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Chemistry 113.1 Experiment 1. Density complete solutions correct answers key
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Chemistry 113.1 Experiment 1. Density complete solutions correct answers key

 

 

Introduction to Chemical Techniques

 

INTRODUCTION

Density (ñ) is defined as the ratio of the mass (m) of a sample to its volume (V):

ñ = m / V

Mass and volume are extensive properties of matter-properties that depend on the quantity of a

substance. Such properties are not in themselves useful in characterizing or identifying

substances. Intensive properties such as density however are useful in identifying substances.

Intensive properties are often determined by taking the ratio of two extensive properties

measured under constant temperature and pressure conditions. As an intensive property, density

can be useful in identifying a substance. Density alone cannot absolutely identify a substance but

can be a useful value contributing to an identification. For example, a colorless liquid found to

have a density of 1.00 g/mL at 4 0C and 1.0 atmosphere pressure could be water, since this is the

known density of water. Additional information would be needed to absolutely identify the

substance. In contrast, a colorless liquid found to have a density of 0.85 g/mL at 4 0C and 1.0

atmosphere pressure could not be water.

In the experiments below, you will use several methods to determine the volume of samples,

both solid objects and liquids. You will use the electronic balance to determine the mass of the

samples to 0.001 g (1 milligram). From these measurements, you will determine the densities of

these samples.

A. DENSITY OF REGULARLY SHAPED OBJECTS

For regularly shaped objects, such as cylinders, the volume can be determined by measuring the

dimensions of the object with a ruler, then applying the proper formula to determine the volume.

In this experiment, you will determine the density of a group of objects (all cylinders)

individually, then by a graphical method.

1. Obtain one set of cylinders from your instructor. Record the CODE on the container in

your laboratory notebook. Record the color and any other distinguishing characteristics as

well.

2. Using the plastic ruler in your kit, measure the diameter(d) and the length (or height-h) of

each cylinder to the nearest 0.5 millimeter (e.g. diameter = 13.5 mm = 1.35 cm).

3. Using the electronic balance assigned to you, determine the mass of each cylinder to the

nearest milligram (0.001 g).

4. Using the measured diameter and length, calculate the volume of each cylinder in cm3.

V= ðr2h = ð(d/2)2h

Chemistry 113.1

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Experiment 1. Density (May 2012)

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5. Calculate the density r of each cylinder. Report the calculated value for each cylinder, as

well as the value of the mean (average). [See the addendum regarding the mean and

mean absolute deviation. These should always be reported whenever three or more

determinations of the same quantity are the result of identical experiments.]

6. Graph the data for the cylinders with the mass (g) as the y-axis and the volume (cm3) as

the x-axis. Using the straight line fitting function of the graphing software, find the

formula corresponding to the best fit of the graphed points to a straight line. The slope of

this line is the density (Δm/ΔV). Report this value. Compare it to the average density

reported in (5) above.

B. DENSITY OF IRREGULARLY SHAPED OBJECTS

If an object has and irregular shape, its volume can be determined using Archimedes’ principle

which states: An insoluble body completely submerged in a fluid displaces its own volume.

Thus, the volume of the displaced fluid is equal to the volume of the irregularly shaped object.

1. Obtain a set of mineral samples from your instructor. Record the code identifying the

sample in your laboratory notebook. Note any distinguishing characteristics of the

minerals such as color, shape etc.

2. Using the electronic balance record the mass of the samples to the nearest milligram

(0.001 g).

3. Place approximately 30.0 mL of water in the 100 mL graduated cylinder. Record the

exact volume to the nearest 0.1 mL.

4. Carefully add the sample or samples to the water in the graduated cylinder without

causing any water to be lost by splashing. Note: It may be best to determine the

combined volume of 2 or more pieces of mineral together. As long as the samples are

completely submerged, the greater the increase in volume for the water in the graduated

cylinder, the more precise the measurement of density will be.

5. Remove the samples from the cylinder, dry them with a paper towel, return them to the

storage container and return them to your instructor.

6. Calculate the density (ñ) of the mineral sample.

C. DENSITY OF LIQUIDS: CONSTRUCTING A CALIBRATION CURVE AND DETERMINING V%

COMPOSITION OF AN UNKNOWN SAMPLE

In this experiment, you will prepare a series of liquid mixtures of known composition (percent by

volume or V%) and determine their densities. Using your graph of density versus V% you will

determine the V% of an unknown sample by measuring its density and comparing it to the

graphed values.

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Experiment 1. Density (May 2012)

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1. Using a small beaker, obtain about 10 mL of alcohol (either methanol or ethanol). Record

the name of the alcohol used.

2. Place a clean dry 10 mL graduated cylinder on the electronic balance and tare it to 0.000

g.

3. Carefully transfer 2 mL of the alcohol into the graduated cylinder. Read and record the

exact volume to the nearest 0.1 mL. [ Note: When reading the volume, the level being

read should be at eye level.]

4. Record the mass of the alcohol. Using the mass and volume, calculate the density and

enter it in the data table.

5. Add 1 mL of distilled water to the alcohol in the cylinder. Record the exact total volume

now in the cylinder and total mass. [Note: Total volume should now be about 3.0 mL]

6. Add 2 additional mL of distilled water to the contents of the cylinder (the total volume at

this point should be approximately 5 mL total). Record the exact total volume and the

total mass.

7. Add 2 additional mL of distilled water to the contents of the cylinder (the total volume at

this point should be approximately 7 mL total). Record the exact total volume and the

total mass.

8. EMPTY the graduated cylinder and dry it.

9. At this point you may need to place the cylinder on the balance and tare to 0.000g once

again. Then-add 2 mL of distilled water to the graduated cylinder and record the exact

volume and the mass to 0.001 g. Divide the recorded mass of the water by its volume.

The value should be 1.00 g/mL, which is the known density of water. If this is not the

case consult with your instructor immediately.

Obtain a sample of alcohol/water of unknown (to you) V% composition . Determine the volume

and mass of two individual 3 to 4 mL portions (known as aliquots) of this sample. Record the

exact volume and mass of these two aliquots.

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Experiment 1. Density (May 2012)

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RECORDING DATA

In your laboratory notebook, you should record all data in a format similar to the suggested

formats below:

A. DENSITY OF REGULARLY SHAPED OBJECTS

Sample Code __________________

Description:

Table A. The dimensions and mass of each object.

Object Diameter (cm) Length or Height (cm) Mass (g)

A

B

C

D

Note: In your laboratory report, this data should be transcribed into a neatly typed table. A final

column should be added giving the calculated value of density ñ, in g/cm3.

Using the tabulated values above prepare a graph of mass(g) (y-axis) versus volume (cm3) (xaxis).

A straight line fit has a slope equal to the density of the sample [slope = Δm/ΔV]. Report

this value of the determined density. Using the table of Materials and their Densities provided,

can you identify the material ?

B. DENSITY OF IRREGULARLY SHAPED OBJECTS

Table B. The mass (or combined masses) of the mineral sample(s), the initial volume of water in

the graduated cylinder [Vi (H2O)], the final volume after addition of the sample(s) [Vf (H2O)],

and the volume of the sample(s), Vmineral.

M (g) [Vi (H2O)] (mL) [Vf (H2O)] (mL) Vmineral (mL)

Note: In your laboratory report, this data should be transcribed into a neatly typed table. A final

column should be added giving the calculated value of density ñ, in g/cm3. [Recall that 1 mL=1

cm3]. Using the provided table of mineral densities and descriptions, can you tentatively identify

the mineral you were assigned ?

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Experiment 1. Density (May 2012)

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C. DENSITY OF LIQUIDS: CONSTRUCTING A CALIBRATION CURVE AND DETERMINING VOLUME

PERCENT (V%) COMPOSITION OF AN UNKNOWN SAMPLE

(1) The alcohol used in this experiment was ____________________

(2) The V% for this alcohol is _______________________

[Note: ethyl alcohol (ethanol) is 95% by volume ethanol and 5% water]

Table C1. Volume of alcohol Valcohol , volume of added water Vwater, total volume Vtotal , total

mass m and volume percent alcohol V%alcohol

Sample Valcohol Vwater added Vtotal mass total V%alcohol

1 At start (only

entry)

2

3

4

5 (water only)

The volume percent (V%) is equal to: 100 x (Valcohol / Vtotal ). For example, if 2.0 mL of 95V%

ethanol is initially present, then for a Vtotal of 5.0 mL (after addition of 1 and then 2 mL of

water):

V%ethanol = 100 x(0.95 x 2.0) / (5.0) =38%

Note: In your laboratory report, this data should be transcribed into a neatly typed table. A final

column should be added giving the calculated value of density ñ, in g/mL for each value of

V%alcohol.

In your report, you will graph the values of density (g/mL) (y-axis) versus V% (x-axis) and

provide a straight line fit to the points on the graph.

Table C2. Volume V and mass m of two aliquots of alcohol/water mixture of unknown V%.

Volume Mass Density (r)

aliquot #1:

aliquot #2:

Using the graph of density (g/mL) (y-axis) versus V% (x-axis) which you have prepared,

determine and report the V% of each of the two aliquots.

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Experiment 1. Density (May 2012)

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A. TABLE OF MATERIALS AND THEIR DENSITIES

Material Density(g/cm3)

aluminum 2.71

teflon 2.20

polyvinyl chloride 1.37

phenolic 1.32

polyurethane 1.23

acrylic 1.17

nylon 1.15

polypropylene 0.90

B. TABLE OF MINERAL DENSITIES AND DESCRIPTIONS

Density

(g/mL) description Probable ID

5.0 metallic, brassy, crystalline FeS2 pyrite

4.94-5.07 metallic gray hematite, a-Fe2O3

2.75-2.79 light blue with white and yellow-brown aquamarine

4.01 dark brown with red highlights Alamndine garnet

3.12 black Schorl tourmaline

2.93 blue with white lapis lazuli

2.83 dark brown with paler gold and white streaks tiger eye, microcrystalline SiO2 with

iron oxide

2.71 purple and hite amethyst, SiO2 with Fe impurities

2.69 light brown with paler gold and white streaks citrine

2.59 light gray with black Labradorite

2.58 turquose (light) with white amazonite

2.53 light blue with white and brown chrysoprase

2.36 royal blue with white veining sodalite, Na4Al3(SiO4)3Cl

Chemistry 113.1

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Experiment 1. Density (May 2012)

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LABORATORY REPORT SUGGESTIONS FOR THIS LABORATORY

In addition to the more general instructions posted on Blackboard, here are some specific tips

with reference to the Density experiments:

Abstract

• One sentence defining density

• One sentence stating the results of the regular shaped object density determination

• One sentence stating the results of the irregularly shaped object(s)-minerals density

determination

• One sentence or two giving the results of the density of liquids experiments, including the

volume percent (V%) of the unknown mixture

Introduction

• Explain the difference between intensive and extensive properties and why one is useful

in characterizing materials

• Explain density with the equations, defining all variables or symbols used

Experimental

• Summarize the procedures for regularly shaped objects

• Summarize the procedures for irregularly shaped objects

• Summarize the procedures for density of liquids, including the unknown sample.

Results/Discussion

• Tables of data should be neatly transcribed from the data sheets in your laboratory

notebook, adding where needed values (such as density) calculated from the data.

• Provide a sample calculation of each type of calculated value (but not every one).

• Show the graphs for density of regular shaped objects (A) and volume percent as a

function of density (C). State the slope for each and what it signifies.

• In the appropriate section, provide the tentative ID for the material of the regularly

shaped objects (cylinders) and explain your reasoning. Do the same for the irregularly

shaped object(s) with appropriate reasoning. In each case provide an estimate of the

uncertainty in the reported density values, with explanation.

• Report the value of the V% for the unknown alcohol-water mixture, with uncertainty.

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ADDENDUM

REPORTING THE MEAN AND MEAN ABSOLUTE DEVIATION

Calculating the Mean density and Mean Absolute Deviation

1) Determine the Mean: Add all numbers and divide by the count (3)

Example: the density of three cylinders, denoted by letters are found to be:

A : 1.6 g/cm3

B : 2.0 g/cm3

C : 1.8 g/cm3

Mean = (1.6 + 1.8 + 2.0)/3 = 1.8 g/cm3

2) Determine deviation of each result from the Mean ( individual value – Mean)

1.6 - 1.8 = - 0.2

1.8 - 1.8 = 0.0

2.0 - 1.8 = + 0.2

3) Eliminate the + or – sign and take the mean of the absolute deviations

Thus the Mean Absolute Deviation is (0.2+0.0 +0.2)/3 =0.13 g/cm3 round to 0.1 g/cm3

Report the density as 1.8 +/- 0.1 g/cm3 [Alternatively, you can use the Excel formula

=AVEDEV(1.6, 2.0, 1.8) to obtain the result. ]

Chemistry 113.1

Introduction to Chemical Techniques

Experiment 2. Hydrate Composition (May 2012)

1

I. INTRODUCTION

The law of definite (or multiple) proportions states that when two or more elements combine to

form a given compound, they do so in fixed proportions by mass. For example, sodium chloride

contains 39.3% by mass sodium and 60.7% by mass chlorine. In these experiments, the law of

definite proportions will be used to determine the empirical formulas of hydrated ionic salts. An

empirical formula expresses the simplest whole number ratio of atoms or units in a compound.

(For ionic compounds or hydrates, the unit can be a polyatomic anion or water.)

Hydrates are substances formed when water combines chemically in definite proportions

with an ionic salt, thereby giving a constant ratio of water molecules to the ions of the salt.

Hydrates are not mixtures, since the water is coordinatively (covalently) bound to either the

cation or anion or both in the salt. In CuSO4 • 5 H2O, for example, the bonding involves four

water molecules coordinatively bound to the Cu2+ ion in a square planar structure and one

molecule of water bound to the sulfate ion by hydrogen bonds. The anhydrous (without water)

form of a hydrated salt is produced when all the waters of hydration are lost. Some examples of

hydrates are listed below:

Formula Common name

2 CaSO4 • H2O plaster of Paris

CaSO4 • 2 H2O gypsum

CuSO4 • 5 H2O blue vitriol

MgSO4 • 7 H2O Epsom salt

Na2CO3 • 10 H2O Washing soda

The • in the formula indicates a kind of chemical bond that usually can be easily broken. For

example, magnesium sulfate heptahydrate can be converted to anhydrous magnesium sulfate by

heating:

MgSO4 • 7 H2O (s) → MgSO4 (s) + 7 H2O (g) .

In this chemical reaction equation (or chemical equation), the (s) indicates a solid and the (g)

indicates a gas. In the appendix, more details about this reaction equation will be given along

with how these equations are balanced and how they can be used to predict products of reactions.

In this experiment, you will heat various hydrated salts to determine the number of water

molecules in the salt.

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Experiment 2. Hydrate Composition (May 2012)

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II. PROCEDURE

Obtain from the instructor a hydrated salt chosen from copper sulfate, calcium sulfate, and

magnesium sulfate. The difference in the mass of the anhydride and the hydrate will then be

used to determine the mass of water in the hydrate and, therefore, the empirical formula of the

hydrate. The procedure, which should be performed on two samples of the same hydrate, is as

follows:

1. Weigh a clean, dry, labeled crucible. Record the weight in your notebook.

2. Introduce about 1 - 2 grams of the pulverized hydrated salt. Note the appearance and

color of the solid.

3. Weigh the crucible and contents. Record this weight in your laboratory notebook.

4. Setup a wire triangle on the iron ring over a Bunsen burner. (Ensuring that the wire

triangle will hold the crucible in an upright position.)

5. Watch the instructor demonstrate how to setup and properly light a Bunsen burner and

how to turn-off the burner after use. (Be sure to record this in your notebook for later

referral.)

6. Heat the crucible and contents in the hottest part of the flame for 5 - 10 minutes. (The

bottom of the crucible should turn a dull red during heating.) Initially, the hydrate should

be heated slowly by waving the burner flame fairly rapidly under the crucible. If the

material begins to boil or crackle, the heating is too intense and splattering may occur.

Within approximately 1 minute, the material should become drier and stronger heat can

be applied. At the end of the 5 - 10 minute period of heating, allow the crucible to cool

slightly before transfer.

7. Using clean crucible tongs, transfer the crucible to a desiccator and allow the crucible to

cool to room temperature.

8. When cool, weigh the dish and the anhydride and record this weight in your notebook.

9. Heat the crucible in the flame again for 5 minutes, place in desiccator and allow the

crucible to cool. Once cool, weigh the sample again. Continue the heat/cool/weigh cycle

until the mass of the sample remains constant. Be sure to record all of your

measurements in your notebook.

10. Place a thermometer in the anhydrous salt and record the temperature.

11. Add a few drops of water to the anhydrous salt near the thermometer and record the

temperature (once the temperature has stopped increasing). This temperature change

represents the change from an anhydrous salt to a hydrated salt.

Remember: Write the experimental procedures that YOU followed while you were doing

the experiment. Be sure to note if the salt splattered or popped out of the crucible while

heating. These types of observations will be important when discussing sources of

experimental error.

Chemistry 113.1

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Experiment 2. Hydrate Composition (May 2012)

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Useful information that should be recorded in the notebook at some point during the experiment:

• Name of salt and formula

• Qualitative description of the salt before and after heating.

• Temperature of the salt before and after the addition of water

Table 4.1. Masses m in grams (g) necessary to determine the composition of the salt from the

first trial.

Object m (g) Notes

Clean, dry crucible

Crucible with salt (before

heating)

Crucible with salt after 1st

heat/cool cycle

Crucible with salt after 2nd

heat/cool cycle

Crucible with salt after 3rd

heat/cool cycle (or until stable)

Crucible with salt after 4th

heat/cool cycle (or until stable)

Crucible with salt after 5th

heat/cool cycle (or until stable)

Chemistry 113.1

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Experiment 2. Hydrate Composition (May 2012)

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Table 4.2. Masses m in grams (g) necessary to determine the composition of the salt from the

second trial.

Object m (g) Notes

Clean, dry crucible

Crucible with salt (before

heating)

Crucible with salt after 1st

heat/cool cycle

Crucible with salt after 2nd

heat/cool cycle

Crucible with salt after 3rd

heat/cool cycle (or until stable)

Crucible with salt after 4th

heat/cool cycle (or until stable)

Crucible with salt after 5th

heat/cool cycle (or until stable)

Chemistry 113.1

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Experiment 2. Hydrate Composition (May 2012)

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III. POST-LABORATORY DISCUSSION AND QUESTIONS

The mass of a single atom is difficult to measure. (For instance, the mass of a single hydrogen

cation (or proton) is 1.67 × 10-24 g.) Therefore, the mole is defined as the number of 12C atoms in

exactly 12 grams of 12C. Moreover, the basic unit of mass for elemental chemistry, namely the

atomic mass unit (amu or dalton) is defined as 1 amu ≡ 1/12 the mass of an atom of 12C = 1.6605

× 10-24 g. Thus,

The constant 6.022 × 1023 atoms (or molecules)/mole is known as Avogadro's number NA. Since

the mole and the atomic mass unit are defined using the same scale, 1 amu × NA = 1 g/mole.

Thus, the masses given on the periodic table can also be expressed as the number of grams of the

element per mole of element. The molar mass M of a compound is obtained by summing the

mass of all of the elements in a compound and, therefore, has units of g/mol. Moreover, the

definition of a mole when combined with the law of definite proportions implies that a sample of

H2O will have 2 moles of atomic hydrogen for every 1 mole of atomic oxygen, while a sample of

MgF2 has a mole ratio of 1 mole of magnesium for 2 moles of atomic fluorine.

Please note that moles are used as the universal conversion factor in chemistry. The

chemical reaction equation written in the introduction, for instance can now be read as follows:

1 mole of solid magnesium sulfate heptahydrate decomposes with heating to generate one

mole of solid magnesium sulfate and 7 moles of gaseous water

Thus, please become more familiar with this difficult concept by reading about the mole in a

standard freshman chemistry book or on Wikipedia. Also, in the future, when in doubt,

convert to moles! The post-laboratory questions below will help guide you in the conversions

between mass in grams and moles of compound and will show you how this allows you to

determine an empirical formula for a hydrate.

Chemistry 113.1

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Experiment 2. Hydrate Composition (May 2012)

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POST-LABORATORY QUESTIONS

Where appropriate, these questions should be answered for each trial. Remember, do not

write Question 1 and then an answer. Also remember to show all work for the calculations

for one of the trials.

1. Determine the mass mh of the hydrated salt by subtracting the mass of the empty crucible

from the mass of the salt and the crucible before heating.

2. Determine the mass ma of the anhydrous salt by subtracting the mass of the empty

crucible from the mass of the salt and the crucible after the heat/cool cycles are complete.

3. The difference in the mass of the hydrated salt and the anhydrous salt is the mass of water

present in the sample. Why is the mass different (i.e., what happened to the water)?

4. Calculate the molar mass of water.

5. Calculate the moles of water in the hydrated salt by dividing the mass of water by the

molar mass of water.

6. Calculate the molar mass of the anhydrated salt. (To do this, use the chemical formula

you wrote from the name of the compound that you used.)

7. Calculate the moles of the anhydride in the sample by dividing the mass of the anhydride

by the molar mass of the anhydride.

8. Determine how many moles of water are associated with a single mole of anhydride by

dividing the moles of water by the moles of anhydride. What is the average value for this

ratio?

9. Using the information, write the formula of the hydrated salt in the form

Anhydride • x H2O , where x is the average value obtained in question 8

10. Is x an integer to the correct precision? If not, why? What sources of error could have

caused x not to be an integer?

11. Lookup your salt on Wikipedia. Is x in Question 9 an appropriate value based on the

possible hydrates that your salt can form? What is the percent error in your value,

assuming that the information on Wikipedia for the hydrate is correct?

Chemistry 113.1

Introduction to Chemical Techniques

Experiment 3. Precipitation reactions (May 2012)

1

I. INTRODUCTION

In Experiment 3, you applied heat from a Bunsen burner to decompose a hydrate into an

anhydrous salt and gaseous water. A decomposition reaction is one of four broader categories of

chemical reactions. The remaining categories are precipitation reactions, acid/base reactions,

and oxidation/reduction reactions. In this experiment, you will investigate precipitation

reactions. In a precipitation reaction, two aqueous solutions of soluble salts are mixed and yield

an aqueous solution of a soluble salt and a solid compound. The formation of the solid is called

precipitation and the solid is called the precipitant.

When ionic compounds dissolve in water, the water interacts with the cation and anion to

weaken the Coulombic interaction holding the two ions together as a solid. Thus, as the ions

break apart because of water surrounding the individual ions in the compound, the solid

dissolves. For example, when CuSO4 dissolves in water, the chemical reaction equation is

CuSO4 (s) → Cu2+ (aq) + SO4

2- (aq) .

Again, the (s) stands for solid, while the (aq) stands for aqueous. Precipitation occurs when

aqueous cations and anions form Coulombic interactions that are strong enough to overcome the

interaction of the water molecules with the separate ions in solution. In other words, the solid is

more stable than an aqueous solution containing the two aqueous ions. Compounds that do not

dissolve in water are called insoluble, while those that do are called soluble. Table 5.1 gives

some basic rules for the solubility of ionic salts. These rules should be memorized, although in

Chemistry 114, you will learn that these rules are not as black and white as they are presented

here.

Table 5.1. Rules for determining the solubility of ionic compounds

1. Compounds of the alkali metal ions (i.e., the column to the far left on the periodic table) are soluble.

2. Compounds containing ammonium ion are soluble.

3. Nitrates, chlorates, perchlorates, and acetates are soluble.

4. Chlorides, bromides, iodides are soluble except when combined with lead(II), silver(I) and mercury(I)

cations. Mercury(II) iodide is also insoluble.

5. All sulfates are soluble except when combined with strontium, barium, calcium, lead(II), mercury(I), and

silver cations. Small amounts of calcium, silver, and mercury(I) sulfates will dissolve in solution (i.e.,

slightly soluble).

6. Carbonates, phosphates, oxalates, and chromates are insoluble unless they fall under the categories 1 and 2

(i.e., alkali metal or ammonium ion salts).

7. Sulfides are insoluble unless they fall under categories 1 and 2. Alkaline earth metals (i.e., calcium,

strontium, and barium) form slightly soluble sulfides.

8. Hydroxides and oxides are insoluble except for those that fall under categories 1 and 2. Alkaline earth

metals form slightly soluble hydroxide and oxide salts.

Chemistry 113.1

Introduction to Chemical Techniques

Experiment 3. Precipitation reactions (May 2012)

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In a precipitation reaction, two solutions containing soluble ionic salts are mixed. However,

some of the ions, when the solutions are combined, can interact to form insoluble salts. When

this occurs, a solid forms (i.e., the precipitant) and produces a cloudy solution and often falls to

the bottom of the container. For example,

Solution A: Aqueous copper sulfate: CuSO4 (aq) or Cu2+ (aq) + SO4

2- (aq)

Solution B: Aqueous sodium carbonate: Na2CO3 (aq) or 2 Na+ (aq) + CO3

2- (aq)

Notice that we can represent aqueous solutions (aq) in two ways, namely one with the ionic

compound formula (aq) and one with the individual ions (aq). Both ways are equally valid.

Also notice that when sodium carbonate goes into solution, we obtain one mole of carbonate

anion for every two moles of sodium cations. Bulk solutions can have no net charge. Thus,

when you sum the charges for all ions in solution, you must obtain zero. The integer number to

the left of the sodium cation insures that the solution remains at zero net (or total) charge. Then,

when the two solutions are mixed, we have the following chemical reaction equation:

Cu2+ (aq) + SO4

2- (aq) + 2 Na+ (aq) + CO3

2- (aq) → CuCO3 (s) + 2 Na+ (aq) + SO4

2- (aq)

The equation written above is called an ionic reaction equation. Notice that the sodium cation

and the sulfate anion appear on both sides of the equation. This is because he combination

Na2SO4 is soluble. These ions are called spectator ions because they do not participate in the

chemical reaction that yields the solid copper(II) carbonate. Thus, the net ionic reaction is

Cu2+ (aq) + CO3

2- (aq) → CuCO3 (s)

Notice how both of these reaction equations have the same number of each atom on the right

hand and left hand side of the equation and that both of these equations also have the same net

charge on both sides of the equation. Insuring that the charge and mass on both sides of a

chemical reaction equation is the same are both part of balancing a chemical reaction equation.

II. EXPERIMENT

1. Given the following aqueous solutions, predict which mixtures of any two of these

solutions would yield precipitants and determine the formula for the precipitant:

a. Copper(II) sulfate b. Barium nitrate

c. Sodium chloride d. Silver nitrate

e. Lead nitrate f. Sodium sulfate

2. Show your predictions to the instructor.

3. Test to see if your predictions where correct by placing a small amount of the appropriate

solution into a clean test tube and adding the second solution. You must test all of the

reactions that you predict would form a precipitant. When the precipitant forms, record

the color of the precipitant. You must also test at least two reactions that, based on Table

5.1, you predict would not form a precipitant.

Chemistry 113.1

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Experiment 3. Precipitation reactions (May 2012)

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III. Post-laboratory discussion and questions

You must always balance chemical reactions before using these reactions to predict the amount

of product one might expect from the reaction (i.e., the theoretical yield of the reaction). In this

experiment, one has probed solubility and, therefore, must now learn how to write a chemical

reaction equation and balance it. When balancing a chemical reaction equation, the mole ratio

of various compounds in the reaction can be adjusted to balance both mass and charge for the

reaction. However, the formula of a compound or a polyatomic ion does not change. Therefore,

do not change atomic symbols or subscripts in chemical formulas while attempting to balance

reactions. Some basic guidelines are

1. Balance atoms other than H and O first

2. Pure Elements [e.g. Fe (s) or Cl2 (g) ] should be balanced last

3. Balance as a unit any polyatomic ions that appear unchanged on both sides of the arrow.

As an example, let us balance the reaction of aqueous copper sulfate with aqueous sodium

hydroxide. The first step is to write the ionic reaction. Thus,

CuSO4 (aq) + NaOH (aq) → products

To determine the products, we need to look at the solubility rules in Table 5.1. Rule 8 states that

hydroxides are insoluble unless that are formed with alkali metal cations or with ammonium.

Since copper is not an alkali metal, this tells us that copper hydroxide will be insoluble in water.

Thus, one of the products will be copper(II) hydroxide. We also know from Rule 5 that almost

all sulfate salts are soluble. Thus, sodium sulfate will not form an insoluble salt. Therefore, the

sodium ion and the sulfate ion are spectator ions in this reaction and can be removed from both

sides of the equation. Using this information, we can now write the unbalanced net ionic

reaction

Cu2+ (aq) + OH- (aq) → Cu(OH)2 (s)

To balance the reaction, we can see that we have one mole of Cu2+ on the left side of the

equation and one mole of Cu2+ on the right side of the equation. However, we only have one

mole of OH- on the left side of the equation and two moles on the right side of the equation.

Therefore, we need to place a 2 in front of the OH- on the left side of the equation. Once we

have done so, we obtain the balanced reaction (check the net charge on both sides)

Cu2+ (aq) + 2 OH- (aq) → Cu(OH)2 (s)

What happens if one mixes aqueous sodium sulfate with aqueous potassium nitrate. Since

potassium sulfate is soluble (Rule 5) and sodium nitrate is soluble (Rule 3), no reaction occurs!

Thus, one would write K2SO4 (aq) + NaNO3 (aq) → no reaction.

Chemistry 113.1

Introduction to Chemical Techniques

Experiment 3. Precipitation reactions (May 2012)

4

POST-LABORATORY QUESTIONS

1. Write the balanced chemical reaction equations for all reactions that yielded a precipitate.

2. Did the color of the solution and the color of the precipitate match? Speculate why or

why not. (This is a very complicated question. We are not expecting four pages to

answer this question after extensive research. We are asking that you think about what

might cause a color difference.)

3. Did any of the solutions that you expected should not have yielded a precipitate actually

yield a precipitate? If so, state why you think this may have happened ?

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Chemistry 113.1 Experiment 1. Density complete solutions correct answers key
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Chemistry 113.1 Experiment 1. Density complete solutions correct answers key Introduction to Chemical Techniques INTRODUCTION Density () is defined as the ratio of the mass (m) of a sample to its volume (V): = m / V Mass and volume are extensive properties of matter-properties that depend on the quantity of a substance. Such properties are not in themselves useful in characterizing or identifying substances. Intensive properties such as density however are useful in identifying substances. Intensive properties are often determined by taking the ratio of two extensive properties measured under constant temperature and pressure conditions. As an intensive property, density can be useful in identifying a substance. Density alone cannot absolutely identify a substance but can be a useful value contributing to an identification. For example, a colorless liquid found to have a density of 1.00 g/mL at 4 0C and 1.0 atmosphere pressure could be water, since this is the known density of water. Additional information would be needed to absolutely identify the substance. In contrast, a colorless liquid found to have a density of 0.85 g/mL at 4 0C and 1.0 atmosphere pressure could not be water. In the experiments below, you will use several methods to determine the volume of samples, both solid objects and liquids. You will use the electronic balance to determine the mass of the samples to 0.001 g (1 milligram). From these measurements, you will determine the densities of these samples. A. DENSITY OF REGULARLY SHAPED OBJECTS For regularly shaped objects, such as cylinders, the volume can be determined by measuring the dimensions of the object with a ruler, then applying the proper formula to determine the volume. In this experiment, you will determine the density of a group of objects (all cylinders) individually, then by a graphical method. 1. Obtain one set of cylinders from your instructor. Record the CODE on the container in your laboratory notebook. Record the color and any other distinguishing characteristics as well. 2. Using the plastic ruler in your kit, measure the diameter(d) and the length (or height-h) of each cylinder to the nearest 0.5 millimeter (e.g. diameter = 13.5 mm = 1.35 cm). 3. Using the electronic balance assigned to you, determine the mass of each cylinder to the nearest milligram (0.001 g). 4. Using the measured diameter and length, calculate the volume of each cylinder in cm3. V= r2h = (d/2)2h Chemistry 113.1 Introduction to Chemical Techniques Experiment 1. Density (May 2012) 2 5. Calculate the density r of each cylinder. Report the calculated value for each cylinder, as well as the value of the mean (average). [See the addendum regarding the mean and mean absolute deviation. These should always be reported whenever three or more determinations of the same quantity are the result of identical experiments.] 6. Graph the data for the cylinders with the mass (g) as the y-axis and the volume (cm3) as the x-axis. Using the straight line fitting function of the graphing software, find the formula corresponding to the best fit of the graphed points to a straight line. The slope of this line is the density (Δm/ΔV). Report this value. Compare it to the average density reported in (5) above. B. DENSITY OF IRREGULARLY SHAPED OBJECTS If an object has and irregular shape, its volume can be determined using Archimedes’ principle which states: An insoluble body completely submerged in a fluid displaces its own volume. Thus, the volume of the displaced fluid is equal to the volume of the irregularly shaped object. 1. Obtain a set of mineral samples from your instructor. Record the code identifying the sample in your laboratory notebook. Note any distinguishing characteristics of the minerals such as color, shape etc. 2. Using the electronic balance record the mass of the samples to the nearest milligram (0.001 g). 3. Place approximately 30.0 mL of water in the 100 mL graduated cylinder. Record the exact volume to the nearest 0.1 mL. 4. Carefully add the sample or samples to the water in the graduated cylinder without causing any water to be lost by splashing. Note: It may be best to determine the combined volume of 2 or more pieces of mineral together. As long as the samples are completely submerged, the greater the increase in volume for the water in the graduated cylinder, the more precise the measurement of density will be. 5. Remove the samples from the cylinder, dry them with a paper towel, return them to the storage container and return them to your instructor. 6. Calculate the density () of the mineral sample. C. DENSITY OF LIQUIDS: CONSTRUCTING A CALIBRATION CURVE AND DETERMINING V% COMPOSITION OF AN UNKNOWN SAMPLE In this experiment, you will prepare a series of liquid mixtures of known composition (percent by volume or V%) and determine their densities. Using your graph of density versus V% you will determine the V% of an unknown sample by measuring its density and comparing it to the graphed values. Chemistry 113.1 Introduction to Chemical Techniques Experiment 1. Density (May 2012) 3 1. Using a small beaker, obtain about 10 mL of alcohol (either methanol or ethanol). Record the name of the alcohol used. 2. Place a clean dry 10 mL graduated cylinder on the electronic balance and tare it to 0.000 g. 3. Carefully transfer 2 mL of the alcohol into the graduated cylinder. Read and record the exact volume to the nearest 0.1 mL. [ Note: When reading the volume, the level being read should be at eye level.] 4. Record the mass of the alcohol. Using the mass and volume, calculate the density and enter it in the data table. 5. Add 1 mL of distilled water to the alcohol in the cylinder. Record the exact total volume now in the cylinder and total mass. [Note: Total volume should now be about 3.0 mL] 6. Add 2 additional mL of distilled water to the contents of the cylinder (the total volume at this point should be approximately 5 mL total). Record the exact total volume and the total mass. 7. Add 2 additional mL of distilled water to the contents of the cylinder (the total volume at this point should be approximately 7 mL total). Record the exact total volume and the total mass. 8. EMPTY the graduated cylinder and dry it. 9. At this point you may need to place the cylinder on the balance and tare to 0.000g once again. Then-add 2 mL of distilled water to the graduated cylinder and record the exact volume and the mass to 0.001 g. Divide the recorded mass of the water by its volume. The value should be 1.00 g/mL, which is the known density of water. If this is not the case consult with your instructor immediately. Obtain a sample of alcohol/water of unknown (to you) V% composition . Determine the volume and mass of two individual 3 to 4 mL portions (known as aliquots) of this sample. Record the exact volume and mass of these two aliquots. Chemistry 113.1 Introduction to Chemical Techniques Experiment 1. Density (May 2012) 4 RECORDING DATA In your laboratory notebook, you should record all data in a format similar to the suggested formats below: A. DENSITY OF REGULARLY SHAPED OBJECTS Sample Code __________________ Description: Table A. The dimensions and mass of each object. Object Diameter (cm) Length or Height (cm) Mass (g) A B C D Note: In your laboratory report, this data should be transcribed into a neatly typed table. A final column should be added giving the calculated value of density , in g/cm3. Using the tabulated values above prepare a graph of mass(g) (y-axis) versus volume (cm3) (xaxis). A straight line fit has a slope equal to the density of the sample [slope = Δm/ΔV]. Report this value of the determined density. Using the table of Materials and their Densities provided, can you identify the material ? B. DENSITY OF IRREGULARLY SHAPED OBJECTS Table B. The mass (or combined masses) of the mineral sample(s), the initial volume of water in the graduated cylinder [Vi (H2O)], the final volume after addition of the sample(s) [Vf (H2O)], and the volume of the sample(s), Vmineral. M (g) [Vi (H2O)] (mL) [Vf (H2O)] (mL) Vmineral (mL) Note: In your laboratory report, this data should be transcribed into a neatly typed table. A final column should be added giving the calculated value of density , in g/cm3. [Recall that 1 mL=1 cm3]. Usi...
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