UNIT 2 1

HONORS CHEMISTRY

HARVARD-WESTLAKE

UNIT 2

Writing and Balancing

Chemical Reactions

A small piece of sodium which lived in a test tube fell in love with a Bunsen burner.

"Oh Bunsen, my flame. I melt whenever I see you . . .", the sodium pined.

"It's just a phase you're going through", replied the Bunsen burner.

UNIT 2 2

UNIT 2 3

Making sense out of Chemical Reactions Wait! Don't throw this away in disgust! It really is possible to do it! Read, think and learn: We have arbitrarily (and conveniently) divided inorganic chemical reactions into THREE large categories. The bottom line is that you need to learn how to recognize, write and/or complete, and balance these reactions. They are the "stuff" of which chemistry is made. Precipitation reactions How to recognize: the ones you will see in this course involve TWO compounds as possible reactants. e.g. sodium sulfate + barium chloride ?????? Na2SO4 + BaCl2 ?????? is possibly a precipitation reaction. _____________________________________________________ iron + hydrochloric acid ???????? Fe + HCl ????????? is definitely not. How to complete: try switching the ions and check products against the solubility rules; if at least one product is insoluble, you have a reaction; if both possible products are soluble, you have a dud (i.e., no reaction) e.g. Na2SO4 + BaCl2 NaCl + BaSO4 (not balanced!!!) according to the solubility rules, BaSO4 is insoluble. So there is a reaction. _____________________________________________________

NaNO3 + BaCl2 Ba(NO3)2 + NaCl (not balanced!!!) according to the solubility rules, both possible products are soluble. NO reaction will occur. Such a problem is best written (in molecular form) as: NaNO3 + BaCl2 N.R. How to balance: make sure you have the correct formulas first by checking the charges of the ions. Then AND ONLY THEN use coefficients in front of the formulas to get the same number

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of atoms on both sides. DO NOT ALTER FORMULAS TO FORCE A REACTION TO BALANCE. e.g. Na2SO4 + BaCl2 2 NaCl + BaSO4 How to write the net-ionic version: re-write the reaction, breaking up all soluble substances into ions, cancel spectators and then check balance. e.g. 2 Na+ + SO4

2- + Ba2+ + 2 Cl- 2 Na+ + 2 Cl- + BaSO4 --OR-- SO4

2- + Ba2+ BaSO4 Acid Reactions How to recognize: at this point, reactants must include one of the five common acids you are supposed to know and a hydroxide compound (a base) or a metal oxide or a metal carbonate. These are all compounds, so be careful not to confuse this with a precipitation reaction e.g. hydrochloric acid + sodium hydroxide ????? HCl + NaOH ????? is an acid-base reaction (neutralization) _____________________________________________________ HNO3 + NaCl ?????? is not, since there is no base present. How to complete: all common acid-base reactions you will be seeing at this point produce water and an ionic compound or salt. The salt may be soluble or insoluble. e.g. HCl + NaOH H2O + NaCl note that this result is obtained the same way as the products for a precipitation reaction--switch the ions. This example is already balanced. How to balance: do this in the same way as you do precipitation reactions How to write the net-ionic version: follow the same procedure as for precipitation reactions, bearing in mind that all acids and bases are electrolytes and water is not. Electrolytes break up into ions in solution. e.g. H+ + Cl- + Na+ + OH- H2O + Na+ + Cl- --OR- H+ + OH- H2O

UNIT 2 5

Other sequences: Remember, acids react with metal oxides to give water and a salt: e.g. HCl + MgO MgCl2 + H2O (not balanced) 2 HCl + MgO MgCl2 + H2O (balanced molecular) Since metal oxides are solids, the net-ionic equation must show this: 2 H+ + 2 Cl- + MgO Mg2+ + 2 Cl- + H2O (total ionic--not balanced) --OR-- 2 H+ + MgO Mg2+ + H2O (balanced net-ionic) For acids reacting with metal carbonates, the script is similar, with the addition of carbon dioxide as a third product. Details vary depending on whether the carbonate is soluble (like Li2CO3) or not. e.g. H2SO4 + Li2CO3 Li2SO4 + H2O + CO2 (balanced molecular) 2 H+ + SO4

2- + 2 Li+ + CO32- 2 Li+ + SO4

2- + H2O + CO2 2 H+ + CO3

2- H2O + CO2 (balanced net-ionic) Redox Reactions How to recognize: if it doesn't look like precipitation or acid, it is possibly redox of some kind. There are several possibilities: 1) synthesis from elements Na + Cl2 ??????? 2) decomposition of a simple compound HgO ?????? 3) displacement a. metals with acids Zn + HCl ????? b. metals with metal compounds Zn + Cu(NO3)2 ??????

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c. halogens with halide compounds Cl2 + KI ????

Note that all of these follow a simple pattern: they involve an element and a compound 4) OTHER (ugh!) (this includes anything else) How to complete: 1 & 2) Only very obvious synthesis and decomposition reactions will be given or else very generous hints will also be given. e.g. Na + Cl2 NaCl (not balanced) HgO Hg + O2 (not balanced)

3) You must check the metal activity series or the halogen family before deciding if a reaction will occur. Remember, in each case, the element which is a possible reactant must be higher in the list than the part of the compound it is supposed to displace. How do you know what will be displaced? "Like displaces like" (i.e., metals displace metals, positive displaces positive, etc.) e.g. Zn + HCl ZnCl2 + H2 (not balanced) This works since zinc is higher than hydrogen in the activity series. Note that in compounds they are both positive. ______________________________________________________ Cu + HCl N.R. This will not work since copper is below hydrogen in the activity series. ______________________________________________________ Cl2 + KI KCl + I2 (not balanced) This works since chlorine is above iodine in the halogen family. ______________________________________________________ I2 + KCl N.R. This will not work since iodine is below chlorine in the halogen column. 4) OTHER are so complicated that we will have to give you the products so you needn't worry about figuring them out by yourself.

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How to balance: this is done in the usual way for all redox reactions except OTHER. See the end of this handout for examples. How to write the net-ionic version: follow the usual procedure, keeping in mind that elements (like halogens) are not electrolytes and should not be broken up. Be sure to check that the charge is balanced as well as the atoms. e.g. Zn + 2 H+ + 2 Cl- H2 + Zn2+ + 2 Cl- Cl2 + K+ + 2 I- I2 + 2 Cl- + K+ Balancing non-trivial redox reactions Most of these reactions occur between ions in solution. The solution is usually either acidic or basic. The good news is that if you follow the method outlined (same one as in your text) you will end up with a balanced net-ionic reaction. You can decide for yourself what the bad news is. example: permanganate ions react with Cl- in acidic solution to produce Mn2+ and Cl2 1. write "skeleton" reaction: MnO4

- + Cl- Mn2+ + Cl2 (H+) fine print: the H+ tells you the solution is acidic 2. break up the reaction into halves, each containing an element that is changing, and balance those elements if needed: MnO4

- Mn2+ 2 Cl- Cl2 3. Since these reactions occur in solution, use H2O to balance oxygen atoms if needed: MnO4

- Mn2+ + 4 H2O 2 Cl- Cl2 4. Use H+ to balance hydrogens if needed: 8 H+ + MnO4

- Mn2+ + 4 H2O 2 Cl- Cl2

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5. Use electrons to balance charges: 5 e- + 8 H+ + MnO4

- ® Mn2+ + 4 H2O 2 Cl- Cl2 + 2 e- 6. Multiply half-reactions to equalize electrons (find LCM): 2 ( 5 e- + 8 H+ + MnO4

- Mn2+ + 4 H2O) 5 (2 Cl- Cl2 + 2 e- ) 7. Add half-reactions together, simplifying: 16 H+ + 2 MnO4

- + 10 Cl- 2 Mn2+ + 8 H2O + 5 Cl2 Whew. Question: What do I do if the reaction takes place in basic solution? Answer: exactly the same thing and then add to each side as many OH- as there are H+. Combine H+ and OH- on the same side to make H2O and simplify. Altering the previous problem (just for make-believe): 16 OH- + 16 H+ + 2 MnO4

- + 10 Cl- 2 Mn2+ + 8 H2O + 5 Cl2 + 16 OH- 16 H2O + 2 MnO4

- + 10 Cl- 2 Mn2+ + 8 H2O + 5 Cl2 + 16 OH- 8 H2O + 2 MnO4

- + 10 Cl- 2 Mn2+ + 5 Cl2 + 16 OH-

UNIT 2 9

Precipitation Reactions

Consider the following reaction:

Sodium phosphate + barium chloride ?

Step 1: to figure out the products, simply change the “dance partners” and remember to keep the name of the metal (or ammonium) ion first. In this case the full reaction would be:

Sodium phosphate + barium chloride sodium chloride + barium phosphate

Step 2: Use the ion charges you memorized in Unit 1 to write the formulas for the compounds. USE CORRECT SUBSCRIPTS!!!!!!!!!!!!! (Remember there is NO SUCH THING as: Na2Cl2 or Pb(II)CO3…… ionic compounds have the lowest values of subscripts that allow for charge balancing, and NO roman numerals are used in formulas EVER). Our example becomes:

Na3PO4 + BaCl2 → NaCl + Ba3(PO4)2

Step 3: Use coefficients to balance the reaction. DO NOT CHANGE SUBSCRIPTS!!!!!!

2Na3PO4 + 3BaCl2 → 6NaCl + Ba3(PO4)2

**This is the balanced molecular equation**

Step 4: Use the solubility rules to determine which of the products are insoluble. Then split all SOLUBLE compounds into ions, and DO NOT split insoluble compounds. In our example barium phosphate is insoluble (rule 8). Remember that ionic compounds will split into individual ions (just as you memorized them). Subscripts will be multiplied by the coefficients to find the total number of each ion. Na3

+, Na2+, Cl2- DO NOT EXIST. PLEASE go back and

memorize the ions and their charges from Unit 1.

6Na+ + 2PO43- + 3Ba2+ + 6Cl- → 6Na+ + 6Cl- + Ba3(PO4)2

**This is the total ionic equation**

Step 5: Cancel out all ions that exist in exactly the same number and charge on both sides of the reaction. Re-write the equation without these “spectator” ions.

6Na+ + 2PO43- + 3Ba2+ + 6Cl- → 6Na+ + 6Cl- + Ba3(PO4)2

2PO43- + 3Ba2+ → Ba3(PO4)2

**This is the net ionic equation**

Double-check to make sure atoms and overall charge are balanced (equal on both sides of reaction). You are all done!!

UNIT 2 10

Reaction Practice

All reactions occur in aqueous solution unless otherwise noted. 1. calcium bromide + silver nitrate 2. silver metal + copper(II) chloride 3. potassium phosphate + calcium nitrate 4. solid strontium oxide + nitric acid 5. sodium sulfate + ammonium chloride 6. potassium hydroxide + hydrochloric acid 7. butane gas (C4H10) + oxygen 8. nickel metal + zinc chloride 9. magnesium nitrate + copper metal 10. lithium iodide + chlorine 11. barium metal + oxygen 12. sodium metal + water 13. sulfuric acid + barium nitrate 14. calcium nitrate + potassium hydroxide 15. potassium carbonate + hydrochloric acid 16. zinc metal + silver nitrate 17. bromine + sodium chloride 18. silver nitrate + copper metal 19. barium nitrate + sulfuric acid 20. nitrogen gas + hydrogen gas 21. nitrate ions + zinc metal in acidic solution; products include Zn2+ and NH4

+ 22. sulfate ions + copper metal in acidic solution; products include Cu2+ and sulfur dioxide gas 23. CrO2

- + ClO- in basic solution; products include chloride ion and chromate ion 24. permanganate ion + bromide ion in acidic solution; products include bromine and MnO2 25. H2S gas + iodine in acidic solution; products include sulfur and iodide ions  

 

 

 

 

 

 

 

 

UNIT 2 11

Reaction Practice Continued

Write balanced reactions (net-ionic where appropriate) for the processes below which are expected to happen. Reactions that do not occur in aqueous solution are maked with *. 1. copper(II) chloride + sodium nitrate 2. solid calcium carbonate + hydrochloric acid 3. zinc metal + sulfuric acid 4. nitric acid + potassium hydroxide 5. strontium chloride + sodium carbonate 6. lithium metal + water *7. bromine + potassium metal

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8. sodium carbonate + hydrochloric acid 9. sodium acetate + potassium sulfate *10. butane gas (C4H10) + oxygen 11. nickel metal + zinc chloride 12. barium nitrate + sulfuric acid *13. iron metal + oxygen [product contains iron(III) ] 14. solid magnesium oxide + hydrochloric acid            

UNIT 2 13

Non-Trivial Redox Examples 15. sulfate ions + copper metal in acidic solution; products include Cu2+ and sulfur dioxide gas 16. dihydrogen sulfide gas + iodine in acidic solution; products include sulfur and iodide ions 17. CrO2

- + ClO- in acidic solution; products include chloride ion and chromate ion  

 

 

 

 

 

 

 

 

 

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And even more reaction practice All of the reactions below occur in aqueous solution unless marked by *. 1. aluminum + hydrochloric acid 2. silver nitrate + dihydrogen sulfide gas 3. iron(III) chloride + potassium hydroxide 4. solid aluminum hydroxide + hydrochloric acid 5. chlorine gas + sodium bromide 6. hydrochloric acid + lead(II) nitrate 7. potassium bromide + mercury(I) nitrate 8. potassium hydroxide + sulfuric acid 9. barium chloride + potassium carbonate 10. silver nitrate + sodium phosphate 11. bromine gas + potassium iodide 12. zinc metal + sulfuric acid 13. barium chloride + aluminum sulfate 14. iron(III) chloride + ammonium sulfide 15. aluminum metal + mercury(II) nitrate 16. lead(II) nitrate + sodium carbonate 17. calcium metal + water 18. iron metal + sulfuric acid (one product contains iron(II)) 19. potassium phosphate + magnesium nitrate 20. potassium sulfide + copper(II) chloride 21. potassium metal + water *22. magnesium + oxygen 23. zinc metal + lead(II) nitrate 24. silver nitrate + magnesium bromide 25. aluminum + copper(II) sulfate 26. solid barium hydroxide + sulfuric acid non-trivial redox: (all in acidic solution) Cr2O72- + Sn2+ Cr3+ + Sn4+

H2S + I2 S + I-

MnO4- + Br- Br2 + MnO2 Ag+ + AsH3 Ag + H3AsO3

UNIT 2 15

LAB: Precipitation Reactions Reactions occurring in solution may produce substances which are insoluble in the solution and thus eventually settle to the bottom or "precipitate" out. This is of interest to chemists because so much of what we do occurs in some type of solution, generally in water. Typically precipitation reactions occur in aqueous solution between ions (i.e., the reactants are electrolytes). Often only one pair of ions actually react or precipitate while the other pair remain in solution unchanged. These ions which remain unchanged during the reaction are called spectators and can be omitted from the final balanced equation written in net-ionic form. As you will see... Although solubility is often related to the position of an element in the periodic table, there are enough exceptions to make for an interesting (or annoying, depending on your point of view) set of rules that describe which combinations are insoluble. One object of this experiment is to arrive at a partial set of rules that describes which combinations of ions are insoluble in water solution. Preparing to experiment Attached to this sheet you will find a grid in which to fill in your observations of what happens when you mix pairs of solutions. A similar paper is waiting for you in the lab laminated in plastic. All of the reactions can be performed with drop quantities of each solution on top of the plastic sheet. After you have recorded your results, carefully lift the plastic sheet and tip the drops into the sink. Rinse the plastic gently and carefully with water and dry with paper towels. Try not to crinkle it. Waste not, want not. The grid is provided for your use as an observation table. Technique What may at first seem obvious is often not [write that down somewhere!]. To complete this experiment in a single period you need to have some kind of organized approach to putting so many drops onto the sheet--especially since you will probably be sharing a set of solutions with another student. Your instructors have given this a GREAT deal of thought and we recommend the following: a. Pick up a solution and place drops of it in each square where it will be used. Remember that simply putting drops in a horizontal row or vertical column will NOT generally accomplish this since most substances are in both rows and columns. For example, if you look at the grid you will see that Al(NO3)3 (about half-way down the left side) turns up when you reach the end of the row. Most substances "turn a corner" like this. b. Mark off the solutions you have used with a small check mark on your data sheet. Most solutions look alike. c. Be careful not to touch the second solution dropper to the first drop of solution that has already been placed (why?) d. Do not add Na2S until after you have recorded all the results for other mixtures. It has an unpleasant smell and will also change the appearance of all your mixtures!

UNIT 2 16

Analysis 1. So much data and so little time.... If you inspect your observation grid carefully, you should see that there is one compound on the grid which did not react at all. This was NH4NO3. Since no combination of ions reacted with either NH4

+ or NO3-, we might say:

1) all compounds containing NH4

+ are soluble 2) all compounds containing NO3

- are soluble These two rules should make it possible for you to figure out the rest. To help you do so, fill in the grid below. This is done by looking for compounds which contain the anion listed in the left column and placing a check mark in the column under a cation which when added caused a precipitate to form [note that the nitrate and ammonium sections have been "grayed-out"--no reactions occur there]. The Chloride ion is done for you as an example.

Hg22+ Ba2+ Na+ Ca2+ K+ Mg2+ Cu2+ NH4

+ Al3+ Ag+ Pb2+ Fe3+

NO3-

PO43-

S2-

CO32-

Br-

SO42-

Cl-

OH-

2. Now select each anion represented in the experiment and write a rule for its solubility with the various cations you used. Write the rule in the most economical way. In other words, if there are more ions which are soluble with the anion, list those with which it is insoluble. Below is a reasoned example to get you started:

The compound KCl contains Cl- ion and forms precipitates with only three compounds: AgNO3, Pb(NO3)2 and Hg2(NO3)2. In each case one of the possible products is KNO3. But we already know that all NO3

- compounds are soluble. So it must be the chloride compounds that are precipitates. Thus: 3) Cl- is insoluble with Ag+, Pb2+ and Hg2

2+ Continue in this way to write rules for the remaining anions (it may be helpful to proceed in a similar fashion, tackling those that show the fewest reactions first before trying the more complicated cases) 3. Many of the solutions you used in the experiment contain the anion nitrate (NO3

-) and many also contain cations from Group I (e.g., Na+, K+). When your instructor prepares water solutions for experiments, why is it a good idea to choose compounds containing one of these ions? 4. Write molecular and net-ionic balanced equations for any 10 of the reactions which resulted in precipitate.

UNIT 2 17

Ba(NO3)2 NaOH Ca(NO3)2 KCl Na2SO4 Mg(NO3)2 Cu(NO3)2 NaBr NH4NO3 Al(NO3)3 AgNO3 Pb(NO3)2 Na2CO3 Na2S Na3PO4 Fe(NO3)3

Hg2(NO3)2

Fe(NO3)3

Na3PO4

Na2S

Na2CO3

Pb(NO3)2

AgNO3

Al(NO3)3

NH4NO3

NaBr

Cu(NO3)2

Mg(NO3)2

Na2SO4

KCl

Ca(NO3)2

NaOH

UNIT 2 18

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

UNIT 2 19

Name:___________________________ Per.:____ Date:_______________

Reactions of Acids Purpose: to observe reactions of acids with a variety of compounds and some metals and use the observations to write reactions for what has happened Method: I will test the behavior of three indicators (bromthymol blue, phenolphthalein and litmus) with acid, base, boiled distilled water, and a 1:1 mixture of acid and base. Then, using one of the liquid indicators to monitor the progress, I will add two different acids (HCl and H2SO4) to NaOH, drop by drop, to see if the drop ratio for neutralization is related to the balanced equation that represents the reaction. Finally, I will observe the reaction of various solids and metals with a variety of more concentrated acids in order to relate the behaviors to the written reactions. I will use my observations to write balanced equations for the reactions observed and to determine the significance of the drop ratios in the acid-base titrations. Observations/Data:  

indicator

HCl

NaOH

boiled H2O

1:1 HCl/NaOH

litmus

bromthymol blue

phenolphthalein

 

 

 

 

UNIT 2 20

drops HCl drops NaOH drops H2SO4 drops NaOH

 

Indicator used: __________________________ 

 

 

Zn Cu Fe Al CaO CaCO3

HCl

H2SO4

HNO3

 

UNIT 2 21

LAB: Reactions of Acids Acids are electrolytes, i.e., their solutions contain ions. The presence of hydrogen ions, H+ (sometimes written as hydronium ions, H3O+) indicates an acid. This is the simplest possible definition of acids that has anything to do with the actual chemistry behind their behavior. Your text gives what are called "operational definitions" for acids which simply describe some of their more obvious properties. Acids react with many elements (principally metals) and many types of compounds--too many to look at in a single experiment. In this exploration you will observe the following kinds of acid reactions:

with bases with metal oxides with metal carbonates with metals

Bases are also electrolytes which can generally be recognized the presence of hydroxide ion (OH-) in their formulas (ammonia, NH3, is an important exception). They react with acids to produce water and another type of electrolyte compound called a salt. This kind of reaction is called neutralization. Neutralization reactions can be followed in a variety of ways. Certain dyes are sensitive to the amount of acid or base in a solution and will change color as the amount of either changes. These dyes are known as indicators. Also, a quantity called the pH changes during a neutralization reaction. We will have more to say about pH later in the year. In this experiment you will use indicators to investigate the neutralization process. Metal oxides and carbonates are compounds which exhibit some basic behaviors in water solution. So it should not be surprising that they will also react with acids. Metal oxides may be considered as "bases without water" or basic anhydrides. For example, if H2O is removed from the formula Ca(OH)2, the result is CaO. Such compounds react with acids in the same way as bases, producing water and a salt. Metal carbonates are the products of the reaction of metal oxides with carbon dioxide, e.g.,

CaO + CO2 CaCO3 When an acid comes in contact with a metal carbonate, water and a salt are produced along with carbon dioxide gas. The behavior of these two types of compounds with acids is important because of the use of both in building and sculptural materials around the world. The increasing acidity of rain in many developed areas poses a problem for these materials as you will see. Finally, many metals react with acids. This is actually another type of chemical reaction known as "redox" (short for oxidation and reduction) which we will look at in more detail in the next experiment. For most metals the products of the reaction are hydrogen gas and a salt. This is a classic "displacement" reaction. However, the noble metals (Cu, Ag, etc.) are not attacked by all acids and when they do react, the products typically do not include hydrogen gas. Just as acid rain is a concern with respect to oxide and carbonate based building materials, ornamental and structural metals are affected as well. Preparing to experiment You will be provided with the following materials: 1. phenolphthalein indicator 2. bromthymol blue indicator 3. neutral litmus paper 4. solutions of hydrochloric acid, sulfuric acid, sodium hydroxide, and a mixture of equal parts hydrochloric acid and sodium hydroxide [these solutions are of equal concentration, i.e., one drop of each contains the same amount of HCl, H2SO4 or NaOH in moles]

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5. more concentrated solutions of HCl, H2SO4, and HNO3 6. samples of the following metals: Zn, Cu, Fe, Al 7. samples of CaO, CaCO3 (marble chips) 8. boiled distilled water 9. 24 well microplate 10. 24/96 well combination microplate Design an experiment to determine the color of each indicator (and litmus paper) when exposed to an acid, a base, boiled distilled water, and a mixture of equal volumes of acid and base. [hint: use 1-2 drops of each substance and one drop of indicator in one of the small wells on the plate--you only need to try one of the acids] Design an experiment to predict the ratio between HCl and NaOH in a reaction and also between H2SO4 and NaOH. [hints: both liquid indicators change color at approximately "neutral"; start each experiment with about 50 drops of NaOH in one of the large wells on the plate] Design an experiment to test the reactions of the three more concentrated acid solutions with the metals, oxide and carbonate sample [hint: use the 24-well plate---do the nitric acid last and in the fume hood--you may leave your plate there when you are finished observing] Pre-lab take-home quiz Answer these questions on a separate piece of paper to be turned in on the day you do this experiment. 1. Describe the difference in products when HCl reacts with Mg metal, MgO and MgCO3. 2. One solution contains H2SO4 and another contains HNO3. If each has equal moles of the respective acid present, which solution contains more H+ ions? 3. Acids are described as strong or weak depending on whether they are strong or weak electrolytes, i.e., whether they break up completely into ions or only partly. HCl is a strong acid. Nitrous acid (HNO2) is weak. In solutions with equal concentrations of these acids, which would have more free H+ ions present? ____________________________________________________ Technique 1. Indicator papers Many indicators can be used in a more "portable" form by applying the dye to paper and allowing it to dry. The color changes in acid or base solution remain the same. To correctly use an indicator strip you should place a drop of the solution to be tested (either with a dropper or the end of a stirring rod) on the paper. DO NOT dip the paper into the solution. 2. Using boiled water in acid/base experiments We generally use distilled water in experiments because it has been purified (by boiling the water and collecting the condensed steam) so that the majority of ions have been removed. However, distilled water stored in plastic bottles and allowed to stand for any period of time gradually collects CO2 from the atmosphere, just as bodies of fresh and salt water do in the environment outside the lab. The following reactions then take place:

H2O + CO2 H2CO3

H2CO3 H+ + HCO3-

UNIT 2 23

The presence of H+ ions (even in small amounts) makes the water slightly acidic. This may result in spurious results. Boiling the distilled water drives out the CO2 (most gases are less soluble in liquids at higher temperatures--in contrast to most solids) and ensures better results. Be sure not to leave the boiled water containers uncapped for any long period of time. 3. Titration When solutions containing substances that react are compared with regard to concentration, the process is called titration. Sometimes the solution mixture signals when the titration is complete by changing color. Often--as in this experiment--another substance is added (an indicator) to show when the reaction is finished. This is generally the case for acid-base reactions since most acids and bases are colorless and react to form colorless products. Very accurate titrations are done with a device called a buret which you will use in later experiments. But the general technique is the same regardless of whether you use expensive glassware or a simple dropper. You must measure the initial volume of one solution and then add the other until the reaction is complete. The added volume of the second solution is then compared to the original volume to determine the ratio between the reactants. In this experiment this comparison is simplified by the fact that the solutions are of equal concentration, but it is possible to use different concentrations and even solids! Rather than use a buret, you will use a pulled beral pipet. These disposable plastic pipets have thin stems which can be made even thinner by gently pulling on the plastic. This process draws the stem into a fine tip which gives a more reproducible drop size, if the dropper is held vertically. To get an accurate result, you will need to use a reasonable number of drops (50 is not too many--more is better). Since the plastic material the droppers is made from is non-wetting, you can use the same pipet to deliver all of the solutions. That way all your drops will be created equal if you are careful. Just be sure to rinse the pipet thoroughly with distilled water between solutions--and stir during the reaction. Oh, and don't forget to add a drop of indicator before you begin titrating! The chemicals Phenolphthalein is a white powder, a complex hydrocarbon derivative that was used in laxatives (once in Ex-Lax

TM

). It also functions very well as an acid/base indicator for most titrations except those involving ammonia. It is very sensitive to CO2 (which can dissolve in water to a small extent) and for very accurate determinations all water for solutions must be boiled to expel CO2. For titrations it is used in a 1% ethanol solution. It is practically insoluble in water. Bromthymol Blue is another very complex hydrocarbon molecule that is useful as an acid/base indicator. Like phenolphthalein it is not very soluble in water and an alcoholic solution is used (1% in 50% alcohol) Sodium hydroxide is commonly known as lye or caustic soda. It is a very hygroscopic white solid (absorbs water from the air rapidly) and also absorbs CO2. It is very corrosive to vegetable and animal matter and aluminum metal, especially in the presence of moisture. Dissolving NaOH in water generates considerable heat. Besides its use in the laboratory, sodium hydroxide is used in commercial drain cleaner preparations, to treat cellulose in the manufacture of rayon and cellophane and in the manufacture of some soaps. It is corrosive to all tissues and can be detected on skin by the "slimy" feeling associated with bases. It should be rinsed off thoroughly upon contact. It can damage delicate eye tissues and cause blindness. Hydrochloric acid is also known as muriatic acid. It is the same liquid acid that is often used in controlling the pH of swimming pool water. It is sometimes colored yellow by iron impurities, traces of chlorine and organic matter. Reagent grade HCl contains about 38% hydrogen chloride gas, close to the limit of its solubility at room temperature. Hydrochloric acid in concentrated form has the sharp, choking odor of hydrogen chloride. It is used in the

UNIT 2 24

production of other chlorides and in refining some ores (tin and tantalum), cleaning metal products, removing scale from boilers and heat-exchange equipment, and as an important laboratory reagent (often in diluted form). Concentrated solutions cause severe burns; permanent visual damage may occur. Inhalation causes coughing, choking; inflammation and ulceration of the respiratory tract may occur. Ingestion can be fatal. Sulfuric acid is a clear, colorless oily liquid in concentrated form (98%). It is highly corrosive and has a high affinity for water, abstracting it from wood, paper, sugar, etc., leaving a carbon residue behind. Dilution of concentrated sulfuric acid generates a tremendous amount of heat. Here in the lab your instructors prepare the dilute sulfuric acid you use by pouring the concentrated acid slowly over ICE while stirring! Even so, the resulting solution is very warm. As with all acid dilutions, acid is added to water, not the reverse, since the heat generated can boil the water at the point of contact and cause spattering. Sulfuric acid is used to make fertilizers, explosives, dyes, parchment paper, and glue. It is used, in concentrated form, in automobile batteries as the electrolyte. It is corrosive to all body tissues and contact with eyes may result in total blindness. Ingestion may cause death. Frequent skin contact with dilute solutions may cause dermatitis. Nitric acid has been called "aqua fortis" (strong water). It is generally produced by the oxidation of ammonia followed by reaction of the gaseous products with water. When pure it is a colorless liquid that fumes in air with a characteristic choking odor. "Concentrated" nitric acid is a water solution containing 70% HNO3. Even dilute solutions will stain woolen fabrics and animal tissue yellow. It is a very strong oxidizing agent, reacting violently with most organic matter. Nitric acid is used in the manufacture of fertilizers, dye intermediates and explosives. Calcium oxide (also called lime or quicklime) is produced commercially by heating limestone (principally CaCO3) . It readily absorbs CO2 and H2O from the air. Soluble in water, it forms Ca(OH)2, liberating much heat. Calcium oxide is used in the manufacture of bricks, plaster, mortar, stucco and other building materials, as well as in the processing of steel, the manufacture of glass and paper, the production of Na2CO3, dehairing hides, clarification of cane and sugar beet juices, in fungicides and insecticides and in water and sewage treatment (whew!). Calcium carbonate is generally produced from limestone but exists in a variety of forms naturally, including the mineral calcite (crystals notes for the double refraction properties), chalk and marble. It is used in the manufacture of paint, rubber, plastics, paper, ceramics, inks and cosmetics; as a filler it is found in matches, adhesives, pencils, crayons and linoleum. It is sometimes used to reduce the acidity of wines. Copper makes up about 0.01% of the earth's crust. It is one of the earliest known metals and is known for its unique reddish color when pure. However it becomes dull when exposed to air, forming oxides of copper, and in moist air becomes coated with green copper carbonate (this is part of the patina that appears on old copper or copper alloy exposed to the elements--like the Statue of Liberty). It is very slowly attacked by dilute hydrochloric acid and sulfuric acid, while nitric acid can readily dissolve it. It also slowly dissolves in aqueous ammonia. Copper is used in the manufacture of bronzes (copper + tin) and brasses (copper + zinc), and is used extensively in electrical conductors (wires, printed circuits, etc.). Of course copper also makes up a percentage of nearly all U.S. coins minted today. Copper itself probably has little or no toxicity, but some of its compounds can be quite hazardous. Iron occurs to the extent of about 5% in the earth's crust. As a refined metal it has been known for centuries. When pure it is silvery white or gray, hard, malleable and slightly magnetic. Stable in dry air, it readily oxidizes (rusts) in moist air. Its chief uses are as a structural alloying agent (in steels) and in the manufacture of some permanent and electromagnets. Aluminum is a white, malleable metal with a somewhat bluish tint. It can take a high polish in dry air but oxidizes

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superficially in the presence of moisture. As a pure metal or in alloys, aluminum is used in aircraft, cooking utensils, electrical apparatus and even in some dental alloys. Powdered aluminum is found in some fireworks, explosives and aluminum paints. ____________________________________________________ Analysis These questions should be answered in your laboratory notebook following your observations and data. 1. What color is phenolphthalein in acid, base, and neutral solutions? 2. Same question as #1, but for bromthymol blue. 3. Same question as #1, but for litmus. 4. Write the balanced molecular equation for the reactions between HCl and NaOH. Do the same for H2SO4 and NaOH. 5. Write balanced net-ionic reactions for the processes in #4. 6. Use your titration data to show that the ratio between HCl and NaOH matches the one shown in your reaction written for #4. Do the same for H2SO4 and NaOH. 7. Which indicator did you choose for your titrations and why? 8. What visual evidence is there for the reaction of the acids with CaO? With CaCO3? Choose one acid and write balanced molecular reactions for that acid with the two compounds. 9. Write balanced net-ionic reactions for the processes in #9. 10. What visual evidence is there that nitric acid does not react with metals in the same way as hydrochloric or sulfuric acid? Choose one metal and one acid (not nitric) and write a balanced molecular reaction for what happened when they were combined. 11. Write a balanced net-ionic reaction for the process in #10.

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UNIT 2 27

Name:___________________________ Per.:____ Date:_______________

Redox Reactions Purpose: to observe some displacement reactions with various metals and also with the halogens, in order to try to develop rules for what will displace what; in addition, to investigate the relationship between the results of a drop titration and the balanced redox reaction Method: I will test metal strips for displacement reactions with metal ion solutions and acid, looking for visual changes which indicate a reaction has occurred. I will also try various combinations of halogens (in water solution) with halide ions, using hexane to help identify the color of the remaining halogen better. These experiments should allow me to develop rules for deciding which metal (or halogen) will displace which in a potential reaction. I will also perform a drop titration with MnO4

- ion and Fe2+ ion (in acidic solution) to compare the result with the balanced equation for the reaction. Observations/Data:  

Cu(NO3)2

AgNO3

Zn(NO3)2

Mg(NO3)2

HCl

Cu

Ag

Zn

Mg

 

 

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NaCl KBr KI

Cl2

Br2

I2

 

 

drops MnO4- drops Fe2+

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LAB: Redox Reactions Reactions involving oxidation and reduction processes are very important in our everyday world; they make batteries work and cause metals to corrode (or helps to prevent their corrosion). They enable us to obtain heat by burning fuels--in factories and in our bodies. Many redox reactions are complex. However, combustion and synthesis (from elements) are two ordinary examples which require very little description. Just a little more involved is the displacement reactions, with which this exercise is mainly concerned. Your text book divides these processes into three categories:

hydrogen displacement metal displacement halogen displacement

You may remember (?) how to tell if a chemical reaction is occurring by looking at the behavior of a combination of chemicals, but can you predict whether or not a reaction like those above will happen? Displacement reactions involve an element, and a compound containing a "similar" element. "Similar" can mean both metals, or it can simply mean both + or both -. In any case, the general rule that applies is like displaces like. Copper and silver are both metals (both positive in compounds) and are both used for coinage and jewelry. So what will happen if a drop of solution containing silver ions is placed on a piece of copper metal? What will happen if a drop of solution containing copper ions is placed on a piece of silver metal? Chlorine and bromine are both non-metals (both negative in compounds) and both are used for disinfecting and bleaching. What will happen if chlorine is added to a solution containing bromide ions? What will happen if bromine is added to a solution containing chloride ions? Permanganate ions (MnO4

-) react with iron(II) ions (Fe2+) in acid solution. What about the ratio in which these ions react? Does the balanced equation correctly predict what happens? Your task is to answer questions like these about the substances available during this experiment. Preparing to experiment You will be provided with the following materials: 1. strips of copper, silver, magnesium and zinc metal 2. synthetic steel wool for cleaning the metal strips (DO NOT WET THIS) 3. solutions of Cu(NO3)2, AgNO3, Mg(NO3)2, Zn(NO3)2, HCl 4. solutions containing Cl2, Br2 and I2 (use about 5 drops) 5. solutions of NaCl, KBr, KI (use about 5 drops) 6. hexane (use about 5 drops) 7. small test tubes 8. disposable plastic beral pipets 9. solutions of MnO4

- and Fe2+ ions in equal concentrations (i.e., 1 drop of MnO4

- solution contains the same number of ions as 1 drop of Fe2+ solution) [start with at least 50 drops of Fe2+] 10. hand lens

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Design an experiment that will enable you to decide which metals will displace each other from solution and which will displace hydrogen from acid solution. (hint: the easiest way to do this is to place a drop of each solution on each cleaned metal strip) Design an experiment that will enable you to decide which halogens will displace each other from solution (see Technique section). Design an experiment that will enable you to predict the ratio between MnO4

- and Fe2+ in a balanced redox reaction for acidic solution. (hint: MnO4

- is purple, Fe2+ is almost colorless and the products of their reaction are colorless) [hint: try a titration!] Pre-lab take-home quiz Answer these questions on a separate piece of paper to be turned in the day you do this experiment. 1. What distinguishes a redox reaction from a non-redox reaction? (hint: it has to do with electrons—check lecture transcripts!!) 2. The evidence that a redox reaction has occurred is similar to the evidence for any chemical reaction. There are at least four things to look for when trying to decide if any kind of chemical reaction has happened. What are they? 3. Are precipitation or acid/base reactions considered redox reactions? Explain. Techniques 1. Extraction Sometimes dissolved material is more soluble in one kind of liquid than another. It is also possible that certain properties are more pronounced in a different liquid. Such is the case with the halogens used in this experiment. Each of the halogens has a characteristic color that is best observed in a liquid like hexane rather than one like water. But water and hexane do not mix. However, if the two liquids are shaken with one another, and a halogen is present in the water, it will move into the hexane layer (where it is more soluble) and show its particular color. In the past such "extractions" have been done by shaking a mixture of the liquids in a stoppered test tube. However, if small amounts (less than 1 mL) of both liquids are used, it is possible to use a plastic pipet or dropper and alternately suction and expel the mixture in rapid sequence, perhaps three times, then let the liquids separate. Hexane is less dense than water and will float to the top where any distinctive color may be noted. The chemicals Silver is one of the few metals that can be found native and is also found associated with copper, gold, or lead. It constitutes 1 x 10-5 % of the earth's crust. It is more malleable and ductile than any metal except gold and when pure is perhaps the best electrical conductor. It is insoluble in acids except nitric or hot concentrated sulfuric. Most silver salts are light sensitive (they darken on exposure to light) and this property makes them useful in the manufacture of photographic film. Small crystals of silver chloride are used in some photochromic eyeglasses. Absorption of light causes the AgCl to dissociate into Ag and Cl. The finely dispersed silver atoms tint the glass gray. The reverse reaction occurs in subdued light. Silver has been used in coinage for centuries as well as for jewelry (most often alloyed with gold or copper) and in the manufacture of some laboratory vessels. It is also an ingredient in dental amalgams for filling teeth. Silver has been used to purify drinking water (about 1 part per 20 million), is used in a process for purifying swimming pool

 

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water (instead of chlorine) and appears to have no serious toxicity. However, prolonged absorption of silver compounds can lead to grayish-blue discoloration of skin. Many silver salts are irritating to the skin. Magnesium is one of the most common elements in the crust of the earth, making up about 2.1% of it. In addition to its occurrence in minerals, it is also found in sea water, and in animal and plant organisms. Magnesium is a silvery-white metal which slowly oxidizes in moist air. It reacts very slowly with water at room temperature, less slowly at 100oC and burns in a current of steam. It reacts readily with dilute acids liberating hydrogen gas. Magnesium also reduces carbon monoxide, carbon dioxide, sulfur dioxide, and nitrogen oxides at red heat. It reacts directly with nitrogen, sulfur, the halogens, and many other elements. Magnesium is a constituent of light metal alloys, used for manufacturing precision instruments, optical mirrors, in fireworks, flash bulbs and flares. Copper(II) nitrate (the source of Cu2+ in this experiment) is generally sold in hydrated form and is deliquescent (it absorbs water from the environment). It is blue and very soluble in water. It is used in light-sensitive reproductive papers, as a coloring in ceramics, as a mordant in dyeing and printing, and in wood preservatives as a fungicide and herbicide. It is irritating to the skin. Silver nitrate (the source of Ag+ in this experiment) forms colorless, transparent crystals. It is stable and not darkened by light in pure air but darkens in the presence of organic matter and H2S. It decomposes at low red heat into metallic silver. It is used in photography and the manufacture of mirrors, silver plating, indelible inks, hair dyes, etching ivory and as an important reagent in analytical chemistry. It has been used as a topical antiseptic in a 0.1 to 10% solution. However, it is caustic and irritating to skin. Silver nitrate stains skin and clothing. These stains will wear off skin in a few days to a week but clothing is generally ruined. Swallowing silver nitrate can cause severe gastroenteritis that may end fatally. Magnesium nitrate (the source of Mg2+ in this experiment) consists of colorless, clear, deliquescent crystals which are readily soluble in water and alcohol. It is used in pyrotechnics. Zinc nitrate consists of colorless, odorless crystals which are readily soluble in water and very hygroscopic (quickly absorbs water from the air). It is used as a mordant in dyeing. Speaking of a mordant, don't you wonder what that is???? Everything seems to be a mordant! Dyes have to be "fixed" in their fabrics or they will wash out. A mordant reacts with the dye to form an insoluble compound, and thus helps prevent it from washing out. Potassium Bromide (the source of Br - in this experiment) is a white solid which is very soluble in water. It is used in the manufacture of photographic papers and in some engraving processes. Potassium iodide (the source of I- in this experiment) is a white solid, slightly deliquescent, and prone to oxidation in air. It is used in the manufacture of photographic emulsions, and in table salt and some drinking water to help prevent iodine deficiency disease. Chlorine is the eleventh most abundant element, making up about 0.19% of the earth's crust. Sea water contains nearly 3% NaCl. It is produced on a large scale by electrolysis of molten chloride or brines. Small amounts for use in the laboratory are often produced by the reaction of MnO2 and HCl. Chlorine is a yellowish-green diatomic gas at room temperature. In dilute water and hexane solutions it is essentially colorless. It has a suffocating odor. It forms explosive mixtures with hydrogen and many finely divided metals will burn in chlorine. It combines with all other elements except the noble gases. It is a member of the halogen family. Chlorine is used for bleaching, purifying water, and making synthetic rubber and plastics. It is a powerful irritant and excessive exposure can cause death. Bromine is a dark reddish-brown fuming liquid at room temperature, consisting of diatomic molecules. In dilute water and hexane solutions its color varies from golden to dark orange. It is a member of the halogen family and has a chemistry similar to chlorine. It attacks all metals and organic tissues and vaporizes readily at room temperature. Fumes are highly irritating to eyes and lungs.

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Bromine is used for bleaching silk, disinfecting spas, and manufacturing anti-knock compounds. Pure liquid bromine on the skin can cause painful, serious burns which heal only slowly. Even aqueous solutions (like in this experiment) should be handled carefully to avoid direct contact. Iodine [see previous experiment on zinc iodide] Sodium chloride (the source of Cl- in this experiment) is, of course, common table salt (the non-iodized version). It occurs in nature as the mineral halite and is produced by mining underground deposits as well as from sea water by solar evaporation. It is white in small granular form but large crystals are translucent. The salt sold in the grocery store usually contains some calcium and magnesium chlorides which help absorb water and prevent caking. Natural salt is the source of essentially all chlorine and sodium as well as of all, or nearly all their compounds (including HCl). It is used for preserving foods (salt curing), in the manufacture of soaps and dyes, in freezing mixtures (for making ice cream!) in dyeing and printing, and in some metallurgy. Potassium permanganate consists of dark purple or bronze-like crystals and is a strong oxidizing agent. It is used extensively in laboratory work and also in dyeing wood, bleaching, photography, and tanning. Dilute solutions are mildly irritating to the skin and high concentrations are caustic. Potassium permanganate stains skin and clothing like silver nitrate. Analysis These questions should be answered in your laboratory notebook following your observations. 1. It is possible to "rank" the metals (and hydrogen) based on their reactions. Look over your observations from the first part of this experiment and list the metals from most reactive to least reactive. Include hydrogen in this list. Compare your list to the activity series on p. 113 in your text and discuss any differences. 2. Write balanced net-ionic equations for the reactions you observed between the metals and their ions (including H+). 3. It is also possible to "rank" the halogens in a manner similar to the metals, with the most reactive halogen followed by less reactive ones. Use your observations to do this [how can you decide whether the color in the hexane is due to the halogen you added or the product of a possible reaction????]. Fluorine was not used in this experiment. Based on your observations and referring to the periodic table, where do you think fluorine would be placed in your series? 4. Write balanced net-ionic equations for the reactions you observed between the halogens and their ions. 5. Write a balanced redox reaction between MnO4

- and Fe2+ in acidic solution [among the products are Mn2+ and Fe3+]. Discuss the ratio between these two ions in the final balanced equation in light of your observations when you mixed the two solutions.

 

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Unit 2 Sample Test The test will include 5 multiple choice questions, 4 required problems, and one essay question. Beginning with this test you will be expected to master the solubility rules gradually. You need to know Rule 1 this time. The others will be given. The activity series will also be provided to assist you in determining whether or not certain redox reactions are likely to occur. And of course a periodic table will be provided. The following are representative of typical multiple choice questions but do not necessarily indicate topics to be addressed on your actual test. _____1. In the reaction below which takes place in aqueous solution

K2SO4 + Pb(NO3)2 PbSO4 + 2 KNO3

how many moles of soluble product are formed for each mole of lead(II) nitrate? a. 0 b. 0.5 c. 1 d. 2 _____2. In a balanced equation which of the following is always conserved? a. ions b. molecules c. moles of molecules d. moles of atoms _____3. In the halogen family, which is the most reactive element (i.e., which will replace all others in a reaction)? a. F b. Cl c. Br d. I _____4. The reaction products of a typical mixture of acid and base include a. water b. a salt c. water and a salt d. none of these _____5. Which substance below is oxidized during the reaction shown?

Fe + H2SO4 FeSO4 + H2

a. Fe b. H2SO4 c. FeSO4

d. H2

UNIT 2 34

The next section consists of representative problems which might be found in the problem section. All students are expected to work on all 4 of the problems. 6. Magnesium metal, like most metals, reacts vigorously with mineral acids such as hydrochloric acid in a redox displacement reaction. a. Write a balanced molecular equation for this reaction: ______________________________________________________________________________ b. Write a balanced net-ionic equation for this reaction: ______________________________________________________________________________ 7. Permanganate ions react with oxalate ions (C2O4

2-) in an acidic solution. Among the products are Mn2+ and CO2. Write the balanced redox reaction between these two ions in an acidic solution: ______________________________________________________________________________ If 50 drops of the C2O4

2- solution was titrated with permanganate solution until the purple color of the permanganate solution persisted, how many drops of permanganate solution was used? (assume that the solutions of ions were of equal concentrations) 8. Some of the substances commonly used in over-the-counter preparations for neutralizing excess stomach acid include Mg(OH)2 and Al(OH)3. Write a balanced net-ionic reaction for the neutralization of each by hydrochloric acid, assuming that the original bases are insoluble solids.

 

______________________________________________________________________________ 

___________________________________________________________________________ 

UNIT 2 35

9. One of the three metals, X, Y, Z is Sn. Consider the data below for displacement reactions and determine which metal is Sn (the activity series--provided on the test--on p. 108 will be helpful)

Co Y X Cu Z H+ x x x x Z2+ x Cu2+ x x x X2+ x Y2+ x

["x" indicates that a reaction has taken place] Write the balanced net-ionic equation for any ONE of the reactions indicated above: _____________________________________________________________________________ 10. In order to decide whether a solution contains SO4

2-, Cl-, and/or S2-, a student adds Pb2+ to one portion and Ba2+ to another. A precipitate forms with Pb2+ but not with Ba2+. Using the solubility rules, what can you determine about the contents of the solution? 11. Write the balanced molecular equation for the reaction of phosphoric acid with potassium hydroxide: ______________________________________________________________________________ If solutions of equal concentration are mixed, how many drops of potassium hydroxide will be required to make 100 drops of phosphoric acid with bromthymol blue added turn green? The final section of the test will consist of one essay question selected from the following topics: electrolyte behavior in reactions taking place in solution lab equipment (glassware, etc.) applications of the activity series

UNIT 2 36