relative % of light transmitted through the sample and converts the % transmittance to an absorbance value using the equation below.

A = 2 – log (% Transmittance)

Procedure: The instructor will demonstrate the use of the spectrophotometer.

Part A. Visible wavelengths Introduction (Optional; done only in class) 1. Insert an empty spectrophotometer tube containing a strip of white paper into the sample

holder; leave the sample holder cover open. Set the wavelength of the spectrophotometer to one of the wavelengths listed below and look down into the tube and record what color you see at each of these wavelengths - 660nm, 600 nm, 560 nm, 440, nm, and 400 nm.

Part B. Preparing the CuSO4 solutions of known concentration

2. Turn on the spectrophotometer. The machine takes about 10 minutes to warm up. 3. Obtain about 30 mL of 0.40 M CuSO4 stock solution in a small beaker. 4. Use a 10 mL graduated pipette and a 10 mL volumetric flask to prepare 10 mL of a 0.08M

solution according to the volumes you calculated in the prelab. Transfer this solution to a clean, dry test tube labeled #1.

5. Repeat for solutions of 0.16M, 0.24M, 0.32M and 0.40 M CuSO4, placed in test tubes # 2-5, respectively. Thoroughly mix each solution with a stirring rod. Clean and dry the stirring rod between stirrings.

Part C. Measuring the absorbance of the CuSO4 solutions

6. Set the wavelength of the spectrophotometer to 635 nm. 7. Prepare a “blank” by rinsing a cuvette (a special “test tube” for spectrophotometers) with

approximately 1 mL of deionized water. Repeat the rinse another time. And then cuvette about 2/3 full with deionized water. Use the “blank” to zero the spectrophotometer according to your instructor’s directions. Do not pour out the water in this cuvette, you will use this blank to zero the machine before measuring the absorbance of every sample.

8. Prepare “sample” for measurement by taking another cuvette and rinsing two times with approximately 1mL of the 0.08 M CuSO4 solution and then filling to approximately 2/3 full. Zero the machine using the blank and then place the sample cuvette in the machine and measure and record the absorbance. Repeat for the other four solutions of known concentration.

9. Obtain an unknown. Record its ID # and the measure and record its absorbance.

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Name: Date due: Spectroscopy: Determination of Concentration Using Beer’s Law Pre-lab Questions: Read the relevant textbook sections and the entire lab. Then complete the following: 1. Use the internet or library to find two practical uses of copper(II) sulfate. Reference the website. 2. Use the internet or library to find two specific health hazards associated with copper(II) sulfate and

outline the safe handling practices of copper(II) sulfate. Reference the website. 3. Use the dilution equation ( M1V1 = M2V2 ) to calculate the volume of 0.40 M copper(II) sulfate

solution needed to prepare 10.00 mL of these 4 solutions - 0.08M, 0.16M, 0.24M, 0.32M. Show your work.

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Name: Date lab performed: Partner(s) name: Date due: Spectroscopy: Determination of Concentration Using Beer’s Law Data Table: Record your observations and any test results.

A. Visible wavelengths Introduction (Optional)

Wavelength (nm) Observed Color 660

600

560

440

400

Insert an empty spectrophotometer tube containing a strip of white paper into the sample holder; leave the sample holder cover open. Set the wavelength of the spectrophotometer to one of the wavelengths listed below and look down into the tube and record what color you see at each of these wavelengths - 660nm, 600 nm, 560 nm, 440, nm, and 400 nm.

Solution Concentration (mol/L) Absorbance

1

2

3

4

5

Unknown ID # ______

XXXXXXXXXX

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Calculations: Make a properly labeled Beer’s Law plot of absorbance versus concentration (absorbance on the y-axis) for the five solutions of known concentration. Draw the best fit line through these five points. Use this graph to determine the concentration of the unknown. Record your result in the results section below. Use either graph paper or a computer program to make the graph. Attach the graph to the back of this lab report.

Results: Unknown ID # _______________ Concentration _______________ Conclusion Questions:

1. Use your graph to predict the absorbance of a 0.20 M CuSO4 solution.

2. Optional Bonus: Calculate the absorbance of a solution that has a % transmittance of 75%.

chem 100 lab manual harrisburg area community college...last updated 8/8/2013 . chem 100 lab manual....

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