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Assignments - Chemistry Honors(Archived)
Chapter 10 suggested problems
Due Date: 11/20/2012
Subject: Chemistry Honors

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Organic Model Project
Due Date: 10/10/2012
Subject: Chemistry Honors

Organic Project

Work in groups of 2-3, chosen at random, to make an organic molecule with household items. The molecule cannot simply be drawn on a poster, but items should be glued or attached together. For example various colored styrofoam balls representing different atoms can be used, while spaghetti or macaroni strung together can serve as bonds between the atoms. Items can be bought at the craft store to represent atoms and bonds, but household items can serve just as well. Keep in mind that 3- dimensional aspects of the molecule will apply.  

You should include a typed description of the chemical, its function, how it works and what effects it may cause in excess or in depletion in the body. You will need to include a legend to indicate the atoms in the chemical, as well as a picture of the molecule. Grades will be based on neatness, accuracy, completeness and ingenuity/inventiveness. A rubric is attached.

Choose a molecule that is present in the body and is necessary for life function, other than DNA, one that is required by the body and is obtained daily through food or medication (vitamins, insulin, etc.), or one commonly consumed (milk, chocolate, soda, etc.), taken for dietary reasons (B12, folic acid, etc.), or as a medicinal drug (cold/flu, etc.). Look at major organ functions in the body- don’t forget the brain!

One of the molecules listed below may be used or you may choose your own, but it must be approved by the teacher. A molecule may be used by only one group in all classes together. Hurry before they’re gone….


Ascorbic Acid       Caffeine                     Galactose               Estrogen

Guaifenesin           Isopentyl Acetate     Maltol                     Norethindrone

Pyrethrin II           Taxol                           Vanillin                   Vitamin E



Funtional Group Identification Worksheet and Answer Keys
Due Date: 10/5/2012
Subject: Chemistry Honors

Please go to forms- worksheets and answers are posted there

Intermolecular Forces Worksheets (1)
Due Date: 10/3/2012
Subject: Chemistry Honors

Intermolecular Forces Worksheet Answers


1)         Using your knowledge of molecular structure, identify the main intermolecular force in the following compounds. You may find it useful to draw Lewis structures to find your answer.


            a)         PF3                 dipole-dipole force


            b)         H2CO              dipole-dipole force


            c)         HF                   hydrogen bonding



2)         Explain how dipole-dipole forces cause molecules to be attracted to one another.

            Polar molecules have partially positive and partially negative sides (which correspond to the side of the molecule which is more or less electronegative). Because opposite charges attract one another, these molecules stick electrostatically.




3)         Rank the following compounds from lowest to highest boiling point: calcium carbonate, methane, methanol (CH4O), dimethyl ether (CH3OCH3).

            By using intermolecular forces, we can tell that these compounds will rank:

            methane (Van der Waals forces), dimethyl ether (dipole-dipole forces), methanol (hydrogen bonding), calcium carbonate (ionic electrostatic forces that are much stronger than intermolecular forces).




4)         Explain why nonpolar molecules usually have much lower surface tension than polar ones.

            Because the molecules aren’t attracted to each other as much as in polar molecules, these molecules are much less likely to have high surface tension.

Intermolecular Forces Worksheet (2)
Due Date: 10/2/2012
Subject: Chemistry Honors

Intermolecular Forces - Key


For questions 1-5, identify the main type of intermolecular force in each compound:


1)         carbon disulfide

            Van der Waal forces


2)         ammonia

            Hydrogen bonding


3)         oxygen

            Van der Waal forces


4)         CH2F2

            Dipole-dipole forces


5)         C2H6

            Van der Waal forces



Rank the following compounds by increasing melting point:


6)         C2H6, C2H5OH, C2H5F

            C26 (-183.30 C), C2H5F (-143.20 C), C2H5OH (-117.30 C)



7)         H2S, H2O, H2

            H2 (-259.30 C), H2S (-85.50 C), H2O (00 C)



8)         BBr3, BI3, BCl3

            BCl3 (-107.30 C), BBr3 (-460 C), BI3 (49.90 C)



All melting points were taken from The Handbook of Chemistry and Physics, 72nd Edition, by the Chemical Rubber Company. If you don’t have a CRC, you need one because it contains all the reference material you’ll ever need!

Intermolecular Forces Worksheet (3)
Due Date: 10/2/2012
Subject: Chemistry Honors

Types of Intermolecular Forces - Solutions

What is the strongest intermolecular force present for each of the following compounds?


1)         water                                                  hydrogen bonding


2)         carbon tetrachloride                        London dispersion forces


3)         ammonia                                           hydrogen bonding


4)         carbon dioxide                                  London dispersion forces


5)         phosphorus trichloride                    dipole-dipole forces


6)         nitrogen                                             London dispersion forces


7)         ethane (C26)                                   London dispersion forces


8)         acetone (CH2O)                                dipole-dipole forces


9)         methanol (CH3OH)                          hydrogen bonding


10)      borane (BH3)                                     dipole-dipole forces

Polarity Worksheet (1)
Due Date: 10/2/2012
Subject: Chemistry Honors

Polarity Worksheet


For each of the following pairs of molecules, determine which is most polar and explain your reason for making this choice:


1)         carbon disulfide                   OR                  sulfur difluoride





2)         nitrogen trichloride              OR                  oxygen dichloride





3)         boron trihydride                    OR                  ammonia





4)         chlorine                                 OR                  phosphorus trichloride





5)         silicon dioxide                      OR                  carbon dioxide





6)         methane                                OR                  CH2Cl2





7)         silicon tetrabromide                         OR                  HCN





8)         nitrogen trifluoride               OR                  phosphorus trifluoride


Polarity Worksheet (1) Answers
Due Date: 10/2/2012
Subject: Chemistry Honors

Polarity Worksheet Answers


For each of the following pairs of molecules, determine which is most polar and explain your reason for making this choice:


1)         carbon disulfide                   OR                  sulfur difluoride

        carbon disulfide is nonpolar




2)         nitrogen trichloride              OR                  oxygen dichloride

            both are polar, but oxygen dichloride is less symmetric than nitrogen trichloride, making it more polar.




3)         boron trihydride                 OR                  ammonia

        boron trihydride is nonpolar.




4)         chlorine                                 OR                  phosphorus trichloride

            chlorine is nonpolar




5)         silicon dioxide                      OR                  carbon dioxide

        It’s a tie, because both are nonpolar




6)         methane                                OR                  CH2Cl2

            methane is nonpolar




7)         silicon tetrabromide                         OR                  HCN

        silicon tetrabromide is nonpolar




8)         nitrogen trifluoride               OR                  phosphorus trifluoride

            Both are polar and equally symmetric, but the difference in electronegativity between N-F is less than that between P-F

Polarity Worksheet (2)
Due Date: 10/2/2012
Subject: Chemistry Honors

Ranking Molecules by Increasing Polarity


In each of the following problems, rank the molecules from lowest to highest polarity:


1)         PF3, LiOH, SF2, NF3








2)         Ni(OH)3, N2H2, CH3OH, C2H5OH








3)         B2F, H2C2O4, CuCl2, CF2O








4)         PH, PF3, NH3, NF3








5)         H2O, H2S, HF, H2


Polarity Worksheet (2) Answers
Due Date: 10/2/2012
Subject: Chemistry Honors

Ranking Molecules by Increasing Polarity



In each of the following problems, rank the molecules from lowest to highest polarity:


1)         PF3, LiOH, SF2, NF3

            NF3 < PF3 < SF2 < LiOH





2)         Ni(OH)3, N2H2, CH3OH, C2H5OH

            N2H2 < C2H5OH < CH3OH < Ni(OH)3






3)         B2F, H2C2O4, CuCl2, CF2O

            B2F4 < H2C2O4 < CF2O < CuCl2






4)         PH, PF3, NH3, NF3

            PH3 < NH3 < NF3 < PH3






5)         H2O, H2S, HF, H2

            H2 < H2S < H2O < HF

Homework #19-32 Give Geometry and Polarity
Due Date: 10/2/2012
Subject: Chemistry Honors

Check under forms- could not load it here

Homework Assignment

#19-32 Lewis Structures

Give Geometry and tell polar/nonpolar

 ·       Note: Some compounds have more than 1 correct Lewis structure, thus more than 1 correct geometry





Lab 4- 1.Reactivity of Groups I and II; 2. Alkaline Earth Metals
Due Date: 9/21/2012
Subject: Chemistry Honors

Periodic Trends of Chemical Reactivity Lab


Background Information

The periodic table, arranged according to the electron configurations of the elements, can be used to predict the physical and chemical properties of elements and their compounds. The vertical columns of the table are referred to as groups; the horizontal rows are called periods. General trends exist within groups and periods on the periodic table. Some of these trends include atomic size, ionization energy, electronegativity, density, melting point, and chemical reactivity. For example, the atomic size tends to increase as you move down a group and decrease as you move across a period. As you move down a group, the number of energy levels increases, so there are more “clouds” around the nucleus. As you move across a period, the energy level remains constant, but the number of protons (atomic number) is increasing. The positive nuclear charge is increasing which pulls the electrons closer to the nucleus and therefore reduces the atomic radius. Valence electrons are the electrons involved in chemical reactions. Metals tend to lose their valence electrons in chemical reactions, and in doing so, produce a full outer energy level like the noble gases. Ionization energy is the amount of energy required to remove an electron from the outer energy level. Nonmetals tend to gain electrons in chemical reactions, and in doing so, produce a full outer energy level like the noble gases. Electronegativity measures the tendency of an atom to gain electrons. This experiment will investigate chemical reactivity as a function of an element’s location within a group and period. You will determine the trend of chemical reactivity of metals as you move down a group and across a period on the periodic table. You will study the reactivities of the alkali metals (Li, Na, K), alkaline earth metals (Mg, Ca), and aluminum (Al).


Pre-Lab Questions: Answer the following in your OWN words.

1. a. Of the metals you are testing, which are the in same group?

   b. Of the metals you are testing, which are in the same period?

2. What are valence electrons and why are they important?

3. a. What is the trend in ionization energy as you move down a group? Explain why in terms of atomic size.

   b. What is the trend in ionization energy as you move across a period? Explain why in   terms of atomic size.4. Why is ionization energy important when discussing metal behavior in a chemical reaction?



1. Predict the trend in chemical reactivity for metals as you move down a group. Explain

your prediction.

2. Predict the trend in chemical reactivity for metals as you move across a period. Explain your prediction.


Procedure (WEAR GOGGLES!!!!)

1. Observe the reactivity of group I metals with water (lithium, sodium, potassium, rubidium, and cesium) by watching the demonstration video shown by your teacher. Record your observations. Observe the reactivity of the following group II and III metals with water and 0.5M HCl(hydrochloric acid): calcium, magnesium, and aluminum.  Put a small piece of each metal in the well plate. Add a few drops of water to each metal with a pipette. Record your observations.   Remember the observations that signify a chemical change (chemical reaction) is occurring—look for these!

2. Dry off any metals that did not react in water. Place each metal in another well and add a few drops of 0.5M HCl. Record your observations.

3. Throw away any metals that did not react (wrapped in a paper towel in the regular trash can).

4. Rank all of the metals in order of their reactivity, including those in the video, from most reactive to least reactive. Record this below your data table.


Data Table

Read through the entire procedure and prepare a data table that will allow you to record all observations for all the metals.


Post-Lab Questions/ Analysis

1. a. Restate your hypothesis about the trend in reactivity as you move down a group.

   b. What is the trend in reactivity as you move down a group? What evidence do you have for this trend from your experiment?

   c. Was your hypothesis supported or rejected by the data?

   d. Explain your observations in reactivity by relating this trend to the trends in atomic size and ionization energy.

2. a. Restate your hypothesis about the trend in reactivity as you move across a period.

   b. What is the trend in reactivity as you move across a period? What evidence do you have for this trend from your experiment?

c. Was your hypothesis supported or rejected by the data?

d. Explain your observations in reactivity by relating this trend to the trends in atomic size and ionization energy.

3. Would you expect the trend in reactivity of the Halogens to be the same as the reactivity trend you observed in metals? Explain your answer. (Hint: Halogens gain electrons in a reaction.)

4. Why did we not study the trends in reactivity of the Noble Gases?

5. What are some important uses of the alkali metals observed?


The Alkaline Earth Metals Lab



The elements in Group 2 of the periodic table are called the alkaline earth elements. Like the elements in Group 1 (the alkali metals), the elements in Group 2 are chemically active and are never found in nature in the elemental state. Like all members of a group, or family, the elements in Group 2 share certain common characteristics. The metallic character—the tendency to donate electrons during chemical reaction of the Group 2 elements increases as you go down the group. The more metallic of these elements typically react with water to form hydroxides and hydrogen gas. An example of such a

reaction would be:

Ca(s) + 2HOH(l) Þ Ca(OH)2(aq) + H2(g)


As metallic character increases (as you go down the group), the tendency for these elements to form ions increase. Also as you go down the group, the solubility of the hydroxides formed by the elements of this group increase. The more active is the metal, the more basic is its saturated hydroxide solution. The solubility of alkaline earth compounds also shows some interesting and useful tendencies. For example, the sulfate compounds of alkaline earth metals show decreasing solubility as you go down the group. This characteristic is used as a means of separating and identifying metallic ions of this group. Carbonates of all alkaline earth metals are quite insoluble. In this experiment, you will observe some of the characteristics of the alkaline earth metals discussed here.



Investigate some reactions of some Group 2 elements and gain some insights into the properties of these alkaline earth elements.



balance                          pH paper                    burner                        stirrer

flame tester                    test tube holder          filter paper                  test tubes (3)

test tube rack                wood splints



calcium turnings (Ca)                                    saturated solutions of:

magnesium ribbon (Mg)                                  calcium hydroxide (Ca(OH)2)

magnesium sulfate crystals                            magnesium hydroxide (Mg(OH)2)

(MgSO4)                                                          barium hydroxide (Ba(OH)2)

calcium sulfate crystals                                  0.1 M solutions of:

(CaSO4)                                                          sodium carbonate (Na2CO3)

barium sulfate crystals(BaSO4)                   magnesium chloride (MgCl2)

distilled water                                                  calcium chloride (CaCl2)

phenolphthalein solution                                barium chloride (BaCl2)









1. Pour about 2.5 mL of distilled water into a clean, dry test tube and place the tube in the test tube rack. Add 1 piece of calcium turning to the water in the tube. To collect the gas being released, invert a clean, dry test tube over the reactant tube with the calcium, holding the inverted tube with a test tube holder (Figure 1).

2. Test for hydrogen gas by inserting a burning wood splint into the upper part of the inverted tube (Figure 2).

3. Add a few drops of phenolphthalein solution to the reactant tube with a pipette. Do not insert the pipette down into your test tube. Hold it above the test tube in order to avoid contamination. After making your observations, discard the contents of the tube and clean and dry the tube. Use a brush if necessary.

Phenolphthalein indicator remains colorless in acid and turns pink in base.

4. Repeat step 1, using a 5-cm piece of magnesium ribbon in place of the calcium. If no visible reaction occurs, heat the water in the test tube to boiling, using a test tube holder to hold the tube over the burner flame. CAUTION: Point the tube away from yourself and others while heating.

5. Once the water is boiling, stand the tube in a test tube rack and, using a test tube holder, invert a collecting tube over the reactant tube. After a few seconds, test for hydrogen gas.

6. Turn off the burner and add a few drops of phenolphthalein to the reactant tube by holding the pipette above your test tube and adding drop wise. Record your observations. Discard the contents of the tube, and clean and dry the tube.



7. Obtain a few drops each of saturated solutions of calcium hydroxide, magnesium hydroxide, and barium hydroxide in a well plate. Be sure to NOT obtain any of the solid settled on the bottom of the beaker, only the liquid solution. Test each solution with a small piece of pH paper. Record the pH of each solution by comparing the color of the tested piece to the color chart.



8. Using the laboratory balance, measure out a 0.25-g sample of magnesium sulfate. Place it in a clean, dry test tube.

9. Repeat step 8 for calcium sulfate and barium sulfate.

10. Add 1.25 mL of distilled water to each tube. Using a glass stirring rod, stir each mixture thoroughly, getting as much of each solid to dissolve as possible. Record your observations of the relative solubility of each of these compounds.

11. Conduct a flame test for calcium ions (Ca2+) and for barium ions (Ba2+). Dip the wire loop of a flame tester into the solution of calcium sulfate. Place the loop in the burner flame. Observe and record the color of the flame. Clean the loop and repeat the test on the barium sulfate solution.



12. Stand 3 clean, dry test tubes in the test tube rack. Add about 2.5 mL of the 0.1M MgCl2 solution to one tube, 2.5 mL of the 0.1M CaCl2 solution to a second tube, and 2.5 mL of 0.1M BaCl2 to the third tube.

13. To each of the solutions in the test tubes, add about 1 mL of the Na2CO3 solution. Record your observations.





Ca + HOH: Result of test for H2 gas




Result of adding phenolphthalein




Mg + HOH: Result of test for H2 gas (before heating)




Result of test for H2 gas (after heating)







Result of adding phenolphthalein





pH readings:

Mg(OH)2  Ca(OH)2  Ba(OH)2



Apparent solubility:

MgSO4  CaSO4  BaSO4


Flame test results:

Ca2+  Ba2+














1. Describe the reactivity of the metals in Group 2 in terms of their location in the group.

2. How does the reactivity of an alkaline earth metal compare with that of an alkali metal (Group 1) in the same period?

3. What ionic charge can the alkaline earth metals exhibit?

4. Why does the metallic character of the alkaline earth metals increase as you go down the group?

5. What are some important uses of the alkaline earth metals?

Chapter 6 Answers Ionic and Metallic
Due Date: 9/17/2012
Subject: Chemistry Honors


1. an attraction between atoms involving e-(-) and nuclei(+)

2. ionic, metallic and covalent

5. if they can lower their potential energy

9. NaCl- more energy released upon forming the compound; more energy would be needed to break the bonds

14. it refers to a full s and p sublevel in the outermost energy level of an atom, which equals 8 e-; this is a more stable arrangement for atoms and results in lower PE

15. a) 1; b) 7; c) 2; d) 6; e) 3; f) 5; g) 4

19. a) 3; b) 2; c) 7 d) 6; e) 4; f) 5; g) 3; h) 6

25. a) formed by the attraction between cations and anions; b) crystalline solid

26. a) smallest ratio of atoms representing the formula of the compound; b) one Ca2+ ion and two F- ions

27. a) the energy released when 1 mole of ionic compound is formed; b) larger LE = stronger ionic bond

31. empty or partially empty orbitals, low IE values and low EN

32. a) attraction between valence e- of metals (sea) and nuclei of metal atoms; b) heat of vaporization

Answers: Properties of Ionic Compounds Worksheet
Due Date: 9/17/2012
Subject: Chemistry Honors

1. The attraction between cations and anions is very strong; in a lattice arrange the cations/anions are surrounded oneach side by oppositely charged ions resulting in bonding in all directions which strngthens the structure

2. ions need to be mobile to conduct electicity

3. smallest repeating unit of lattice that represents overall structure of crystal

4. Ions are held rigidly in place in lattice by attractive forces- if these forces are disrupted by movement then the lattice breaks

5. 1) not ionic- b.p. is too low

   2) not ionic- solid ionic compounds do not conduct

   3) not ionic- ionic compounds are not gelatinous

Answers: Metallic Bonding
Due Date: 9/17/2012
Subject: Chemistry Honors

A metallic bond can exist between ...metallic elements

Bonding occurs between pool of valence e- (sea) and nuclei of metal atoms (picture)

Metallic Bonds are stronger for metals having more valence e- and smaller ion size (closer packing).

Physical Properties

1. Because bonding occurs in such a way that shifting the crystal structure does not disrupt the attraction between e- and nuclei- it occurs in all directions

2. mobile sea of e- within solid structure

3. many overlapping orbitals very close in energy leads to absorption and emission of many wavelengths of light- emission results in photons of light = shiny appearance

4. Soft = fewer valence e- and larger ion size; hard = more valence e- and smaller ion size

5. Groups I and II have fewer valence e- and are larger in size so they cannot pack as tightly resulting in a weaker bond.

6. mobile sea of e- in molten state

7. metallic bond is very strong- requires large input of energy (heat)

Alloys are a blend of 1 or more metals in a mixture.

Advantages: hardness, color, malleability, ductility, conductivity, strength, corrosion resistence, higher m.p.

Properties: 1) very reactive

                  2) low IEs

                  3) low ENs

Lab 3- Making Micro-Hindenburgs
Due Date: 9/14/2012
Subject: Chemistry Honors

Making Micro-Hindenburgs



The Hindenburg was the largest airship ever built. It had a record of 54 successful flights. However, what people remember most about the Hindenburg is its tragic end at Lakehurst, New Jersey. On May 6, 1937, the hydrogen-filled Hindenburg burst into flames as the pilot was attempting to land it. Thirty-six people were killed while horrified onlookers watched. The exact cause of the explosion has never been determined, but the Hindenburg disaster essentially ended wide-scale development and use of airships.


            Why was the Hindenburg filled with hydrogen? Would another gas have been better? In this lab you will make hydrogen gas and investigate its properties so that you can answer these questions.


Pre-Lab Questions

Read the entire procedure and relevant pages in your textbook, then answer the following questions.

1. Find hydrogen in the periodic table and explain why it is separated from the Group 1A.

2. What chemicals are used to produce hydrogen gas in this lab?

3. Why should you handle hydrochloric acid (HCl) with care?

4. What should you do if you accidentally spill some HCl?

5. What is the purpose of the bubble solution?

6. What do you predict will happen when you hold the flame near the bubble of gas produced in this investigation? On what information is your prediction based on?



What properties of hydrogen prevent its use in modern-day airships?



chemical splash goggles; test tube; test-tube rack; hydrochloric acid (HCl), 3.0 M

scissors; micropipette; solution bubble; one-hole rubber stopper; wooden splint; granulated/mossy zinc      



Wear your goggles, at all times during the investigation. Hydrochloric acid is corrosive. If you spill any acid, immediately wash the area with plenty of cold water and notify your teacher. Tie back loose hair and clothing when working with a flame.




1.    Put on goggles, gloves, and lab apron.

2.    Make a microfunnel, using the scissors to cut off the top of the plastic micropipet bulb.

3.    Insert the microfunnel into the one-hole rubber stopper as shown and set aside.

4.    Place about 1g of granulated zinc in the test tube. Place the test tube in the test tube rack. Then carefully pour 5 mL of 3.0 M hydrochloric acid (HCl) into the test tube.

CAUTION: Hydrochloric acid is corrosive. Avoid spills and splashes. If you do spill acid, immediately rinse the area with plenty of cold water and report the spill to your teacher.

5.    Record your observations.

6.    Insert the rubber stopper with the microfunnel into the test tube so that no gas can escape except by way of the microfunnel. With a micropipet provided by your teacher, place 5-10 drops of bubble solution into the microfunnel, as shown in Figure 18-2.

7.    Record your observations.

8.    Let bubbles accumulate in top of microfunnel.

9.    Light the splint and hold next to your bubbles. CAUTION: Tie back loose hair and clothing when working with the flame. Carefully bring the flame close to the bubbles rising from the microfunnel.

10.  Record your observations. (If generation of gas slows or ceases, extinguish the flame. Remove the stopper and add more zinc and HCl. Then reinsert the micro funnel and stopper, relight the splint, and test the bubbles.)



11.  Disassemble the apparatus and dispose of the reaction products in a container provided by your teacher. CAUTION: The product in the test tube, zinc chloride (ZnCl2), is a severe skin irritant. Avoid direct contact. If spills occur, wash the area with plenty of water.

12.  Clean up your work area and wash your hands before leaving the laboratory.



zinc with hydrochloric acid                                           ____________________________

bubble solution in microfunnel                                  ____________________________

flame held near bubbles                                           ______________________________


Post-Lab Questions

1. What evidence in this investigation suggests that a chemical reaction has occurred?

2. Based upon your data, what properties of hydrogen are demonstrated in this Investigation? (Hint: Why did the bubbles float in the air?)

3. What purpose did the bubble solution serve?

4. Since the German zeppelin Hindenburg filled with hydrogen gas was destroyed in a violent fire:

a.    Write a balanced chemical reaction for the combustion of hydrogen.

b.    Is this reaction exothermic or endothermic?

5. Though more dangerous, a given volume of hydrogen gas will lift more weight than an equal volume of helium. This is because hydrogen is less dense than helium. Use Avogadro’s hypothesis to explain why hydrogen is less dense than helium.

6. Why do you think that hydrogen was used to fill the Hindenburg?

7. Modern airships are filled with helium. Unlike hydrogen, helium does not burn. Use your knowledge of valence electron configurations to explain why helium is safer than hydrogen for use in airships.

8. Balloons are usually filled with simple hot air instead of helium. Use the ideal gas law to explain why a hot air balloon floats.

9. What other gas would have been a better choice than hydrogen to fill the Hindenburg?

10. Why do you think that this investigation is entitled “Making Micro-Hindenburgs”?

11. Why would it be dangerous to do this investigation on a larger scale?                                                              

Chapter 6 problems Ionic and Metallic
Due Date: 9/13/2012
Subject: Chemistry Honors

p. 209 1,2,5,9,14,15,19,25,26,27,31,32

Chapter 5 Problems
Due Date: 9/10/2012
Subject: Chemistry Honors

9/6 p.166 4,6,8,9,13

9/7 p.167 22,23,25,26,27,30,48

Chapter 5 Answers
Due Date: 9/10/2012
Subject: Chemistry Honors

9/6 p.166

4.a) outermost energy level is the same

   b) 8 electrons in outermost energy level- complete energy level

6. period corresponds to outermost energy level

8.a) group 1

   b) most reactive metals, silvery-white, soft, react vigorously with water

9.a) group 2

   b) not as reactive, harder, denser, stronger, higher m.p.

13.a) group 7

     b) most reactive nonmetals, react with metals, most electronegative elements

9/7 p.167

22.a) half distance between 2 like nuclei

     b) get smaller

     c) increasing nuclear charge pulls e- closer

23.a) get larger

     b) energy levels are being added

25.a) across- increase; down- decrease

     b) across- atoms get smaller so e- being removed is closer to nucleus; down- atoms get larger so e- being removed gets farther from nucleus

26.a) energy released when adding an e-

     b) negative- adding e- is favorable/ energy is released; zero-  adding electron is not favorable; positive- energy is needed to add the electron/ adding e- is not favorable

 27.a) cation- atom that has lost e-; anion- atom that has gained e-

30.a) ability to attract e- in a bond

    b) F is most electronegative element- all others are assigned values based on F

48.a)IE,EA(generally), and EN

     b) AR and IR

     c) nothing

Chapter 4 Answers
Due Date: 9/4/2012
Subject: Chemistry Honors

Answers: Chapter 4

2.wave:interference, diffraction, reflection, refraction, wavelength/frequency

particle:absorption/emission of light, line emission spectra

3.λ: 400-700nm; ν: 4.3-7.5x 1014Hz


6. a) inversely b) directly c) directly

8. ground- lowest energy state, closest to nucleus

excited- higher energy than ground

9. by e- falling from high energy state to lower energy state and emitting light as photons

10. 7.05x1016Hz

11. 2.35x10-16J

12. E=hc/λ

13. v=d/t .... c=d/t ....t=8.00x1010m/3.00x108m/s = 267sec

15. a) didsn't explain chemical behavior; b) only worked for H

16. a)main energy level b)n c) energy levels d) 2n2

17. a) number of sublevels and shape of orbitals; b) holds orbitals

18. a) 1-0 b) 2- 0,1 c) 3- 0,1,2 d) 4- 0,1,2,3 e) 7

19. a) number of orbitals b) s-1,p-3,d-5,f-7 c) lie along different axes (x,y,z)

20. a) n2 b) 3rd- 9; 5th-25

21. a) allowed values for e-; spin state b) +1/2; -1/2

22. a) 2 b) 18 c) 32 d) 72 e) 98

Chapter 4 problems
Due Date: 8/31/2012
Subject: Chemistry Honors

8/23/12  p.124  2,3,4,6,8,9,10,11,12,13,15

8/29/12  p.124  16,17,18,19,20,21,22


Lab 2- Flame Tests
Due Date: 8/31/2012
Subject: Chemistry Honors

Flame Test Lab


In this lab students will learn about atomic energy levels, emission spectroscopy, and flame tests for element identification.



Students will use small samples of 7 chloride salts of different metals. These they will place into a flame in order to observe the colors produced. These colors come from the excitation of electrons which then resume their ground states by emitting light of very specific colors.


The electrons in an atom occupy different energy levels, as you know. When all of the electrons are at the lowest possible energy level they are said to be in the ground state. Electrons do not always stay in the ground state. Sometimes they can be promoted to a higher-energy electron shell. This can happen in two ways. First, the electron can absorb a photon of just the right amount of energy to move it from one quantum shell to another. Second, when atoms are heated or energized with electricity their electrons can gain energy. This promotes them to the higher-energy shell. When an electron is in a higher-energy shell it is said to be in an excited state.

Electrons in excited states do not usually stay in them for very long. When electrons lose their energy they do so by emitting a photon of light. Photons are particles with energy but no mass. Their energy is directly proportional to the frequency of the light (remember: E = hν). The photons emitted precisely match the quantum energy difference between the excited state and the ground state.

The light produced by very hot atoms in the gaseous state is a unique spectrum for each element. To observe the spectrum requires the use of a prism, diffraction grating, or spectroscope. Before complex instruments were invented to observe elemental spectra chemists sometimes identified metals in compounds by doing a flame test. Salts are a type of compound that include a metal and a non-metal. Sodium nitrate (NaNO3) is the most familiar example of a salt but others include calcium nitrate (Ca(NO3)2) and copper(II) nitrate (Cu(NO3)2). In flame tests salts that are dissolved in water are evaporated using a hot flame. In the flame the metal atoms become excited and produce their characteristic spectrum of light. However, since the observer does not use a spectroscope only one color is observed. It turns out that many metals produce a unique single color under these conditions. Some metals do produce very similar colors but a practiced eye can often distinguish them. It is a traditional art of the chemistry laboratory to use these colors to identify specimens of compounds that contain unknown metals.

This ability of metal atoms to produce these colors is put to use by practitioners of the art of fireworks manufacture. By including different metal salts, or mixtures of metal salts, in the exploding shell of a firework, these artists can produce beautiful displays in nearly all the colors of the rainbow.


Pre-lab Questions

Answer the following questions and perform the following calculations.

  1. Describe the energy levels of an atom and how an electron moves between them. Be sure to describe how the electron absorbs and emits energy when it does so.
  2. What physical change happens to atoms and molecules in a solution that is strongly heated in a flame? Is there a phase change? If so, what is it?
  3. What is going on inside atoms of a metal when a metallic salt of that metal is dissolved in water and then placed in a hot flame?
  4. What would be required in order to observe a spectrum when viewing the flame test of a metallic salt? Why?
  5. What is required in order to be able to identify an element based on its flame test color?
  6. Why do different elements have different flame test colors?
  7. How to the designers of fireworks make the explosions have different colors?
  8. What are the wavelengths of light that are representative of the following colors: Violet, Blue, Blue-green, Green, Yellow-green, Yellow, Orange, and Red?



1.      well plate

a.          2 wells for 6M HCl

b.      7 wells for the metal nitrates, 1 each

  1. 1 inoculation loop
  2. 1 Bunsen burner
  3. sharpie for labeling
  4. distilled water
  5. a series of metal nitrate solutions such as Ca(NO3)2, Cu(NO3)2, Ba(NO3)2, Li(NO3), K(NO3), Na(NO3), and Sr(NO3)2
    (these will be provided in dropper bottles)
  6. 2 unknown metal nitrates




  • Wear goggles
  • Treat all chemicals in this lab as toxic. Do not touch any of them with your bare hands.
  • Wash well with water immediately if you touch chemicals accidentally
  • Use caution with the burner
    • Do not leave burner unattended
    • Place burner near middle of lab bench
    • Tie back long hair
    • Do not wear baggy clothing in the lab
    • Hot objects look like cold objects: be cautious!
  • Copper(II) nitrate is moderately toxic by ingestion; avoid contact with eyes, skin and mucus membranes.
  • Lithium nitrate is highly toxic by ingestion; avoid contact with eyes, skin and mucus membranes.
  • Wash your hands with soap and water after you complete the day’s lab work, even if you didn’t touch any chemicals directly



Remember to record your observations in your lab notebook or on a piece of paper in your binder before you leave class. When making observations be sure to use all of your senses except taste. Never taste anything in the chemistry lab. Chances are good you will regret it if you do.


Wavelength (nm)

Region (nm)



400 - 440



440 - 470



470 - 490



490 - 560



560 - 570



570 - 585



585 - 630



630 - 700

  1. You will share a set of metal salt solutions with the people at your lab station.
  2. Collect a small sample (a few drops) of each of the known metal salt solutions which your teacher has provided and carry them all to your lab bench.
  3. Obtain an inoculation loop for your group.
  4. Obtain 10 - 20 mL of 6M HCl in your labeled beaker.
  5. Each group member must record information in a neat table with the following columns. Make this table before you even turn on the gas.
    1. Name & Formula of Metal Nitrate
    2. Metal Ion
    3. Color of Flame
    4. Approx. Wavelength (nm)
    5. Approx. Wavelength (m)
  6. Clean the inoculation loop by swirling it gently in the 6M HCl. Then, once you light the burner, heat the loop until it glows red hot. This step removes any ions clinging to the loop from previous experiments.
  7. Light and adjust your Bunsen burner. Be sure to clean your loop carefully. Do not leave the loop in the flame too long as it can cause the loop to degrade and break.
  8. To do a flame test with each metal salt get a film of the solution of a salt inside the loop and bring it into the hottest part of the flame. If this produces poor color then try the edge of the burner flame. Repeat the dip into the salt solution as often as necessary to see the flame test color. Be sure not to over-heat the loop.
  9. Carefully note the color of each metal salt when it is put in the flame. Use the chart on the previous page to estimate the approximate wavelength of the color you see. Use the Representative Wavelength values. Record all data in the table you made earlier.
  10. Clean the inoculation loop using hydrochloric acid (HCl) and heat each time you change from one metal salt to another. Failing to do so will result in mixed flame test colors. Again, do not over heat the loop.
  11. Your teacher has prepared two solutions with two of the known metal salts. They are labeled Unknown 1 and Unknown 2.
  12. After testing all of the salt solutions place a few drops of each unknown in the well plate.
  13. Test the 2 unknown solutions as you did the known solutions and record all information regarding the 2 unknown solutions in the data table. Test the unknowns in the flame as many times as needed. You may also need to test one or more of the known solutions for comparison.
  14. Identify the 2 unknowns according to your observations. Write the name of each unknown down below your data table.
  15. Clean out the well plate using the method recommended by your instructor (hazardous wastes must be disposed of properly). Usually, all leftover solutions will be collected in designated waste containers for hazardous waste disposal.
  16. Wash all equipment carefully and thoroughly using soapy water.
  17. Put all equipment back where you found it.


Post- Lab Questions

  1. Why do different metals have different characteristic flame test colors?
  2. Most salts contain a metal and a non-metal. Look at the compounds we tested and determine whether it is the metal or the non-metal that is responsible for the color produced in the flame test for that salt. How can you be sure your answer is correct?
  3. What colors did the unknowns produce in the flame? What are the unknowns?
  4. Why do the chemicals have to be heated in the flame before the colored light is emitted?
  5. Could flame tests be useful in determining identities of metals in a mixture of two or more salts? If so, what problems might arise? If not, why not? Explain your answer.
  6. Which method is better for precisely identifying elements: examining the full spectrum using a spectroscope or using a flame test? Use your experience in the lab with both of these methods in answering this question. Justify your answer.


Chapter 21 problems*
Due Date: 8/21/2012
Subject: Chemistry Honors

8/14 p.703 2,9,10,14

8/15 p.703-704 11,12,13,16,18,19,20,21,23,24,26-29

8/20 p.704 30-32, 35-39

Chapter 3 problems
Due Date: 8/9/2012
Subject: Chemistry Honors

8/6/12 pg. 89-90 1,2,3,5,32,33

8/7/12 pg. 89-90 7,8,9,10,12

8/8/12 pg. 90 19,20




Lab 1- Law of Definite Proportions
Due Date: 8/9/2012
Subject: Chemistry Honors

Law of Definite Proportions Lab


Objective: To measure and calculate the ratio of magnesium to oxygen in magnesium oxide. To compare the lab ratio to the percent composition calculation based on the formula.



            The Law of Definite Composition states that the elements that form a compound always combine in the same proportion by mass. The compound water H20 always is a chemical combination of hydrogen and oxygen in a 1:8 ratio by mass. If a mixture of hydrogen and oxygen in some other ratio, say 1:2, were reacted, there would be water formed, but there would also be some unreacted hydrogen, because water always forms in the 1:8 ratio by mass.

            In this experiment, you will examine the reaction between magnesium metal, Mg, and the oxygen in the air, O2. The magnesium will be heated strongly in a crucible for several minutes. The mass of magnesium will be compared with the mass of the material produced.


Pre-lab Questions:

  1. Why is it important to begin this experiment with a clean and dry crucible?
  2. What is the purpose of making sure the outside of the magnesium ribbon is shiny?
  3. With what element does the magnesium combine when it is heated?
  4. Why must you reheat the crucible repeatedly until the last two masses agree within 0.03g?
  5. Suppose a compound of sodium and chlorine is formed in the ratio of 1.54g of chlorine for each gram of sodium. How much sodium would you need to completely react with 45.0g of chlorine?
  6. List any safety concerns associated with this lab.



Crucible and Lid           Wire Gauze w/ Ceramic Center            Magnesium Ribbon

Crucible Tongs Safety Goggles                         Distilled Water Bottle

Centigram Balance        Bunsen Burner                         Clay Triangle   

Ring Stand and Ring



                             Put on your goggles. Hot crucibles and magnesium can cause burns, so

                             use with caution. Handle hot crucibles with tongs and place the hot

                             crucible on the wire gauze to cool. Obtain a clean and completely dry

                           crucible and cover. Find the mass of the crucible and cover and record

                             it on the data table. Obtain a piece of magnesium ~12-15 cm

                             and roll the magnesium into a loose coil and place it in the crucible.

                             Find the mass of the crucible, cover, and magnesium. Record it on the        

                           data table.

  1. Set up the ring stand, ring, burner, and clay triangle as pictured above. Place the crucible on the triangle. Begin heating the crucible gradually with the lid completely on. Heat strongly by moving the flame around underneath the crucible. Remove the heat temporarily if a large amount of smoke comes out of the crucible.
  2. After about four minutes of direct heating with no smoke, remove the lid slightly. Heat the crucible for four more minutes. Finally, remove the lid completely and heat strongly for four more minutes.
  3. Turn off the burner and put the lid back on the crucible. Allow the crucible and cover to cool to a temperature low enough so that you can touch the crucible. Find the mass of the crucible, contents, and cover. Add ten drops of distilled water. Smell cautiously, note any odor. Put the crucible back on the ring-stand setup and heat again for four minutes with the lid on. Allow to cool again.
  4. Find the mass of the crucible, cover, and product. Record it on the Report Sheet.
  5. If enough time remains, reheat the crucible for four minutes, allow it to cool, and again find the mass. If this mass differs by more than 0.03g from the mass you found in Step 8, repeat this procedure for a second trial.
  6. If enough time remains, repeat the whole procedure for a second trial.
  7. Clean and put away all of the materials.
  8. Wash your hands thoroughly with soap and water.


Data and Observations:

Mass of Crucible and Cover


Mass of crucible, cover, & Mg


Mass of crucible, cover, & product (1st)


Mass of crucible, cover, & product (2nd)


Mass of crucible, cover, & product (3rd)



Calculations: SHOW WORK

  1. Calculate the mass of the Mg that reacted.
  2. Calculate the mass of the magnesium oxide that was produced.
  3. Calculate the mass of oxygen that reacted.
  4. Calculate the ratio of the mass of magnesium to the mass of oxygen.
  5. The accepted ratio for the mass of magnesium to oxygen is 1.52:1. Calculate your percent error.

% error = (calculated value – accepted value) x 100

                             accepted value


Mass of Magnesium reacted.


Mass of Magnesium Oxide produced.


Mass of Oxygen reacted.


Ratio of Magnesium to Oxygen.


Percent Error







Post-Lab Discussion:

The magnesium metal is an element that combines with another element, oxygen gas, to form the compound magnesium oxide. The ratio of the mass of magnesium oxide to the mass of magnesium should be constant for all of your trials, regardless of the mass of the magnesium that you started with.

            The strong heating insured that all of the magnesium reacted with the oxygen in the air to form magnesium oxide. Since some magnesium nitride (magnesium + nitrogen) could have formed, the addition of water and subsequent heating were done to remove that product from the crucible.

            In order to calculate the ratio, you must first find the masses of magnesium oxide alone and of magnesium alone by subtracting the mass of the crucible from the masses that you recorded. The ratio is then calculated by:


            Ratio = Mass of Magnesium Oxide / Mass of Magnesium



  1. How would your results be affected if all of the magnesium did not react?
  2. Using your ratio, determine the formula of magnesium oxide. Remember the Law…atoms combine in whole number ratios- 1.5 : 1 is the same as ?
  3. Use the accepted ratio to determine the mass of magnesium that would combine with exactly 16.0g of oxygen.
  4. Suppose you tried to combine 42g of Mg with 45g of oxygen.
    1. Which of the substances would have some left after the reaction?
    2. How much magnesium oxide would be formed?



When 1.0lb of gasoline is burned in an automobile, approximately 3lbs of carbon dioxide is given off. Carbon dioxide is one of the gases contributing to global warming. What information from this experiment helps to explain how one pound of gasoline can give off approximately 3 times as much CO2?