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Experiment 9: Double Replacement Reactions
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Chemical reactions is the main focus of chemistry because that's where the action is. That's when something gets made, changed, or destroyed. So a lot of emphasis is placed on chemical reactions.

Chemical reactions can be understood by these three areas of focus. A reaction is basically shuffling around the building blocks of chemistry (atoms/ions/molecules/electrons).  Force & energy is what decides if the reaction occurs and how fast. Mathematics helps you keep an inventory of all the starting and ending materials.

Types of Chemical Reactions
Synthesis (Combination)
Decomposition
Single Replacement/Displacement
Double Replacement/Displacement
Oxidation (Combustion)
Chemical reactions are often classified into the five categories listed on the left panel.  Lab 4 focused on physical and chemical changes.  The chemical changes included decomposition, oxidation, and double replacement reactions.  The previous lab (Lab 8) focused solely on single replacement reactions.   This lab will focus just on double replacement reactions.
   
Objectives

1. Review the general pattern for double replacement reactions

2. Predict if a reaction will occur based on a few simple rules.

3. Carry out several double-replacement reactions used for various applications

a. Tap Water Purification:  Removal of iron ions using a double replacement reaction

b. Drinking Water Purification:  Removal of copper ions using a double replacement reaction

c. Photography: Synthesis of several light sensitive silver halides using a double replacement reaction.

d. A double-replacement reaction because of an unstable product.

Principles and Applications of Double Replacement (displacement) reactions
double replacement pattern

General Pattern for Double Replacement Reactions:

Double Replacement Pattern: AB + CD → AD + CB
Two compounds exchange ions to form two new compounds.  In the generic compound "AB", the "A" element, which is usually a metal, exchanges places with another metal in compound "CD". To the right side of the arrow shows their final arrangement.

Note: The letters are just letters that represent any ion. The "B" is not boron and "C" is not carbon.

As a specific example, we are showing barium hydroxide, Ba(OH)2, reacting with calcium sulfate, CaSO4.   When the reaction occurs, the barium ion (Ba2+) will take the place of the calcium ion (Ca2+). 

 

Even though Ba(OH)2(aq) is written as if the barium and hydroxide are connected, the "(aq)" is saying that they are dissolved in water.  That means they are floating independent of each other.  The same thing is true for CaSO4(aq).  The calcium and sulfate ions are floating independent of each other.  When the two solutions of barium hydroxide and calcium sulfate are mixed, the ions are free to move around and interact with each other.   The positive ions will be attracted to the negative ions.   So the calcium ion (Ca2+) will be attracted to the negative hydroxide ions (OH-); however, water will pull them apart.  Barium ions (Ba2+) will be attracted to the sulfate ions, but this time water can't pull them apart because their attraction is too strong.  So Ba2+ will bond with the SO42- to form solid barium sulfate, which will settle to the bottom of the container.  In other words, these two solutions form a precipitate of barium sulfate. Roll mouse over the image to see the final arrangement.

Below is the full ionic equation:

Ba2+(aq) + 2OH-(aq)  + Ca2+(aq)+ SO42-(aq)→ BaSO4(s) + Ca2+(aq)+ 2OH-(aq) 

If we remove the ions that didn't react (the spectator ions), we get the below net reaction:

Ba2+(aq)  + SO42-(aq)→ BaSO4(s)

So the only reaction that happened was the positive barium ions bonded with the negative sulfate ions. Bonds that are formed because of opposite charges are called ionic bonds.  The only reason this happened is that barium sulfate is not soluble in water.  The reason it's not soluble is that the ionic bonds in barium sulfate are too strong for water to pull apart.  If you knew that barium sulfate was not soluble in water, you could predict that this reaction will happen.  Consulting a table that shows what is soluble or not, will tell you if these type of double-replacement reactions will occur.  You can also consult with the rules of solubility.

X-Ray Image Enhancement:  Below shows an example of using the barium hydroxide and calcium sulfate double replacement reaction.   Let's say you work for a rural hospital and you need a patient to some barium sulfate to help an x-ray show what's happening in the intestines.   Unfortunately, your rural hospital doesn't have any barium sulfate, but you do have some of the cold packs, which contain barium hydroxide.  You also have some plaster of Paris which is used to make casts.  That is made of calcium sulfate.  By mixing the barium hydroxide from the cold pack with some of the plaster, you will make barium sulfate.
cold pack, leg cast and xray of intestines
Toxic Metal Cleanup: Below is a reaction very similar to the one above, but the situation is different.  Let's say a pond is poisoned by barium chloride which could have come from waste water from various industries that use barium chloride.  Barium is classified as a heavy metal (like lead), which is toxic.  Above you learned that you can use sulfate ions to cause barium ions to precipitate out as solid barium sulfate.   Instead of using calcium sulfate (plaster of Paris) to supply sulfate ions, we are using Epsom salts (a magnesium sulfate salt) to supply the sulfate ions.   So adding a few bags of Epsom salts to the pond causes the toxic barium ions in the water to lock onto sulfate ions and settle to the bottom of the pond as barium sulfate.  Even if the barium sulfate gets stirred up and drank, it won't hurt anyone, because it won't dissolve in a person's or animal's digestive tract.
equation to use Epsom salts to get rid of barium

For single replacement reactions, the reaction took place if the metal was more active (higher on the list) than the metal ion that it was in contact with.   Double-replacement reactions take place for different reasons.  One reason is based on solubility.  In other words, whether or not certain positive ions form a solid (a precipitate) if they contact certain negative ions.   Knowing which positive ions will form solids with which negative ions can be learned from a SOLUBILITY TABLE or SOLUBILITY RULES.

If a compound forms a solid precipitate(is insoluble), it will form upon mixing two solutions, one having the positive ion and the other containing the negative ion.  Solid silver chloride, for example, precipitates on mixing any solution containing chloride ions with any other solution of silver ions.

SOLUBILITY TABLE
"aq" stands for aqueous meaning those two ions are soluble and stay in solution.  In other words, there is no reaction (no precipitate) when those ions come in contact with each other.
 
chloride
bromide
iodide
sulfide
hydroxide
carbonate
nitrate
sulfite
sulfate
phosphate
acetate
 
Cl-
Br-
I-
S2-
OH-
CO32-
NO3-
SO32-
SO42-
PO43-
C2H3O2-
H+
aq
aq
aq
solid
liquid1*
gas2*
aq
gas3* 
aq
aq
liquid
Na+
aq
aq
aq
aq
aq
aq
aq
aq
aq
aq
aq
K+
aq
aq
aq
aq
aq
aq
aq
aq
aq
aq
aq
NH4+
aq
aq
aq
aq
gas4*
aq
aq
aq
aq
aq
aq
Ag+
solid
solid
solid
solid
solid
solid
aq
solid
solid
solid
solid
Mg2+
aq
aq
aq
solid
solid
solid
aq
solid
aq
solid
aq
Ca2+
aq
aq
aq
solid
solid
solid
aq
solid
solid
solid
aq
Ba2+
aq
aq
aq
solid
aq
solid
aq
solid
solid
solid
aq
Fe2+
aq
aq
aq
solid
solid
solid
aq
solid
aq
solid
aq
Fe3+
aq
aq
solid
solid
solid
solid
aq
solid
aq
solid
solid
Co2+
aq
aq
aq
solid
solid
solid
aq
solid
aq
solid
aq
Ni2+
aq
aq
aq
solid
solid
solid
aq
solid
aq
solid
aq
Cu2+
aq
aq
solid
solid
solid
solid
aq
solid
aq
solid
aq
Zn2+
aq
aq
aq
solid
solid
solid
aq
solid
aq
solid
aq
Hg2+
aq
solid
solid
solid
solid
solid
aq
solid
solid
solid
aq
Cd2+
aq
aq
aq
solid
solid
solid
aq
solid
aq
solid
aq
Sn2+
aq
aq
solid
solid
solid
solid
aq
solid
aq
solid
aq
Pb2+
solid
solid
solid
solid
solid
solid
aq
solid
solid
solid
aq
Mn2+
aq
aq
aq
solid
solid
solid
aq
solid
aq
solid
aq
Al3+
aq
aq
aq
solid
solid
solid
aq
solid
aq
solid
aq

The solubility table above can also be described using the below solubility rules.

Solubility Rules:

1.   All compounds of Na+, K+, and NH4+ are soluble.
2.   All compounds of NO3- are soluble.
3.   All compounds of C2H3O2- are soluble except for silver or iron(III).
4.   All compounds of Cl-, Br-, and I- are soluble except  Ag+, Hg22+, and Pb2+.
5.   All compounds of SO42-  are soluble except Ca2+, Ag+, Pb2+, Sr2+, and Ba2+.
6.   All metal hydroxides (OH-) are insoluble except those of Groups IA. Group IIA metal hydroxides are moderately soluble.
7 .  All carbonates (CO32-), phosphates (PO43-), and sulfides (S2-) are insoluble except those of Group IA and ammonium (NH4+).
8.   All metal oxides (O2-) are insoluble except those of Group IA, that dissolve to form hydroxides.  Oxides of Group IIA are slightly soluble.

cigarette smoke

Using the solubility chart and solubility rules, we will see how we can make silver acetate.  Acetate is in the far right column.  As we come down we see that the silver ion will form a solid (a precipitate) with the acetate ion.   That means if we find a salt of silver that is soluble (aqueous) and a salt of acetate that is soluble, then we can mix them and the silver ions will bump into the acetate ions to form the insoluble (solid) silver acetate we want.  In the solubility table and #2 of the solubility rules, the only ion that keeps silver dissolved is the nitrate ion.  So we locate some silver nitrate (like what's in the kit).   Now we look for a soluble (aqueous) salt of acetate.  According to our table and #3 of the solubility rules, we can use about any positive ion other than iron(III), Fe3+.   For example, sodium acetate, potassium acetate, ammonium acetate, and many others would provide us with the acetate ions we need.  This the double-replacement reaction we want.  Picking sodium acetate as one source, we see that silver takes the place of sodium (Na) to make silver acetate.

NaC2H3O2(aq) + AgNO3(aq) → AgC2H3O2(s) + NaNO3(aq)

By the way, silver acetate was once used in small quantities (2.5 milligrams) in lozenges and chewing gum to deter smoking.  Apparently some ingredient in the cigarette smoke would react with the silver acetate to produce a compound that tasted very bad (a strong metallic taste). 

Another way a double-replacement reaction occurs is if the reaction produces an unstable compound that decomposes into a gas and water. 

The three most common unstable compounds which give gases are:

                           Carbonic acid:                           H2CO3    CO2(g)   + H2O(l)

                           Sulfurous acid:                          H2SO3   SO2(g)   + H2O(l)

                           Ammonium hydroxide:             NH4OH    NH3(g)   + H2O(l)

For example, if hydrochloric acid and sodium bicarbonate are mixed we get the below reaction. 
HCl(aq) + NaHCO3(aq) → H2CO3(l) + NaCl(aq)

You can see that the H in HCl exchanges places with Na in NaHCO3.  So this looks like a double-replacement reaction; however, the H2CO3 formed is not a solid that takes itself out of solution.  H2CO3 stays in the solution and the H in H2CO3 can easily come off, meaning that H and Na can rather easily go back to where they came from. So there would be no double-replacement.  However, H2CO3 is unstable and it will decompose into CO2 gas and H2O.  After CO2 gas escapes, it can not longer combine with water to reform H2CO3.  With H2CO3 decomposing, it can't give back the hydrogen atom that it took from HCl.  That makes the double-replacement reaction possible.

bubbles from  H2CO3

Here is the double-replacement reaction again, but with the follow-up decomposition reaction.

HCl(aq) + NaHCO3(aq) → H2CO3(l) + NaCl(aq)

                                            ↓ 
                                       H2CO3  ®   CO2(g)   + H2O(l)

Since H2CO3 decomposes into CO2 and H2O, you can see bubbles forming.  That's a sign that the double-replacement reaction is occurring.

H+ + OH- → H2O:   There is one other way a double-replacement reaction can happen.  It's when an acid (H+ ions) neutralizes hydroxide ions (OH-) to form water (H2O).    There are no precipitate or bubbles to show this occurred; however, heat is given off and the pH of both liquids move closer to a neutral pH.
neutralizing bases

Neutralizing Corrosive Chemicals:
Lye (sodium hydroxide) is used to make biodiesel and for unclogging a drain.  Let's say a jar of it fell and spilled all over the floor.  Pure sodium hydroxide is very corrosive and dangerous.  You sweep up what you can, but you need to neutralize the remaining powder that can't be swept up.   Using vinegar (dilute acetic acid) to cause a double-replacement reaction will neutralize the sodium hydroxide.

Notice the acidic hydrogen on acetic acid is shown in red.  It's called the "acidic hydrogen" because it is the hydrogen which comes off of acetic acid to form H+ ions.  That hydrogen ion replaces the sodium ion on sodium hydroxide.  It does that because the H+ ion will bond very strongly with the hydroxide (OH-) ion to form water (HOH = H2O).

Note:  Sometimes acetic acid is written as CH3COOH.   The acidic hydrogen is written last.  In that case, the H+ ion is "B" in the generic "AB" compound.  However, it still is a double-replacement reaction.

a. Lab 9 Experiment 1: Tap Water Treatment: Removal of iron ions
copper sulfate to dissolve
Later in this experiment you will be needing some dissolved copper(II) sulfate.  As you learned in the previous lab, copper(II) sulfate crystals take a while to dissolve.  Put a crystal or two of copper(II) sulfate into a test tube and add purified water.   Fill the test tube about half full of water.  Let it sit while you do the first experiment with iron(II) sulfate.
iron stained tub

Excess iron in tap water will stain sinks, tubs, showers, and clothes.  Certain salts of iron are soluble and some are not.   The iron salts that are soluble pass through filters and a stay in the tap water until they come in contact with the air.   Then they form the reddish-brown iron stains. 

Also, if the level or iron ions are too high, the taste of the water is affected.

One strategy in getting trapping the dissolved iron salts is to change them to iron salts that are not soluble.  That way solid iron particles can be filtered out.

iron sulfate in test tube

Your kit has some iron(II) sulfate.  This iron salt does dissolve in water and is one source of iron in tap water.  

Your kit also has some sodium carbonate (washing soda).  This will provide us with a source of carbonate ions (CO32-() which will combine with the iron ions to form the insoluble iron(II) carbonate.

FeSO4(aq) +  Na2CO3(aq) → FeCO3(s) +  Na2SO4(aq)

The above reaction shows the double-replacement going on.  Iron and sodium swap places.  The below reaction shows the total ionic equation, which better reflects what is going on.

Fe2+(aq) + SO42-(aq) + 2Na+(aq) + CO32-(aq) → FeCO3(s) + 2Na+(aq)+ SO42-(aq)

If we don't write the ions that don't react, we get the below simpler reaction, which is iron bumping into carbonate ions and forming solid iron(II) carbonate.

Fe2+(aq) + CO32-(aq) → FeCO3(s)

iron sulfate and sodium carbonate in test tubes

Use your microspatula to place a little of the iron(II) sulfate powder into a empty clean test tube.   Also, place a little sodium carbonate into another empty clean test tube.  See outer test tubes in image for approximate amounts.

Note:  The iron(II) ion Fe2+ is also referred to as "ferrous".  So iron(II) sulfate is also called ferrous sulfate. The iron(III) ion with plus 3 charge (Fe3+) is referred to as "ferric".  So iron(III) sulfate would also be called ferric sulfate. 

The "-ous" endings were were used for the ion with the lower positive charge.  The "-ic" ending was for the ion with the larger positive charge.  For example, Cu+ is indicated with "-ous" ending and Cu2+ had the "-ic" ending.  So copper(I) chloride is called cuprous chloride, and copper(II) chloride is called cupric chloride.

The use of the "-ous" and "-ic" endings was an earlier nomenclature rule compared to using Roman numerals to express the charge. However, that older nomenclature is still in widespread use.

 

 

add water to salts
Add water to the test tubes with the small amounts of iron(II) sulfate and sodium carbonate.  Fill the test tubes to about half full or to the height of the white plastic posts.
caps on test tubes
Place caps on the test tubes with the water and shake to dissolve.   The caps don't have to be these colored ones, but the clear plastic caps are fine.
transfer from sodium carbonate test tube to iron tube
Use a clean disposable plastic pipette to transfer some of the sodium carbonate solution to the test tube with the dissolved iron(II) sulfate.
iron carbonate precipitate

You should see a blue-green precipitate of iron(II) carbonate.  The carbonate salt of iron is not soluble so it forms solid particles in the solution.

Fe2+(aq) + CO32-(aq) → FeCO3(s)

A simple filter like a filter paper will not remove Fe2+ ions.  However, when those iron ions are bonded with carbonate ions to form solid iron(II) carbonate, those particles can be filtered out of the water using filter paper.

 

 

small amount of iron sulfate in test tube

The iron(II) carbonate precipitate above was formed with a rather concentrated solution of iron sulfate.  Iron levels won't be that high in tap water.  So let's try a more realistic level of iron that might be in tap water. 

Transfer just a small amount of the iron(II) sulfate powder to an empty test tube. See image for approximate amount.

Add water to these few granules of iron(II) sulfate to dissolve them.  Fill the test tube to about half full with purified water.

iron carbonate green color in test tube

Like before, transfer some of the sodium carbonate solution to the test tube that has just a small amount of dissolved iron(II) sulfate. 

For us, the solution went from clear to a light green. There also seems to be some small green particles in the solution.

a1) What did you see when you first added the sodium carbonate solution?

iron carbonate precipitates from weak and strong

Here we have both precipitates from the strong and weak solutions of iron(II) sulfate.

In a few minutes, the light green color of the weak solution has turned into dark green particles that are starting to settle out.   This is good news.  This means that even small amounts of iron ions can be turned into solid particles.  As particles, they can be trapped using standard filters.

Take a photo of your test tubes that show the reaction of carbonates with both the strong and weak iron(II) sulfate solutions.

a2) Instead of iron(II) sulfate (FeSO4) to make iron(II) carbonate, what other iron compound could have been used as a source of iron ions?  (consult the solubility table).

a3) Instead of sodium carbonate (Na2CO3) as the source of carbonate (CO32-) ions, what other carbonate compound could have been used? (consult the solubility table)

closeup of iron carbonates

Here is a close up of the precipitates.   Notice the reddish-orange color.  That's probably some of the unreacted iron(II) ions that are reacting with oxygen in the air to form iron oxides and possibly iron hydroxides.  These have the characteristic rust color that we associate with iron stains.

b. Lab 9 Experiment 2: Drinking Water Treatment: Removal of copper ions
copper sulfate to dissolve
Earlier you should have dissolved a crystal or two of copper(II) sulfate into a test tube.  If not, do that now.  Fill the test tube about half full of water.   Stir with glass rod to speed up dissolving or let it sit for 10 minutes or so.
pool with algae

Removal of toxic copper ions:

Algae are very hardy plants that are difficult to kill, but copper ions are commonly used to kill algae in pools and ponds.  

Excess iron ions in water are more of a nuisance; however, excess copper ions are toxic when ingested.  Anything over 0.002 grams of copper ions per liter of water is considered unsafe. 

This next experiment takes the same strategy that was used to remove iron ions as the way to take out toxic copper ions.

copper sulfate and sodium carbonate in test tubes

By now, the copper(II) sulfate crystals should be dissolved.  Add some more sodium carbonate to the test tube that already has some dissolved sodium carbonate.  Then add some more purified water to dissolve it.

Also transfer a small amount of the dissolved copper(II) sulfate solution to an empty test tube.  Add more water to dilute it.

Like before you will add some dissolved sodium carbonate to both strong and weak solution of dissolved copper(II) sulfate.   Here is the reaction you should see.  It's very similar to the double-replacement we did with iron(II) sulfate and sodium carbonate.

CuSO4(aq) +  Na2CO3(aq) → CuCO3(s) +  Na2SO4(aq)

The above reaction shows the double-replacement going on.  Copper and sodium swap places.  The below reaction shows the total ionic equation, which better reflects what is going on.

Cu2+(aq) + SO42-(aq) + 2Na+(aq) + CO32-(aq) → CuCO3(s) + 2Na+(aq)+ SO42-(aq)

If we leave out the ions that don't react, we get the below simpler reaction, which is copper bumping into carbonate ions and forming solid copper(II) carbonate.

Cu2+(aq) + CO32-(aq) → CuCO3(s)

Copper carbonate precipitates in test tubes

Add some of the sodium carbonate solution to both the strong and weak copper(II) sulfate solutions.  Look for any precipitation.    The strong solution created a lot of blue-green precipitation made of copper(II) carbonate.  Even the weak solution shows precipitation.  This is good news. That means these particles of copper(II) carbonate can be trapped with a standard filter.  So using this technique can remove the toxic copper ions from tap water.

b1) What did your precipitate of copper(II) carbonate look like?

Our next step is to use a paper filter to trap the particles of copper(II) carbonate.

If a filter wasn't available, allowing the precipitate to settle to the bottom and only using the water above the precipitate would be another way to get water that was free of copper ions. 

It seems that having some sodium carbonate might be a handy thing to have on a trip.  It can cause several toxic metals in the water to precipitate out.

Filter ready to trap copper carbonate

Get the plastic funnel from the kit and set it in an empty test tube.   Also, get a sheet of filter paper. 

First fold the filter paper in half.  

folded filter paper

Then you fold the filter paper again, but not in half.  See image below. 

final folded filter paper

put water on filter paper
Pull open the filter paper so one of the folds makes a pouch.  Place it into the funnel.  Add a little purified water will help hold the filter paper in place.
filtering copper carbonate

Pour the contents of the test tube that had the sodium carbonate added to the strong copper(II) sulfate solution.  That will have the most copper(II) carbonate precipitate.

Take a photo of your filter setup for trapping copper(II) carbonate.

b2) Instead of copper(II) sulfate (CuSO4) as the source of copper ions to make copper(II) carbonate, what other copper compound could have been used as a source of copper ions?  (consult the solubility table).

b3) Instead of sodium carbonate (Na2CO3) as the source of carbonate (CO32-) ions, what other carbonate compound could have been used? (consult the solubility table)

 

copper carbonate filtered out

As the solution passes through the filter, it appears that the filter traps the copper(II) carbonate precipitate.  The solution that is collected is clear.    Warning: If there wasn't enough sodium carbonate added to react with all of the copper(II) sulfate, there will still be copper ions in this solution.  Only by adding a little more sodium carbonate solution can one be sure.

b4)  Go ahead and add a little more sodium carbonate solution to this liquid that passed through the filter.  Did more blue-green precipitate appear?

 

 

c. Lab 9 Experiment 2:  Silver Salts for Photography
First portrait using photography

In lab 4: Physical and Chemical Changes, you learned about silver nitrate and its reaction with sodium chloride.  At that time, we didn't mention it was a double-replacement reaction.  For this experiment, we will use silver nitrate to produce double-replacement reactions with three different halides (chloride, bromide, and iodide).    All of these have been used as the light sensitive compounds that are used in photographs.

The photo, as you may recall, is the first person ever to be photographed.  It was back in 1839.  The photo was that of the photographer, Robert Cornelius.

test tubes with sodium chloride, potassium bromide, and potassium iodide

Locate the test tubes with sodium chloride, potassium chloride, and potassium iodide.  Also, get 3 clean test tubes.  If you need to, wash out the ones used earlier in this lab.  Use purified water to rinse them, but it is OK if there's a little water left in the test tubes.  They don't have to be dry.

Add a small amount of each of these salts to an empty test tube.  See photo for amounts that are to be used.

Add purified water to each of these to dissolve the salts.   Add enough water to bring the level up to the top of the white plastic posts.  That's about 4 mL.

After adding water, shake or swirl the water to dissolve the salts.

 

adding silver nitrate to test tubes with halide salts

Add about 7 drops of silver nitrate to each of the salt solutions.

Here is the double-replacement reaction we expect after adding silver nitrate to dissolved sodium chloride. Sodium takes silver's place and silver takes sodium's place.

NaCl(aq) + AgNO3(aq) → AgCl(s) + NaNO3(aq)

Because these are dissolved in water, the ions are not together but floating around separately.  That is more apparent with a full ionic equation shown below. 

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

silver halides

If we leave out the ions that don't react, we get the below simpler reaction, which is silver ions bumping into sodium ions and forming solid silver chloride.

Ag+(aq)+ Cl-(aq))→ AgCl(s)

Doing the same for potassium bromide and potassium iodide.

Ag+(aq)+ Br-(aq))→ AgBr(s)

Ag+(aq)+I-(aq))→ AgI(s)

Take a photo of your silver halide precipitates.

silver halides

There is a bit of a color difference between these silver halides.

c1) What color would you call the silver chloride precipitate?

c2) What color would you call the silver bromide precipitate?

c3) What color would you call the silver iodide precipitate?

If exposed to the sunlight (or UV light), all of these will become black in color due to the silver ion breaking away from the silver salt to form small particles of silver metal.

c4) Instead of sodium chloride (NaCl) as a source of chloride ions to make silver chloride, what other chloride compound could have been used as a source of chloride ions?  (consult the solubility table).

c5) Instead of potassium bromide (KBr) as the source of bromide (Br-) ions to make silver bromide, what other bromide compound could have been used? (consult the solubility table)

c6) Instead of potassium iodide (KI) as the source of iodide (I-) ions to make silver iodide, what other iodide compound could have been used? (consult the solubility table)

d. Lab 9 Experiment 4: Double-Replacement reaction because of unstable product
HCl and bicarbonate

In the previous experiments of this lab, the double-replacement occurred because one of the products formed a precipitate, which prevents the reaction from reversing.   Earlier it was mentioned that another way a double-replacement reaction would proceed is if one of the products decomposes into a gas and water.  The decomposition prevents the reaction from reversing.

Locate the 0.1M HCl solution and the test tube with sodium bicarbonate (NaHCO3).  Transfer a small amount of the sodium bicarbonate to an empty test tube.  See image for amount to use.

This time don't add any water to the sodium bicarbonate that you transferred.  There is already water in the 0.1 M HCl.

 

Adding HCl to sodium bicarbonate

Add 0.1 M HCl to the sodium bicarbonate powder.

Bring the level up to about halfway up the white posts.

 

bubbles from bicarbonate

Here is the double-replacement reaction with the follow-up decomposition reaction.

HCl(aq) + NaHCO3(aq) → H2CO3(l) + NaCl(aq)

                                            ↓ 
                                       H2CO3  ®   CO2(g)   + H2O(l)

Since H2CO3 decomposes into CO2 and H2O, you can see CO2 bubbles forming.  That's a sign that the double-replacement reaction is occurring.

d1) Instead of hydrochloric acid (HCl) as a source of hydrogen ions (H+)to make carbonic acid (H2CO3), what other acid could have been used as a source of hydrogen ions?  (consult the solubility table).

Summary of Data to Report:

If you wish, you can copy the below summary into your email (or Word document) and type your answers after the descriptions.  The required photos can either be attached to the email or inserted in the Word document if going that route.  Try to keep each image under 2 megabytes.   If you are in section 20598, which has a start date of Aug. 19th, email your lab report to Quinn Thacker at QRT2004@yahoo.com.  If you are in section 20930 which has a start date of Sept. 9th, then email your lab report to Loree Cantrell-Briggs at lor2060912@phoenixcollege.edu.  Be sure to title the email "Lab 9".

a. Lab 9 Experiment 1: Tap Water Treatment: Removal of iron ions

a1) What did you see when you first added the sodium carbonate solution?

Attach a photo of your test tubes that show the reaction of sodium carbonate with both the strong and weak iron(II) sulfate solutions.

a2) Instead of iron(II) sulfate (FeSO4) to make iron(II) carbonate, what other iron compound could have been used as a source of iron ions?  (consult the solubility table).

a3) Instead of sodium carbonate (Na2CO3) as the source of carbonate (CO32-) ions, what other carbonate compound could have been used? (consult the solubility table)

b. Lab 9 Experiment 2: Drinking Water Treatment: Removal of copper ions

b1) What did your precipitate of copper(II) carbonate look like?

Attach a photo of your filter setup for trapping copper(II) carbonate.

b2) Instead of copper(II) sulfate (CuSO4) as the source of copper ions to make copper(II) carbonate, what other copper compound could have been used as a source of copper ions?  (consult the solubility table).

b3) Instead of sodium carbonate (Na2CO3) as the source of carbonate (CO32-) ions, what other carbonate compound could have been used? (consult the solubility table)

b4)  Go ahead and add a little more sodium carbonate solution to this liquid that passed through the filter.  Did more blue-green precipitate appear?

c. Lab 9 Experiment 2: Silver Salts for Photography

Attach a photo of your silver halide precipitates.

c1) What color would you call the silver chloride precipitate?

c2) What color would you call the silver bromide precipitate?

c3) What color would you call the silver iodide precipitate?

c4) Instead of sodium chloride (NaCl) as a source of chloride ions to make silver chloride, what other chloride compound could have been used as a source of chloride ions?  (consult the solubility table).

c5) Instead of potassium bromide (KBr) as the source of bromide (Br-) ions to make silver bromide, what other bromide compound could have been used? (consult the solubility table)

c6) Instead of potassium iodide (KI) as the source of iodide (I-) ions to make silver iodide, what other iodide compound could have been used? (consult the solubility table)

d. Lab 9 Experiment 4: Double-Replacement reaction because of unstable product

d1) Instead of hydrochloric acid (HCl) as a source of hydrogen ions (H+)to make carbonic acid (H2CO3), what other acid could have been used as a source of hydrogen ions? (consult the solubility table).

Post-Lab Questions:
Post-Lab questions and problems are on the Sapling Learning website.    http://www2.saplinglearning.com/

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