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Transcript
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 Na O (#8) (#11) (#20) (#9) (#54) (#49) g) argon (#18) h) lithium (#3) i) sulfur (#16) j) aluminum (#13) k) bromine (#35) l) barium (#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 lose 1 ___ 1+ _______ h) lithium gain 2 ___ 2_______ g) argon _______ ___ i) sulfur _______ ___ j) aluminum _______ ___ k) bromine _______ ___ l) barium _______ ___ _______ ___ _______ ___ _______ ___ _______ ___ _______ ___ 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 oxide ion _____________ g) argon sodium ion _____________ h) lithium _____________ i) sulfur ____________ j) aluminum ____________ k) bromine ____________ l) barium _____________ _____________ _____________ ____________ ____________ ____________ 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: _________________ sodium oxide Na Na O [Na]1+ [ O ]2- 1+ Na2O [Na] c) lithium & fluorine name: __________________ d) sodium & nitrogen name: __________________ e) magnesium & sulfur name: ___________________ f) aluminum & fluorine name: __________________ g) aluminum & phosphorus name: _______________ h) barium & phosphorus name: __________________ 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! Positive Ions (Cations) Fixed-Charge 1+ Na1+ sodium 1+ K potassium 1+ H hydrogen Ag 1+ silver 1+ NH4 ammonium Positive Ions (Cations) Variable-Charge Cu1+ Cu2+ copper(I) copper(II) Fe2+ Fe3+ iron(II) iron(III) Co2+ Co3+ cobalt(II) cobalt(III) Pb2+ Pb4+ lead(II) lead(IV) Negative ions (Anions) Monatomic 1F1fluoride 1Br bromide 1Cl chloride I1iodide 2- 2+ 2+ Mg Ca2+ Zn2+ magnesium calcium zinc O2S23- N P3- nitride phosphide Positive Ions (Cations) Fixed-Charge 1+ Li1+ lithium 1+ Rb rubidium 1+ Cs cesium Fr 1+ francium 2+ Be2+ Sr2+ Ba2+ Ra2+ Cd2+ beryllium strontium barium radium cadmium 3+ Sc B = boron CO32SO42CrO42- carbonate sulfate chromate aluminum Also know the following prefixes & suffixes: hypo- -ite, -ite, -ate, per- -ate, bi-, and dihydrogen- 3+ (AKA: hydrogen carbonate) 23- 3+ Al3+ oxide sulfide Negative ions (Anions) Polyatomic 1OH1hydroxide 1NO3 nitrate 1ClO3 chlorate BrO31bromate 1IO3 iodate 1C2H3O2 acetate HCO31- bicarbonate scandium Positive Ions (Cations) Variable-Charge Au1+ Au3+ gold(I) gold(III) Cr2+ Cr3+ chromium(II) chromium(III) Mn2+ Mn3+ manganese(II) manganese(III) Ni2+ Ni3+ nickel(II) nickel(III) 2+ Hg2 Hg2+ mercury(I) mercury(II) Sn2+ Sn4+ tin(II) tin(IV) Negative ions (Anions) Monatomic 1H1hydride 22Se selenide Te2telluride 3As3- arsenide 3PO4 3- phosphate Negative ions (Anions) Polyatomic 1CN1cyanide 1CNO cyanate 1SCN thiocyanate HCO21- formate MnO41- permanganate NH21amide N31azide 1Al(OH)4 aluminate 2- Also: here are the nonmetals (and some metalloids): C = carbon N = nitrogen O = oxygen F= fluorine Si = silicon P = phosphorus S = sulfur Cl = chlorine As = arsenic Se = selenium Br = bromine Te = tellurium I = iodine At = astatine 2 Cr2O72S2O32C2O42MoO42SiO32C4H4O62Ne = neon Ar = argon Kr = krypton Xe = xenon Rn = radon dichromate thiosulfate oxalate molybdate silicate tartrate 3AsO43BO33- arsenate borate 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-, SO42- 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 S ____________ 80. Ag2CO3 silver carbonate ______________________________________ 50. potassium sulfide K2 S ____________ 81. K2O potassium oxide ______________________________________ 51. copper(I) oxide Cu2O ____________ 82. Mg2+ magnesium ion ______________________________________ 52. aluminum Al ____________ 83. Cu2+ copper(II) ion ______________________________________ 53. aluminum ion Al3+ ____________ 84. Cu copper ______________________________________ 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 ______________________________________ ____________ 4+ 95. Pb ______________________________________ 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______________________________________ 64. barium chloride ____________ 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 ______________________________________ 4 Covalent Bonding Worksheet 1. Explain how covalent bonding occurs: Name: ___________________ 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. oxygen difluoride a) oxygen & fluorine name: _________________ F O FO F hydrogen monobromide b) hydrogen & bromine name: _________________ H OF2 Br H Br HBr F 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: _________________ 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 ______________________________________ 54. Hg2+ ______________________________________ 42. cadmium ion ____________ 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 _________________________________ NO21- 15. iron(III) sulfide _______ 45. 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 _________________________________ 56. Cu _________________________________ 26. ammonium chromate 27. xenon _______ _______ _________________________________ _________________________________ _________________________________ 1+ 57. Cu _________________________________ 2+ 28. hypoiodous acid _______ 58. Mg _________________________________ 29. phosphorus pentafluoride _______ 59. Be(NH2)2 _________________________________ 30. nitric acid _______ 60. HBrO4(aq) _________________________________ 8 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. 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 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 F F 1. F2 NH3 8. Br2 2. BF3 9. CO 3. H2O 10. SO32- 4. BrF3 11. CH4 5. SF6 12. OF2 6. CO2 13. SF4 (Write small!) 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 H C N 19. NH41+ 26. CH2O 20. CH2F2 27. C2H2 O H C H H C C H 10 Molecular Shapes Note Sheet Hybrid. Name: ______________________ 1. Learn these thirteen molecular shapes; know them by heart – their names, their bond angles and how to draw them. ex AXE 3-D sketch of molecular shape Name of shape b.a. 2. 3. 4. 5. 6. 7. 8. 9. 11. 12. 13. 10. 14. 11 b.a. = bond angle (not something else) 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 bonded atoms 0 NEPs 120* ba 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 4 bonded atoms 2 NEPs _____ ba __ bonded atoms __ NEPs ________ ba 6 bonded atoms __ NEPs 90* ba trigonal planar __ bonded atoms __ NEPs _____ ba trigonal pyramidal 2 bonded atoms 1 NEP __ _____ ba see-saw __ bonded atoms __ NEPs ______ ba __ bonded atoms __ NEPs ______ ba T-shaped 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. F F FCF F C F F a. CF4 F b. SF2 c. BF3 F FBF FS F tetrahedral 12 d. SO2 e. BeBr2 OS O Br Be Br g. SF4 g. NH41+ F Fs F F f. PH3 H HPH i. CS2 H 1+ [H N H] H j. CH4 S C S k. XeF2 H HCH H i. H2O H OH F Xe 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. F F FCF F C F F a. CF4 F d. CO2 b. SeF2 c. BeF2 e. SO3 f. SF6 tetrahedral write small g. PH5 h. OF2 i. BrF41+ j. NH3 k. SO32- l. SiH4 13 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. 1. + H O + H 2. PB NPB H H C H PB NPB - H water (H2O) 6. PB NPB Br Br methane (CH4) 7. bromine (Br2) 11. PB NPB Cl Cl B O C O PB NPB PB NPB O S O 8. PB NPB 4. F HH 16. PB NPB H C C O H PB NPB HH C O Cl H 21. Cl P Cl PB PB F C H Cl NPB NPB H Cl phosphorus pentachloride (PCl ) 20. 5 H Br hydrogen bromide (HBr) 9. 10. PB NPB PB NPB H N H H 13. HH 14. PB NPB H C C C H PB NPB HH HH F N F F nitrogen trifluoride (NF3) ammonia (NH3) O F 17. H C C C H PB F S F NPB F F HH HH F acetone (C3H6O) ethanol (C2H5OH) PB NPB hydrogen (H2) F HH HH HH HH H C C C C C C C C H HH HH HH HH propane (C3H8) 15. 5. H H PB NPB carbon monoxide (CO) sulfur dioxide (SO2) boron trichloride (BCl3) F F C 4.0 3.5 3.0 3.0 2.8 2.5 2.5 2.5 2.1 2.1 2.0 carbon tetrafluoride (CF4) carbon dioxide (CO2) 12. Cl 3. PB NPB F O N Cl Br C S I H P B octane (C8H18) 18. PB NPB F F S F F 19. PB NPB H C N hydrogen cyanide (HCN) sulfur hexafluoride (SF6) sulfur tetrafluoride (SF4) 22. PB NPB H H S 23. PB NPB fluoromethane (CH3F) hydrogen sulfide (H2S) O H C H 24. PB NPB formaldehyde (CH2O) F F I F F F iodine pentafluoride (IF5) 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 H toward one side of the molecule. These molecules therefore have a - side and a + side. Such moleO cules are called “polar molecules” (they have a dipole moment). Find all the molecules above that are + H 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. + - 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) polar nonpolar H FC F H 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.) 15 SF3Br3 S S SF4Br2 S S SF2Br2Cl2 S S SF6 S S 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. 2. 3. N2(g) + 3 H2(g) + 2 H2O(l) 2 NH4OH(aq) Mg(s) + 2 H2O(l) Mg(OH)2(aq) + H2(g) 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 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 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 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 ___ C’s ___ H’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 wholenumbers). 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 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 19 #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 ________. 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 __ 40. Al + Br2 __ 27. Na + F2 __ 41. Al + CaCl2 __ 28. AgF + CaCl2 __ 42. NBr3 __ 29. C2H4 + O2 __ 43. AlBr3 + CaO __ 30. K2S __ 44. NH3 __ 31. O2 + Mg __ 45. C5H10O4 + O2 __ 32. F2 + AlBr3 __ 46. CaO + HBr __ 33. C2H6O + O2 __ 47. I2 + Mg __ 34. Li2S + MgCl2 __ 48. NaF + CaBr2 __ 35. HCl + Zn __ 49. K2SO4 + Ca(NO3)2 __ 36. BaBr2 __ 50. CuI2 __ 37. Na2O __ 51. Na3N + Ca __ 38. O2 + C6H12 __ 52. Na3N + Cl2 __ 39. S8 + Na 20 __ 53. O2 + C2H2 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. Mg Cu 2) Place the strips on a piece of wax paper in the order shown at right: Zn Pb Zn(NO3)2 x x x x Mg(NO3)2 x x x x Cu(NO3)2 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.) Pb(NO ) x x x x x x x x 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. 3 2 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 NO31- 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 Zn2+ completely. Then make sure to wash your hands with soap. Also, place out a fresh sheet of wax paper for the next group. Zn Mg Cu Pb Mg2+ 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? ____ Cu2+ 2. Which solution (row) was most reactive? ___ Which was least?____ Pb2+ 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.) 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! _______________________________ Li Rb K Ca Na Mg Al Mn Zn Fe Ni Sn Pb Cu Ag Pt Au _______________________________ _______________________________ 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! _______________________________ _______________________________ _______________________________ 22 F Cl Br I 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: Br1- cesium sulfide solution Cs1+ Cs1+ S 2- Cs1+ Cs1+ Cs1+ S Cs1+ S Ba2+ Br1- Br1- barium bromide solution Br1- Br1- 2- Cs1+ Ba2+ 2 Ba2+ Br1- Br1- Br1- Cs1+ Cs1+ Cs1+ Br1Cs1+ Br1- Br1- Cs1+ - Ba2+ S2 Ba2+ - S2 Ba2+ S2 cesium & bromide spectator ions (still dissolved) Br1- insoluble barium sulfide precipitate 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 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. Ca2+ sodium hydroxide ________ lead(II) nitrate ________ silver acetate ________ sodium iodide ________ strontium chloride ________ ammonium sulfate ________ potassium carbonate ________ sodium nitrate ________ C2H3O21- calcium acetate ________ SO42- potassium chloride ________ ammonium chloride ________ Na1+ Sr2+ Ag1+ NH41+ K1+ Pb2+ Cl1- S OH1I1NO31- CO32- I light grey 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 24 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 except 1) Group I1+, NH41+, NO31-, C2H3O21- (no exceptions) 2) Cl1-, Br1-, I1- except Pb2+, Ag1+, Hg22+ 3) SO42- except Pb2+, Hg22+ Ca2+, Sr2+, Ba2+ 4) Group I1+, NH41+ except O2-, S2-, OH1-, CO32-, CrO42-, PO43- 1+ 1+ 1+ NH41+, NO31-, 1+ 2+ Pb2+, Ag 1- , Hg2 Rule #1: all compounds that include group I ions (Li , Na , K …), C2H3O2 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+, Hg22+ 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-, CrO42-… 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 __ 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) + CO32-(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 g) CuBr + Li2CrO4 b) Pb(NO3)2 + AlBr3 h) K2S + (NH4)3PO4 c) NaOH + BaI2 i) K2S + Al2(SO4)3 d) Li3PO4 + AgC2H3O2 j) VBr3 + KOH e) SrCl2 + (NH4)2CO3 k) Na2SO4 + AgNO3 f) CrBr3 + Sn(NO3)2 l) FeCl3 + AgC2H3O2 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 = _________ 26 D = _________ Single Replacement WS Name: _______________________________ Using the activity series at the right, complete the following reactions, or write “NR” for the reactions that wouldn’t Li 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 Ca 2) Al(NO3)3 + 3) MgBr2 + 9) I2 4) AgC2H3O2 + 5) F2 + 6) Na + 7) FeF2 + 8) NaI + Zn KCl Fe(NO3)3 + Zn FeI3 10) Br2 + 11) MgCO3 + 12) Ni2(SO4)3 + Mg Al 13) Pb(NO3)4 + Mn2(SO4)3 Ca 14) AlF3 + F2 Pt 15) PtI2 + Cl2 Cl2 16) Al + AuBr3 Rb K Ca Na Mg Al Mn Zn Fe Ni Sn Pb Cu Ag Pt Au F Cl Br I 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 27 24) KOH + Mn(NO3)2 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 NO2ions [ 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: H+ H Cl- + ClH H+ Br- H+ + Cl- K+ Cl- HCl(aq) + K+ K BrBr- + Br Br- ClK H+ K+ KBr (aq) - K+ K+ Br Cl- H+ H+ BrK+ + - H+ Cl- H Cl- + Cl- ClH Br- no reaction H+ F- H+ + Cl- K+ Cl- K+ K F- F - ClK H F- F- + H K F + Cl- F H K+ HCl(aq) + KF (aq) + + Cl- K+ K+ Cl- F reaction - Net ionic eq: H (aq) + F (aq) HF(aq) Net ionic eq: NR 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 + 2+ 2example, forms when H ions react with carbonate (CO3 ) ions: 2 H (aq) + CO3 (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 KNO2 + HCl NH4Br + NaOH K2SO4 + CuBr2 (because PbBr2 is insoluble in water – check the solubility rules) (because HNO2 is a weak acid – it’s not one of the six strong ones listed above) (because NH4OH is one of the 3 that decomposes spontaneously) (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? ___________ And how many of them would there be in that one mole sample? __________ 20. How much would one mole of F2 weigh? __________ Would F2 be made up of atoms or molecules? ___________ And how many of them would there be in that one mole sample? __________ In the space below, write down everything you just learned about the mole: 29 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 gggggg amu amu amu amu amu amu 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 viceversa. 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: ________________________ Names: ________________________ ________________________ ________________________ Results for part II (fill in the blanks with the correct chemical formulas) Results for part I (fill in the blanks with the correct letters A-F) 1 = __ 2 = __ 3 = __ 4 = __ 5 = __ 6 = __ 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) 6.022 x 1023 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? _______ 32 Ans: (IRO+1) 0.0133 0.244 0.544 8.95 15.6 39.2 52 76.1 190 306 320 22 4.5 x 10 22 5.68 x 10 24 1.51x10 units: (IRO+1) ggggg g moles moles moles moles moles cules atoms atoms 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 1025 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? _______ Ans: (IRO+1) 1.32 4.96 9.34 13.0 30.8 62.4 76.1 952.72 24 2.018 x 10 24 2.09 x 10 24 3.5 x 10 24 5.81 x 10 25 2.58 x 10 27 4.7 x 10 units: (IRO+1) g g g g cm moles moles moles cules cules cules atoms atoms atoms 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 1022 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? ________ 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 d) Mg(NO3)2 b) Fe2O3 c) Na2CO3 e)C4H10 2. a) What is the percent silver in Ag2SO4? f) C8H20 g) N2 b) What is the percent phosphorus in Ca3(PO4)2? Ans: ______ c) What is the % aluminum in aluminum oxide? (Hint: write the correct formula first) Ans: ______ d) What is the percent sodium in sodium borate? 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) 34 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 (%) 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: _____________________ 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 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 = __________ 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 Br 2. 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 Br 2 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. 36 STOICHIOMETRY WORKSHEET Name: _____________________________ SHOW ALL WORK IN FACTOR LABEL FORM ALL THE WAY THROUGH! 1. 4 Na + O2 2 Na2O Ans: (IRO+1) a) How many moles of sodium (Na) would be needed to react with 3.82 moles of oxygen (O2)? Ans: ______ b) How many moles of Na2O can be produced from 13.5 moles of Na? Ans: ______ c) How many moles of O2 are needed to produce 5.62 moles of Na2O? Ans: ______ d) What mass of Na would react with 0.785 moles of O2? Ans: ______ e) How many moles of O2 are needed to produce 34.7 g of Na2O? Ans: ______ 2. C2H4 + 3 O2 2 CO2 + 2 H2O a) When 0.624 moles of O2 are reacted, how many moles of carbon dioxide are produced? Ans: ______ b) How many grams of ethene (C2H4) are needed to produce 3.7 moles of water? Ans: ______ c) How many grams of oxygen (remember, that’s O2) are needed to react with 2.56 g of C2H4? Ans: ______ d) How many molecules of carbon dioxide are produced along with 3.7 moles of water? Ans: ______ e) What mass of carbon dioxide is produced from 7.15 x 1021 molecules of C2H4 Ans: ______ 37 0.280 0.416 1.05 2.81 4.19 6.75 8.78 11.6 15.3 52 72.2 2.2 x 1024 units: moles moles moles moles moles g g g g g cules 3. N2 + 3 F2 2 NF3 a) When 62.0 g of fluorine are reacted, how many moles of NF3 will be formed? Ans: ______ b) How many molecules of N2 are needed to produce 2.85 g of nitrogen trifluoride? Ans: ______ c) 7.61 x 1022 molecules of nitrogen will react with how many grams of fluorine? Ans: ______ d) What mass of nitrogen would react with 15.4 moles of fluorine? Ans: ______ e) How many moles of F2 are needed to produce 1.00 x 1025 molecules of NF3? Ans: ______ 4. 4 NH3 + 7 O2 4 NO2 + 6 H2O a) What mass of NO2 can be produced from 3.56 x 1022 molecules of oxygen? Ans: ______ b) 13.8 g of NH3 would be able to produce how many moles of H2O? Ans: ______ c) What mass of NH3 would react with 10.0 g of oxygen? Ans: ______ d) How many moles of water can be produced from 9.62 x 1026 molecules of O2? Ans: ______ e) What mass of oxygen would react with 16.3 moles of NH3? 38 Ans: (IRO+1) 1.09 1.22 1.55 3.04 8.11 14.4 24.9 73.5 144 913 1370 1.21x1022 units: moles moles moles moles moles g g g g g cules Baking Soda Stoichiometry Lab Name: ________________ Partner: ________________ In this lab you will combine your powers of observation, reasoning, equation balancing and knowledge of stoichiometric calculations to earn a perfect 10/10 (hopefully). Procedure: Data Table 1. Obtain a large Pyrex test tube and weigh it on one of the scales in the front of the room. Record this mass in the table at right. Pyrex is a kind of glass that can be subjected to very high (and low) temperatures without shattering! 2. Go back to your lab station and place one large scoop of baking soda (NaHCO3) into the test tube, then using the same scale as before, weigh the test tube with the baking soda. Record this mass in the table at right. (You should be able to figure out what mass of baking soda is in the test tube.) 3. Holding the test tube nearly horizontal, shake the baking soda gently so that it spreads out a bit as shown at right: 4. Then tighten the test tube clamp securely around the test tube, just below the lip so that it is positioned nearly horizontally about 20 cm above the lab desk as shown: 5. Light a burner and adjust it to a large cool flame hitting the bottom half of the test tube as shown: Record the time you started heating: __________ This heat will initiate a chemical change (a sort of decomposition reaction) that breaks the NaHCO3 down, not into its elements but into three separate compounds. 6. Put one drop of the green universal indicator solution on the small “paint brush”. Then carefully insert this end into the mouth of the test tube as shown at right: See if you can observe a distinct color change. If a metal oxide like K2O, Na2O or MgO is being produced it will create a basic solution and turn the drop blue. If a nonmetal oxide like NO2, SO3 or CO2 is being produced, it will create an acidic solution and turn the drop yellow or even orange. What color does the drop turn? ___________ So, is the reaction producing a metal or a nonmetal oxide? ___________ Look at the chemical formula of the substance you are heating: NaHCO3… what common oxide is being produced in the test tube? 7. What do you observe happening in the upper half of the test tube? ___________________________________ What common substance appears to be a second product of this reaction? 8. Move the burner occasionally to a different spot to ensure a thorough heating of the entire bottom half of the test tube. Consider the substance that is left in the test tube: it may look just like the baking soda you put in the test tube, but it has actually been converted into something else: sodium carbonate. This is the third product. What is the correct formula for sodium carbonate: Now go down and answer questions 1-4 below, but keep an eye on the time: After you have heated the test tube for 8-10 minutes, turn off the burner and let the test tube cool for 5-6 minutes. Questions: 1. You should have figured out from #6, 7 and 8 above what the three products are. Write the chemical equation (unbalanced) for the reaction that just took place: Check it with the teacher to make sure you have it right, then go back and balance it. (Hint: it is a very easy one to balance: the coefficients are all pretty small.) 2. Look back at your data table above. What mass of the NaHCO3 did you start within the test tube: ___________ 3. Starting with that mass of NaHCO3, use stoichiometry (and your balanced equation above) to figure out what mass of sodium carbonate you should have ended up with in the test tube: Show your work in factor label form: 39 This is your official prediction. Make sure it is correct. Your grade depends on it! 4. So... assuming all the baking soda you started with got converted into sodium carbonate, what should the test tube (and contents) weigh right now? If your test tube has been cooling for 5-6 minutes, it should be ready for the official weigh-in! Bring the test tube, along with this sheet containing your prediction above, up to the instructor at the front table. They will weigh it – on the same scale you used – but not show you the weight. They will tell you your grade based how close your prediction was to the actual weight (see the table at right). If you are satisfied with your grade, congratulations! You are done. If you are not satisfied, you can go back, correct your mistake and change your prediction for a second attempt. This second attempt will cost you one point, and you may end up with a lower score, so only try the second attempt if you are fairly sure you can correct whatever mistake you may have made the first time. If your then your prediction grade will is within… be…. 0.03 g 10/10 0.10 g 9/10 0.20 g 8/10 0.30 g 7/10 0.50 g 6/10 1.00 g 5/10 5.00 g 4/10 10.00 g 3/10 20.00 g 2/10 50.00 g 1/10 Otherwise you get a 0/10 5. Observe the substance that is left in the test tube; compare it to the fresh test tube of NaHCO3 at your lab station. Do you notice any slight difference between the two? _______________________________________________. After you have finished all of the above, rinse out the test tube into the sink, then place it in the used bin at the front of the room. Take a fresh (dry) test tube and place it in the clamp for the next group. Follow-up Questions: 6. If you hadn’t heated the test tube long enough, would that make your prediction come out too high or too low? _____ Explain: 7. CO2 is more dense than air. But the CO2 produced from the reaction rose upward out of the mouth of the test tube? Why? _______________________________________________________________________________ 8. Why did the water only condense on the upper half of the test tube? ________________________________ 9. Using your original mass of baking soda (NaHCO3), determine what mass of H2O was produced: ___________ show work: 10. Using your original mass of baking soda (NaHCO3), determine what mass of CO2 was produced: ___________ show work: 11. Add the two masses from #9 and #10 above along with the calculated mass of Na2CO3 produced (#3 above). What total mass of products does this give: _____________ 12. How does this mass compare with the initial mass of NaHCO3 you put in the test tube? ___________________ Explain why this makes sense:_______________________________________________________________ 13. If a person accidentally leaves a pan with cooking oil heating on the stove, it might get so hot that it catches on fire. This is known as a grease fire, Pouring water on a grease fire is a bad idea. Why? 14. Pouring baking soda on a grease fire is a much better idea. Why? 15. Chemical reactions can be categorized as either exothermic (heat is given off by the reaction) or endothermic (heat is taken in by the reaction). What type of reaction is the decomposition of NaHCO3? ____________ How do you know this? 40 Hydrate Lab Name: ________________ Partner: ________________ Based on the introductory discussion, what exactly is a “hydrate?” Write the chemical formulas for these hydrates: Write the names for these hydrates: nickel(II) nitrate hexahydrate _____________ BaBr2·2H2O ____________________________________ sodium acetate trihydrate _____________ FeSO4·4H2O ____________________________________ Billy took a sample of magnesium sulfate ?-hydrate with an unknown coefficient of hydration (MgSO4·XH2O) and put it in a test tube – which he had already found to weigh 18.45 g. The magnesium sulfate ?-hydrate and test tube together weighed 24.66 g. He then heated it gently over a cool flame for several minutes, then let it cool down and weighed it again. It weighed 21.57 g. Just to make sure, he heated again for few more minutes, let it cool and it was then down to 21.48 g, so he heated and cooled it again, and again got 21.48 g. So he figured he had driven off all the water molecules. He quickly calculated that pure (anhydrous*) magnesium sulfate left in the test tube weighed 3.03 g (you should be able to see how he figured that out). Converted to moles, that would be 0.02517 moles MgSO4. He also determined that 3.18 g of water were present in the original hydrate (again, you should be able to see how he figured that out). Converted to moles, that would be 0.1765 moles H2O. From this information, he was able to calculate that the coefficient of hydration (X) must be 7, and thus that the compound was MgSO4·7H2O or magnesium sulfate heptahydrate. Show your work below to demonstrate that you do indeed understand each of these calculations. Objective: Obtain a sample of copper(II) sulfate ?-hydrate (CuSO4·XH2O), and determine X. Procedure: Based on the paragraph above, develop a step by step procedure to accomplish this objective. Write neatly, have the steps numbered, in command form, and include a labeled sketch at right of your set-up. Do a rough draft on scrap paper first. 41 Conduct the experiment, and construct a data table at right to record your raw data. Calculations – done NEATLY in a logical format: Data Table: Results: Based on your results… what was the coefficient of hydration: CuSO4·___H2O (and name: ____________________________________________) Extension: When your test tube has completely cooled, use a squirt bottle to fill it about half-way with distilled water. Feel the bottom of the test tube. What do you feel? _______________________________________________________________________________ Now stopper and shake the test tube for 20-30 seconds. What do you observe? _____________________________________________________________________________ Now find an empty petri dish lid, wipe it dry and write your name on the edge of it. Then weigh it: _________ Pour the copper sulfate solution through a filter into the petri dish. If it didn’t all dissolve, add a Bit more water, stopper and shake and then add this to the filter. When it is done filtering, carefully carry the petri dish to a designated place on the window sill and leave it there. The next day, observe the contents of the petri dish. Draw a sketch at right: After several days, reweigh the dish: _____________. What percent of the original hydrate did you recover? __________ Show work: Follow-up questions: 1. Why did you heat the hydrate in the first place? ________________________________________________________ 2. Why did you reheat the hydrate over and over? ________________________________________________________ 3. How did you know when you were done heating it? _____________________________________________________ 4. Where the reaction was occurring, you could not see the water that was produced. Why not? ___________________ 5. Near the mouth of the test tube, you did see the water. Why? _____________________________________________ 6. Write a complete balanced equation for the reaction that took place in the test tube. Include all descriptors (s, l, g…) 7. A test tube weighs 31.47 g. A sample of sodium acetate trihydrate is placed in the test tube and the combined weight is 38.05 g. The sample is then heated to a constant mass. How much will the test tube and contents weigh now? Show your work: Ans: _________ 42 8. A beaker weighs 245.67 g. Some aluminum sulfate ?-hydrate is added and the combined mass is 398.31 g. It is placed in an oven for 20 minutes, then removed and cooled. It now weighs 353.07 g. What is a) the formula and b) the name of the hydrate? a) _________________ b)_____________________________________ 9. What is the percent water in a) calcium nitrate tetrahydrate? _____ b) sodium sulfide nonahydrate? ______ Ans 7-12 IRO: 1 6 8 30.5 35.44 67.5 L L L L L H H H H H O O units: g % % 10. A cobalt(II) nitrate hydrate is found to be 37.14% water by weight. What is a) its formula and b) its name? a) _________________ b)_____________________________________ 11. A calcium chloride hydrate is found to be 13.97% water by weight. What is a) its formula and b) its name? a) _________________ b)_____________________________________ 12. Decide whether the following error sources would cause your value for X to come out too high (H) too low (L) or not have any effect (0). __ Some of the hydrate compound spilled out of the test tube before the initial weighing. __ Some of the anhydrous compound spilled out before the final weighing. __ The test tube wasn’t heated over a hot enough flame. __ The test tube was heated over a flame that was so hot, it started driving the O’s off the CuSO4 (changing it to CuS). __ The test tube had a slight chip on its rim from the very beginning. __ The test tube had a few drops of water in it from the very beginning. __ The test tube had a few drops of water in it at the end of the experiment. __ The final weighing was done too soon and the test tube was so hot, it had convection currents lifting upward on it. __ The final weighing was done too late and the anhydrous compound had started to absorb moisture from the air. __ The original hydrate was somewhat wet. __ The original hydrate was contaminated with some copper – that oxidized in the heat to form CuO. __ The original hydrate was contaminated with some completely dry dirt – that didn’t do anything in the heat. 43 Unit 4 Review WS (20 pts) Name: ________________________ 1. Name the following acids: Write formulas for the following acids: HF (aq) ____________________________________ hydrobromic acid_________ H2CO3(aq) ____________________________________ oxallic acid _________ HClO4(aq) ____________________________________ boric acid _________ Name the following compounds: Write formulas for the following compounds: K2S2O3 ____________________________________ iron(II) hypoiodite _________ Sn3(PO4)4 ____________________________________ zinc acetate _________ S2Br4 ____________________________________ nitrogen trichloride _________ 2. Draw Lewis structures for each of the following. a) KrF2 b) SiH2F2 c) CO32- d) BF3 For the next three questions, consider the following six Lewis structures: a b c d e f 3. Draw the correct 3-D shapes NEATLY, and name the shapes in the spaces below: a b c d e f Matching from above (a-f) hybridization: sp __ bond angles: 87*__ 90*__ 2 sp __ 106*, 109* & 111*__ 3 sp __ __ 117*__ 3 sp d __ 105*__ 3 2 sp d ___ 180*__ Which of the above (a-f) are polar? __________ nonpolar? ___________ neither? __________ 44 F 4.0 O 3.5 N 3.0 Cl 3.0 Br 2.8 C 2.5 S 2.5 Se 2.4 H 2.1 P 2.1 4. Balance the following equations: __ C5H8 + __ O2 __ CO2 + __ H2O __ K3N + __ Br2 __KBr + __ N2 __(NH4)2C2O4 + __ Fe(C2H3O2)3 __ NH4C2H3O2 + __ Fe2(C2O4)3 __ AlBr3 + __ H2O __ HBr + __ Al(OH)3 __ Ba + __ Al(NO3)3 __BaO + __ Al + __NO2 __N2O5 + __ PCl5 __ NCl3 + __ P4 + __ O2 (this one uses some big #s) 5. Label the following equation: 2 HCl (aq) + Na2CO3(s) 2 NaCl(aq) + CO2(g) + H2O(l) Write a balanced equation like the one above (complete with descriptors like s, l, g, etc) for the following: Crystals of iron(III) iodide are mixed with liquid chlorine to produce iodine vapors and a solution of iron(III) chloride. 6. Complete the following equations. (Assume they all react. No need to balance the overall equations.) a) Al + F2 b) HNO3 + Ca c) O2 + Mg3P2 d) C3H7OH + O2 e) NH4C2H3O2 + HClO4 f) H2 + O2 7. Predict whether the following reactions will occur or not. If they do occur, write the correct product, if they do not, simply write “NR.” a) CuCl2 + H2 d) Zn + MnBr2 b) Br2 + AlF3 c) Ca + Al(NO3)3 e) F2 + MgI2 f) Ca + HC2H3O2 8. Predict whether the following reactions will occur or not. If they do occur, write a balanced net ionic equation beneath the problem. If they do not occur, write “NR.” a) HI + Pb(NO3)2 b) (NH4)3PO4 + KOH c) Cu(C2H3O2)2 + KOH d) HNO3 + Zn(BrO2)2 e) Na3PO4 + (NH4)2SO4 f) Fe(C2H3O2)2 + Na3PO4 45 F Cl Br I Li Rb K Ca Na Mg Al Mn Zn Fe H Ni Sn Pb Cu Ag Pt Au 9. a) How many moles are there in a block of cobalt containing 3.45 x 1025 atoms? __________ b) How many molecules would be present in a 74.5 g sample of silicon tetrafluoride? ___________ 10. a) What mass of lithium phosphate would contain 2.34 kg of lithium? _________ b) A compound is 38.8% calcium, 19.9% phosphorus and the rest oxygen. What is its empirical formula? ________ c) A compound’s empirical formula is C3H5O and 1.00 g of this compound contains 2.64 x 1021 molecules. What is the molecular formula of the compound? Ans: ________ 46 11. For the next few questions, use the equation: 4 NH3(g) + 7 O2(g) 4 NO2(g) + 6 H2O(g) a) What mass of water could be produced from 3.62 x 1025 molecules of NH3? ________ b) If only 1.48 kg of water are produced, what percent yield would this be? ________ When 6.78 moles of ammonia reacts with 9.25 moles of oxygen… c) What mass of water is produced? __________ d) How many moles of ammonia react? _______ e) How many molecules of ammonia are left over? ________ f) How many moles of oxygen react? _______ g) How many molecules of oxygen are left over? ________ 12. An empty test tube weighs 34.56 g. Some Ni(NO3)3 * X H2O is added and it weighs 38.98 g. After heating once it weighed 37.89 g, after a 2nd heating it weighed 37.63 g and after a third heating it weighed 37.63 g. Determine a) the formula and b) the name of the hydrate. a) ______________ b) ________________________________ 47
