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Representing Changes inChemical Systems with Equations

IntroductionIn this experiment you will carry out some changes in chemical systems and

record your observations on the report sheet. All of the changes can be representedby balanced chemical equations that you will be asked to write. Pay careful attentionto the physical state of each substance (solid(s), liquid(l), gas(g), or dissolved in aqueoussolution(aq)) and use subscripts as indicated. Also, if ions are present in aqueoussolution, show each ionic species with its charge and with the (aq) subscript toindicate that it is solvated by the solvent water molecules. For example, solid NaCl(s)

dissolves in water to give an aqueous solution of Na+(aq) and Cl-

(aq) ions. Thissodium chloride solution can also be written in shorthand as NaCl(aq) which standsfor Na+

(aq) + Cl-(aq), but in this experiment use the latter complete ionic notation andnot the shorthand. The changes you will observe in this experiment include simplephase changes (e.g., the melting of ice, the evaporation of water, or the sublimationof “dry ice”, CO2(s)), the process of a solid dissolving in water or precipitating fromaqueous solution, and more complicated chemical reactions that involve gas evolution,pH change, temperature change, color change, or formation of new compounds withdifferent properties. As an example, the reaction of solid magnesium metal with anaqueous solution of sulfuric acid gives the following observations: evolution of acolorless gas, increase in temperature (heat is produced), decrease in the amount ofmagnesium metal (visually detected), and decrease in the acidity (increase in pH) ofthe solution as detected by pH indicator. The balanced chemical equation thatrepresents this change is:

Mg(s) + 2H+(aq) + SO4

2-(aq) → Mg2+

(aq) + SO42-

(aq) + H2(g)

All equations must be mass and charge balanced and the complete ionic notation is tobe used. The (aq) subscript means this species is completely dissolved in the solventwater.

The background information for this experiment is contained in a high schoolchemistry course and in chapters 1 - 4 of your text. You can do this experiment andmake the observations without much background review, but you will have to referto your text to write the chemical equations. The following is a brief summary ofsome of the key concepts that will help you do this experiment.

Classification of Substances. Solid pure substances are described by the typeof bonding present: ionic, molecular, or metallic. A solid ionic compound (or salt)contains positive and negative ions that are held together by strong electrostaticforces (ionic bonding). These compounds are generally high melting, brittle solids.

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Most ionic compounds dissolve in water and dissociate completely to give aqueoussolutions of solvated ions as illustrated for table salt:

NaCl(s) + excess H2O(l) → Na+(aq) + Cl-(aq).

Aqueous solutions of ions conduct electricity. You will explore this property inExperiment 2. Note that the above equation appears not to be balanced since H2O(l) is notexplicitly listed as a product. Since water is the solvent in a dissolution reaction, it is notgenerally written as a reactant or product but is assumed to be present. The (aq) indicatesthat the product ions are solvated and the term excess H2O(l) is included to make clear it isthe solvent. The term excess H2O(l) could also have been written above the arrow.

In a molecular compound the bonding is due to the sharing of electrons to givecovalent bonds. Some examples are oxygen O2(g), iodine, I2(s), carbon dioxide CO2(g),water H2O(l), methane CH4(g), ammonia NH3(g) and sucrose (table sugar)C12H22O11(s). The molecular formula represents the distinct molecule where theatoms are held together by strong covalent bonds. At room temperature, molecularcompounds can be gases, liquids, or solids depending on the number of atoms in themolecule and on the details of their bonding. All gases and essentially all liquids atroom temperature are molecular in nature. Only very “polar” molecules are watersoluble (sucrose is an example) but many dissolve in molecular “organic” solventssuch as carbontetrachloride (CCl4(l)), a dry cleaning solvent. An example is I2(s)

dissolved in CCl4(l) to give a deep red-purple solution of solvated I2 molecules.Many molecular compounds can be made to undergo phase changes by adding orremoving heat, such as the melting of ice (H2O(s) → H2O(l)) or the sublimation of“dry ice” (CO2(s) → CO2(g)). Note: The reverse of this reaction is also calledsublimation. In solid molecular compounds, the distinct molecules are held togetherby weaker forces of attraction. For example, in ice, the distinct water molecules areheld together by weaker hydrogen bonds and dipole-dipole interactions. It'simportant to realize that it is the distinct molecular formula that represents themolecule. When H2O(l) boils it becomes H2O(g). The strong covalent bonds are notbroken in a phase change or dissolution process. When C12H22O11(s) dissolves inwater it becomes C12H22O11(aq) as represented in the equation:

C12H22O11(s) + excess H2O(l) → C12H22O11(aq).

There is a special subclass of molecular solids where the entire solid is onegiant or macromolecule (macromolecular compound). An example is a crystal ofdiamond which consists of an array of C atoms held together by strong covalentbonds. The crystal is one molecule so its impossible to write a true molecularformula for diamond (sometimes we write Cx to represent this). We will not dealwith macromolecular compounds in this experiment.

Metals are characterized by metallic bonding, a type of bonding in which theelectrons are much freer to move around than in other kinds of substances. This“delocalized” bonding leads to solids that are good conductors of electricity, have ashiny appearance, are soft (not brittle), and do not dissolve in solvents. As with

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macromolecular substances, we do not write a molecular formula for metals butsimply use the notation for the element, such as Fe(s) for iron.

The terms ionic, molecular, macromolecular, and metallic are somewhatarbitrary, and some substances have properties that place them in a borderlinecatagory, somewhere between one bonding type and another. You should thereforeexpect some substances not to fit neatly in one of the above catagories. It is veryuseful, however, to consider the general characteristics of typical ionic, molecular,and metallic substances, since many common substances can be readily assigned toone catagory or another.

Representing Substances by Chemical Formulas. Formulas convey a lot ofinformation about a compound (such as the types and number of atoms and thephysical state of the substance) but you must use some chemical knowledge tointerpret the formula. For example, comparison of the formulas NaCl(s) and I2(s)does not immediately tell you that one is ionic and the other molecular (or covalent).How do we know that NaCl(s) is an ionic salt and I2(s) is a molecular compound? Youcould do experiments and determine that NaCl(s) dissolves in water to give aconducting solution of ions and it has other typical ionic properties (high melting,brittle), while I2(s) is not soluble in water but readily dissolves in organic solventssuch as carbon tetrachloride, it is low melting, and volatile (sublimes to I2(g)).

Experience will tell you that most covalent compounds contain a combinationof nonmetal elements (elements of groups 3A - 7A of the periodic table above darkdiagonal line) or compounds formed between identical elements. Ionic substancesthat contain two types of elements (binary ionic compounds) result from thecombination of atoms from opposite ends of the periodic table, i.e. the combination ofa group 1A or 2A element with a group 6A or 7A element. Examples are calciumoxide, CaO(s), rubidium bromide, RbBr(s), and cesium sulfide, Cs2S(s). These are theelements that form stable ions of the following charges: group 1A elements form 1+charged ions, 2A form 2+, 6A form 2-, and 7A form 1-. The metallic elements ingroup 3A also commonly form ions with 3+ charge, such as Al3+. The overallformula of an ionic compound must be neutral.

? Write the formula for magnesium iodide: .

Many ionic substances also contain polyatomic ions such as nitrate, NO3-, sulfate, SO4

2-,carbonate, CO3

2-, and ammonium, NH4+. The bonding within the polyatomic ions is

primarily covalent so these ions do not dissociate in aqueous solution and remainintact. These polyatomic ions can be combined with ions from the groups mentionedabove {e.g. calcium nitrate = Ca(NO3)2 and sodium sulfate = Na2SO4} or with eachother {e.g. ammonium carbonate = (NH4)2CO3} to give neutral ionic substances. Thefollowing is a listing of the most common polyatomic ions you will encounter in thiscourse. These formulas, charges, and names must be memorized. A more extensivelist is given in your text.

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Some Common Polyatomic Ions Memorize the names and formulas!

ion formula name ion formula nameNH4

+ ammonium NO2- nitrite

NCS- thiocyanate NO3- nitrate

SO32- sulfite CO3

2- carbonateSO4

2- sulfate HCO3- bicarbonate

OH- hydroxide ClO2- chlorite

CN- cyanide ClO3- chlorate

PO43- phosphate ClO4

- perchlorateCrO4

2- chromate MnO4- permanganate

Cr2O72- dichromate O2

2- peroxideC2H3O2

- acetate

? Write the formula for calcium permanganate: .

? Write the equation that represents the dissolution and complete ionic dissociation ofsolid calcium permanganate in water:

Some Elementary Acid-Base Chemistry. You will need to know some of thefundamentals of acids and bases to do this experiment. According to the Arrheniusdefinition, acids are substances that produce H+

(aq) ions when they are dissolved inwater. If the dissociation is complete, the acid is called strong (e.g. HCl, HNO3,H2SO4, see eq 1) and if the dissociation is only partial, the acid is called weak (e.g.acetic acid HC2H3O2, see eq 2).

HNO3(aq) + H2O(l) → H+(aq) + NO3

-(aq) eq 1

HC2H3O2(aq) + H2O(l) H+

(aq) + C2H3O2-(aq) eq 2

The single arrow in eq 1 means the reaction is essentially complete. The doublearrows in eq 2 means that the reaction is reversible (can go in both directions) andreaches an equilibrium position not complete in either direction. In this case thepositon of equilibrium lies to the left and the acetic acid only partially dissociates intoions.

Bases are substances that produce OH-(aq) ions when they dissolve in water.

When a salt that contains OH- dissolves in water it dissociates into ions and theresulting solution is a called basic. Salts that contain OH- ions are strong basesbecause they completely dissociate into ions. An aqueous solution of ammonia alsoresults in a basic solution by reacting with water (eq 3).

NH3(aq) + H2O(l) NH4+ + OH-

(aq) eq 3

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An ammonia solution is often called an ammonium hydroxide solution for obviousreasons. The position of the equilibrium in eq 3 lies to the left so ammonia is a weakbase.

It is convenient to use the concept of pH to put the acid-base concept on morequantitative grounds. Pure deionized (DI) water undergoes a partial self-ionization(eq 4).

H2O(l) H+

(aq) + OH-(aq) eq 4

The position of equilibrium of eq 4 lies far to the left and at room temperature theconcentrations of H+

(aq) and OH-(aq) both equal 1 x 10-7 moles/liter or [H+] = [OH-] =

1 x 10-7 (the brackets indicate molar concentration). This is defined as a neutralsolution. The pH is defined as the -log10[H

+] so the pH of a neutral solution where[H+ ] = 1 x 10-7 is -log10[1 x 10-7] = 7.0. The pOH is similarly defined (pOH = -log10[OH-]) so the pOH of a neutral aqueous solution is also 7.0. It will be shownlater in the course that for an any aqueous solution at room temperature thefollowing relationship will always hold: pH + pOH = 14.0. An acidic solution is onethat contains an excess of H+ ions so [H+] > 1 x 10-7 and the pH < 7.0. A basic solutioncontains an excess of [OH-] ions so [OH-] > 1 x 10-7 and pOH < 7.0 and pH > 7.0. ThepH of a solution can vary depending on the strength and concentration of the acid orbase present. In this experiment you will measure pH qualitatively with the aid of auniversal indicator. Later in the course you will use a pH meter to determine pHquantitatively.

Examples of Some Reactions that Occur in Solution. In this experiment youwill carry out a variety of reactions that can be represented by chemical equations.Most of these will fall into one of the following classifications of chemical reactions.

a) A precipitation reaction occurs if upon mixing two salt solutions a solidprecipitates from the solution. A solid will precipitate if any one of the possiblecombinations of the ions present forms an insoluble salt. The solubility rules given inyour text will enable you to determine which salts are insoluble. An example of aprecipitation reaction is the appearance of the white solid AgCl(s) upon mixingaqueous solutions of AgNO3(aq) and NaCl(aq) as shown by the following equation(AgCl(s) is insoluble in water):

Ag+(aq) + NO3

-(aq) + Na+

(aq) + Cl-(aq) → AgCl(s) + Na+(aq) + NO3

-(aq).

The net ionic equation summarizes the key change by eliminating the spectator ionswhich are common to both sides of the equation:

Ag+(aq) + Cl-(aq) → AgCl(s).

If no precipitate forms upon mixing two salt solutions, no reaction takes place andthere is no net ionic equation.

b) Acid-base neutralization reactions occur when you mix a solution of anacid with a solution of a base. An example is the reaction of an aqueous solution ofHCl(aq) (H+

(aq) + Cl-(aq)) with an aqueous solution of NaOH(aq) (Na+

(aq) + OH-(aq))

shown in the following equation:

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H+(aq) + Cl-(aq) + Na+

(aq) + OH-(aq) → H2O(l) + Na+

(aq) + Cl-(aq).

When exactly equal molar amounts of acid and base are reacted (equivalence orneutralization point) the solution contains only a salt and water, in this case Na+

(aq) +

Cl-(aq). If the acid and base are both strong (completely dissociated into ions), thefinal pH will be 7.0. If either acid or base is weak, the pH may not be exactly 7.0 atthe equivalence point. The net ionic equation for a neutralization reaction between astrong acid and a strong base is:

H+(aq) + OH-

(aq) → H2O(l).

c) Oxidation-reduction or redox reactions occur when one or moreelectrons are transferred from one reactant to another. These reactions can beidentified by noting a change in oxidation state of an element during the reaction.An example is the reaction of magnesium with sulfuric acid:

Mg(s) + 2H+(aq) + SO4

2-(aq) → Mg2+

(aq) + SO42-

(aq) + H2(g).

Mg changes oxidation state from 0 to +2 (loss of two electrons, oxidation) and thetwo H+ ions change from +1 to 0 in H2 (gain of two electrons, reduction). In thisreaction, Mg(s) is oxidized and H+

(aq) is reduced. Redox reactions are common andoften involve gas evolution, color changes, and the production of heat.

? What is net ionic equation for this reaction?

ProcedureFor each procedure below, use the report form to record your observations as

they are made. Check for temperature changes by touching the outside of the testtube with your fingers. Write out the complete and balanced chemical equation whichrepresents each change. Usually one equation will represent the change in each partbut in some cases several equations may be used if several changes occur. In suchcases, one equation that represents the overall change (sum of the separate steps) canbe written, but it is acceptable to leave the individual equations as long as each iscorrect and balanced.

If you work in pairs, each of you must make and record all observations in your own wordsand independently write the chemical equations and answers to questions, but you maydivide up the individual experimental procedures. This experiment is expected to requireapproximately one and a half lab periods and it must be turned in by the end of the secondperiod. Therefore you will need to write the corresponding chemical equation immediately aftereach observation.

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Safety• All chemicals should be treated with respect. Avoid contact with your skin

and clean spills promptly. • If a chemical comes in contact with your skin, rinse it off immediately with

plenty of water and check with your TA for further treatment.• Some chemical reactions can be unexpectedly rapid and evolve gas and heat,

so never point an open reaction vessel at yourself or another person whencarrying out a reaction.

• ALWAYS wear your safety goggles in the laboratory.

Waste disposalCarefully follow all directions for waste disposal or treatment to prevent any

hazardous waste from entering the drainage system. Set aside two large beakers(400 - 600 mL) for collecting waste. Label one beaker acid/base waste and the otherExpt. 1 waste. The Expt. 1 waste will be placed in the appropriate container at thecompletion of the experiment. The acid/base waste will be neutralized at thecompletion of the experiment and rinsed down the drain.

Be responsible and minimize waste! For each of the following procedures, take onlya minimal amount of the prepared solutions from the dispensing bottles. In mostcases you only need 1 mL so determine the approximate level of 1 mL in a 6” testtube and take only slightly more. Use 6” test tubes for obtaining the solutions and 4”test tubes for carrying out most of the reactions. All test tubes must be clean butthey can be slightly wet with DI water since none of the experiments require highlyaccurate concentrations.

NOTE: If you are already familiar with writing complete and net ionic equations,this Experiment will provide an excellent review. If you are just learning thisimportant skill, you will need to make an extra effort to master writing ionicequations by the completion of this lab report! Seek help from your TA!

The order in which you do the following procedures is not important, so it is best foreach person or group to begin with a different one.

A. Silver nitrate, AgNO3(aq), and calcium chloride, CaCl2(aq)

Place about 1 mL (20 drops) of 0.1M AgNO3(aq) in a clean 4” test tube. Add upto 8 drops of 0.1M CaCl2(aq), dropwise, and record your observations. The chemicalequation that represents the change in this reaction is given below as an example.Write this on your report sheet. Dispose of the mixture/solution into your Expt. 1waste beaker and rinse the contents of your test tube into the waste beaker using aminimal amount of water.

2Ag+(aq) + 2NO3

-(aq) + Ca2+

(aq) + 2Cl-(aq) →→→→ 2AgCl(s) + Ca2+(aq) + 2NO3

-(aq)

Note that the above equation is mass and charge balanced, ionic species are shownwith their correct charges, and all species have subscripts indicating their physicalstate ((aq) means the species is soluble in aqueous solution). Use this notation inwriting all equations in this experiment.

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B. Sodium carbonate, Na2CO3(s) (washing soda) and hydrochloric acid solution,HCl(aq) (The products of the this reaction are sodium chloride, water andcarbon dioxide.)Place a small (pea sized) sample of solid Na2CO3(s) in a clean 4” test tube. Addup to 1 mL of 1M HCl(aq), dropwise, and record your observations. Disposeof this solution in your acid/base waste beaker.

C. Copper(II) sulfate, CuSO4(aq), and excess sodium nitrite, NaNO2(aq)

Place about 1 mL of 0.1M CuSO4(aq) in a clean 4” test tube. Add up to 8 dropsof 1M NaNO2(aq), dropwise, with mixing. Record your observations. Disposeof this solution into your Expt. 1 waste beaker.

D. Ammonium nitrate, NH4NO3(s), and excess waterPlace about 1 mL of deionized (DI) water in a clean 4” test tube. Add a sampleof solid NH4NO3(s) about the size of two peas and shake to dissolve. Recordyour observations. Dispose of this solution into your Expt. 1 waste beaker.

E. Iron(III) chloride, FeCl3(aq), and sodium hydroxide, NaOH(aq)

a) Place 10 drops of 0.1M FeCl3(aq) in a clean 4” test tube. Add up to 6drops of 1M NaOH(aq), dropwise, and record your observations.

b) Now add up to 10 drops of 1M HCl(aq), dropwise with mixing, andrecord your observations. Dispose of this final solution into your Expt.1 waste beaker.

F. Ammonia gas, NH3(g), and hydrochloric acid gas, HCl(g)Ammonia and hydrochloric acid are both gases under normal conditions and

readily escape their concentrated aqueous solutions. Carefully and “indirectly” smellthe 6M NH3(aq) solution and you will experience ammonia gas.

? Where in your everyday life have you experienced this smell?

HCl(g) also escapes its aqueous solution, but do not smell this gas as its very irritating.In order to observe the reaction of the two gases, place one drop of 6M HCl(aq) andone drop of 6M NH3(aq) about 2 cm apart on a watch glass. Cover the two dropswith a small inverted beaker to minimize air currents. Observe the space around thedrops for several minutes and record your observations. Rinse the drops into theacid/base waste beaker.

CCAAUUTTIIOONN:: 6M HCl(aq) is harmful to your skin and clothing. If you get some onyour skin, flush it thoroughly with water and ask your TA about further treatment.If you spill some on your clothing or on your desk, you can neutralize it with asolution of sodium bicarbonate and flush with water.

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G. Aluminum metal, Al(s), and hydrochloric acid solution, HCl(aq)Place 10 drops of DI water in a clean 6” test tube and add 10 drops of 6MHCl(aq). Mix the solution. Add a small piece of Al(s) and push it down intothe acid solution with a glass stirring rod. Record your observations. Disposeof the mixture/solution into your Expt. 1 waste beaker and rinse the contentsof your test tube into the waste beaker using a minimal amount of water.

H. Use of universal pH indicator to observe the reaction of acetic acid,HC2H3O2(aq), with the base, ammonium hydroxide, NH4OH(aq)A few drops of liquid universal indicator added to a colorless solution willpermit the determination of the approximate pH by observing the color of theindicator in the solution. A neutral solution such as pure DI water has a pH =7.0. Test a sample of 1 mL of DI water by adding 2 drops of universalindicator and observing the color. Record the color on your report sheet.Later you will have to neutralize your acid/base waste solution to this color.You may want to save a test tube of DI water plus indicator for future colorcomparison. Place 15 drops of 1M NH4OH(aq) and two drops of universalindicator in a different clean 4” test tube and observe the color. To thisNH4OH(aq) solution add up to 21 drops of 1M HC2H3O2(aq) dropwise, andafter each addition record the colors in the table provided on the report sheet.Mix the solution after each addition before noting the color. Dispose of thesolution into the acid/base waste beaker. Note: The concentrations of the givensolutions of NH4OH(aq) and HC2H3O2(aq) must be close to 1M or neutralization willnot occur at the correct point and the full range of color changes may not be observed.

I. Aluminum metal, Al(s), and copper(II) chloride solution, CuCl2(aq)

Place about 1 mL of 0.5M CuCl2(aq) in a clean 4” test tube. Add a 1 cm piece ofaluminum wire making sure that the metal is mostly in the solution. Let thereaction proceed for a few minutes, remove the aluminum metal, andobserve. Record your observations on the report sheet. Dispose of themixture/solution into your Expt. 1 waste beaker and rinse the contents ofyour test tube into the waste beaker using a minimal amount of water.Recover the aluminum metal.

J. Copper metal, Cu(s), and silver(I) nitrate solution, AgNO3(aq)

Place about 1 mL of 0.1M AgNO3(aq) in a clean 4” test tube. Add a 2” piece ofpolished copper wire and observe for about 10 minutes. Record yourobservations. Dispose of the mixture/solution into your Expt. 1 waste beakerand rinse the contents of your test tube into the waste beaker using a minimalamount of water. Recover the copper wire.

K. Sublimination of iodine, I2(s) (The sublimation of I2(s), described below, will bedemonstrated by your TA.)Place a few crystals of iodine in a 100 mL beaker, cover the beaker with asmall watch glass, concave side up, and place a few chips of ice on the watchglass. Place a hotplate under a benchtop hood, fill a larger beaker about two-

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thirds full of water and heat it. Briefly hold the bottom of the beakercontaining the iodine crystals in the water to gently warm them. Observe anycolor change in the beaker. Remove the beaker from the water, let it cool, andexamine the underside of the watch glass. Record your observations. Place thewaste directly into the gallon size Expt. 1 waste container.

CC AA UU TT II OO NN :: Iodine vapor is toxic so avoid releasing it into the lab. Heat gently

and never heat the I2(s) without a watch glass with ice completely covering the top ofthe beaker. Do this experiment under the bench top fume hood to avoid gettingiodine vapor into the lab atmosphere.

Waste DisposalTransfer the contents of your Expt. 1 waste container to the appropriately

labeled container in the lab. Place the Al and Cu metal pieces in the container labeledUSED Al or USED Cu.

You will need to neutralize the solution in your acid/base waste beaker sothat it can be safely flushed down the sink drain. Refer to the universal indicatorcolors for neutral, acidic, and basic solutions from part H, and add either strong acid(1M HCl(aq)) or strong base (1M NaOH(aq)) dropwise until the color of the wasteliquid is yellow-green, indicating a neutral solution. Rinse the neutralized contents ofthe acid/base waste beaker down the drain.

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Name Date

Name of your TA:

Pre-Laboratory Questions for Experiment 11. Write the chemical formula for ammonium phosphate. Do not write chemical

equations on the Report Sheets before you come to lab. The following ReportSheets are not part of the Pre-Laboratory Questions.

2. Write the balanced complete ionic equation that represents the reaction in whichsolid ammonium phosphate dissolves in excess water to give an aqueoussolution. Use the complete ionic equation notation that shows the actual formsof all reactants and products. This notation will be used throughout thisexperiment.

3. Write the balanced complete ionic equation that represents the reaction of asolution of barium nitrate with a solution of sodium sulfate. (Hint: check thesolubility rules in Table 4.1 of your text.)

4. Write the net ionic equation for the reaction in question 3.

5. Write the balanced complete ionic equation that represents the reaction of a nitricacid solution with a solution of calcium hydroxide. (Even though Ca(OH)2 isnot a very soluble compound, dilute solutions are possible.)

6. Write the net ionic equation for the reaction in question 5.

7. Which one of the two reactants is undergoing oxidation in the followingREDOX reaction? Circle the reactant of your choice.

2Na(s) + Cl2(g) → 2NaCl(s)

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Name Date

Partner’s name

Name of your TA:

Experiment 1 REPORT SHEETRepresenting Changes in Chemical Systems with Equations

In the spaces provided, record your observations for each of the changes/reactions asyou carry them out and write out the balanced complete ionic equation that representseach change. In some cases you will also write the net ionic equation. Answer thequestions if present. Use descriptive detail to describe your observations. Forexample, describe color, odor (if any), physical form of solids such as powder orcrystals, change in temperature, evolution of a gas, formation of a precipitate,dissolving of a solid, change in pH, etc. Solutions may be described as colorless orassigned a color, but not as clear, since all solutions are assumed to be clear.

A. Silver nitrate, AgNO3(aq), and calcium chloride, CaCl2(aq)Record your Observations:

Write the Complete Ionic Equation: (given as an example)

2Ag+(aq) + 2NO3

-(aq) + Ca2+

(aq) + 2Cl-(aq) →→→→ 2AgCl(s) + Ca2+(aq) + 2NO3

-(aq)

Write the Net Ionic Equation:

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B. Sodium carbonate, Na2CO3(s) (washing soda) and hydrochloric acid solution,HCl(aq)Record your Observations:

Write the Complete Ionic Equation:

Write the Net Ionic Equation:

C. Copper(II) sulfate, CuSO4(aq), and excess sodium nitrite, NaNO2(aq)Record your Observations:

Write the Complete Ionic Equation: (Hint: Cu2+(aq) ions often bind four anionic

species called ligands, L-(aq), giving highly colored complex polyatomic ions CuL4

2-(aq).

The NO2-(aq) ion is an example of a ligand that binds Cu2+

(aq), so the key product ofthis reaction is Cu(NO2)4

2-(aq)). Complete and balance the following complete ionic

equation for this reaction:

Cu2+(aq) + SO4

2-(aq) + Na+

(aq) + NO2-(aq) →

Write the Net Ionic Equation:

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D. Ammonium nitrate, NH4NO3(s), and excess waterRecord your Observations:

Write the Complete Ionic Equation:

This reaction is used in a commercial product called a Cold Pack. This productconsists of solid NH4NO3(s) and a breakable plastic tube that contains water, allsealed in a flexible plastic pouch. The reaction is initiated by breaking the plastic tubethus releasing the water and dissolving the salt. Look for this product at a drugstore.

E. (Part A) Iron(III) chloride, FeCl3(aq), and sodium hydroxide, NaOH(aq)Record your Observations:

Write the Complete Ionic Equation:

Write the Net Ionic Equation:

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E. (Part B) (addition of HCl(aq))Record your Observations:

Write the Complete Ionic Equation: (Hint: What’s the product of the reactionbetween OH-

(aq) and H+(aq)?)

Write the Net Ionic Equation:

Tap water contains dissolved iron salts and if the pH gets too high a rust coloredsolid can form. This is a problem in swimming pools. Based on your observationsabove, what would you add to a swimming pool to dissolve the solid?

F. Ammonia gas, NH3(g), and hydrochloric acid gas, HCl(g)Record your Observations:

Complete the Equation:

NH3(g) + HCl(g) →

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G. Aluminum metal, Al(s), and hydrochloric acid solution, HCl(aq)Record your Observations:

Write the Complete Ionic Equation: (Hint: Aluminium and magnesium are both“active” metals and react similarly with acid solution (see introduction). What’s thecommon positive charge or oxidation state of aluminum ions?)

H. Use of universal pH indicator to observe the reaction of acetic acid,HC2H3O2(aq), with the base ammonium hydroxide, NH4OH(aq)

Write the complete ionic Equation for the reaction between acetic acid, HC2H3O2 (aq),and ammonium hydroxide, NH4OH(aq):

Record your Results: (Write the colors of the following solutions that contain twodrops of universal indicator.)

solution + universal indicator color of DI waterDI water

# of drops of 1M HC2H3O2(aq) color of solutionadded to 15 drops of 1M NH4OH(aq)

0 drops 2 drops 4 drops 6 drops 8 drops 10 drops

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11 drops 12 drops 13 drops 14 drops 15 drops 16 drops 17 drops 18 drops 19 drops 20 drops 21 drops

What is the color of the indicator for a “basic” solution?

What is the color of the indicator for an “acidic” solution?

What is the color of the indicator at the neutralization or equivalence point?

How many drops of 1M HC2H3O2(aq) were required to reach the equivalence orneutralization point?

Is the pH of the final solution (after addition of all 21 drops of acid) >7, = 7, or <7?

Write the formulas for the products of the reaction:

I. Aluminum metal, Al(s), and copper(II) chloride solution, CuCl2(aq)Record your Observations:

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Write the Complete Ionic Equation: (Hint: There are actually two reactions occuringhere. One produces H2 gas and is similar to the reaction you observed above in partG. We are interested in the other reaction in this part, so ignore the fact that H2 gasis formed. Focus on the reaction between the Al(s) and the Cu2+

(aq) ion in solution.)

J. Copper metal, Cu(s), and silver(I) nitrate solution, AgNO3(aq)Record your Observations:

Write the Complete Ionic Equation (hint: Cu2+(aq) ion is blue in color):

K. Sublimation of iodine, I2(s)Record your Observations (note that there are two processes):

Write the Chemical Equations for both processes:

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Name Date

Name of your TA:

Post-Laboratory Questions for Experiment 11. List by letter all of the reactions in this experiment that involve acid-base

neutralization. (Hint: Look for reactions where water plus a salt are produced from thereaction of H

+ with OH

-, but watch out for disguised bases such as NH3 and CO3

2-.

2. List by letter all of the reactions in this experiment that involve oxidation-reduction.

3. List by letter all of the reactions in this experiment that involve precipitation ofan insoluble salt from aqueous solution.

4. You will notice that throughout this experiment you are given explicitinstructions for disposing of the chemicals you use. You either treated thechemical waste, if necessary, and flushed it down the drain or collected it forproper packaging (including labeling) and disposal. Based on the wastetreatment procedures in this experiment, answer the following questions.NOTE: H+ and OH- may be safely "sewaged" over a very restrictedconcentration or pH range.

a) List three cations in this experiment that may be safely disposed of down thedrain:

b) List four cations in this experiment that are not discarded down the drain:

c) List two anions in this experiment that are disposed of down the drain:

d) What is the approximate pH of non-toxic aqueous solutions that may bediscarded into the sewage system without damaging the environment? _____

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e) The narrow pH range allowed for solutions discarded by industry andchemistry laboratories is routinely exceeded in household wastes. Describethese common household chemicals as acidic or basic.Coca Cola __________________ household ammonia __________washing soda (Na2CO3) _________ lemon juice ________________

Environmental Questions: For a course to be approved as meeting theEnvironmental theme requirements it must contain a significant amount of materialrelevant to the environment as defined by the Liberal Education Taskforce. A fairlycomplete description of these requirements follows this experiment. The importantthing to notice is the breadth of the requirement and its intention that the course, inthis case chemistry, provide a specific background that is then used to look at largerissues involving social, cultural, legal, political, economic, religious and otherconsiderations. The implication is strong that because of the importantinterdependency of humans and the natural environment, understanding thiscomplex relationship and a concern for environmental degradation are goals for theenvironmental portion of this course. This gives a certain bias to the questions, buteach person must ultimately make his or her own decisions on environmental issues.Answering these questions and participating in the Poster Sessions will help you tomeet the Environmental theme goals for this course and to make informed decisionsin the future.

How green is your chemistry? Every experiment in this lab manual has been writtenwith the goal of producing zero or minimal amounts of hazardous chemical waste.The designing of experiments or industrial processes to eliminate or dramaticallyreduce the production of hazardous waste is called “green chemistry”. An exampleof green chemistry at the consumer level is the development of new dry cleaningprocesses that use water or liquid CO2 as the solvent rather than an organic liquid.Organic liquids require special handling and disposal and some, such as carbontetrachloride, CCl4, are hazardous to human health and to the environment. Severalgoals of “green chemistry” in industry include:

•developing alternative, low waste, methods to prepare chemical compounds•using less fuel by lowering the reaction temperature through the use of chemical

catalysts (ch. 12 in your text)•choosing the least toxic chemicals for industrial and consumer use whenever

several choices are available.

5. Working alone or with a partner (partner’s name _______________________ )go back through this experiment and specifically and clearly describe at least 4ways in which the experimental procedure is “green”, i.e., is designed to reduceor eliminate the amount of hazardous waste either generated or requiringspecial treatment at another facility