lab experiment 1 making observations · lab experiment 1 making observations objectives: 1. to make...

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LAB EXPERIMENT 1 MAKING OBSERVATIONS OBJECTIVES: 1. To make observations while watching materials interact and undergo change. 2. To understand, recognize and record qualitative and quantitative observations. 3. To make interpretations based upon observations and data. PREDICTIONS: Make sure you read through the lab and understand the objectives

before making your predictions. MATERIALS: Apparatus Reagents 250 ml beaker aluminum foil (about a 10 cm X 10 cm sheet) Thermometer 1 M copper (II) chloride solution Mass balance PROCEDURES: 1. Get a piece of aluminum foil and record the mass. 2. Place 75.0 ml of copper (II) chloride solution into a clean 250 ml beaker, and

record the temperature of the solution. Make and record observations about the chemicals that you are using before you start step 3.

3. Crumple the aluminum foil into a loose ball and place it into the beaker

containing the copper (II) chloride. 4. Immediately start recording your observations in Table 1. As well, make and

record observations about the chemicals as the reaction is taking place. Make sure you record the highest temperature achieved.

5. When no further changes appear to be happening, make and record observations

about the chemicals once the reaction is complete. 6. Place all the contents of the beaker in the waste disposal jar provided by your

instructor. 7. Before you leave the lab, wash your hands thoroughly with soap and water.

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OBSERVATIONS AND DATA: Volume of copper (II) chloride used: __________ Mass of aluminum foil used: __________ TABLE 1: Temperature of the solution before, during and after the reaction

TIME (minutes) TEMPERATURE (°C) 0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0 4.5 5.0 5.5 6.0

Maximum temperature achieved: __________ QUESTIONS AND CALCULATIONS: 1) Knowing the temperature change that occurred over a four minute span, calculate

the rate of temperature change per minute? 2) Where should the chemicals from the beaker be placed after the reaction is

complete? Why? 3) If you used a stronger (more concentrated) solution of copper (II) chloride with

the same piece of aluminum foil, what changes in observations might you observe?

CONCLUSIONS: 1) Write out one significant qualitative and one quantitative observation from this lab. 2) What is the difference between these two types of observations? 3) Write out an interpretation of the data that you could take from this experiment. 4) State any sources of error which may have affected this lab.

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LAB EXPERIMENT 2 COOLING AND HEATING CURVES OF A PURE SUBSTANCE

OBJECTIVES: 1. To investigate the cooling process for liquid paradichlorobenzene. 2. To investigate the heating process for solid paradichlorobenzene. 3. To determine and compare the melting and freezing temperatures of

paradichlorobenzene. PREDICTIONS: Make sure you read through the lab and understand the objectives

before making your predictions. MATERIALS: Apparatus Reagents Figure 1 Ring stand and support paradichlorobenzene Test tube clamp Test tube (18 mm x 150 mm) 1000 mL beaker Thermometer PROCEDURES: PART 1: THE COOLING PROCESS 1. Obtain a test tube of paradichlorobenzene from your instructor. Remove the stopper

and save until the end of the experiment. 2. You will need to melt the paradichlorobenzene in order to place the thermometer into

the test tube and record the temperature. To do this, place the test tube clamp around the test tube and attach it to the ring stand, immersing the test tube in the hot water bath (It should be around 65-70°C) until all the substance has melted. Make sure the water level is above the level of the paradichlorobenzene.

3. Insert the thermometer into the test tube and record the temperature as the zero time

of your cooling curve in your data table. Remember to hold the thermometer off of the glass. See figure 1.

4. Quickly remove the test tube from the hot water bath and move over to the warm

water bath (It should be around 40-45°C). Attach your clamp to the ring stand and immerse the test tube in the water, again making sure the water level is higher than the paradichlorobenzene. Begin taking temperature readings and make observations every 30 seconds until all of the paradichlorobenzene has solidified. You will have to hold the thermometer until the substance has solidified around it. Keep taking the readings until the temperature has reached around 49°C. (The thermometer should be embedded in the solid)

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PART 2: THE HEATING PROCESS 1. Once your thermometer has been embedded in the solid, take a final temperature

reading. This will be your zero temperature reading for the heating curve data. 2. Remove the test tube from the warm water bath and move back to the hot water bath.

Immerse the test tube in it (making sure the water level is above the paradichlorobenzene). Begin taking temperature readings and observations every 30 seconds until all of the paradichlorobenzene has melted and a temperature of 60°C has been reached.

3. When the temperature has reached 60°C, remove the thermometer from the liquid

paradichlorobenzene and immediately wipe it with a paper towel. 4. Raise the test tube out of the hot water bath and allow the paradichlorobenzene to

resolidify. When it has cooled, put the stopper back on and return it to the instructor. 5. Before you leave the lab, wash your hands thoroughly with soap and water. OBSERVATIONS AND DATA: TABLE 1: Cooling and heating of paradichlorobenzene Cooling Process Heating Process Time (min)

Temperature (°C)

Observations Temperature (°C)

Observations

0 0.5 1.0 1.5 2.0 . . .

.

.

.

.

.

.

.

.

.

.

.

. ANALYSIS OF DATA: PART 1: THE COOLING PROCESS 1. Using the data you obtained during the cooling process, construct a graph of

temperature versus time. Use small circles for these data points, and sketch a smooth curve.

2. Indicate on the graph where solidification began and ended.

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PART 2: THE HEATING PROCESS 1. On the same graph as the cooling curve, plot temperature versus time for the

heating process. Use small squares for these data points, and sketch a smooth curve.

2. Indicate on the graph where melting began and ended. QUESTIONS: 1. From your cooling curve, determine the freezing point of paradichlorobenzene. 2. From your heating curve, determine the melting point of paradichlorobenzene. 3. Compare your freezing and melting points with those of two other lab groups and

explain any similarities or differences. 4. What can you conclude about the melting and freezing points of a pure substance? 5. How would you explain the plateaus in your heating and cooling curves? 6. Suppose more paradichlorobenzene had been used in Part 1. What would be the

appearance of the new cooling curve? CONCLUSIONS: Make sure to answer your objectives, summarize your results and tell what you have learned in this lab. Don’t forget any sources of error that may have occurred.

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LAB EXPERIMENT 3 CHEMICAL AND PHYSICAL CHANGE OBJECTIVES: 1. To observe some changes in the laboratory. 2. To infer whether each is a chemical or physical change. 3. To record some recognizable characteristics of chemical changes. PREDICTIONS: Make sure you read through the lab and understand the objectives

before making your predictions. MATERIALS: Figure 1 Apparatus Reagents 1 spot plate or glass square set of 4 unknown solutions PROCEDURES: 1. Obtain a clean spot plate or glass square. Place a piece of coloured paper underneath

the glass square if necessary. On the paper draw a grid as seen in Figure 1. If you use the spot plate, label the chemicals you will be combining above the well of the spot plate.

2. On the glass square or spot plate, combine the solutions by mixing A with B, A with

C, A with D, B with C, B with D, and C with D. Record your observations in your data table (there are only 6 observations necessary). Make sure you use separate, clean droppers for each of the solutions.

3. Once you have your data, rinse the solutions down the sink with lots of water. Before

you leave the lab, make sure you wash your hands thoroughly with soap and water. OBSERVATIONS AND DATA: Table 1 Unknown A B C D

A

B

C

D

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QUESTIONS: 1. State whether a physical or a chemical change occurred in each of the six

combinations of solutions. 2. Describe two chemical changes that you might observe occurring in everyday life. 3. Describe two physical changes that you might observe occurring in everyday life. CONCLUSIONS: Make sure to answer your objectives, summarize your results and tell what you have learned in this lab. Don’t forget any sources of error that may have occurred.

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LAB EXPERIMENT 4 SEPARATION OF A MIXTURE BY PAPER CHROMATOGRAPHY

PRE-LAB READING AND QUESTIONS

Chromatography is one technique used by chemists to separate mixtures of chemical compounds in order to identify or isolate their components. In chromatography, mixtures are separated according to the different solubilities of the components in liquids, or their adsorptions on solids. Chromatography has many applications, including detection and measurement of pesticides in foods, and drugs and urine specimens. It is also used extensively in biological research to separate alcohols, amino acids, and sugars. As well, the pharmaceutical industry relies on chromatography for the production of high-purity chemicals. There are a variety of chromatographic techniques, but all share two features: a moving carrier phase, and a stationary phase. In the stationary phase of paper chromatography, the sample to be analyzed is spotted onto a piece of chromatography paper. The sample is carried along this stationary phase by a solvent which acts as the moving carrier. The components of the sample are carried different distances along the paper, depending on their individual solubilities. After a length of time, the original spot is spread out into a series of bands. These bands are then analyzed to determine their identities. In paper chromatography, one method of identifying these separated components of a mixture is to calculate the Rf value of each. Rf stands for “ratio of fronts”. An Rf value is simply the ratio of the distance travelled by the solute to the distance travelled by the solvent: Rf = d1/d2 where d1 = distance travelled by solute d2 = distance travelled by solvent The Rf value of a substance is a characteristic of that substance for a specific solvent. A substance having a high solubility in the moving phase will be carried further and will have a higher Rf value. By definition, Rf values vary from 0 to 1. 1. On the basis of what principle is chromatography used to separate mixtures? 2. What industry uses chromatography to produce very pure chemicals? 3. a) What constitutes the moving phase in this experiment?

b) What constitutes the stationary phase in this experiment? 4. a) What does Rf mean?

b) Write the mathematical equation for calculating an Rf value and explain all the terms.

5. What substances and mixtures will be tested in this experiment? 6. Why do you need a pencil for this experiment?

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OBJECTIVES: 1. To assemble and operate a paper chromatography apparatus. 2. To study the meaning and significance of Rf values. 3. To test various food colourings and to calculate their Rf values. 4. To compare measured Rf values with standard Rf values. 5. To separate mixtures of food colourings into their components. 6. To identify the components of mixtures by means of their Rf values. PREDICTIONS: Make sure you read through the lab and understand the objectives

before making your predictions. MATERIALS: Apparatus Reagents Scissors Set of food colourings Ruler (yellow, red, blue, green) Pencil Unknown mixture Chromatography paper 3 large test tubes (20mm x 200mm) 3 Erlenmeyer flasks PROCEDURES:

PART 1: SETTING UP 1. Obtain 3 large test tubes and 3 Erlenmeyer flasks. Place a test tube in each of the

flasks and label the test tubes A, B, and C. 2. Obtain a 60 cm length of chromatography paper and cut it into 3 strips of 20 cm each.

Then cut each strip so that they are 2.0 cm in width. Using a pencil, lightly draw a line across each strip 4.0 cm from one end. Use scissors to trim this end of the strip into a point, as shown in the figure below. Label these A, B, and C with a pencil.

20.0 cm 2.0 sample dye spot cm

4.0 cm

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3. Your instructor will assign you one colour (red, yellow, or blue) for you to test. Using a capillary tube, spot your first strip of chromatography paper (labelled A) on the pencil line with the colour assigned (see the figure above). The spots should not exceed 0.5 cm in diameter.

4. For the strip labelled B, spot with the colour green. For the strip labelled C, spot with

the unknown mixture. PART 2: Rf VALUES OF INDIVIDUAL FOOD COLOURINGS AND SEPARATION

OF MIXTURES INTO THEIR COMPONENTS 1. Add about 2 cm of water to each of the test tubes. Put strip A into test tube A so that

the tip just touches the bottom of the test tube. Do not allow the dye spot to be immersed in the water. Do not allow the flat surface of the strip to rest against the walls of the test tube.

2. Do the same for strips B and C into test tubes B and C respectively. Observe what

happens to your sample as the water moves slowly up the paper as a result of capillary action.

3. See whether or not your samples separate into component colours. After about 30

minutes have elapsed, remove strip A from the test tube and immediately draw a pencil line across the top edge of the solvent front (the water line) and the top of the colour line.

4. Repeat Step 3 for strip B and strip C, making lines across all the component colours

and the water line. 5. Measure with a ruler the distances d1 (the distance the colour travelled) and d2 (the

distance the water travelled) on strip A. Record these results in Table 1 and calculate the Rf value. Record the class data in Table 2.

6. Repeat Step 5 for strip B, measuring all the component colour distances and

calculating Rf values for each component colour in the green food colouring. Record your results in Table 3.

7. Repeat Step 5 for strip C, measuring all the component colour distances and

calculating Rf values for each component colour in the unknown mixture. Record your results in Table 3.

8. Identify the dyes used in this lab by comparing your calculated Rf values with those in

Table 4. Record those results in Table 3. 9. Clean up your materials. The strips may be placed in the garbage.

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OBSERVATIONS AND DATA: TABLE 1: Assigned Food Colouring Colour Tested Solute Distance (cm) Solvent Distance (cm) Rf Value TABLE 2: Class Data for Individual Food Colourings Tested

Station Red Rf Yellow Rf Blue Rf 1 2 3 4 5 6 7 8 9 10 11 12

Average Rf TABLE 3: Separation of Mixtures Into Their Component Colours Component Colours d1 d2 Rf Dye ID

Green Colouring

Unknown Mixture

TABLE 4: Some Approved Dyes for Food Colouring

Dye Red #2 Red #3 Red #4 Yellow #5 Yellow #6 Blue #1 Blue #2 Rf 0.81 0.41 0.62 0.95 0.77 1.00 0.79

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QUESTIONS: 1. a) Which, if any, of the colours you tested in this experiment appeared to contain one

or more of the approved dyes listed in Table 4? b) Which, if any, of the colours you tested did not correspond to any of the approved

dyes? 2. From your results, what are the components of green food colouring? Support your

answer both qualitatively and quantitatively. 3. What can you conclude about the identity of the components in the unknown

mixture? What qualitative and quantitative evidence supports your answer? 4. What might happen if ink, rather than pencil, were used to mark the sample line on

the chromatography paper? 5. Why should green food colouring be classified as a mixture, whereas red, blue or

yellow should not? 6. Identify the dyes that appear on the chromatogram in Figure 1 (see Table 4). The

original sample was orange food colouring. 7. A pharmaceutical chemist runs a chromatography test on a substance and identifies

two of its components by comparing their Rf values. If the two components have Rf values of 1.00 and 0.41, and the solvent front has travelled 12.0 cm from the sample’s origin, what is the separation distance between the components on the chromatogram?

8. A chemist performs an Rf calculation, obtains a value of 1.20, and decides that the

answer is unacceptable. Why? Figure 1 CONCLUSION: Make sure to answer your objectives, summarize your results and tell what you have learned in this experiment. Don’t forget any sources of error that may have occurred.

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LAB EXPERIMENT 5 GRAPHING AS A MEANS OF SEEKING A RELATIONSHIP

OBJECTIVES: 1. To make measurements of mass and volume for three different liquids. 2. To analyze the data by means of graphing techniques. 3. To determine a mathematical relationship between mass and volume for each liquid. PREDICTIONS: Make sure you read through the lab and understand the objectives

before making your predictions. MATERIALS: Apparatus Reagents 250 mL Erlenmeyer flask water mass balance methanol several burettes salt (sodium chloride) solution PROCEDURE: 1. Your instructor will assign you a volume of liquid to test. You will then use this

volume for all three liquids. The data will be shared with the class and therefore you will be depending on each other for good results.

2. Determine and record the mass of a clean, dry 250 mL Erlenmeyer flask. 3. Go to one of the burettes that contains water, methanol, or salt solution (the order you

do these in is not important) and get your assigned volume of liquid as accurately as possible. If you do not get the precise amount it does not matter. What does matter is that you record exactly the volume that you did get in Table 1.

4. Mass out the total mass of the flask and the liquid. Subtract the mass you found in

Step 2 in order to determine the mass of the liquid. Record this in Table 1. 5. Now repeat Steps 3 and 4 for the other two liquids. Do not empty the flask each

time you add a different liquid – just keep determining the mass of each volume by subtracting the previous mass balance reading. Record your results in Table 1.

6. A data table similar to Table 2 will be on the chalkboard. Record your results from

Table 1 into the chart on the board. 7. Once all groups have finished, copy the completed Table 2 into your lab report. 8. Clean up all your materials and pour the contents of the flask down the sink with

plenty of water. Wash your hands with soap and water before leaving the lab.

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OBSERVATIONS AND DATA: Table 1: Results for Lab Station _____

Water Alcohol Salt Solution Volume (mL) Mass (g) Volume (mL) Mass (g) Volume (mL) Mass(g)

Table 2: Class Results Water Alcohol Salt Solution

Lab Station

Volume (mL)

Mass (g)

Volume (mL)

Mass (g)

Volume (mL)

Mass (g)

1 2 3 4 5 6 7 8 9 10 11 12

ANALYSIS OF DATA: 1. Following the rules of good graphing, plot a graph showing mass vs. volume for each

liquid using the class data. Plot the results for all three liquids on the same graph, making sure to differentiate between each liquid.

2. Draw best-fit lines for each graph and calculate the slope of each. Remember, slope is

calculated by: m = slope = ∆y (change in y values)

∆x ( change in x values)

Drawing ‘deltas’ on your graph and subtracting your values to find your change in y and your change in x does this.

3. Determine the mathematical relationship between the mass and volume for each

liquid. Remember: y = mx + b where y = the y variable

m = the slope of the line x = the x variable b = the y-intercept

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4. The slope of your graph is actually the density. Compare these slope values with the

accepted values and do a ‘percent error’ on them.

Percent Error = experimental value – accepted value X 100 accepted value

The actual (accepted) density of water is 1.00 g/mL. The actual density of methanol is 0.78 g/mL. The actual density of salt solution is 1.10 g/mL. QUESTIONS: 1. Use your graph to predict the mass of 6.5 mL of methanol. 2. Use your mathematical relationship to calculate the mass of 6.5 mL of methanol. 3. Compare your answers to Questions 1 and 2. Explain why they might not be identical. CONCLUSIONS: Make sure to answer your objectives, summarize your results and tell what you have learned in this experiment. Don’t forget any sources of error that may have occurred.

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LAB EXPERIMENT 6 INVESTIGATING MASS CHANGES IN CHEMICAL REACTIONS

OBJECTIVES: 1. To observe a chemical reaction in a sealed flask. 2. To determine the change in mass that occurs during a chemical reaction. PREDICTIONS: Make sure you read through the lab and understand the objectives

before making your predictions. MATERIALS: Apparatus Reagents 250 mLErlenmeyer flask One of the following pairs of solutions: rubber stopper for flask 1. A. barium chloride 2 test tubes (13 mm × 100 mm) B. sodium sulphate tongs 2. A. lead acetate mass balance B. potassium iodide 3. A. iron (III) nitrate

B. potassium thiocyanate 4. A. calcium chloride

B. sodium carbonate PROCEDURES: 1. Get 2 test tubes and label them A and B. 2. Your instructor will assign you a set of solutions to react. Half fill test tube A with

solution A. Half fill test tube B with solution B. 3. Pour solution B into the Erlenmeyer flask. 4. Carefully lower test tube A into the flask using a pair of tongs so that you do not spill

any into solution B. Place the rubber stopper on the flask. 5. Determine the mass of your assembled apparatus and record this value in Table 1. 6. After making certain that the stopper is secure, gently turn the flask upside down and

mix the chemicals. 7. Determine the mass of your apparatus after you have mixed the chemicals and record

this value in Table 1. 8. Calculate the mass gained or lost during the reaction and record this in Table 1.

Record your results in Table 2 on the blackboard and into your lab report. You will be analyzing the class data as a whole when answering the questions.

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9. Dispose of the reacted chemicals in the chemical waste disposal container. Make sure

to wash your hands with soap and water before leaving the lab. OBSERVATIONS AND DATA: Table 1: Results for Lab Station ______. Identity of solution A _________________________ Identity of solution B _________________________ Mass of apparatus and contents before reaction (g) Mass of apparatus and contents after reaction (g) Change in mass as a result of reaction (+ or -) (g) Table 2: Class Results Lab Station Solution A Solution B Mass Change (g)

1 2 3 . . .

QUESTIONS:

1. Why is it important that the flask be sealed for this experiment, even after the flask is returned to an upright position after mixing?

2. What observations lead you to believe that a chemical reaction occurred in the

flask?

3. In general, what overall mass change results from a chemical reaction?

4. Suppose that a reaction was carried out in an open flask and the final mass was significantly greater than the initial mass. What would you conclude?

5. If a reaction was carried out in an open flask and the final mass was significantly

less than the initial mass, what would you conclude?

6. Write balanced chemical equations to show what happened in each reaction. Assume all reactions are double replacement reactions.

CONCLUSIONS: Make sure to answer your objectives, summarize your results, state what you have learned and discuss any sources of error which may have affected your results.

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LAB EXPERIMENT 7 FINDING MASS AND COUNTING

PARTICLES

This lab illustrates how you can find relative masses of particles. Remember a particle can contain a single atom such as an atom of sodium or a particle can be thought of as a single molecule such as a molecule of water (made up of atoms). The particles you will use today are paper clips. Part B illustrates how the relative masses found in Part A can be used to count out equal numbers of particles. (Note: This is done without counting out paper clips in the beginning.)

PREDICTIONS: Make sure you read through the lab and understand the objectives

before making your predictions.

MATERIALS:

1 box of small paper clips 1 box of large paper clips

PROCEDURE:

PART A: FINDING RELATIVE MASSES OF PAPER CLIPS

1. Obtain a palm full of large paper clips and a palm full of small paper clips.

2. Clip one large paper clip to one small paper clip and set them aside.

3. Continue pairing one large with one small paper clip, putting them with the other pairs, until all of one size of paper clip is used up. Set the unused ones aside.

4. Now separate the pairs of clips and put the large clips in one pile and the small clips in another pile.

5. Find the mass of large clips and record. Find the mass of the small clips and record.

6. Using the data from step 5, calculate the relative mass of large paper clips to one gram of small paper clips. Show this calculation as a ratio or a fraction.

PART B: COUNTING PAPER CLIPS BY MEASURING MASS

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7. Using the results from Part A, predict the mass of large clips that would be necessary to pair up one for one with 10 grams of small clips.

8. Weigh out 10 grams of small clips and your predicted mass of large clips and see if they pair up one for one.

9. If the number of large clips did not pair up evenly with the number of small clips in step 8 (within ± 1 clip), recalculate step 7 and repeat step 8.

QUESTIONS: 1. Show the calculations for step A-6 and B-7. 2. What mass of large clips would be necessary to pair up one for one with 50.0 g of

small clips? 3. What mass of small clips would be necessary to contain the same number of clips

as in 8000 g of large clips? 4. What mass of large clips would be necessary to contain the same number of clips

as in 300 kg of small clips? 5. If you were to clip two small clips to each large clip, what mass of small clips

would be necessary if you had 70.0 g of large clips? 6. Based on question 5, what would the new relative mass ratio be (mass of large

paper clips to one gram of small paper clips)? 7. Based on your answer to question 6, what would the mass of large clips be if you

had 12.0 g of small clips? 8. If you can, mass out 12.0 g of small clips and the mass of large clips calculated in

question 7. What should be true about the numbers of large clips and the number of small clips present?

CONCLUSION:

Answer your objectives, state what you have learned, and list any sources of error which may have affected your results.

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LAB EXPERIMENT 8 MOLES OF IRON AND COPPER OBJECTIVES: 1. To determine the number of moles of copper produced in the reaction of iron and

copper (II) chloride. 2. To determine the number of moles of iron used up in the reaction of iron and copper

(II) chloride. 3. To determine the ratio of moles of iron to moles of copper. 4. To determine the number of atoms and formula units involved in the reaction. PREDICTIONS: Make sure you read through the lab and understand the objectives

before making your predictions. MATERIALS: Apparatus Reagents 250 mL beakers copper (II) chloride wash bottle 2 iron nails stirring rod 1 M hydrochloric acid mass balance distilled water steel wool tongs PROCEDURE: DAY 1: 1. Label a clean, dry 250 mL beaker with your name. Find the mass of the beaker and

record it in your data table. 2 Add approximately 8 g of copper (II) chloride crystals to the beaker. Find the mass

and record it in your data table. 3 Add 50 mL of distilled water to the beaker. Stir the solution with a stir rod until all

the crystals have dissolved. 4 Obtain 2 nails and clean them with a piece of steel wool until the surface of the nails

is shiny. Find the mass of the clean nails and record it in your data table. 5 Place the nails into the solution and leave them undisturbed for 20 to 30 minutes.

During this time copper should be forming in the solution.

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6 Once the time has elapsed, use tongs to carefully pick up one of the nails and use the

wash bottle to rinse off any copper from the nail. You may have to use the stir rod to scrape any excess copper off. Place the nail to dry on a piece of paper towel. Repeat this step for the other nail.

7. Carefully decant the liquid from the solid into another beaker. Decant means to

pour off only the liquid from a container that contains both a solid and a liquid. Use your stir rod for this process (see figure 1). If you lose some solid, you need to start decanting it all again.

Figure 1

8. Rinse the solid with about 10 mL of distilled water. Decant. Repeat this process four more times.

9. Wash the solid with 25 mL of hydrochloric acid, decant, and then wash one last

time with 25 mL of distilled water. Decant one last time.

10. Place the beaker in the drying oven until next day.

11. Once the nails are completely dry, find the mass of them and record it in your data table.

12. Pour the decanted solution down the sink with lots of water and clean up. Wash

your hands before leaving the lab. DAY 2: 1. Get your dry beaker with the copper in it from the drying oven. 2. Find the mass of the beaker and the copper and record it in your data table. 3. Place the copper in the garbage and clean up your beaker. Wash your hands before

leaving the lab. OBSERVATIONS AND DATA: Table 1: Before the Reaction Mass of empty, dry beaker (g) Mass of beaker and copper (II) chloride (g) Mass of two iron nails

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Table 2: After the Reaction Mass of two iron nails (g) Mass of beaker and dry copper (g) QUESTIONS: 1. Find the following masses:

a) mass of iron used in the reaction b) mass of copper (II) chloride used c) mass of copper produced in the reaction

2. Find the number of moles of the following:

a) moles of iron used b) moles of copper produced

3. Find the number of atoms of each substance involved in the reaction.

a) atoms of iron used b) atoms of copper produced

4. Calculate the ratio of moles of copper produced to moles of iron used. 5. Suppose you have an unlimited supply of copper (II) chloride to react with iron. How

many moles of copper would be produced by reacting 34.0 g of iron with copper (II) chloride solution?

6. How many moles of iron would have been used up if 45.0 g of copper were to be

produced? 7. How many atoms of copper would be involved in question 6? 8. How many atoms of iron would be involved in question 6? 9. How many grams of copper would be produced from the reaction of 456 g of iron? CONCLUSIONS: Make sure to answer your objectives, summarize your results and tell what you have learned in this experiment. Don’t forget any sources of error that may have occurred.

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LAB EXPERIMENT 9 CALCULATIONS WITH A CHEMICAL REACTION

OBJECTIVES: 1. To observe the reaction between solution of calcium chloride and sodium carbonate,

forming insoluble calcium carbonate. 2. To calculate the number of moles of each of the starting materials present in the

solution. 3. To determine the reactant that is in excess. 4. To determine the theoretical amount of calcium carbonate that could be produced. 5. To compare the theoretical amount to the actual amount of calcium carbonate and

calculate the percent yield. PREDICTIONS: Make sure you read through the lab and understand the objectives

before making your predictions. MATERIALS: Apparatus Reagents 2 - 250 mL beakers 0.60 M sodium carbonate, Na2CO3

wash bottle 0.40 M calcium chloride, CaCl2 stirring rod with rubber scraper distilled water mass balance 2 - 100 mL graduated cylinders ring stand with funnel holder funnel filter paper PROCEDURES: DAY 1: 1. Pour approximately 75 mL of sodium carbonate into a graduated cylinder and record

the exact volume in your data table. Pour the sodium carbonate into a clean, dry 250 mL beaker.

2. Pour approximately 50 mL of calcium chloride into the other graduated cylinder,

again recording the exact amount in your data table. 3. Pour the calcium chloride into the beaker and describe the resulting reaction in your

data table. Stir the contents for about one minute.

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4. Obtain a piece of filter paper and label it with your name in pencil. Find the mass of

the filter paper and record it in your data table. 5. Set up your ring stand with the funnel holder. Place the funnel in the holder. Fold

your filter paper into a funnel shape and place it in the funnel. Wet the filter paper with a small amount of distilled water to hold it in the funnel (your instructor may show you how).

6. Stir your solution and begin carefully pouring it into the filter paper. Be careful so

that none of the solid flows out of the filter paper or funnel. Continue filtering until you get as much of the solid out of the beaker as possible (you may have to use the rubber scraper).

7. Rinse the inside of the beaker with some distilled water and filter. Repeat this two or

three more times to remove as much solid as possible. 8 Once all the solid is on the filter paper and the liquid has all drained through into the

beaker, carefully remove the filter paper from the funnel and unfold it onto a piece of paper towel.

9. Place the filter paper and paper towel in the drying oven until next day. 10. Clean up your apparatus and pour the filtered solution down the sink with plenty of

water. Wash your hands with soap and water before leaving the lab. DAY 2: 1. Get your filter paper with your solid on it from the drying oven. 2. Find the mass of the filter paper and the solid and record it in your data table. 3. Place the solid in the garbage and clean up. Wash your hands before leaving the lab. OBSERVATIONS AND DATA: Table 1: Volume of sodium carbonate solution (mL) Volume of calcium chloride solution (mL) Describe what happens after mixing the solutions

Mass of dry filter paper (g) Mass of filter + dry solid (g)

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QUESTIONS: 1. Calculate the following values:

d) moles of sodium carbonate used e) moles of calcium chloride used f) mass of calcium carbonate produced g) moles of calcium carbonate produced

2. Write a balanced chemical reaction for the equation that you observed in this lab. 3. Determine which of the reactants was in excess in this reaction. 4. Calculate the amount of calcium carbonate that should theoretically form from the

amount of the limiting reactant. 5. Calculate the percent yield in your reaction. 6. Suppose you wanted to add just enough 0.40 M calcium chloride solution to 75.0 mL

of 0.60 M sodium carbonate solution for all of the solutions to react.

a) What volume of 0.40 M calcium chloride solution would be required? b) What mass of calcium carbonate would be produced assuming 100% yield? c) What mass of sodium chloride would be produced assuming 100% yield?

7. How would you be able to recover the sodium chloride from the solution? 8. A similar reaction occurs when barium chloride is mixed with sodium carbonate

solution. Write a balanced chemical equation for the reaction. 9. Calculate each of the following for the reaction that occurs when 56.0 mL of 0.50 M

barium chloride is mixed with 78.0 mL of 0.75 M sodium carbonate solution:

a) moles of barium chloride added b) moles of sodium carbonate added c) mass of barium carbonate produced, if the yield is 78.0%

CONCLUSIONS: Make sure to answer your objectives, summarize your results and tell what you have learned in this experiment. Don’t forget any sources of error that may have occurred.

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LAB EXPERIMENT 10 THE PERIODIC TABLE

OBJECTIVES: 1. To understand the relationship between electron configuration and the location of an

element within the period. 2. To examine and graph periodic trends in atomic radii and the first ionization energies. 3. To construct the periodic table published by Mendeleev in 1871 according to a list of

clues and your knowledge of the modern periodic table. PREDICTIONS: Make sure you read through the lab and understand the objectives

before making your predictions. MATERIALS: Apparatus Graph paper Blank periodic table sheet for electron configurations Blank 1871 version of the periodic table Reference book (Heath Chemistry) PROCEDURES: PART 1: ELECTRON CONFIGURATION AND THE PERIODIC TABLE 1. Obtain blank periodic table for the first 4 rows. 2. For each of the elements listed, write the electron configurations in the space

provided. PART 2: TRENDS OF THE PERIODIC TABLE 1. Obtain 2 pieces of graph paper. 2. Using the data listed in Table 1, construct a graph of atomic radius vs. atomic number

on the first piece of graph paper. 3. Using the data listed in Table 1 and the second piece of graph paper, construct a

graph of first ionization energy vs. atomic number.

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Table 1:

ELEMENT ATOMIC NUMBER

ATOMIC RADIUS (nm)

FIRST IONIZATION ENERGY (kJ/mol)

Hydrogen Helium Lithium

Beryllium Boron Carbon

Nitrogen Oxygen Fluorine

Neon Sodium

Magnesium Aluminum

Silicon Phosphorous

Sulphur Chlorine Argon

Potassium Calcium

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20

0.0357 0.050 0.152 0.111 0.088 0.077 0.070 0.066 0.064 0.070 0.186 0.160 0.143 0.117 0.110 0.104 0.099 0.094 0.231 0.197

1312 2372 519 900 799 1088 1406 1314 1682 2080 498 736 577 787 1063 1000 1255 1519 418 590

PART 3: MENDELEEV’S PERIODIC TABLE 1. Obtain a blank 1871 version of the periodic table. 2. Using the clues below, place the name of the 67 known elements of 1871 in the

correct space. You will need to use the reference book provided by your instructor. 1A has a single electron in the 1s sublevel 1B derived its name from the Latin word for stone, lithos 1C can be collected as a silver liquid in the electrolysis of salt 1D has a first ionization energy of 418 kJ/mol 1E has a density of 8.96 g/cm3 1F is the first alkali metal with a completed 3d sublevel 1G was originally identified by its Latin name, argentum 1H is an alkali metal located in period 6 of the modern periodic table 1I derived its name from the Latin word for dawn, aurora 2A was used as a target substance in the experiments by Irene Joliot-Curie 2B is an alkaline earth metal located in period 3 2C is the metallic component of the substance limestone 2D is a transition metal with 30 protons 2E is represented by the symbol Sr

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2F possesses a nuclear charge of +48 2G is an alkaline earth metal found in period 6 of the modern periodic table 2H is a liquid metal, originally called hydrargyrum 3A has a single electron in the 2p sublevel 3B is a lightweight metal with a molar mass of 27.0 g/mol 3C is a transition metal with an atomic number of 39 3D is a metal represented by the symbol In 3E is the first metal of the lanthanide series 3F is found in group 13 and period 6 of the modern periodic table 4A is the element whose common isotopic form is the basis of the atomic mass unit 4B has a second ionization energy of 1577 kJ/mol 4C is located between scandium and vanadium on the modern periodic table 4D is represented by the symbol Zr 4E derived its symbol from the Latin word stannum 4F is a very dense metal with an atomic mass of 207.2 g/mol 4G is the second member of the actinide series 5A is the most abundant element in the atmosphere 5B has 3 electrons in its 3p sublevel 5C is a byproduct of fossil fuel oxidation and represented by the symbol V 5D is a period 4 nonmetal known since 1250 5E is a member of both group 5 and period 5 of the modern periodic table 5F was originally called stibium 5G is located in period 6 of the modern periodic table, below niobium 5H is a metal with atomic number 83 6A is the most abundant element in the Earth’s crust 6B is a member of group 16, known during the time of the Roman Empire 6C is the first member of group 6 on the modern periodic table 6D is represented by the symbol Se 6E has 42 protons within its nucleus 6F is a halogen whose crystals sublime 6G was originally called wolfram 6H is the fourth member of the actinide series 7A has an atomic radius of 0.099 nm 7B is represented by the symbol Mn 7C forms a diatomic gas with the molar mass of 159.8 g/mol 7D has a molar mass of 127.6 g/mol 8A is a group 8 metal known during the time of the Roman Empire 8B is an element named after the German word for Satan 8C has an average atomic mass of 58.9 g/mol 8D is located between iron and osmium on the modern periodic table 8E has a nuclear charge of +45 8F has an atomic mass of 106.4 g/mol 8G has 114 neutrons within its nucleus 8H is named after the Latin word for rainbow, iris 8I is an inert metal often used in electrodes and has an atomic number of 78

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QUESTIONS: 1. Examine the placement of electron configurations on your periodic table. What

relationship can be seen in an element’s placement within a group and its electron configuration?

2. Examine the graph of atomic radius vs. atomic number. What trend do you observe as

you go across a period? 3. Which group appears to have members with the largest atomic radii? Which group

has the smallest? 4. Examine the graph of ionization energy vs. atomic number. What trend do you

observe as you go across a period? 5. No members of group 18 of the modern periodic table can be found on Mendeleev’s

classification chart. Suggest a reason for their absence. 6. What factor may account for the observed trend in atomic radii as one proceeds

across a period? CONCLUSIONS: Make sure to answer your objectives, summarize your results, and state what you have learned.

30

LAB EXPERIMENT 11 POLAR AND NON-POLAR SOLUTES AND SOLVENTS

OBJECTIVES: 1. To determine the type of solvent that generally dissolves ionic compounds. 2. To determine the type of solvent that generally dissolves polar covalent compounds. 3. To determine the type of solvent that generally dissolves non-polar covalent

compounds. 4. To investigate the effect of adding a polar liquid solute to a non-polar liquid solvent. PREDICTIONS: Make sure you read through the lab and understand the objectives

before making your predictions. MATERIALS: Apparatus Reagents Tweezers sodium chloride crystals Test tube rack sucrose crystals 6 test tubes (13 mm x 100 mm) iodine crystals 6 stoppers for test tubes 3 unknown solid solutes paint thinner glycerol PROCEDURES: PART 1: Solubility Tests on Known Solutes 1. Obtain 6 clean, dry test tubes and place them in a test-tube rack so that you have two

rows of 3 test tubes each. 2. Half fill one set of 3 test tubes with room temperature water and half fill the other 3

test tubes with paint thinner. 3. Into the first pair of test tubes (one with water and one with paint thinner) add

enough salt crystals to just cover the bottom of the test tubes. 4. Stopper the test tubes and hold it on with your thumb while you shake each test tube

to see if each substance dissolves. When you are sure no further changes are taking place, record your observations in Table 1. Dispose of the chemicals in the waste disposal jar.

5. Repeat Steps 2, 3 and 4 with crystals of sugar in each. Dispose of the chemicals in

the waste disposal jar. 6. Repeat Steps 2, 3 and 4 with a single crystal of iodine in each. Dispose of the

chemicals in the waste disposal jar.

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PART 2: Solubility Tests on Unknown Solutes 1. Repeat Part 1 of this experiment, but refer to Table 2 and test the solubilities of

unknown solutes A, B, and C. PART 3: Mixing Two Liquids 1. Fill a clean test tube one quarter full with paint thinner, then add twice as much water

to the same test tube. 2. Stopper the test tube and shake the test tube. Examine what happens to the liquids

after shaking, and record your observations in Table 3. 3. Add one iodine crystal to the test tube, stopper and shake. Make a sketch of the test

tube and its contents in your observations. 4. In a second test tube, fill it one quarter full with glycerol. It is a polar liquid. Add

twice as much water, stopper and shake. Record your observations in Table 3 and make a sketch in your observations.

5. Dispose of all chemicals in the waste disposal jar. Clean up all your apparatus and

make sure you wash your hands with soap and water before leaving the lab. OBSERVATIONS AND DATA: Table 1: Known Solutes with Known Solvents

Solutes Solvents

Salt

(NaCl – Ionic) Sugar

(C12H22O11 – Polar Covalent)

Iodine (Non – Polar Covalent)

Water (Polar – Covalent)

Paint Thinner (Non – Polar Covalent)

Table 2: Unknown Solutes with Known Solvents

Solutes Solvents A B C

Water (Polar – Covalent)

Paint Thinner (Non – Polar Covalent)

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Table 3: Liquid Combinations Combinations of Liquids Covalent Types Results Water and Paint Thinner

Water and Glycerol QUESTIONS: 1. a. What general trend appears in Table 1 with regard to which type of solute

dissolves in which type of solvent? b. This general solubility trend is sometimes expressed as “Like dissolves like.

Explain

2. a. Attempt to classify each of the unknown solutes from Part 2 as ionic, polar covalent, or non – polar covalent. b. What problem do you encounter in making this classification? c. How could you overcome this problem?

3. Explain the terms “miscible” and “immiscible” and use these terms to explain the

results from Part 3. 4. How did the addition of an iodine crystal help in identifying the layers of liquids in

the water – paint thinner combination? 5. Explain how many layers you would expect to see if water, paint thinner, and

glycerol were combined in one test tube. 6. Explain which solvent from this experiment you would use to remove road salt

stains from a pair of jeans. CONCLUSIONS: Make sure to answer your objectives, summarize your results and tell what you have learned in this lab. Don’t forget any sources of error that may have occurred.

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LAB EXPERIMENT 12 ACID-BASE TITRATION OBJECTIVES: 1. To titrate a hydrochloric acid solution of unknown concentration with standardized

sodium hydroxide solution. 2. To utilize the titration data to calculate the molarity of the hydrochloric acid solution. PREDICTIONS: Make sure you read through the lab and understand the objectives

before making your predictions. MATERIALS: Apparatus Reagents 250 mL Erlenmeyer flask standardized sodium hydroxide solution (approx. 0.50 M) volumetric pipette (10 mL) unknown hydrochloric acid solution burette phenolphthalein solution burette clamp and stand pipette bulb PROCEDURE: 1. Obtain about 50 mL of the unknown hydrochloric acid solution and about 50 mL of

the standardized sodium hydroxide solution. Your instructor will provide you with the exact molarity of the NaOH. Record this value in Table 1.

2. Using the pipette bulb and volumetric pipette, rinse out the pipette and then transfer

10.0 mL of the HCl solution into a 250 mL Erlenmeyer flask. 3. Add 3 drops of phenolphthalein solution.

4. Rinse out a burette with some of the standardized NaOH solution and drain it,

keeping the tip filled with the NaOH. Refill the burette and record your initial volume in Table 1.

5. Gradually add the NaOH into the Erlenmeyer flask containing the HCl. Swirl the

flask continuously. Continue adding NaOH, noting any changes in the flask.

6. As the equivalence point is reached, a pinkish colour will appear. As you approach this point, the colour will take longer to disappear as you swirl it. Add the NaOH drop by drop until the colour no longer disappears. Record the final volume of the burette in Table 1. The most accurate reading is one in which the solution is a very faint pink.

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7. Repeat steps 2 through 6 after cleaning out your Erlenmeyer flask (the contents can be poured down the sink with lots of water). Continue repeating until you get two trials that are within ± 0.10 mL of one another.

8. Clean up all your apparatus and make sure you wash your hands with soap and water

before leaving the lab. OBSERVATIONS AND DATA: Table 1: Volume of NaOH Needed to Neutralize 10.00 mL of Unknown HCl Molarity of NaOH =_______________

Trial 1 Trial 2 Trial 3 (if necessary)

Initial Volume of NaOH Final Volume of NaOH Volume of NaOH used Average Volume of NaOH QUESTIONS: 1. Calculate moles of NaOH from the average volume used, and the given molarity. 2. Calculate moles of HCl present originally. 3. Calculate the molarity of the HCl solution. 4. What was the reason for cleaning out the burette with NaOH before starting the

titration? 5. What is the concentration of a NaOH solution when it requires 30.0 mL of 0.50 M

HCl to neutralize 50.0 mL of the NaOH? 6. What is the concentration of acetic acid (HCH3COO) when 32.5 mL of 0.56 M NaOH

is required to neutralize 15.0 mL of the acid? 7. When 6.25 g of oxalic acid (H2C2O4) are placed in water containing phenol red, the

suspension is yellow. What is the molarity of the KOH if 32.2 mL of KOH is required to change the colour to red?

8. A 5.0 g tablet of TUMS (Mg(OH)2) neutralizes 450 mL of stomach acid (HCl). What

is the molarity of the stomach acid? CONCLUSIONS: Make sure to answer your objectives, summarize your results and tell what you have learned in this experiment. Don’t forget any sources of error that may have occurred.

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LAB EXPERIMENT 13 MOLECULAR MODEL BUILDING OBJECTIVES: 1. To represent molecular structures with electron dot and structural (line) diagrams. 2. To construct molecular models of simple substances. 3. To construct molecular models illustrating the different types of isomers. PREDICTIONS: Make sure you read through the lab and understand the objectives

before making your predictions. MATERIALS: Apparatus Molecular model kit PROCEDURE: 1. Construct models for each of the following alkanes. In your lab, draw the structural

diagram and the electron dot diagrams for each of these molecules.

a. methane, CH4 b. ethane, C2H6 c. propane, C3H8

2. Construct models for all structural isomers for each of the following compounds.

Draw the structural diagram, and name each one.

a. butane, C4H10 (there are 2 of them) b. pentane, C5H12 (there are 3 of them) c. hexane, C6H14 (there are 5 of them) d. cyclohexane, C6H12 (there is only 1) e. methylhexane, C7H16 (there are 2 of them)

3. Construct models for all structural isomers of butene, C4H8. Draw the structural

diagrams for each (there are 4 of them) and then name each one. Note that 2 of them are geometric isomers.

4. Construct models for all structural isomers of propyne (there is only 1), C3H4, and

butyne (there are 2 of them), C4H6. Draw the structural diagrams for each, and then name each one.

5. Construct models for all structural isomers of n-hexanol, C6H13OH. Draw the

structural diagrams for each (there are 3 of them) and then name each one.

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6. Construct models for the following structural isomers. Compare their structures and

the placement of the oxygen atom. Draw the structural diagram for each isomer.

a. ethanol, C2H5OH b. dimethyl ether, CH3OCH3

QUESTIONS: 1. Can cyclopentane (C5H10) be considered an isomer of pentane? Explain. 2. What is the difference between a cis and trans isomer? 3. Explain why there are such a huge number of organic compounds. 4. How many isomers are possible for the following structures?

a. methane b. ethane c. propane d. butane e. pentane f. hexane


37

LAB EXPERIMENT 14

FACTORS AFFECTING REACTION RATES OBJECTIVES: 1. To observe and record the effect of reactant concentration on the reaction rate. 2. To observe and record the effect of reactant surface area on the reaction rate. 3. To observe and record the effect of reactant temperature on the reaction rate. MATERIALS: Apparatus Reagents 11 - 100 ml beakers 0.5 M hydrochloric acid Thermometers 1.0 M hydrochloric acid Mass balance 3.0 M hydrochloric acid Hot plate 6.0 M hydrochloric acid Stop watch Magnesium ribbon Scissors Calcium carbonate pieces Ruler Mortar and pestle 4 – 250 mL beakers PROCEDURES: ● You will work with 2 other lab groups, each group performing one part of

the lab and then sharing the information with the other groups. PART 1: Effect of Concentration on Reaction Rate 1. Obtain an 8 cm piece of magnesium ribbon from the roll. Use steel wool to

remove any coating from the outside of the metal and then record its mass in Table 1. Divide the total mass by 8 to find the average mass of a 1 cm strip and record this value in your table.

2. Cut the strip in half and give one half to the lab group performing Part III. 3 Cut the remaining strip into 1 cm lengths 4. Label four 100 mL beakers with the appropriate concentration of acid to be added

(0.5 M, 1.0 M, 3.0 M, 6.0 M) 5. Carefully measure 15 mL of each acid and place them in the appropriate beakers.

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6. Obtain a stopwatch. When ready, start the stop watch and place a 1 cm strip of

magnesium into each of the beakers all at the same time. As each piece of magnesium disappears, take note of the time and record your results in Table 1.

7. Once all reactions have been completed, dispose of the chemicals into the waste

disposal jars and clean up. PART II: Effect of Surface Area on Reaction Rate 1. Select three small chunks of calcium carbonate as equal in size as possible and

record the total mass of all three chunks in Table 2. Divide the total mass by 3 to get the average mass of each piece and record it in the table.

2. Put aside one chunk to use as is. 3. Take the second chunk and crush it just a bit into six or seven smaller pieces with

the mortar and pestle and keep all the pieces together. 4. Take the third chunk and crush it into a fine powder and keep this all together. 5. Obtain three 100 mL beakers and put 20 mL of 6.0 M hydrochloric acid into each. 6. Obtain a stopwatch. When ready, start the stop watch and place each trial into a

separate beaker all at the same time. When the reaction is complete (either all the solid has disappeared or it has stopped bubbling) note the time and record the results in Table 2.

7. Once the reactions are complete, dispose of the chemicals into the waste disposal

jars and clean up. PART III: Effect of Temperature on Reaction Rate 1. Obtain a 4 cm strip of magnesium from the lab group performing Part I. 2. Cut the length into 1 cm strips. 3. Label four 250 mL beakers A, B, C, and D and add about 150 mL of water to

each. These will be your water baths needed to get your chemicals to the proper temperature before you perform each trial.

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4. The temperatures for the water baths are as follows:

Beaker A – 5 ºC (use ice if needed) - record exact value! Beaker B – Room temperature - record exact value! Beaker C – 50 ºC (use hot plate) - record exact value! Beaker D – 80 ºC (use hot plate) - record exact value!

5. Place 20 mL of 3.0 M hydrochloric acid into each of four 100 mL beakers. Once the water baths are ready, place these beakers into the water baths for about 5 minutes to get the chemicals to the assigned temperatures. Record these temperatures in Table 3.

6. Obtain a stop watch. When ready, start the stop watch and place a 1 cm piece of

magnesium into each of the beakers containing acid all at the same time. As each piece of magnesium disappears, note the time and record the results in Table 3.

7. Once the reactions are complete, dispose of the chemicals in the waste disposal

jars and clean up. DATA AND OBSERVATIONS: Table 1 Mass of 8 cm strip of Mg: _______________________ Average mass of 1 cm strip of Mg: _____________________ Concentration of Acid (M) Reaction Time (s) Reaction Rate (g Mg/s)

0.5 1.0 3.0 6.0

Table 2 Mass of 3 chunks of CaCO3: ____________________ Mass of one chunk of CaCO3: ____________________

Type of CaCO3 Reaction Time (s) Reaction Rate ( g CaCO3/s) Beaker A – rock Beaker B - pieces Beaker C - powder

40

Table 3

Temperature (ºC) Reaction Time (s) Reaction Rate (g Mg/s) Beaker A – Beaker B – Beaker C – Beaker D – QUESTIONS AND CALCULATIONS: 1. On graph paper, graph the reaction rate vs. concentration of HCl (Part I). On a second

piece of graph paper, graph the reaction rate vs. the temperature (Part III). 2. Complete the following statements and explain why.

a) As concentration increases, reaction rate … b) As surface area increases, reaction rate … c) As temperature increases, reaction rate …

3. Look at your data from Part I. Does doubling the concentration of HCl double the

reaction rate? Explain. 4. Look at your data from Part II. Which reaction rate was the fastest? Explain. 5. Look at your data from Part III. Which reaction has the slowest rate? Explain. 6. Use your graphs to predict the reaction rates and then calculate the reaction times for

each of the following:

a) reaction of 1 cm magnesium strip with 4.0 M HCl b) reaction of 1 cm magnesium strip with 3.0 M HCl and a temperature of 15 ºC

7. Does doubling the temperature double the rate of the reaction in Part III? Explain. 8. What effect would the addition of a catalyst have on a reaction? 9. Explain why you will be more successful lighting a fire made from kindling wood

rather than lighting a log directly. 10. Explain why blowing on a smoldering fire may make it burn better. CONCLUSIONS: Make sure to answer your objectives, summarize your results and tell what you have learned in this experiment. Don’t forget any sources of error that may have occurred.

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LAB EXPERIMENT 15

MEASURING REACTION RATES USING VOLUMES OF GAS PRODUCED

OBJECTIVES: 1. To measure the volume of a gas produced from a reaction mixture at regular time

intervals during the reaction. 2. To interpret the results and obtain the overall rate of reaction. 3. To observe how the rate changes at different temperatures and concentrations. MATERIALS: Apparatus 250 mL Erlenmeyer flask Trough Thermometer 50 mL glass measuring tube Stopper with glass tubing and hose 25 mL and 10 mL graduated cylinders Hot plate Ice Stop watch Burette, stand and clamp Reagents sodium hypochlorite solution (bleach – 5.25%) 0.10 M cobalt (II) nitrate, Co(NO3)2 PROCEDURES: 1. Refer to Figure 1 to help understand how to set up the apparatus. Figure 1

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2. Half fill the trough with water. Fill the glass measuring tube with water, and invert it into the trough without letting any water come out. Hold it in the vertical position with the burette clamp attached to the stand.

3 Place the hose end of the rubber stopper apparatus into the gas measuring tube. 4. Measure out 15 mL of bleach solution into a 25 mL graduated cylinder and then

pour it into the Erlenmeyer flask. 5. Measure out 5 mL of 0.10 M cobalt (II) nitrate solution into the 10 mL graduated

cylinder. 6. Obtain a stopwatch. When ready, pour the cobalt nitrate solution into the

Erlenmeyer, immediately attach the rubber stopper and start the timer. Swirl the flask continuously until the end of the reaction.

7. Record the total volume of oxygen that has collected in the gas measurement tube

at 30 second intervals until a volume of 50 Ml is attained. Record the exact time 50 mL of oxygen gas has been reached.

8. Pour the chemicals into the waste disposal and clean the Erlenmeyer flask. 9. Repeat steps 2-8, but have the reactants at a temperature 10°C above room

temperature (record the exact value). You can get this ready by placing your chemicals in a water bath for about 10 minutes on a hot plate, then mixing the chemicals, and swirling your flask inside the water bath.

10. Repeat steps 2-8 but bring the reactants to a temperature 10°C below room

temperature (record the exact value) by using ice. 11. Repeat steps 2-8 at room temperature but add 20 mL of water to the bleach

solution before mixing, so the overall concentration is one half its original concentration.

12. Repeat steps 2-8 at room temperature but add 60 mL of water to the bleach

solution before mixing, so the overall concentration is on quarter its original concentration.

13. Clean up all equipment and wash your hands thoroughly before leaving the lab.

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DATA AND OBSERVATIONS: You will need 5 tables similar to the one below. Conditions: Temperature: __________ Bleach Concentration: __________

Time (s) Volume of Oxygen Gas Produced (mL) 0 30 60 90 . . .

QUESTIONS AND CALCULATIONS: 1) On a single sheet of graph paper, plot the volume of gas produced vs. time

elapsed. Label each graph with the conditions under which the results being graphed were obtained.

2) For each trial, calculate the overall rate of 50 mL of oxygen gas production (in

mL/min ). 3) Compare the calculated values of the rates with the temperatures used. By what

multiplication factor did the rate change with a 10°C increase in temperature? By what factor did it change with a 10°C decrease in temperature?

4) Compare the calculated values of the rates at the different concentrations of

bleach. By what multiplication factor did the rate change when the concentration of bleach was halved? By what factor did it change when the concentration of bleach was quartered?

5) Bleach is made by the action of chlorine gas on sodium hydroxide:

Cl2 (g) + 2OH-1 (aq) → Cl-1 (aq) + ClO-1

(aq) + H2O (l) However, if an acid is added to bleach, the reverse process occurs:

Cl-1 (aq) + ClO-1 (aq) +2H+1 (aq) → Cl2 (g) + H2O (l)

Why should you never mix bleach with any cleaner or other household product which may contain acid?

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6) Special cleaning agents such as those used for cleaning mold and mildew off bathroom tiles may contain 10% sodium hypochlorite. Predict how the shape of the rate curve with this concentration will differ from that of regular strength bleach.


45

LAB EXPERIMENT 16

SOLUBILITY TRENDS AND PRECIPITATE FORMATION OBJECTIVES: 1. To mix several pairs of solutions together and note if precipitates form. 2. To deduce from the results which combinations of ions form precipitates. 3. To write a balanced formula equation for each precipitation reaction. 4. To write a complete ionic equation for each precipitation reaction. 5. To determine and write the net ionic equation for each precipitation reaction. MATERIALS: Apparatus Reagents Glass square Trays of 8 solutions PROCEDURES: 1. Obtain the 8 solutions from the instructor. 2. Place 2 or 3 drops of one solution onto the glass square. Add 2 or 3 drops of a

second solution into the first solution. 3 If a precipitate forms, record the result in the copy of Table 1 by placing “ppt” in

the appropriate square. If no precipitate forms, simply mark a dash ( - ) in the square.

4. Repeat steps 2 and 3 until all possible combinations have been tested. 5. All solutions may be washed down the sink with plenty of water. 6. Make sure to wash your hands before leaving the lab.

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DATA AND OBSERVATIONS: Table 1

Solution (8) Ag+1

NO3-1

(7) Na+1

NO3-1

(6) Al+3

Cl-1

(5) Ba+2 NO3

-1

(4) Ba+2 Cl-1

(3) Sr+2 NO3

-1

(2) Al+3 SO4

-2

(1) Na+1 SO4

-2 (1)

Na+1 SO4

-2

(2) Al+3 SO4

-2

(3) Sr+2 NO3

-1

(4) Ba+2 Cl-1

(5) Ba+2 NO3

-1

(6) Al+3

Cl-1

(7) Na+1

NO3-1

(8) Ag+1

NO3-1

QUESTIONS AND CALCULATIONS: 1) What observations led you to believe that precipitates formed? 2) How many precipitation reactions did you observe? 3) How many different precipitates did you observe?

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4) Construct a solubility table summarizing your results similar to Table 2, which

shows the combination of ions that formed precipitates and those that did not. Table 2 Example

This Negative Ion Plus These Positive Ions Formed S-2 Na+1, Ca+2 No precipitate S-2 Cu+2 Precipitate

5) For each different precipitate, write the following:

a) balanced formula equation b) complete ionic equation c) net ionic equation


48

LAB EXPERIMENT 17

DETERMINATION OF A SOLUBILITY PRODUCT CONSTANT OBJECTIVES: 1. To prepare a number of solutions of each of Pb+2 and I-1, of differing concentrations. 2. To mix combinations of the above solutions and note whether a precipitate occurs. 3. To obtain an approximate value of the Ksp for PbI2 at room temperature. 4. To obtain the approximate Ksp for PbI2 at temperatures higher than room temperature. MATERIALS: Apparatus Reagents 12 large test tubes 0.010 M Pb(NO3)2 2 graduated cylinders (10 mL) 0.020 M KI Medicine dropper 600 mL beaker 2 – 100 mL beakers Hot Plate Thermometer PROCEDURES: 1. Obtain in separate 100 mL beakers about 40 mL each of 0.010 M Pb(NO3)2 and

0.020 M KI and label the beakers. 2. Obtain 12 large test tubes and arrange them in 2 rows, each labeled near the top of

the test tube A to F. 3 Into the first set of test tubes place 10.0 mL, 8.0 mL, 6.0 mL, 4.0 mL, 3.0 mL, and

2.0 mL of 0.010 M Pb(NO3)2. Use your 10 mL graduated cylinder and medicine dropper to get the precise amount.

4. Add an amount of distilled water to each tube to make the total volume in each

test tube 10 mL. (that is, add 0 mL, 2 mL, 4 mL, 6 mL , 7 ml, and 8 mL respectively)

5. Repeat Steps 3 and 4 using 0.020 M KI. 6. Mix the contents of test tube A from the lead nitrate solutions to test tube A from

the potassium iodide solutions.

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7. Repeat Step 6 for the other solutions B through F. 8. Record in which test tubes a precipitate occurred in Table 1. 9. Add about 400 mL of tap water to the 600 mL beaker, and place the beaker on a

hot plate with a thermometer in it. Place all the test tubes which contain a precipitate into the beaker and begin heating the water bath. All others can go into waste disposal.

10. When the precipitate in each test tube dissolves, note and record the temperature

in Table 1. 11. Once the precipitate has dissolved, it may be removed and put in the waste

disposal jars. 12. After cleaning up, make sure to wash your hands before leaving the lab. DATA AND OBSERVATIONS: Table 1

Test Tube A B C D E F Volume of 0.010 M Pb(NO3)2 (mL) 10.0 8.0 6.0 4.0 3.0 2.0

Volume of water added (mL) 0.0 2.0 4.0 6.0 7.0 8.0 Volume of 0.020 M KI (mL) 10.0 8.0 6.0 4.0 3.0 2.0 Volume of water added (mL) 0.0 2.0 4.0 6.0 7.0 8.0

Precipitate or no precipitate (room temperature)

Temperature at which precipitate dissolves (°C)

QUESTIONS AND CALCULATIONS: 1) For each test tube, calculate the [Pb+2] in the final mixed solution. Use dilution

calculations – total volume is 20 mL. 2) For each test tube, calculate the [I-1] in the final mixed solution. 3) Calculate the value for the trial ion product (Trial Ksp) for each of the test tubes

(found by [Pb+2] [I -1]2 ) 4) State the range of values where the experimental Ksp must lie. This will be

between the two test tubes when you got a precipitate and when you did not get a precipitate.

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5) From the results of the precipitates that were heated, make up a table showing: the

temperature at which the precipitate dissolved, the Trial Ksp, and the solubility of the test tube (given by the [Pb=2]).

6) Using the table from question 5, plot a graph of solubility (mol/L) vs.

temperature. 7) What is the trend in solubility as the temperature is increased? 8) After doing this experiment, a student finds that the test tubes have a yellow

coating on the inside. On the basis of your experimental results, suggest the best method for removing this coating.

9) If you were given a saturated solution of lead iodide and asked to determine the

Ksp, explain the procedure you would use to determine it. 10) Will lead iodide be more soluble or less soluble in a solution of 0.01 M KI than it

will be in pure water? Explain. CONCLUSIONS: Make sure to answer your objectives, summarize your results and tell what you have learned in this experiment. Don’t forget any sources of error that may have occurred.

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LAB EXPERIMENT 18

BRONSTED-LOWRY ACID AND BASE EQUILIBRIA OBJECTIVES: 1. To obtain an understanding of the equilibria which involve acids and bases. 2. To observe the colour changes that occur with a number of different acid-base

indicators in several different solutions. 3. To arrange all the Bronsted-Lowry acids involved in this experiment in order of

decreasing strength. MATERIALS: Apparatus Reagents Sheet of acetate 1 M HCl and NaOH Mixing grid 6 HA/A- acid solutions 5 HIn/In- indicator solutions PROCEDURES: 1. Obtain a piece of acetate paper and grid supplied by the instructor. 2. Obtain the trays of acid and indicator solutions. 3 Place 4 drops of 1 M HCl into each of the squares as indicated on the grid. 4. Add 1 drop of each indicator solution to the designated area, and record the colour

in Table 1. This gives you the colour of the acid form of each indicator. 5. Repeat Steps 3 and 4 using 1 M NaOH instead of HCl. This gives you the colour

of the base form of each indicator. 6. Repeat Steps 3 and 4 using the unknown acid solutions (HA1/A

-1 through to

HA6/A-6). Be specific when recording colours and shades of colours. You are

finished once all 30 possible combinations are complete. 7. Wash all the chemicals down the sink with plenty of water. Make sure you wash

your hands before leaving the lab.

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POST LAB DISCUSSION: The results with the HCl and NaOH with each indicator allow us to determine the acid and base forms (colours) of the indicator through the equation:

HIn + H2O ↔ H3O+ + In-

(acid form) (base form) Adding an acid (HCl) means we are adding H3O

+. This means the reaction shifts left and HIn is formed (which has a colour associated with it) In order to interpret the results and deduce a list of acid strengths, we will use the following equation:

HA + In- ↔ A- + HIn If the acid form of the indicator (HIn) is showing its colour, then the HA will be a stronger acid (remember that the stronger the acid means the more it wants to donate its proton, which in turn favours the opposite side of the equation). If the base form of the indicator (In-) is showing its colour, then the HIn is a stronger acid. DATA AND OBSERVATIONS: Table 1 HIn 1/In1

- HIn 2/In2- HIn 3/In3

- HIn 4/In4- HIn 5/In5

- HCl

NaOH HA1/A1

- HA2/A2

- HA3/A3

- HA4/A4

- HA5/A5

- HA6/A6

- QUESTIONS AND CALCULATIONS: 1) Make up another table like Table 1 but leave out HCl and NaOH. From your

results and the Post Lab Discussion, fill in each box indicating the relative strengths of the two acids combining. For example: HIn1 > HA1 or HIn1 < HA1.

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2) Using the results from question 1, arrange the 11 acids (6 HA’s and 5 HIn’s) with the strongest at the top and the weakest at the bottom. As well, write the ionization equations of each and label each side of the equations with vertical arrows indicating the increasing strength of the acids and the increasing strength of the conjugate bases.

3) The five indicators used in this lab were bromcresol green, bromthymol blue,

indigo carmine, orange IV and thymolphthalein. Write the pH range over which each indicator changes its colour.

4) Use your results to identify each of the indicators (HIn’s) used in the lab. 5) The unknown acids (HA’s) were all whole number pH units in the range of 0-14.

From your results, determine the pH of each unknown acid. CONCLUSIONS: Make sure to answer your objectives, summarize your results and tell what you have learned in this experiment. Don’t forget any sources of error that may have occurred.

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LAB EXPERIMENT 19

ACID-BASE TITRATION

OBJECTIVES: 1. To titrate a hydrochloric acid solution of unknown concentration with standardized

sodium hydroxide solution. 2. To utilize the titration data to calculate the molarity of the hydrochloric acid solution. MATERIALS: Apparatus Reagents 250 mL Erlenmeyer flask standardized sodium hydroxide solution (approx. 0.50 M) volumetric pipette (10 mL) unknown hydrochloric acid solution burette phenolphthalein solution burette clamp and stand pipette bulb PROCEDURE: 1. Obtain about 50 mL of the unknown hydrochloric acid solution and about 50 mL of

the standardized sodium hydroxide solution. Your instructor will provide you with the exact molarity of the NaOH. Record this value in Table 1.

2. Using the pipette bulb and volumetric pipette, rinse out the pipette and then transfer

10.0 mL of the HCl solution into a 250 mL Erlenmeyer flask. 3. Add 3 drops of phenolphthalein solution. 4. Rinse out a burette with some of the standardized NaOH solution and drain it,

keeping the tip filled with the NaOH. Refill the burette and record your initial volume in Table 1.

5. Gradually add the NaOH into the Erlenmeyer flask containing the HCl. Swirl the

flask continuously. Continue adding NaOH, noting any changes in the flask. You may wash any splashes down into the flask with distilled water.

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6. As the equivalence point is reached, a light pinkish colour will appear. As you approach this point, the colour will take longer to disappear as you swirl it. Add the NaOH drop by drop until the colour no longer disappears. Record the final volume of the burette in Table 1. The most accurate reading is one in which the solution is a very faint pink.

7. Repeat steps 2 through 6 after cleaning out your Erlenmeyer flask (the contents can

be poured down the sink with lots of water). Continue repeating until you get two trials that are within ± 0.10 mL of one another.

8. Clean up all your apparatus and make sure you wash your hands with soap and water

before leaving the lab. OBSERVATIONS AND DATA: Table 1: Molarity of NaOH =_______________


Initial Volume of NaOH Final Volume of NaOH Volume of NaOH used Average Volume of NaOH QUESTIONS: 1. Calculate moles of NaOH from the average volume used in the experiment.

2. Calculate moles of HCl present originally.

3. Calculate the molarity of the HCl solution.

4. What was the reason for cleaning out the burette with NaOH before starting the

titration? 5. What is the concentration of a NaOH solution when it requires 30.0 mL of 0.50 M

HCl to neutralize 50.0 mL of the NaOH? 6. What is the concentration of acetic acid (HCH3COO) when 32.5 mL of 0.56 M

NaOH is required to neutralize 15.0 mL of the acid? 7. When 6.25 g of oxalic acid (H2C2O4) are placed in water containing phenol red,

the suspension is yellow. What is the molarity of the KOH if 32.2 mL of KOH is required to change the colour to red?

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8. A 5.0 g tablet of TUMS (Mg(OH)2) neutralizes 450 mL of stomach acid (HCl). What is the molarity of the stomach acid?


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LAB EXPERIMENT 20

HYDROLYSIS – THE REACTION OF IONS WITH WATER

OBJECTIVES: 1. To measure the pH of a large number of salts, and identify those which have

undergone hydrolysis. 2. To explain why hydrolysis occurs (or does not occur) and to write a net ionic

equation for each hydrolysis. 3. To deduce which is greater for some amphiprotic anions, the Ka for the further

ionization of the ion, or the Kb for the anionic hydrolysis of the ion. MATERIALS: Apparatus Reagents Spot plate 0.10 M solutions of 19 various salts Universal indicator paper PROCEDURE: 1. Obtain 2 spot plates and label the wells from 1 to 19. 2. Obtain some universal indicator paper. Tear off 1 cm pieces and place them in each

of the wells in the spot plates. 3. Add 2 or 3 drops of each solution into its corresponding well. Using the colour chart

provided with the universal indicator paper, record the colour and the corresponding pH of each solution in Table 1 and Table 2 (because the solutions in Table 2 are amphiprotic, the results will be interpreted differently).

4. Once finished, put the paper in the garbage and wash all solutions down the sink with

plenty of water. 5. Clean up and make sure you wash your hands before leaving the lab.

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OBSERVATIONS AND DATA: Table 1:

Solution Colour of Universal Indicator

pH Type of Hydrolysis (anionic, cationic, both, or neither)

1. NaCH3COO 2. NaCl 3. NH4Cl 4. (NH4)2SO4 5. AlCl3 6. Ca(NO3)2 7. Fe2(SO4)3 8. Na2CO3 9. Na3PO4 10. K2SO4 11. KBr 12. (NH4)2C2O4 13. NH4CH3COO 14. (NH4)2CO3 Table 2:

Solution Colour of Universal

Indicator pH Type of Reaction

(further ionization or anionic hydrolysis)

15. K2HPO4 16. KH2PO4 17. NaHCO3 18. KHSO4 19. NaHSO3 QUESTIONS AND CALCULATIONS: 1. For each of the ammonium salts write dissociation equations and then, if any of

the species will undergo hydrolysis, write the hydrolysis reactions . 2. For each of the salts from question 1, compare the Ka for ammonium with the Kb

of the anion of each salt. State which reaction should be predominant (occur to the greater extent) and then verify if your calculations agree with your experimental observations.

3. For each of the amphiprotic salts from Table 2, write the two net ionic

(hydrolysis) equations that could occur.

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4. For each of the substances in question 3, compare Ka with Kb and state which reaction should be predominant. Verify if your calculations agree with your experimental observations.


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LAB EXPERIMENT 21

ACID-BASE TRENDS OF METAL AND NON-METAL OXIDES

OBJECTIVES: 1. To test the acid-base properties of metal oxide solutions. 2. To test the acid-base properties of non-metal oxide solutions. 3. To determine what, if any, periodic trend exists in acid-base properties of oxide

solutions. MATERIALS: Apparatus Reagents 10 large test tubes universal indicator solution Rubber stoppers CaO, MgO, ZnO, Al2O3 powders 2 medicine droppers 2 unknown non-metal oxide solutions 2 small beakers 6 M HCl 6 M NaOH PROCEDURE: Part I: Preparing and Testing Solutions of Metal Oxides 1. Add a small amount of each of the four powdered oxides to separate, labelled test

tubes. The powder should just cover the bottom of the test tube. 2. Half fill the test tubes with distilled water, stopper them, and shake the contents for

about 2 minutes. These will form slurries, as the oxides are only slightly soluble. 3. Add 2 or 3 drops of universal indicator solution to each slurry, stopper and shake

again. Compare the colours formed with the indicator colour wheel and record your results in Table 1.

Part II: Testing for Amphiprotic Behaviour 1. Shake each test tube thoroughly again, then divide each slurry into two portions by

pouring half of the sample into a clean test tube.

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2. Obtain about 10 mL of 6 M HCl and 10 mL of NaOH in separate labelled beakers. Place a clean medicine dropper in each beaker.

3. Test one portion of each slurry by adding 6 M HCl drop by drop, shaking the test tube

occasionally. Continue adding the acid and shaking until no further changes occur. Pay close attention to what happens to the cloudiness of each slurry. Note all changes in Table 2.

4. Test the other portion of each slurry by adding 6 M NaOH drop by drop. Shake the

test tube occasionally and record any changes in Table 2. 5. Once finished, dispose of the chemicals in the waste disposal jar. Part III: Testing Solutions of Non-metal Oxides 1. Half fill two labelled test tubes with unknowns A and B. Solution A is a hydroxide

solution of an oxide of sulphur. Solution B is a hydroxide solution on an oxide of nitrogen.

2. Test both solutions with universal indicator solution and record the results in Table 1. 3. Place the chemicals in the waste disposal jar. 4. Make sure to clean up and wash your hands before you leave the lab. OBSERVATIONS AND DATA: Table 1:

Hydroxide Solution Tested

Universal Indicator Results

Acidic, Basic, or Neutral

Mg(OH)2 Ca(OH)2 Zn(OH)2 Al(OH) 3

SO2(OH)2 NO2(OH)

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Table 2:

Hydroxide Solution Tested

Addition of Acid or Base

Universal Indicator Results Before After

Other Changes

Mg(OH)2 Mg(OH)2 Ca(OH)2 Ca(OH)2 Zn(OH)2 Zn(OH)2 Al(OH) 3 Al(OH) 3

QUESTIONS: 1. Classify each of the six hydroxide solutions you examined as acidic, basic,

neutral, or amphiprotic. 2. a. Write the formula for solution A (SO2(OH)2) in a different, more familiar form.

(Hint: Put the elements in groups and place hydrogen at the beginning) b. What is the acid name of this substance?

3. Repeat question 2 for solution B (NO2(OH)). 4. Carbon dioxide gas, an oxide of carbon, is present in varying amounts in the

atmosphere. Predict whether rainwater containing dissolved CO2 would be acidic or basic.

5. How would acid rain be formed from the pollutants SO2 an NO? (Hint: React

these with water) CONCLUSIONS: Make sure to answer your objectives, summarize your results and tell what you have learned in this experiment. Don’t forget any sources of error that may have occurred.

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LAB EXPERIMENT 22

PREPARATION AND STANDARDIZATION OF ACID AND BASE SOLUTIONS AND THE TESTING OF UNKNOWNS

OBJECTIVES: 1. To prepare a standard solution of oxalic acid and use it to standardize an unknown

solution of sodium hydroxide solution. 2. To determine the molar mass of an unknown solid acid by titration with standardized

NaOH solution. 3. To determine the pH and molarity of an unknown acid solution and calculate the Ka

from the results. 4. To analyze the acid concentration in soft drinks. MATERIALS: Apparatus Reagents 250 mL volumetric flask oxalic acid crystals 250 mL Erlenmeyer flask unknown sodium hydroxide solution Volumetric pipettes (10 & 25 mL) phenolphthalein indicator Burette unknown solid acid crystals Burette clamp and stand unknown weak acid solution Suction bulb soft drink Funnel Mass balance Beakers (100 mL, 250 mL) Wash bottle PROCEDURE: Part I: Preparation of a Primary Standard Acid 1. Before beginning the lab, calculate the mass of oxalic acid (H2C2O4•2H2O) that you

will need to make up 250.0 mL of a 0.0500 M solution and show it to your instructor. 2. Accurately determine the mass of a clean and dry 100 mL beaker and record it in

Table 1.

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3. Measure into the beaker the amount of oxalic acid you have calculated, determine the

mass of the oxalic acid and beaker, and record it in Table 1. The important thing is to record the mass you have, not to get exactly the mass calculated in step 1.

4. Dissolve the acid in about 60 mL of distilled water, and pour the solution through a

funnel into a 250 mL volumetric flask. Wash the beaker with distilled water twice, and add these to the flask. Now add water to the flask until you reach the volumetric mark.

5. Stopper the flask, and shake to ensure that the solution is thoroughly mixed. This is

your standard solution of oxalic acid. Part II: Standardization of an Unknown NaOH Solution 1. Obtain about 100 mL of NaOH solution and label the beaker. Add about 15 mL to a

burette, rinse it back and forth, and discard it through the tip into the sink. 2. Fill up the burette to the top and allow some to drain through the tip to remove any air

bubbles. 3. Pour about 50 mL of oxalic acid into a beaker and label it. Using a suction bulb,

withdraw about 5 mL into the 25 mL pipette, rinse it around and discard it down the sink.

4. Withdraw 25 mL of the oxalic acid into the pipette and transfer it to a 250 mL

Erlenmeyer flask using the release mechanism on the suction bulb. The correct volume is delivered when you have touched the tip of the pipette to the side of the flask after you have released all the oxalic acid.

5. Add 3 drops of phenolphthalein to the acid in the Erlenmeyer flask and place it under

the burette. 6. Read the initial volume of NaOH on the burette as accurately as you can and record

the value in Table 2. Then open the valve on the burette and allow the NaOH to run into the Erlenmeyer flask, swirling constantly to ensure through mixing.

7. After a time, you will notice a pink colour appearing. When this colour takes a longer

time to disappear, slow down the rate of addition of the NaOH until you are adding it drop by drop. Stop the titration when the faintest pink possible stays in the flask for at least 30 seconds. Read the final volume of NaOH on the burette as accurately as possible and record it in Table 2.

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8. If you are in doubt as to whether you have a pale pink, add one more drop. If the colour turns a darker pink, then you have reached the endpoint of the titration. Discard the solution down the sink and rinse the Erlenmeyer with distilled water.

9. Pipette another 25 mL of oxalic acid into the Erlenmeyer flask and again add 3 drops

of phenolphthalein. Refill the burette if necessary, take your initial volume reading and repeat the titration. You will continue this process until you have two ‘volume of NaOH used’ readings that are within ± 0.10 mL.

Part III: Determination of the Molar Mass of an Unknown Solid Acid 1. Obtain a vial containing an unknown solid acid. Record the identifying number in

Table 3. 2. Weigh out about 0.75 g of the solid acid into a clean, dry beaker and record the exact

mass you obtained in Table 3. (It does not have to be exactly 0.75 g, just record exactly how much you got)

3. Dissolve the acid in about 40 mL of water and transfer all the solution into an

Erlenmeyer flask. Rinse the beaker twice to ensure all the acid is transferred. Add 3 drops of phenolphthalein to the flask.

4. Fill the burette with NaOH, record the initial volume, and titrate the solution until the

endpoint is reached. Record the final volume in Table 3. Pour the solution down the sink and rinse the Erlenmeyer with distilled water.

5. Repeat steps 2 to 4 until you get two readings in close agreement. If you did not have

the masses exactly the same each time, check whether the results agree by determining the ratio of volumes to masses of each trial.

Part IV: Determination of Ka for an Unknown Monoprotic Weak Acid 1. Obtain approximately 100 mL of the unknown weak acid provided and record its pH

in Table 4. 2. Use the suction bulb to withdraw about 5 mL of the unknown acid into a 25 mL

pipette, rinse, and discard it down the sink. 3. Pipette a 25 mL sample of the unknown acid and release it into an Erlenmeyer flask

and add 3 drops of phenolphthalein. 4. Fill the burette with NaOH and titrate as you did in Parts II and III, until you get two

results agreeing to within ± 0.10 mL, recording your results in Table 4.

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Part V: Determination of Acid Concentration in Soft Drinks 1. Obtain about 50 mL of soft drink in a clean, dry beaker from your instructor. 2. Obtain a 10 mL pipette and suction bulb. Clean out your pipette by withdrawing

about 3 mL of soft drink into the pipette, rinsing, and discarding into the sink. 3. Pipette a 10 mL sample of the soft drink and release it into an Erlenmeyer flask. Add

3 drops of phenolphthalein. 4. Fill the burette with NaOH and titrate as in Part IV until you get two results agreeing

to within ± 0.10 mL, recording your results in Table 5. 5. Once finished, clean all apparatus, and make sure to wash your hands before leaving

the lab. OBSERVATIONS AND DATA: Table 1: Preparation of Oxalic Acid Solution Calculated Mass of Oxalic Acid (g) Mass of Beaker (g) Mass of Beaker + Oxalic Acid (g) Mass of Oxalic Acid Used (g) Table 2: Standardization of NaOH Trial 1 Trial 2 Trial 3

(if necessary) Initial reading of burette (mL) Final reading of burette (mL) Volume of NaOH used (mL) Average Volume of NaOH (mL) Table 3: Determination of Molar Mass of Unknown Solid Acid

Unknown Solid Acid # ______


Mass of Beaker (g) Mass of Beaker + Acid (g) Mass of Acid (g) Initial Volume of NaOH (mL) Final Volume of NaOH (mL) Volume of NaOH used (mL)

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Table 4: Determination of Ka for an Unknown Monoprotic Weak Acid

pH of unknown acid solution __________


Initial reading of burette (mL) Final reading of burette (mL) Volume of NaOH used (mL) Table 5: Determination of Acid Concentration in Soft Drinks Trial 1 Trial 2 Trial 3

(if necessary) Initial reading of burette (mL) Final reading of burette (mL) Volume of NaOH used (mL) Average Volume of NaOH (mL) QUESTIONS AND CALCULATIONS: Part I: Preparation of a Primary Standard Acid 1. Calculate the molar mass of the oxalic acid (H2C2O4•2H2O). 2. Calculate the number of moles of oxalic acid used based on your measured mass. 3. Calculate the molarity of the oxalic acid solution you made up. Part II: Standardization of an Unknown NaOH Solution 1. Calculate the number of moles of oxalic acid in the 25.00 mL sample. 2. Write the neutralization equation and calculate the number of moles of sodium

hydroxide necessary to neutralize this amount of oxalic acid (from question1). Remember, oxalic acid is diprotic.

3. Using the average volume of NaOH used in this titration, calculate [NaOH] in the

standardized solution.

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Part III: Determination of the Molar Mass of an Unknown Solid Acid 1. Calculate the number of moles used in each trial. 2. For each trial, calculate the number of moles of unknown acid neutralized by this

amount of NaOH. Remember the unknown acid is monoprotic. 3. Using the relationship: molar mass = mass of substance (g) number of moles For each trial, calculate the molar mass. 4. Calculate the average molar mass of the unknown solid acid. Part IV: Determination of Ka for an Unknown Monoprotic Weak Acid 1. Using the pH, calculate the equilibrium [H3O

+] present in the unknown weak acid solution.

2. Calculate the average volume of NaOH required for the titrations. 3. Calculate the number of moles of NaOH needed to neutralize the acid solution.

Remember the acid is monoprotic. 4. Calculate the original concentration of the unknown acid. 5. Using an ICE chart, calculate the Ka for this weak monoprotic acid. Part V: Determination of Acid Concentration in Soft Drinks 1. Using the data from Table 5 and knowing that the acid in soft drinks is citric acid

(which is triprotic ), calculate the concentration of the acid in soft drinks. 2. Other ingredients of soft drinks include sugar or other artificial sweeteners. What

property of acids is overcome by using sugar or sweetener? 3. Every day, a manufacturing plant produces 5.0 × 103 L of 3.0 M NaOH. In order to

comply with environmental regulations, this NaOH must be neutralized before being discharged as effluent. What mass of 12 M HCl will be required to neutralize it? The density of HCl is 1.2 kg/L.


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LAB EXPERIMENT 23

ELECTROCHEMICAL CELLS

OBJECTIVES: 1. To become familiar with the construction and operation of electrochemical cells. 2. To predict the reactions and voltages that should result. 3. To construct three electrochemical cells and measure their voltages. 4. To observe the effect of non-standard conditions on voltage. MATERIALS: Apparatus Reagents 3 – 250 mL beakers metal strips of copper, zinc, and lead U-tube 0.5 M copper (II) nitrate solution Cotton balls 0.5 M zinc nitrate solution 2 wire leads with clips 0.5 M lead (II) nitrate solution Voltmeter 0.5 M potassium nitrate solution Steel wool PROCEDURE: Part I: Making a Zinc-Lead Cell 1. Obtain about 80 mL of zinc nitrate solution and 80 mL of lead (II) nitrate solution in

separate labeled beakers. 2. Prepare a salt bridge by filling a glass U-tube with potassium nitrate and plugging

both ends with cotton balls. Do not allow air to become trapped in the tube. Carefully invert the tube and place it into the two beakers of solutions. You will use this prepared U-tube throughout the experiment.

3. Obtain a zinc strip and a lead strip and clean the surfaces with steel wool. Place the

zinc strip in the zinc solution and the lead strip in the lead solution. 4. Connect a wire lead to the Zn strip and another to the Pb strip. Connect the other

ends to a voltmeter so as to get a positive reading on the voltmeter scale. Record the measured voltage and make a labeled sketch of your electrochemical cell. Identify the anode and cathode as well as the electron flow on your diagram.

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5. Use the table of standard reduction potentials to calculate and record the theoretical

cell voltage. 6. Observe the effect of removing the salt bridge or any electrode from the cell. 7. Take the cell apart and save the solutions and salt bridge for Parts II and III. Part II: Making a Lead-Copper Cell 1. Obtain 80 mL of copper (II) nitrate solution and a copper strip. Clean the strip with

steel wool. 2. Use the procedure from Part I to study a lead-copper electrochemical cell. Record the

measured voltage and make a labeled sketch of the cell, as in Procedure 4 in Part I. 3. Use the table of standard reduction potentials to calculate and record the theoretical

cell voltage. 4. Take the cell apart. The lead (II) nitrate can be returned to the original container

unless it is contaminated. If it is contaminated, dispose of it in the chemical waste jar. The copper solution will be used in Part III.

Part III: Making a Zinc-Copper Cell 1. In this part, you will construct and study a zinc-copper cell using the procedure from

Part I. Record the measured voltage and make a labeled sketch of the cell, as in Procedure 4 in Part I.

2. Use the table of standard reduction potentials to calculate and record the theoretical

cell voltage. 3. Take the cell apart and put the salt bridge contents in the waste chemical jar. The

other solutions can go back into their original containers unless they are contaminated. If contaminated, place in the waste disposal jar.

4. Clean up all apparatus and wash your hands before you leave the lab.

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QUESTIONS: 1. For each cell constructed, write the equations for the following:

a. the anode half reaction b. the cathode half reaction c. the overall cell reaction

Calculate the theoretical standard state cell voltage

2. What effect did removing the salt bridge have on the operation of each electrochemical cell? Explain.

3. How did the theoretical voltages compare with the experimental voltages? Was

there any general trend or pattern? 4. Which cations should migrate into the salt bridge of a zinc-copper cell? How

could you test this? 5. Predict the voltage reading of a lead-copper cell when the cell reaches

equilibrium. 6. Explain why you should not expect an electrochemical cell to operate for an

unlimited amount of time. 7. Design and draw a diagram of an electrochemical cell that has a magnesium

anode and a theoretical standard state voltage of 3.17 V. Label all parts of the diagram as you did in the experiment.


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LAB EXPERIMENT 24

ELECTROLYTIC CELLS

OBJECTIVES: 1. To electrolyze a KI solution using carbon electrodes. 2. To electrolyze a ZnSO4 solution using carbon electrodes. 3. To copper plate a metal object. 4. To interpret the products of each electrolytic cell with anode and cathode half

reactions. MATERIALS: Apparatus Reagents Ring stand and burette clamp metal strip of copper U-tube 2 cylindrical carbon rods 2 - 250 mL beakers bare copper wire 2 wire leads with clips 1.0 M KI Electrode holder 1.0 M ZnSO4 DC power supply 1.0 M CuSO4 Steel wool phenolphthalein solution PROCEDURE: Part I: Electrolysis of 1.0 M KI 1. Fill a clean U-tube with 1.0 M KI solution so that the solution level is about 2.0 cm

from the top. Add 5 drops of phenolphthalein to each side of the tube. 2. Obtain 2 carbon rods and clean with steel wool. Attach a wire lead to each of the

carbon rods and place them in each side of the U-tube. Do not allow the metal ends of the wire touch the solution.

3. Connect the other end of the wire leads to a DC power supply. Observe what

happens at each electrode during the next several minutes and record these observations in Table 1.

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4. Use the table of standard reduction potentials to determine the anode and cathode half reactions and record these in Table 1.

5. Turn off the power supply and remove the carbon rods from the solution. Dry the

rods for use in Part II and dispose of the U-tube chemicals in the waste disposal jar. Part II: Electrolysis of 1.0 M ZnSO4 1. Obtain 150 mL of 1.0 M ZnSO4 solution in a clean 250 mL beaker. 2. Place an electrode holder on the beaker and insert a carbon rod into each side of the

electrode holder. 3. Attach the wire leads to the carbon rods and then to the power supply. Turn on the

power supply record your observations in Table 2. Do not operate the cell longer than 5 minutes, as the carbon rods may begin to flake apart.

4. Use the table of standard reduction potentials to determine the anode and cathode half

reactions and record these in Table 2. 5. Turn off the power, remove the rods and place the solution back into its original

container unless it is contaminated. If contaminated, dispose of it in the waste disposal jar.

Part III: Copper Plating 1. Obtain 200 mL of 1.0 M CuSO4 solution in a clean 250 mL beaker. 2. Prepare your metal object for plating by polishing it with steel wool. Wear gloves

when doing so as the idea is to remove all oxides and oils so that the copper plating will adhere more effectively. After cleaning the object, make sure you do not touch it with your fingers.

3. Place an electrode holder onto an empty 250 mL beaker. On one side of the electrode

holder attach a copper strip. Get a length of bare copper wire and attach it to the metal object you are going to plate. Attach the end of the copper wire to the other side of the electrode holder ensuring that the metal object is completely below the 200 mL mark

4. Attach the wire leads to the power supply. Connect the negative lead of the power

supply to the bare copper wire and the positive lead to the copper strip. 5. Lift the whole apparatus off of the empty beaker and place it in the 200 mL of CuSO4

solution. Turn on the power supply for several minutes and record your observations in Table 3.

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6. Use the table of standard reduction potentials to determine the anode and cathode

half reactions and record these in Table 3. 7. Once plating is complete, turn off the power supply and clean up. The CuSO4 can go

back into its original container. Dry off your plated object with paper towel. 8. Make sure you wash your hands before leaving the lab. DATA AND OBSERVATIONS: Table 1: Electrolysis of 1.0 M KI

Electrode Observations Half Reactions Anode

Cathode Table 2: Electrolysis of 1.0 M ZnSO4


Cathode Table 3: Copper Plating


Cathode QUESTIONS: 1. For Part I, how do the anode and cathode half reactions explain your

observations? 2. For Part II, how do the anode and cathode half reactions explain your

observations? 3. In Part II, what anode and cathode reactions would occur if the overpotential

effect did not exist? 4. For Part III, how do the anode and cathode half reactions explain your

observations?

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5. The electrolytic cell in Part I produces reactions similar to those in a chlor-alkali

industrial plant, where a NaCl solution is electrolyzed using inert electrodes.

a. Predict the anode and cathode half reactions for a chlor-alkali plant. b. Suggest a reason why it would be unwise to do this reaction at home? c. What common acid is often manufactured as a byproduct using the products

of the reactions?

6. If the overpotential effect did not exist, which common electroplating process would be impossible? Why?

7. A 1.0 M Na2SO4 solution is to be electrolyzed in a U-tube containing carbon

electrodes. Before beginning, bromthymol blue indicator solution is added to the colourless solution, and the colour is adjusted to a neutral green. Write the anode and cathode half reactions that should occur and predict the colours that will result at each electrode once the power is turned on.


lab experiment 1 making observations · lab experiment 1 making observations objectives: 1. to make...

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