Experiment 2: Measuring Density

Name _____________ Lab Partner(s) ________________________ Date __________

All matter has mass, volume, and thus also has density: mass/volume. In this experiment, you will measure the density of matter in various solid and liquid states and shapes.

For a reference on density, see your text: Zumdahl, Introductory Chemistry: A Foundation, Chapter 2.8.

Prelab Question 1: How many cm³are in 1 mL?

Prelab Question 2: How many cm³ are in 1 cc?

Procedure A: Measuring the Density of Regular Solids

From our collection of metal cubes and cylinders, choose one cube and one cylinder.

Observation A1: Find the dimensions of the cube, to the nearest 0.1 cm:

Width (W): Length (L): Height (H):

Question A1: The volume of a rectangular solid = WLH. Show how to calculate the volume of the cube, in cm³.

Observation A2: What is the mass of the cube, in g?

Question A3: What is the density of the cube in g/cm³? Show your calculation.

Observation A3: Find the dimensions of the cylinder, to the nearest 0.1 cm:

Diameter (d): Height (h) :

Question A4: The volume of a cylinder = π (0.5 d)² h. Show how to calculate the volume of the cylinder, in cm³.

Observation A4: What is the mass of the cylinder, in g?

Question A5: What is the density of the cylinder in g/cm³? Show your calculation.

Procedure B: Measuring the Density of Water

Place a dry 25 mL graduated cylinder on a balance and tare it. From a wash bottle, add reagent water to a volume between 20 and 25 mL. If you spill any on the balance, wipe it off before recording the mass.

Observation B1: Record the mass of the water to the nearest 0.01 g:

Observation B2: Record the volume of the water to the nearest 0.2 mL:

Question B1: From Observations B1 and B2, calculate the density of reagent water:

Observation B3: Measure the temperature of the water and record, in ºC:

Question B2: From Table 1, p8, for the temperature nearest recorded in Observation B3, find the density of water, in g/mL:

Question B3: Look at the table. As the water gets colder, how does the density change?

Procedure C: Measuring the Density of Methanol

CAUTION! Methanol is flammable, and poisonous if ingested. Dispose of down the drain, flushing with tap water.

Empty and shake as dry as possible the 25 mL graduated cylinder. Again, place it on a balance and tare it. From a dropper bottle, add methanol (methyl alcohol) to a volume between 20 and 25 mL. If you spill any on the balance, wipe it off before recording the mass.

Observation C1: Record the mass of the methanol to the nearest 0.01 g:

Observation C2: Record the volume of the methanol to the nearest 0.2 mL:

Question C1: From Observations C1 and C2, calculate the density of methanol:

Procedure D: Measuring the Density of Sodium Chloride (NaCl) Solutions

Here you will prepare three aqueous or water solutions of sodium chloride, table salt grade, of three different concentrations, 5 %, 15 %, and 25 % (w/v). ” (w/v)” means “weight/volume”, the weight of the solute NaCl divided by the total volume of the solution you end up with. You will weigh out the required amount of NaCl directly into a tared 25 mL graduated cylinder, then add the water and dissolve the salt. You’ll then record the total mass of the solution and calculate the density.

Here’s how to make a 5 % (w/v) NaCl solution. Bring the 25 mL graduated cylinder and the NaCl to the balance. Put the cylinder on the balance, tare it, then weigh 1.25 g NaCl into the cylinder. Why 1.25 g NaCl? Look at this:

5 % (w/v) means 5 g NaCl solute . But you only want 25 mL of solution, so
100 mL solution

5 g NaCl solute = 1.25 g NaCl = 5 % (w/v)
100 mL solution 25 mL solution

After you weigh out the NaCl, from a wash bottle add reagent water, put Parafilm over the end, and invert the cylinder until the NaCl is dissolved. Then, add reagent water to the 25.0 mL mark and invert again. Remove the Parafilm, place the cylinder back on the balance and record the total weight of the solution.

Observations

Complete the following table. % Error = Measured Density – Handbook Density x 100. This kind of error can be + or –. Handbook Density

Look up the “Handbook Density” in Table 2, p8.

NaCl concentration, % (w/v)	Mass of NaCl solute (g)	Mass of solution (g)	Volume of solution (mL)	Density of solution (g/mL)	Handbook density, g/mL	% Error
0.00 % (pure water, from Procedure B)
5.0 %
15 %
25 %

Now, plot this data on page 7.

Question E1: On the graph paper above, plot your measured density vs. concentration of NaCl; density on the y-axis (vertical axis) and concentration, % NaCl (w/v), on the x-axis (horizontal axis). Hint: start the y-axis at the lowest density you measured, not zero. Label the two axes. Plot the four points from the table. Then, use a transparent ruler to draw one straight line. Draw it as close as you can to all four points at once.

Question E2: Using your plot, find the expected density of a 20 % (w/v) solution of NaCl. On the plot, show how you got this.

Temperature of Water, °C	Density of water, g/mL
0	0.99987
3.98	1.00000
5	0.99999
10	0.99973
15	0.99913
18	0.99862
20	0.99823
25	0.99707
30	0.99567
35	0.99406
38	0.99299

Concentration of NaCl, % (w/w)	Density, g/mL at 20°C
5.0	1.03
6.0	1.04
7.0	1.05
8.0	1.06
9.0	1.06
10.0	1.07
11.0	1.08
12.0	1.09
13.0	1.09
14.0	1.10
15.0	1.11
16.0	1.12
17.0	1.12
18.0	1.13
19.0	1.14
20.0	1.15
21.0	1.16
22.0	1.16
23.0	1.17
24.0	1.18
25.0	1.19

Table 1. Density of water as a function of temperature

Table 2. Density of sodium chloride solutions as a function of concentration

Date taken from Handbook of Chemistry and Physics, 52^nd ed., CRC Press (1971).

Experiment 3: Measuring Boiling Points

Name _____________ Lab Partner(s) ________________________ Date___________

In this experiment you will first calibrate an alcohol thermometer by measuring a known temperature, the boiling point of water, 100°C. Then, you will use your calibrated thermometer to measure an unknown temperature, the boiling point of an alcohol.

For a reference on temperature, see your text: Zumdahl, Introductory Chemistry: A Foundation, Chapter 2.7.

Prelab Question 1: A typical lab temperature is 70°F. What is this temperature in ºC? Show that the units cancel:

Prelab Question 2: When water is heated to higher and higher temperatures, it eventually boils, turning into a gas or vapor. Is this a physical or a chemical change? Why?

Procedure A: Calibrating a Thermometer

The lab thermometers contain a red-dyed alcohol in a glass ‘capillary’ or narrow tube. They are inexpensive and not particularly accurate. When you calibrate your thermometer, you will find the difference between what it reads and what it is supposed to read. You will record the difference between what it reads and what the actual temperature is, and then calculate and apply a correction factor to all later temperature readings obtained using this thermometer. This is called ‘calibration’.

Since you will later use the thermometer to measure boiling points between 60°C and 100°C, you will calibrate the thermometer using the boiling point of water. Water boils at 100ºC when the atmospheric pressure (P_atm) pushing down on the water is 760 mm Hg (torr) or 1 atm. The actual boiling temperature of any liquid, including water, depends on the atmospheric pressure pushing down on the liquid today, in this room. Convert to mm (millimeter) Hg, or torr using this unit-factor or conversion factor calculation:

Atmospheric pressure, in Hg x 25.4 mm = Atmospheric Pressure, mm Hg = Atmospheric Pressure, torr
1.00 in

Observation A1: From the barometer, record today’s atmospheric pressure using the units in (inches) Hg:

Question A1: Show how to calculate today’s atmospheric pressure in mm Hg, using the conversion factor shown above:

Now, using today’s atmospheric pressure (P_atm), look up today’s boiling point of water (T_bp) in Table 3, p12. The boiling temperatures of water range from 97.7ºC to 101.4 ºC. If P_atm does not appear in the table, estimate T_bp by “interpolating” or guessing in between the nearest P_atm values. Ask the instructor to help you read the table if you’re confused.

Question A2: From the table, record today’s boiling point of water, to 0.1 ºC:

Prelab Question 3: As the atmospheric pressure rises, what happens to the boiling point of water?

Observation A2: Obtain an alcohol thermometer. Using it, read and record the room temperature (ºC): Then, have your instructor check the reading and initial here:

Attach a ring to a ring stand and place a wire gauze on the ring. There may be a model apparatus for you to look at in the lab. The ring should be high enough so that you can place a bunsen burner underneath it and still have about 5 cm of clearance for the flame. Place a 250 mL beaker about ¾-full of reagent water on the wire gauze. Place the thermometer in a clamp and suspend it so that its bulb is immersed in the water. The thermometer bulb should not touch the sides or bottom of the beaker.

Light the bunsen burner with the flint lighter or a match. Open the air holes at the bottom by rotating the collar so that you get a blue, hot flame. Heat the water until it is boiling steadily.

CAUTION! Keep an eye on the flame all the time! Remove packs, any extra books, binders and papers from your lab bench and store them away to reduce the chance of catching them on fire! Tie long hair back so that it does not fall into the flame! The burner base may get hot.

Observation A3: If the temperature is above 105ºC or keeps rising, make sure that the thermometer bulb is not touching the sides or bottom of the beaker. Wait until it stops rising, then record the thermometer reading of the boiling water (ºC).

Question A3: Now, calculate the difference between the table or true boiling point of water you recorded in Question A2, and the reading you recorded in Observation A3 by subtracting the reading from the true temperature:

Thermometer Reading – True (table) Temperature =

Question A4: Does the thermometer read too high or too low?

The ‘correction factor’ is the negative of the difference in Question A3. For example if the thermometer reads too high by 5ºC, the correction factor will be -5ºC; you’ll need to subtract 5ºC from every reading. But, if the thermometer reads too low by 5ºC, the correction factor will be +5ºC; you’ll need to add 5ºC from every reading.

Question A5: What is the correction factor for your thermometer (+ or - ºC)?

Procedure B: Measuring the Boiling Point of an Alcohol

In this experiment, you will measure the boiling point of an alcohol. The alcohol will be placed in a small test tube, and the test tube placed in a heated oil bath. The oil will be in a glass tube called a ‘Thiele tube’. By heating the bottom of the Thiele tube with the flame of a bunsen burner, you can slowly and carefully increase the temperature of the alcohol in the test tube.

CAUTION! The alcohols are flammable! If you heat the alcohol too rapidly, or above 100ºC, it will boil over into the flame and catch on fire! If this happens, stand back and allow the alcohol to burn away. Tell the instructor!

Obtain a Thiele tube; they are already set up in clamps on ring stands, and are already filled with mineral oil. Leave them like this at the end of the lab. There may be a complete demonstration assembly with thermometer and glass tube in place in the lab. Refer to this as you set up your own apparatus.

Obtain a small test tube (10 mL) and attach it to the bottom of the thermometer with two of the special rubber bands. Double loop the rubber bands to make sure they are tight. The bottom of the thermometer bulb and the bottom of the small test tube should be about even.

Fill the small test tube about half full with one of the two alcohols: 2-propanol (isopropanol) or ethanol.

Observation B1: Record the name of the alcohol you obtained:

Obtain a glass capillary tube. The tubes are closed at one end, and open at the other. Snap the capillary in half.

CAUTION! When breaking the glass, protect your hands with a cloth or paper towel.

Discard the half that is now open on both ends. Make sure you keep the one that is closed (sealed) on one end, or the experiment will fail.

With the broken, open end down, drop the half capillary into the alcohol in the test tube. If the open end is up, the experiment will fail.

Place the thermometer and test tube assembly into the thermometer clamp and then lower the test tube into the oil of the Thiele tube until the alcohol level is about even with the oil level. Do not immerse the test tube any more than this or, when the oil expands, it will flow into the test tube and the experiment will fail.

The boiling points of both alcohols are above 65ºC but less than 100ºC.

CAUTION! Do NOT heat the oil above 100ºC !

Adjust the gas tap so that you have a small flame, about 2 cm (1 inch) in height. Place the flame at the bottom ‘elbow’ of the Thiele tube and heat the oil up to about 65ºC to start with. After you reach 65ºC, remove the flame and make sure the temperature is not rising rapidly.

Now, heat the oil slowly, by about +1ºC every 10 s. Keep an eye on the thermometer! You will see air bubbles slowly leaving the capillary tube and coming out of its open end, which should be down. If you don’t see this, you may have the open end up; stop and correct this if so.

CAUTION! Do NOT heat the oil above 100ºC ! Heat the oil slowly, with a small flame!

As you approach the boiling point of the liquid, you will suddenly see a rapid, steady stream of bubbles emerge from the broken, open end of the capillary. This is not air, it is the vapor of the liquid. You are now above the boiling point. Remove the flame and allow the oil and liquid to cool.

CAUTION! If you have not seen the stream of bubbles by 100ºC, stop and remove the flame; there is something wrong with your experimental setup; consult with your instructor. Do NOT heat the oil above 100ºC !

As the liquid cools back through its boiling point, the steady stream of alcohol vapor bubbles will suddenly stop. It is very dramatic.

Observation B2: (Trial 1) Record the uncorrected temperature at which the stream of bubbles stops (ºC):

Cool the oil to about 10ºC below this boiling temperature. Drop another broken half capillary into the alcohol in the test tube, again with the broken open end down. Repeat and obtain another measurement of the boiling point:

Observation B3: (Trial 2) Record the uncorrected temperature at which the stream of bubbles stops (ºC) :

Compare the boiling points measured in the two trials. If they differ by more than 2ºC, do a third trial, again using a newly broken capillary.

Observation B4: (Trial 3) Record the uncorrected temperature at which the stream of bubbles stops (ºC) :

Allow the oil to cool to about 40ºC. Remove the thermometer from the Thiele tube, wipe the oil off with a paper towel, and remove the test tube and rubber bands. Discard the alcohol into the sink, with the tapwater running. Put the broken capillaries into the special broken glass waste container. Remove and save the rubber bands. Wipe off the outside of the test tube and return it, the rubber bands, and the thermometer to storage. Return the Thiele tube assembly still assembled.

Question B1: Calculate the average boiling point from all the above trials:

Question B2: To this average boiling point, apply the thermometer correction factor you measured and recorded in Question A5. What is the corrected average boiling point of your alcohol?

Table 3: Boiling point of water (T_bp) as a function of atmospheric pressure (P_atm)

P_atm, torr (mm Hg)	T_bp,°C
700	97.7
705	97.9
710	98.1
715	98.3
720	98.5
725	98.7
730	98.9
735	99.1
740	99.3
745	99.4
750	99.6
755	99.8
760	100.0
765	100.2
770	100.4
775	100.5
780	100.7
785	100.9
790	101.1
795	101.3
800	101.4

From Handbook of Chemistry and Physics, 52^nd ed., CRC Press (1971).

Experiment 4: Chemical Reactions I

Name _____________ Lab Partner(s) ________________________ Date __________

In this lab experiment you will perform some of the classical chemical reactions beloved by students throughout the generations. Pay attention to the names and formulas in this lab. These reactions are a good start at learning formulas, names, and to practice the new skill of writing chemical reactions.

Procedure A: Reactions of the Alkali Metals (Group 1) and Alkaline Earth Metals (Group 2) with Water

Elements in the first two columns of the periodic table, Group 1 and Group 2 are all metals. Group 1 metals are called the ‘alkali metals’ and Group 2 metals are called the ‘alkaline earth metals’.

Most of these Group 1 and 2 metals are reactive and do not occur commonly in their metallic or elemental state. Many of these metals react with water to form the corresponding cation or positively-charged version of the neutral metal. The metals are much more reactive than the cations; in fact, Group 1 and 2 metal elements are usually found in their cation state. For example, sodium metal is highly reactive, but the sodium cation is very inert and is found in some very unreactive, common solids such as sodium chloride (NaCl) and sodium carbonate (Na₂CO₃). There is probably nothing more exciting than sodium metal and nothing more boring than the sodium cation.

Here is a ‘prototype’ unbalanced chemical equation for the reaction of sodium metal with water:

REACTANTS PRODUCTS

Na°(s) + H₂O (l) = NaOH (aq) + H₂ (g) (Equation I)

Solid sodium metal + liquid water = aqueous sodium hydroxide + hydrogen gas

The left side of Equation I shows a list of ‘reactants’ or what chemical substances you started with, and the right side shows a list of products, the substances you ended up with. The ‘=’ sign separates the reactants from the products, and each item in the reactants and products lists is separated by a ‘+’ sign. In Equation I, the sodium metal reactant is emphasized by the symbol Na°; the little ° says that this element is in the metallic or elemental, uncharged form. The product NaOH (aq), sodium hydroxide, is in the ‘aqueous’ or ‘dissolved in water’ form, and consists of the sodium cation, Na⁺, and the hydroxide anion, OH^-. You’ll learn more about cations and anions later.

Prelab Question A1: List the formulas of the products for the reaction of sodium metal with water:

Prelab Question A2: List the formulas of the reactants for the reaction of sodium metal with water:

Equation I is ‘unbalanced’ because the number of atoms of some elements on the left side is different than the number on the right side. Here is the balanced version of Equation I:

2Na°(s) + 2H₂O (l) = 2NaOH (aq) + H₂ (g)

Prelab Question A3: In the balanced equation, how many atoms of hydrogen are there on the reactants side?

Prelab Question A4: In the balanced equation, how many atoms of hydrogen are there on the products side?

Prelab Question A5: In the balanced equation, how many atoms of oxygen are there on the reactants side?

Prelab Question A6: In the balanced equation, how many atoms of oxygen are there on the products side?

Prelab Question A7: What is the difference between an unbalanced and a balanced chemical equation?

Procedure A1: Your instructor will demonstrate the reactions of lithium, sodium, and potassium metals with water. Then, you will obtain samples of magnesium and calcium metals and perform the same reaction at your bench. Pay attention to which metal reacts most vigorously and which one reacts least vigorously. Then, either your instructor or you will measure the pH of the NaOH (aq) solution that results. A pH > 7 means the solution is ‘basic’, a pH < 7 means that the solution is ‘acidic’. Pure water has a pH = 7 and is ‘neutral’.

Observation A1: Describe the reaction of lithium metal with water:

Observation A2: What was the pH of the resulting solution? Is it acidic, basic, or neutral?

Question A1: First see Equation I, p13, for sodium metal. Then write an unbalanced chemical equation for the reaction of lithium with water:

Question A2: Is LiOH an acid, base or neutral salt?

Observation A3: Describe the reaction of sodium metal with water:

Observation A4: What was the pH of the resulting solution? Is it acidic, basic, or neutral?

Question A3: First see Equation I, p13, for sodium metal. Then, write an unbalanced chemical equation for the reaction of sodium with water:

Question A4: Is NaOH an acid, base or neutral salt?

Observation A5: Describe the reaction of potassium metal with water:

Observation A6: What was the pH of the resulting solution? Is it acidic, basic, or neutral?

Question A5: First see Equation I, p13, for sodium metal. Then, write an unbalanced chemical equation for the reaction of potassium with water:

Question A6: Is KOH an acid, base or neutral salt?

Observation A7: Describe the reaction of magnesium metal with water:

Observation A8: What was the pH of the resulting solution? Is it acidic, basic, or neutral?

Question A7: First see Equation I, p13, for sodium metal. Then, write an unbalanced chemical equation for the reaction of magnesium with water. Note that magnesium, as a Group 2 metal, forms Mg(OH)₂, not MgOH.

Observation A9: Describe the reaction of calcium metal with water:

Observation A10: What was the pH of the resulting solution? Is it acidic, basic, or neutral?

Question A9: First see Equation I, p13, for sodium metal. Then, write an unbalanced chemical equation for the reaction of calcium with water. Note that calcium, as a Group 2 metal, forms Ca(OH)₂, not CaOH.

Question A10: Is Ca(OH)₂ an acid, base or neutral salt?

Question A11: Using your reactions with water as the measurement, which alkali metal (Group 1) is most reactive ? Which is least reactive? As you go down Group 1’s column in the periodic table, how does the reactivity of these elements change?

Question A12: Using your reactions with water as the measurement, which alkaline earth metal (Group 2) is most reactive ? Which is least reactive? As you go down Group 2’s column in the periodic table, how does the reactivity of these elements change?

Question A13: Which has the more reactive metals, Group 1 or Group 2?

Procedure B: Two Reactions of Magnesium Metal

You saw above that magnesium metal, a Group 2, alkaline earth metal, is not as reactive as some other Group 1 or Group 2 metals. It needs more than water to get it reacting. Here, you’ll do two other reactions with magnesium metal as the reactant. Like the reactions above, in these reactions as well, hydrogen gas, H₂ (g), is a product, and the metal forms an aqueous salt, dissolved in the water.

Prelab Question B1: Find and give names and symbols for two other Group 2 alkaline earth metal elements besides calcium and magnesium:

Procedure B1. Reaction of magnesium metal with hydrochloric acid. Obtain two or three strips of magnesium metal. Fill the bottom of a watch glass with dilute (6 M) hydrochloric acid, HCl (aq). Drop one of the strips into the dilute hydrochloric acid. Repeat if you like.

Observation B1: Describe the reaction of magnesium metal with hydrochloric acid:

Question B1: Write an unbalanced chemical equation for this reaction:

Question B2: What is the name of the magnesium product?

Procedure B2. Air combustion of magnesium metal. Air combustion was a very confusing reaction for the early chemists. In the mid-18^th century, chemists finally figured out that combustion was actually a reaction with the air gas oxygen, O₂ (g), and not something that came out of the metal. When metals combust in air, they form ‘oxides’. Light a bunsen burner to get about a 5 cm flame. With your tongs, hold a strip of magnesium metal in the flame until it ignites or starts to combust. Repeat if you like.

Observation B2: Describe the air combustion of magnesium metal. Describe the product:

Question B3: Write an unbalanced chemical equation for this reaction. Guess the formula of the product.

Question B4: What is the name of the magnesium product?

Procedure C: Flame Tests: Atomic Emission of Group 1 and Group 2 Metals

When excited with heat or other forms of energy, every element has a characteristic ‘atomic emission’ of light. In this procedure, you will excite chloride salts of several Group 1 and 2 metals with a methane flame and observe the visible atomic emission. The characteristic color shows the presence of this element in any solid mixture or solution.

Prelab Question C1: Write the names of the six metal chlorides whose formulas are given in Procedure C1:

Procedure C1. Flame atomic emission. On a piece of a blank paper draw six circles about the size of a half dollar. Label the circles ‘LiCl’, ‘NaCl’, ‘KCl’, ‘CaCl₂’, ‘BaCl₂’, and ‘SrCl₂’. Obtain samples about the size of a dime and place each sample within the corresponding labeled circle. Light a bunsen burner and get a good, roaring flame. Take a nichrome wire inoculating loop and clean it by placing the wire part into the flame until nothing but the blue methane flame is seen. Then, for each of the six salts, dip the wire first into deionized (reagent) water, then into the solid salt, then place the salt into the hottest part of the flame, just above the lighter blue inner cone of the flame. Observe the atomic emission and its color.

Observation C1: Describe the color and intensity of the atomic emission for each of the Group 1 and 2 metals:

Metal	Formula of metal chloride	Observation of atomic emission
lithium
sodium
potassium
calcium
barium
strontium

Procedure D: Two Sublimations

A sublimation is when a solid substance turns directly into a gas, without going first through the liquid state. Here, you will sublime two pure substances, carbon dioxide (CO₂) and molecular iodine (I₂)

Prelab Question D1: Is a sublimation a physical or a chemical change?

Prelab Question D2: In a sublimation, would you expect to see any new compounds form as products?

Procedure D1. Sublimation of solid carbon dioxide. Obtain a chunk of dry ice. Dry ice is very cold (-78°C) and will ‘burn’ your skin. Handle it with tongs. Drop the piece into some water.

Observation D1: Describe the sublimation of solid carbon dioxide:

Question D1: Write an equation for the sublimation of solid carbon dioxide.

Procedure D2. Sublimation of solid iodine. All halogens (Group 7 elements) occur as diatomic molecules, molecules with two of the same atom. So, solid iodine is I₂. Shake about five crystals of solid iodine into the bottom of a dry 250 mL beaker. Cover the beaker with a watch glass, convex side down. Heat with a small flame until you see something happen.

Observation D2: Describe the sublimation of solid iodine:

Question D2: Write two equations: a) an equation for the sublimation of solid iodine, and b) an equation for the change that occurs on the underneath side of the watch glass.

Experiment 6: Precipitation Reactions

Name _____________ Lab Partner(s) ________________________ Date __________

In a ‘precipitation’ reaction, the reactants are metal cations and non-metal anions These aqueous ions come from soluble substances dissolved in water. The product, however, is an insoluble solid or insoluble salt, a new substance which is formed from the combination of the aqueous reactant ions. The insoluble product forms a cloudy suspension of solid particles, which may or may not settle out; the product has been ‘precipitated’ from solution, like rain or snow precipitates to the ground.

In this experiment, you will combine a number of solutions of soluble compounds already dissolved in water, two solutions at a time. Every reaction will thus have two soluble reactants with two cations and two anions. But, only one of the cations and one of the anions will react! The other two just stay in solution and ‘watch’ the reaction; they are called the ‘spectator ions’. You will write observations, identify the insoluble product, and write a ‘net ionic reaction’ which has as reactants only the one cation and one anion that formed the precipitate. After doing all of this, you’ll be an expert on precipitation reactions!

Since there are four possible reactant ions, how can you tell which of the two possible products is the insoluble one? First, remember, if you don’t see a precipitate, you didn’t get a reaction. Next, if you do see a precipitate, check the solubility rules summarized in Table 7.1, inside back cover of Zumdahl, Fundamentals of Chemistry 6th ed., your textbook. These rules will help you identify the insoluble product. Table 5.1 to 5.4 in the same place will help you to name anions and cations.

Procedure: Do each of the following reactions by combining 3-5 drops of each of the two solutions on a watch glass. Rinse off the watch glass between each reaction with deionized (reagent) water, not tapwater.

Reaction A: Combine silver nitrate and sodium chloride solutions: (Answers are given for Questions A1, A2, A4)

Question A1: Write the formulas for the four ions, two cations and the two anions, in the two solutions.

Ag⁺ NO₃^- Na⁺ Cl^-

Question A2: What are the formulas for the two reactants?

AgNO₃ and NaCl

Observation A1: What did you see when you mixed the two solutions?

(You answer this one!)

Question A3: If you saw an insoluble product form, what is its name and formula?

(You answer this one!)

Question A4: If you saw a reaction, write a balanced net ionic chemical equation showing only the cation and anion that reacted to form the insoluble product. Make sure you use the (aq) and (s) symbols to show what’s in solution (aq) and what’s not (s). Ag⁺(aq) + Cl^-(aq) = AgCl(s)

Reaction B: Combine silver nitrate and sodium carbonate solutions:

Question B1: Write the formulas for the four ions, two cations and the two anions, in the two solutions.

Question B2: What are the formulas for the two reactants?

Observation B1: What did you see when you mixed the two solutions?

Question B3: If you saw an insoluble product form, what is its name and formula?

Question B4: If you saw a reaction, write a balanced net ionic chemical equation showing only the cation and anion that reacted to form the insoluble product. Make sure you use the (aq) and (s) symbols to show what’s in solution (aq) and what’s not (s).

Reaction C: Combine silver nitrate and potassium chromate solutions:

Question C1: Write the formulas for the four ions, two cations and the two anions, in the two solutions.

Question C2: What are the formulas for the two reactants?

Observation C1: What did you see when you mixed the two solutions?

Question C3: If you saw an insoluble product form, what is its name and formula?

Question C4: If you saw a reaction, write a balanced net ionic chemical equation showing only the cation and anion that reacted to form the insoluble product. Make sure you use the (aq) and (s) symbols to show what’s in solution (aq) and what’s not (s).

Reaction D: Combine silver nitrate and barium chloride solutions:

Question D1: Write the formulas for the four ions, two cations and the two anions, in the two solutions.

Question D2: What are the formulas for the two reactants?

Observation D1: What did you see when you mixed the two solutions?

Question D3: If you saw an insoluble product form, what is its name and formula?

Question D4: If you saw a reaction, write a balanced net ionic chemical equation showing only the cation and anion that reacted to form the insoluble product. Make sure you use the (aq) and (s) symbols to show what’s in solution (aq) and what’s not (s).

Reaction E: Combine potassium chromate and lead(II) acetate solutions:

Question E1: Write the formulas for the four ions, two cations and the two anions, in the two solutions.

Question E2: What are the formulas for the two reactants?

Observation E1: What did you see when you mixed the two solutions?

Question E3: If you saw an insoluble product form, what is its name and formula?

Question E4: If you saw a reaction, write a balanced net ionic chemical equation showing only the cation and anion that reacted to form the insoluble product. Make sure you use the (aq) and (s) symbols to show what’s in solution (aq) and what’s not (s).

Reaction F: Combine potassium chromate and barium chloride solutions:

Question F1: Write the formulas for the four ions, two cations and the two anions, in the two solutions.

Question F2: What are the formulas for the two reactants?

Observation F1: What did you see when you mixed the two solutions?

Question F3: If you saw an insoluble product form, what is its name and formula?

Question F4: If you saw a reaction, write a balanced net ionic chemical equation showing only the cation and anion that reacted to form the insoluble product. Make sure you use the (aq) and (s) symbols to show what’s in solution (aq) and what’s not (s).

Reaction G: Combine lead(II) acetate and sodium carbonate solutions:

Question G1: Write the formulas for the four ions, two cations and the two anions, in the two solutions.

Question G2: What are the formulas for these two reactants?

Observation G1: What did you see when you mixed the two solutions?

Question G3: If you saw an insoluble product form, what is its name and formula?

Question G4: If you saw a reaction, write a balanced net ionic chemical equation showing only the cation and anion that reacted to form the insoluble product. Make sure you use the (aq) and (s) symbols to show what’s in solution (aq) and what’s not (s).

Reaction H: Combine lead(II) acetate and sodium sulfate solutions:

Question H1: Write the formulas for the four ions, two cations and the two anions, in the two solutions.

Question H2: What are the formulas for the two reactants?

Observation H1: What did you see when you mixed the two solutions?

Question H3: If you saw an insoluble product form, what is its name and formula?

Question H4: If you saw a reaction, write a balanced net ionic chemical equation showing only the cation and anion that reacted to form the insoluble product. Make sure you use the (aq) and (s) symbols to show what’s in solution (aq) and what’s not (s).

Reaction J: Combine barium chloride and sodium carbonate solutions:

Question J1: Write the formulas for the four ions, two cations and the two anions, in the two solutions.

Question J2: What are the formulas for the two reactants?

Observation J1: What did you see when you mixed the two solutions?

Question J3: If you saw an insoluble product form, what is its name and formula?

Question J4: If you saw a reaction, write a balanced net ionic chemical equation showing only the cation and anion that reacted to form the insoluble product. Make sure you use the (aq) and (s) symbols to show what’s in solution (aq) and what’s not (s).

Reaction K: Combine barium chloride and sodium sulfate solutions:

Question K1: Write the formulas for the four ions, two cations and the two anions, in the two solutions.

Question K2: What are the formulas for the two reactants?

Observation K1: What did you see when you mixed the two solutions?

Question K3: If you saw an insoluble product form, what is its name and formula?

Question K4: If you saw a reaction, write a balanced net ionic chemical equation showing only the cation and anion that reacted to form the insoluble product. Make sure you use the (aq) and (s) symbols to show what’s in solution (aq) and what’s not (s).

Reaction L: Combine sodium chloride and sodium carbonate solutions:

Question L1: Write the formulas for the four ions, two cations and the two anions, in the two solutions.

Question L2: What are the formulas for the two reactants?

Observation L1: What did you see when you mixed the two solutions?

Question L3: If you saw an insoluble product form, what is its name and formula?

Question L4: If you saw a reaction, write a balanced net ionic chemical equation showing only the cation and anion that reacted to form the insoluble product. Make sure you use the (aq) and (s) symbols to show what’s in solution (aq) and what’s not (s).

Experiment 7: Oxidation-Reduction (Redox) Reactions

Name _____________ Lab Partner(s) ________________________ Date __________

In an “oxidation-reduction” or “redox” reaction, there are usually two reactants. One reactant gets oxidized, and the other reactant gets reduced. The products are the reduced and oxidized form of the two reactants, respectively. This happens by transfer of electrons. In other words:

Oxidized Reactant One + Reduced Reactant Two = Reduced Product One + Oxidized Product Two

[Oxidized Reactant One and Reduced Product One] are called a “redox pair”. And,
[Reduced Reactant Two and Oxidized Product Two] form another redox pair.

The redox pairs can be put into two ‘half-reactions’. When a reduced compound or ion is oxidized, one of its atoms has an increase in oxidation state or number. It loses or gives up electrons. When an oxidized compound or ion is reduced, one of its atoms has an decrease in “oxidation state” or “oxidation number”. It gains or gets electrons. Gaining electrons is gaining something negative; that’s why it’s called a ‘reduction’ because its oxidation state is reduced to a lower number:

Reduced Reactant = Oxidized Product + electrons

[This is an Oxidation, Loss of electrons, Oxidation number goes up]

Oxidized Reactant + electrons = Reduced Product

[This is a Reduction, Gain of electrons, Oxidation number goes down]

You can remember which is which using this advice from Leo the Lion: “LEO says GER” (‘Grrrrrr!’, as in ‘Grrrrrrl’!)
LEO = Loss of Electrons is Oxidation, GER = Gain of Electrons is Reduction

There are always two redox pairs and two half-reactions in every redox reaction. In this experiment you’ll do four redox reactions, identify the two redox pairs in each reaction, learn rules for and figure out the oxidation state or number for the products and reactants, and identify the oxidizing and reducing agents.

For reference, you might want to consult Ch. 18 of Zumdahl, Fundamentals of Chemistry 6th ed., your textbook.

Procedure A: Oxidation of Magnesium Metal

Obtain a 4 cm-long strip of magnesium metal and, using tongs, place it in your burner flame. Carefully observe the product that is formed by this combustion or air oxidation reaction.

Observation A1: What did the product of this reaction look like?

Question A1: Write a balanced chemical equation for this air oxidation of magnesium:

Question A2: The oxidation number or state of an element by itself, or combined only with itself is zero (0). The oxidation number or state of a one-element cation or anion is just the charge on the ion.

Above the balanced chemical equation in Question A1, write the oxidation state directly above each of the Mg and O atoms in both the reactants and the product.

Question A3: Which of the two reactants gets oxidized?

Question A4: Which of the two reactants gets reduced?

Question A5: When a compound or ion is oxidized or reduced, the number of electrons gained or lost is equal to the change in its oxidation state. Did the magnesium atom gain or lose electrons? How many electrons?

Question A6: Did the oxygen atoms gain or lose electrons? How many electrons for each oxygen atom?

Question A7: Which reactant was the “oxidizing agent”, the thing that oxidized the other reactant?

Question A8: Which reactant was the “reducing agent”, the thing that reduced the other reactant?

Procedure B: Oxidation-Reduction of Copper and Zinc

Place a thin piece of metallic zinc in a small test tube. Add sufficient 1.0 M copper(II) sulfate solution to immerse the zinc metal. Note the appearance of the zinc and the solution. Let stand for 20 min. Then again note the appearance of the solution and the zinc.

Observation B1: What changes did you see?

Question B1: In this reaction, the Zn metal dissolves to form aqueous zinc sulfate. At the same time, the reactant copper(II) sulfate plates out on the surface of the Zn as copper metal. Write a balanced equation for this reaction:

Question B2: The oxidation number or state of an element by itself is zero (0). The oxidation number or state of a one-element cation or anion is just the charge on the ion.

Above the balanced chemical equation in Question B1, write the oxidation state directly above each of the metal atoms in both the reactants and the products.

Question B3 Which of the two reactants gets oxidized?

Question B4 Which of the two reactants gets reduced?

Question B5 When a compound or ion is oxidized or reduced, the number of electrons gained or lost is equal to the change in its oxidation state. Did the zinc atom gain or lose electrons? How many electrons?

Question B6 Did the copper atom gain or lose electrons? How many electrons?

Question B7 Which reactant was the “oxidizing agent”, the thing that oxidized the other reactant?

Question B8 Which reactant was the “reducing agent”, the thing that reduced the other reactant?

Procedure C: Oxidation-Reduction of Ammonium Dichromate

Make a small pile of ammonium dichromate, (NH₄)₂Cr₂O₇ in the bottom of your white porcelain evaporating dish. It should be heaped up in a pile, with a base diameter of about 3 cm. Place the dish in a sink. Light a kitchen match and place the lit match into the center of the ammonium dichromate pile.

Observation C1: What did the reaction look like? It should be rather dramatic.

Dispose of the product by collecting it with a wet sponge, then rinsing down the lab sink drain.

Question C1: In this reaction, the nitrogen atom from the ammonium cation forms nitrogen gas. The hydrogen atom and the oxygen from the dichromate anion (Cr₂O₇^2-) combine to form water vapor. And, the chromium atoms combine with the oxygen atoms to form chromium(III) oxide. None of this requires oxygen gas as a reactant! Write a balanced equation for this reaction:

Question C2: It’s not easy to see what gets oxidized or reduced in this reaction. Let’s find the oxidation numbers or states of the atoms that react. Look at the cation, ammonium, first. The oxidation state of the highly-reduced nitrogen in ammonium
is -3. The oxidation number of an element by itself is zero (0).

Above the balanced chemical equation in Question C1, write the oxidation state directly above each of the nitrogen atoms in both the reactants and the products.

Question C3: Does the nitrogen atom get oxidized or reduced?

Question C4: Now look at the chromium atom in the dichromate anion, Cr₂O₇^2-. In this oxyanion, oxygen has an oxidation state of -2, just like its charge as an oxide. The sum of each atom times its oxidation state must equal the overall charge of the anion. What’s the oxidation state of each of the Cr atoms in the dichromate anion?

Question C5: What is the oxidation state of each of the chromium atoms in the product?

Question C6: Above the balanced chemical equation in Question C1, write the oxidation state directly above each of the chromium atoms in both the reactants and the products.

Question C7: Does the chromium get oxidized or reduced?

Question C8: Which reactant ion was the “oxidizing agent”, the thing that did the oxidizing?

Question C9: Which reactant ion was the “reducing agent”, the thing that reduced the other ion?

Procedure D: Oxidation-Reduction of Halogens

Add 5 mL of chlorine water, Cl₂(aq), to about 1 mL of methylene chloride in a small test tube. This will form two layers because methylene chloride does not mix, or is not ‘miscible’, with water. Shake the tube with a whipping motion. The lower layer is methylene chloride as it’s more dense than water. What color is the lower, methylene chloride layer? Now add about 1 mL of 10 % potassium iodide. Shake again.

Observation D1: Observe the lower, methylene chloride layer against a white background. What does it look like now?

Question D1: In this reaction, the iodide in the potassium iodide reacts with the chlorine water to form the iodine molecule. This iodine is extracted from the water by the methylene chloride layer and has a strong characteristic color in this layer which you may recognize. The methylene chloride is just a solvent and does not react. Write a balanced equation for this reaction:

Question D2: The oxidation number or state of an element by itself is zero (0). The oxidation number or state of a one-element cation or anion is just the charge on the ion.

Above the balanced chemical equation in Question D1, write the oxidation state directly above each of the halogen atoms in both the reactants and the products.

Question D3: Which of the two reactants gets oxidized?

Question D4: Which of the two reactants gets reduced?

Question D5: When a compound or ion is oxidized or reduced, the number of electrons gained or lost is equal to the change in its oxidation state. Did the reactant iodide anion gain or lose electrons? How many electrons?

Question D6: Did the reactant chlorine atom gain or lose electrons? How many electrons for each atom?

Question D7: Which reactant was the “oxidizing agent”, the thing that oxidized the other reactant?

Question D8: Which reactant was the “reducing agent”, the thing that reduced the other reactant?

Question D9: Look at a periodic table. What are the six possible oxidation states of chlorine? Which one is the most ‘reduced’? Which one is the most ‘oxidized’?

Experiment 9: An Assay of Vinegar

Name _____________ Lab Partner(s) ________________________ Date __________

In this experiment, you will ‘assay’ vinegar for its concentration of acetic acid, the main ingredient of vinegar. One does an ‘assay’ when one measures how much of the concentration of the major ingredient or major ‘constituent’ is in a mixture. For example, if you assay gold ore, you are measuring how much gold is in the ore. If you assay Tylenol, you’re measuring how much of the active ingredient, acetaminophen, there is in it.

You will assay the vinegar by titrating it with a known concentration of aqueous sodium hydroxide, NaOH (aq). In this titration reaction, one reactant’s amount is known (NaOH) and the other reactant’s amount is the unknown (acetic acid).

The NaOH solution is called the ‘titrant’ and has a concentration of about 0.1 M. The NaOH titrant is delivered from a buret, a long, skinny graduated cylinder with a valve or ‘stopcock’ at the end of it. The ‘M’ in ‘0.1 M’ stands for ‘molar’, a unit of concentration, and means mol/L. The ‘molarity’ of a solution is its concentration in M or mol/L. So the NaOH solution has a concentration or molarity of 0.1 mol NaOH/L.

Vinegar is a dilute aqueous acid solution, about 5 % (w/v) acetic acid, CH₃COOH. The acetic acid is called the ‘analyte’ because you are measuring the unknown amount of acetic acid in the vinegar sample.

Below is the balanced chemical reaction for the titration reaction, a neutralization of the acetic acid in vingear with the base titrant sodium hydroxide (NaOH). The net ionic reactants, OH^- and H⁺ are shown enlarged. In a neutralization reaction, an acid reacts with a base to form a salt and neutral (pH = 7) water:

NaOH (aq)+ CH₃COOH (aq) = NaCH₃COO (aq) + H₂O (l)

Titrant, + Analyte, = a sodium salt + water
sodium hydroxide acetic acid

Prelab Question 1: Write the net ionic chemical equation for the above neutralization reaction:

You will add an acid-base indicator, a dilute aqueous solution of phenolphthalein (the active ingredient in many laxatives), to the vinegar sample. The vinegar is an acidic solution, and the phenolphthalein indicator is colorless in acidic solutions. At the point where the amount of NaOH titrant you will add uses up all the analyte, acetic acid, the very next drop of titrant will be an excess of NaOH. This excess drop of the strong base NaOH makes the vinegar solution quickly turn from acidic to basic.

As soon as the solution turns basic, the indicator will turn bright pink, its basic color. This color change is called the endpoint. The color change you can see; what you can’t see is exactly when all the acetic acid is neutralized by the sodium hydroxide; this is called the ‘equivalence point’. If you are careful not to add too much excess of titrant, the endpoint will be very close to the equivalence point. But it’s always necessary to go a little beyond it to get an endpoint.

Prelab Question 2: Which reactant is in excess before the endpoint?

Prelab Question 3: Which reactant is in excess after the endpoint?

Procedure A: Preparation of the 0.1 M NaOH Titrant:

CAUTION: Solid sodium hydroxide is very corrosive! Do not touch the pellets! Tranfer them using a spatula or by shaking them. If you drop one on the bench or floor, wipe it up with a wet sponge, then rinse the sponge into the sink.

To a 600 mL beaker, add about 2 g NaOH, to the nearest pellet. Don’t try to divide the pellets.

Observation A1: Record the actual mass of NaOH you obtained (g): ____________________________________

Dissolve the NaOH in about 100 mL of reagent water, then dilute with reagent water to 500 mL. Stir with a glass rod to mix.

Question A1: The sample calculation below was done for a mass of 1.0 g NaOH. Use your mass of NaOH, similarly calculate your actual molarity or actual concentration of your 500 mL of NaOH titrant:

1.0 g NaOH x 1000 mL x 1 mol NaOH = 0.050 mol NaOH = 0.050 M NaOH
500 mL 1 L 40.0 g NaOH 1 L

Procedure B: Buret Practice

Mount a buret clamp on a ring stand and then mount the buret in the buret clamp. Arrange the clamp so that it clamps on to the buret just above the stopcock at the buret’s bottom; most of the buret should be above the clamp. Your instructor may have an example set up in the lab; if so, take a look at it.

Close the stopcock. Place a small beaker below the stopcock. Fill the buret about ¾ full with reagent water from your wash bottle. Open the stopcock and deliver some water with the buret into the small beaker. Watch how fast it comes out. Try regulating the flow rate of this water ‘titrant’ by turning the stopcock. When you are far from the endpoint, you can deliver titrant quickly to save time, but when you get near the endpoint, you’ll want to deliver the titrant more slowly.

Close the stopcock and read the buret using the buret reading card supplied. The black part of the reading card is placed just below the ‘meniscus’, the downward-curved surface of the water, and the buret is read at the bottom of the meniscus. Look straight across the buret at the card, not down on it, or up at it.

Observation B1: Record the buret reading to the nearest 0.l mL: __________________________________

Observation B2: Have your instructor check this reading, and then initial here: _______________________

Procedure C: Titration of Vinegar

1. Empty the water from the buret. Put a small beaker below the buret.

2. Pour about 10 mL of your NaOH titrant into the buret using a funnel. Rinse the buret with this titrant, then discard this titrant.

Prelab Question 4: What would happen if you did not rinse the buret with the NaOH titrant?

3. Fill the buret carefully with your NaOH titrant to above the 0.0 mL mark using a funnel. If it overflows, rinse off the outside and dry before using.

4. ‘Zero’ the buret by opening the stopcock until the titrant in the buret drops to or below the ‘0.0 mL’ mark. It does not have to be at ‘0.0 mL’. Use the buret reading card. Rinse off the tip of the buret with reagent water from a wash bottle.

Observation C1: Record the initial buret reading (mL): ________________________________

5. Pipet 5.00 mL of vinegar from the automatic repipettor into a 125 mL erlenmeyer flask. Your instructor will demonstarte the repipettor. This is the sample of vinegar you will assay.

6. Add 2 drops of phenolphthalein indicator to the flask.

7. Place a piece of white paper under the flask to see the endpoint clearly. Now, titrate the vinegar until the indicator turns pink. Swirl the flask as you add titrant. The indicator will be locally pink as the titrant hits the vinegar solution, but turn colorless as the titrant is mixed in and neutralized by the acetic acid. As you near the endpoint, the pink will swirl away more and more slowly. Slow down the titrant delivery as you near the endpoint. Try not to overshoot the endpoint. When the pink color persists, even after swirling, you’ve passed the endpoint.

Observation C2: Record the final buret reading at the endpoint (mL): ________________________________

Question C1: Calculate the volume of titrant delivered to the endpoint (= final volume – initial volume):

Trial 1 titrant volume (mL): __________________________________________________________________

8. Rinse the erlenmeyer flask out with reagent water, then take another 5.00 mL sample of vinegar and repeat Procedure C. Record the buret readings below:

Observation C3: Record the initial buret reading (mL): ___________________________________________

Observation C4: Record the final buret reading at the endpoint (mL): ________________________________

Question C2: Calculate the volume of titrant delivered to the endpoint (final volume – initial volume (mL):

Trial 2 titrant volume (mL): __________________________________________________________________

9. The Trial 1 and Trial 2 titrant volumes should agree to within about 2 %. If they are off, or if you want to make extra sure you are getting an accurate assay, do a Trial 3:

Observation C5: Record the initial buret reading (mL): ___________________________________________

Observation C6: Record the final buret reading at the endpoint (mL): ________________________________

Question C3: Calculate the volume of titrant delivered to the endpoint (final volume – initial volume (mL):

Trial 3 titrant volume (mL): __________________________________________________________________

10. Calculate the mean (average) titrant volume using the trials of the vinegar assay that agree to 2 %:

Question C4: Mean titrant volume (mL) ________________________________________________________

11. The sample calculation below was done for a mean titrant volume of 20.0 mL and a titrant concentration of 0.100 M NaOH (0.100 mol NaOH/L). Using your own mean titrant volume data and your own titrant concentration, similarly calculate the percent acetic acid in vinegar by weight or ‘% acetic acid (w/v)’:

[Mean titrant volume] [Titrant concentration]

20.0 mL NaOH x 1 L x 0.100 mol NaOH x 1 mol CH₃COOH x 60.0 g acetic acid = 0.12 g acetic acid

1000 mL 1 L 1 mol NaOH 1 mol CH₃COOH

0.12 g acetic acid x (100 %) = 2.4 % (w/v) acetic acid solution.
5.00 mL vinegar

Question C5: Show the same calculation using your own data. Solve for your measured % acetic acid in the vinegar, the assay of the vinegar:

Empty the buret into the sink and rinse it with reagent water. Make sure the stopcock is open and that you rinse reagent water through the tip or else it will clog when it dries. Return the buret after rinsing it.

Question C6: On the vinegar bottle label the manufacturer usually gives its acetic acid assay. Using your measurement, calculate the percent relative difference from the label concentration: % difference = [(measured – label) / label ] x 100 %

Question C7: Which is the ‘right’ concentration, the one on the label or the one you measured? Explain.

Experiment 10: Classes of Organic Compounds

Name _____________ Lab Partner(s) ________________________ Date __________

An organic compound is a compound which has carbon in it. Because carbon can covalently bond with many atoms, including itself, oxygen, and hydrogen, there are over 3,000,000 known organic compounds, containing from one to thousands of carbon atoms. In order to make sense of this huge variety, it’s useful to classify organic compounds into classes or families of compounds. In this lab, you will get an introduction to some classes of organic compounds by observing a few physical and chemical properties of each type. In each class, every organic compound molecule has a characteristic ‘functional group’ which you will identify by drawing the ‘structural formula’ for the molecule.

For reference and help with the Prelab Questions, refer to Chapter 20, ‘Organic Chemistry’, and Ch 21.7 ‘Carbohydrates’, in your textbook (Zumdahl, Introductory Chemistry, A Foundation, 6th ed). Two useful web sites are www.chem.purdue.edu/gchelp/molecules/index.html and chemfinder.cambridgesoft.com.

Properties and Reactions of some Hydrocarbons: Alkanes and Alkenes

A hydrocarbon is an organic compound which contains only carbon and hydrogen atoms. The carbons may be bonded to each other in single (C— C), double (C=C), or triple (C≡C) bonds. The C— H bond is always a single bond. Hydrocarbons with only single bonds are called ‘alkanes’. Hydrocarbons with double bonds are called ‘alkenes’, and those with triple bonds are called ‘alkynes’. Both alkenes and alkynes are ‘unsaturated’ hydrocarbons. Alkenes and alkynes, with their double bonds, can undergo reactions that produce alkanes. Alkanes, with only single bonds, cannot undergo such reactions.

A six-carbon ring , with alternating double and single bonds, is called an ‘aromatic hydrocarbon’. Benzene is the simplest aromatic hydrocarbon and is the molecule on which all the other aromatic hydrocarbons are built.

Prelab Question 1: Draw the structural formula for hexane. Is it an alkane or an alkene?

Prelab Question 2: Draw the structural formula for 1-hexene. Is it an alkane or an alkene?

Prelab Question 3: Draw the structural formula for cyclohexane. Is it an alkane or an alkene?

Prelab Question 4: Draw the structural formula for cyclohexene. Is it an alkane or an alkene?

Prelab Question 5: Draw the structural formula for toluene. What kind of hydrocarbon is toluene?

Procedure A: Air Combustion of Hexane

Observation A1: Your instructor will ignite hexane. What do you observe?

Question A1: Write a balanced equation for the complete air combustion of hexane

Procedure B: Bromination of Double Bonds in Unsaturated Hydrocarbons

Place 1 mL samples of hexane, cyclohexane, cyclohexene, and toluene. Shake the ends of each tube to mix. CAUTION! Bromine water is aqueous Br₂, a corrosive liquid and gas. Hydrocarbons with double bonds will ‘decolorize’ or bleach some or all of the color from the bromine water by reacting with the Br₂ . This reaction forms colorless ‘brominated’ hydrocarbons which are saturated. This is called an ‘addition’ reaction and is a fast reaction.

Observation B1: Compare the solution in each tube to the water blank, which will not react. What did you see? Which hydrocarbons reacted?

Question B1: When an unsaturated compound reacts in a ‘bromine addition’ reaction, the Br₂ molecule ‘adds across the double bond’: One Br atom from Br₂ attaches to a carbon on one side of the double bond, and the other Br atom attaches to the carbon on the other side of the double bond. The double bond becomes a single bond. Remembering that carbon always has four bonds, first write the structural formula for 1-hexene, and then write the structural formula for the bromine addition product 1,2-dibromohexane:

Properties and Reactions of Six Alcohols: Functional Groups

The class or family of organic compounds called ‘alcohols’ all have something in common. They all contain the hydroxyl ‘functional group’ —OH. A functional group is part of a molecule. In general, the formula of an alcohol is R—OH where R stands for the Rest of the molecule. For example, the formula for ethanol, an alcohol, is CH₃CH₂OH. ‘R’ is CH₃CH₂— , the ‘ethyl’ group. Functional groups always have the –yl ending: hydroxyl, ethyl, methyl, etc. They aren’t molecules, but important and recurring pieces of molecules that define the organic compound classes.

Here are the six alcohols you will examine: methanol, ethanol, isopropanol, 1-butanol, 1-pentanol, and 1-octanol.

Prelab Question 6: Draw the structural formula for methanol (methyl alcohol). Circle the hydroxyl functional group.

Prelab Question 7: Draw the structural formula for ethanol (ethyl alcohol). Circle the hydroxyl functional group.

Prelab Question 8: Draw the structural formula for 2-propanol (isopropanol or isopropyl alcohol). Circle the hydroxyl functional group.

Prelab Question 9: Draw the structural formula for 1-butanol (n-butyl alcohol). Circle the hydroxyl functional group.

Prelab Question 10: Draw the structural formula for 1-pentanol or n-pentyl alcohol. Circle the hydroxyl functional group.

Prelab Question 11: Draw the structural formula for 1-octanol or n-octyl alcohol. Circle the hydroxyl functional group.

Procedure C: Water Solubility of Alcohols

Obtain 1 mL samples of the above six alcohols and determine which are soluble or ‘miscible’ in water.

Observation C1: Which alcohols are miscible with water?

Question C1: Looking at the formulas for these alcohols, can you see a trend in which molecules are soluble and which are insoluble?

Question C2: For the insoluble alcohols, is the top or bottom layer the alcohol? Why?

Question C3: What does the water molecule have in common with all alcohols?

Procedure D: Air Combustion of Ethanol

Observation D1: Your instructor will ignite ethanol. What do you observe? How does this combustion differ from the combustion of hexane?

Question D1: Write a balanced equation for the complete air combustion of ethanol

Procedure E: Lucas Test for Alcohols

If you look at the C to which the –OH group is bonded, in a primary (1°) alcohol there is only one other C atom bonded to this C, in a secondary (2°) alcohol there are two other C atoms bonded to this C, and in a tertiary (3°) alcohol there are three other C atoms bonded to this C. The Lucas reagent, conc. HCl and ZnCl₂, will react with 2° and 3° alcohols to form a white, solid precipitate, but will not react with 1° alcohols.

Place 5 drops of ethanol, 2-propanol (isopropanol), and 2-methyl-2-propanol (t-butyl alcohol) into three dry test tubes. Into a fourth tube place 5 drops of a water as a blank. To each tube add 1 mL of the Lucas reagent. Shake the ends of the tubes to mix. Cover the ends of the tubes with Parafilm. Heat the tubes in a 60 °C water bath for 15 min.

Observation E1: Compare the solution of each alcohol to the water blank, which will not react. Which of these alcohols gave a positive test? Which of these alcohols are primary alcohols?

Structures of Five Aldehydes and Ketones

The similar classes or families of organic compounds called ‘aldehydes’ and ‘ketones’ all have something in common.
                                                                                    O
                                                                                    ║
They all contain the carbonyl functional group:   —C—

O
║
In general, the formula of an aldehyde is R—C—H, or R—(CHO), where R stands for the Rest of the molecule.

O
║
The formula for a ketone is R₁ —C— R₂, or R—(CO)—R, where the Rs can be the same or different.

Here are the aldehydes and ketones you will examine: formaldehyde, acetaldehyde, acetone, glucose and fructose. Glucose and fructose are simple 6-carbon sugars, called ‘hexoses’. Glucose is an ‘aldose’, and fructose a ‘ketose’.

Prelab Question 12: Draw the structural formula for formaldehyde. Is it an aldehyde or ketone? Circle the functional group.

Prelab Question 13: Draw the structural formula for acetaldehyde. Is it an aldehyde or ketone? Circle the functional group.

Prelab Question 14: Draw the structural formula for acetone. Is it an aldehyde or ketone? Circle the functional group.

Prelab Question 15: Draw the structural formula for the straight-chain form of glucose. Is it an aldehyde or ketone? Circle the functional group.

Prelab Question 16: Draw the structural formula for the straight-chain form of fructose. Is it an aldehyde or ketone? Circle the functional group.

Properties of Two Organic Acids

The class or family of organic compounds called ‘organic acids’ or ‘carboxylic acids’ have, like all acids, a removable proton (H⁺) to donate. The proton is found on the end of the ‘carboxyl’ functional group:

║
—C— OH or —COOH

When an organic acid donates its proton, it turns into its conjugate base, its ‘carboxylate anion’:

                                                 O                                                   O
                                                    ║                                                   ║
                                             R—C— OH       + H₂O      =      R—C— O ^—    + H₃O ⁺

Many organic acids are soluble in water and have pleasant tastes; acids like citric, ascorbic, and acetic acids are found in many foods. You will examine two organic acids: acetic acid and propionic acid.

Prelab Question 17: Draw the structural formula for acetic acid. Circle the carboxyl group.

Prelab Question 18: Now, draw the formula for its conjugate base, the acetate anion:

Prelab Question 19: Draw the structural formula for propionic acid. Circle the functional group.

Prelab Question 20: Now, draw the formula for its conjugate base, the propionate anion.

Procedure F: pH of Organic Acids

Observation F1: Measure the pH of each of the two organic acid solutions by dipping a glass rod into them and touching the wetted rod to the pH paper. What pHs did you measure for each organic acid?

Observation F2: Carefully note and record the odors of the acids:

Properties of an Organic Base or Amine

The class or family of organic compounds called ‘organic bases’ have, like all bases, the capability to accept a proton (H⁺). All organic bases are related to the base ammonia, NH₃ If you remove one or more H’s from ammonia and replace with an R--, you will get an organic base, or an ‘amine’. An organic base or amine has an ‘amino’ group, —NH₂. Like in ammonia, the amino group has a lone pair of electrons on the N atom that ‘likes’ protons and can accept one in this reaction:

R—NH₂ + H⁺ = R—NH₃⁺ or

R—NH₂ + H₂O = R—NH₃⁺ + OH^-

This means that most organic bases, like organic acids, are soluble in water. You will examine two bases in dilute aqueous solution: dilute ammonia, and a solution of the organic base n-butylamine. Both bases, and most amines have unpleasant, fishy smells. In fact, fish that’s less than fresh is full of ‘histamines’, organic bases which cause allergic symptoms.

Prelab Question 21: Draw the structural formula for ammonia. Circle the amino group.

Prelab Question 22: Now, draw the structural formula for its conjugate acid. What’s its name?

Prelab Question 23: Draw the structural formula for n-butylamine (1-aminobutane) Circle the amino group.

Prelab Question 24: Now, draw the formula for its conjugate acid, the n-butylammonium cation.

Procedure G: pH of Organic Bases.

Observation G1: Measure the pH of the two organic base solutions. What pHs did you measure for the ammonia and for the n-butylamine?

Observation G2: Carefully note and record the odors of the bases:

Experiment 11: Organic Synthesis of Aspirin

Name _____________ Lab Partner(s) ________________________ Date __________

In this lab, you will prepare, isolate and purify a compound in an organic ‘synthesis’ or ‘preparation’. You will synthesize or prepare the ester aspirin from the organic acid salicylic acid. ‘Aspirin’ is the Bayer Company’s former brand name for acetylsalicylic acid, a ‘derivative’ of salicylic acid, and is one of the oldest synthetic pharmaceuticals still in production, by many firms. After further purification and drying, you could put the aspirin you will make in a bottle, label it and use if it met FDA and USP (U.S. Pharmacopeia) standards. Pharmaceutical companies doing similar preparations on a large scale must follow ‘Good Manufacturing Practices’ (GMP) in order to comply with the terms of their FDA applications. Since you won’t have the time or resources to comply with these regulations, you also will not be able to sell or consume the aspirin you make!

In this synthesis, salicylic acid undergoes an ‘esterification’ reaction by reacting with acetic anhydride. Upon reaction with water, acetic anhydride forms acetic acid. You could just use acetic acid as the esterification reagent. However, acetic anhydride reacts more quickly with salicylic acid.

The reaction for the synthesis of aspirin is:

salicylic acid + acetic anhydride = acetylsalicylic acid (‘aspirin’) + acetic acid

C₇H₆O₃ + C₄H₆O₃ = C₉H₈O₄+ C₂H₄O₂

Sulfuric acid is used as a ‘catalyst’, to speed up the reaction, but it is not a reactant. The reaction is shown here using names and the empirical formulas.

Prelab Question 1: Below, write this equation showing the structural formulas for the reactants and products. See Zumdahl, Introductory Chemistry: A Foundation, (your text), Ch. 19.15, and the Merck Index for the formulas. Two useful web sites are www.chem.purdue.edu/gchelp/molecules/index.html and chemfinder.cambridgesoft.com.

Prelab Question 2: Is this equation balanced?

Prelab Question 3: What’s the molar mass or ‘molecular weight’ (MW) of salicylic acid? Show your calculation.

Prelab Question 4: What’s the molar mass or ‘molecular weight’ (MW) of aspirin? Show your calculation.

Prelab Question 5: 1.0 g of salicylic acid reacts completely with an excess of acetic anhydride. What’s the theoretical yield of aspirin in g? Show your calculation.

Lab Preparation: Make sure the water baths are set and warmed up to 70°C and that the fume hoods are on. An ice bath should also be prepared.

Procedure A: Preparation of Aspirin

On the balance, tare a 125-mL erlenmeyer flask; that is, place the flask on the balance and press the ‘Tare’ button so that the balance reads ‘0.00 g’. Then add about 1.0 g of salicylic acid to the flask.

Take the flask to a fume hood. In the fume hood, add about 3 mL of acetic anhydride to the salicylic acid in the flask.

CAUTION! Acetic anhydride causes severe skin burns. The vapors are highly irritating; do not remove it from the fume hood. If you spill any on your skin, flush the skin with tapwater immediately! If you spill any in the fume hood, tell the instructor.

To the above mixture, add 2 or 3 drops of 50 % sulfuric acid.

CAUTION! Sulfuric acid causes severe skin burns. If you spill any on your skin, flush the skin with tapwater immediately! If you spill any in the fume hood, tell the instructor.

Swirl the solution until it is homogeneous. Now place the flask in the water bath set to 70°C and heat for 20 minutes. Make sure the flask does not float in the water! If it does, some water needs to be removed. Swirl the solution about every 5 min. Make sure the water bath is not hotter than 70°C using a thermometer.

After the 20 min of heating, the aspirin product has formed, but is dissolved in the solution. You can carefully note the odor of acetic acid, the other product, by wafting the air above the flask towards your nose. Remove the flask from the heating bath and place it into an ice bath, a mixture of melting ice cubes and water at 0°C. Cooling the solution reduces the solubility of the aspirin in the solution and, eventually, if the synthesis was successful, you will start to see crystals of your product form. Often, gently scratching the surface of the flask with a glass stirring rod helps start the crystallization.

Observation A1: What do you eventually see happen as you cool and scratch the flask?

Once all the product has crystallized out of solution, add 50 mL of reagent water. This reacts with any excess acetic anhydride to form acetic acid.

Set a long-stemmed funnel in a large erlenmeyer flask. Place a piece of fluted filter paper into the funnel. Your instructor will shown you how to fold a piece of filter paper into a fluted shape.

Stir up the crystallized aspirin and pour the suspension through the filter paper. You can use small amounts of reagent water from a wash bottle to help transfer the product into the filter paper. Once all the filtrate has passed through the filter paper, rinse the crude product with about 20 mL of water.

The wet solid left in the filter paper is your wet aspirin product. If you were to purify it and dry it, this would take another lab period, so you will stop at this stage of the synthesis.

Procedure B: Bicarbonate Test for an Acid

To see if you did get some kind of acid as a product, you’ll do some tests on the crude, wet product. Place a small amount of the crude product in a test tube, add to it a few milliliters of sodium bicarbonate solution, and mix. Bicarbonate (hydrogen carbonate) will act as a base if there is acid present, forming carbon dioxide gas.

Observation B1: What do you see when you do this test?

Question B1: The reaction has three acids in it; one reactant and both products are acids. Does this test confirm that your crude product is aspirin?

Procedure C: Iron (III) Test for Salicylic Acid

If the synthesis has been unsuccessful, there will be a lot of unreacted salicylic acid present in the crude product. Also, when pure aspirin decomposes over time, it reverts to salicylic acid and acetic acid. Old bottles of aspirin may smell like acetic acid if this has happened. Salicylic acid is irritating to the stomach so it’s important to see if the aspirin contains significant concentrations of this impurity.

A quick USP ‘spot test’ for the presence of salicylic acid is to add a solution of iron (III) cation to the aspirin. If there is salicylic acid present, a dark purple color will form as the phenolic –OH group of salicylic acid reacts with the iron (III).

Question C1: Why wouldn’t you expect to see the purple color upon testing pure aspirin?

Set up a rack with three test tubes. Into the first, place a small amount of your crude, wet aspirin product. Into the second, place a small amount of crushed, commercial aspirin. Into the third, place a small amount of pure salicylic acid as a control. To each tube, add 5 mL of iron (III) chloride (ferric chloride) solution and shake. Let settle for 1 min, then observe the solution layer.

Observation C1: Results for your crude product:

Observation C2: Results for commercial aspirin:

Observation C3: Results for pure salicylic acid:

Question C2: Interpret each of the three test results. What can you conclude?

Procedure D: Triboluminescence of Methyl Salicylate

Methyl salicylate, like aspirin, is another ester of salicylic acid. Methyl salicylate is found in nature in the wintergreen plant and gives it its distinctive flavor. Most esters, in fact, have pleasant and interesting flavors, and the field of perfumes and fragrances is one of ester chemistry.

After the lab, go out and buy some wintergreen Lifesavers or Altoids; get the ones with sugar in them, not the sugar-free kind. These contain the flavoring methyl salicylate. Take them into a dark bathroom, close the door and wait five minutes to get your eyes used to the dark. Stand in front of the mirror and place a Lifesaver between your back teeth. With your mouth open, crunch down hard. You will see a flash of light emitted from the Lifesaver. This is called ‘triboluminescence’ and is a property of the methyl salicylate, sucrose or glucose, and nitrogen from the air reacting to the sudden pressure of your teeth. All three compounds must be present for the triboluminescence light emission to be seen. Keep trying if you don’t see it the first time. If you do it with another person instead of a mirror, it’s even more fun.

Object No.	Description of object	Gross Weight = object + container (g)	Tare Weight of container (g)	Net Weight or mass of object (g)
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