Name: · 2020. 10. 8. · Ionic Bonding Work sheet Name: _____ 1. Use the periodic table to write...

Ionic Bonding Work sheet Name: ___________________

1. Use the periodic table to write appropriate electron-dot structure for each of the following:

a) oxygen b) sodium c) calcium d) fluorine e) xenon f) indium

(#8) (#11) (#20) (#9) (#54) (#49)

g) argon h) lithium i) sulfur j) aluminum k) bromine l) barium

(#18) (#3) (#16) (#13) (#35) (#56)

2. For each of the above atoms, predict how many electrons the atom would gain (or lose) to achieve a stable octet and indicate what charge that would give it as an ion: a) oxygen b) sodium c) calcium d) fluorine e) xenon f) indium

_______ ___ _______ ___ _______ ___ _______ ___ _______ ___ _______ ___


_______ ___ _______ ___ _______ ___ _______ ___ _______ ___ _______ ___

3. Name each of the ions you formed in #2 above: (Hint: positive ions keep the same name: sodium -> sodium. Negative ions get an -ide ending: oxygen -> oxide) a) oxygen b) sodium c) calcium d) fluorine e) xenon f) indium

_____________ _____________ _____________ ____________ ____________ ____________


_____________ _____________ _____________ ____________ ____________ ____________

4. Use electron-dot structures to diagram the formation of the ionic compounds that would form between each of the following pairs, also name the ionic compound: a) sodium & oxygen name: __________________ b) cesium & bromine name: _________________

c) lithium & fluorine name: __________________ d) sodium & nitrogen name: __________________

e) magnesium & sulfur name: ___________________ f) aluminum & fluorine name: __________________

g) aluminum & phosphorus name: _______________ h) barium & phosphorus name: __________________

Na

Na O

[Na]1+ [ O ]2-

[Na]1+ Na2O

sodium oxide

Na O

gain 2 2- lose 1 1+

oxide ion sodium ion

1

The following 36 ions are very common and will be used often throughout the year. So you will need to memorize them

forward & backward – their names, their formulas and their charges. Flash cards are recommended! Deadline: Nov 22!

2

Positive Ions (Cations) Variable-Charge

Cu1+ copper(I) Cu2+ copper(II) Fe2+ iron(II) Fe3+ iron(III) Co2+ cobalt(II) Co3+ cobalt(III) Pb2+ lead(II) Pb4+ lead(IV)

Positive Ions (Cations) Fixed-Charge

1+ Na1+ sodium K1+ potassium H1+ hydrogen Ag 1+ silver NH4

1+ ammonium 2+ Mg2+ magnesium Ca2+ calcium Zn2+ zinc 3+ Al3+ aluminum

Negative ions (Anions) Polyatomic

1- OH1- hydroxide NO3

1- nitrate ClO3

1- chlorate BrO3

1- bromate IO3

1- iodate C2H3O2

1- acetate HCO3

1- bicarbonate 2- CO3

2- carbonate SO4

2- sulfate CrO4

2- chromate 3- PO4

3- phosphate

Negative ions (Anions) Monatomic

1- F1- fluoride Br1- bromide Cl1- chloride I1- iodide 2- O2- oxide S2- sulfide 3- N3- nitride P3- phosphide

Positive Ions (Cations) Variable-Charge

Au1+ gold(I) Au3+ gold(III)

Cr2+ chromium(II) Cr3+ chromium(III)

Mn2+ manganese(II) Mn3+ manganese(III)

Ni2+ nickel(II) Ni3+ nickel(III)

Hg22+ mercury(I)

Hg2+ mercury(II)

Sn2+ tin(II) Sn4+ tin(IV)

Positive Ions (Cations) Fixed-Charge

1+ Li1+ lithium Rb1+ rubidium Cs1+ cesium Fr 1+ francium 2+ Be2+ beryllium Sr2+ strontium Ba2+ barium Ra2+ radium Cd2+ cadmium 3+ Sc3+ scandium

Negative ions (Anions) Polyatomic

1- CN1- cyanide CNO1- cyanate SCN1- thiocyanate HCO2

1- formate MnO4

1- permanganate NH2

1- amide N3

1- azide Al(OH)4

1- aluminate 2- Cr2O7

2- dichromate S2O3

2- thiosulfate C2O4

2- oxalate MoO4

2- molybdate SiO3

2- silicate C4H4O6

2- tartrate 3- AsO4

3- arsenate BO3

3- borate

Negative ions (Anions) Monatomic

1- H1- hydride 2- Se2- selenide Te2- telluride 3- As3- arsenide

Also know the following prefixes & suffixes:

hypo- -ite, -ite, -ate, per- -ate, bi-, and dihydrogen-

Also: here are the nonmetals (and some metalloids):

B = boron C = carbon N = nitrogen O = oxygen F= fluorine Ne = neon Si = silicon P = phosphorus S = sulfur Cl = chlorine Ar = argon As = arsenic Se = selenium Br = bromine Kr = krypton Te = tellurium I = iodine Xe = xenon At = astatine Rn = radon

(AKA: hydrogen carbonate)

Ionic Naming Work sheet Name: _____________________

Write appropriate formulas/names for the following ionic compounds

1. sodium fluoride ____________ 7. KCl ________________________________________

2. potassium sulfide ____________ 8. Ag2O ________________________________________

3. lithium oxide ____________ 9. AlI3 ________________________________________

4. silver chloride ____________ 10. Li2S ________________________________________

5. magnesium bromide ____________ 11. Na3P ________________________________________

6. aluminum sulfide ____________ 12. Zn3N2 ________________________________________

The following ionic compounds involve variable-charge metal ions. For them, Roman numerals are necessary to indicate the charge on the ion. For example: CuO is not just “copper oxide;” it is “copper(II) oxide,” so that we don’t get it confused with copper(I) oxide [Cu2O]

13. copper(I) fluoride ____________ 19. FeN ________________________________________

14. copper(II) fluoride ____________ 20. FeO ________________________________________

15. cobalt(III) sulfide ____________ 21. Au2S ________________________________________

16. iron(II) oxide ____________ 22. NiBr3 ________________________________________

17. iron(III) oxide ____________ 23. PbI2 ________________________________________

18. mercury(II) chloride ____________ 24. SnS2 ________________________________________

The following ionic compounds contain polyatomic ions (ions like NO31-, SO4

2- or OH1- which are made up of several atoms bonded together). Whenever a polyatomic ion needs to be doubled or tripled in a formula, parentheses must be used to avoid confusion. For example: magnesium nitrate = Mg(NO3)2 [not MgNO32] and aluminum hydroxide = Al(OH)3 [not AlOH3] If a polyatomic ion does not need to doubled or tripled, there is no need for parentheses.

25. sodium carbonate ____________ 30. Fe(CN)3 ________________________________________

26. copper(II) cyanide ____________ 31. KNO2 ________________________________________

27. zinc sulfate ____________ 32. Ag2SO3 ________________________________________

28. ammonium phosphide ____________ 33. Au(C2H3O2)3 _____________________________________

29. iron(III) chromate ____________ 34. Co2(CO3)3 ______________________________________

The following problems just contain more practice of the problems above:

35. iron(III) phosphate ____________ 42. Cu3BO3 ______________________________________

36. iron(II) phosphate ____________ 43. SnI2 ______________________________________

37. zinc bicarbonate ____________ 44. AgC2H3O2 ______________________________________

38. magnesium fluoride ____________ 45. HgS ______________________________________

39. silver nitrate ____________ 46. Na2O ______________________________________

40. calcium hydroxide ____________ 47. Mg3(PO4)2 ______________________________________

41. copper(II) nitride ____________ 48. KF ______________________________________

3

and... because you can never have too much practice... (Careful... some easy elements and ions have been mixed in.) The first five on each side are done for you: 49. sulfur ____________ 80. Ag2CO3 ______________________________________

50. potassium sulfide ____________ 81. K2O ______________________________________

51. copper(I) oxide ____________ 82. Mg2+ ______________________________________

52. aluminum ____________ 83. Cu2+ ______________________________________

53. aluminum ion ____________ 84. Cu ______________________________________

54. aluminum nitrate ____________ 85. FeO ______________________________________

55. iron(II) nitrite ____________ 86. CdF2 ______________________________________

56. aluminum acetate ____________ 87. Ba(NO3)2 ______________________________________

57. ammonium phosphate____________ 88. NH4C2H3O2 ______________________________________

58. calcium iodide ____________ 89. ZnBr2 ______________________________________

59. sodium carbonate ____________ 90. PO43- ______________________________________

60. iron(III) ion ____________ 91. Fe(OH)2 ______________________________________

61. iron(III) sulfite ____________ 92. Fe(OH)3 ______________________________________

62. copper(II) bromide ____________ 93. Al2S3 ______________________________________

63. mercury ____________ 94. Na2SO3 ______________________________________

64. barium chloride ____________ 95. Pb4+ ______________________________________

65. ammonium sulfate ____________ 96. Cu3(BO3)2 ______________________________________

66. iron(III) phosphate ____________ 97. Cu3BO3 ______________________________________

67. iron(II) phosphate ____________ 98. SnI2 ______________________________________

68. zinc bicarbonate ____________ 99. Cr ______________________________________

69. magnesium acetate ____________ 100. HgSO4 ______________________________________

70. sulfite ion ____________ 101. Na2O ______________________________________

71. lithium hydroxide ____________ 102. Mg3(PO4)2______________________________________

72. copper(II) phosphate____________ 103. KHCO3 ______________________________________

73. calcium nitrite ____________ 104. S ______________________________________

74. barium ion ____________ 105. S2- ______________________________________

75. aluminum iodide ____________ 106. CuCO3 ______________________________________

76. sodium chloride ____________ 107. FeBO3 ______________________________________

77. magnesium sulfite ____________ 108. ZnBr2 ______________________________________

78. phosphate ion ____________ 109. Al(NO2)3 ______________________________________

79. zinc sulfide ____________ 110. SnBr2 ______________________________________

silver carbonate

potassium oxide

magnesium ion

copper(II) ion

copper

S

K2S

Cu2O

Al

Al3+

4

Covalent Bonding Worksheet Name: ___________________

1. Explain how covalent bonding occurs:

2. How is covalent bonding different than ionic bonding:

3. Answer: I) ionic bonding C) covalent bonding B) Both ionic & covalent N) neither.

__ Occurs between two nonmetals

__ Occurs between two metals

__ Occurs between a metal and a nonmetal

__ Involves only the outermost (valence) electrons

__ Positive and negative particles are formed

4. Neon (Ne) is a nonmetal. Why does it not tend to react with other nonmetals to form covalent bonds?

5. Do you think that Ne would react with a metal to form an ionic bond? ____

Explain your answer.

6. For each of the following pairs of atoms, state whether I) an ionic bond would form C) a covalent bond would form or

N) no bond at all would form.

___ Na & Br ___ Br & S ___ Br & Br ___ Mg & Ar ___ Mg & O ___ S & Ar ___ C & H

7. Label the following depictions as I) ionic C) covalent or N) neither

___ I’ll swap you mine for yours

___ Let’s pool together what we’ve got

___ You can have it; I really didn’t want it that much

___ I feel we both gain something from this relationship

___ Does this completed octet make me look fat?

___ As long as we stick together, we’ll be OK.

5

8. Use electron-dot structures to diagram the formation of the covalent compounds that would form for each of the following pairs. Also name the compound. If no compound forms, write NA in the box. The first two are done for you.

a) oxygen & fluorine name: _________________ b) hydrogen & bromine name: _________________

c) carbon & fluorine name: _________________ d) nitrogen & hydrogen name: _________________

e) hydrogen & oxygen name: _________________ f) chlorine & chlorine name: _________________

g) phosphorus & fluorine name: _________________ h) carbon & hydrogen name: _________________

i) hydrogen & fluorine name: _________________ j) sulfur & hydrogen name: _________________

k) hydrogen & hydrogen name: _________________ l) carbon & chlorine name: _________________

F O OF2

F

F O F

H Br H Br HBr

oxygen difluoride hydrogen monobromide

6

Covalent Naming Work sheet Name: _______________ Write appropriate formulas/names for the following covalent compounds. Use the lower left hand side of the chart for the name of the first component, and the upper right hand side for the name of the second component. Also use the Greek prefixes listed to describe how many there are of each atom. For example: N2O = dinitrogen monoxide

1. sulfur difluoride ____________ 16. HCl ______________________________________

2. dinitrogen tetroxide ____________ 17. H2S ______________________________________

3. nitrogen tribromide ____________ 18. AsI3 ______________________________________

4. selenium hexafluoride ____________ 19. CS2 ______________________________________

5. carbon tetrachloride ____________ 20. NH3 ______________________________________

6. dihydrogen monoxide ____________ 21. SF6 ______________________________________

7. phosphorus triiodide ____________ 22. SeBr2 ______________________________________

8. boron trichloride ____________ 23. XeF4 ______________________________________

9. bromine heptafluoride ____________ 24. N2O3 ______________________________________

10. diarsenic pentoxide ____________ 25. P2Br3 ______________________________________

11. tricarbon octabromide ____________ 26. As2I5 ______________________________________

12. sulfur dioxide ____________ 27. S2F8 ______________________________________

13. chlorine monofluoride____________ 28. Si3Cl6 ______________________________________

14. sulfur tetrachloride ____________ 29. NO2 ______________________________________

15. silicon dioxide ____________ 30. SO3 ______________________________________

And now... to mix in some ionic compounds, ions, and elements

31. ammonium phosphate____________ 43. NF3 ______________________________________

32. iron(III) hydroxide ____________ 44. CO32- ______________________________________

33. iron(III) ion ____________ 45. Cu3BO3 ______________________________________

34. iron ____________ 46. SnI2 ______________________________________

35. dicarbon hexabromide____________ 47. Al2O3 ______________________________________

36. zinc fluoride ____________ 48. Au2O3 ______________________________________

37. silver carbonate ____________ 49. P2O3 ______________________________________

38. carbon monoxide ____________ 50. Cr3(PO4)2 ______________________________________

39. bicarbonate ion ____________ 51. N2O ______________________________________

40. sulfur trioxide ____________ 52. K2O ______________________________________

41. sulfite ion ____________ 53. Cu2O ______________________________________

42. cadmium ion ____________ 54. Hg2+ ______________________________________

7

Write appropriate formulas/names for the following acids:

1. hydrobromic acid _______ 31. H3BO3(aq) _________________________________

2. nitrous acid _______ 32. HBrO2(aq) _________________________________

3. hypoiodous acid _______ 33. H2C2O4(aq) _________________________________

4. carbonic acid _______ 34. HI(aq) _________________________________

5. hydrocyanic acid _______ 35. HSCN(aq) _________________________________

6. chromous acid _______ 36. HIO4(aq) _________________________________

7. sulfuric acid _______ 37. H3PO4(aq) _________________________________

Write appropriate formulas/names for the following elements, ions, compounds & acids:

1. sulfur tetrafluoride _______ 31. AuF _________________________________

2. lead(II) chlorite _______ 32. Li3N _________________________________

3. potassium bicarbonate _______ 33. AsI3 _________________________________

4. selenium hexachloride _______ 34. HClO(aq) _________________________________

5. barium acetate _______ 35. PH3 _________________________________

6. dihydrogen monoxide _______ 36. SeBr6 _________________________________

7. chlorous acid _______ 37. HF(aq) _________________________________

8. chromium(III) oxide _______ 38. XeF4 _________________________________

9. zinc sulfate _______ 39. N2O3 _________________________________

10. diarsenic pentasulfide _______ 40. Mn2O3 _________________________________

11. dinitrogen monosellenide _______ 41. Al2O3 _________________________________

12. phosphorous acid _______ 42. C3F8 _________________________________

13. iron(II) perbromate _______ 43. Sn(NO3)4 _________________________________

14. iron(III) bromide _______ 44. NO2 _________________________________

15. iron(III) sulfide _______ 45. NO21- _________________________________

16. dichromic acid _______ 46. HC2H3O2(aq) ________________________________

17. aluminum hydroxide _______ 47. CO32- ________________________________

18. carbon monoxide _______ 48. Cu3PO4 _________________________________

19. cobalt _______ 49. H3AsO4(aq) _________________________________

20. cobalt(II) ion _______ 50. HgO _________________________________

21. sulfur trioxide _______ 51. Ni2(S2O3)3 _________________________________

22. sulfite ion _______ 52. P2S3 _________________________________

23. calcium cyanide _______ 53. Cr(MnO4)2 _________________________________

24. oxalic acid _______ 54. HMnO4(aq) _________________________________

25. lead(IV) chromate _______ 55. Ag3P _________________________________

26. ammonium chromate _______ 56. Cu _________________________________

27. xenon _______ 57. Cu1+ _________________________________

28. hypoiodous acid _______ 58. Mg2+ _________________________________

29. phosphorus pentafluoride _______ 59. Be(NH2)2 _________________________________

30. nitric acid _______ 60. HBrO4(aq) _________________________________

B = boron C = carbon N = nitrogen O = oxygen F = fluorine Ne = neon Si = silicon P = phosphorus S = sulfur Cl = chlorine Ar = argon As = arsenic Se = selenium Br = bromine Kr = krypton Te = tellurium I = iodine Xe = xenon

Two new things to learn/memorize:

1) This system for naming acids: -ate -ic -ite -ous -ide hydro- -ic

2) the nonmetals listed below (their symbols and names); these will no longer be given to you on quizzes or tests.

8

Lewis Structures (Electron-Dot Structures) worksheet Name: __________________

Write Lewis structures for each of the following molecules or ions. Practice in the left-hand box. Rewrite your final answer NEATLY in the right-hand box.

Ex: F2 7. HF

1. F2 NH3 8. Br2

2. BF3 9. CO

3. H2O 10. SO32-

4. BrF3 11. CH4

5. SF6 (Write small!)

12. OF2

6. CO2 13. SF4

F F

continued on next page

9

Lewis Structures worksheet (continued). In the last 3 problems, it is not clear exactly which atom (or atoms) are in the center. For these problems, the atomic arrangements are given. All you need to do is fill in the electrons correctly.

14. H2S 21. SO3

15. BeI2 22. BrH41+

16. CS2 23. N2

17. IBr21- 24. H2

18. O2 25. HCN

19. NH41+ 26. CH2O

20. CH2F2 27. C2H2

H C N

O H C H

H C C H

10

Molecular Shapes Note Sheet Name: ______________________

1.

2.

4.

7.

11.

3.

5.

8.

12.

6.

9.

13.

10.

14.

Hybrid.

Name of shape b.a.

ex AXE

3-D sketch of

molecular

shape

b.a. = bond angle

(not something else)

Learn these thirteen molecular shapes; know them by heart – their names, their bond angles and how to draw them.

11

Molecular Shapes Worksheet Name: ____________________

1. What does VSEPR theory stand for: _________________________________________________________ According to this theory, which of the following influences the shapes of molecules? (circle one) a) repulsive forces between the central atom’s protons b) repulsive forces between the central atom’s electrons c) attractive forces between the central atom’s protons and electrons ... and what are the two specific factors that determine the precise shape of a molecule. (circle two) a) the number of atoms bonded to the central atom b) the total number of protons in the c. a. c) the total number of electrons on the c. a. d) the number of valence electrons on the c. a. e) the number of valence electron groups on the c.a. 2. Based on the notes you took in class, fill in the spaces below with the missing information: # bonded atoms, # NEPs (NEP = nonbonding electron pairs – AKA: “ghosts”), drawing of shape, name of shape. 3. For each of the following molecules or ions, draw and name the shape. (Rather than use balls for the atoms, use the elements symbols). The first one is done for you.

3 bonded atoms 0 NEPs 120* ba

trigonal planar

2 bonded atoms 0 NEPs ______ ba

__ bonded atoms __ NEPs ____ ba

2 bonded atoms 3 NEPs ______ ba

__ bonded atoms __ NEP 90* & 120* ba

__ bonded atoms __ NEPs _____ ba

trigonal pyramidal

__ bonded atoms __ NEPs _____ ba

4 bonded atoms 2 NEPs _____ ba

__ bonded atoms __ NEPs ________ ba

see-saw

6 bonded atoms __ NEPs 90* ba

2 bonded atoms 1 NEP __ _____ ba

__ bonded atoms __ NEPs ______ ba

__ bonded atoms __ NEPs ______ ba

T-shaped

F C

F

F F F

a. CF4 b. SF2 c. BF3

12

tetrahedral

C F F

F

S F F F B F

F

4. Putting it all together... For each of the following molecules or ions, first draw the correct Lewis structure; then draw and name the shape. The first one is done for you.

a. CF4 b. SF2 c. BF3

d. SO2 e. BeBr2 f. PH3

g. SF4 g. NH41+ i. CS2

j. CH4 k. XeF2 i. H2O

a. CF4 b. SeF2 c. BeF2

d. CO2 e. SO3 f. SF6 write small

g. PH5 h. OF2 i. BrF41+

j. NH3 k. SO32- l. SiH4

13

O S O Be Br Br H P H H

F s F F

F H N H

H

H [ ]

1+ S C S

H C H H

H F Xe

F H O

H

F C

F

F F F

tetrahedral

C F F

F

Polar Bonds/Polar Molecules Worksheet Name: ________________________ 1. Consider the following molecules. Every line represents an electron pair being shared between the two atoms. But not all these electron pairs are being shared equally. At right is a list of electronegativity values (EV). The EV shows how much an atom tends to “hog” the electrons in a bond. If two bonded atoms have more or less the same EV (not more than 0.4 apart), then we say the electrons are being shared evenly between the two atoms and the bond is nonpolar -- that is, the electrons are pretty much evenly spread within the bond. But if one atom has a significantly greater EV than the other (with a difference of 0.5 or more), then we say the electrons are unevenly shared between the two atoms, spending more time around the atom with the higher EV and giving it a

partial negative (-) charge and giving the other atom a partial positive ( +) charge. This is what is called a polar

bond. Consider each molecule below. If it contains a polar bond, circle “PB” below the number, and label the +

and - ends of the bond. If a molecule does not contain any polar bonds, circle the “NPB” below the number. (#1 and #2 are done for you). Hint: Ignore the nonbonding electron pairs (AKA “ghosts”): About 17 of the 24 molecules contain polar bonds. 2. Now look back over the molecules above. Those that contain no polar bonds (like #2) have a fairly even distribution of electrons and the molecules as a whole are said to be nonpolar (they have no dipole moment). Now consider those that you found to contain polar bonds. In many of these (like #11), the polar bonds all cancel each other out – because they are symmetrically arranged around the central atom. These molecules are also considered to be nonpolar. In others (like #1), the polar bonds do not cancel out: that is, the electrons are being hogged

toward one side of the molecule. These molecules therefore have a - side and a + side. Such molecules are called “polar molecules” (they have a dipole moment). Find all the molecules above that are

polar, and draw a circle around the entire molecule. Then indicate which end of the molecule is - and

which is + as shown at right. Hint: about 12 of the 24 are polar; the rest are nonpolar molecules.

Cl

H

O H

H

H H

C H H H

C C C

1. 2. 3. 4. 5.

6. 7. 8. 9. 10.

Cl

B

11. 12. 13. 14.

C

15. 16. 17. 18. 19.

20. 21. 22. 23. 24.

H Br PB NPB

PB NPB

PB NPB

PB NPB

PB NPB

PB NPB

PB NPB

PB NPB

PB NPB

PB NPB

PB NPB

PB NPB

PB NPB

PB NPB

PB NPB

PB NPB

PB NPB

PB NPB

PB NPB

PB NPB

PB NPB

PB NPB

PB NPB

PB NPB

F 4.0 O 3.5 N 3.0 Cl 3.0 Br 2.8 C 2.5 S 2.5 I 2.5 H 2.1 P 2.1 B 2.0

F

F F

C F

H

F H

C H

H

S H

Br Br O C C O O H

H N H F

F N F

Cl Cl S

O O H H

H H

H H H H

C C C C H C

C C C H

H H H H H H H H

H H H H H H H H

water (H2O) methane (CH4) carbon tetrafluoride (CF4)

hydrogen (H2) hydrogen bromide (HBr)

bromine (Br2) carbon dioxide (CO2) carbon monoxide (CO) ammonia (NH3) nitrogen trifluoride (NF3)

boron trichloride (BCl3) sulfur dioxide (SO2)

propane (C3H8) octane (C8H18)

+

+ -

H

O H

+

+ -

+

-

C O H H

H H

H H ethanol (C2H5OH)

C C C H H

H H H H acetone (C3H6O)

O F

F S F

sulfur hexafluoride (SF6) F

F F

F

S

sulfur tetrafluoride (SF4) F

F F H C N

hydrogen cyanide (HCN)

fluoromethane (CH3F) hydrogen sulfide (H2S)

Cl

Cl P

phosphorus pentachloride (PCl5) Cl

Cl F

F I F

iodine pentafluoride (IF5)

F F

C H H

O

formaldehyde (CH2O)

14

3. Define electronegativity: __________________________________________________________________ 4. What is a polar bond: ______________________________________________________________________ 5. Two atoms, A and B, have formed a covalent bond between them. A has an electronegativity of 2.5, and B has an electronegativity of 2.7. We know that: (circle the two statements that apply) e- = electron a) the e-s in bond are being shared exactly evenly b) the e-s in the bond are being shared more or less equally c) A is hogging the e-s substantially more than B d) B is hogging the e-s substantially more than A e) A will acquirte a partial negative charge f) A will acquire a partial positive charge g) B will acquirte a partial negative charge h) B will acquire a partial positive charge i) the bond will be considered nonpolar j) the bond will be considered polar 6. Two atoms, C and D, have formed a covalent bond between them. C has an electronegativity of 3.0 and D has an electronegativity of 2.1. We know that: (circle the four statements that apply) a) the e-s in bond are being shared exactly evenly b) the e-s in the bond are being shared more or less equally c) C is hogging the e-s substantially more than B d) D is hogging the e-s substantially more than A e) C will acquire a partial negative charge f) C will acquire a partial positive charge g) D will acquire a partial negative charge h) D will acquire a partial positive charge i) the bond will be considered nonpolar j) the bond will be considered polar 7. What two things make a molecule polar? ________________________ _____________________________ 8. Is it possible for a molecule to contain polar bonds but not be a polar molecule? ______ Explain. Give an example and use a diagram. 9. Is it possible for a molecule to contain nonpolar bonds but still be a polar molecule? ______ Explain. Give an example and use a diagram. 10. a) Explain why the Lewis structure below for CH2F2 makes it look as though the molecule would be nonpolar. b) In fact, the molecule is polar. Explain. (hint: consider the actual shape of the molecule) 11. For each of the eight sketches at right, place the specified F and Br atoms around the sulfur atom to create molecules that are polar and nonpolar. (All molecules are octahedral. Not all answers are possible. When impossible, just draw a big “X” across it.)

F F C H

H SF3Br3

SF4Br2

SF2Br2Cl2

SF6

S

S

S

S

S

S

S

S

polar nonpolar

15

Reaction Terminology WS Name: __________________ Some definitions: Reactants: The substances (elements and compounds) that a reaction starts out with. Reactants are always written on the left side of a chemical equation. Products: The substances (elements and compounds) that are produced by a reaction. (They are what the reactants turn into, and they are chemically different than the reactants.) Products are always written on the right side of a chemical equation. Yields: This means “forms” or “reacts to produce.” It is represented by an arrow (*) in a chemical equation. This arrow separates the reactants on the left from the products on the right. Solid: A state of matter having a definite volume and shape. Its particles vibrate in place but do not move around one another. Most of the volume of a solid is due to the particles themselves, not to space between the particles. If a solid is heated up enough it will melt into a liquid, and then boil into a gas. Solid is represented by (s) in a chemical equation. Liquid: A state of matter having a definite volume but not a definite shape. It flows and takes on the shape of whatever container it is placed in. Like a solid, most of the volume of a liquid is due to the particles themselves, not to space between the particles. A liquid’s particles are moving fast enough that they move around each other but always stay close to neighboring particles. If a liquid is cooled down, it will freeze into a solid; if a liquid is heated up it will boil into a gas. Liquid is represented by (l) in a chemical equation. Gas: A state of matter having no definite volume or shape. It flows and expands and contracts to take on the shape and volume of any container it is placed in. Its particles move quickly and randomly, constantly bouncing into one another. Unlike a solid or a liquid, most of the volume of a gas is due to the space between the particles, not to the particles themselves. If a gas is cooled down, it will condense into a liquid and then eventually freeze into a solid. Gas is represented by (g) in a chemical equation. Aqueous: When a substance is dissolved in water it is said to be “aqueous.” Like a gas, aqueous particles are spread far apart and move randomly, but rather than moving through empty space as they do in a gas, they move through the water. Aqueous is represented by (aq) in a chemical equation. Chemical Formula: This is a way of representing what elements are present in a compound, and in what atom ratio they occur. So H2O is the chemical formula for water and it tells us that H (hydrogen) and O (oxygen) are present in a 2:1 atom ratio. Subscripts: These are the small numbers that are used in chemical formulas to show how many atoms of each element are present. Like the 2 in H2O. Note that it comes immediately after the atom it is referring to. Thus, the 2 in H2O means that there are two H’s. In Al2O3, there are 2 Al’s and 3 O’s. “1” is never used as a subscript. Instead, if there is no subscript, that counts as a “1.” Coefficients: These are the large numbers in a chemical equation that are written before a chemical formula to show how many of those atoms of molecules are involved in the reaction. As with subscripts, “1” is never used as a coefficient. Instead, if there is no coefficient, that counts as a “1.” Unlike subscripts, a coefficient comes immediately before the substance it is referring to. So in the equation: Mg + 2 H2O Mg(OH)2 + H2 the large 2 in front of the H2O is the coefficient and it shows that there are 2 H2O molecules involved in the reaction. All the other substances have no coefficients, so that means there is just one of each of them.

16

1. For the following three equations fill in the blanks with the following vocabulary words you just learned: reactants, products, reacts to produce, solid, liquid, gas, aqueous, chemical formula, subscript, coefficient: (not all of these will apply to each equation)

1. N2(g) + 3 H2(g) + 2 H2O(l) 2 NH4OH(aq)

2. Mg(s) + 2 H2O(l) Mg(OH)2(aq) + H2(g)

3. N2(g) + 2 O2(g) 2 NO2(g) 4. In each of the three equations above, circle one element, and put a box around one compound. 5. Write an equation like the ones above for each of the following described reactions. Include (s), (l), (g) & (aq) where appropriate. A) A chunk of calcium and a solution of hydrochloric acid react to produce a solution of calcium chloride and bubbles of hydrogen gas. (You will need a coefficient of 2 on the hydrochloric acid to balance the equation.)

B) Liquid bromine and a piece of aluminum react to produce a hard coating of aluminum bromide. (Use two 2’s and a 3 to balance the equation. Can you tell where to put them?)

C) A solution of lithium fluoride is mixed with a solution of lead(II) nitrate. They react to produce some tiny crystals of lead(II) fluoride and a solution of lithium nitrate. (You will need to use two 2’s as coefficients to balance the equation. Can you tell where to put them?)

17

Balancing Equations WS Name: _________________

Before we can learn to balance equations, we must be proficient at “counting atoms.” For 1-8, count the

number of each atom present. The first is done for you.

1. K3PO4 ___ K’s ___ P’s ___ O’s 5. 3 Ca(BrO3)2 ___ Ca’s ___ Br’s ___ O’s

2. Ba(NO3)2 ___ Ba’s ___ N’s ___ O’s 6. 2 C4H10O2 + 11 O2 ___ C’s ___ H’s ___ O’s

3. C3H7OH ___ C’s ___ H’s ___ O’s 7. 8 CO2 + 10 H2O ___ C’s ___ H’s ___ O’s

4. Al2(SO4)3 ___ Al’s ___ S’s ___ O’s 8. 3 Sn3(PO4)4 + 6 V(NO3)3 __ Sn’s __P’s __ V’s __ N’s __ O’s

For 9-16, take inventory of each side and determine whether the equation is balanced (Y) or not (N):

9. H2 + Cl2 2 HCl ___ 13. H2 + O2 2 H2O ___

10. 3 F2 + N2 2 NF3 ___ 14. 2 KClO3 2 K + Cl2 + 3 O2 ___

11. 3 Na + 3 H2O 3 NaOH + H2 ___ 15. 3 K2CO3 + 2 Al(OH)3 6 KOH + Al2(CO3)3 ___

12. C2H6 + 5 O2 2 CO2 + 3 H2O ___ 16. 2 C3H7OH + 9 O2 6 CO2 + 8 H2O ___

For the remaining problems, balance the equation by writing in the appropriate coefficients (lowest whole-

numbers). Check your answers by taking inventory. Hint: use pencil or erasable pen!

17. ___ K + ___ I2 ___ KI 24. ___ Li + ___ O2 ___ Li2O

18. ___ N2 + ___ O2 ___ N2O 25. ___ N2 + ___ H2 ___ NH3

19. ___ Fe + ___ O2 ___ Fe2O3 26. ___ KBr ___ K + ___ Br2

20. ___ MgCl2 ___ Mg + ___ Cl2 27. ___ Ag + ___ CuCl2 ___ Cu + ___ AgCl

21. ___ Al2O3 ___ Al + ___O2 28. ___ FeBr3 + ___ F2 ___ FeF3 + ___ Br2

22. ___ CO + ___ O2 ___CO2 29. ___BaO + ___ HCl ___ BaCl2 + ___ H2O

23. ___ NH4OH ___ NH3 + ___ H2O 30. ___ Na + ___ H2O ___ NaOH + ___ H2

31. ___ N2 + ___ F2 ___ NF3

32. ___ Al + ___ Fe2O3 ___ Al2O3 + ___ Fe

33. ___ CF4 + ___ Br2 ___ CBr3F + ___ F2

34. ___ Sn + ___ Fe2O3 ___ Sn3O4 + ___Fe

1 1 1 2 2 3

3 3 3 4 6 6

6 8 8 8 9 12

12 18 18 20

20 26 26 102

18

35. ___ N2 + ___ O2 ___ N3O5

36. ___ Sn3(BO3)4 ___ Sn + ___ B + ___ O2

37. ___ Ca + ___ HF ___ CaF2 + ___ H2

38. ___ NH3 + ___O2 ___NO + ___H2O

39. ___ K3PO4 + ___ Ca(OH)2 ___ Ca3(PO4)2 + ___ KOH

40. ___ Na2CO3 + ___ HCl ___ NaCl + ___H2O + ___ CO2

41. ___ C5H12 + ___O2 ___ CO2 + ___H2O

42. ___ Sn3P4 + ___ MgCl2 ___ SnCl4 + ___ Mg3P2

43. ___ Al2O3 + ___ C + ___ Cl2 ___ AlCl3 + ___ CO

44. ___ SiF4 + ___ H2O ___ H4SiO4 + ___ H2SiF6

45. ___ HNO3 + ___ P4O10 ___ N2O5 + ___ H3PO4

46. ___ C6H14 + ___ O2 ___ CO2 + ___H2O

47. ___ C2H6S + ___ O2 ___ CO2 + ___ H2O + ___ SO3

48. ___ Li2O2 + ___ H2O ___ LiOH + ___ O2

49. ___ Fe2O3 + ___ CO ___ CO2 + ___ Fe

50. ___ H3BO3 ___ H4B6O11 + ___ H2O

51. ___ C2H5NO2 + ___ O2 ___ CO2 + ___ H2O + ___ NO

52. ___ C2H3OF + ___ O2 ___ CO2 + ___ H2O + ___ HF

53. ___ C4H10O2 + ___ O2 ___ CO2 + ___ H2O

54. ___ Br2 + ___ H2O + ___ SO2 ___ HBr + ___ H2SO4

55. ___ ClO2 + ___ O3 ___ Cl2O6 + ___ O2

56. BONUS: ___ N3H7O + ___ NO2 ___N2O + ___H2O

#56 is very tricky. You can actually use

algebra to solve it. 100 bonus points will be

divided up among those students who show

how to do this and turn in their solution on

a separate sheet of paper by ________.

19

Reaction Types & Predicting Products WS Name: ____________ For 1-25, specify what all five reaction of that type have in common, then complete the reactions by writing in the predicted products. Remember: when you form an element, check “P S Br I N Cl H O F” and when you form a compound, cancel charges (refer to ion sheet for ions you don’t know).

Composition Reactions Decomposition Reactions Single Replacement Reactions

1. Na + Cl2 6. MgO 11. AgCl + Mg

2. O2 + K 7. AlCl3 12. Ca + FeF3

3. H2 + F2 8. H2O 13. HCl + Al

4. Li + N2 9. CaS 14. KBr + O2

5. Ca + S8 10. NF3 15. Cl2 + Al2O3

Double Replacement Reactions Combustion Reactions

16. CaCl2 + Al2O3 21. CH4 + O2

17. LiCl + Pb(NO3)2 22. C5H12 + O2

18. Na2SO4 + CaCl2 23. O2 + C6H6

19. HCl + K3PO4 24. C2H5OH + O2

20. HBr + Ca(OH)2 25. C12H22O11 + O2

For the remaining problems, identify the reaction type, then complete the reaction. (CP = composition, DC = decomposition, SR = single replacement, DR = double replacement, CB = combustion) __ 26. Na + CaF2

__ 27. Na + F2

__ 28. AgF + CaCl2

__ 29. C2H4 + O2

__ 30. K2S

__ 31. O2 + Mg

__ 32. F2 + AlBr3

__ 33. C2H6O + O2

__ 34. Li2S + MgCl2

__ 35. HCl + Zn

__ 36. BaBr2

__ 37. Na2O

__ 38. O2 + C6H12

__ 39. S8 + Na

__ 40. Al + Br2

__ 41. Al + CaCl2

__ 42. NBr3

__ 43. AlBr3 + CaO

__ 44. NH3

__ 45. C5H10O4 + O2

__ 46. CaO + HBr

__ 47. I2 + Mg

__ 48. NaF + CaBr2

__ 49. K2SO4 + Ca(NO3)2

__ 50. CuI2

__ 51. Na3N + Ca

__ 52. Na3N + Cl2

__ 53. O2 + C2H2

20

Activity Series Lab: When do single replacement reactions occur? Name: __________ Partner: __________

Consider the two sets of reactants shown at right: AlCl3 + Na NaCl + Al Although we have learned how to complete this type of single replacement reaction, we have not learned which of these single replacement reactions actually occur and which ones produce no reactions at all. In all single replacement reactions, either a metal is being pitted against another metal (as is the case in the above reactions), or a nonmetal is being pitted against another nonmetal. In either case, only the more reactive element will be able to successfully replace the other. For example, in the above reactions, it turns out that Na is a more reactive metal than Al. Remember that what makes metals reactive is their capacity to lose their electrons to some other substance. Na atoms are much better at getting rid of their electrons than are Al atoms. In the first reaction we have Na atoms trying to lose their electrons to Al3+ ions. Since Na is more reactive than Al, this reaction will in fact happen, and as the electrons get lost by the Na atoms and gained by the Al3+ ions, it produces Na1+ ions and Al atoms: AlCl3 + Na NaCl + Al. In the second reaction, however, we have just the opposite: Al atoms trying to lose their electrons to Na1+ ions. Since Al is less reactive than Na, this reaction does not occur: NaCl + Al NR (NR = no reaction). This means that only about half of all single replacement reactions actually occur; the rest are NR’s. In this lab you will be comparing four different metals: Mg, Zn, Cu and Pb, to discover which one is the most reactive, which one is second most reactive, which one is third, and which one is the least reactive. You will essentially be determining what is called an “activity series” for these four metals: an activity series is a list that ranks these metals from most reactive on top to least reactive on bottom. Procedure: 1) To begin with, most metal samples have some form of oxide coating that forms on their surface as the metal reacts with oxygen in the air. To get rid of this coating, place the four strips of metal on a piece of cardboard and then lightly sand each one, just on one side. Be careful not to sand off the symbol label. Also, the lead strip is very soft and tears easily. 2) Place the strips on a piece of wax paper in the order shown at right:

3) Then starting with the Zn(NO3)2 solution, place a tiny drop of it on each of the four strips, just below the symbol, on the top X’s shown at right.

4) Then place tiny drops of Mg(NO3)2 below the Zn(NO3)2 drops. (Make sure to leave enough room so they don’t mix with the previous drops.)

5) Then place drops of Cu(NO3)2, and finally drops of Pb(NO3)2 on the last two rows of X’s Keep in mind that these reactions are all about the metals competing with one another; the NO3

1- ions are not really part of the competition. Observe all 16 drops for a reaction. When a metal sheet gets darker it is because it is having ions of the other metal gaining electrons and turning into atoms on top of it. So, fill out the table at right describing each reaction you see, or NR if you don’t see any change at all. (Careful: one of the Zn(NO3)2 reactions in the top row is very faint and very slow, but it should be visible after a minute or two.) You should have a total of six reactions and ten NR’s. Clean up: Use a paper towel to soak up the drops, throw it away along with the wax paper, then use a second paper towel to wipe off each of the strips completely. Then make sure to wash your hands with soap. Also, place out a fresh sheet of wax paper for the next group. Follow-up questions: 1. Look at the atoms (columns in the table at right), not the ions. Which metal was most reactive? ____ 2nd? ___ 3rd? ____ least reactive? ____ 2. Which solution (row) was most reactive? ___ Which was least?____ 3. In general, the more reactive an atom is, the _______ reactive its ion is. (This should make sense: if a metal is good at losing electrons then its ion wouldn’t be very good at gaining them back.)

Zn Cu Pb x x x x

x x x x

x x x x

x x x x

Zn(NO3)2

Mg(NO3)2

Cu(NO3)2

Pb(NO3)2

Zn Mg Cu Pb

Zn2+

Mg2+

Cu2+

Pb2+

Mg

21

4. Between the two reactions mentioned in the introduction, the only one that occurred was: AlCl3 + Na NaCl + Al. Write equations for each of the six reactions that took place. _______________________________ _______________________________ _______________________________ _______________________________ _______________________________ _______________________________ 5. In the first box at right, use the four metals from the lab to construct an activity series. This is a vertical list that ranks the metals from most reactive on top to least reactive on bottom. (Use the metal atoms – not the ions for this list.) 6. Use the information below to add four more metals (K, Ni, Al and Ag) to the activity series, and rewrite it in the second (longer) box at right. A solution of KNO3 would not react with Zn or even Mg metal. Ni metal reacts with Pb(NO3)2 and Cu(NO3)2 solutions, but not with Zn(NO3)2. A piece of silver metal will not react with any of the solutions from the lab. Aluminum is better at losing its electrons than Zn, but not as good as Mg or K. 7. Look at the more extensive activity series at right. Does your answer for #6 above agree with it? _____ Explain: 8. Use the activity series at right to complete the following twelve reactions. If no reaction occurs, write NR. (That will apply to about half of them.) a) and e) are done for you. a) AgNO3 + Ca e) MgBr2 + Mn i) K2CO3 + Li b) Ca(NO3)2 + Ag f) Al + Zn(BrO3)2 j) Pb(C2H3O2)2 + Na c) Zn + FeF3 g) SnBr2 + Pt k) Mn + Mn(OH)2 d) Au + NiSO4 h) Mg + Au2S l) Cu + Ca(HCO3)2 9. Take any three reactions above (not NR’s) that are not already balanced, and balance them! _______________________________ _______________________________ _______________________________ 10. At right is the nonmetal activity series – at least one that include the four halogens. Remember that what makes a nonmetal atom reactive is its ability to gain electrons and become a negative ion. It should make sense that F is the most reactive nonmetal. Why? 11. Use the nonmetal activity series to complete the following eight reactions. If no reaction occurs, write NR. (That will apply to about half of them.) a) and d) are done for you. a) MgBr2 + I2 d) CrBr3 + F2 g) NaBr + Br2 b) AlI3 + Cl2 e) KF + Br2 h) Br2 + AuCl3 c) F2 + FeCl3 f) Cl2 + CrI3 12. Take any three reactions above (not NR’s) that are not already balanced, and balance them! _______________________________ _______________________________ _______________________________

Li Rb K Ca Na Mg Al Mn Zn Fe Ni Sn Pb Cu Ag Pt Au

F Cl Br I

22

Solubility Rules Lab Name: _________________________

Some ionic compounds are soluble in water; others are not. Two compounds that are soluble are cesium sulfide (Cs2S)

and barium bromide (BaBr2). When they dissolve in water, they make solutions that appear clear and colorless (just like

water), but when you mix these two solutions together, an instantaneous white cloud appears. This cloud is the

compound barium sulfide (BaS) which forms when the Ba2+ ions (from the BaBr2) combine with the S2- ions (from the

Cs2S). The reason it appears as a cloudy substance is because barium sulfide is not soluble in water. When two clear

solutions are mixed and they form an insoluble compound like this, we call that compound a “precipitate,” because – if

it is left long enough – the cloudy granules will eventually settle out to the bottom of the container. You may be

wondering what happened to the other two ions in the mixture described above – the Br1- ions and the Cs1+ ions. These

two ions form a compound cesium bromide (CsBr), but since

cesium bromide is soluble in water, these two ions stay

dissolved and invisible, just like they were before the reaction.

We call such ions “spectator ions.” They are not really

involved in the reaction; they just sit around “watching”

the other two react!

One way of representing this reaction is with the following diagrams:

Now, if the barium bromide above had been mixed with cesium iodide (instead of cesium sulfide), then no cloudy

precipitate would have formed – in fact, nothing would have happened. That is because potential products (the cesium

bromide and the barium iodide) are both soluble. There would be no reaction, because all the ions would simple stay

dissolved in the water just as they were before. They would all be considered spectator ions, and just as is the case with

people, when everyone is sitting around watching, nothing really gets done!

So, which compounds are soluble, and which ones are not? Does the solubility of ionic compounds follow a set of rules?

Knowing the answers to these questions would help us predict whether or not a precipitate reaction will occur when

two ionic solutions are mixed.

But… rather than just giving you the solubility rules, this lab will help you figure them out for yourself.

You will be given eleven different ionic solutions. Mix them together two at a time – one drop of each, and record

whether or not a precipitate forms. If nothing happens (no precipitate forms), then you will know that both potential

products are soluble. If a precipitate does form, then you will know that one of the potential products is insoluble – but

you won’t know which one until you do more mixings to narrow down the possibilities!

The eleven ionic solutions are listed in the column below. Next to each, write the correct charge balanced formula.

Note that many of the ions are repeated two or three times. This should be useful.

Objective: The table at right below represents forty nine different ionic compounds. For example, the upper left-hand

box represents the compound formed from Ca2+ ions and Cl1- ions – CaCl2 (calcium chloride). Likewise, the lower right

Ca2+ Na1+ Sr2+ Ag1+ NH41+ K1+ Pb2+

Ba2+ Cs1+ S2

-

Cs1+

Cs1+

Cs1+

Cs1+ Cs1+ S

2-

S2-

Br1-

Br1-

Br1-

Br1-

Br1-

Br1-

Ba2+ Ba2+

Cs1+ Cs1+

Cs1+

Cs1+ Cs1+

Cs1+

Br1-

Br1-

Br1-

Br1-

Br1-

Br1-

Ba2+

S2-

S2-

S2-

Ba2+

Ba2+

barium

bromide

solution

cesium

sulfide

solution insoluble

barium sulfide

precipitate

cesium & bromide

spectator ions

(still dissolved)

23

hand box represents PbCO3 (lead(II) carbonate). Your objective is to fill in this entire table with S’s and I’s: S’s for all the

compounds that are soluble in water and I’s for all that are insoluble. Once you have it filled out, look for similarities

and patterns that might help you establish some generalized solubility rules.

Procedure:

1) Since each of the solutions listed at left has already been dissolved in water, you know they are all soluble, so place

S’s in each appropriate box. (The one for sodium hydroxide has been done for you. You do the other ten.)

2) Take the sodium hydroxide vial (F) and squeeze a drop of it out onto a plastic tray. Then squeeze out a

drop from the ammonium chloride vial (2) on top of the first drop. As you do this BE CAREFUL NOT TO

PLACE THE TIP OF ONE VIAL IN THE DROP OF THE OTHER since that might contaminate the solution

in the vial. To avoid this, ALWAYS MAKE SURE THE DROP FALLS FROM A DISTANCE OF ABOUT 1 CM

as shown at right: Observe what happens. You should see nothing. You already knew that

sodium hydroxide and ammonium chloride were soluble; now you know that sodium chloride

and ammonium hydroxide are too, so place S’s in the two appropriate boxes above.

3) Keeping track of your observations: Because it is easy to lose track of what you have done, each of you should

document your observations: Take out a separate sheet of paper and write:“1) NaOH + NH4Cl NR”

4) Now try taking the calcium acetate solution and mixing it with the potassium carbonate solutions. This time you

should observe a distinct reaction: the formation of a light grey milky substance. (On your observation sheet write:

“Ca(C2H3O2)2 + K2CO3 grey ppt.”) That is what a precipitate looks like, and that tells you that one of the two potential

products – either potassium acetate or calcium carbonate – is insoluble, but which one? To narrow it down, try mixing

the potassium chloride solution with the calcium acetate. (Again, record what you observe.) If no precipitate forms, you

know that both potassium acetate and calcium chloride are soluble, and that would mean it would have to be the

calcium carbonate that was insoluble in the previous reaction. (If this is the case, write “I” in the box for calcium

carbonate, and write “light grey” beneath the “I” to indicate the color of this insoluble compound.) If, however,

another precipitate forms, then you may think that it was the potassium acetate that was insoluble, but that would not

be conclusive. Do you understand why?

5) Use this kind of logic to mix and match the various solutions to fill out the rest of the table. Every time you try a

combination, make sure you both record it on your observation sheet. And every time you discover that a compound is

sodium hydroxide ________

lead(II) nitrate ________

silver acetate ________

sodium iodide ________

strontium chloride ________

ammonium sulfate ________

potassium carbonate ________

sodium nitrate ________

calcium acetate ________

potassium chloride ________

ammonium chloride ________

Ca2+ Na1+ Sr2+ Ag1+ NH41+ K1+ Pb2+

Cl1-

OH1-

I1-

NO31-

C2H3O21-

SO42-

CO32-

S

24

I light grey

soluble, write “S” in the box. Whenever you discover that a compound is insoluble, write “I” in the box along with the

color or appearance of the insoluble precipitate. Try to fill out the entire table using as few mixings as possible.

Important note: One of the hydroxide compounds forms a very faint precipitate that is easy to miss.

Clean-up: Since some of the solutions you used contain heavy metal ions (lead and silver), do NOT rinse the solutions

down the sink. Instead, use a squirt bottle to rinse them into a heavy metal waste recovery jar; then wipe the tray with

a paper towel.

Follow-up Questions:

1. List the ions (pos. and neg.) for which every compound is soluble (hint, there should be 5): ___________________

2. Although rubidium ions (Rb1+) were not used in this lab, do you predict they would generally be soluble or not? _____

Explain this prediction: _________________________________________________________

3. Most chloride compounds are (soluble/insoluble). What are the exceptions to this rule? ___________________

What other negative ion forms compounds with the same solubility rule as chloride compounds? _______________

Why do you think these two groups of compounds follow the same solubility rules? ____________________________

4. Although bromide ions were not used in this lab, predict whether you think the following compounds would be S or I:

NaBr ___ CaBr2 ___ AgBr ___ PbBr2 ___ Explain these predictions: _______________________________________

5. Most sulfate compounds are (soluble/insoluble). What are the exceptions to this rule? ___________________

6. Below is a chart that summarizes many of the solubility rules including the compounds you used in the lab along with

many other. You will not need to memorize this chart, but you do need to know how to use it.

Soluble Insoluble

Rule #1: all compounds that include group I ions (Li1+, Na1+, K1+…), NH41+, NO3

1-, C2H3O21- are soluble – no exceptions

So these would all be soluble: KOH, Na2CO3, LiBr, Cu(NO3)2, (NH4)3PO4… get it?

Rule #2: All compounds that include Cl1-,Br1- & I1- are soluble except those of Pb2+, Ag1+ and Hg22+

So these would all be soluble: BaCl2, AlBr3, CsI... but these would be insoluble: AgBr, PbCl2, Hg2I2…

Rule #3: All compounds that include SO42- are soluble except those of Pb2+, Hg2

2+ Ca2+, Ba2+, Sr2+

So these would all be soluble: CuSO4, Al2(SO4)3, Ag2SO4... but these would be insoluble: PbSO4, BaSO4 …

Rule #4: All compounds that include O2-, S2-, OH1-, CO32-, CrO4

2-… are insoluble except those already covered by rule #1.

So these would all be insoluble: MgO, Cr2S3, Al(OH)3, Ag2CO3... but these would be soluble: K2CO3, Li2CrO4, (NH4)2S

Use the rules above to predict whether the following compounds would be soluble (S) or insoluble (I): (If you did this correctly, you should have had 4 S’s and 6 I’s.)

Na2SO3 __ CuS __ V(OH)3 __ CaSO4 __ PbCl2 __ Fe(NO3)3 __ NiSO4 __ Al2(CO3)3 __ (NH4)2CrO4 __ AuPO4 __

Group I1+, NH41+, NO3

1-, C2H3O21-

Cl1-, Br1-, I1-

SO42-

(no exceptions)

O2-, S2-, OH1-, CO32-, CrO4

2-, PO43-

Pb2+, Ag1+, Hg22+

Pb2+, Hg22+ Ca2+, Sr2+, Ba2+

Pb2+, Ag1+, Hg22+

Group I1+, NH41+

1)

2)

3)

4)

except

except

except

except

25

7. Note: the rules above hold true most of the time, but sometimes compounds break the rules. Check the S’s and I’s

from the table you completed on page 2. List any compounds you found to be soluble that shouldn’t be according to the

rules? _________ List any compounds you found to be insoluble that shouldn’t be according to the rules? __________

8. Net ionic equations are simply equations that show just the ions that are reacting -- leaving out any spectator ions.

For example, for the precipitate reaction described in the introduction between Cs2S and BaBr2, the net ionic equation is

simply Ba2+(aq) + S2-(aq) BaS(s). [The Br1- and Cs1+ ions are left out because they are spectator ions.] When

aluminum chloride solution – AlCl3(aq) – is mixed with lithium carbonate solution – Li2CO3(aq), the aluminum carbonate

– Al2(CO3)3 – forms an insoluble precipitate. The net ionic equation for this would be Al3+(aq) + CO3

2-(aq) Al2(CO3)3(s)

Balanced, this would be 2 Al3+(aq) + 3 CO32-(aq) Al2(CO3)3(s).

Write a balanced net ionic equation (or NR) for each of the following combinations: (The first two are done for you).

(If done correctly, there should be a total of 4 NR’s and 8 reactions. Coefficients you will need to use: 2, 2, 2, 2, 3, 3, 3)

a) KBr + Mg(C2H3O3)2

b) Pb(NO3)2 + AlBr3

c) NaOH + BaI2

d) Li3PO4 + AgC2H3O2

e) SrCl2 + (NH4)2CO3

f) CrBr3 + Sn(NO3)2

9. You are given an unknown solution to analyze and told that it contains either lead or silver ions, what solution might

you mix it with that could help you figure out which one it contains? _____ If it contained lead ions what would

happen? ______________ If it contained silver ions, what would happen? _______________________

What other solution might you mix the unknown with that could help you figure out which one it contains? ________

If it contained lead ions what would happen? _________________ If it contained silver ions, what would happen?

_______________________

10.* You are given four different solutions labeled A, B, C and D, and you are told that they correspond to solutions of

KC2H3O2, BaCl2, AgNO3 and Li2SO4 – only not in that specific order. Your objective is to try to determine which one is

which. You mix A + B and get a precipitate. What can you conclude? _______________________________________

What will you observe when you mix C + D: NR or ppt? How do you know? _________________________________

You mix all the other combinations and get: A+C NR, A+DNR, B+CNR, and B+Dprecipitate.

What specific match-ups can you be certain of? __________________________________________________________

You then take solutions A and C and mix them both with some Na2CO3 solution and neither one reacts with it.

Identify all four solutions: A = _________ B = _________ C = _________ D = _________

g) CuBr + Li2CrO4

h) K2S + (NH4)3PO4

i) K2S + Al2(SO4)3

j) VBr3 + KOH

k) Na2SO4 + AgNO3

l) FeCl3 + AgC2H3O2

26

Single Replacement WS Name: _______________________________

Using the activity series at the right, complete the following reactions, or write “NR” for the reactions that wouldn’t

occur. Then balance each reaction. (5 NR’s, coefficients: 2,2,2,2,2,2,2,2,2,2,2,2,2,2,2, 3,3,3,3,3,3,3,3,3,3,3, 4,4, 6)

1) Cu + NiBr3 9) Fe(NO3)3 + Zn

2) Al(NO3)3 + Ca 10) Br2 + FeI3

3) MgBr2 + I2 11) MgCO3 + Mg

4) AgC2H3O2 + Zn 12) Ni2(SO4)3 + Ca

5) F2 + KCl 13) Pb(NO3)4 + Al

6) Na + Mn2(SO4)3 14) AlF3 + F2

7) FeF2 + Pt 15) PtI2 + Cl2

8) NaI + Cl2 16) Al + AuBr3

Double Replacement Reactions WS

Read the next page; then write balanced net ionic equations (or NR) for each of the following. Include (s), (aq), (g), etc.

(There are 6 NR’s)

1) AgNO3 + NiBr3 13) Zn(NO3)3 + H2SO4

2) Al(NO3)3 + HBr 14) Zn(NO3)3 + KOH

3) Mg(OH)2 + HI 15) Na2CO3 + AlCl3

4) K2SO3 + HNO3 16) Ni2(SO4)3 + BaBr2

5) Ca(NO2)2 + HCl 17) HNO3 + Ca(C2H3O2)2

6) NaBr + Mn2(SO4)3 18) NaOH + HCl

7) FeBr3 + Pb(NO3)2 19) CaCl2 + Li3PO4

8) NaHCO3 + H2SO4 20) (NH4)3P + CsOH

9) NaOH + (NH4)2SO4 21) AlF3 + HBr

10) NH4I + AgC2H3O2 22) CuCl2 + MgSO4

11) NaI + CaCl2 23) LiHSO3 + HCl

12) Li2SO3 + HClO4 24) KOH + Mn(NO3)2

Li Rb K Ca Na Mg Al Mn Zn Fe Ni Sn Pb Cu Ag Pt Au

F Cl Br I

27

H+ H+

H+

H+

H+

H+ H+ H+ Br- Br

- Br-

Br- Br

-

Br-

Br-

Br-

K+ K+

K+ K+

K+

K+

K+ K+

Cl-

Cl-

Cl-

Cl-

Cl-

Cl-

Cl-

Cl-

H+ H+

H+

H

H+

H H H

+ F- F

- F-

F- F

F-

F

F

K+ K+

K+ K+

K+

K+

K+ K+

Cl-

Cl-

Cl-

Cl-

Cl-

Cl-

Cl-

Cl-

As you learned in the Solubility Rules Lab, for double replacement reactions, if both possible products are soluble, then

no reaction occurs. If one of the possible products is insoluble, however, then a precipitate will form and the mixture

will turn instantly cloudy (and sometimes colored).

As it turns out, formation of a precipitate is only one factor that can cause a double replacement reaction to occur.

There are two others:

The second factor that can cause a DR reaction to occur is the formation of a weak acid. Acids are compounds

comprised of H+ ions bonded to negative ions: HF, HNO3, H2SO4, H3PO4… These acids are all soluble in water, and when

they dissolve, they dissociate to a greater or lesser extent into ions: For example HNO2 dissociates into H+ ions and NO2

-

ions [ HNO2(aq) H+(aq) + NO2

- (aq)]. Likewise HClO4 dissociates into H+ ions and ClO4

- ions [ HClO4(aq) H+(aq) +

ClO4- (aq)]. The difference is that most acids dissociate like this only to a small extent – about 5% or less. They are all

considered weak acids. There are however six acids that dissociate like this 100%: HCl, HBr, HI, HNO3, HClO4 and H2SO4.

These are the six strong acids. If a double replacement reaction forms one of these six strong acids, then that is like

forming a soluble ionic compound: it does not cause any reaction to happen. But when a double replacement reaction

forms a weak acid (any acid that is not in the list above), then a reaction will happen. These weak acids that form are

still dissolved, but they are no longer dissociated to the extent they were before. Although you would not notice the

solution getting cloudier, there is still a chemical reaction taking place which is turning ions into neutral molecules:

HCl(aq) + KBr (aq) no reaction HCl(aq) + KF (aq) reaction

Net ionic eq: NR Net ionic eq: H+(aq) + F-(aq) HF(aq)

Notice in the first reaction above, since a strong acid (HBr) is formed, all the ions simply stay dissolved and dissociated,

so nothing happens – only mixing, and that never counts as a chemical reaction. In the second reaction above, however,

a weak acid (HF) is formed. See how most of the HF in the final beaker is bonded into neutral molecules? This is the

reaction: H+ ions and F- ions coming together and forming neutral molecules of HF.

The third factor that can cause a DR reaction to occur is the formation of a compound that decomposes spontaneously.

There are three of these compounds that you need to be familiar with: H2CO3, H2SO3 and NH4OH. Each of these

decomposes into water (H2O) and a gas:

H2CO3(aq) H2O(l) + CO2(g) H2SO3 H2O(l) + SO2(g) NH4OH(aq) H2O(l) + NH3(g) H2CO3, for

example, forms when H+ ions react with carbonate (CO32-) ions: 2 H+(aq) + CO3

2-(aq) H2CO3(aq) H2O(l) + CO2(g)

or when H+ ions react with bicarbonate (HCO3-) ions: H+(aq) + HCO3

-(aq) H2CO3(aq) H2O(l) + CO2(g).

So here are some sample problems:

NaBr + Pb(NO3)2 (because PbBr2 is insoluble in water – check the solubility rules)

KNO2 + HCl (because HNO2 is a weak acid – it’s not one of the six strong ones listed above)

NH4Br + NaOH (because NH4OH is one of the 3 that decomposes spontaneously)

K2SO4 + CuBr2 (because neither KBr nor CuSO4 is insoluble, forms a weak acid or decomposes spontaneously)

28

Mole Tutorial Worksheet Name: ___________________

1. How much does one mole of He weigh? ________

2. How much does one atom of He weigh? ________

3. How much does one atom of Ca weigh? ________

4. How much does one mole of Ca weigh? ________

5. How many atoms are present in one mole of He? ________

6. How many atoms are present in one mole of Ca? ________

7. How much would one mole of CO2 weigh? ________

8. What would one molecule of C2H6 weigh? _________

9. What would one mole of (NH4)2S weigh? _________

10. What would one molecule of NF3 weigh? _________

11. How many molecules are there in one mole of SO2? _________

12. How much would one mole of SO2 weigh? ________

13. How much would one mole of Ca(NO3)2 weigh? __________

14. How much would one atom of Co weigh? ___________

15. How much would 6.02 x 1023 atoms of Co weigh? _________

16. How much would 6.02 x 1023 molecule of CS2 weigh? ____________

17. How much would one molecule of CS2 weigh? _________

18. How much would one mole of C6H12O6 weigh? ___________

Would C6H12O6 be made up of atoms or molecules? ___________

And how many of them would there be in that one mole sample? __________

19. How much would one mole of Ne weigh? __________

Would Ne be made up of atoms or molecules? ___________


20. How much would one mole of F2 weigh? __________

Would F2 be made up of atoms or molecules? ___________


In the space below, write down everything you just learned about the mole:

Ans (IRO) 4.0026 4.0026 20.18 30.0 38.0 40.08 40.08 44.0 58.9 58.9 64.0 64.0 68.1 71.0 76.1 76.1 164.1 164 180.0 6.02x1023 6.02x1023 6.02x1023 6.02x1023 6.02x1023 6.02x1023 molecules molecules atoms atoms units: g g g g g g g g g g g amu amu amu amu amu amu

29

6 x 6 mystery solutions lab Name: _______________ partner: __________________

This lab has you again mixing and matching drops of solutions, but this time, rather than knowing before-hand what the

solutions are and then seeing how they react (or not), this time you will be mixing them and observing what happens, and then using that information to identify the solutions!

One of you will be given a set of six different unknown solutions labeled A, B, C, D, E and F. The other will be given a corresponding set of six solutions labeled 1, 2, 3, 4, 5 and 6 – but in scrambled order. In part I of this lab, you and your

partner need to figure out which numbered solution matches which lettered solution (Perhaps 1 = B, 2 = F…). For this you really do not need any chemistry know-how, just good logic and problem solving skills.

Procedure: The one with the lettered solutions should use only those solutions, none of the numbered ones, and vice-

versa. Each person should work quickly quietly and very carefully, collecting his/her data and recording it in the space

below. The format you use to record that data is up to you. You may want to discuss this with your partner before you even start.

Once you have collected all the data you need, then clean up and wash your hands, then go sit down and figure out

which letter corresponds with which number. Complete the slip in the lower left-hand corner of this page and turn it in to

the teacher, just one slip per pair. As you do this, the teacher will hand you a list of six compounds. They are the identities of these six solutions, but again they will be in scrambled order. Part II of this lab requires you to use your

understanding of what causes double replacement reactions to occur to figure out which compound corresponds to which number. Complete the slip in the lower right-hand corner of this page and the turn that in to the teacher. Then answer

the follow–up questions on the back side. The two slips must be turned in during class time. What you don’t finish of the follow-up questions can be finished for homework.

Names: ________________________

________________________

Results for part I (fill in the blanks with

the correct letters A-F)

1 = __

2 = __

3 = __

4 = __

5 = __

6 = __

Names: ________________________ ________________________

Results for part II (fill in the blanks with the correct chemical formulas)

1 = ________

2 = _________

3 = _________

4 = _________

5 = _________

6 = _________

30

Follow-up questions:

1. What is the name of the compound in the lab that produced the most precipitates? ____________________________

Write net ionic equations for each reaction this compound was involved in? (omit any repeats)

2. One reaction you observed produced a compound which decomposed spontaneously. What are the names of the two

compounds that were involved in this reaction? _____________________________ ____________________________

Write the balanced net ionic equation for this reaction: __________________________________________

3. Two reactions occurred but produced no evidence of a reaction. Write their net ionic equations for these two

reactions: _______________________________________ ________________________________________

Why were you unable to see any signs of these reactions?

4. A seventh solution (X) was mixed with each of the six solutions in the lab and

the only reaction was a bubbling with one of the solutions. What might be a possible identity of X? ________

5. An eighth solution (Y) was mixed with each of the six solutions in the lab and it produced a

precipitate with one of the solutions and a pungent odor with another. What might be a possible identity of Y? ________

6. A ninth solution (Z) was mixed with each of the six solutions in the lab and it

produced no reactions of any kind – visible or not. What might be a possible identity of Z? ________

7. A similar lab, involving only three lettered solutions (J, K & L) and three corresponding numbered solutions (10, 11 & 12) was performed by Joey and Alice. Joey observed that J + K produced no reaction at all, and K + L produced a red

precipitate. Alice observed that 10 + 11 gave off a stinky smell. Then time ran out before they could make any more

observations. They figured that all was lost, and that there was no way they could possibly figure out what the three match-ups were. But then Alice, desperate for any points she could get, realized that at least they could make one

definite match. What was that one match-up which she made?

____ = ____ Explain how she knew that match was correct:

31

Moles WS: Name: _______________________ Remember: 1 mole = 6.022 x 1023 atoms (or molecules) and 1 mole = X g (where X comes from the masses in the PT)

1. How much would 3.27 moles of C weigh? _______ 2. What would be the mass of 0.390 moles of Ca? _________ 3. What would 6.95 moles of CO2 weigh? _________ 4. How many atoms are present in 2.50 moles of neon? _________ 5. How many molecules are there in 0.075 moles of water? _________ 6. How many moles are there in a 34.6 g sample of Cu? ________ 7. How many moles are there in a 34.6 g sample of sodium sulfate? _________ 8. A sample of aluminum contains 4.58 x 1025 atoms. How many moles is that? _________ 9. A flask contains 7.98 x 1021 molecules of dinitrogen monoxide. How many moles is that? _________ 10. How many grams of iron would it take to make 5.8 moles? ________ 11. How many grams of oxygen would it take to make 5.8 moles? ________ 12. How much would 7.8 x 1024 atoms of helium weigh? ________ 13. How many atoms would be present in a sample of calcium weighing 3.78 g? _______

Ans: (IRO+1) 0.0133

0.244 0.544

8.95 15.6 39.2

52 76.1 190 306 320

4.5 x 1022

5.68 x 10

22

1.51x1024

units: (IRO+1) g g g g g g

moles moles moles moles moles cules atoms atoms

6.022 x 1023

32

14. How many moles are there in 1790 g of C4H10? ________ 15. What is the mass of 0.0150 moles of carbon tetrafluoride? ________ 16. How many moles are there in a 79.6 g sample of CH4? ________ 17. How many molecules are there in a 114.2 g sample of H2S? _________ 18. How many atoms of Bi would it take to make 5.8 moles? ________ 19. A ring contains 3.98 x 1022 atoms of Au. How much would it weigh? _________ 20. A sample of O2 contains 4.58 x 10

25 molecules. How many moles is that? _________ 21. How many molecules are there in 7.8 x 103 moles of CO? ________ 22. How many atoms would be present in a sample of Fe weighing 539 g? _______ 23. How many molecules are there in 2490 g of C4H10? ________ 24. Propane is C3H8. A propane torch weighs 957.32 g. It is used and 6.28 x 10

22 molecules of propane are burned. How much does the torch now weigh? ________ 25. Aluminum’s density is 2.70 g/mL. How many atoms would be present in a 34.7 mL chunk of aluminum? _______ 26. What would be the height of a cube of pure ice (D=0.9167 g/mL) containing 2.50 x 1025 molecules? ________

Ans: (IRO+1) 1.32 4.96 9.34 13.0 30.8 62.4 76.1

952.72 2.018 x 10

24

2.09 x 1024

3.5 x 10

24

5.81 x 1024

2.58 x 10

25

4.7 x 1027

units: (IRO+1) g g g g cm

moles moles moles cules cules cules

atoms atoms atoms

33

PERCENT COMPOSITION WORKSHEET Name: ____________________ SHOW WORK DETERMINING PERCENT BY MASS OF EACH ELEMENT IN A COMPOUND 1. Determine (to 3 sig figs) the percent composition for each element in the following substances: a) FeO b) Fe2O3 c) Na2CO3 d) Mg(NO3)2 e)C4H10 f) C8H20 g) N2 2. a) What is the percent silver in Ag2SO4? b) What is the percent phosphorus in Ca3(PO4)2? Ans: ______ Ans: ______ c) What is the % aluminum in aluminum oxide? d) What is the percent sodium in sodium borate? (Hint: write the correct formula first) Ans: ______ Ans: ______ USING PERCENT COMPOSITION 4. How many grams of iron would be present in 68.5 g of Fe2O3? Ans: ______

5. How much calcium could be extracted from 185 g of CaBr2?

Ans: ______

6. If you had 10.0 g of magnesium, how much MgCO3 could you make?

Ans: ______

7. What mass of Ca(NO3)2 would need to be decomposed to produce 14.6 g of calcium?

Ans: ______

8. What mass of aluminum could be extracted from 406 g of Al2(CO3)3?

Ans: ______

9. How many grams of cobalt(III) sulfide could be produced from 39.4 g of cobalt?

Ans: ______

10. What mass of phosphorus is needed to make 45.8 g of diphosphorus trioxide?

Ans: _____

(cont on back side)

Ans: 11.3 16.4 17.4 17.4 18.9 20.0 22.3 25.8 27.8 30.1 34.7 37.1 43.4 45.3 47.9 52.9 54.0 59.8 64.7 69.2 69.9 71.6 77.7 82.6 82.6 93.7 100.0 units: 7(g) 19 (%)

34

USING PERCENT COMPOSITION TO FIND EMPIRICAL FORMULAS:

11. A copper oxide compound is 88.82% copper, 11.18% oxygen by mass. Empirical formula = ________

name= __________________

12. A carbon fluoride compound is 13.64% carbon (and the rest is fluorine). Emp. Form. = ________

name= _________________

13. A compound is 16.39% Mg, 18.88%N and the rest O. Emp. form = _________ name: _____________________

14. A compound is 43.64% phosphorus and the rest oxygen. Emp. form = _________ name: _____________________

15. A compound is 89.14%gold and the rest oxygen. Emp. form = _________ name: _____________________

16. A compound is 36.62% Cr, 12.68% C and the rest O. Emp. form = _________ name: _____________________

17. A compound is 39.13% C, 8.69% H and the rest O. Emp. form = _________ name: (google it!)_____________________

18. Tetramethylcyclobutadiene has an emp. formula of C2H3. Its molar mass is 108 g/mol. Molecular formula = _______

19. Glucose has an emp. formula of CH2O and 4.27 moles of it weighs 769 g. Mol form = _________

20. Difluorobutane has an emp. formula of C2H4F and 4.19 x 1022 molecules of it weigh 6.55 g. Mol form = _________

21. Diethylether has an emp. formula of C4H10O. A 1.00 g sample contains 8.14 x 1021 molecules. Mol form= _________

22. Trinitrobenzene has an emp. Formula of C2HNO2. One molecule of it weighs 213 amu. Mol form = __________

Answer subscripts (other than 1’s): 2 2 2 2 2 2 3 3 3 3 3 3 4 4 4 5 6 6 6 6 6 8 8 8 9 10 12 12

35

Stoichiometry Helpful Hints! 1. How to recognize a stoichiometry question. A stoichiometry problem is one that requires you to change from one substance to another. In this problem: “How many molecules are there in 34.5 g of CO2?” There is no changing of substances – it’s all just CO2, nothing else – so this is not a stoichiometry problem, it’s just a regular mole problem, and you would solve it like this: In this problem: “How many molecules of Br2 would react would 34.5 g of Al?” We are definitely changing substances: Our given is in grams of Al, and the question asks about molecules of Br2. To change substances, you must have a balanced equation: (2 Al + 3 Br2 2 AlBr3), and you must make the change from one substance to another in moles (because that’s what the equation recipe is written in). This problem would be solved like this: 2. Steps to solving a stoichiometry problem: (Steps 1 & 3 may or may not be needed; steps 0 and 2 are always needed) 0. Always start by writing down your given: number, units and substance! For example 34.5 g Al 1. If your given is not in moles, change it to moles. In this step, you will always be putting a “1” in front of “mole.” (If you’re given is already in moles, go straight to step 2.) 2. The stoich step! Change moles of what you’re given into moles of what you want. It’s really quite simple: moles of what you’re given go on bottom, moles of what you want goes on top & the numbers you put in front of them are just the coefficients from the balanced equation. 3. If the question just asks for moles, you are done. If it asks for grams or atoms or molecules, then you have a finishing step: put moles of the new stuff on bottom, and g (or atoms or molecules) of the new stuff on top. Like in step 1, this step will always have a “1” in front of “mole.” 3. Things to look for: Every stoichiometry problem will require the mole to mole step (step 2). Any other steps will have “mole” somewhere in it (either top or bottom) with a “1” in front of the “mole”, and something other than a “1” in front of the other unit. You cannot change substances and units in the same step. For example, you cannot go from grams of Br2 to mol of Al. If you change the units, you must keep the substance the same: grams of Br2 to mol of Br2 is OK. If you change the substance, you must keep the units the same, and those units must be moles: mol of Al to mol of Br2 is OK.

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