Unit 6

- 17 -

Ionic and Covalent Bonding Activity

Purpose:

To use several models to find a formula for an ionic compound and covalent molecule.

Materials:

Paper and pencil

Procedure: Ionic Bonds

1] Draw a diagram of a lithium atom. 2] Next to it draw a diagram of a chlorine atom. 3] Draw an arrow to show that the outer electron of lithium moves to the outer energy

level in chlorine. 4] Draw a diagram of a calcium atom. 5] Next to it draw a diagram(s) of a bromine atom.

6] Draw an arrow to show that the outer electrons (of calcium) moves to the outer energy level in bromine.

Answer the questions that follow . . .

Covalent Bonds

1] Draw a diagram of a hydrogen atom. 2] Next to it draw a model of a fluorine atom . . . draw them so that their outer energy

levels overlap. One electron from hydrogen should be paired so they belong to either hydrogen or fluorine.

3] Indicate on your diagram which electrons are being shared between the two atoms. 4] Draw a diagram to show how carbon can form four bonds with four hydrogen atoms.

Answer the questions that follow . . .

- 18 -

Questions:

Ionic Bonding

1] What kind of bond is formed between lithium and chlorine? 2] What kind of bond is formed between calcium and bromine?

3] What is the formula for the compound of lithium and chlorine?

4] What is the formula for the compound of calcium and bromine?

5] What is the formula for the compound of magnesium and chlorine? Draw the diagram to help you . . . if needed

Covalent Bonding 1] How many pairs of electrons are shared in the hydrogen and fluorine compound?

2] What is the formula for the compound?

3] How many pairs of electrons are shared in the hydrogen to carbon compound? 4] What is the formula for that hydrogen to carbon compound?

Unit 6

- 19 -

Ionic Binary Formula Names

Purpose: To write binary formulas for the ions given To learn how to name binary formulas

Materials: Chart

Procedure: 1] Write the cation and the anion from the list. 2] Make sure that the oxidation numbers cancel out.

3] Write the name of the cation 4] Add the root name of the anion, then add an (-ide) to the ending

Data:

Cation / metal Anion / non-metal Sodium Na+1

Chlorine Cl-1

potassium K+1 Bromine Br-1

Magnesium Mg+2 Oxygen O-2

Iron Fe+3 Phosphorous P-3

Cation Anion Formula Binary name Na+1

Cl-1 NaCl Sodium Chloride

Unit 6

- 20 -

Formulas and Oxidation Numbers

Background Information

Oxidation numbers and the charges of ions give the information needed to write the

formulas of many chemical compounds. Only a few guidelines are needed: 1. In a neutral compound, the charges on ions, or the oxidation numbers, balance out to

zero.

2. One positive charge balances one negative charge. 3. Atoms with positive charges or positive oxidation numbers are written first. 4. Subscripts show the relative numbers of atoms or ions in a compound.

5. To show more than one of a polyatomic ion, the symbol is enclosed in parentheses and the subscript follows, for example Al2(SO4)3.

The purpose of this activity is to use paper models to show how chemical formulas are derived from oxidation numbers.

Purpose • write formulas of chemical compounds.

• name chemical compounds. Procedure

1] Make a copy and cut out each ion square, found on the next page 2] Assemble the ions for a compound containing magnesium and chlorine. To do this,

place the Mg+2 ion on a piece of paper. Place enough Cl- ions along side the magnesium ion to balance the charges.

3] Record the formula and name of the compound of magnesium and chlorine in the data table like the one shown.

4] Assemble the ions for five compounds from the list below and record their formulas and

names in your data table. Use the rules listed in the background information and in

your reference book: Aluminum and bromine; sodium and oxygen; iron

(II) and sulfur; aluminum and nitrate ion; potassium and sulfate ion; iron (III) and chlorine ammonium ion and sulfur aluminum and oxygen iron (II) and

sulfate sodium and phosphate

- 21 -

- 22 -

Formulas and Oxidation Numbers - continued

Observations

Substances Chemical Formula Compound Name magnesium + chlorine MgCl2 magnesium chloride

Questions

1] Some compounds are described as “binary compounds.” What does this mean?

2] List the formulas and names of any binary compounds you have constructed. 3] Which element on your list forms ions in two different oxidation states?

4] Parentheses must be used to show more than one polyatomic ion. List the formulas

of any compounds on your list where this was necessary.

Extension

1] Some elements have more than one oxidation number. To show the oxidation

number of such elements in a compound, either a prefix or a roman numeral is given in the name of the compound. Give both names for the following

compounds: a. UF5 b. UF6 c. SeCl2 d. Se2Cl2

2] Manganese has an oxidation number of 4+ in a number of compounds. Write the

formulas and names of two such compounds of manganese with oxygen and with

bromine.

Unit 6

- 23 -

Acid Naming Activity

Purpose: To correctly write the formulas for binary acids and oxyacids. To correctly write the

names of each binary acid or oxyacid.

Materials: Ions chart

Periodic table Procedure:

1] Cut out the ions and paste them into your lab book 2] Record the ions and add the correct number of hydrogens to each

3] Name each of the binary acids using the prefix hydro- and ending each with the –ic ending.

4] Name each oxyacid using the correct ending, -ic for an –ate ending and –ous for

an –ite ending. 5] If oxyacid has more or fewer oxygens then the -ate or -ite forms, be sure to use

a hypo- or per- prefix.

Data:

Hydrogens anions Formula Acid Name

H+1 N-3 H3N hydronitric acid

H+1 CrO4-2 H2CrO4 Chromic Acid

- 24 -

ions

F- I-

H+

Cl- S-2 Br-

ClO4- CO3

-2 IO3-

ClO3- SO3

-2 SO4-2

ClO- PO4-3 PO3

-3

ClO2- NO3

- NO2-

BrO3- CH3COO-

Unit 4

- 25 -

Tests for Iron (II) and Iron (III) In this experiment, the complex hexacyanoferrate (II) ion (ferrocyanide), Fe(CN)46-, and the

hexacyanoferrate(III) ion (ferricyanide), Fe(CN)36-, will be used in identification tests for Fe2-, and Fe3- ions. The charges on the two complex ions clearly indicate the difference in the oxidation state of the iron present in each. The (CN) group in each complex ion has a charge of -1. Thus, iron(II) is present in the ferrocyanide ion, [Fe2+(CN-)6]4-. Iron (III) is present in the ferricyanide ion group, [Fe3+(CN-)6]3-. A deep blue precipitate results when either complex ion combines with iron in a different oxidation state from that present in the complex. The deep blue color of the precipitate is caused by the presence of iron in both oxidation states. The color provides a means of identifying either iron ion. If the deep blue precipitate is formed on addition of the [Fe2+(CN-

)6]4- complex, the iron ion responsible must be the iron(III) ion. Similarly, a deep blue precipitate formed with the [Fe3+(CN-)6]3- complex indicates the presence of the iron(II) ion.

Both of the deep blue precipitates are known to have the same composition. The potassium salt of the complex ion has the formula KFe(CN)6•H2O.

The thiocyanate ion, SCN-, provides a test for confirming the presence of Fe3+ ion. The soluble FeSCN2- complex imparts a blood red color to the solution.

OBJECTIVES Observe tests of known solutions containing iron (II) or iron (III) ions. Compare results for the two ions and infer conclusions. Design a procedure for identifying the two ions in one solution.

MATERIALS • FeCl3, 0.1 M • lab apron • Fe(NH4)2(SO4)2, 0.2 M • microplate • gloves • safety goggles • K3Fe(CN)6, 0.1 M • sheet of paper, white • K4Fe(CN)6, 0.1 M • thin-stemmed pipets (5) • KSCN, 0.2 M

Do not touch any chemicals. If you get a chemical on your skin or clothing, wash the chemical

off at the sink while calling to your teacher. Make sure you carefully read the labels and follow the precautions on all containers of chemicals that you use. If there are no precautions stated on the label, ask your teacher what precautions to follow. Do not taste any chemicals or items used in the laboratory. Never return leftover chemicals to their original containers; take only small amounts to avoid wasting supplies.

Acids and bases are corrosive. If an acid or base spills onto your skin or clothing, wash the

area immediately with running water. Call your teacher in the event of an acid spill. Acid or base spills should be cleaned up promptly.

- 26 -

Procedure

1. Put on safety goggles, gloves, and a lab apron. 2. Check that you have a labeled pipet for each solution listed in the materials. 3. Place the plastic-wrap rectangle on a white sheet of paper. 4. Along the top of the plastic wrap, place 5 drops of a freshly prepared iron (II)

ammonium sulfate solution in each of the three locations shown in Figure 1. 5. Along the bottom of the plastic wrap, place 5 drops of a freshly prepared iron (III)

chloride solution, as shown in Figure 1.

6. Add 1 drop of 0.1 M K4Fe(CN)6 solution to the first sample of iron(II) ions at the

top of the plastic wrap and l drop to the first sample of iron(III) ions at the bottom. Record your observations in Table 1.

7. Add 1 drop of 0.1 M KSCN solution to the second sample of iron(II) ions and 1 drop to the second sample of iron (III) ions. Record your observations in Table 1.

8. Add 1 drop of 0.1 M K3Fe(CN)6 solution to the third sample of iron (II) ions and 1 drop to the third sample of iron (III) ions. Record your observations.

9. Clean all apparatus and your lab station. Return equipment to its proper place. Dispose of chemicals and solutions in the containers designated by your teacher. Do not pour any chemicals down the drain or in the trash unless your teacher directs you to do so. Wash your hands thoroughly before you leave the lab and after all work is finished.

- 27 -

Tests for Iron(II) and Iron(III) - continued

TABLE 1: RESULTS OF TESTS FOR IRON(II) AND IRON(III)

Iron ion Hexacyanoferrate

(II) ion [Fe2+(CN-)6]4-

Hexacyanoferrate

(II) ion [Fe3+(CN-)6]3-

Thiocyanate

ion SCN-

Fe2+

Fe3+

Observations in step 6

Observations in step 7

Observations in step 8

Analysis 1. Explain specifically how you would make a conclusive test for an iron (III) salt. 2. Which test for iron(II) ions is conclusive?

3. When iron (II) ammonium sulfate was mixed with the [Fe2+(CN-)6]4- ion, the

precipitate was initially white but turned blue upon exposure to air. What happened to the iron(II) ion when the precipitate turned blue?

Unit 6

- 28 -

Three-Dimensional Models of Molecules

Background Information A single covalent bond is formed when two atoms share a pair of electrons. Each

atom provides one of the electrons of the pair. If the two atoms are alike, the bond is

said to be non-polar covalent. If the atoms are unlike, one exerts a greater attractive force on the electrons, and the bond is polar covalent. More than one pair of electrons can be shared. This results in a double or triple bond.

A group of atoms held together by covalent bonds is called a molecule. Molecules can be either polar or non-polar. If bonds are non-polar, the molecule is non-polar. If bonds are polar, molecules can still be non-polar if the charge distribution through-out the

molecule is symmetrical. A molecule’s symmetry depends on its shape, that is the positions in space of the atoms making up the molecule. Some possible shapes are linear, angular (bent), pyramidal, and tetrahedral.

Although we represent molecules on paper as being two-dimensional for convenience, they are actually three-dimensional. By building molecular models, chemists

come to understand the bonding, shapes, and polarity of even the most complex molecules.

Purpose:

Build three-dimensional models of some simple covalent molecules. Predict their shapes and polarities from knowledge of bonds and molecule polarity rules.

Materials:

white – hydrogen blue - nitrogen

black – carbon green – chlorine, fluorine red – oxygen, sulfur orange – bromine purple – iodine

Procedure:

1] Begin by creating a Lewis diagram for each of the formulas on the data table.

2] Identify the AXE, for each Lewis diagram to help you solve for the shape and polarity.

Your VSEPR theory chart can be used. 3] Build the models using your Lewis diagrams as a reference. Use the colors above to

represent the elements needed. Short connectors are to be used for single bonds and long connectors are for multiple bonds only.

- 29 -

Observations and Data:

NAME FORMULA AXE SHAPE POLARITY

hydrogen

H2

water

H2O

methane CH4

chlorine Cl2

ammonia NH3

hydrogen floride HF

ethyne C2H2

dichloromethane CH2Cl2

nitrogen N2

carbon dioxide CO2

methanol CH3OH

hydrogen peroxide H2O2

oxygen O2

hydrogen sulfide H2S

- 30 -

Unit 7

- 31 -

Double Replacement Reactions

Background Information:

Some of the most impressive chemical reactions are double-replacement reactions, also called ion-exchange reactions. In a double-replacement reaction, a clear solution of an ionic compound is added to a clear solution of another ionic compound. The positive

ions of one compound react with the negative ions of the other compound to form a precipitate, a gas, or water. When a precipitate forms, the result can be very dramatic. The precipitate is an insoluble solid substance that may produce a sudden bright color in

the remaining clear solution.

Purpose:

In this investigation you will

• perform double-replacement reactions, and • observe and interpret the results

Materials:

potassium iodide KI

sodium chloride NaCl sodium iodide NaI Copper (II)sulfate Cu (SO4)

lead nitrate Pb (NO3)2 silver nitrate Ag (NO3) Copper (II) nitrate Cu (NO3)2

Ammonium hydroxide NH4(OH)

Procedure: 1] On the microplate, place a few drops from the first three sample bottles listed in

the materials list. See the data table 2] On the table, record the color of each chemical used in the appropriate reactant

box. Use clear if it is colorless. 3] Now place a few drops of the next three samples, with its appropriate reactant.

See the data table 4] Now record the color of each chemical used in the appropriate reactant box.

Again use clear if it is colorless.

5] After mixing, record any color changes on the product row of the table 6] After recording all results, use the table on page 11 to figure out which reactant

will form the precipitate.

- 32 -

Observations:

Reactants: KI NaCl Pb (NO3)2 NaI Cu (SO4) Cu (SO4)

Color:

Reactants: Pb (NO3)2 Ag (NO3) Ag (NO3) Cu (NO3)2 NH4(OH) Cu (NO3)2

Color:

Product Color:

ppt:

Questions:

1. What did you observe when you mixed the two solutions together? [general]

2. A double replacement reaction is also called an ion-exchange reaction. Describe the exchange of ions that occurred in this lab investigation.

Unit 6

- 33 -

Single Replacement Reactions

Background Information:

In nature, elements can occur either free, meaning uncombined with other elements, or chemically combined in a compound. The tendency of a particular element to combine with other substances is a measure of the activity of that element. The more active an

element is, the more likely it is to combine. In a single-replacement reaction, an uncombined element replaces a less active element that is combined in a chemical compound. The less active element is then freed from the compound.

For example, in the reaction

Zinc + Copper sulfate Zinc sulfate + Copper

Zinc replaces the less active copper, and combines with sulfate, which frees the copper from the compound.

Purpose:

In this investigation you will observe

• how various metals undergo single-replacement reactions when placed in acid. • if the metat is more active than hydrogen, it will replace it and release the hydrogen

gas.

Materials:

Magnesium - Copper -

Aluminum - Zinc - Iron - Nickel -

Procedure:

1] Label each test tube with the name of one of the metals listed in the materials

2] With your safety goggles on, carefully pour 5 mL of hydrochloric acid into each

test tube.

3] Observe the color, appearance and textures of the metal samples in the material

list and record them in the data table.

4] One at a time, place the appropriate metal in each test tube. Observe what happens to the metal in each test tube. Record your observations in the data table – be sure to record how fast the reaction was, and to feel each test tube as

the reaction proceeds. • If a reaction does not occur, you may record “no reaction to acid” in the

data table

5] Dispose of all materials as directed by your instructor.

- 34 -

Observations:

1. Magnesium

-

2. Aluminum -

3. Iron -

4. Copper -

5. Zinc -

6. Nickel -

Conclusions:

1] Were these reactions endothermic or exothermic? ___________________

Explain. ______________________________________________________

____________________________________________________

2] Which of the metals are more active than hydrogen? ___________________ 3] Which of the metals are less active than hydrogen? ___________________

4] What could you do to prove that hydrogen gas was produced as a result of these

reactions?

____________________________________________________

____________________________________________________

____________________________________________________

5] The rate at which hydrogen gas is produced as a result of these single-replacement

reactions is an idicator of the relative activity of the metals. List the metals in order of their activity from most active to least active.

Unit 7

- 35 -

Toward More Efficient Combustion

Purpose: In this investigation you will • prepare an alkyne and • observe complete and incomplete combustion

Materials:

4 or 5 pieces of calcium carbide, CaC2 250-mL beaker

water, H2O 3 test tubes forceps wood splints

CAUTION: The solution formed when calcium carbide reacts with water is Ca(OH)2 a strong

base. Ca(OH)2 is corrosive; avoid skin contact. Rinse spills with plenty of water.

Procedure: 1] Fill a 250-mL beaker full with tap water. Fill the test tube

completely full of water. Place your thumb over the mouth of

the test tube and invert it into the filled beaker. Remove your hand from the beaker.

2] Using forceps, place a piece of calcium carbide, CaC2 into the beaker. As it reacts with the water to produce a gas, place the mouth of the inverted test tube over the CaC2 and collect the gas

by water displacement. If the test tube does not fill completely with gas, use a second piece of CaC2.

3] When the test tube is filled with gas, remove it from the beaker. Keeping the mouth of the tube down, place a burning splint in the mouth of the tube. Record your observations. Label the test

tube and put aside for future reference and comparison with other test tubes.

4] Repeat the above procedure using test tubes that are full and 1/10 full of water. Record your observations.

5] Before cleaning the test tubes with detergent, record your

observations of the amount of carbon on each. Do not place

burned splints in the sink. Extinguish them under a stream of water and then discard them properly. Dispose of the liquid into a container provided by your instructor.

- 36 -

Analysis:

1] Observations:

2] Calcium carbide reacts with water to produce ethyne gas and calcium oxide. Write a balanced equation for this reaction. ( Calcium oxide reacts with water to form calcium hydroxide ).

3] Describe the amount of soot on each test tube after the ethyne gas was burned.

Test Tubes Observations

full 100% ethyne

full 50% ethyne

1/10 full 10% ethyne

4] Write and balance the chemical equation for the complete combustion of ethyne gas.

Conclusions:

1] Does the amount of oxygen gas (air) in relation to the amount of fuel (ethyne gas) affect the rate of combustion? Which test tube mixture gave you the fastest rate of combustion?

2] Which gas-air mixture produced the greatest amount of air pollution?

3] Which gas-air mixture produced the greatest amount of energy?

4] Another name for ethyne is acetylene. Look this name up in a chemical

dictionary and list some of its uses.

Unit 7

- 37 -

Balancing Equations Using Models

How can molecular models and formula-unit ionic models be used to balance chemical

equations and classify chemical reactions?

Materials

• large and small gumdrops in at least four different colors • toothpicks

Procedure

Examine the partial equations in Groups A–E. Using different-colored gumdrops to

represent atoms of different elements, make models of the reactions by connecting the appropriate “atoms” with toothpicks.

Use your models to (1) balance equations (a) and (b) in each group,

(2) determine the products for reaction (c) in each group, and (3) complete and balance each equation (c). Finally, (4) classify each group of reactions

by type.

Observations

Group A

a. H2 + Cl2 HCl ______________________________________

b. Mg + O2 MgO ______________________________________

c. BaO + H2O ______________________________________

reaction type: ______________________________________

Group B

a. H2CO3 CO2 + H2O ______________________________________

b. KClO3 KCl + O2 ______________________________________

c. H2O ______________________________________

reaction type: ______________________________________

- 38 -

Observations continued

Group C

a. Ca + H2O Ca(OH)2 + H2 ______________________________________

b. KI + Br2 KBr + I2 ______________________________________

c. Zn + HCl ______________________________________

reaction type: ______________________________________

Group D

a. AgNO3 + NaCl AgCl + NaNO3 ______________________________________

b. FeS + HCl FeCl2 + H2S ______________________________________

c. H2SO4 + KOH

______________________________________

reaction type: ______________________________________

Group E

a. CH4 + O2 CO2 + H2O ______________________________________

b. CO + O2 CO2 ______________________________________

c. C3H8 + O2 ______________________________________

reaction type: ______________________________________

Unit 7

- 39 -

Mole and Mass Relationships

Pre-Lab Discussion: In a balanced equation, all reactants and products must be represented by symbols or formulas. The total number of atoms of each element must be the same on each side of

the equation to satisfy the law of conservation of mass. A calculation of the formula mass of a reactant or product enables a researcher to

convert from grams of a particular substance taking part in a reaction to moles of that

substance. The mole relationship given by the coefficients of the balanced equation then allows the researcher to calculate how many moles of every other substance will take part in the reaction.

In this experiment, you will investigate the quantitative relationships in the reaction:

NaHCO3 s( ) + HCl aq( ) NaCl aq( ) + CO2 g( ) + H 2O g( )

A known mass of sodium bicarbonate will be reacting with excess hydrochloric acid. Knowing the mass of NaHCO3(s) that reacts, you can determine from the balanced

equation the mass of NaCl that should be produced. You can compare this theoretical value with the actual experimental mass of NaCl produced.

This experiment should aid in the understanding of the mole-mass relationships that exist in a chemical reaction and in the interpretation of a balanced chemical equation.

Purpose: Compare the experimental mass of a product of a chemical reaction with the mass predicted for that product by calculation.

Materials:

6M hydrochloric acid (HCl) burner

sodium bicarbonate (NaHCO3) ringstand evaporating dish 10 mL graduated cylinder

watch glass disposable pipet

CAUTION: Handle hydrochloric acid with care. Flush any spills with cold water and a dilute solution of sodium bicarbonate and report to your instructor. Do not

lean over the apparatus.

Procedure: 1] Flame dry a clean evaporating dish by heating it in the hot part of a burner flame

for a few minutes. Allow the dish to cool.

2] Find the combined mass of the evaporating dish plus a watch glass. This is mass (a) in your data list.

- 40 -

3] Leaving the watch glass and evaporating dish on the balance. Tare the balance

and add 2.50 g of sodium bicarbonate using a microspatula. Record this mass as

(b) in your data list. 4] Set up the ring stand and burner as shown in figure 1 below. If the dish is

smaller than the iron ring, use wire gauze. Place the watch glass on top of the evaporating dish in the iron ring.

5] Obtain about 5 mL of the 6M hydrochloric acid in a clean, graduated cylinder.

Using the disposable pipet, slowly add HCl to the NaHCO3 in the evaporating

dish, a few drops at a time. See figure 2 below.

Continue adding acid until the reaction (bubbling) stops. Carefully tilt the

evaporating dish back and forth a couple of times to make sure that the acid has

contacted all the NaHCO3. After making sure that all bubbling has stopped,

remove the watch glass and place it curved side up on the lab bench. 6] Holding the burner in your hand, gently heat the evaporating dish. Use a low

flame and move the burner back and forth to avoid splattering. When almost all

of the liquid is gone replace the watch glass and heat gently again until no liquid remains. Allow the dish to cool.

7] Find the combined mass of the watch glass, evaporating dish, and contents

(NaCl). Record this mass (c), in you data list.

- 41 -

Mole and Mass Relationships - continued

Observations:

a] evaporating dish + watch glass __________ g

b] evaporating dish + watch glass + NaHCO3 __________ g

c] evaporating dish + watch glass + NaCl __________ g

Calculations:

1] Find the mass of the NaHCO3 reactant [b-a] __________ g

2] Find the mass of the NaCl reactant [c-a] __________ g

3] Write out the balanced equation:

4] Calculate the mass of NaCl you expected to get when 2.5 g of NaHCO3 was reacted

with HCl. This value is the theoretical yeild.

given

conv.mol req.mol given

req. conv.= required

5] Calculate the percent yield using your theoretical yield above, and your experimental

mass from number 2.

actual

theoretical=% yeild

Conclusion: List and discuss at least three sources of error which occurred in this lab. They should

be based on the percent yield (causes for high or low yield).

Unit 7

- 42 -

Mass and Mass Relationships Pre-Lab Discussion:

Given a balanced equation and the mass of one of the substances in the reaction, the

mass of any other substance can be calculated. Calculations in which a known mass is used to find an unknown mass in a chemical reaction are called mass-mass calculations.

In this experiment, a double replacement reaction will occur when an aqueous solution of either potassium iodide or potassium chromate is mixed with aqueous lead nitrate solution. There are two products of this reaction. One is an insoluble solid, which

will precipitate out of solution. The other is a soluble salt which will remain in solution (see page 11 for a solubility chart). The insoluble solid will be separated from the liquid and dried, and its mass determined. The value of the experimentally measured mass of

the compound will be compared to the theoretical mass of the compound predicted by a mass-mass calculation.

This experiment further emphasizes the importance of mass-mass calculations in the

chemistry laboratory. Purpose:

To compare the theoretical mass of one of the products of a double replacement

reaction with the experimentally determined mass of the sameproduct.

Materials:

potassium iodide (KI) graduated cylinder

potassium chromate (K2CrO4) 2 250-mL beakers lead nitrate (Pb(NO3)2) 250-mL erlenmeyer flask

filter paper / funnel hot plate

Procedure: Using (KI)

1] Measure out exactly 1.70g of potassium iodide (KI). Record this mass as (a) in

your data list. Place the KI in a clean 250-mL beaker and add 50 mL of water. Stir thoroughly to make sure all crystals are dissolved. A hot plate may be used. Label this beaker as “A”.

or

Using (K2CrO4)

2] Measure out exactly 1.00g of potassium chromate (K2CrO4). Record this mass as (aa) in your data list. Place the K2CrO4 in a clean 250-mL beaker and add 50

mL of water. Stir thoroughly to make sure all crystals are dissolved. A hot plate may be used. Label this beaker as “A”.

- 43 -

3] Measure out exactly 2.00g of lead nitrate (Pb(NO3)2). Record this mass as (b) in

your data list. Place it in a clean 250-mL beaker and add 50 mL of water. Stir thoroughly to make sure all crystals are dissolved. A hot plate may be used.

Label this beaker as “B”. 4] Record the appearance of beakers A + B, into the data list as (e).

5] Pour the potassium solution from beaker-A carefully into the lead nitrate solution

of beaker-B. Record the appearance of the beaker, into the data list as (e).

6] Find the mass of a piece of filter paper. Record this as mass (c). Fold the filter

paper and place it in the funnel. Wet the tip with a small amount of tap water.

7] Pour the mixture from the 250-mL beaker into the funnel ( as shown in figure

17-1 ). POUR SLOWLY. Do not allow the liquid to rise above the edge of the

filter paper in the funnel.

8] Rinse the beaker with about 20 mL of water. Pour the rinse water through the

filter. Repeat the rinsing and filterings until all of the precipitate is out of the beaker.

9] Wash the precipitate by pouring about 10 mL of clean water through the filter.

Record the appearance of the filter paper, into the data list as (f).

10] Remove the filter paper and precipitate from the funnel and allow it to dry over

night at about 45°C. You will now be able to find the mass of the dry precipitate

+ filter paper [mass (d) in your data list].

- 44 -

Mass and Mass Relationships – continued

Data :

a. mass of KI g

aa. mass of K2CrO4 g

b. mass of Pb(NO3)2 g

c. mass of filter paper g

d. mass of filter paper + precipitate g

Observations

e. Before filtering:

• what each reactant’s appearance is, before they are mixed together (beaker-A

+ beaker-B),

• and when they are mixed together (final contents of beaker-B)

f. After filtering:

• As you filter the mixture discuss it’s appearance (in the filter paper)

• and the appearance of the filtrate collected (in the flask).

Calculations:

1] Write out the balanced equation: 2] Which material is the limiting reactant? _____________________

3] Calculate the mass that you attained experimentally? [c-b] ___________________

4] Use a mass-mass calculation to solve for the theoretical mass of the material which should have precipitated out, when your limiting amount reacts completely.

given

conv.mol req.mol given

req. conv.= required

5] Calculate the percent yeild:

actual

theoretical=% yeild

Conclusion: List and discuss at least three sources of error which occurred in this lab. They should

be based on the percent yield (causes for high or low yield).

Unit 7

- 45 -

Mole Calculation of Magnesium Oxide

Purpose: In this investigation, you will

• prepare magnesium oxide, • determine the theoretical yield of your product • calculate the percent yeild of your product

Materials: 70 cm magnesium ribbon burner distilled water bottle crucible and cover clay triangle tongs ring stand microspatula

Procedure: 1] Record the mass of a clean, dry crucible and cover in the data table. Obtain a

piece of magnesium ribbon approximately 70 cm long from your instructor. Roll it into a loose ball and mass it in the crucible, with the lid on. Record the mass in the data table.

2] Place the crucible, cover, and magnesium ball on a clay triangle as shown in figure 5-1. Begin to heat the crucible while holding the cover just above the crucible.

3] When the magnesium begins to burn, replace the cover on the crucible. While you

continue to heat the crucible, lift the cover briefly about every ten seconds. Continue this process until the magnesium no longer burns when the cover is lifted.

4] Adjust the crucible and cover on the clay triangle so that the lid is ajar as shown in figure 5-2. This position will allow a steady flow of air into the crucible. Heat strongly for 3 to 5 minutes. Then allow the crucible to cool for 5 minutes.

5] Crush the contents of the crucible into powder using a glass stirring rod. Slowly add 15 drops of water. Note any odor as you add the water.

6] Heat the crucible strongly, without the cover, for 3 to 5 minutes to dry the residue. Cool the crucible and contents for 5 minutes. Mass the crucible, cover, and contents. Record the mass in the data table.

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Mole Calculation of Magnesium Oxide:

Observations: a. mass of crucible, cover, Mg before heating __________ g

b. mass of empty crucible and cover __________ g

c. mass of crucible, cover, Mg after heating __________ g

d. mass of residue produced (MgO) __________ g

e. mass magnesium reacted (Mg) __________ g

f. observations?

•

•

•

Calculations:

1. Write the balanced equation for this reaction.

2. What is the mass of the residue produced experimentally in the reaction? 3. Using your lab data, calculate the theoretical yeild for magnesium oxide being

produced in the reaction.

given

conv.mol req.mol given

req. conv.= required

4. Using your lab data, calculate the percent yeild of magnesium oxide produced.

actual

theoretical=% yeild

Conclusions: Describe and discuss three sources of error encountered in the experiment. They should

be based on the percent yield (causes for high or low yield).

Unit 7

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Aluminum to Alum

Purpose: In this inestigation you will

• prepare potassium aluminum sulfate, and • demonstrate the mass-mass relationship in chemical reactions

Materials:

aluminum foil about 1g filter paper / funnel

4g of potassium hydroxide (KOH) ring stand / iron ring 12 mL concentrated sulfuric acid (H2SO4) hot plate ice bath 2 250-mL beakers

CAUTION: The preparation of KOH and H2SO4 solutions are very exothermic.

These substances are corrosive: avoid skin contact. Rinse spills with plenty

of water. Dispose of materials as directed by your instructor.

Procedure: 1] Obtain a piece of aluminum foil having a mass of 1 g, and record its mass to the

nearest 0.01 g. Place about 4 g of potassium hydroxide in a beaker, and add 50

mL of water to the KOH in the beaker. Stir to dissolve. 2] Tear up the aluminum foil into small pieces and place them into the beaker with

the KOH solution. Label the beaker and place it in the fume hood for about 15

minutes. 3] While you are waiting, place 13 mL of waterin a second small beaker and slowly

add 12 mL of concentrated H2SO4. REMEMBER: add acid to water. Add very

small amounts, stir then add more very small amounts. Carefully place the beaker containing the hot H2SO4 solution in an ice water bath.

4] After the aluminum foil has dissolved in the KOH solution,

place the 250 mL beaker into an ice water bath to cool the hot solution. Allow the black residue to settle to the bottom of the

beaker. This reaction is

2Al + 2KOH + 6H 2O 2KAl OH( )4 + 3H 2

5] Set up the filtration apparatus while you are waiting. Filter the

KAl(OH)4 into the cold sulfuric acid solution. Slowly stir the white precipitate until all of it dissolves. Label the beaker with your group’s information and set it aside overnight.

6] Decant the liquid from the crystals. Wash the liquid down the

drain with large amounts of water. Rinse the crystals once

with water. Decant. Set the crystals aside to dry. 7] Remove the alum crystals from the beaker. Obtain the mass

of the dry crystals.

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Observations:

Day 1 :

Day 2 :

Calculations:

a. mass of Aluminum [given] g

b. mass of Alum crystals [recovered] g c. Formula mass of Alum g

d. Theoretical :

given

conv.mol req.mol given

req. conv.= required

e. Percent Yeild :

actual

theoretical=% yeild

Conclusion:

List and discuss at least three sources of error which occurred in this lab. They should be based on the percent yield (causes for high or low yield).