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Name________________________________________ SC_____-_____ Date___________ ACIDS AND ACID ANHYDRIDES

Equipment: 3 test tubes, 1 gas collecting bottle, 1 deflagrating spoon, stopper, charcoal Materials: pH papers, phenolphthalein, red litmus paper, blue litmus paper, methyl orange, 20 mL of 3M HCl (hydrochloric acid), 10 mL 3M H2SO4 (sulfuric acid), 10mL HC2H3O2 (CH3COOH) (acetic acid), 4-5 pieces of Zn(s). Aim: What are some important properties of acids? Common household substances containing acids are vinegar (acetic acid) and lemon juice (citric and ascorbic acids). A. Test 3M hydrochloric acid (HCl) with red and blue litmus papers and pH paper as follows:

Put 10 mL of water into a test tube. Add a drop of the acid. Use a glass rod to transfer a drop of the diluted acid to the litmus paper. Record the result in table I below.

B. Pour 5 mL. of 3M HCl into a test tube and add 2 drops of phenolphthalein. C. Pour 5 mL. of 3M HCl into another test tube and add a few drops of methyl orange. Record

the results below. D. Repeat "A" "B" and "C" above with 3M H2SO4 (sulfuric acid) and 3M HC2H3O2 (acetic

acid). Record the results below. Table I 3M HCl 3M H2SO4 3M HC2H3O2 Blue Litmus

Red Litmus

pH Paper

Phenolphthalein

Methyl Orange

E. Place a small piece of zinc in the bottom of a test tube and cover this with 3M HCl. Test the

ensuing gas with a burning splint and record the results below. Repeat, using 3M H2SO4 and 3M HC2H3O2 in separate test tubes. Caution: This experiment should be performed in the hood.

Table II Zn metal + Gas Evolved Rate of Reaction

Fast, Moderate, Slow Reaction Products

HCl

H2SO4

HC2H3O2

F. Put a very small piece of zinc in a test tube. Add 3M HCl. When all of the zinc is

consumed in the reaction transfer the solution to a beaker and evaporate to near dryness. Caution: Watch out for spattering!

1. How does the residue compare in appearance with the original zinc? 2. What is the name of the compound formed? Teacher Demonstrations G. Lower a deflagrating spoon containing burning charcoal into a gas collecting bottle one third

full of water. After half a minute remove the spoon with the burning charcoal, stopper the bottle and shake. Test with litmus and pH paper.

3. Describe the results. 4. What type of compound must have been formed? Summary Questions Complete and Balance the following equations: 1. C + O2 __________ __________ + H2O H2CO3 2. S + O2 __________ __________ + H2O H2SO3

3. _____ + ___ P + ___ O2 H3PO4

8. In Part E one of the acids produced hydrogen slower than the other two. Explain. 9. What element is present in all acids? 10. How can we often tell, just by looking at a formula, if a substance is an acid? Conclusions Acids are compounds containing the element __________ which is easily replaced by an

active __________. Acids have a __________ taste, and change litmus from __________ to

__________. An acid anhydride is a __________ oxide that reacts with water to form an

__________ solution. Based on Reference Table J the four least active metals are

__________, __________, __________ and __________.

Additional Experimentation On a separate piece of paper create a chart similar to Table I above. Test the pH of the household substances that you teacher instructed you to bring with you to class. Use the spot wells provided and record you results. Allow one point for each answer given correctly in Tables I and II = ________ (24 max). Allow three points for each correct answer in questions one through ten = ________ (30 max). Allow three points for each correct answer in the conclusion section = ________ (33 max). Student brought in a product from home and tested it as done in parts A through D = ________ (13 max). Deductions for lateness to class, report handed in late, safety and participation = ________. Total Grade = ________.

Name________________________________________ SC_____-_____ Date___________ BASES AND UNKNOWN IDENTIFICATION

Equipment: 3 test tubes, stopper. Seven unknown solutions as below. Conductivity set up with light bulb. Materials: pH papers, litmus papers, methyl orange, charcoal 5-6 mL of the following: 3M NaOH, 3M KOH, sat'd Ca(OH)2, NH3(aq),

NaOH(dilute), con. NaOH Aim: What are some important properties of bases? Some common household substances which contain bases are lye (NaOH), milk of magnesia (Mg(OH)2), and

ammonia water. A. Test 3M NaOH, 3M KOH, saturated Ca(OH)2, and NH3(aq) with pH paper. Do this by pouring a few drops of

the base into a test tube, dip a glass rod into the solution, then touch the pH paper with the tip of the glass rod. Test the bases with phenolphthalein, methyl orange, and litmus. Record the results in the Table below.

B. Put a drop of very dilute sodium hydroxide on the tip of your finger and rub it with your thumb. WASH IT

OFF IMMEDIATELY!!! Repeat using dilute solution of the other bases. Record the results in the Table below.

Table of Results - Indicators of Bases NaOH KOH Ca(OH)2 NH3(aq)

pH

Litmus – Red

Litmus – Blue

Phenolphthalein

Methyl Orange

Feel

Action on Acids

1. State four characteristics of bases. C. Identification of unknown solution. On the teachers’ desk there are seven solutions labeled A through G. There is one of each of the following: a strong acid, a weak acid, a strong base, a weak base, a sodium chloride solution, an alcohol solution and a sample of pure water. How would you go about identifying the solutions from each other? Examine the available materials on the desk with your classmates and create a procedure together. Then check with your teacher to see if your procedure is safe and feasible. Your teacher will work with the class as a group in identifying the solutions as per your suggested procedures.

Plan and record your predictions below: Solution: Test you are going perform: Results you expect:

Strong Acid

Strong Base

Weak Acid

Weak Base

Pure Water

Salt Solution

Alcohol Solution

Your conclusion: Solution A is

Because you observed:

Solution B is


Solution C is


Solution D is


Solution E is


Solution F is


Solution G is


Conclusions According to the Arrhenius Theory, bases are _______________ compounds which

_______________ in water to release _______________ion. This ion in solution causes certain

indicators to change _______________. Fairly concentrated basic solutions feel _______________ and

they can _______________ acid solutions. A basic anhydride is a(n) _______________ of a metal that

reacts with water to form a _______________ solution.

One point for each box in part B = ________ (28 max). Two points for each box in part G = ________ (56 max). Two points for each answer in the conclusion questions = ________ (16 max). Deductions for lateness to class, report handed in late, safety and participation = ________. Total Grade = ________.

Name________________________________________ SC_____-_____ Date___________ BLUEPRINTING: A PHOTOCHEMICAL PROCESS

Equipment: 3 test tubes, photoflood lamp, pen, print frames. Materials: 10mL FeCl3, 5mL H2C2O4 (oxalic acid), drops (3 mL) K3Fe(CN)6 (potassium ferri cyanide), ferric ammonium citrate, India ink, white paper, tracing paper Aim: How can we use a redox reaction to prepare a blueprint? A. Add equal (5 mL) amounts of oxalic acid solution and FeCl3 to a test tube and shake. Pour half of the resulting solution into a second test tube. Pour 5mL of FeCl3 into a third test tube. Place one of the first two test tubes under the hood, or some place to keep it out of the light as far as possible. Expose the second and third (containing only the FeCl3) tubes, for 3 to 5 minutes to sunlight or a 500 Watt photoflood lamp. Test the contents of each test tube by adding a few drops of K3Fe(CN)6. Note the color of just the FeCl3. __________. Note the color of just the H2C2O4 __________. 1. In which test tubes does a blue precipitate appear? 2. Which ion does this show is present, Fe+++ or Fe++ ? 3. How can the presence of this ion be explained? 4. Complete and balance the equation: FeCl3 + H2C2O4 → HCl + CO2 + __________ (oxalic acid) B. Coat a piece of plain white paper (about 3" X 4") with a solution of ferric ammonium citrate. This may be accomplished by holding the paper by one corner and drawing it over the surface of the ferric ammonium citrate solution. Dry the paper by waving it in the air, and as far as possible keeping it out of strong light. While waiting for the paper to dry, draw a design, or write, using India ink, on tracing paper. When the paper is perfectly dry, place the tracing paper onto the paper coated with the ferric ammonium citrate and place both in a printing frame. The tracing paper should be next to the glass in the frame. Expose the paper in the frame to the light for 3 to 5 minutes depending upon how strong the light is. Then take the paper out and dip it in a solution of K3Fe(CN)6 for a moment, and immediately wash the paper with running tap water. Let it dry and attach your blueprint to your report. 5. Describe and explain the result. 6. Why was the washing with water necessary? 7. In this experiment what substance was oxidized? CONCLUSIONS

The blueprint process depends upon the of Fe+++ to Fe++ by means of a agent in the presence of light.

Name___________________________________________ SC_____-_____ Date___________ THE BUNSEN BURNER

Equipment: Bunsen burner, tongs, chalk, matches or flint-lighter

Aim: What is the procedure for using a Bunsen Burner properly? A. Use of the Bunsen Burner. In your laboratory work you may use two types of burners, the Tirrell burner and the Meeker burner. While details of construction vary, each has the following parts:

1. a gas inlet in the base, to which a rubber tube is usually attached, 2. a barrel, a vertical tube in which the gas is mixed with air, 3. air ports, adjustable openings at the base of the barrel for admitting air into the gas stream. 4. Tirrell and Meeker burners also have an adjustable needle valve, sometimes called the spud, that regulates the supply of gas. Examine your burner and identify the four parts mentioned above. Draw a labeled diagram of your burner.

B. Attach the burner to the gas jet. Partially close the air ports. Check that the needle valve is not closed. Turn

the gas on the table to the full on position, (the handle will be parallel to the jet) and hold a lighted match about a half-inch above the top and to the side of the burner. Then turn the needle valve to regulate the flame.

1. Describe how turning the needle valve affects the flame. Be sure to answer this question and every question in the lab manual in complete sentences!

C. After obtaining a medium sized flame, open the air ports completely. (Caution: If a very low flame is needed,

the air ports should remain partially closed. Otherwise the flame will strike back and burn at the base of the barrel. If this happens shut off the gas, decrease the amount of air admitted and re-light the burner.)

2. What is the color of the flame? 3. This flame produces the highest possible temperature. Why? D. CHECK THE FUME HOODS FOR AIR BEING DRAWN. IF THE FUME HOODS ARE NOT WORKING

THEN SKIP PARTS D, E and H. A teacher can demo this part at the front fume hood instead. Working under the fume hoods, close the air ports. 4. What is the color of the flame now?

5. This flame does not produce the highest possible temperature. Why? E. The gas that is burning is mainly a compound of carbon and hydrogen (methane). When it burns completely

carbon dioxide and water vapor are formed. With a pair of tongs hold a piece of chalk in the yellow flame for a few seconds.

6. Draw, label and describe what you see on the chalk. 7. What could this deposit be? Explain how it could have been formed.

SAVE THIS PIECE OF CHALK for part H.

F. Open the air ports again. This non-luminous flame is the flame you should use most of the time. 8. Make a diagram of this flame showing its different regions.

To determine which part of the flame is the hottest, hold a piece of wire with tongs in the various regions of the flame until it glows red hot. Count the number of seconds it takes for the wire to glow. After each trial dip the wire in water so that each trial will start at the same temperature.

9. On the diagram you made in Part F, above, label the hottest and coolest parts of the flame. H. Hold the piece of chalk (use tongs) from Part E in the hottest part of the flame for at least one minute. 10. Describe what happened to the deposit on the chalk. 11. Write a possible equation for the process you observed. SUMMARY QUESTIONS 12. How should you adjust the SIZE of the flame? 13. How should you adjust the TEMPERATURE of the flame?

14-15. Methane gas is made of molecules of CH4. Write equations to for reactions when burning methane with both yellow and blue flames. Yellow______________________________________________________________________________ Blue________________________________________________________________________________ 16. Write an equation to show how carbon monoxide can form when methane is burned. What problems are caused by carbon monoxide? Give students full credit for parts D, E and H if the fume hoods were not working and these parts of the experiment were not completed. Number of questions correct for 1 through 16 times five points each = ________ (80 max). Drawing and labeling the Bunsen Burner in Part A times 4 points each = ________ (20 max). Deductions for lateness to class, report handed in late, safety and participation = ________ Total Grade = ________

Name___________________________________________ SC_____-_____ Date___________ CALORIMETRY AND HEAT OF COMBUSTION

Equipment: balance, can lid, chimney, empty can, glass rod, ring, ring stand, thermometer Materials: candle, matches Aim: How can we determine the heat of combustion for a substance?

A. Attach a candle to a can lid, and weigh them to the nearest 0.01 gram. Record this weight in the data table.

B. Weigh an empty can to the nearest .01 gram. Record

this weight in the data table.

C. Set up the equipment (as in illustration) so that the flame will almost, but not quite, touch the bottom of the can. DO NOT LIGHT THE CANDLE YET! A large can will serve as a chimney.

D. Fill the weighed can about 2/3 full of water and

re-weigh. Record this weight in the data table.

E. Record the temperature of the cold water to the nearest 0.1oC in the data table.

F. Light the candle and heat the water, stirring gently with the thermometer until the water temperature is about 15 degrees above its initial temperature (as measured in Part E). Extinguish the candle by blowing the flame out and continue to stir until the temperature stops rising. Record this highest temperature (to 0.1oC) in the data table.

G. Re-weigh the candle and lid to the nearest 0.01 gram; record this weight in the data table.

DATA TABLE Weight of candle + lid, before heating grams Weight of candle + lid, after heating grams Weight of wax burned “a” grams

Weight of can + water grams Weight of empty can grams Weight of water “b” grams

Temperature of water after heating ºC Temperature of water before heating ºC Change in temperature of water “c” ºC CALCULATIONS 1. Determine how many joules were used to warm the water. Multiply the weight of the water in grams (b) by the change in temperature (c) in oC by the specific heat of water (4.18 joules/grams·ºC). [joules = b · c · 4.18]

2. Determine the heat of combustion of wax. Divide the joules (calculated in Step 1 above) by the weight of the wax burned in (a) in grams. Heat of combustion of wax =

b · c · 4.18 ----------------------------------- = _______________ a

Summary Questions 3. How many joules would be given off if 100 grams of wax were burned? How many kilojoules is this? 4. What are some inherent sources of error in the procedure? 5. How is the heat of combustion of practical significance in the following areas? a. Choosing foods to maintain a proper diet. b. Choosing a fuel to heat your home. 6. Find the number of joules absorbed when the temperature of 20 grams of water is heated from 22ºC to 31ºC. 7. 5 grams of a substance are burned and the heat given off raises the temperature of 50 grams of water from 30ºCto 55ºC. Find the heat of combustion of the substance.

8. Draw a fully labeled potential energy diagram for the burning of wax. 9. On this diagram, label ∆H. From your calculation, what is the value of this number? 10. How was the activation energy for this reaction supplied? Indicate this on your diagram.

All nine spaces in the data table are entered using the correct number of decimal places and with the correct number of significant figures = ________(36 max). Calculations from questions one and two at ten points each = ________ (20 max). Eight summary questions times five points each = ________ (40 max). Total from above plus four point curve = ________ (100 max). Deductions for lateness to class, report handed in late, safety and participation = ________ Total Grade = ________

Name____________________________________________ SC_____-_____ Date___________ SOME CHEMISTRY OF THE THIRD-ROW ELEMENTS

Equipment: test tubes, rack, glass rods. Materials: pH indicator paper, hydroxides of Na, Mg, Al, P, S, and Cl

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

1

H He

2

Li Be B C N O F Ne

3

Na Mg Al Si P S Cl Ar

Aim: Why are some hydroxides acidic while others are basic? Let us show the model structural formula for each of these hydroxides as X-1-O-2-H where X stands for any of the Period 3 elements. (We omit other OH groups and oxygen atoms in order to keep this model formula simple). What we are going to test the water solutions of these hydroxides for pH by using a pH indicator paper. Water molecules are strong dipoles and so we can picture their positive or negative ends pulling element X or H out of our model formula by breaking bond 1 or 2. If bond 1 breaks, OH- ion is liberated from the compound and we obtain a base. On the other hand, if bond 2 breaks, H+ ion is liberated and we obtain an acid. What we are trying to find out for each hydroxide in this experiment is a) does bond 1 or bond 2 break, and b) to what extent does it break. For example, if we get a strong base (high pH), then many OH- ions were liberated and bond 1 must be weak; but if we get a strong acid (low pH), then many H+ ions were liberated and bond 2 must be weak. Of course there are other possibilities between these two extremes. A. Determine the approximate hydrogen ion concentration of each of the aqueous solutions or slurries of the

hydroxides using appropriate indicators. List the pH in the table below. (The Al(OH)3 must be freshly prepared).

1. Which of the hydroxides solutions were acidic? 2. Which of the hydroxides solutions were alkaline? 3. From the approximate pH, estimate the strength of the acid or base, using the following terms: "very

strong", "strong", "moderate", "weak", "very weak". Record your estimate in column 5 of Table B below. 4. How does the strength of the acid or base correlate with the position of element X in Period 3?

SUMMARY QUESTIONS 1. Compare the electronegativity difference between two atoms to the strength of the bond between them? 2. Could you have predicted the acidic or basic properties of each of the Period 3 hydroxides by using

electronegativity differences alone? Try it, and see if your predictions agree with the experimental results! Explain any differences.

B. Table of Results

1 2 3 4 5 6 Hydroxide Indicator

Color pH Which bond

breaks, 1 or 2?

Estimated strength of acid or base

Is it an acid or base

solution?

Na──O──H

H──O──Mg──O──H

H──O──Al──O──H

│ O │ H

O ║

H──O──P──O──H │ O │ H

O ║

H──O──S──O──H ║ O

O ║

O══Cl══O │ O │ H

Questions one, two and four were answered at ten points each = ________ (30 max). Summary Questions are answered at fifteen points each = ________ (30 max). Two points deducted for each incorrect answer given on the chart (do not deduct for answers that are consistent for incorrect answers) = ________. Final Grade = _________.

Name________________________________________ SC_____-_____ Date___________ CONVERSION OF STARCH TO GLUCOSE

Equipment: Bunsen burner, one 250mL beaker, 5 test tubes, rack, ring, ringstand, wire gauze. Materials: 5mL glucose(aq)), starch, Lugol's iodine solution, 10mL, Benedict's solution, 2mL HCl, 5-10mL Na2CO3, splints. Aim: How can we test for glucose that has been made from starch? Wear safety goggles. Do all heating in a water bath. Do not heat anything directly. PART I. HOW TO TEST FOR GLUCOSE A. To about 5mL of Benedict's solution in a test tube add 5ml of glucose solution. Heat in a water bath. This is the method used to test for glucose. (5 mL = 1/6 of the test tube) 1. Describe the result. PART II. HOW TO TEST FOR STARCH B. Heat in a water bath a pinch (enough to cover the tip of a splint) of starch in a test tube containing about 20mL of water until it becomes clear. Pour 5mL of this solution into another test tube. Cool it, and add a drop of iodine solution. This is the method used to test for starch. 2. Describe the result. C. Pour into a test tube another 5 mL. of the starch solution prepared in "B" and heat in the water bath with 5 mL of Benedict's solution. 3. . Using Benedict’s solution, how can you chemically distinguish between starch and glucose? PART III. CONVERSION OF STARCH TO GLUCOSE D. To the remainder of the starch solution prepared in "B", add 2 mL of concentrated HCl and heat in the water bath for five to ten minutes. Now add enough Na2CO3 to just neutralize the HCl. E. Pour 5 mL of the solution prepared in "D", into a test tube, cool the solution, and add a drop of iodine solution. 4. What does the result indicate?

F. To the other portion of the solution prepared in "D", add about 5mL of Benedict's solution and heat in the water bath. 5. What does the result indicate? 6. The simplest formula for starch is (C6H10O5)n; of glucose C6H12O6. Write the equation for the conversion of starch of glucose. PART IV. TESTING COMMERICAL PRODUCTS FOR GLUCOSE G. Test a sample of candy for glucose. 7. What does the test indicate about the sugar used in some candies? Summary Questions 8. Why is it important to chew such foods as bread and potatoes fully? 9. What is the purpose of the HCl in this experiment? 10. Starch is a polymer of 11. How may starch be converted to glucose? 12. State some important applications of this process.

Questions one through five at nine points each = ________ (45 max). Questions six through twelve at eight points each = ________ (56 max). Deductions for lateness to class, report handed in late and participation = ________. Total Grade = ________.

Name________________________________________ SC_____-_____ Date___________ FLAME TESTS

Aim: How does the structure of an element effect the spectral lines it produces? I. Examining Sunlight and Light bulbs.

A. Sign out a spectrometer or prism from your teacher. If you have a spectrometer be sure to point the slit on the right side at the light source. Keeping the slit pointed at the light source, look to the left of the slit and see which colored bands are produced on the scale inside the device. Use these items to examine sunlight. What colors do you see when sunlight is passed through a spectrometer or prism?

B. Using a spectrometer examine the white light emitted from the light bulbs above your head and on the teacher’s desk. Take note of the spectra produced and draw them in the space provided below. II. Examining Gas Tubes

C. Using the spectroscopes examine the spectral lines produced by hydrogen, helium and then by neon. Draw the spectral lines in the space provided below:

Hydrogen: Helium:

Neon: 1. How can you account for difference in the complexity of the spectral lines produced by these three elements? III. Flame Tests – WORK UNDER A FUME HOOD!

D. Return the spectroscopes to your teacher and cross your name off the list that you signed. Provide your teacher with an ID card and obtain a pair of diffraction gradient glasses.

E. Light a Bunsen burner and open the air ports as to produce a strong blue flame with an inner cone. Using tongs, obtain a splint (or toothpick, popsicle stick) that has been soaked in a solution containing potassium, calcium, strontium or sodium.

F. Hold the slit near the edge of the flame and take note of the color produced. Record the results of the flame test for each element in the chart below. If possible, include a rough sketch of the spectral lines produced by each element. Colors: Potassium

Calcium

Strontium

Sodium

Spectral Lines Produced (You may research this if needed): Potassium

Calcium

Strontium

Sodium

IV. Discussion 2. What is happening at the atomic level that causes spectral lines to be produced by an element? 3. How are we exciting the atoms for our flame tests? 4. How do fireworks get their color? 5. List the colors of the visible spectrum from lowest energy to highest energy. Spectra from parts I and II at ten points each = ________ (40 max). Eight answers from part III at five points each = ________ (40 max). Five questions are answered at four points each = ________ (20 max). Total from above = ________ (100 max). Deductions for lateness to class, report handed in late, safety and participation = ________ Total Grade = ________

Name___________________________________________ SC_____-_____ Date___________

ORGANIC AND INORGANIC CHEMICALS IN HOUSEHOLD ITEMS

Materials: Attached food labels.

Aim: What are the structures of the chemicals found in household items?

In this laboratory exercise you will be presented with a number of labels from household items and products that we use on a daily basis. For each label, identify the structures and/or molecular formulas of as many items that you can. Be sure to clearly label which substance is in question and which label it came from. Submit this to your teacher on a separate sheet of paper with your group members’ names on it.

Scope Mouthwash Original Flavor: Water, Alcohol (15 wt%), Glycerin, Flavor, Polysorbate 80, Sodium Saccharin, Sodium Benzoate, Cetylpyridinium Chloride, Domiphen Bromide, Benzoic Acid, Blue 1, Yellow 5 Softsoap Ultra Rich Shea Butter: Water, Sodium Laureth Sulfate, Acrylate Copolymer, Cocamidopropyl Betaine, Fragrance, Sodium Chloride, Glycol Distearate, Sodium Hydroxide, Laureth 4, DMDM Hydantoin, Butyrospermum Parkii (Shea Butter), Tetrasodium EDTA, Gelatin, Acacia Senegal Gum, Mica, Citric Acid, Titanium Dioxide, Iron Oxides Crest Toothpaste: Active Ingredients: Sodium Fluoride (0.15% w/v fluoride ion) Inactive Ingredients: Water, Hydrated Silica, Sorbitol, Glycerin, Tetrapotassium Pyrophosphate, PEG 6, Disodium Pyrophosphate, Tetrasodium Pyrophosphate, Sodium Lauryl Sulfate, Flavor, Xanthan Gum, Sodium Saccharin, Carbomer 956, Titanium Dioxide, FD&C Blue 1 Pantene 2 in 1 Shampoo and Conditioner: Water, Ammonium Laureth Sulfate, Ammonium Lauryl Sulfate, Sodium Chloride, Glycol Distearate, Cocamide MEA, Panthenol, Panthenyl Ethyl Ether, Lysine HCl, Methyl Tyrosinate HCl, Histidine, Dimethicone, Fragrance, Cetyl Alcohol, Sodium Citrate, Sodium Benzoate, Guar Hydroxypropyltrimonium Chloride, Disodium EDTA, Hydrogenated Polydecene, Trimethylolpropane Tricaprylate/Tricaprate, Citric Acid, Methylchloroisothiazolinone, Methylisothiazolinone, Ammonium Xylenesulfonate

Capsaisin HP: Active Ingredients: Capsaicin (.1%) Inactive Ingredients: Benzyl Alcohol, Cetyl Alcohol, Glyceryl Monostearate, Isopropyl Myristate, PEG 100 Stearate, Petrolatum, Sorbitol, Water Speed Stick Deodorant: Active Ingredients: Aluminum Zirconium Tetrachlorohydrex Gly (15.5%) (Antiperspirant) Inactive Ingredients: Cyclomethicone, Stearyl Alcohol, Mineral Oil, Talc, Hydrogenated Castor Oil, PPG 14 Butyl Ether, Steareth 100, Steareth 2, Behenyl Alcohol LA Looks Hair Gel: Aqua (Water), VP/VA Copolymer, Methacryloyl Ethyl Betaine/Acrylates Copolymer, Carbomer, Propylene Glycol, Panthenyl Ethyl Ether, Keratin Amino Acids, Glycerin, Aminomethyl Propanol, Glyceryl Polymethacrylate, Palmitoyl Oligopeptide, Rahnella Soy Protein Ferment, Sorbitol, Alcohol, PEG 75 Lanolin, PEG 40 Hydrogenated Castor Oil, PPG 5 Ceteth 20, PEG 8, Disodium EDTA, Benzophenone 4, Methylparaben, Methylchloroisothiazolinone, Methylisothiazolinone, Parfum (Fragrance), Benzyl Benzoate, Benzyl Salicylate, Limonene, Linalool, Hexyl Cinnamal, CI 19140 (Yellow 5), CI 15510 (Orange 4) Rubbing Alcohol: Isopropyl Alcohol (70%)

Name________________________________________ SC_____-_____ Date___________ THE FORMULA FOR A HYDRATE

Equipment: triple beam balance, 6" test tube, 150mL beaker, clamp, ringstand, Bunsen burner, tongs, test tube rack Materials: 10g blue Cu2SO4 crystals Aim: How can we determine the percent hydration of a salt? Many salts exist in two or more forms with respect to water content. Ionic crystals containing covalently bonded water (called HYDRATES) may be converted, by heating, to the form lacking water molecules (ANHYDRATES). A DOT indicates the water is not directly part of the compound but is bonded to it. For example: CaBr2•3H2O(s) + HEAT = CaBr2(s) + 3H2O(g). In this experiment you'll determine the value of x in one or both of the following compounds: CuSO4•xH2O and

BaCl2•xH2O SAFETY WARNING: All chemicals in this lab are toxic if eaten. Wash your hands before leaving. Wear eye protection at all times! Procedure Note: All readings should be between two or three decimal places depending on your balance. I Find the mass (weight) of a salt sample. 1) Zero your balance and weigh a 6" test tube in

a 150 mL beaker (both must be dry). Record the weight in spaces (B) and (D) (lines 2 and 5).

2) Fill the test tube to a height of about 1.5" with

the copper compound (less than 10g. of the blue crystals). Weigh the test tube again (in the beaker) and record the weight in space (A) (line 1).

3) Subtract the container mass (weight) from the

total to get the sample mass. (A-B=E). II Dehydrating the Salt 4) Clamp the test tube horizontally on a ringstand

in the hood. The clamp should be near the mouth of the tube and the mouth should point into the hood. Tap the test tube to spread the crystals out over the test tube so that more surface area is exposed. Be careful not to spill any crystals.

Data Table – Record Measurements here.

CuSO4∙XH2O W = wt of beaker + test tube

1. (A) W + CuSO4∙XH2O g

2. (B) - W g

3. (E) CuSO4∙XH2O g

4. (C) W + CuSO4 g

5. (D) - W g

6. (F) CuSO4 g

7. (E) CuSO4∙XH2O g

8. (F) - CuSO4 g

9. (G) XH2O g

5) Using the hottest flame (see exp. 1 Part F), heat the compound vigorously for at least 10 minutes. Hold the burner by the base and keep the flame moving under the sample in the test tube to avoid "hot spots." Also heat the empty part of the tube to drive out any water that may have condensed there. Caution: Avoid heating the clamp. What changes do you see as the crystals are being heated? What do these changes indicate?

6) Use the clamp to transfer the hot test tube to the 150 mL beaker. Determine the weight and enter in space (C)

(line 4). III Clean-up and Disposal 7) When the tube and sample have cooled, empty your anhydrate into the designated container in the teacher's

fume hood. Wash and invert the test tube you used. Put your balance near the center of the table with a major rider off zero.

IV Calculating the Hydrate Formula 8) To complete the calculations use the readings from the data table in the calculation box.

CuSO4∙XH2O

(F) CuSO4 g

gfm CuSO4 = 160g/mole

(F) 160

= moles CuSO4

mole

(G) XH2O g

gfm H2O =18g/mol

(G) 18

mole

X = Moles H2O Moles CuSO4

mole

3. Calculate the percent error of your experiment (refer to Reference Table T). Show all work. Observed value – Accepted value Percent error = ------------------------------------------------- x100% = Accepted value Observed value = __________ Accepted Value = __________ Percent error = __________

VI Summary Questions 1. What are the main sources of error in this experiment? 2. What evidence is there for a chemical change? 3. Suppose you purchased $100 dollars of this hydrated copper sulfate salt. How much of your money was spent on purchasing water? Data chart on page one shows correct number of decimal places and significant figures used = ________ (36 max). Four calculations were completed on page two at ten points each = ________ (40 max). Three summary questions were answered at eight points each = ________ (24 max). Deductions for lateness to class, report handed in late, safety and participation = ________ Total Grade = ________.

Name___________________________________________ SC_____-_____ Date___________ Laboratory Exercise on Graphing Periodic Trends

Aim: What are some trends that we observe across the periodic table? In this exercise you will be using Reference Table S to discover the trends that occur in the Periodic Table by constructing four graphs. The construction of these graphs should support the idea that “the chemical properties of elements are periodic functions of their atomic number.” Properties to be investigated: 1. Electronegativity – the ability to attract electrons. There is no unit for electronegativity. 2. First Ionization Energy – energy required to remove the outermost electrons from an atom. The

units are kilojoules/mole. 3. Atomic Radius – The units are in picometers (10-12 meters). 4. Density – Mass per unit volume. Units are grams per milliliter (g/ml) or grams per cubic centimeter

(g/cm3). Procedure: You will be preparing four separate graphs of the above properties. • The horizontal axis for all four graphs will be atomic numbers 1-20, 31-38 and 49-56. Yes, the

transition elements have been skipped over. However, you will not leave any gap between the break in numbers on the x-axis.

• Each vertical axis will be one of the four given properties listed above. 1. For the graph of Atomic Number v. Electronegativity: The numbers on the y-axis range from 0 to 4.0

with a change in the value of y by 0.2 units per line. 2. For the graph of Atomic Number v. First Ionization Energy: The numbers on the y-axis range from 0

to 2400 with a change in the value of y by 80 units per line. 3. For the graph of Atomic Number v. Covalent Radius: The numbers on the y-axis range from 0 to 250

with a change in the value of y by 10 units per line. 4. For the graph of Atomic Number v. Density: The numbers on the y-axis range from 0 to 7.5 with a

change in the value of y by 0.5 units per line. Be sure each graph has a title, labeled axes and that the units of the properties are included! Lab Questions: 1. Name the following groups: Group 1 (IA) _________________________ Group 2 (IIA)_________________________ Group 17 (VIIA)______________________ Group 18 (0) _________________________

2. How is the modern Periodic Table arranged? 3. What are the vertical columns called? 4. What are the horizontal rows called? 5. What three things does the atomic number of an element tell us? 6. What are the majority of the elements on the Periodic Table classified as? 7. What are the “middle elements” using the “d” orbitals called? 8. What are the charges (sign and value) of neutrons, protons and electrons? 9. What two names do we have for the elements that show the properties of both metals and nonmetals?

10. Which family is considered to be “highly inactive”? Why is this so? 11. On the back of each graph that you constructed, explain the trends you see for that particular graph

from: a) Left to right in a period b) Top to bottom in a group

Grading for the project follows a one point deduction for each error made on the graphs and for errors made while answering the questions. For example, a student can lose up to six points for questions number eight above and two points for question nine. Deduct a point for each missing unit on a graph, a missing title, etc. Deduct 25 points for a missing graph. Four graphs attempted at twenty-five points each = ________ (100 max). Deductions for errors in graphs and questions at one point each = ________. Lateness deduction = ________. Total grade = ________.

Name____________________________________________________________Date___________ FACTORS THAT AFFECT RATES OF REACTION

Materials: Beakers, ring stand, iron ring, gauze, 0.1 M HCl, 0.5 M HCl, Magnesium Strips, Food coloring, Alka Seltzer Tablets, tea bags, sugar cubes. Aim – What are some factors that can alter the rate of reaction? A. Obtain two beakers. In one beaker place about ¾ cup of cold water. Place an equal amount of water into the other beaker and heat gently. At the same time, drop a sugar cube into each beaker. Allow to stand for five minutes. What do you observe? In terms of the collision theory, why does this occur? B. Obtain two beakers. In one beaker place about ¾ cup of cold water. Heat an equal amount of water in the other beaker. At the same time drop a tea bag into each beaker. Allow to stand for three to five minutes. What do you observe? In terms of the collision theory, why does this occur? C. Be prepared to time this experiment. Obtain two beakers. Place an equal amount of tap water into each beaker. Into one beaker drop a sugar cube. At the same time drop a sugar cube that has been crushed up into the other beaker. Allow to stand. Which one dissolves faster? How long did it take for the sugar to dissolve in each beaker? In terms of the collision theory, why does this occur? D. Obtain two beakers. Place an equal amount of tap water into each beaker. Drop a sugar cube into each beaker at the same time. Stir the contents of one beaker with a stir rod for two minutes. Which sugar cube dissolves faster? In terms of the collision theory, why does this occur?

E. Examine the substances on the teachers’ desk. How would you design three different experiments using these items? Complete the write up for these experiments in the spaces provided below. Your teacher will demonstrate some of the experiments deemed safe. 1. 2. 3. Conclusions: 1. What were the factors that affected the rate of reaction in this experiment? How did each factor affect the

rates? 2. What are some other factors that affect rate of reaction? How do they alter rates of reaction? 3. What measures would you take as to slow down a reaction? Name at least five methods.

Name________________________________________ SC_____-_____ Date___________ HYDROLYSIS

Equipment: 5 test tubes, rack, glass rod. Materials: red and blue litmus paper, 1-2g of the following: NaCl, Na2CO3, Na3PO4, Al2(SO4)3, AlCl3, NH4Cl, (NH4)2SO4, wood

splints (as spatulas).

In this experiment you will learn why some highly basic household cleaning agents do not have to contain KOH or NaOH and why "excess stomach acid" can be neutralized with NaHCO3.

A. Dissolve a pinch of NaCl in about 1/3 test tube of water. Dissolve similar amounts of Na2CO3 and

Na3PO4 in each of two other test tubes. Test each of these solutions with red and blue litmus paper and with pH paper. Enter your findings in the table below.

1. Which ion must be in excess in order to cause the results observed in two of the test tubes? 2. What is the source of this ion? B. Dissolve a pinch of Al2(SO4)3 or AlCl3 in about 1/3 test tube of water and a similar amount of NH4Cl

or (NH4)2SO4 in another test tube of water. Test each solution with red and blue litmus paper and with pH paper. Enter your findings in the table below.

3. Which ion must be in excess in order to cause the results observed? 4. What is the source of this ion?

Salt Solutions Used Effect On Litmus pH Ion In Excess

NaCl

Na2CO3

Na3PO4

5. It is often convenient to have a fast and simple way of predicting the result of hydrolysis of a given salt solution. One way is to consider hydrolysis the reverse of neutralization and then to compare the

strengths of the acid and base produced according to the Arrhenius Theory. Neutralization: acid + base = salt + water Hydrolysis: salt + water = acid + base Complete and balance the following equations. If applicable, circle or underline the stronger of the acid and base. Is the solution expected to be acidic or basic? __Na2CO3 + __HOH = ___ ________ + ___ _______ Acidic or Basic? __________ __Na3PO4 + __HOH = ___ ________ + ___ _______ Acidic or Basic? __________ __KNO3 + __HOH = ___ ________ + ___ _______ Acidic or Basic? __________ __NH4Cl + __HOH = ___ ________ + ___ _______ Acidic or Basic? __________ 6. Even though the Arrhenius Theory correctly predicts the pH of a given salt solution, chemists think the proton transfer ideas of the Brönsted-Lowry Theory are closer to what is really happening in hydrolysis. Use your knowledge of Brönsted-Lowry Theory to predict the products of the following reactions. You may also want to consult the previous version of Reference Table L from previous versions of Reference Tables for Chemistry. CO3

-2 + HOH = _______________ + _______________ PO4

-3 + HOH = _______________ + _______________ NH4

+ + HOH = _______________ + _______________ C. Use both Arrhenius Theory and the Brönsted-Lowry Theory to predict the pH of solution due to the hydrolysis of the following salt solutions. Test your predictions with pH paper and litmus paper.

Table II Salt Solution

Used Prediction pH of Solution Reason / Explanation

NaNO3

CuSO4

FeSO4

K2SO4

Conclusions According to the Brönsted-Lowry Theory, hydrolysis is the slight reaction between the

_____________________ or _____________________ of a salt, and water to give an excess of

_____________________ ions or _____________________ ions.

Fifteen spaces correctly answered in parts A and B = ________ (30 max). Twenty eight answers to question number five at one point each = ________ (28 max).

Six answers to question number six at two points each = ________ (6 max).

Twelve answers to Table II at two points each = ________ (24 max).

Conclusion questions at three points per space = ________ (12 max).

Deductions for lateness to class, report handed in late, safety and participation = ________. Total Grade = ________.

Name________________________________________ SC_____-_____ Date___________ RATE OF A CHEMICAL REACTION: THE IODINE CLOCK

Equipment: 5 test tubes, test tube rack, and one of each of the following: 50 mL and 25 mL graduated cylinders, rubber stopper, 250 mL beaker Materials: 75 mL 0.01M KIO3, 75mL sat'd H2SO3 + starch, H2O at room temp. Aim: What are some factors that effect reaction rate? Part I – Procedure A. Prepare a set of test tubes of varying amounts of 0.01M KIO3 (one of the reactants) as outlined in the table

below. B. To test tube 1, add 15 mL of saturated H2SO3 and stir several times, recording the exact time of addition of

H2SO3. Use a watch with a second hand or other timing device. (All solutions and water should be at room temperature).

1. Record the number of seconds needed for the solution to become blue. 2. Calculate the rate, for each trial, which would be proportional to the time, 1000/t (t=time in seconds).

Record in the table below. C. Repeat the same procedure with the remaining test tubes. Using 15 mL of H2SO3 in all cases. It is advisable to

complete the work with one test before going on the next. (If time is short, however, one partner of each pair of students could do one test tube so that two would be done simultaneously).

Test Tube

0.01M KIO3 H2O Saturated H2SO3

Molarity of KIO3 after mixing.

Time (sec.) Rate = 1000/t

1

5 mL 30 mL 15 mL 0.001 M

2

10 mL 25 mL 15 mL 0.002 M

3

15 mL 20 mL 15 mL 0.003 M

4

20 mL 15 mL 15 mL 0.004 M

5

25 mL 10 mL 15 mL 0.005 M

D. Prepare a graph by plotting the concentration of KIO3 (0 to 0.05 with 0.01 intervals) as the abscissa and rate (1000/t) as the ordinate. Be sure to plot the data with a best fit line.

0 0.001 0.002 0.003 0.004 0.005 Molarity of KIO3

PART II- DISCUSSION Equations involved in the above reaction are: (a) 3H2SO3 + HIO3 → 3H2SO4 + HI This equation (a) is a slow reaction and will be completed when the H2SO3 is depleted. The concentrations of

the reactants being such that the KIO3 is in excess in all the tubes.

(b) 5HI+ HIO3 → 3I2 + 3H2O This reaction (b) occurs the moment that the H2SO3 has been removed and is an instantaneous reaction.

Starch is present in the H2SO3 solution and immediately detects the presence of iodine. Although this reaction is instantaneous, I2 cannot exist if any H2SO3 is still present because of the following reaction:

(c) H2SO3 + H2O + I2 → H2SO4 + 2HI Summary Questions 1. From the line drawn in your graph, what is the relationship between the concentration of KIO3 and the rate of the

chemical reaction?

2. Define “rate of reaction”. 3. Explain why the results of the above experiment may have been distorted if all parts of the experiment were not

performed at the same temperature. 4. What are some other factors that could affect the rate of reaction? 5. In any chemical reaction, as time continues, what should happen to the speed of the forward reaction as time

goes on? Explain. Five rates were correctly calculated from the number of seconds recorded = ________ (25 max). The y-axis was set up correctly as to accommodate the rates obtained = ________ (5 max). Five points were plotted correctly = ________ (25 max). A best fit line was drawn = ________ (5 max). Five summary questions at eight points each = ________ (40 max). Deductions for lateness to class, report handed in late, safety and participation = ________. Total Grade = ________.

Name___________________________________________ SC_____-_____ Date___________ LABORATORY TECHNIQUES

Equipment: triple beam balance, beaker, graduated cylinder, glass rod, funnel, Bunsen burner, coin, Materials: 2g NaCl, sand, matches Aim: What are some ways we can separate mixtures, measure weights, and determine volumes? Part I. Weighing (or massing). In most of your lab work you will use triple beam

balances. The platform type is accurate to 0.01g. The swinging pan type is accurate to 0.001g. Both are sensitive instruments and must be handled with care. Your teacher will demonstrate the proper technique.

1. Never place hot objects or chemicals on the pan.

2. Weigh solid chemicals on a piece of paper or dry container.

3. Place objects on the pan and change weights gently.

4. Leave the balance clean. 5. Move a large rider mass off zero when

there is no arresting mechanism to stop swinging.

Fill in one column based upon the balance that you have.

.01 grams .001 grams

1

2

3

4

5

Average

A. Weigh a coin five times, record each reading, and then calculate the average weight. Place your data in the appropriate column above. Why are five readings needed?

Five readings with correct decimal places ______ (15) Average calculated with correct Sig Figs ______ (5) Question Answered ________ (5)

B. On a small piece of properly folded paper,

weigh out approximately 2 grams of a mixture of a salt and sand. Record the weight of the mixture to 0.01g or 0.001g, depending on the type of balance you are using. Transfer the mixture to a clean beaker and save it for later.

.01 or .001 grams Gross

Tare (Wt. of paper)

Net

Three numbers with correct decimal places calculated with correct significant figures ______ (5) Part II. Measuring the Volumes of Liquids. Liquids are most easily measured by volume instead of by weight. (Why?) Common measuring devices are:

graduated cylinders, burets, pipettes, and volumetric flasks. You will be using the graduated cylinder most often. Your teacher will demonstrate how to use this device and how to read a meniscus.

C. Measure 25 mL of water.

Part III. Separation. Using the materials that are present on your teacher’s desk and at your laboratory station, design an

experiment that would allow you to separate the sand and salt mixture. CHECK YOUR PROCEDURE WITH YOUR TEACHER BEFORE PERFORMING YOUR EXPERIMENT.

My procedure is to: (20 max). D. Now that you have finished the separation of the sand and salt answer the following: 1. What is a residue? What does the residue consist of? 2. What is a filtrate? What does the filtrate consist of? 3. How did you separate the components of the filtrate? Three questions at five points each equals ________ (15 max). Part IV. Calculating Density. Obtain five pennies from the teachers’ desk. How would you calculate the density of the five pennies? Design an experiment that uses the equipment you have used today in lab. CHECK YOUR PROCEDURE WITH YOUR TEACHER BEFORE PERFORMING YOUR EXPERIMENT. Perform the experiment when cleared to do so. My procedure is: My data collected is:

My calculation of the density of a penny is: Three steps at five points each = ________ (15 max). SUMMARY QUESTIONS Each question is worth four points = ________ (20 max). 1. Make a check list for leaving your laboratory station in proper condition after use. 2. What is the meniscus of a liquid? Describe it. Draw an example. 3. Why should a volume measuring device be read at eye level? 4. Describe, in detail, how one may separate the components of rubbing alcohol (a mixture of two liquids named 2-

propanol and water.) 5. Behind the teacher's desk is a container of NaHCO3 (baking soda, sodium carbonate). What is it used for? Total points from above = ________ (100 max) Deductions for lateness to class, report handed in late, safety and participation = ________ Total Grade = ________

Name __________________________________________________ Lab Section ____________

In

PhysicalState Solid

Densitying/mL 7.31

Hardness Verysoft

Conductivity Medium

Meltingpointin°C 157

Solubilityinwater None

Color Silverywhite

Pb

PhysicalState Solid

Densitying/mL 11.35

Hardness Somewhatsoft

Conductivity Poor



Color Silverywhite

Ar

PhysicalState Gas

Densitying/mL 0.00178

Hardness None

Conductivity Verypoor

Meltingpointin°C -189


Color Colorless

Ga

PhysicalState Solid

Densitying/mL 5.904

Hardness Soft

Conductivity Medium



Color Silvery

Cs

PhysicalState Solid

Densitying/mL 1.87

Hardness Soft

Conductivity Good


Solubilityinwater Reactsviolently

Color Silverywhite

Unknown1

PhysicalState Solid

Densitying/mL 2.33

Hardness Brittle

Conductivity Intermediate



Color Gray

Unknown2

PhysicalState Gas


Hardness None



Solubilityinwater Slight

Color Paleyellow

Unknown3

PhysicalState Solid

Densitying/mL 1.53

Hardness Soft

Conductivity Good


Solubilityinwater Reactsviolently

Color Silverywhite

Unknown4

PhysicalState Gas


Hardness None




Color Colorless

Unknown5

PhysicalState Solid

Densitying/mL 19.3

Hardness Soft

Conductivity Excellent



Color Gold

Unknown6

PhysicalState Solid

Densitying/mL 2.54


Conductivity Good


Solubilityinwater Reactsrapidly

Color Silverywhite

Unknown7

PhysicalState Solid

Densitying/mL 5.32

Hardness Fairlybrittle

Conductivity Fairtopoor



Color Gray

Unknown8

PhysicalState Solid

Densitying/mL 1.74

Hardness Medium

Conductivity Good


Solubilityinwater Reactsslowly

Color Silverywhite

Unknown9

PhysicalState Solid

Densitying/mL 11.85

Hardness Verysoft

Conductivity Medium



Color Silverywhite

I2

PhysicalState Solid

Densitying/mL 4.93

Hardness Soft



Solubilityinwater Negligible

Color Bluish-black

Li

PhysicalState Solid

Densitying/mL 0.534

Hardness Softclaylike

Conductivity Good


Solubilityinwater Reactswithwater

Color Silver

Ag

PhysicalState Solid

Densitying/mL 10.50





Color Silver

Cu

PhysicalState Solid

Densitying/mL 8.96





Color Orange-Brown

C

PhysicalState Solid

Densitying/mL 2.10

Hardness Softandbrittle

Conductivity Good



Color Black

Cl2

PhysicalState Gas


Hardness none



Solubilityinwater Slight

Color Greenishyellow

He

PhysicalState Gas


Hardness None




Color Colorless

Na

PhysicalState Solid

Densitying/mL 0.971

Hardness Softandclaylike

Conductivity Good



Color Silver

Ca

PhysicalState Solid

Densitying/mL 1.57

Hardness Medium

Conductivity Good


Solubilityinwater Reacts

Color Silverywhite

Be

PhysicalState Solid

Densitying/mL 1.85

Hardness Brittle




Color Gray

Sn

PhysicalState Solid

Densitying/mL 7.31


Conductivity Good



Color Silvergray

Ne

PhysicalState Gas


Hardness None




Color Colorless

Br2

PhysicalState Liquid

Densitying/mL 3.12

Hardness None


Meltingpointin°C -7.2


Color Reddishbrown

K

PhysicalState Solid

Densitying/mL 0.86

Hardness Softandclaylike

Conductivity Good



Color Silver

Ba

PhysicalState Solid

Densitying/mL 3.6

Hardness Soft

Conductivity Good


Solubilityinwater Reactsstrongly

Color Silverywhite

Xe

PhysicalState Gas


Hardness None




Color Colorless

Name________________________________________ SC_____-_____ Date___________ THE NATURE OF BURNING

Equipment: glass plate, Bunsen burner, wood splints, funnel, bottles, trough, generator, deflagrating spoon Materials: charcoal, 6% H2O2 MnO2, 20mL Ca(OH)2 Aim: How can we prepare oxygen gas? What are some properties of oxygen gas? Put enough manganese dioxide in a generating bottle to cover the bottom to a depth of ¼ - ½ inch. Add enough hydrogen peroxide solution to obtain a steady flow of oxygen. Make sure the solution level is higher than the bottom end of the thistle tube - why? Collect 3 bottles of gas by water displacement. Let the gas in the first bottle escape.

1. Why do we let it escape?

2. Write the equation for this reaction.

3. Was it exothermic or endothermic? How do you know? B. Into an empty bottle (1), pour about 20 mL (about an inch) of a saturated solution of Ca(OH)2

(limewater). Cover the bottle with a glass plate and shake the solution for about 10 seconds. 4. Do you observe any change in the solution? 5. What does this show? C. Ignite a splint and let it burn in the bottle you used in B (1). Remove the splint, cover the bottle, and

shake the solution again for ten seconds. 6. Describe any change in the solution. Note: this effect is used to test for the presence carbon dioxide

gas. D. Add 20 mL of saturated Ca(OH)2 solution to one of the bottles (2) of oxygen gas you prepared. Let an

ignited splint burn in the bottle for a few seconds, as before, cover, and shake.

7. Did the splint appear to burn the same way in air as in pure oxygen? Describe your observation and

explain why there might be a difference?

8. Compare the solutions in the bottles from Part D and Part C.

9. What gas was the product of burning in each case? How can you tell? E. Put a few pieces of charcoal in a deflagrating spoon. USE FUME HOOD! Hold the spoon in the HOTTEST part of the Bunsen flame until the charcoal glows red. Lower the deflagrating spoon midway into the second bottle (3) of oxygen you prepared. When the reaction seems to have stopped, remove the spoon, add 20 mL of Ca(OH)2 solution cover, shake, and observe. 10. What gas is produced when charcoal burns in oxygen? 11. Write an equation for this reaction. Part F is optional on the teacher’s behalf. Do not hold students accountable for this part if it is not performed. F. TEACHER DEMONSTRATIONS: THESE MUST BE DONE INSIDE THE FUME HOOD! Burning metal: Tare steel wool (at least 7g) in a 2L beaker on a triple beam balance. What do you think will happen to the mass after burning? What happens to the mass? Explain why there is a difference.

SUMMARY QUESTIONS 12. When a substance burns it combines chemically with the element to form compounds known as . These reactions are always (exothermic, endothermic). 13. A substance burns in oxygen than in air. 14. Our atmosphere is approximately 78% nitrogen. What part does this element play in burning? 15. Draw and completely label a potential energy diagram for the reaction used to make oxygen gas. Be sure to show the effect of the catalyst. Fifteen questions at six points each plus a ten point curve = ________ (100 max). If teacher demo was completed: Eighteen questions at five points each plus a ten point curve = ________ (100 max). Deductions for lateness to class, report handed in late, safety and participation = ________ Total Grade = ________

Name________________________________________ SC_____-_____ Date___________ ESTERS

Equipment: 4 small test tubes, one 150 mL beaker, ring, ringstand, Bunsen burner. Materials: 50mL CH3OH, 10mL of the following: 2 methyl butanol, acetic acid, sulfuric acid, ethanol, butanol, salicylic acid, wooden splints. Aim: How can we prepare and organic ester? What are some characteristics of organic

esters? WARNING: Sulfuric acid is a powerful dehydrating agent. It can injure your eyes and skin

and damage clothing. AS ALWAYS, Safety goggles must be worn. Rinse with plenty of water. PROCEDURE: Set-up a hot water bath inside your hood. Keep your test tubes inside the hood

as much as possible. A. In a small test tube place 10 drops of methyl alcohol. Add some salicylic acid, as much as

you can put on the end of a wooden splint. Add two drops of concentrated sulfuric acid. Note the odor, if any, of the three materials. Place the test tube in a hot water bath for five minutes.

1. How does the odor compare with that of the original materials? 2. What substance has a similar odor? 3. What is the chemical name for this material? B. Prepare the following mixtures in small test tubes and heat for five minutes in a hot water

bath. a. 10 drops of 2-methyl-1-butanol (amyl alcohol) + 5 drops of ethanoic acid (acetic acid) + 2

drops of concentrated H2SO4 b. 10 drops of ethanol (ethyl alcohol) + 5 drops of ethanoic acid + 2 drops of concentrated

H2SO4 c. 10 drops of butanol + 5 drops of ethanoic acid + 2 drops of concentrated H2SO4 4. Describe the products you made in a, b, and c above, in terms of familiar materials.

a. b. c. SUMMARY QUESTIONS

5. What is the general formula for alcohols?

6. What is the general formula for organic acids?

7. Complete the following:

8. What is the name of inorganic compound always produced in a reaction between an alcohol and an acid?

9. What kind of an organic compound is produced in a reaction between an alcohol and an acid?

10. Complete the following equations: alcohol + carboxylic acid → 2-methyl-1 butanol + ethanoic acid → methanol + salicylic acid → butanol + ethanoic acid → ethanol + butanoic acid → 1, 2, 3-propantriol + nitric acid → glycerol + nitric acid →

11. Why is sulfuric acid used in making esters?

Questions one through four at ten points each = ______ (40 max). Questions five through nine and eleven at six points each = ________ (30 max). Question ten at four points each = ________ (28 max). Two points added to the above. Deductions for lateness to class, report handed in late and participation = ________. Total Grade = ________.

Name________________________________________ SC_____-_____ Date___________ MOLECULAR MODELS OF ORGANIC MOLECULES

Equipment: chemistry model kit Aim: How can organic models help us become familiar with the many different compounds that can

be formed. Carbon has four valence electrons and is able to form single, double and triple covalent bonds with many elements and itself. SOME RULES TO FOLLOW IN MAKING MOLECULAR MODELS Use black spheres for carbon atoms, yellow for hydrogen, green for chlorine and orange for bromine. One covalent bond (one pair of shared electrons) is represented by one stick. Use two or three steel springs when double or triple bonds are constructed. Each carbon atom must share four pairs of electrons (have four bonds). PART I: CHAIN HYDROCARBONS (ALIPHATICS) The hydrocarbons are divided into homologous series. This grouping makes their study easier. A homologous series is a series of compounds in which each member differs from the next by the same number of atoms. A. Alkanes: These hydrocarbons contain only single covalent bonds.

# of C’s # of H’s Molecular Formula Structural Formula Methane

1 4 CH4

Ethane

Propane

Make a model of methane and the other hydrocarbons listed. Determine the molecular and structural formulas and fill in the chart. Consult your Reference Tables for help. 1. What are the molecular formulas for n-butane and n-pentane? 2. The general formula for an alkane is_______________. B. Alkenes: These hydrocarbons contain one double bond.

# of C’s # of H’s Molecular Formula

Structural Formula

Ethene

Propene

Construct models of the alkenes listed above (use the springs for the double bonds) and fill in the chart.

3. The general formula for an alkene is _______________. C. Alkynes: These hydrocarbons contain one triple bond.

# of C’s # of H’s Molecular Formula

Structural Formula

Ethyne

Propyne

Construct models of the alkynes listed above (use the springs for the triple bonds) and fill in the chart. 4. The general formula for an alkyne is _______________. 5. Another name for ethyne is ___________________________. PART II: AROMATIC HYDROCARBONS Benzene is an example of an important ring compounds. Its formula is C6H6. Make a model of benzene and draw its structural formula below. (2) PART III: ISOMERS 6. Define: Isomers (2) D. Construct a model of n-butane. How many different compounds can you make using the same number of atoms of Carbon and Hydrogen? ______ Draw the isomers below. You teacher will help you with any new names. (6)

PART IV: ALKYL HALIDES A halogen may replace a hydrogen atom. Construct and draw the following: H. Chloromethane (2) I. Dichloromethane (2) J. Bromoethane (2) K. Construct models for the compounds whose molecular formula is C2H4Cl2. How many isomers are there ______? Draw their structures. (6) M. Draw structures for the isomers of C3H6BrCl and name them. (20)

E. Construct a model of n-pentane. How many different compounds can you make using the same number of atoms of Carbon and Hydrogen? ______ Draw the isomers below. You teacher will help you with any new names. (8) F. Construct the two molecules that have the molecular formula of C2H6O (oxygen is red). Draw their structures below. (4) G. Construct a model of 2,2-dimethylbutane.

Draw its structural formula. (2)

PART V: POLYMERS Construct a model of ethene. Pull one end of one of the springs out of a carbon atom. Your neighbors will do the same. Now connect your molecule to your neighbors’ molecule. Draw what you have made. Place the models back in their appropriate kit. (2) 7. What is the difference between a polymer and a monomer? (4)

One point deduction for an incorrect answer from Part I: A, B and C (max 31 points lost) = ________ points lost. Deduct a maximum of the number of points listed in Parts II through IV for incorrect answers to each question = ________ (31 parts at two points each, 62 max lost points). Deductions for lateness to class, report handed in late and participation = ________. Total Grade = ________.

Name ____________________________________________________Class_____________ Regents Chemistry Spring Term

Aim: How can we determine if a precipitate will form in a double replacement reaction?

In this lab you will mix together various different ions to determine how the solubility rules listed in Table F were created. Using your dropper bottles and well plate, you will mix together each of the substances in the table. Be sure to thoroughly CLEAN your well plate. Look at the table below and mix one substance from the left side of the table with one substance from the top of the table. Note the five exceptions. When using your dropper bottles be sure to hold the bottle 1-2 cm ABOVE the plate so that the tip DOES NOT touch the solution in the plate – if there is contamination in the bottles you will not be able to properly determine the solubility rules. Be sure to write a complete reaction and identify the solid precipitate that forms, if any. In each box write the number for each reaction for the double replacement the two substances combined. Make note of the precipitates formed, if any. NaOH CaCl2 NaI Na3PO4 AgNO3

BaCl2

XXXXXXXXXXXX

CuSO4

Na2CO3 XXXXXXXXXX

XXXXXXXXXXXX

XXXXXXXXXXXX

NH4OH XXXXXXXXXX

In each line write a complete reaction for the double replacement the two substances combined corresponding to the number in the box 1. 2. 3. 4 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15.

1. Which ions always dissolved without exception?

2. Which ions dissolved most of the time? What were the exceptions?

3. Which ions precipitated without exception?

4. Which ions formed a precipitate most of the time? What were the exceptions?

5. Name all of the precipitates that formed in this lab with appropriate IUPAC nomenclature.

6. Provide the formulas for all of the precipitates that formed in this lab.

7. Draw a picture of KCl dissolved in water. Be sure to show at least five water molecules and 3KCl showing proper orientation of the molecules and ions.

8. How is the way that sugar dissolves in water different from the way salt dissolves in water? In which process are bonds broken? In which process are intermolecular attractions broken?

9. Suppose you wanted to prepare 25 mL of a 1.5 M solution of copper (II) sulfate pentahydrate. How many moles of the salt are needed? How many grams is this?

One point for each box with correct notation for precipitates = ________ (15 max). Three points for each correct equation = ________ (45 max). Four point four(4.4) for each answer in the conclusion questions = ________ (40 max). Deductions for lateness to class, report handed in late, safety and participation = ________. Total Grade = ________.

Name________________________________________ SC_____-_____ Date___________ PREPARATION AND PROPERTIES OF CARBON DIOXIDE

Equipment: generator set (made from: 6oz. bottle, 2 holed rubber stopper, thistle tube, delivery tube), test tube, ringstand, clamp.

Materials: 2-3 marble chips (CaCO3), 10mL HCl, 10mL Ca(OH)2 (limewater), blue litmus paper Aim: What are some of the chemical and physical properties of carbon dioxide? A. Fit a six-ounce bottle with a two-hole rubber stopper, thistle tube, and delivery tube. Set up the apparatus for the collection of a gas by the upward displacement of air. Put some marble chips into the bottle to a depth of about one inch. Add enough dilute HCl to cover the marble chips and the bottom of the thistle tube. Add small amounts of HCl if the reaction slows down. 1. Which reactant is the source of the CO2? 2. Write an equation for the formation of CO2 from HCl and CaCO3. 3. Consider your answer to question 2. Relate this to acid rain. B. Collect a bottle of CO2 by upward displacement of air. This bottle should be dry. The bottle is filled with gas

when a lighted splint goes out as soon as it is lowered into the neck of the bottle. 4. Why is the gas collected by the upward displacement of air? 5. Explain the effect of CO2 on the lighted splint. 6. How do the answers to questions 4 and 5 show the usefulness of CO2 in firefighting? C. Allow the gas to bubble through 5 mL of water in a test tube for at least 5 minutes. Test the solution with blue

litmus paper. 7. What is the effect of the solution on blue litmus paper?

8. Write an equation for the reaction of CO2 and water. 9. Explain the slight effect on blue litmus paper

D. Fill 1/4 of a test tube with a saturated solution of Ca(OH)2 ("limewater") and allow CO2 to bubble through it

until a precipitate forms. Then, remove the delivery tube. The precipitate is CaCO3. To make that substance, you need Ca+2 ions (which are in the "limewater"), and CO3

-2 ions. 10. Write an equation for the formation of the precipitate CaCO3. 11. Write an equation to show where the CO3

-2 ions come from. 12. Part D will not work with a CaCl2 in place of Ca(OH)2 solution. Why not? 13. Put the delivery tube back into the solution containing the CaCO3 precipitate and allow the gas to bubble

through for a while. Describe what happens and explain the result. Thirteen questions at seven points each plus a nine point curve = ________ (100 max). Deductions for lateness to class, report handed in late, safety and participation = ________ Total Grade = ________

Name________________________________________ SC_____-_____ Date___________ PREPARATION AND PROPERTIES OF HYDROGEN

Equipment: 3 bottles, thistle tube, hydrogen generator, splints, glass plates, pneumatic trough, ring stand, universal clamp, GOGGLES MUST BE WORN Materials: Zinc (mossy, metal), dilute HCl (3N) Aim: How can we prepare hydrogen gas? What are some important characteristics of hydrogen gas? Generating Hydrogen A. Set up the hydrogen generator and arrange the apparatus to collect 3 bottles

of hydrogen by water displacement. Put enough Zn into a generator to fill the bottom fifth (1/5); add enough dilute HCl (3N) to cover the bottom of the thistle tube. Collect one bottle of gas immediately. Cover with glass plate and store it mouth down. Collect two more bottles of gas.

Pour (decant) the acid from the generator into the sink. Rinse twice with water. Keep the Zinc in the generator bottle for the next class. NO ZINC IN THE SINK!!

DO NOT LIGHT FLAMES UNTIL ALL GENERATORS ARE STOPPED AND RINSED!

Testing Hydrogen B. WHEN IT IS SAFE TO DO SO, light a Bunsen burner. Use the burner to light a splint. Hold the SECOND

bottle of hydrogen you collected mouth downward and put the splint halfway into the bottle and hold it there for 30 seconds. DO NOT JERK YOUR HAND AWAY.

1. Describe what happens to the splint inside of bottle and, separately, at the mouth. 2. Does hydrogen burn? What is the evidence? Write an equation for any reaction that occurs. 1. Does hydrogen support burning (can you burn things in hydrogen)? What is the evidence?

C. Hold the FIRST bottle of hydrogen you collected mouth downward and insert a burning splint. 4. Compare this reaction with the one in Part B.

5. Explain any differences. D. Rinse an empty bottle (#1 or #2 from above). Stand it mouth upward. Place the THIRD bottle of hydrogen you collected mouth down on top. Remove the glass plate. Let stand for two minutes. Bring each bottle, separately, mouth downward, to a Bunsen flame. 6. Describe what happens to bottle #3 and to the other bottle as well. 7. If hydrogen gas is less dense than air, how could hydrogen gas get from the top bottle into the bottom bottle? Summary Questions 8. What property of hydrogen enables collection by water displacement? 9. Why do mixtures of hydrogen and air burn faster than pure hydrogen? 10. Complete and balance the equation: Zn + HCl 11. For the reaction: 2H2 + O2 2H2O, what is the sign of ∆H? According to this sign for ∆H, how does the chemical potential energy in water (product) compare to the chemical potential energy in a mixture of hydrogen gas and oxygen gas (reactants)?

12. Draw and completely label a potential energy diagram for burning hydrogen in oxygen. PE in KJ/mole

Reaction Coordinate Total from above at eight points per question = ________ (96 max) + four points curve = ________ (100 max). Deductions for lateness to class, report handed in late, safety and participation = ________ Total Grade = ________

Name________________________________________ SC_____-_____ Date___________ EXPERIMENT 23: RELATIVE REACTIVITY OF METAL IONS

Equipment: 3 test tubes, rack, gloves Materials: 15 mL Cu+2

(aq), 15 mL Pb+2(aq), 15 mL Ag+

(aq), 3 strips each of Cu and Zn, sand paper Aim: Why was reference Table J set up in a particular order?

CAUTION! BE SURE TO POLISH THE METALS IN THE SINK WITH WET SAND PAPER AND RUNNING WATER! USE GLOVES!

A. Pour about 1 inch (25 mm), 5 mL of Cu+2 ion solution into a clean test tube (label it #1), 5 mL of

Pb+2 ion solution into another (label it #2) and 5 mL Ag+(aq) ion solution to a third (label it #3). Add

a strip of polished Zn to each test tube. Do not allow the solutions to completely cover the zinc strip. After 4 minutes, examine each Zn strip.

1. In which solutions do you find deposits on the Zn strip? 2. Identify the solid deposited and record your observation in the table on the next page. If there was

no deposit write none. 3. For the reaction occurring in test tube #1, write an equation showing how copper ion changes to

copper metal. (Reduction half reaction) 4. Where did the electrons used in the reduction of copper ion come from? Write an equation to

illustrate your answer. (Oxidation half reaction) 5. Combine the equations you wrote above to obtain a net redox equation for the reaction which took

place in test tube #1. 6. What is the spectator ion? 7. Why is it not included in the net redox equation? Use this information to fill in the table on the next page.

B. Apply 1 through 6 above to the reactions in test tubes 2 and 3. Enter your equations and observations in the Table of Results.

C. Repeat procedure A substituting strips of Cuo for Zn and test tubes containing solutions of Zn+2, Pb+2

and Ag+. Enter your equations (if reactions took place) in the Table of Results.

SAFETY NOTE: BE SURE YOUR HANDS ARE THOROUGHLY WASHED BEFORE LEAVING THE LABORATORY!

Sol’n Metal

Added

Metal Deposited

Reduction Half-Reaction

Oxidation Half-Reaction

Net Redox Reaction

Cu2+ Zn

Pb2+ Zn

Ag+ Zn

Zn2+ Cu

Pb2+ Cu

Ag+ Cu

EVALUATING THE RESULTS 1. Which metal ions can take electrons from Zno ? 2. Which metal ions can take electrons from Cuo ? 3. Based on your answers to questions 1, 2, and 3 above, rank the metal ions Pb+2, Cu+2, and Zn+2 in decreasing order of ability to take electrons (oxidizing agents). 1. How does your ranking compare with their relative order on Reference Table J? Summary Questions 1. When Zn is placed in a solution of HCl, H2(g) is produced. Use Table J

to identify three metals that will not react with HCl to produce hydrogen gas. 2. The best taker of electrons (oxidizing agent) on Table J is ___________. Where on Table J is this oxidizing agent located? 3. The best donor of electrons (reducing agent) on Table J is ___________. Where on Table J is this reducing agent located? Twenty four spaces in the table are answered = ________ (48 max). First seven questions answered at three points each = ________ (21 max). Evaluation questions at four points each = ________ (16 max). Summary questions at five points each = ________ (15 max). Deductions for lateness to class, report handed in late, safety and participation = ________. Total Grade = ________.

Name___________________________________________ SC_____-_____ Date___________ Laboratory Exercise on Significant Figures

Aim: How can we perform calculations involving significant figures? In this exercise you will learn the rules for determining the correct number of significant figures to use when you perform a calculation. It is expected that you will apply these rules to the calculations that you make in lab for the entire year. On a calculator divide 9653.258 by 5874. What is the answer that your calculator displays? ___________________________________. How many decimal places should be in your answer? A. Rules For Determining Significant Figures. 1. All non-zero digits are significant. Examples: 1, 2, 3, 4 … 2. All zeros between non-zero digits are significant. Examples, 304, 2007, etc… In these cases the

zeros are “holding” the tens and hundreds places. 3. Initial zeros are not significant. For example, in 0.0203 the underlined zeros can be replaced with

scientific notation and the number can be written as 2.03x10-2. This example has only three significant figures. They are the 2, the 3 and the 0 located between the 2 and 3.

How many significant figures are there in the following? a) 11254cm _____ b) 1003mL _____ c) 0068L _____ d) 0.002m ______ 4. Final zeros are not significant if there is no decimal present. For example, in 2,500g only the two and

the five are significant. The final two zeros can be replaced and the number can be written as 2.5x103g.

5. Final zeros are significant if there is a decimal present. For example 200kg. and 2,500kg. have three

and four significant figures, respectively. 6. Final zeros after a decimal are significant. For example, the numbers 2.50L and 3.7770L have three

and five significant figures respectively.

Complete the problems on the next page.

For each number below, indicate the number of significant figures. Number Number of

Significant Figures Number Number of

Significant Figures 2057m

10.77g

60720cm

40770.g

0.209L

0.3090kg

0.00407kg

0.0029cm

27.0036m

8.00g

2.5mL

37,000,000km

B. Calculations Involving Significant Figures

1. When adding and subtracting numbers, the answer has the same number of decimal places as the number in the calculation with the fewest number of decimal places. Round off your answer.

Example, 34.34cm 2.6cm has the fewest number of decimal places (one decimal place). + 1.341cm 38.281cm becomes 38.3cm (has one decimal place and was rounded off).

2. When multiplying and dividing numbers, the answer has the same number of significant figures as the number in the calculation with the fewest number of significant figures. Once again, round off.

Example, 6.24m has three significant figures x 4.3m has two significant figures 26.832m becomes 27m! In this case the answer must have two significant figures. Rounding off, 26.832m becomes 27m. Solve using significant figures: a) 531.46mL - 86.3mL

b) 24.24cm x 43.9cm

c) 5.1 / 21.3

d) 1167.2 + 37.28 + 17.019

e) Go back to the very first problem under the “Objective” of this exercise. What is the correct answer to this problem?

Each question correct from page one times four points each = ______ (64 max). Each question correct from page two times seven points each = ______ (35 max). One point curve plus the above equals _________________ (100 max). Deduction for lateness = ________. Total grade = ________.

BTHS, Mr. Randy Asher, Principal Regents Chemistry 1 Mr. T. Evangelist, A.P., Science Heating Curve Project Due Date: __________________________________ Directions: In this project you will be looking at another substance other than (water. This substance is Benzene (C6H6). The information given below in the chart gives you the various properties of benzene. You must use this information to answer questions 1-10. Write the questions and give answers in complete sentences. When answering questions mathematically write the "Given”, “Formula” and show all work. This project must be typed and provide a good illustration for the graph.

BENZENE (C6H6)

Melting Point

Boiling point

Hf Hv Sample size

Sp, Heat (solid)

Sp. Heat (liquid)

Sp. Heat (gas)

Rate of heating

5.5°C 80°C 127J/g 551J/g 100g 1.5J/g°C 1.7J/g°C 1.0J/g°C 400J/min.

Questions: 1. How much heat is required to melt the sample at the melting point? (10 pts.) 2. How much heat is required to boil the sample at the boiling point? (10 pts)

3. How much heat is required to raise the temperature of the sample from 20°C below melting point to

the melting point? (10 pts) 4. What is the difference in temperature between melting point and boiling point? (5 pts) 5. How much heat is required to raise the temperature of the sample of the liquid phase from the end of melting to the beginning of boiling? (10 pts)

6. How much heat is required to raise the temperature of the sample of the gas phase from the end of boiling to 20ºC above the boiling point? (10 pts) The following questions are based on the above: 7.(20 pts) a. How many minutes are needed to reach the melting point? b. To heat liquid phase to the boiling point? c. How many minutes are needed to change phase at the melting point? d. Boiling point? e. How many minutes are needed to reach the maximum temperature from the boiling point? 8. What is the total time, in minutes, from beginning of melting to the end of boiling? (5 pts)

9. What phase is the substance at 50ºC? (5 pts) 10. Graph data in a precise heating curve. Label all coordinates. (15 pts) Deduction for lateness = ________. Total grade = _________ (100 max).

Name________________________________________ SC_____-_____ Date___________ ICE CREAM MAKING: APPLYING COLLIGATIVE PROPERTIES

Equipment (per group of 2): 1 small zip-lock bag, 1 large zip-lock bag or tennis ball container, 1 spoon Materials (per group of 2): 1/2 cup of milk, 1/4 teaspoon of vanilla or other flavoring, 1 tablespoon sugar, ice, salt (NaCl).

Aim: How can we take advantage of colligative properties to make ice cream? Safety: It is best to do this experiment in a non-laboratory setting. Never eat or drink in the laboratory! Be sure to take a digital image or provide proof that the experiment was completed. Attach this picture or evidence to the lab report when handing it in. Procedure: (Record your data and observations as you work.) • Put 1/2 cup of milk, vanilla and sugar into the small zip-lock bag. • Squeeze out the air as you seal the bag. • Put ice into a tennis ball container (or large bag) to about 1/3 (250 mL). • Roll the small bag of milk into a cylinder and place it into the tennis ball container (or large bag). • Fill the container with ice (about 250 mL more). • Add 6 tablespoons (about 50 mL) of NaCl (Kosher or table salt). • Cap/seal the container. • Gently agitate the mixture by shifting the ice from side to side or rolling the container on the table

until the milk mixture becomes hard. Observations and Analysis: 1. Describe what happens to the ice and salt mixture. 2. Describe how the temperature of the ice mixture and the milk mixture change. Conclusions: 3. Explain why the ice cream forms (solidifies) in terms of the endothermic and exothermic processes

that occur. 4. Which way is the heat flowing?

5. What do you observe to defend your answers to 11 and 12? Student provided evidence of completion of project = ________ (40 max). Questions one through five at twelve points each = ________ (60 max). Deductions for lateness = ________. Total Grade = ________.

BTHS, Mr. R. Asher, Principal Name:_________________________ Mr. T. Evangelist, A. P., Science Solutions Project

Rock Candy Recipe

When you make rock candy, you can see the shape of sugar crystals on a giant scale. The key is giving them lots of time (about 7 days) to grow. As the water evaporates, sugar crystals form on the string or stick, and the shapes that they form reflect the shape of individual sugar crystals.

What Do I Need? • 2 cups sugar • 1 cups water • a small saucepan • a wooden spoon • a small, clean glass jar or any glass vessel • a measuring cup • cotton string • a weight to hang on the string (such as a screw or galvanized washer) • aluminum foil or saran wrap • a pencil or popsicle stick (to suspend the string in the jar) What Do I Do?

Did You Know? Rock candy is one of the oldest and purest forms of candy. It was originally used by pharmacists to make medicines for many kinds of illnesses. 1. Sterilize the glass jar and paper clip by cleaning both and rinsing with hot water for a few minutes. 2. Tie a short piece of kitchen string to the middle of a pencil or Popsicle stick. 3. Attach the paper clip to the end of string. 4. Moisten string lightly and roll in sugar. Place pencil across top of jar, letting the paper clip and string hang into jar (see photo). 5. Heat the water in the saucepan over medium-high heat until it comes to a boil. 6. Completely dissolve the sugar in the boiling water, stirring continuously with the wooden spoon until the solution grows clear and it reaches a rolling boil – do not burn!

CAUTION When making candy, the syrup gets very hot. Get the help of an adult if possible! 7. Remove the solution from the heat, and then carefully pour it into the jar. Cover the jar making sure the clip and string are about 2/3 submerged. 8. Let sit at room temperature, undisturbed, for several days. You can check each day to see how much your crystals have grown. If you can, take pictures of the crystals as they change from day to day. It’s tempting, but don’t touch the jar until the experiment is finished—it usually takes about seven days. 9. At the end of the week, the crystals on your string should be clearly defined, with sharp right angles and smooth faces of various sizes. In the field of crystallography, these are called monoclinic crystals. Their shape is determined by the way the individual sugar molecules fit together, which is similar to the way the shape of a pile of oranges is determined by the shape of the individual oranges and the way they stack together. Provide your Teacher with evidence of completion of project by submitting a sample of the candy with your report (40 points). The following is to be typed or printed NEATLY on separate paper for submission to your teacher. Discussion: (30 points): Describe problems you may have had during the procedure. Did you get your expected results? Do you need to revise your procedure at all based on your results? Be sure to discuss the effect of rate of cooling on crystal formation.

Conclusion Questions (30 points): a. Why does the string need to be soaked?

b. What makes the crystals grow?

c. What is chemical formula of table sugar (sucrose)?

d. What is the chemical structure of sucrose?

e. Draw the structural formula of sucrose.

f. How does the amount of sugar that dissolves change as you change the temperature of the mixture?

g. Why do you need to heat the sugar solution to high temperatures? What type of solution are you trying to

create as the solution cools?

h. What is a crystal and how does it form?

i. Is there another method to form crystals in the solution besides adding seed crystals?

j. Discuss colligative properties in terms of your experiment.

Deduction for lateness = ________. Total grade = ________. Bring in your candy to share with your friends!

Name________________________________________ SC_____-_____ Date___________ NEUTRALIZATION AND TITRATION

Equipment: 2 burets (marked acid and base), 2 small funnels (marked A and B), one 100mL or 150 mL beaker, glass stirring rod, iron ring, ring stand, buret holder, wire gauze, Bunsen burner Materials: 20 mL 0.20M NaOH, phenolphthalein, 20mL HCl of unknown concentration, matches Aim: How can a titration help us to determine the molarity of an acid? A. 1. Fill a buret (labeled BASE; with a pinch clamp) with approximately 20 mL of a solution of NaOH of

known molarity (approx. 0.20 M). 2. Fill another buret (labeled ACID; with a glass stopcock) with 20 mL of a solution of HCl of unknown

molarity. Be sure the tips of the burets are filled. 3. Record the initial volumes in the table below. The meniscus (curved surface of the liquid) should always

be read at eye level, and the readings should always be made at the bottom of the meniscus. B. 4. Run 10 mL of the base into a 150 mL beaker.

5. Add 1 drop of phenolphthalein indicator, and stir. 6. Add the acid drop-wise, and stir after each drop, until the solution remains clear/colorless. 7. Read the meniscus on each buret and record the final volumes in the table below.

C. 8. Repeat the entire procedure in A and B at least once in order to test the precision of the method. Save the

final solution from step B6 for part "D". D. 9. Place the beaker containing a few mL of the neutralized solution on a wire gauze over a burner.

10. Boil until all but a few drops of liquid have evaporated. Then shut off the burner, and allow the remainder to evaporate to dryness. Caution: Overheating may crack the beaker.

Questions: What do you observe? What properties are different in the products that allowed them to be separated in this manner?

Write an equation for the reaction between the acid and the base.

Trial 1 Trial 2 Trial 3 Trial 4 Base: final reading

Base: initial reading

mL of base used

Acid: final reading

Acid: initial reading

mL of acid used

3. Calculate the molarity of the acid for each trial (to the correct number of significant figures). (M acid)(V acid) = (M base)(V base) Trial #1 Trial #2 Trial #3 Trial #4 Average Molarity of the Acid = __________ M 4. Obtain the actual Molarity of the acid from your teacher and calculate your percent error (Reference Table T).

5. Why is it necessary to repeat the experiment four times (at least)? 6. List some possible sources of experimental error. 7. Complete and Balance: _____ NaOH + _____ H2SO4 ____ ______ + ____ ______ _____ KOH + _____ HNO3 ____ ______ + ____ ______ 8. How many milliliters of 0.50M HNO3 would be needed to neutralize 15.0 mL of 1.00M KOH? Conclusions In a neutralization reaction an __________ and a __________ react to form a __________ and

__________. Phenolphthalein, litmus, methyl orange and methyl red are often used in titrations as acid /

base __________.

Three questions from Part D are answered = ________ (10 max.) Questions four through eight at six points each = ________ (30 max). Four trials are completed with correct number of decimal places and correct significant figures = ________ (40 max). Average molarity is calculated using correct number of significant figures = ________ (10 max). Conclusions are answered = _______ (10 max). Deductions for lateness to class, report handed in late, safety and participation = ________. Total Grade = ________.

Name________________________________________ SC_____-_____ Date___________ WEIGHING WITHOUT A BALANCE

Equipment: 8" test tube, 1000mL beaker, rubber stopper Materials: a strip of Mg, HCl, paper strip or rubber band, water at room temperature Aim: How can we find the mass of a strip of Mg without weighing it? A. Place a strip of Mg in an 8" test tube. Fill the test tube with water (at room temp.) to within 1" of its

mouth. Place 400 mL of water (also at room temp.) in a 1000 mL beaker. Add enough HCl solution to the test tube to fill it. Cover the test tube with the LARGE end of a rubber stopper and, while holding the stopper in place with your finger, quickly invert the test tube into the beaker of water. With the mouth of the test tube well below the surface of the water in the beaker, remove the stopper carefully so that the Mg does not fall out. Let the mouth of the test tube rest on the bottom of the beaker. Observe the reaction. [Rinse your hand with tap water.]

1. Describe the reaction: 2. Write a balanced equation for the reaction. B. When all the Mg has reacted, add (room temperature) water to almost fill the beaker. Raise the test

tube so that the liquid level in the beaker is even with the liquid level in the test tube. Put a strip of wet paper on the test tube to mark the level. Remove and drain the test tube. Add water to the level marked by the paper, and then empty the contents of the test tube into a graduated cylinder. Record the volume.

3. VOLUME: mL. 4. What is the reason for equalizing the water levels? C. Obtain the following information: 5. ROOM TEMPERATURE: C = K 6. VAPOR PRESSURE OF WATER AT ABOVE TEMPERATURE: kPa. 7. TODAY'S ATMOSPHERIC PRESSURE: in Hg = mmHg = kPa. D. The hydrogen gas collected in the test tube was actually a mixture of water vapor and H2 gas. 8. Considering your answers to #4 and #7, what is the pressure inside the test tube after the water levels

have been equalized?

P (inside) = P (due to H2 (g)) + P (water vapor) = kPa.

9. What is the partial pressure of hydrogen only? kPa. D. We now know the pressure and the volume of the hydrogen gas at room temperature. We need to

know what the volume WOULD be at S.T.P. Use the Combined Gas Law (P1V1/T1 = P2V2 /T2) to determine this.

10. VOLUME OF HYDROGEN GAS AT STP: _____ mL = __________Liters F. Since one mole of an ideal gas takes up 22.414 L of space at S.T.P., how many moles of hydrogen

gas are represented by the volume you calculated in #10? 11. MOLES OF HYDROGEN GAS: ___________________ G. Use the balanced equation (question #1) and your knowledge of mole ratios to determine how many

mole of Mg must have reacted to produce that many moles of hydrogen (#11). Then use the gram atomic weight of Mg to find the weight of that many moles of Mg. Compare that with the known weight (provided by your teacher) and calculate your percent error.

12. MOLES OF Mg USED: __________ moles 13. WEIGHT OF Mg USED: __________grams 14. ACTUAL WEIGHT: __________grams (from teacher) 15. Observed value – Accepted value Percent error = ------------------------------------------------- x100% = Accepted value Percent error = __________ Refer to Reference Table T and experiment #13.

Questions: 1. Given the reaction: C4H10 + O2 CO2 + H2O a. Balance the equation: b. Based upon part “a”, how many moles of carbon dioxide are produced when 6.5 moles of C4H10 are burned? c. Based upon “b” how many grams of carbon dioxide would be produced? Questions one, two and four at five points each plus one point = ________ (15 max). Calculations for questions 3, 5-13 and 15 at five points each = ________ (55 max). Summary questions were answered at ten points each = ________ (30 max). Deductions for lateness to class, report handed in late, safety and participation = ________. Total Grade = ________.

name sc - date acids and acid · pdf fileconsumed in the reaction transfer the solution to a...

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