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General Chemistry Lab I and II

Summer Quarter 2008

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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 Cabasco-Cebrian

CLASS TIME

Section A 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) tcabas@sccd.ctc.edu

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. -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, which is scheduled on Friday Aug 29.

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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 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.

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Conversion scale Transcript Grade Point* A 95 - 100% Outstanding 3.8 - 4.0 4.0 A- 90 – 94.9% 3.5 – 3.7 3.7 B+ 87 – 89.9% 3.2 - 3.4 3.3 B 83 – 86.9% Above Average 2.9 - 3.1 3.0 B- 80 – 82.9% 2.5 - 2.8 2.7

C+ 77 – 79.9% 2.2 – 2.4 2.3 C 73 – 76.9% Average 1.9 – 2.1 2.0 C- 70 – 72.9% 1.5 – 1.8 1.7 D+ 67 – 69.9% 1.2 – 1.4 1.3 D 63 – 66.9% Below Average 0.9 – 1.1 1.0 D- 60 – 62.9% Lowest Passing Grade 0.6 – 0.8 0.7 F Below 60% Failing 0.0 – 0.5 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

7/11 Sec A

Check-in Read Safety Precaution Experiment 1: Measurements and Observations

7/18 Sec A

Experiment 2: Covalent Bonds

7/25 Sec A

Experiment 3 Cations and Anions

8/1 Sec A

Experiment 4: Stoichiometry

8/8 Sec A

Experiment 5: Types of Chemical Reaction

8/15 Sec A

Experiment 6: Solutions and Clock Reaction

8/22 Sec A

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

8/29 Sec A Experiment 8: Make up lab Lab report due Sat, Aug 30 @ 8:00 PM by email only.

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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 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. Turn in Schedule: Pre-lab questions are due at the beginning of each lab period. Data sheets and post lab questions are due the following lab period.

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LAB REPORT Formal lab reports will be required for General Chemistry lab 2 only. For General Chemistry lab 1, a clean and legible data sheet 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 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. 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.

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

MASS

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In the laboratory the terms mass and weight are used interchangeably; however, there 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 essentially no weight 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 an object and the weight of an object are interchangeable as long as the force of gravity remains constant. The basic unit of mass in the metric system is the gram (g). When we "weigh" an 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 operates by 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 widely used balances in chemical laboratories today (Fig. 2). Commonly, the top loading balances measure mass to the nearest centigram, that is, it weighs to 100th of a gram. Another common usage is to weigh to the nearest 0.01. This simply instructs you to use a top loading balance that is capable of weighing to the two significant numbers after the decimal. When using an electronic balance, you record the mass that balance is capable of giving. Do not round.

Figure 2 Electronic Top-loading Balance

TEMPERATURE Heat is the most common form of energy used in general chemistry laboratories. Heat flows between objects that are at different temperatures. Thus, temperature is a measure of the heat (energy) of an object or the flow of heat between two objects. A thermometer is used to measure temperature in the laboratory. A thermometer is a capillary tube that is generally filled with mercury or colored alcohol. These liquids are used in thermometers because they readily expand or contract with small changes in temperature. Thermometers in the laboratory are generally calibrated in degrees Celsius. Use utmost care when using mercury thermometers. Mercury vapors are extremely toxic.

Report mercury thermometer breakage to your instructors immediately.

VOLUME The basic unit of volume in the metric system is the liter (L). Graduated glassware used to

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quantitatively measure volume in the laboratory includes graduated cylinders, pipets, volumetric flasks and burets. Beakers and Erlenmeyer flasks are graduated to approximate volumes only. Errors in the graduation can be > 10% and should not be used for quantitative measurements. For example, a 50-mL beaker may actually hold 45-55 mL. Graduated cylinders are commonly used to make more precise measurements of 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 and major divisions of 5 or 10 mL. Estimates of volume can be made to the nearest 0.1 mL using a graduated cylinder. Volumetric flasks are not graduated and measure only fixed volume. Volumetric flasks come in different volume sizes. There 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 the volumetric 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 mL. It is most commonly used for titration.

WASHING AND CLEANING OF GLASSWARE

Washing and brushing with a detergent can clean most pieces of laboratory glassware. After they have been thoroughly cleaned, they are rinsed with tap water and finally with a spray of distilled water. If the surface is clean, the water will wet the surface uniformly.

READING THE MENISCUS Volumetric flasks, burets, pipets and graduated cylinders are calibrated to measure volumes of liquids. When a liquid is confined in a narrow tube such as a buret or a pipet, the surface is found to exhibit a marked curvature, called a meniscus. It is common practice to use the bottom of the meniscus in calibrating and using volumetric ware. Special care must be used in reading this meniscus. By positioning a black-striped white card behind the meniscus, which is transparent, it becomes more distinct.

Figure 3 Meniscus

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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 **

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.

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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 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

GRAVITY FILTRATION

REPARING FILTER PAPER FOR A FILTER FUNNEL

Figure 6

Folding of a Filter Paper

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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.

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.

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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.

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

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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 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.

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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.

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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Measurement

Name: _____________________________________________________

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.

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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.

Measurement

Name __________________________________________

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 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.

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 require 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 of electron dots against the sum of all valence electrons.

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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.

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.

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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.

General Chemistry Covalent Bonding Name________________________________________________ DATA TABLE

Molecular Models with Single Bonds

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Covalent Bonds

Prelab Questions Name ______________________________________ 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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General Chemistry Lab Cations and Anions

Ions are nothing more than atoms or molecules that have gained or lost an electron. Those that have lost an electron are called positive ions, while those that have gained an electron are negative ions. Ions are formed when an electron is detached from a neutral molecule (or atom). The molecule losing an electron becomes a positive ion and the molecule gaining an electron becomes a negative ion. In an ionic solution both the cation and anion exist.

In a chemical reaction ions give distinctive flame tests, undergo color changes, and produce gas or insoluble solid products. In this lab, you will be given one aqueous solution which contains two unknown ions: one cation (either K+, Ca2+, NH4

+, or Fe3+) and one anion (Cl-, SO42-, PO4

3-, or CO32-). By the end of

the lab, you should be able to identify both the cation and anion in your unknown solution. You’ll then use the same methods to analyze a common household product, and to make some conclusions about which ions it contains. To do all this, you’ll use characteristic chemical properties of the different ions, comparing the reactivity of your unknown to that of the provided reference solutions. For example, the reference solution for sodium, Na+, is 0.1 M NaCl, because it contains sodium cations. The presence of the anion, Cl-, is not important since it will not affect the cation tests. (The “0.1 M” represents the solution strength, or concentration, which we’ll learn about later in the quarter.) The specific tests you’ll perform include the following:

1. The Flame Test. Different elements emit characteristic colors when heated in a flame. This is due to electron transitions. When an atom is heated, it absorbs energy, and an electron can “jump” to a higher energy level. This “excited” electron will later return to its initial spot, and this “drop” releases light of different colors. Barium’s flame color is characteristically green/blue, for example. In this lab, you’ll observe the characteristic colors of sodium, calcium and potassium.

2. The Iron (III) Test. Iron (III) cations are easy to identify based on their reaction with

thiocyanate, SCN-, shown by the reaction below. Both reactants are clear, colorless liquids, but the product is a distinct blood-red color. Any sample that turns red when KSCN is added is positive for iron (III).

Fe3+ + 3SCN- Fe(SCN)3

3. Ammonium Test. Ammonium cations can be identified by reaction with hydroxide: NH4

+ + OH- NH3(g) H2O

The NH3 produced by this reaction is a gas, and you can notice its presence as a characteristic ammonia smell. Because smell is unreliable, you will use litmus paper to scientifically detect its presence. Litmus paper held above the test tube containing NH4

+ will turn blue (after hydroxide is added), as NH3 vapors diffuse up from the reaction mixture to the litmus paper.

There are several other tests, besides these three examples. The reactions may seem complicated, but the overall approach is simple: If you observe matching results between a reference solution and your unknown, you can be confident that your unknown contains the ion in that reference solution. You can say

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that your unknown is “positive” for the one cation and one anion it contains, and “negative” for all the ions it does not contain. PROCEDURE General notes:

• Be as clear and detailed in your observations as possible. • Keep all your test tubes well-labeled. White labeling tape is available. • If you get ambiguous results for a test, it is worthwhile to repeat it. • All reagents are in dropper bottles. You will add the reagents directly from the dropper bottles into your test tubes, by squeezing. No need to use graduated cylinders: there are about 40 drops in 2 mL of liquid. To save time, you can make a “measuring test tube” by putting 2 mL of water in a tube and then using it as a visual comparison to roughly measure 2 mL of any substance into a new test tube.

• A “community water bath” will be set up. Disposal and safety instructions: Please tie back long hair and secure loose clothing when using the Bunsen burners. Follow instructions on proper use of the Bunsen burners carefully. Silver nitrate (AgNO3) will stain your skin, and NaOH and HCl can cause burns: Use running water to wash off these chemicals if spilled on your skin, and then notify an instructor. All waste should be put in the labeled waste jar for the class. Part A -- Cation tests A1. Flame tests for K

+, and Ca

2+

Part A -- Cation chemical tests A1. Flame tests for, K

+, and Ca

2+

You will conduct flame tests to observe the flame emission colors for the following ions: Na+

, K+

, and Ca2+

. Sodium will be demonstrated by the instructor, because it takes a very long time to clean it off of the wire loop. Record your observations for sodium’s characteristic flame color in the table on your “Lab Report” pages. Safety precaution: Please tie back long hair and secure loose clothing when using the Bunsen burners. Do not touch the chemicals with your fingers! Wash your hands immediately if you accidentally touch the chemicals.

1. Get a wire loop. Light your Bunsen burner with a striker. Adjust the flame so that you have a distinct inner and outer cone. Your instructor will check to make sure that your flame is adjusted properly for the activity. 2. Place about 20 drops of 6 M HCl in one test tube. Fill a second test tube half-way with DI water. These will be your cleaning solutions for the wire loop. 3. Dip the wire loop in the 6 M HCl and place the loop at the tip of the inner cone of the flame. Burn the loop until the flame remains blue, indicating that the wire loop is clean. If it’s not clean, repeat the dipping and burning a few times. Consult with your instructor if it seems that there is a persistent contaminant on your loop. 4. To conduct the test for K

+

, transfer a few grains of KCl to a watch glass and a 1 drop of DI water. Roll the loop in the solid KCl. Place the wire loop in the hottest part of the flame. Observe and carefully record the color of the flame on the data table, being as descriptive as possible. In addition to the color, you should record observations about the duration, intensity, and shape of the

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flame. You may repeat the test several times, to be certain about your results. 5. Clean the wire loop between samples by dipping it in the DI water, then into a fresh 6 M HCl, and then burning it until it is clean. Pour out the water cleaning solution between samples, because it easily becomes contaminated. Rinse the water tube with water before putting fresh DI water in it. Repeat the test with solid CaCl2, and record your results for calcium in the table. 6. Next, obtain a vial with an unknown solid from the stockroom. Repeat the flame test with your unknown. (Make sure you have cleaned your wire loop beforehand, as described in step 6.) For the remaining tests, your unknown must be in the form of a solution, so add deionized water to the vial containing your unknown solid, and shake the vial to dissolve your solid, making sure not to splash or lose any of your unknown.

A2. Oxalate test: Test for Ca2+

1. Place 2 mL of 0.1 M CaCl2 (the reference solution) in a labeled test tube and 2 mL of your

unknown solution in another test tube. 2. Add 15 drops of ammonium oxalate solution, 0.1 M (NH4)2C2O4, to each and mix. Look for the formation of a cloudy, white solid (a “precipitate”). 3. Sometimes, heat is required for the precipitate to appear. If no precipitate forms in step 2, put the test tube in the warm water bath for about five minutes, and check again for a precipitate. If there is still no precipitate, there is no Ca2+ in that sample. Record your observations.

A3. Test for ammonium ion, NH4+

1. Place 2 mL of 0.1 M NH4C1 (the reference solution) in a test tube and 2 mL of your unknown in another test tube. Place two red litmus paper on a watch glass and moisten the litmus papers with DI water.

2. Add 15 drops of 6 M NaOH to each tube without letting any of the NaOH touch the mouth of the test tube and mix. If the NaOH touches the lip of the test tube, discard and start again.

3. Place a strip of moistened red litmus paper across the top of each test tube and set both test tubes (with the strips still balanced on top) in a warm water bath. If there is NH4

+ in the sample, ammonia gas (NH3) will be released and diffuse up the tube to the litmus paper and turn the paper a distinct blue.

A4. Test for iron (III) ion, Fe3+ 1. Place 2 mL of 0.1 M FeCl3

(the reference solution) in one test tube and 2 mL of your unknown in another test tube.

2. To each tube, add 5 drops of 6 M HNO3 and 2-3 drops of potassium thiocyanate, 0.1 M

KSCN and mix. An intense blood-red color indicates that Fe3+ is present in that sample. (A

faint pink color is not a positive test for iron.) Record the results.

Part B: Anion tests

B1. Test for chloride ion, Cl-

1. Place 2 mL of 0.1 M NaCl (the reference solution) in a test tube and 2 mL of your unknown in another test tube. 2. To each sample, add 5-10 drops of 0.1 M AgNO3

and 10 drops of 6 M HNO3. Gently mix the test tube contents by “flicking” the bottom of the tube (without spilling the contents!) A white solid that appears briefly, but then disappears with the HNO3, is NOT a positive result. A positive result for chloride is the appearance of a white solid that remains even after the HNO3 is added and mixed. Record the results for both your known and unknown.

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B.2 Test for sulfate ion, SO4

2- 1. Place 2 mL of 0.1 M Na2SO4 (the reference solution) in a test tube and 2 mL of your unknown in another test tube. 2. Add 1 mL (20 drops) of BaCl2 and 5-6 drops of 6 M HNO3 to each test tube and mix. As before, you are looking for a white solid that persists after everything is added; temporary appearance of a white solid that redissolves when HNO3 is not a positive result. Record your results for both the known and unknown.

B.3 Test for phosphate ion, PO4

3-

1. Place 2 mL of 0.1 M Na3PO4 (the reference solution) in a test tube and 2 mL of your unknown in another test tube. Add 10 drops of 6 M HNO3 to each. Warm the test tubes in a hot water bath (60°C) for a couple minutes, and then add 5 drops of ammonium molybdate solution, (NH4)2MoO4 and mix. The formation of a yellow precipitate indicates the presence of PO4

3-. Record the test results of the known and the unknown. B.4 Test for carbonate ion, CO3

2-

1. Place 2 mL of 0.5 M Na2CO3 (the reference solution) in a test tube and 2 mL of your unknown in another test tube. 2. While carefully observing the solution, add 10 drops of 3 M HCl to one sample and swirl. Watch for bubbles of CO2 gas as you add the HCl. The gas bubbles are formed quickly, and may be missed. If gas bubbles are not observed, add another 15-20 drops of HCl as you watch the solution. Repeat with the second sample – it’s better to do them one at a time.

DATA AND OBSERVATIONS Affix your unknown

sticker here

Part A: Cation tests A1. Flame tests:

Ion tested Flame Color

Potassium (K+)

Calcium (Ca2+)

Unknown

A2, A3, A4. Other tests of cations:

Test Observation for reference

solution Observation for unknown

solution

Oxalate test

(test for Ca2+)

Ammonium

test

Iron test

Conclusion from Part A: What cation does your unknown solution contain?

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Part B: Anion tests

Test Observation for reference

solution Observation for unknown

solution

Chloride test

Sulfate test

Phosphate

test

Carbonate

test

Conclusion Part B: What anion does you solution contain?

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Cations and Anions Prelab Questions Name:_____________________________________________________________

1. Describe how you will clean the wire loop between flame tests.

2. Suppose your unknown sample turns deep red when KSCN is added. What cation does the unknown contain?

3. What gas is produced when NaOH is added to NH3Cl and how is it detected?

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General Chemistry Lab

STOICHIOMETRY OF HYDROGEN PEROXIDE AND BLEACH

Hydrogen peroxide, H2O2, is a disinfectant for cuts and sold as a 3% (weight/weight) solution in Bleach, whose active ingredient is sodium hypochlorite, NaOCI is sold as a 6% (w/w) solution in water. When hydrogen peroxide and sodium hypochlorite are mixed a chemical reaction resulting in bubbles of oxygen gas is formed. Two factors that affect the yield of the products in chemical reactions are the amounts of the starting materials, the reactants, and the percent yield of the reaction. Chemicals react according to fixed mole ratios. Products are formed are limited by the amount of starting materials. A good allegory is the making of a cake. The main ingredients for a two layer cake are 4 cups of flour, 2 cups of sugar, ¼ cup oil, 2 sticks of butter and 1cup of milk. However, in gathering of the ingredients you find out that you have more than enough of all the ingredients except for there is only one stick of butter. You can decide not to make the cake or use all the butter and make just one layer. In this scenario, the available butter determined how many layers of cake you can make. The butter was the limiting ingredient. You are making the same cake just less of it. In chemistry, the ratios of reactions are based on moles. A mole is the quantity of a substance that contains 6.02 x 1023 units. A "unit" is the smallest measurable entity in the substance, generally either an atom or a molecule. One mole of a substance is equal to the substance's molar mass or molecular weight, in grams/mole. In a given chemical reaction, moles of reactants are mixed to give moles of product. In the reaction sample below one mole of sodium hydrogen carbonate, NaHCO3(s) (baking soda), reacts with one mole of acetic acid, C2H3O2 (vinegar) to form a mole of products of carbon dioxide gas, CO2; water, H2O, and sodium acetate, NaC2H3O2 .

NaHCO3(s) + HC2H3O2(aq) → CO2(g) + H2O(l) + NaC2H3O2(aq)

Just as the cake, the amount of product formed is based on the mole ration and amount of reactants available. For example: The available NaHCO3 is only 2 grams, but there is more 100 mL of vinegar. Vinegar is 5 % (w/w)acetic acid. How much products can be made? Remember that reactions are based on mole ratio, so we need to convert the mass of the NaHCO3 and HC2H3O2 to moles. To do this we need the molar mass of the two compounds. NaHCO3 has a molar mass of 83 grams/mole and HC2H3O2 is 33 grams per mole. How many moles are in 2 gram of NaHCO3 and 100 mL of HC2H3O2? Assume that the density of vinegar is 1 gram/mL 1. Mole of NaHCO3 in 2.0 grams

2.0 gram NaHCO3 x 1 mole NaHCO3 = 0.0241 mole NaHCO3 83 gram NaHCO3

2. Mole of HC2H3O2 in 25 grams How many grams of HC2H3O2 are in 100 mL of vinegar? We can convert from grams to mL using the density of vinegar which is 1 gram/mL. 200 mL vinegar x 1gram vinegar = 100 gram vinegar mL vinegar How many moles of HC2H3O2 are in the 100 grams vinegar? Remember vinegar is only 5% HC2H3O2. Another way to write 5% w/w is 5 gram HC2H3O2/100 gram solution.

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100 gram vinegar x 5 gram HC2H3O2 = 5.00 gram HC2H3O2 100 gram vinegar To find how many moles of HC2H3O2 is in 5.00 grams of HC2H3O2 we would need the molar mass of HC2H3O2 which is 33 grams/mole. 5.00 gram HC2H3O2 x 1 mole HC2H3O2 = 0.152 mole HC2H3O2 33 gram HC2H3O2

Available to make the products: CO2, H2O, and NaC2H3O2 are 0.0241 mole NaHCO3 and 0.152 mole HC2H3O2. Although there is considerably more vinegar, the amount of products we can make is limited by the 0.0241 moles of NaHCO3. Because the mole ration of the reaction is 1 to 1 for the reactants, we will be only using 0.0241 moles of the HC2H3O2. If we were to write the reaction based on the limiting reagent, AgNO3, the equation will look like the equation below.

0.0241 NaHCO3(s) + 0.0241 HC2H3O2(aq) → 0.0241 CO2(g) + 0.0241 H2O(l) + 0.0241 NaC2H3O2(aq)

You can measure the volume of the CO2 gas using a set up where the CO2 gas displaces the water. Once the volume is determine the moles of CO2 can be calculated using the ideal gas law, PV = nRT. P is ambient pressure, V is the volume of the gas, n is mole of the gas, T is temperature of the gas and R is a constant. Pressure is measured in atm, Volume is measure in L, Temperature is measure in oK (273oK/1oC). Lastly R is equal to 0.0821 L·atm K·mol Let look at the example above. When 100 ml of vinegar is added to 2 grams of NaHCO3, the reaction yielded 58.0 mL of CO2. We can calculate the number of moles in a 58.0 mL volume using the ideal gas law. P from the barometric reading = 1 atm T of the reaction = 24 oC = 297o K (24 + 273) V= 55.0 mL = .055 L (5.0 mL x 1L/1000 mL) R= 0.0821 L·atm K·mol PV= nRT this equation can be setup so that you are solving for n. PV = n RT 1 atm x .055 L = 0.0240 moles CO2 00821 L atm x 279 K K mol CO2 How close is calculated moles CO2 to the expected or theoretical result? In this experiment, the product of the reaction is a gas. You will measure how much oxygen is produced when known amounts of hydrogen peroxide solution are mixed with known amounts of Clorox bleach. The chemical equation for the reaction of NaOCl and H2O2 is below. The mole ratio for the reaction is 1.

NaOCl(aq) + H2O2 → O2(gas) + NaCl(aq) + H20

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Because oxygen is a gas, the volume will be determined by displacement. As the gas is being evolved, it will displace its volume in water. Based on the amount of reactants used (bleach and hydrogen peroxide) and the amount of product formed (oxygen) you will determine the stoichiometry of the chemical reaction. Part of the calculation will involve the conversion of the % wt/wt to moles using the density of each of the reactant. You will be walked through the calculation in the prelab questions. PROCEDURE You will run two sets of reactions for this experiment, in Reaction Set A, the volume of bleach will remain constant (5 mL) and the volume of hydrogen peroxide will vary. In Reaction Set B, the volume of hydrogen peroxide will remain constant (5 ml) and the volume of bleach will vary. Listed below are the volumes of reactants to be used for each set of reactions Set A

Volume of Bleach

Volume of Hydrogen Peroxide

Run 1 5ml_ 1.5mL

Run 2 5ml_ 2.5mL

Run3 5 ml 3.5 ml

Run 4 5mL 4.5 mL

Run5 5 ml 5.5 mL

Run 6 5mL 6.5 mL

Run? 5 ml 7.5mL

Set B Volume of Bleach

Volume of Hydrogen Peroxide

Run 1 1.5mL 5mL

Run 2 2.5 mL 5mL

Run3 3.5mL 5mL

Run 4 4.5mL 5mL

Run 5 5.5 mL 5mL

Run 6 6.5mL 5mL

Run 7 7.5mL 5mL

1. Label two clean, dry 100 mL beakers, one for bleach and one for hydrogen peroxide. Obtain

approximately 60 mL of bleach and 60 mL of hydrogen peroxide.

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2. Obtain 2 -10 mL graduated cylinder and label one for bleach and one for hydrogen peroxide. With the aid of a plastic pipet, transfer the designated amount of bleach (5 mL for Set A, reaction 1) into the 10 mL graduated cylinder. Since the experiments are to be semi- quantitative, it is necessary to measure the quantities exactly. Pour the bleach into the Erlenmeyer flask.

3. Measure the appropriate volume of hydrogen peroxide for the trial and pour into the vial. Rinse off any residual hydrogen peroxide on the outside of the vial and wipe dry. Set aside.

4. Fill a water trough with tap water. Completely immerse a 100 mL graduated cylinder (remove the plastic end from the cylinder) in the water trough, filling it with water. Turn the cylinder upside down, keeping the mouth below the surface of the water in the trough. At this point the up-ended graduated cylinder should be full of water (NO AIR BUBBLES) and held in place so that the opening is on the “undivided indentation” on the bottom of the trough. This is the oxygen-measuring vessel; gas formed by the reaction will bubble into the up-ended cylinder, where it will displace some of the water, and you can read the volume displaced directly from the graduations on the cylinder.

5. Place the flask containing the vial near the trough. Using a forcep, very carefully place the vial containing the hydrogen peroxide on the center of the flask.

6. Place the the rubber tubing with the bent glass into the mouth of the up-ended graduated cylinder. Place the rubber stopper end of the tubing on the flask tightly so it is leak free. Gently, jiggle the flask so that all the contents of the vial is transferred out and swirl the flask gently. The reaction will generate gas (bubble). Swirl the flask gently until no more bubbles evolve.

7. Once the setup is complete, record the amount of gas that was produced by reading the graduated cylinder.

8. Get a 600 mL beaker to contain your spent reactions. Rinse the flask and vial (including the outside) at least three times with DI water and proceed to the next trial.

9. Repeat all of the reactions noted above for Reaction Set A and Reaction Set B.

DATA AND CALCULATIONS 1. Graph A: Plot the volume (ml) of oxygen produced versus the volume (mL) of hydrogen peroxide solution used for the reaction set where bleach is held constant.

2 Graph B: Plot the volume (mL) of oxygen produced versus the volume (mL) of bleach solution used for the reaction set where hydrogen peroxide is held constant.

3 From each graph determine at what point does adding more reagent fail to produce more oxygen. For example (and not a correct one), when you are reacting bleach with 4 ml of bleach the curve levels off and you get the same amount of oxygen no matter how much additional bleach is added. Were the limiting reactants determined for both experiments?

Data: Graph A ___________ Graph B _____________

Follow the steps below to calculate the mole ratio of bleach to hydrogen peroxide for each graph.

NOTE: Assume the density of bleach and the density of hydrogen peroxide solution to be 1 .00 g/mL. 5. Calculate the number of moles of NaOCI in 5 mL of bleach [bleach is 6% (w/w) NaOCI.

6. Calculate the number of moles of H2O2 above volume recorded from Graph A required to consume 5

mL of bleach, NaOCl. H2O2 is 3% (w/w). 7. What is the mole ratio of the reactants. 8. Repeat these calculations for the data obtained from Graph B.

9. Indicate on Graph A where hydrogen peroxide is the limiting reagent in the reaction. Indicate on Graph A where hydrogen peroxide is in excess (not all of the reactant is consumed).

10. What can you conclude with Graph B?

11. Calculate the moles of O2 collected. What is the percent difference between the experimental result and theoretical. Hint: absolute value of the difference between theoretical and experimental result divided by the theoretical multiplied by 100.

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Stoichiometry Prelab Questions Name:_______________________________________________________________ 1. Calculate the moles of H2O2 in a 30 mL of 6% solution (w/w). Assume that the density of H2O2? A 6% w/w solution means that 6 grams of H2O2 is dissolve water to a total of 100 gram solution or 6 gram H2O2/100 gram solution

a. Calculate the molar mass of H2O2. ______________________gram/mole

b. Calculate the how many grams of H2O2 is in 30 mL of H2O2 solution. Assume that the density of H2O2 is 1 gram/mL. How would you set this up the equation so that you end up with an answer with a unit in grams? __________________grams

c. To calculate the moles H2O2 in 30 mL of H2O2, set up the equation using 15 grams/100 gram solution, grams of H2O2 in 15 mL solution which is the answer form 1b, and molar mass of H2O2 in grams/mole from 1a. Remember to have the units cancel so that your answer has the mole as a unit.___________________mole

2. Calculate how many moles of NaOCl is in 10 mL of bleach. Bleach is 6% NaOCl (w/w). Assume that its density is also 1 gram /mL. _________________mole NaOCl.

3. An experiment shows that increasing the volume of H2O2 past 2.5mL showed no increase in the volume of O2 generated when added to 5.0 mL of bleach.

a. Determine the number of moles of H2O2 in 2.5 mL solution.

b. Determine the number of moles NaOCl in 10.0 mL solution.

b. Which of the two is the limiting reagent?

c. What is the mole ratio of the reactants?

General Chemistry Lab 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

Symbol___ ___Translation 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 (l) liquid (g) gas (aq) aqueous solution

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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. Combination 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.

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. Heat 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.

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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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Chemical Reactions Name ________________________________________________________________ DATA TABLE _____________________________________________________________________ Procedure __________________________Evidence of Reactions________________ A. Combination Reactions Demonstration 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) B. 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. sulfuric acid (aq) + sodium hydroxide (aq) sodium sulfate (aq) + water 3. phosphoric acid (aq) + sodium hydroxide (aq) sodium phosphate (aq) + water

Chemical Reactions NAME ____________________ 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 __________ 3. Convert the following word equations to balanced chemical equations, (a) copper metal (S) + oxygen(g) copper(II)oxide (s) (b) iron(m) carbonate(S) iron(in) oxide(S) + carbon dioxide(g) (c) sodium metal (S) + water(l) — sodium hydroxide (aq) + hydrogen (g) (d) aluminum metal (s) + sulfuric acid (aq) aluminum sulfate (aq) + hydrogen (g) (e) copper(II) sulfate (aq) + lithium chromate(aq) copper(II) chromate(s) + lithium sulfate(aq)

(f) acetic acid (aq) + barium hydroxide(aq) barium acetate(aq) + water

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Chemical Reactions Pre-lab Questions Name:_____________________________________________________________________ 1. In your own words, define the following terms: aqueous solution - catalyst - precipitate (ppt) - product - reactant - 2. Explain the meaning of the following symbols: ,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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General Chemistry Lab

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 1a 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.

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.

3. 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

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4. When the color changes, stop the stopwatch and record the time on the data sheet. 5. Measure the temperature of the blue solution with thermometer and record on the

data sheet. 6. Conduct two additional, identical trials (Trials 2 and 3) 7. Clean and dry the 50-mL beaker thoroughly, so you can re-use it. 8. 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.

B. Trial #5

1. Again, clean and dry the 50-mL beaker thoroughly. 2. 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.

3. 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.

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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The Clock Reaction

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 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: ____________

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B. Reaction at a lowered temperature:

Volume of 0.10 M potassium iodate used: _____________

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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The Clock Reaction Pre-lab questions NAME___________________________________________________________

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?

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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Titration of a Weak Acid with a Strong Base

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

Label three 125 mL Erlenmeyer flask numbers 1 to 3. 3. 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. 4. Add 50 mL of distilled water each of the flask to dissolve. Swirling the contents

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

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

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6. 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.

7. Place Erlenmeyer flask under the buret and begin titration by dropping the NaOH into the flask in steady drops until a light pink (not fuchsia) 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.

8. 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.

9. Throw out data from titrations that have endpoints with fuchsia color. Only endpoints with light pink endpoints will be considered as acceptable trial.

10. 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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Titration Pre-Laboratory Questions Name____________________________________________________________

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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