general chemistry 1 and 2 lab packet - seattle central...

General Chemistry

1 and 2

Lab Packet

Summer 07

06/26/07

BASTYR UNIVERSITY

COURSE INFORMATION FOR STUDENTS

Summer Quarter 07

COURSE NUMBER BC 2112 and BC 2114 COURSE TITLE General Chemistry Laboratory I and II INSTRUCTOR Tess Cebrian

CLASS TIME Section A, Thursday , 1:30 – 6:50 Section B Friday, 1:300 PM to 6:50PM

CREDITS 1 each STUDENT ADVISING HOURS

1-1:30 PM Friday by appointment

PHONE (W)206 587-4075 off campus

(H) (E-mail) [email protected]

Website: www.chemsccc.org

Students are responsible for knowing and adhering to Academic Policies and Procedures as outlined in the Student Handbook.

*Listed are the major areas to cover. Please see Course Syllabus Instructions for more details on content*

1. Table of Contents Faculty Listing Requirements Course Overview Evaluation Course Objective Lab Report Lab Schedule 2. Course Overview • Course Description The course is designed as a practical application of the theories learned in lecture. Experiments include techniques of volumetric measurements and titration, stoichiometric application, reaction and qualitative analysis. The experiments will vary in degree of difficulty. As the quarter progresses the experiments will be more challenging but also more interesting and relevant to your program. It is important to come to lab prepared, on time and with a positive attitude. Remember the definition of “experiment." Although the experiments have been tried and are known to work, sometimes they do not. Failure of an experiment is just as important a learning tool as one that works, it allows you to examine the procedure more thoroughly to determine what went wrong. A degree of enthusiasm and the willingness to learn and work hard are the key to a successful lab experience.

• Major Course Competencies -Learn chemistry lab techniques.

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http://www.chemsccc.org/

-Develop data analysis skill. -Apply theories learned from lecture to interpret data gathered from the experiments. -Learn to write a complete, clear and concise lab report. • Organization & Requirements

1. Experiment/s will performed each week. Completion of all experiments is required. 2. Experiments will be performed in pairs unless otherwise instructed. 3. One lab can be made up on Friday Aug 31. 4. A lab notebook (composition type- not binders or spiral type) is required for the course. All

notes and record keeping during the experiment will be done in the lab notebook not in loose pieces of paper. Write in pen only. Record all data on the lab notebook. Cross-out unwanted data. Do not erase or block-out data. Transfer data to the printed data sheet in the lab book to be included in the lab report. There have been many data sheets misplaced or lost. Having the data recorded in the notebook reduces the probability of loosing important data.

5. Complete pre-lab questions prior to coming to lab. The lab lecture should give addition information for you to complete the prelab questions that you not able to answer. Prelab questions will be collected after lab lecture. You are expected to have read the experiment prior to coming to lab.

6. Lab reports and postlab questions are due the next lab period after the completion of the experiment. A 10-point deduction per week will be applied for late lab reports.

7. Compliance to the laboratory rules is essential for the safety of class. Instructional Materials and Resources

Lab Manual: Lab Notebook Experiment Hand-out (8) 3. GRADING • Evaluation Standards with Criteria for Passing and Remediation

Your lab grade will be based on the following Pre-lab questions 10pts each Data and post lab questions 100pts each and Laboratory performance Laboratory performance will be an evaluation of lab technique, lab preparedness, lab etiquette; which includes cleaning after oneself, keeping track of lab glassware, putting away chemicals, following safety instructions and effort given to achieve a successful experiment. Courtesy to lab instructor and classmates is expected. Undergraduate Grade Descriptions The following are general parameters for letter grades in the A-F system in graded undergraduate courses at Bastyr University: A = 90-100% This grade is given for excellence in completing course work whose development, presentation, scope and reasoning are considerably beyond the requirements outlined for a question, project or assignment. Outstanding work, extra effort, and excellent use of written or spoken language will contribute to this grade.

B = 80-89.9% This grade reflects good work that exceeds the minimal standards for a course. All key assignments must be accomplished and submitted on time and the work must be competently done and well presented. Deductions for late work may influence this grade. C = 70-79.9% This grade reflects average work which reflects minimum standards for success in a course in some way. Some aspects of coursework that would contribute to a “C” grade might include: weak development of a theme in a paper; inadequate

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consideration of an essay question on an exam; insufficient research or lack of evidence supporting a point or series of points in an assignment; or lack of organization and language use, including incorrect use of grammar. This grade may also reflect mixed work that includes some high quality aspects and some low quality aspects. Deductions for late work may influence this grade. D = 60-69.9% This grade reflects work that has not met the major requirements of the course. Grades awarded at this level often represent thinking that is disorganized and indicates only rudimentary understanding of basic concepts and principles, as well as poor organization and language use, including poor use of grammar. Deductions for late work may influence this grade. F = 59.9% or below This grade reflects serious problems with the work that prevents a higher grade. This may include work that is consistently unclear, imprecise and poorly reasoned. This may also include lack of even basic understanding of concepts and principles including the fundamental principles of language use and grammar. Deductions for late work and failure to submit required work may influence this grade. Conversion scale: Grade Point*

Transcript Grade Point*

A 4.0 95 - 100% Outstanding 3.8 - 4.0 A- 3.7 90 - 94.9 B+ 3.3 87 – 89.9% 3.2 - 3.4 B 3.0 83 – 86.9% Above Average 2.9 - 3.1 B- 2.7 80 – 82.9% 2.5 - 2.8 C+ 2.3 77 – 79.9% 2.2 – 2.4 C 2.0 73 – 76.9% Average 1.9 – 2.1 C- 1.7 70 – 72.9% 1.5 – 1.8 D+ 1.3 67 – 69.9% 1.2 – 1.4 D 0.7 63 – 66.9% Below Average 0.9 – 1.1 D- 1.0 60 – 62.9% Lowest Passing Grade 0.6 – 0.8 F 0.0 Below 60% Failing 0.0 – 0.

*This is the figure used to calculate GPA; see page 20 of catalog for additional grading information.

5. Course Outline & Time Schedule

DATE EXPERIMENT Pre-lab

Questions Post lab Questions

7/12 Sec A Check-in

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7/13 Sec B Read Safety Precaution Experiment 1: Measurements and observations

none All

7/19 Sec A 7/20 Sec B

Experiment 2: Identification of Compounds and their Chemical Properties. Physical observations elements and compounds.

All All

7/26 Sec A 7/27 Sec B

Experiment 3 Parts 1 and 2: Ionic and Covalent Bonds All Part A All Part A

8/2 Sec A 8/4 Sec B

Experiment 4: Stoichiometry and Limiting Reagent All All

8/9 Sec A 8/10 Sec B

Experiment 5: Solutions and Clock Reaction All All

8/16 Sec A 8/17 Sec B

Experiment 6: Types of Chemical Reaction All All

8/23 Sec A 8/24 Sec B

Experiment 7 : Titration of a Weak acid with a Strong Base

All All

8/31 Sec A and B

Experiment 8: Make up lab Lab report due Sat, Sept 1 @ 8:00 PM by email only.

All All

GLOBAL STUDENT COMPETENCIES

TO BE INCORPORATED INTO PROGRAM CURRICULUM (Please indicate which of these competencies are assessed in your class.)

1. COMMUNICATION SKILLS x Writing: Express self clearly, concisely and effectively for various purposes (political, teaching, scientific, clinical, and public affairs); adhere to grammar and syntax. x Listening: Listening without interrupting, accurate paraphrasing, clarification, and focus on speaker. Respond to verbal and nonverbal cues with congruence and empathy.

Speaking: Determine audience for appropriate language, content and delivery. Clearly articulate concepts and how they apply through organized thought (intro, body, ending).

Information Literacy Public Speaking

2. CRITICAL THINKING x Synthesis & Integration: Ability to gather and assess relevant information from many sources and divergent points of view. Ability to arrive at well-reasoned conclusions and solutions based on consideration of information from divergent points of view. Ability to apply solutions and test their effectiveness against relevant criteria and standards. Ability to generate new knowledge from assimilated knowledge. x Reflective Evaluation: Ability to understanding one's own assumptions and biases/point of view. Ability to understanding of the role of one's own inferences and

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interpretations. Ability to reconsider or reflect about one's own thinking and decision making processes. x Problem Solving: Ability to break the problem apart into its elements, analyze the problem, and estimate reasonableness of the proposed solution. Ability to find and execute a solution in order to achieve a goal using appropriate technologies and techniques. Ability to consider the ethical implications of the proposed solution. x Analytical Skills: Ability to make inferences based on understanding of many perspectives. Ability to recognize and analyze multiple perspectives, including quantitative and qualitative patterns. Ability to construct a claim and support it with logic and evidence. x Intuitive Skills

Research Skills: Research is the ability to conduct field or literature-based inquiry using available technology/techniques and producing a result in the discipline-appropriate form. Ability to understand, design and apply research strategies; evaluate sources of information in terms of relevancy, accuracy and bias; demonstrate knowledge of how information is obtained, analyzed and communicated in a discipline-appropriate manner; interpret and/or apply the results of the research strategy in an ethical manner.

3. PROFESSIONAL BEHAVIOR

Medical & Professional Ethics: Confidentiality and sharing of information, plagiarism and cheating, fairness and equality, and doing no harm.

Compassionate Caring Behaviors: Do no harm, active listening, honesty, and clear expectations (i.e. a syllabus). x Respectful Communication: Openness to new ideas and information, being proactive vs. reactive, respectful communication with/for students, faculty, and staff.

Personal Health & Wellness: In order to be present for patients (modeling), walk your talk.

Professional Boundary Skills: Knowing the limit of self and others. Students are not health care practitioners with a right to practice.

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LAB REPORT Lab report will be required for General Chemistry lab 2 only. For General Chemistry lab 1, a clean and legible data sheet and prelab and post lab question is all that is required. Lab reports will include a cover page and the data sheet provided in the lab book. Word process the cover page. Data sheet should be kept neat and writing legible. Your lab report will contain the following: 1. Cover Page

• Objective: There may be more than objective in an experiment. • Data: Refer to the data sheet. Use the data sheet provided in the lab book • Procedure: Refer to the page and experiment #. • Discussion and Conclusion: A brief discussion of what you observed in the

course of the experiment. Were the results what you expected? What may have led to the unexpected results. The conclusion should address the questions posed in the objectives.

2. Data Sheet

• Use the data sheet provided with the hand-out 3. Post lab questions

• Check the syllabus to find out the assigned questions.

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Sample Lab Report

Theresa Avalon Partner: Frankie Valle July 18, 2000

Experiment 3 Sink or Float?

Objective: To determine the density of various liquids and solids by measuring their volumes and masses. Data: Please refer to attached data sheet. Procedure: The instructor demonstrated the relative density of methylene chloride, water, hexane by adding the three liquids one at a time into the a graduated cylinder. Solid objects; glass marble, rubber stopper, cork and ice, were then added to the layered liquids. Where the solid objects either floated or sank was noted. A 10.00 mL volume water was weighed, same was done to an unknown liquid. The volume of a rubber stopper, an irregularly shaped object, was determine by volume displacement. The regularly shaped object, a wood block, was weighed. The height, length and width of the wood block was measured to determine its volume. The thickness of a foil was determined using the formula: Density= mass/( thickness x length x width) Thickness=mass/(density x thickness x width) The density of aluminum is 2.70 g/mL and the length and width of the aluminum was measured and its mass weighed on the top loading balance. Dicussion and Conclusion: The demonstration showed that methylene chloride was the most dense and hexane was the least dense. The glass marble sank to the bottom of the graduated cylinder, the rubber stopper floated in the methylene chloride layer, the ice floated in the water layer and the cork floated in the hexane layer. The calculated density of distilled water was 0.993 g/mL. The unknown liquid's density was 0.804 g/mL. The rubber stopper had a volume of 6.5 mL and a mass of 8.67 g. The density of the rubber stopper was 1.3 g/mL. The volume of the rectangular object was 1.55 cm3 and its density was 2.80 g/mL. The thickness of the aluminum was calculated to be 0.00158 cm.

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General Chemistry Lab

Experiment 1

Measurements

INTRODUCTION The metric system uses a basic set of units and prefixes. The basic unit of mass is the gram, the basic unit of length is the meter, and the basic unit of volume is the liter. Metric prefixes make these basic units larger or smaller by powers of 10. For example, a kilogram is a thousand times larger than a gram, and a milligram is a thousand times smaller than a gram. In the laboratory, the most common unit of mass, length, and volume are the gram (symbol g), centimeter (symbol cm), and milliliter (symbol mL) respectively. Scientific measurements have gradually progressed to a high state of sensitivity. However, it is still not possible to make an exact measurement. The reason for this is that all measurements utilize instruments that possess a degree of uncertainty—no matter how sensitive. The amount of uncertainty is shown by the significant digits in the measurement. For example, Ruler A below has major divisions of 1 cm. One can accurately measure to the 1 cm. However, if the item measured falls between the major divisions, as shown below, one estimates the number between the major divisions. The line below measures 8.5 cm. The 8-cm is an exact measurement, and the 0.5-cm is an estimate and is a significant. What would the length of the line be using Ruler B? The general rule is to estimate one more decimal place than the smallest division.

Figure 1 Centimeter and millimeter rulers

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MASS boratory the terms mass and weight are used interchangeably; however,

tially no

n

he basic unit of mass in the metric system is the gram (g). When we "weigh"

tes

und.

In the lathere is quite a difference in the two. Mass is a measure of the amount of matter in an object. Weight is a measure of the force of gravity on an object. Therefore, an object that weighs 50 pounds on earth would have essenweight in space. The mass of an object remains constant regardless of its location (an object has the same mass on earth or in space). The mass of aobject and the weight of an object are interchangeable as long as the force ofgravity remains constant. Tan object in the laboratory, we are determining "the mass of" the object in grams. The mass of an object is measured using a balance. A balance operaby comparing the mass of the object being weighed to the mass of a standard reference weight. Top-loading electronic balances are, by far, the most widelyused balances in chemical laboratories today (Fig. 2). Commonly, the top loading balances measure mass to the nearest centigram, that is, it weighs to100th of a gram. Another common usage is to weigh to the nearest 0.01. Thissimply instructs you to use a top loading balance that is capable of weighing tothe two significant numbers after the decimal. When using an electronic balance, you record the mass that balance is capable of giving. Do not ro

Figure 2 Electronic Top-loading Balance

EMPERATURE mon form of energy used in general chemistry

tures. of

se

elsius.

THeat is the most comlaboratories. Heat flows between objects that are at different temperaThus, temperature is a measure of the heat (energy) of an object or the flowheat between two objects. A thermometer is used to measure temperature in thelaboratory. A thermometer is a capillary tube that is generally filled with mercury or colored alcohol. These liquids are used in thermometers becauthey readily expand or contract with small changes in temperature. Thermometers in the laboratory are generally calibrated in degrees C

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Use utmost care when using mercury thermometers. Mercury vapors are

OLUME it of volume in the metric system is the liter (L). Graduated

es

n the

5-55 mL. s

st

.

e

. It

ASHING AND CLEANING OF GLASSWARE pieces of laboratory

EADING THE MENISCUS and graduated cylinders are calibrated to

ch a

ding ,

extremely toxic. Report mercury thermometer breakage to your instructors immediately. VThe basic unglassware used to quantitatively measure volume in the laboratory includgraduated cylinders, pipets, volumetric flasks and burets. Beakers and Erlenmeyer flasks are graduated to approximate volumes only. Errors igraduation can be > 10% and should not be used for quantitative measurements. For example, a 50-mL beaker may actually hold 4Graduated cylinders are commonly used to make more precise measurementof volume (0.5 -1.0% error) in the laboratory. The graduated cylinders that have measuring capability of 50 ml and above have divisions of 1.0 mL andmajor divisions of 5 or 10 mL. Estimates of volume can be made to the neare0.1 mL using a graduated cylinder. Volumetric flasks are not graduated and measure only fixed volume. Volumetric flasks come in different volume sizesThere are two different types of pipets, the volumetric pipet and the measuring pipet. The volumetric pipet, like the volumetric flask, only measures fixed volume. The manufacturer determines the uncertainty or accuracy of both thvolumetric pipet and flask. The measuring pipet is a graduated pipet and can be used for increment measurements. The buret is one of the more precise measuring glassware. It is graduated and is accurate to the nearest 0.01 mLis most commonly used for titration. WWashing and brushing with a detergent can clean mostglassware. After they have been thoroughly cleaned, they are rinsed with tapwater and finally with a spray of distilled water. If the surface is clean, the water will wet the surface uniformly. RVolumetric flasks, burets, pipets measure volumes of liquids. When a liquid is confined in a narrow tube suas a buret or a pipet, the surface is found to exhibit a marked curvature, called meniscus. It is common practice to use the bottom of the meniscus in calibrating and using volumetric ware. Special care must be used in reathis meniscus. By positioning a black-striped white card behind the meniscuswhich is transparent, it becomes more distinct.

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Figure 3 Meniscus

VOLUMETRIC MEASUREMENTS Location of the eyes in reading any graduated glassware is important.

1. With the eye above the meniscus, too small a volume is observed. 2. With the eye at the same level as the meniscus, the correct volume is

observed. 3. With the eye below the meniscus, too large a volume is observed.

The eye must be level with the meniscus of the liquid to eliminate parallax errors. Read the top of the black part of the card with respect to the graduations on the buret.

TOOLS OF VOLUMETRIC ANALYSIS Pipets, burets and volumetric flasks are standard volumetric equipment. Volumetric apparatus calibrated to contain a specified volume is designated TC, and apparatus calibrated to deliver a specified amount, TD. PIPETS Pipets are designed for the transfer of known volumes of liquid from one container to another. Pipets, which deliver a fixed volume, are called volumetric or transfer pipets. Other pipets, known as measuring pipets, are calibrated in convenient units so that any volume up to maximum capacity can be delivered. DIRECTIONS FOR THE USE OF A PIPET

** NEVER DRAW LIQUIDS INTO THE PIPET BY MOUTH, USE A PIPET PUMP **

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1. Clean pipet thoroughly with soap and rinse with distilled water. 2. Drain completely. Condition the pipet by rinsing three times with

the solution to be measured. 3. Keep the tip of the pipet below the surface of the liquid. 4. Draw the liquid up beyond the calibration mark. Lift the pipet

above the liquid and adjust to volume. 5. For volumetric pipets, transfer the pipet to the container to be used

and remove the pipet pump from the pipet or press the dispensing bar. Allow the solution to drain completely. Remove the last drop by touching the drop to the wall of the container. The calibrated amount of liquid has been transferred.

6. For measuring pipets, transfer the pipet to the container to be used. Using the pipet pump release the volume needed to transfer. If the measuring pipet is a blowout pipet, use the pipet pump to blow out the remaining drops. In case of color-coded measuring pipet, a frosted ring indicates complete blowout. NOTES: Pipets should be thoroughly rinsed with distilled water after each use.

BURETS

Burets, like measuring pipets, deliver any volume up to their maximum capacity. Burets are designed to measure the volume of solutions dispensed. The calibration marks start at 0.00 mL and end at 50.00 mL. Fifty mL burets are calibrated so that measurements can be carried to 2 significant numbers after the decimal. DIRECTIONS FOR THE USE OF A BURET Before being placed in service, a buret must be scrupulously cleaned. In addition, it must be established that the stopcock is liquid-tight. NOTE:

When using the buret, dispense the solution down from the 50.00 mL mark only. Above the 50.00 mL mark does not have any measurement.

FILLING THE BURET Test the buret for cleanliness by clamping it in an upright position and allowing it to drain. No water drops should adhere to the inner wall. If they do, clean the buret again. Make certain that the stopcock is closed. Condition the pipet with 5 to 10 mL of solution by carefully rotating the buret to wet the wall completely; allow the liquid to drain through the tip. Repeat this procedure two more times. Then fill

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the buret above the zero mark. Free the tip of air bubbles by rapidly draining the solution through while gently tapping on the buret. Finally, lower the level of the solution to or somewhat below the zero mark; after allowing about a minute for drainage, take an initial volume reading. After dispensing the necessary volume from the buret take the final volume reading. The final volume minus the initial volume is the amount of solution dispensed from the buret. Clean the buret with soap and water. Rinse with distilled water before storage. HOLDING THE STOPCOCK Always push the plug into the barrel while rotating the plug during a titration. A right-handed person points the handle of the stopcock to the right, operates the plug with the left hand and grasps the stopcock from the left side as shown.

Figure 5

Use of the Buret

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GRAVITY FILTRATION REPARING FILTER PAPER FOR A FILTER FUNNEL

Figure 6

Folding of a Filter Paper

If the solid is separated from the liquid through a filtering process, then the filter paper must be properly prepared. For a gravity filtration procedure, first fold the filter paper in half, again fold the filter paper to within about 10° of a 90° fold, tear off the corner unequally, and open. The tear enables a close seal to be made across the paper’s folded portion when placed in a funnel.

Place the folded filter paper snugly into the funnel. Moisten the filter paper with the solvent of the liquid/solid mixture being filtered (most likely this will be deionized water) and press the filter paper against the top wall of the funnel to form a seal. Support the funnel with a clamp or in a funnel rack.

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Figure 7

Gravity Filtration

TRANSFERING THE LIQUID

The tip of the funnel should touch the wall of the receiving beaker to reduce any splashing of the filtrate. Fill the bowl of the funnel until it is less than two-thirds full with the mixture. Always keep the funnel stem full with the filtrate; the weight of the filtrate creates a slight suction on the filter in the funnel, and this hastens the filtration process.

Flush a precipitate from a beaker with the mixture’s solvent (usually deionized water) contained in a wash bottle, while holding the beaker over the funnel or receiving vessel.

LABORATORY BURNER

NOTE: Before attempting to light a Bunsen burner, make sure that you are successful in generating sparks out of the striker. If you are not successful in getting the burner lit after two attempts, TURN OFF THE GAS FROM THE GAS JET.

Always check the rubber tubing for holes. Some heating in your chemistry course is done with a gas burner. In this laboratory you will use a burner of the Bunsen type. The burner has an air inlet just above the gas inlet, which can be adjusted by screwing or unscrewing the barrel of the burner. This adjustment determines the amount of air mixing with the gas. The larger the air opening is, the hotter the flame gets.

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The fuel used for the burner is natural gas. You will find a natural gas jet at each work area. Always be sure the gas jet is shut off completely when the burner is not lit.

Before lighting the burner, adjust the barrel of the burner so that you see an air opening. Turn the gas jet 90o. Light the burner with a striker. Adjust the air control to get a blue, nearly transparent flame.

Figure 8

Bunsen burner

• If the air inlet is closed and the gas is lit, the flame will be large and luminous. The

light is the radiation given off by the hot carbon particles that are burned only partially. This luminous flame is not very hot and dangerously flimsy. This very cool flame type will never be used in this lab.

• If the air control is adjusted so that air is mixed with the gas before it gets to the flame, the flame will become less luminous, and finally blue. When the air is adjusted correctly to give the hottest flame, it will look something as shown in the picture. The inner cone of the flame is pale blue, and the outer cone is pale violet. The inner cone contains the unburned gas that is hot enough to radiate light. The hottest point is just above the inner cone. SIGNIFICANT NUMBERS AND CALCULATIONS Rules:

1. Non-zero digits are always significant. 2. Zeroes:

a. Zeroes at the beginning of a number used just to position the decimal are never significant. 0.3560

b. Zeroes between non-zeroes are always significant. 20001 c. Zeroes at the end of a number that contains a decimal point are

always significant. 369.0 and 40.00 d. Zeroes at the end of a number without a decimal may or may not be

significant. 2000 (may be for significant number) or 200 x 101

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3. Exact numbers can be considered as having an unlimited number of significant figures. This applies to defined quantities. An example is conversion factors.

1 foot = 12 inches, both 1 and 12 are significant numbers. 4. A calculated number can never be more precise than the numbers used to calculate it.

a. Addition and Subtraction. The last digit retained in the sum or difference is determined by the position of the first doubtful digit.

358.986 595.71 1.4 956.096

b. Multiplication and Division. The answer should contain no more significant numbers than the least number of significant numbers in the operation.

Is the example below correct? (359.25) (40.2580) = 14462.6865

EXPERIMENT

The laboratory experiment below incorporates the techniques discussed above. For each step write an observation. Apply the significant measurements of each of the volumetric glassware when performing volumetric measurements. Follow the rules of significant number when asked to calculate.

1. Weigh 0.5 g of Potassium Hydrogen Phthalate (KHP) to the nearest

milligram (0.001). Record the exact mass in the data sheet.

2. Transfer the KHP to a 125 mL Erlenmeyer flask. 3. Using a 50 or 100 mL graduated cylinder, add 25 mL of distilled water to

the KHP. Swirl gently. 4. Adjust the ring approximately 9 inches from the bottom of a ring stand.

Place wire gauze on the ring. 5. Place a burner under the ring.

6. Place the Erlenmeyer flask containing the KHP solution on the wire gauze. Heat the KHP solution to 800C. Don't boil. Stir the solution with a

stirring rod to dissolve all the KHP. 7. Remove the Erlenmeyer flask and let the KHP solution cool close to

room temperature. While the KHP solution is cooling, fill the buret to just below the zero mark with the NaOH solution. Remember to fill the buret

tip and check to make sure it is free of air bubbles. 8. Record the initial buret volume on the data sheet. Remember your

significant numbers.

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9. To the cool KHP solution, add two drops of phenolphthalein solution.Place the Erlenmeyer flask under the buret and add NaOH

slowly to the KHP solution until the entire solution of KHP is pink. Swirl the solution as you add the NaOH.

10. Record the final buret volume on the data sheet. 11. Proceed with calculation as instructed on the data sheet using the

significant number rules.

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Name: _____________________________________________________

Class: _______________________________________________________

DATA SHEET

Mass of KHP ______________________g

Buret

Final Volume _______________________mL

Initial volume _______________________mL

Volume of

NaOH used ________________________mL (final volume - initial volume)

Calculation: This calculation will walk you through to determine the concentration (strength) of the NaOH solution. Calculate using the significant number rules.

1. Divide the mass of KHP used by 204.2 (gram/mole).

Answer: ________________moles KHP

2. Divide the answer above by the volume of NaOH used.

Answer: _________________M (moles/L)

Next week turn in this data sheet along with the post lab questions below. Also write a narrative incorporating your observations with your data.

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Name __________________________________________

Class _____________________________

POST LAB QUESTIONS 1. What is the physical quantity expressed by the following measurement: 10.0 cm? (a) length (b) mass (c) temperature (d) time (e) volume 2. What is the physical quantity expressed by the following measurement: 10.0 g? (a) length (b) mass (c) temperature (d) time (e) volume 3. What is the physical quantity expressed by the following measurement: 10.0 mL? (a) length (b) mass (c) temperature (d) time (e) volume 4. What is the physical quantity expressed by the following measurement: 10.0 s? (a) length (b) mass (c) temperature (d) time (e) volume 5. What is the physical quantity expressed by the following measurement: 10.0°C? (a) length (b) mass (c) temperature (d) time (e) volume 6. What is the length of the object shown on the metric ruler? (a) 0.75 cm (b) 0.8 cm (c) 0.80 cm (d) 7.5 cm (e) 7.50 cm

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7. What is the length of the object shown on the metric ruler? (a) 6 cm (b) 6.0 cm (c) 6.6 cm (d) 6.60 cm (e) 6.55 cm

8. What is the length of the object shown on the metric ruler? (a) 3 cm (b) 3.0 cm (c) 3.00 cm (d) 30 cm (e) 30.0 cm

9. What is the volume of liquid shown in the graduated cylinder?

10. What is the temperature shown by the thermometer?

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General Chemistry Lab Experiment 2

Reactivity of Metals

Introduction

About three-fourths of the 112 elements are metals. Of these, at least 20 are fairly common and are of major importance. To get acquainted with some physical properties of various metals peruse table 1. You will study some of the chemical properties of metals in today's experiment. Metals in their elemental form are commonly described as lustrous. They are good conductors of heat and electricity and are capable of being deformed without breaking. The fusing (melting together) of particular metals gives alloys which have a wide variety of properties and uses. Steels are alloys of iron with other metals and carbon. Stainless steels contain nickel and chromium. Brass contains mostly copper and zinc, whereas bronze is primarily copper and tin.

Metals and nonmetals combine in many ways to make compounds, some of which are soluble in water. In most compounds the metal exists in the form of positive ions and the nonmetal as negative ions. Generally those compounds of the metals whose negative ions are oxides, O-2, or hydroxides, OH- are called bases. The others are called salts. Examples, showing the charges on the ions, are sodium chloride, Na+ Cl-

, and copper sulfate, Cu2+ SO-2. Keep in mind the distinction between the free metal (the element) and the combined form of the metal (the compound); the former consists of neutral metal atoms and the latter contains positively charged metal ions (along with negative nonmetal ions). There is a vast difference between copper metal (red, lustrous, solid, insoluble in water, somewhat soft and flexible) and copper ions in copper sulfate (blue crystal, brittle, soluble in water). When you use copper sulfate solution in this experiment, note how copper (as ions) in solution is very unlike copper (as atoms) in metallic form. This experiment focuses on the properties of free metals. In nature metals usually are found in combination with oxygen or other nonmetals. Only a few of the less reactive ones, such as gold and sometimes silver and copper, are found in the free metallic (elemental) state.

One of the most important characteristics of a metal is its activity (reactivity, or ability to react to form compounds). Metals range widely in activity, from vigorously reactive cesium, potassium, and sodium, to quite inactive (inert) platinum, gold, and silver. Since the latter resist oxidation, they are often called the noble metals. The coinage metals include gold and silver along with some metals of lesser value, such as copper and nickel. Many metals become oxidized by reacting with oxygen in the air to form a tarnish or rust (oxide). Examples of metallic oxides are sodium oxide, Na2O, aluminum oxide, A12O3, ferric oxide, Fe2O3, and cupric oxide, CuO.

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TABLE 1 Some Metals and Their Properties

Metal Symbol Property Uses

Iron Fe Strong, tough, corrodes Main ingredient in steel Nickel Ni Resists corrosion Coinage, alloy for stainless steel Chromium Cr Shiny, resists corrosion Chrome plating Zinc Zn Forms protective coating of ZnO Galvanizing coating Radium Ra Radioactive Cancer treatment Aluminum Al Light and strong Airplanes, window frames Uranium U Fissionable Energy source Magnesium Mg Light and strong Auto wheels, luggage Mercury Hg Dense liquid at room temperature Thermometers, barometers Sodium Na Very active, soft Heat transfer medium Copper Cu Red color, good electrical conductor Electrical wiring Silver Ag Excellent electrical conductor Electrical contacts, mirrors, jewelry Gold Au Yellow metal, soft Coinage, instruments, dentures Platinum Ft Inert, high melting Jewelry, instruments Tantalum Ta Resists attack by acids Synthetic skulls and bone parts Tungsten W \7ery high melting Light-bulb filament Tin Sn Resists corrosion Coating for steel cans Lead Pb Low melting, dense, soft Plumbing, fishing weights

Replacement Series

In this experiment you will be ranking some of the metals according to their activities, from most to least active. One way to do this is to observe the relative vigor of their reactions (tendency to react) with nonmetals such as oxygen or chlorine. Another method, which differentiates more clearly, is to note whether one metal can replace another in a chemical compound. The general rule is that the more active metal replaces the less active one. This is easily observed in reactions between metals and metal ions (cations) in solution. Iron, for example, replaces copper from cupric chloride, CuO2, because it is more active than copper.

Fe + CuCl2 FeCl2 + Cu

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In this example two electrons are transferred from the neutral iron atom to the positive copper ion. This is more apparent if the equation is written in net ionic form.

Fe + Cu+2 Fe+2 + Cu

This reaction, which involves the transfer of electrons, can be divided into two half-reactions which show exactly where electrons are gained and lost.

Fe Fe+2 + 2e- and Cu2+ + 2e- Cu

The general reaction for the replacement of a metal ion by another metal may be written as follows. (We are arbitrarily showing only one electron per atom being transferred.)

MA + MB+ X- MA+ X- + MB

Free metal A Salt of metal B Salt of metal A Free metal B

Or, in net ionic form, MA + MB

+ MA+ + MB

Replacement of Hydrogen from Acids and Water

Acids are an important group of compounds, which, in water, produce hydronium ions, H3O+. Hydronium ions enter into replacement reactions with metals; so hydrogen appropriately is included in an activity or replacement ranking among metals. Those metals more active than hydrogen will replace hydrogen from acids, but those less active are unable to do so.

Zn + 2HC1 ZnCI, +H2(g)

Or, in ionic style, Zn + 2 H+ Zn2+ + H2(g)

You may consider water to be a special case of a very weak acid. Metals which replace hydrogen from acids also liberate hydrogen from water, but drastic conditions may be required. Only the most active metals, such as sodium and potassium, replace hydrogen from water readily at room temperature. The reaction of sodium with water is one you will observe today This vigorous reaction can be sufficiently exothermic to ignite the explosive hydrogen gas, so you will use only a small amount of sodium. Some reactions of metals with water are

2Na + 2H2O NaOH + H2 @ 25oC

3 Fe + 4 H2O Fe3O4 + 8H2 @ 800oC

Metal Activities and the Periodic Chart

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The activity of a metal is related to the ease with which its outer electrons can be removed. Generally, atoms having the fewest electrons in the outer shell and those with the largest radii lose their electrons most easily. In the periodic chart, elements with the fewest outer electrons are located toward the left side, and those with the largest radii are found toward the bottom. In general, the most active metals are located toward the lower left corner of the chart.

In this experiment you will become familiar with the properties of various metals, and especially with the differences in their reactivities, learn to relate the activity of a metal to its position in the periodic chart and understand clearly the distinction between free metals and metals in the combined state.

Procedure 1. Write your observations on the physical appearance of the metals listed in Table 1 of the data sheet. 2. Reaction of Metals with Oxygen

In this part of the experiment you will make a variety of observations on the reactivity of oxygen in the air with several metals. Record your observations in the table and answer the questions.

A. Reaction at Room Temperature

DEMONSTRATION: Your instructor will place small strips of metals on a sheet of paper and label them. Each metal will be scratched to expose a clean surface to the air. Suggested metals are iron, tin, lead, copper, magnesium, aluminum, and zinc. Observe the metals throughout the laboratory period. Do they react with oxygen, that is, do they become tarnished?

Your instructor will cut a piece of sodium. Notice how easily sodium can be cut with a knife. Observe a freshly cut piece and see what happens when air comes into contact with the metal. Does a reaction occur between sodium and oxygen in the air?

B. Reaction at Elevated Temperature

Most of your metal samples, which you placed on the paper towel, probably have not reacted noticeably with air. Reactions that are sluggish at room temperature often can be accelerated by raising the temperature. Collect another set of metal strips used in #1. Heat the samples by cautiously holding them in the flame with tongs. Do any of the metals burn (react) with oxygen? Could you suggest an ingredient for making flares?

2. Reaction of Metals with Water

To determine whether the metals of part 1 react with water, place a piece of each, the size of a match head, in 1 mL of water in a test tube. Remember

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that water acts as an extremely weak acid and that the evidence for a reaction is the evolution of hydrogen gas (bubble formation). DEMONSTRATION: Your instructor will cautiously add a small piece of sodium to water in a beaker. Record in the table your observations for sodium and the other metals and answer the questions.

3. Reaction of Metals with an Acid

Safety Caution: HCl is corrosive. Wash off the acid immediately is you get it on your skin and report to your instructor any major spills.

Now repeat the reaction of part 2 with three different concentrations of a strong acid, such as hydrochloric acid, HC1. The concentrations are 1 M, 3 M and 6 M; 3M is three times more concentrated than 1 M and 6 M is twice as concentrated as 3 M. Place a small piece of each metal which did not react with water in 1 mL of dilute (1 M) HC1 in test tubes. Observe the rates of reaction (evolution of hydrogen gas in form of bubbles). Those metals which cause the evolution of hydrogen gas can be presumed to be more active than hydrogen. Slow reactions often can be speeded by increasing the concentration of a reactant. If any of the metals reacts extremely slowly with the 1 M acid, or fails to react, test its reactivity in 3 M hydrochloric acid, HC1. Finally, try any metal in 6 M HC1 if it fails to react in 3 M HC1. Record in the table the data on the reactivity of each metal. Indicate if 3 M or 6 M HC1 was required and what you observed. Answer the questions following the table.

4. Reactions of Metals with Other Metal Ions

Compare the reactivity of iron, copper, and silver in the following ways, in separate test tubes place 5 mL of each of the specified salt solutions containing the metal cation, and add a piece of the indicated metal (a strip or a bright nail). Allow at least 15 minutes for a reaction to occur. The more active metal will replace the less active one from its solution. If a replacement reaction occurs you will easily recognize it by the deposition of the new metal on the nail or strip. The metals and salt solutions are: a. iron (nail) in 0.1 M CuSO4 solution b. copper (strip) in 0.1 M FeSO4 solution. This solution must be freshly made by the student. Weigh 0.13 grams of FeSO4 and dissolve in 10 mL of water measure with a 10 mL graduated cylinder c. copper (strip) in 0.1 M AgNO3 solution.

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NAME:________________________________________________________ DATA SHEET

1-3. Reactions of Metals with Oxygen, Water, and Acid Before performing the reactions below write you observation on the appearance of each of the metal below.

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In the table below record your observations on the reactivity of the metals used in the experiment. Use qualitative terms such as fast(f), slow(s), vigorous(v), sluggish(sl), or no reaction(nr) to describe what happened. Make a separate record for each concentration of acid.

Table 1

Metal With Air Heated in Air

RXN WITH Water

RXN WITH 1M HCl

RXN WITH 3M HCl

RXN WITH 6M HCl

Iron

tin

lead

copper

magnesium

aluminum

zinc

sodium Write your observations below for each of the metal when reacted. Iron Tin Lead Copper

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Magnesium Aluminum Zinc Sodium Post lab Questions 1a Did you observe any tarnishing on the surface of the freshly cut sodium? If so, what is the name of the corrosion product? Was sodium metal a reactant?

1b. What reacted with sodium, Na to form the tarnish Na2O?

Attempt to write an equation for the reaction.

1c. Why is sodium normally stored under a liquid such as toluene or

mineral oil?

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2. Write equations for the reactions that occurred when you heated zinc and aluminum in the presence of oxygen. Did the zinc and/or aluminum react in the presence of oxygen, O2.

Write equations for the reactions that occurred.

3. What is the formula of the compound that was formed when magnesium was heated in air? 4a. Make a sketch showing valence electrons for an atom of the free magnesium metal and for the magnesium ion in the resulting compound. 4b Did the atoms of magnesium lose electrons when the sample was heated?

4c. Did metals react with oxygen more rapidly at flame temperature than at room temperature? Can you explain your observation?

5. Would any of the burning metals make a good night flare?

6. Write an equation showing the oxidation (rusting) of iron to ferric oxide.

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7. Write equations for the reactions (if any) that occurred between hydrochloric acid and aluminum, and hydrochloric acid and copper.

Of the metals you have examined in this experiment, which would be the most suitable for use in jewelry? Which would be the least suitable?

Zinc is more active than copper. On the basis of this statement only, which of the following statements are true?

_____ Zinc has a greater tendency than copper to exist in the combined form (as a cation).

_____ It is easier to get copper metal from copper sulfate than to get zinc metal from zinc sulfate

_____ Copper cyanide is more easily decomposed than zinc cyanide. Zinc tarnishes faster than copper.

____ Zinc replaces hydrogen from acids but copper does not.

4. Reactions of Metals with Other Metal Ions What, if anything, did you observe in each of the following trials? Write an. equation for the reaction if one occurred.

Iron with CuSO4 solution:

Copper with FeSO4 solution:

Copper with AgNO3 solution:

When a metal atom becomes an ion, or an ion becomes an atom, there is a transfer of electrons. In the preceding reactions, draw circles around the metal atoms which lost electrons and boxes around the metal ions which gained them (reactants). Which is the more active: (a) iron or copper, (b) iron or silver, (c) silver or copper? Arrange the three metals in order of decreasing activity.

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Arrange all of the metals tested in all parts of this experiment, including hydrogen, in order of decreasing activity.

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General Chemistry Lab Experiment 3 Part 1

Ionic Bonds

INTRODUCTION Ionic compounds are collections of ions held together by the force of electrostatic attraction between its positive and negative components, the cation and anion respectively. When dissolved in water, the ionic compound exists as charged positive and negative ions. Qualitative analysis is a systematic procedure for the separation and identification of ions and compounds present in a sample. Cation analysis involves the separation and identification of a positively charged cation present in a sample. If you have a solution containing several cations, it is usually possible to select a reagent that will form a precipitate with one of the cations but not with the others. If you place the test tube and precipitate into a centrifuge, the solid particles in solution are forced to the bottom of the test tube. Thus, the precipitated cations are separated from the remaining cations in solution. The solution is then separated and the analysis continues for the remaining cations. For example, a solution containing Barium, Ba2+ , Calcium, Ca2+, and Magnesium, Mg2+, ions can be separated using ammonium sulfate. The sulfate ion, SO4

-2, will precipitate Ba2+ but does not react with either Ca2+ or Mg2+ ions (see Figure 1). There is no reaction between Ca2+ and SO4

2-, or Mg2+ and SO42-, because CaSO4 and MgSO4 are soluble.

Figure 1

Precipitation of Ba2+ as BaSO4

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When a barium cation and a sulfate ion are together in the same solution, a precipitate forms because barium sulfate, BaSO4, is insoluble. The reaction is shown in the equations below. Molecular Equation: The (aq), aqueous, indicates an ionic solution.

BaCl2 (aq) + Na2SO4 (aq) BaSO4 (solid) + 2NaCl(aq) Ionic Equation: The ions are shown as they exist in the reaction. The charged ions are in solution.

Ba2+ + 2Cl- + 2Na+ + SO42- BaSO4 (solid) + 2Na+ + 2Cl-

Net Ionic Equation: The spectator ions are cancelled out.

Ba2+ + SO42- BaSO4 (solid)

If calcium or magnesium ions are in the solution, no precipitate forms because calcium sulfate, CaSO4, and magnesium sulfate, MgSO4, are soluble in water. Ca2+ is separated from Mg2+ by precipitating the cation with ammonium oxalate, (NH4)2C2O4 to form CaC2O4 as shown in the net ionic equation below. Ca2+ + C204

-2 CaC2O4 (solid) The remaining Mg+2 is precipitated with NaOH. 3Mg2+ + 2HPO4

2- + 2OH- Mg3(PO4) (solid) + 2H20 In this experiment you will separate and identify Ba2+, Ca2+, and Mg2+. First, a known solution containing all three cations will be analyzed to develop the necessary techniques. Second, an unknown solution with one or more of the three cations will be analyzed to determine the cations present. Flame testing is a technique you will use to confirm the presence of an ion. A flame-test is performed by dipping a wire into a solution and then holding the wire in a hot flame while observing the color produced (Figure 2). Many elements produce colored flames. For example, sodium is yellow, potassium is violet, and copper is green. Since sodium is always present as an impurity, the yellow sodium flame invariably contaminates flame-tests.

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Figure 2

Flame-test Technique A drop of solution is placed on the tip of a wire, which is held above a hot flame. The color of the flame is characteristic for the presence of a given element. This will not be done in today's experiment, but a demonstration will be performed on pure solutions of barium and calcium. Litmus paper can be used to determine whether a solution is acidic or basic. A stirring rod is placed into the solution and touched to the litmus paper. Acidic solutions turn blue litmus paper red. Basic solutions turn red litmus paper blue (Figure 3).

Figure 3

Litmus-paper Technique If the solution is neutral, there is no change to red or blue litmus papers. The qualitative cation analysis begins with a known solution containing Ba2+, Ca2+, and Mg2+. Figure 4 presents an overview of the analysis.

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Figure 4 Cation Analysis

The systematic separation and identification of Ba2+, Ca2+, and Mg2+ cations.

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PROCEDURE General Directions • Clean three test tubes and a stirring rod with distilled water. • Label the test tubes #1, #2, and #3. As a solution is analyzed, record the color of each

precipitate in the Data Table. • Use a plastic or Pasteur pipet to transfer supernates from one test tube to another.

A. Analysis of a Known Solution 1. Separation and Identification of Ba2+ in a Known Solution • Place 10 drops of the known solution into test tube #1. Add 10 drops of ammonium

sulfate, (NH4)2SO4, and mix with a glass-stirring rod. Note: A white precipitate suggests Ba2+ is present.

• Centrifuge and then test for completeness of precipitation by adding another drop of ammonium sulfate. If the solution precipitated after adding the drop of ammonium sulfate, add another 5 drops, mix, and recentrifuge. If no precipitate is observed, all the Ba+2 is completely precipitated. Proceed to the next step.

• Pipet only the supernate into test tube #2 and save for Step 2. 2. Separation and Identification of Ca2+ in a Known Solution • Add 10 drops of ammonium oxalate, (NH4)2C2O4, to the solution in test tube #2.

Note: A white precipitate suggests Ca2+ is present. • Centrifuge and then test for completeness of precipitation by adding another drop of

ammonium oxalate. If precipitation occurs add another 5 drops and recentrifuge. • Pipet the supernate into test tube #3 and save for Step 3. 3. Identification of Mg2+ in a Known Solution • Add 10 drops of sodium monohydrogen phosphate, Na2HPO4, to the solution in test

tube #3. • Add 1 drop of sodium hydroxide, NaOH, and stir with a glass rod.

Note: A white precipitate suggests Mg2+ is present.

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B. Analysis of an Unknown Solution

1. Separation and Identification of Ba2+ in an Unknown Solution • Place 10 drops of unknown solution into test tube #1. Add 10 drops of ammonium

sulfate, (NH4)2SO4, and mix with a glass stirring rod. Note: If there is no precipitate, Ba2+ is absent. Go directly to Step 2.

• Centrifuge and then test for completeness of precipitation by adding another drop of ammonium sulfate. Pipet the supernate into test tube #2 and save for Step 2.

2. Separation and Identification of Ca2+ in an Unknown Solution • Add 10 drops of ammonium oxalate, (NH4)2C2O4, to the solution in test tube #2.

Note: If there is no precipitate, Ca2+ is absent. Go directly to Step 3.

• If a precipitate forms, centrifuge and then test for completeness of precipitation by adding another drop of ammonium oxalate. Pipet the supernate into test tube #3 and save for Step 3.

3. Identification of Mg2+ in an Unknown Solution • Add 10 drops of sodium monohydrogen phosphate, Na2HPO4, to the solution in test

tube #3. • Add 1 drop of sodium hydroxide, NaOH, and stir with a glass rod.

Note: If there is no precipitate, Mg2+ is absent. 4. Based upon the observations in steps 1-3, identify the cation(s) present in the unknown solution.

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Name _________________________________________ Class ____________________________ DATA Part A Known Analysis

Test Tube #1

Mixture of Ba+2, Ca+2, Mg+2

(NH4)2SO4 Centrifuge

Supernate Precipitate: Yes No Test Tube #2 Color ________ Cation Present? ______ (NH4)C2O4 Centrifuge Supernate Precipitate: Yes No Test Tube #3 Color _______ Cation Present? ________ Na2HPO4 NaOH Centrifuge Precipitate Yes No Color _______ Cation Present?

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Part B Unknown Analysis

Test Tube #1 Mixture of Ba+2, Ca+2, Mg+2

(NH4)2SO4 Centrifuge

Supernate Precipitate: Yes No Test Tube #2 Color ________ Cation Present? ______ (NH4)C2O4 Centrifuge Supernate Precipitate: Yes No Test Tube #3 Color _______ Cation Present? ________ Na2HPO4 NaOH Centrifuge Precipitate Yes No Color _______ Cation Present?

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Name __________________________________ Class ___________________________ POSTLABORATORY ASSIGNMENT 1. An unknown solution of the barium, calcium, and magnesium group did not give a

precipitate with (NH4)2SO4 solution. Upon addition of (NH4)2C2O4 solution to the unknown, a white precipitate formed. To the supernatant liquid was added a solution of Na2HPO4 and NaOH, but no precipitate resulted. State the ion(s) present in the unknown solution.

2. Write the formula for each of the following cations.

(a) aluminum _________ (b) stannic _________ (c) calcium _________ (d) copper (II) _________ (e) hydrogen _________ (f) ferric _________ (g) magnesium_________ (h) mercury (I) _________ (i) potassium _________ (j) zinc _________

3. Complete the table below as shown by the example. Combine the ions into a correct formula and name the compound.

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Name: _______________________________________________________ Class: ________________________________________________ PRELABORATORY ASSIGNMENT 1. In your own words, define the following terms: cation centrifuge precipitate (ppt) supernate 2. Why is it necessary to use distilled water throughout this experiment? 3. How do you test for the completeness of precipitation? 4. An unknown cation solution is analyzed for barium, calcium, and magnesium ions.

• The unknown solution plus (NH4)2SO4 gives a white precipitate.• The supernate is poured into test tube #2. • Test tube #2 plus (NH4)2C2O4 remained a solution. • The supernate is transferred into test tube #3. • Test tube #3 plus Na2HPO4 and NaOH gives a white precipitate.

Refer to Figure 4 and determine which of the following cations are present in the unknown solution: Ba2+, Ca2+, and Mg2+.

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General Chemistry Lab Experiment 3Part 2

Covalent Bonds

INTRODUCTION The attraction between two atoms in a molecule is called a chemical bond. In a covalent bond, two nonmetal atoms are attracted to each other by sharing valence electrons. The valence electrons are the electrons furthest from the nucleus and occupy the highest s and p sublevels. The number of valence electrons is found from the periodic table. The group number of an element indicates the number of valence electrons. For example, fluorine is in Group VIIA/17 and has seven valence electrons (7 e-). Example Exercise 1 Refer to the periodic table and find the number of valence electrons for the following elements: (a) H; (b) C; (c) N; (d) O; (e) Cl, Br, I. Solution: (a) The element hydrogen is in Group IA/1. Since the group number is 1, hydrogen has

one valence electron, (b) Carbon is in Group IVA/14; thus, carbon has four valence electrons. (c) Nitrogen is in Group VA/15; thus, nitrogen has five valence electrons. However,

under ordinary conditions only three of nitrogen's valence electrons are shared. The remaining two electrons do not usually bond and are referred to as nonbonding electrons.

(d) Oxygen is in Group VIA/16; thus, oxygen has six valence electrons. (e) Chlorine, bromine, and iodine are in Group VIIA/17; thus, each of the halogens has

seven valence electrons. In this experiment we will write the structural formula and electron dot formula for molecules after building a model. A model is constructed from spherical balls and connectors where each ball represents an atom and each connector a single bond. Since a single bond shares two electrons, each connector represents an electron pair. A double bond shares two pairs of electrons. A molecular model is constructed using two connectors to represent the double bond. A triple bond shares three pairs of electrons. A molecular model is constructed using three connectors to represent the triple bond. The following example exercises illustrate the structural formula and electron dot formula for molecular models having single, double, and triple bonds.

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Example Exercise 2 The model of a water molecule is sketched below. Draw (a) the structural formula, (b) the electron dot formula corresponding to the model and (c) verify the electron dot formula by checking the total number of electron dots against the sum of all valence electrons.

Solution: (a) Each connector represents a single bond; the structural formula is

(b) A dash in the structural formula indicates an electron pair, thus

Each hydrogen atom shares a maximum of two electrons. However, each oxygen requires an octet of electrons and in the above diagram shares only four. Therefore, we must add two more pairs of electrons to oxygen in order to complete the octet. The electron dot formula is

.. H:O: ..

H (c) To verify the above formula we will add up the valence electrons from each atom in the molecule. Recall that hydrogen is in Group IA/1 and oxygen is in Group VIA/6.

2 H(2 x 1 e-) = 2 e- l O(l x 6 e-) = 6 e-

sum of valence electrons = 8 e- There are eight dots used to write the electron dot formula. Since this equals the number of valence electrons, the electron dot formula is correct. Example Exercise 3 The three-dimensional model of chloroform is sketched below. Draw (a) the structural formula and (b) the electron dot formula. Each atom (excluding H) should be surrounded by an octet of electrons. (c) Verify the electron dot formula by checking the total number

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of electron dots against the sum of all valence electrons.

Hydrogen shares two electrons and is complete. Carbon shares a total of eight electrons and satisfies the octet rule. However, each chlorine also requires an octet, which we will complete as follows:

(c) To verify the above electron dot formula we will find the sum of all valence electrons.

1H (1 x 1 e-) = 1 e-

1C (1 x 4 e-) = 4 e-

3Cl (3 x 7e-) = 21 e-

sum of valence electrons = 26 e-

Example Exercise 4 A molecular model of formaldehyde is sketched below. Draw the (a) structural formula, (b) electron dot formula and (c) find the sum of all valence electrons to verify the electron dot formula.

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Solution: (a) Two connectors joining the carbon and oxygen atoms represent a double bond. The

structural formula can be shown as

(b) Each single bond contains one electron pair and the double bond two electron pairs.

Hydrogen shares two electrons and is complete. Carbon shares a total of eight electrons and satisfies the octet rule. Oxygen has only four of the eight electrons necessary to complete the octet. Therefore, we will add two unshared electron pairs.

(c) We can verify the above electron dot formula as follows:

The 12 valence electrons equal the 12 electron dots and verify the formula. Example Exercise 5 A molecular model of hydrogen cyanide is sketched below. Write (a) the structural formula, (b) the electron dot formula and (c) verify the electron dot formula.

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Solution: (a) The three connectors linking the carbon and nitrogen represent a triple pair of

electrons.

H-C≡N (b) We can write an electron dot formula after realizing the triple bond contains three

electron pairs.

H:C:::N In the above formula nitrogen shares only six electrons. Therefore, we must add one unshared electron pair.

H:C:::N:

(c) Let's verify the preceding electron dot formula.

1 H (1 x 1 e-) = 1 e- 1 C (1 x 4 e-) = 4 e- 1 N (1 x 5 e-) = 5e-

sum of valence electrons = 10e-

The 10 valence electrons verify the 10 electron dots. Directions for Using Molecular Models • When constructing a model, a hole in a ball represents a missing electron that is

necessary to complete an octet. • If two balls are joined by one connector, the connector represents a single bond

composed of one electron pair. • If two balls are joined by two connectors, the two connectors represent a double bond

composed of two electron pairs. • If two balls are joined by three connectors, the three connectors represent a triple

bond composed of three electron pairs.

one rigid connector — single bond (one electron pair) two flexible connectors — double bond (two electron pairs) three flexible connectors — triple bond (three electron pairs)

A molecular model uses different color balls to represent hydrogen, carbon, oxygen, chlorine, bromine, iodine, and nitrogen atoms. The color code for each ball is as follows:

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white ball — hydrogen (one hole) black ball — carbon (four holes) red ball — oxygen (two holes) green ball — chlorine (one hole) orange ball — bromine (one hole) purple ball — iodine (one hole) blue ball — nitrogen (three holes)

Note: If the blue nitrogen ball has more than three holes, use a small peg or tape to fill the additional hole(s). All the holes in each ball must have a connector for a model to be built correctly. PROCEDURE 1. Construct models for each of the molecules listed on page. Sketch the molecular model in the Data Table showing its three-dimensional structure. 2. Draw the structural formula corresponding to the molecular model. 3. Draw the electron dot formula corresponding to the structural formula. Complete the octet by surrounding each atom with 8 electrons (2 electrons for a hydrogen atom). 4. Verify each electron dot formula by summing the valence electrons for the molecule using the periodic table. This sum should equal the total number of dots in the electron dot formula.

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Name________________________________________________ Class____________________________ DATA TABLE

Molecular Models with Single Bonds

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Name ______________________________________ Class_______________________________________ PRELABORATORY ASSIGNMENT 1. Refer to the periodic table in order to predict the number of valence electrons for each of the following elements: H, C, O, Cl, and N. 2. Draw the structural formula and electron dot formula for each of the following.

3. Perform a valence electron check on each of the examples in the preceding question.

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Limiting Reactant

INTRODUCTION Two factors affect the yield of products in a chemical reaction: (1) the amounts of starting materials (reactants) and (2) the percent yield of the reaction. Many experimental conditions, for example, temperature and pressure, can be adjusted to increase yield of a desired product in a chemical reaction. The reactant determining the amount of product generated in a chemical reaction is called the limiting reactant. Sometimes only a limited amount of the one of the reactants needed for the reaction is available, or perhaps it is easier to carry out a reaction by adding an excess of one of the reactants. The maximum amount of product that can be formed is determined by the amount of reactant that is used up first. Chemicals react according to fixed mole ratios (stoichiometrically), so only a limited amount of product can form from given amounts of starting materials. To better understand the concept of the limiting reactant, let us look at the situation where a toy store is trying to assemble bicycles. Suppose that each bicycle requires six nuts, six bolts, and six washers. If the shipment of parts includes 60 nuts, 60 washers and only 59 bolts, then the bolts become the limiting factor and only 59 bicycles can be assembled. In his experiment, the reaction of sodium phosphate dodecahydrate, Na3PO4·12H2O, and barium chloride dihydrate, BaCl2·2H2O, in an aqueous system produces solid barium phosphate, Ba3(PO4)2. The molecular form of the equation for the reaction in aqueous so-lution is

2 Na3PO4·12H2O(aq) + 3 BaCl2·2H2O(aq) → Ba3(PO4)2(s) + 6 NaCl(aq) + 30 H2O(l)

The two reactant salts and sodium chloride are soluble in water but barium phosphate is insoluble. The ionic equation for the reaction is 6 Na+ + 2 PO4

3- + 24 H2O(l) + 3 Ba2+ + 6 Cl- + 6 H2O(l) → Ba3(PO4)2(s) + 6 Na+ + 6 Cl- + 30 H2O(l)

Presenting only the ions that show evidence of a chemical reaction occurring, i.e., the formation of a precipitate (and removing spectator ions from the equation), the net ionic equation for the observed reaction is

2 PO43- + 3 Ba2+ → Ba3(PO4)2(s)

Spectator ions, cations or anions, do not participate in any observable or detectable chemical reaction.

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Net ionic equation is an equation that includes only those ions that participate in the observed chemical reaction. From the balanced net ionic equation, 2 moles of phosphate ion (from the 2 mol of Na3PO4·12H2O, molar mass = 380.12 g/mol, or 760.24 g) react with 3 moles of barium ion (from 3 mol of BaCl2·2H2O, molar mass = 244.27 g/mol, or 732.81 g), if the reaction proceeds to completion. The equation also predicts the formation of 1 mole of Ba3(PO4)2 (molar mass = 601.93 g/mol), or 601.93 g. In this experiment the solid salts Na3PO4·12 H2O and BaCl2·2 H2O form a heterogeneous mixture of unknown composition. The mass of the solid mixture is measured and added to water, then insoluble Ba3(PO4)2 forms. The Ba3(PO4)2 precipitate is collected, via gravity filtration and dried, and its mass is measured.

The percent composition of the salt mixture is determined by first testing for the limiting reactant. The limiting reactant for the formation of solid barium phosphate is determined from two precipitation tests of the solution: (1) the solution is tested for an excess of barium ion with a phosphate reagent. The formation of a precipitate indicates the presence of an excess of barium ion (and a limited amount of phosphate ion) in the salt mixture. (2) The solution is also tested for an excess of phosphate ion with a barium reagent. The formation of a precipitate indicates the presence of an excess of phosphate ion (and a limited amount of barium ion) in the salt mixture. CALCULATIONS The calculations required for the analysis of the data in this experiment are involved. The question, after collection of all of the data, becomes, "How do I proceed to determine the percent composition of a salt mixture containing the salts Na3PO4·12H2O and BaCl2·2H2O from the data of a precipitation reaction?" Consider the following example: A 0.942 g sample of the salt mixture is added to water and 0.188 g of Ba3(PO4)2, precipitate forms. Tests reveal that BaCl2·2H2O is the limiting reactant. What is the percent composition of the salt mixture?

• BaCl2·2H2O is the limiting reactant • The stoichiometry of the reaction indicates that 1 mole Ba3(PO4)2 precipitate

requires 3 moles Ba2+ and therefore 3 moles BaCl2·2H2O. • To find the number of grams of BaCl2·2H2O in the original mixture that produced

0.188 g Ba3(PO4)2, the following calculation is performed.

0.188g Ba3(PO4)2 x 1 mole Ba3(PO4)2/601.93g Ba3(PO4)2 x 3 mole Ba2+/1 mole Ba3(PO4)2

= 9.37 x 10-4 mole Ba2+

9.37 x 10-4 mole Ba2+ x 1 mole BaCl2·2H2O/1 mole Ba2+ x 244g BaCl2·2H2O/1 mole BaCl2·2H2O = 0.228 g BaCl2·2H2O is a part of the original salt mixture

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The mass of the Na3PO4·12H2O in the salt mixture must be the difference between the total mass of the original salt sample and the mass of the BaCl2·2H2O, or (0.942 g - 0.229 g =) 0.713 g. The percent BaCl2·2H2O in the salt mixture is 0.229 g/0.942 g x 100 = 24.3% BaCl2·2H2O The percent Na3PO4·12H2O in the salt mixture is 0.713 g/0.942 g x 100 = 75.7% Na3PO4·12H2O PROCEDURE Part A

1. Weigh a clean, dry 400 mL beaker. Record this mass for Trial 1 on the Report Sheet.

2. Transfer all the unknown salt mixture to the 400-mL beaker and reweigh. Record the mass on the Report sheet. Add 200 mL (±0.2 mL) of deionized water.

3. Stir the mixture with a stirring rod for about 1 minute and then allow the precipitate to settle. Leave the stirring rod in the beaker.

4. Digest the precipitate by covering the beaker with a watchglass and warm the solution (80-90°C) over a hotplate.

5. While the precipitate is being kept warm, proceed to set up the gravity filtration apparatus.

• Place your initials (in pencil) on a piece of filter paper. • Fold the filter paper and tear off its corner. Determine its mass (±0.001 g). • Seal the filter paper into the filter funnel with a small amount of deionized

water. • Discard the deionized water from the receiving flask. Have your

instructor inspect your apparatus before continuing. Periodically check on the progress of the heating solution.

6. After 30 minutes, remove the heat and allow the precipitate to settle; the solution does not need to cool to room temperature.

7. While the precipitate is settling, heat (80-90°C) about 30 mL of deionized water for use as wash water.

8. Once the supernatant has cleared, while still warm, filter the precipitate. Transfer any precipitate on the wall of the beaker to the filter with the aid of a rubber policeman.

9. Transfer two 50 mL volumes of the filtrate (measure with a graduated cylinder) into separate 100 mL beakers, labeled Beaker I and Beaker II. Save for Part B.

10. Wash the Ba3(PO4)2 precipitate on the filter paper with two additional 5 mL portions of warm water.

11. Remove the filter paper and precipitate from the filter funnel. 12. Air-dry the precipitate on the filter paper over the weekend. Return on Monday at

your convenience to weigh. 13. Determine the combined mass (±0.001 g) of the precipitate and filter paper. Record.

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B. Determination of the Limiting Reactant

1. Add 2 drops of the test reagent 0.5 M BaCl2 to the 50 mL of supernatant liquid in Beaker I. If a precipitate forms, the PO4

3- is in excess and Ba2+ is the limiting reactant in the original salt mixture.

2. Add 2 drops of the test reagent 0.5 M Na3PO4 to the 50 mL of supernatant liquid in Beaker II. If a precipitate forms, the Ba2+ is in excess and PO4

3- is the limiting reactant in the original salt mixture. An obvious formation of precipitate should appear in only one of the tests.

Disposal: Dispose of the barium phosphate, including the filter paper, in the "Waste Solids" container. Dispose of the waste solutions in the "Waste Liquids" container. CLEANUP: Rinse each beaker with small portions of warm water and discard in the "Waste Liquids" container. Rinse twice with tap water and twice with deionized water and discard in the sink.

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Name _____________________________________ Class ________________________ Prelaboratory Assignment

1. A 1.146 g mixture of the solid salts Na2SO4 and Pb(NO3)2 forms an aqueous solution with the precipitation of PbSO4. The precipitate was filtered and dried, and its mass was determined to be 0.672 g. The limiting reactant was determined to be Na2SO4.

a. Write the molecular form of the equation for the reaction.

b. Write the net ionic equation for the reaction.

c. How many moles and grams of Na2SO4 are in the reaction mixture?

d. What is the percent by mass of each salt in the mixture?

2. The Ba3(PO4)2 (molar mass = 601.93 g/mol) precipitate that formed from a salt mixture has a mass of 0.667 g. Experimental tests revealed that Na3PO4·12H2O (molar mass = 380.12 g/mol) was the limiting reactant in the formation of the precipitate and that BaCl2·2H2O was the excess reactant in the salt mixture. Determine the mass of Na3PO4·12H2O in the salt mixture.

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Name ________________________________ Class _____________________________ Data Sheet Unknown number ___________ Part A

Trial 1 1. Mass of 400 mL beaker + salt mixture (g) _______________ 2. Mass of empty 400 mL beaker _______________ 3. Mass of salt mixture (#1- #2) _______________ 4. Mass of filter paper + Ba3(PO4)2 (g) after drying _______________ 5. Mass of filter paper _______________ 6. Mass of Ba3(PO4)2(g) (#4- #5) _______________ Part B 1. Limiting reactant in salt mixture (write complete formula) 2. Excess reactant in salt mixture (write complete formula) Data Analysis 1. Moles of Ba3(PO4)2 precipitated (mol) ________________ 2. Moles of limiting reactant in salt mixture (mol)

• formula of limiting hydrate ________________ 3. Mass of limiting reactant in salt mixture (g)

• formula of limiting hydrate ________________ 4. Mass of excess reactant in salt mixture (g)

• formula of excess hydrate ________________ 5. Percent limiting reactant in salt mixture (%)

• formula of limiting hydrate ________________ 6. Percent excess reactant in salt mixture (%)

• formula of excess hydrate ________________ Show calculations for the trial below.

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General Chemistry Lab

Experiment 5 The Clock Reaction

Chemical reactions vary greatly in speed or rate. Some reactions, such as the explosion of methane are extremely rapid. Others, like geological process, may be so slow that centuries pass before the reaction has proceeded noticeably. Between these extreme rates, there are many reactions of moderate speed that can be studied in the laboratory. Factors that Influence Rates The concentrations of various reactants affect the rates of most reactions. Higher concentrations increase the reaction rate, and lower concentrations decrease it. Sometimes the concentration affects the rate in a direct linear relationship; doubling the concentration doubles the rate. Often it is a direct square relationship, in which case doubling the concentration quadruples the rate and tripling the concentration increases the rate by three squared or nine times. In some instances, a change in concentration has no effect on the rate. Each reactant has its individual effect independent of other reactants. Almost all chemical reactions proceed more rapidly as the temperature increased. The rate increase depends on the reaction. For many reactions, especially in organic chemistry, a rule of thumb is that the rate is approximately doubled for each 10oC rise in temperature. A catalyst is defined as substance that affects the rate of reaction but emerges unchanged from the reaction. Catalysts are usually thought of in a positive sense-as increasing the reaction's speed. Enzymes in the body are examples of some of nature's catalysts. Some reactions such as rusting of iron, take place at a surface. The rate of reaction is dependent on, among other things, the amount of surface area in contact with the other reactants. Colliding Molecules One can better understand the factors affecting the rates of reactions by considering the behavior of individual molecules involved. Most chemical reactions occur because two or more molecules collide effectively and bond together or because unstable molecules come apart by breaking bonds. In order for collisions to result in the formation of a new compound, the particles must collide with sufficient force and they must be in proper orientation toward each other. When substances are mixed, particularly in case of gases or liquids, the number of collisions is astronomically high. Only a fraction of collisions are likely to be effective. For slow reactions, the fraction of effective collisions is generally much less than for fast reactions.

Changes in conditions that make particle collisions more frequent or more effective speed the reaction. The increase in rate owing to higher concentrations of reactants is explained by the increase in the number of collisions. The more particles there are in a given volume, the more they will bump into one another. Raising the temperature causes the particles to mover faster. An increase in speed of the particles increases both the number and the effectiveness of collision, and so the reaction goes faster. A catalyst helps by providing a more effective pathway or mechanism for the reaction

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The Iodine Clock Reaction The "clock reaction" is a reaction is famous for its dramatic clear-to-blue color change, often used in introductory chemistry courses to explore the rates of reactions. The clock reaction is actually three reactions that occur in sequence: (1) IO3

- (aq) + 3 HSO3

- (aq)

I- (aq) + 3SO4

2- (aq) + 3 H+

(aq) (2) IO3

- (aq)

+ 5 I- (aq) + 6 H+

(aq) 3 I2 (aq) + 3 H2O (3) I2 (aq) + starch {a dark blue iodine/starch complex}

For this experiment, it's not important to understand how all three steps of the entire reaction work together. The iodate ion (IO3

-) reacts with the bisulfite ion (HSO3-). When

the bisulfite is used up, the excess iodate immediately signals the end by reacting with iodide (I-) to produce iodine (I2). The iodine reacts with the starch indicator to give a dark blue iodine/starch complex marking the completion of the reaction. The time required for the blue color to appear is related to the rate of the reaction. Reaction rate is defined as the change in concentration of a reactant or product per unit time. Rate = change of concentration = C Change in time t

The rate of a chemical reaction is smoothly changing quantity because it is dependent upon concentration, which changes as reactants are consumed and products are formed. For this lab, you will investigate how two factors affect the rate of the above reaction, the concentration of the reactants and temperature of the reaction. You will also prepare the reagents to be used in the experiment. PROCEDURE: Part 1 Preparation of solutions. The first part of the experiment will be the preparation of the reagents, KIO3 and NaHSO3. A. Preparation of KIO3 Calculate the molar mass of KIO3 and record below. 1. Molar Mass of KIO3 _____________grams/mole 2. Calculate the moles KIO3 needed to make 0.1 L (100mL) 0.1 M KIO3. Step 1 M=moles/liter Calculate the mass of KIO3 needed to make a 0.1M solutions (0.1 M KIO3 in moles/L) (Molar Mass of KIO3 in grams/mole) = _________g KIO3/L The answer above is mass of KIO3 needed to make a 1 Liter solution of 0.1M KIO3

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Step 2 Since you only need 100 mL (0.1 L) and 0.1L is 1/10 of l L, dividing the answer from a in Step 1 will give you the mass needed to make 100 mL of 0.1M KIO3. (_________g KIO3/L)/10) = __________g KIO3 /0.1L B. Preparation of 0.01 M NaHSO3 Calculate the volume of 0.1M NaHSO3 needed to prepare a 100 mL solution of 0.01M of NaHSO3. 1. Volume of 0.1 NaHSO3 needed to make 0.1L of 0.01 M NaHSO3 ________________mL Show calculation. (0.1L )(0.01M NaHSO3) = (V)(0.1M NaHSO3 )

Procedure:

1. Weigh the calculated amount of KIO3 from Step 2 above. 2. Rinse a volumetric flask with distilled water three times. This process of rinsing is called

conditioning. 3. Fill the volumetric flask half way with distilled water. Add the weighed KIO3 into the

volumetric flask. Swirl to dissolve. 4. Add water to bring the volume to the 100 mL mark. Cover the flask and mix by inverting

the flask back and forth. 5. Transfer the solution into a labeled 250 mL beaker. 6. Wash the volumetric flask with soap and water and condition the volumetric flask three

times with distilled water 7. Measure the calculated volume of 0.10 M NaHSO3 using a graduated cylinder. Half fill

the volumetric flask with distilled water. Add the 0.10 M NaHSO3. Complete the volume to 100 mL and mix thoroughly.

8. Transfer the NaHSO3 solution into a second labeled 250 mL beaker.

Part 2 Kinetics Before starting, clean and dry all the required glassware thoroughly. Trace contaminants can significantly affect your results in these reactions. Part I: A. Trials # 1-3.

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1. Label one a clean, dry 10-mL graduated cylinder. "sodium bisulfite, NaHSO3 ". Obtain a second clean, dry 10-mL graduated cylinder and label it "potassium iodate, KIO3 "

2. Measure 10 mL of 0.01 M sodium bisulfite using the labeled graduated cylinder. Add two drops of 4 % starch indicator to the sodium bisulfite solution. Measure 10 mL of 0.10 M potassium iodate using the appropriately labeled graduated cylinder. Place a dry 50-mL beaker on a white sheet of paper on the lab top. Have one person start the stopwatch, as a second person pours the two 10 mL solutions into the beaker. Mix once with a glass stirrer

3. When the color changes, stop the stopwatch and record the time on the data sheet.

4. Measure the temperature of the blue solution with thermometer and record on the data sheet.

5. Conduct two additional, identical trials (Trials 2 and 3) 6. Clean and dry the 50-mL beaker thoroughly, so you can re-use it. 7. Repeat steps 2-5, above, for trial #2 and trial #3

Part II: A. Trial #4 1. Clean and dry a 50-mL beaker thoroughly. 2. Repeat steps 2-5 in Part IA above. However, measure only 5 mL of the

0.10 M potassium iodate solution into the graduated cylinder, and then add 5 mL of water to bring the total to 10 mL. Then proceed as before.

C. Trial #5

3. Again, clean and dry the 50-mL beaker thoroughly. 4. Repeat steps 2-5 in Part I A above. However, measure only 2 mL of the

0.10 M potassium iodate solution into the graduated cylinder, and then add 8 mL of water to bring the total to 10 mL. Then, proceed as before.

5. Remember to add the two drops of starch indicator to the sodium bisulfite solution.

Part III: Reaction rate and temperature Using the same amounts of reactants used in Part I, you will conduct the reaction at two different temperatures. You will investigate the effect of temperature on reaction rate. A. Elevated temperature 1. Pour 10 mL of 0.010 M sodium bisulfate and 2 drops of starch indicator

into one small test tube, and 10 mL of 0.10 M potassium iodate into a second small test tube.

2. Put both test tubes into the warm water bath and allow them to sit for at least 5 min. (begin Part III B, while you're waiting.)

3. Record the temperature of the bath ( should be close to 45oC). 4. Conduct steps 3 and 4 from Part IA above to run the reaction.

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B. Lowered temperature 1. Prepare an ice bath by filling your 600-mL beaker 1/3 of the way with ice, and then adding tap water to fill ½ way. 2. Repeat steps 1-4 in Part III A, above, using the ice bath instead of the warm water bath. When you mix the two solutions in the 50-mL beaker, you can place the beaker gently in the ice bath while the reaction occurs, to keep the temperature constant.

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NAME______________________________________________ DATA SHEET Part I: Volume of 0.10 M potassium iodate used for each trial:__________ Volume of 0.010 M sodium bisulfite used for each trial: __________

Trial # Time (seconds) Temperature (oC) 1

2

3

Average time for trials 1-3: ___________ Part II :

Trial # Volume 0.01 M Sodium

Bisulfite

Volume 0.10 M

Potassium Iodate

Volume Water

Time (Seconds)

Temperature (oC)

4

10 mL 5 mL 5 mL

5

10 mL 2 mL 8 mL

Part III: A. Reaction at an elevated temperature: Volume of 0.10 M potassium iodate used: ____________

Volume of 0.010 M sodium bisulfite used: ____________ Temperature of warm water bath: ____________ Time for reaction to turn blue: ____________ B. Reaction at a lowered temperature:

Volume of 0.10 M potassium iodate used _____________

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Volume of 0.010 M sodium bisulfite used _____________ Temperature of ice water bath: _____________ Time for reaction to turn blue: _____________

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POST LAB QUESTIONS: Write up answers to the following on a separate sheet of paper.

1. In part I of this lab, you ran the reaction three times under the same conditions (that is, the same reactant concentrations and the same temperature.) Looking at your data, what can you say about the reproducibility of your results? How much uncertainty would you report in your average time of reaction? (For example, 100 ± 100 seconds, 120 ± 10 seconds, 125 ± 1 seconds, or 125.4 ± 0.1 seconds…)

2. Discuss the data you obtained in part II, in comparison to the average time calculated in part I. What conclusions would you draw about how the amount of potassium iodate affects the rate of the reaction?

3. Compare the data you obtained in part III, in comparison to the average time calculated in part I. What conclusions would you draw about how the temperature of the reaction solution affects the rate of the reaction?

4. According to "collision theory," two reactant molecules must collide for a reaction to occur between them. If the frequency of collisions is increased, the rate of a reaction then increases. Does your answer to questions 2 and 3 make sense, according to this theory? How would the two factors (amount of reactant and temperature) affect the number of collisions?

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NAME____________________________________________________________ PRELAB QUESTIONS

1. Why is important that all test tubes used in this experiment be clean and dry? What is affected if the test tubes are clean but wet?

2. What is the function of a catalyst? 3. What is the visual evidence which signals the completion of the iodine-clock reaction?

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Types of Chemical Reaction Introduction Most ordinary chemical reactions can be classified as one of five basic types. The first type of reaction occurs when two or more substances react to form a single compound. This type is called a combination reaction.

A + Z AZ A second type of reaction occurs when a single compound breaks down into two or more simpler substances, usually by the application of heat. This type is called a decomposition reaction.

AZ A + Z A third type of reaction occurs when one element displaces another element from a compound or aqueous solution. For this reaction to occur, the element that is replaced must be lower in the activity series. This type is called a single-replacement reaction.

A + BZ AZ + B A fourth type of reaction occurs when two substances in aqueous solution switch partners; that is, an anion of one substance exchanges with another. This type is called a double-replacement reaction.

AX + BZ AZ + BX A fifth type of reaction occurs when an acid and a base react to form a salt and water. This type is called a neutralization reaction.

HX + BOH BX + HOH Notice the hydrogen ion in the acid neutralizes the hydroxide ion in the base to form water. If water is written as HOH, the neutralization is more obvious and the equation may be easier to balance. In this experiment, we will carefully observe and record evidence for a chemical reaction. Evidence for a reaction may include any of the following: (1) a gas is produced; (2) a precipitate is formed; (3) a color change is observed; (4) an energy change is noted. In order to describe the reaction, we use various symbols in the chemical equation. Table 13.1 lists some of these. Table 1 Symbols in Chemical Equations

produces, yields (separates reactants from products) + added to, reacts with (separates two or more reactants or products)

heat (written above —>) NR no reaction (written after —>) (s) solid or precipitate (f) liquid

Symbol___ ___Translation

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(g) gas (aq) aqueous solution In order to write an equation, it is necessary to predict the products from a given reaction. Initially, this is a difficult task. To aid you in writing equations, word equations are supplied for each reaction. However, it is necessary to translate the word equations into balanced chemical equations. The following examples will illustrate. A. Conbination Reaction iron(s) + oxygen iron (III) oxide 4Fe + 3O2 2Fe2O3 B. Decomposition Reaction lithium hydrogen carbonate(s) lithium carbonate(s) + steam(g) + carbon dioxide(g)2 LiHCO3(S) Li2C03(s) + H20(g) + CO2(g) C. Single-Replacement Reaction tin(s) + hydrochloric acid(aq) tin(II) chloride(aq) + hydrogen(g)Sn(s) + 2 HCl(aq) SnCl2(aq) + H2(g)____________________________________________________________________ D. Double-Replacement Reaction potassium carbonate(aq) + calcium chloride(aq) calcium carbonate(s) + potassium chloride(aq) K2C03(aq) + CaCl2(aq) CaCO3(s) + 2 KCl(aq) ____________________________________________________________________ E. Neutralization Reaction — nitric acid(aq) + barium hydroxide(aq) barium nitrate(aq) + water 2 HNO3(aq) + Ba(OH)2(aq) Ba(NO3)2(aq) + 2HOH(l)____________________________________________________________________ PROCEDURE General Directions: For Procedures A-E, record your observations in the Data Table. A. A. Combination Reactions - Instructor Demonstration 1. Hold a 2 cm strip of magnesium ribbon with crucible tongs and ignite the metal in a hot burner flame. B. Decomposition Reactions 1. Put a few crystals of copper(II) sulfate pentahydrate in a dry test tube. Grasp the

test tube with a test tube holder or a two prong clamp. Heat the side of the test tube with a burner. Use a soft blue flame. Note the color change and observe the inside wall of the test tube.

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2. Add sodium hydrogen carbonate (baking soda) into a 250 mL Erlenmeyer flask so as to sparsely cover the bottom. Support the flask on a ring stand using a wire gauze.

a. Hold a flaming splint in the mouth of the flask for 10 seconds. Does the splint continue to burn or does the flame blow out. b. Now eat the flask strongly with the laboratory burner until moisture is observed on the side of the flask; quickly hold a flaming splint in the mouth of the flask and record how long it burns.

C. Single-Replacement Reactions 1. Put 20 drops of silver nitrate solution into a test tube and add a small piece of

copper wire. Allow a few minutes for reaction and then record your observation. 2. Put 20 drops of hydrochloric acid into a test tube and add a small piece of

magnesium metal. Record your observation. 3. Put 20 drops of distilled water into a test tube and add a small piece of calcium

metal. Record your observation. D. Double-Replacement Reactions 1-3. Put 10 drops of silver nitrate, copper(II) nitrate, and aluminum nitrate solutions into separate test tubes #1-3. Add a few drops of ammonium carbonate solution into test tubes #1, #2, and #3. Observe and record your observations. 4-6. Put 10 drops of silver nitrate, copper(II) nitrate, and aluminum nitrate solutions into separate test tubes #4-6. Add a few drops of sodium phosphate solution into test tubes #4, #5, and #6. Observe and record your observations. E. Neutralization Reactions 1. Put 10 drops of nitric acid, sulfuric acid, and phosphoric acid into separate test

tubes # 1-3. Add one drop of phenolphthalein into each of the test tubes. Add drops of

dilute sodium hydroxide solution into test tube #1 until a permanent color change is observed.

Note: Phenolphthalein is an acid-base indicator that is colorless in acidic and neutral solutions and pink in basic solutions.

2. Add drops of dilute sodium hydroxide solution into test tube #2 until a permanent color change is observed.

3. Add drops of dilute sodium hydroxide solution into test tube #3 until a permanent color change is observed.

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Name ________________________________________________________________ DATA TABLE Write your observations. Look for evidence of reacations. A. Combination Reactions - Instructor Demonstrations 1. Mg + O2 __________________________________________________ B. Decomposition Reactions 1. CuSO4.5H2O __________________________________________________

2. NaHCO3 __________________________________________________ C. Single-Replacement Reactions 1. Cu + AgNO3 __________________________________________________ 2. Mg + HC1 __________________________________________________ 3. Ca + H2O __________________________________________________ D. Double-Replacement Reactions 1. AgNO3 + (NH4)2CO3 ____________________________________________ 2. Cu(NO3)2 + (NH4)2CO3 ____________________________________________ 4. A1(NO3)3 + (NH4)2CO3 ____________________________________________ 5. AgNO3 + Na3PO4 ____________________________________________ 6. Cu(NO3)2 + Na3PO4 ____________________________________________ 7. A1(NO3)3 + Na3PO4 ____________________________________________ E. Neutralization Reactions 1. HNO3 + NaOH ____________________________________________ 2. H2SO4 + NaOH ____________________________________________ 3. H3PO4+ NaOH ____________________________________________

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Translate Each Word Equation into a Balanced Chemical Equation A.Combination Reactions - Instructor Demonstrations 1. magnesium(S) + oxygen(g) magnesium oxide(s) A. Decomposition Reactions 1. copper(II) sulfate pentahydrate(S) copper(II) sulfate(S) + water(g)

2. sodium hydrogen carbonate (s) sodium carbonate(S) + water(g) + carbon dioxide(g) C. Single-Replacement Reactions 1. copper(s) + silver nitrate (aq) copper(II) nitrate (aq) + silver(s) 2. magnesium(S) + hydrochloric acid(aq) magnesium chloride (aq) + hydrogen (g 3. calcium(s) + water(l) calcium hydroxide (s) + hydrogen (g) D. Double-Replacement Reactions 1. silver nitrate(aq) + ammonium carbonate(aq) silver carbonate(S) + ammonium nitrate(aq) 2. copper(II) nitrate(aq) +ammonium carbonate(aq) copper(II) carbonate(s) +ammonium nitrate(aq) 3. aluminum nitrate(aq) +ammonium carbonate(aq) aluminum carbonate(S) + ammonium nitrate(aq) 4. silver nitrate(aq) + sodium phosphate(aq) silver phosphate(S) + sodium nitrate(aq) 5. copper(II) nitrate (aq) + sodium phosphate(aq) copper(II) phosphate(S) + sodium nitrate(aq)

6. aluminum nitrate(aq) + sodium phosphate(aq) aluminum phosphate(S) + sodium nitrate(aq) E. Neutralization Reactions 1. nitric acid (aq) + sodium hydroxide (aq) sodium nitrate(aq) + water 2. sulf uric acid (aq) + sodium hydroxide (aq) sodium sulfate (aq) + water 3. phosphoric acid (aq) + sodium hydroxide (aq) sodium phosphate (aq) + water

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NAME ____________________ DATE _____________ POSTLABORATORY ASSIGNMENT 1. Provide the chemical formula for the following substances produced during the experiment. Refer to pages 148-149 for the substances produced from the chemical reactions. (a) the white smoke produced from reaction A.1 __________ (b) the colorless liquid produced from reaction B.1 __________ (c) the flame-extinguishing gas produced from reaction B.2 __________ (d) the gray solid produced from reaction C.I __________ (e) the colorless gas produced from reaction C.2 __________ (f) the yellow ppt produced from reaction D.I __________ (g) the blue ppt produced from reaction D.2 __________ (h) the white ppt produced from reaction D.3 __________

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Name:_____________________________________________________________________ PRELABORATORY ASSIGNMENT 1. In your own words, define the following terms: activity series - aqueous solution - catalyst - precipitate (ppt) - product - reactant - 2. Explain the meaning of the following symbols: , +, A, NR, (s), (l), (g), (aq)

3. List four observations that are evidence of chemical reaction.

4. What color is phenolphthalein indicator in (a) an acidic solution? (b) a basic solution?

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Experiment 7 Standardization of Base and Titration of Vinegar

Objective: Standardized NaOH solution. Determine the concentration of vinegar Reactions: Part A Standardization of NaOH Reaction

KHC8H4O4 + OH- H2O + KC8H4O4-

NaOH are solid pellets that readily absorbs water from the atmosphere. Consequently, the concentrations of unstandardized NaOH solutions are approximate. To determine the exact concentration of NaOH, it is titrated against a primary standard. For the experiment potassium hydrogen phthalate, KHC8H4O4, is the primary standard. The acronym, KHP is commonly used to call potassium hydrogen phthalate. The exact mass of KHP is weighed and the moles KHP is calculated.

(Mass KHP in grams)(1 mole KHP/204.4 g KHP) = mole KHP

The KHP is dissolved and phenolphthalein indicator is added. The KHP is the analyte. The NaOH solution is placed in the buret. The NaOH is the titrant. The NaOH solution is added to the KHP until the whole solution changes from colorless to a permanent faint pink color that last for 30 seconds. The reaction for the neutralization of KHP by NaOH is

KHC8H4O4 + OH- H2O + KC8H4O4-

At end point of titration the moles of KHP and NaOH is equal.

Moles KHP = Moles NaOH (at end point)

The molar concentration of the NaOH can now be calculated.

Molar concentration = moles NaOH/volume (L) of NaOH used Procedure

1. Label three 125 mL Erlenmeyer flask numbers 1 to 3. 2. Weigh three – 2 grams samples of dried KHP to the nearest mg into the three

labeled flasks. Make sure that you record the masses in the corresponding flasks. 3. Add 50 mL of distilled water each of the flask to dissolve. Swirling the contents

will aid in the dissolution. Add two drops of phenolphathlein to each of the flask and set aside.

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4. Rinse the buret with distilled water once. Follow by rinse with 5mL portion of NaOH. Repeat the NaOH two more times.

5. Fill the buret with the NaOH making sure that the tip is also filled and free of air bubbles. Record your buret’s initial volume.

6. Place Erlenmeyer flask under the buret and begin titration by dropping the NaOH into the flask in steady drops until a light pink (not fuschia) color stays for the whole solution for 30 seconds. Swirl the flask gently as the NaOH solution is being added. You are nearing endpoint when the pink color remains longer after each drop of NaOH. Record your buret’s final volume.

7. Titrate the two remaining samples. The approximately volume of NaOH solution it will take to reach endpoint is now known. There is no need to go as slowly as you did the first titration. Add the NaOH to within 3 mL of the first titration’s endpoint and titrate slowly from there.

8. Throw out data from titrations that have endpoints with fuchia color. Only endpoints with light pink endpoints will be considered as acceptable data.

9. Calculate and record the concentration of the standardized NaOH. Two titration must have concentrations that agree within one percent. If none of the three titration agree within 1 %, perform another trial until two titrations within 1% is achieved.

Calculation for 1 % agreement. M trial1 – M trial2 x 100% M average of trial 1 and 2

Part B Titration of Acid Solution A 10.0 mL volume of acid is titrated against the standardized NaOH. Again, at end point the moles OH- is equal to the moles H+ as shown below.

Moles NaOH = (Molarity of NaOH in moles/L)(volume NaOH used in L) = Moles H+

To calculate the concentration of H+:

Molar concentration of H+ = moles H+/volume (L) of H+ used

Procedure 1. Measure 10.0 mL of the vinegar solution using a clean, dry pipet and transfer to a 125

mL Erlenmeyer flask. Add two drops of phenolphthalein. 2. Refill the buret with the NaOH solution. Record the initial volume. 3. Titrate the acid until a light pink endpoint is reached. Record the final volume.. 4. Calculate the concentration of the unknown acid.

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Data Table Standardization of NaOH Trial 1 Trial 2 Trial 3 Mass KHP Moles KHP Final Buret Reading Initial Buret Reading Volume NaOH Molarity of NaOH

Average Molarity of NaOH __________________________

Show calculations below.

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Titration of Unknown Acid. Unknown #_____________

Trial 1 Trial 2 Trial 3 Volume of Acid Final Buret Reading Initial Buret Reading Volume NaOH Moles NaOH Moles Acid Molarity of Acid

Average Molarity of unknown acid ______________________ Show calculations below.

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PreLaboratory Questions.

1. How can you tell when the endpoint is near?

2. If a 2.051 grams KHP sample requires 27.30 mL of NaOH solution to reach endpoint, what is the concentration of the base? Show your calculation.

3. A 20 mL volume of unknown acid was titration against 0.09899 M NaOH. It took 50.00 mL of NaOH to reach end point. What is the concentration of the acid? Show your calculation.

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general chemistry 1 and 2 lab packet - seattle central...

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