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Chemistry 20 Unit 0: Chemistry Review 1 Chemistry 20 Unit 0: Chemistry Review WHMIS workplace hazardous materials information system all chemicals are treated with respect WHMIS has been developed to provide guidelines for handling, storage and disposal of reactive materials MSDS is a material safety data sheet 2 Chemistry 20 Unit 0: Chemistry Review The Atom proton p+ positive charge neutron n0 zero charge electron e- negative charge electrons are found in a cloud region around the nucleus nucleus contains the proton and neutrons which make up most of mass of the atom mass number = protons + neutrons neutrons = mass number - protons isotope = atoms with the same of protons but a different of neutrons (different mass numbers) Examples: carbon-12 carbon-14 3 Chemistry 20 Unit 0: Chemistry Review Periodic Table arranged in groups (columns) and periods (rows) group number = number of outer level (valence) e- period number = number of energy levels occupied by eExample: Na Group Number = 1 Period Number = 3 This information is helpful for drawing Energy Level Diagrams Energy Level Diagrams atoms are electrically neutral which means that the of protons = of electrons maximum number of e-: 3rd level = 8 e2nd level = 8 e1st level = 2 e- 4 Chemistry 20 Unit 0: Chemistry Review Examples: - Lesson #R1: Review of Atomic Structure Draw the energy level diagrams for each of the following elements. 1. hydrogen atom 6. aluminum atom 2. carbon atom 7. phosphorus atom 3. helium atom 8. chlorine atom 4. oxygen atom 9. argon atom 5. sodium atom 10. calcium atom 5 Chemistry 20 Unit 0: Chemistry Review 11. What is the relationship between group number and number of valence (outermost) electrons? 12. What is the relationship between period number and the number of energy levels occupied by electrons? Ions ions are particles or groups of particles that have a net charge (either positive of negative) neutral atoms are unstable if their valence level is not full atoms will strive to satisfy the octet rule in order to become stable…in other words, they strive to have a full valence level and do so by giving away or taking e- Metals - metals give away e- and become positive ions Na+, Ca2+, Fe3+ cations: positively charged ions have fewer electrons than protons naming: metal name gets ion added to it sodium turn to sodium ion Example: 8 e- Chemistry 20 Unit 0: Chemistry Review Na a Na+ + + electron e- Chemistry 20 Unit 0: Chemistry Review Non-Metal - non-metals take e- and become negative ions Cl-, P3-, O2- anions: negatively charged ions have more electrons than protons naming: non-metal name has –ine dropped and –ide added chlorine turn to chlroide Example: A chloride ion, Cl- Chemistry 20 Unit 0: Chemistry Review Cl + electron + e- Cl- Lesson #R2: Review of Ionic Structure Draw the energy level diagrams for each of the following ion. 1. lithium ion 6. nitride ion 2. fluoride ion 7. sodium ion 3. aluminum ion 8. sulphide ion 4. chloride ion 9. calcium ion 5. magnesium ion 10. oxide ion Chemistry 20 Unit 0: Chemistry Review 11. What is the relationship between the electron configuration of an ion of one of the representative elements and the electron configuration of the nearest noble gas? 12. What problem arises when trying to predict the charge on an ion in Group 14? Lesson #R3: Review of Atoms vs. Ions Complete the following chart: Name Example: calcium ion Symbol # of Protons # of Electrons Net Charge Ca2+ 20 18 2+ 1. oxygen atom 2. fluoride ion 3. C 4. Cl 5. 12 2+ 6. 16 2 7. 18 1+ 8. 10 0 9. Ba2+ 10. helium atom 11. H+ 12. 7 13. Fe3+ 14. Sn4+ 10 Chemistry 20 Unit 0: Chemistry Review 15. sodium ion Al3+ 16. 17. 29 18. 2+ 54 1 54 1+ 19. gold atom 20. Elements metals exist as single atoms Example: Li(s), Cu(s), Hg(l) nonmetals and hydrogen do not exist as single atoms – 7 starting at 7 that make a 7 or flag pole! H2 N2 O2 F2 P4 S8 Cl2 Br2 I2 Practice: 1. Cu(s) = 4. fluorine gas = 2. O2(g) = 5. barium = 3. Al(s) = 6. nitrogen gas = Ionic Compounds Cl Cl Chemistry 20 Unit 0: Chemistry Review metals + non-metals or polyatomic ions monovalent (K+, Be2+) or polyatomic metals (Fe3+, Fe2+) charges on the ions are the result of taking or giving e- to go from formula to name: name of first ion, then brackets for charge if multivalent in roman numerals, then name for the second ion - first element ( ) second element-ide Examples: AlCl3 = Fe2O3 = Practice: 1. Zn3P2 = 2. NaNO3 = 3. NiF3 = 4. MnO2 = 5. Cr2(SO4)3 = to go from name to formula: write the symbol for each ion, then add subscripts to balance charges remember to reduce to lowest whole numbers Examples: calcium sulphide = iron (II) hydroxide = Chemistry 20 Unit 0: Chemistry Review Practice: 1. lithium bromide = 4. ammonium sulphate = 2. sodium phosphate = 5. calcium phosphate = 3. magnesium nitride = Hydrated Compounds ionic compounds containing water in their structure water is represented by “ x H2O” in the formula where x is the number of water molecules prefixes: 1 = mono 6 = hexa 2 = di 7 = hepta 3 = tri 8 = octa 4 = tetra 9 = nona 5 = penta 10 = deca Chemistry 20 Unit 0: Chemistry Review to go from formula to name: give the ionic name for the first part of the compound, then name the “ x H2O” part as prefix + “hydrate” Examples: NaF3H2O = CuSO45H2O = to go from name to formula: first part is the same as before…look up the symbol for each ion then balance the charges using subscripts, then for the hydrate part…add “ x H2O” where x is the number given in the prefix Examples: iron (III) nitrate nonahydrate = sodium chlorate tetrahydrate = nickel (II) sulphite heptahydrate = Lesson #R4: Review of Elements and Ionic Nomenclature Complete the following chart: Formula 1. IUPAC Name CdO(s) 2. sodium fluoride 3. chlorine gas 4. AlP 5 H2O(s) copper (II) oxide 5. 6. Mg(OH)2(s) 7. Na2CO3 3 H2O(s) magnesium sulphate nonahydrate 8. 9. N2(g) 10. lithium chloride 11. sodium chlorate 12. 13. K3PO4(s) calcium metal Chemistry 20 Unit 0: Chemistry Review nickel (III) bromide 14. 15. MnO2(s) 16. ammonium sulphite 17. zinc sulphide 18. NaHSO3(s) ammonium sulphate 19. 20. Au(s) copper (II) chloride 21. 22. SnF2(s) 23. phosphorus 24. sodium hypochlorite 25. KMnO4(s) 26. SrF2(s) 27. RbCl(s) 28. Li2O(s) 29. iron (III) sulphide 30. zinc chloride 31. aluminum sulphide 32. CoCl2(s) 33. Au(NO3)3(s) 34. Cu2O(s) 35. lead (IV) acetate trihydrate 36. chromium (II) oxide 37. magnesium iodide 38. KC6H5COO(s) 39. Na2S2O3(s) 40. NH4HCO3(s) 41. ammonium sulphide 42. barium sulphite 43. magnesium hydroxide Chemistry 20 Unit 0: Chemistry Review 44. FeSO4 4 H2O(s) 45. LiCl 2 H2O(s) sodium phosphate decahydrate 46. 47. TiO2(s) 48. bismuth (V) sulphate 49. tin (IV) sulphide 50. NaOH(s) Molecular Compounds non-metals only e- are shared therefore no ions are formed no charges involved use prefixes in naming to go from formula to name: name of first element (including prefix if necessary), then name for second element with “-ide” ending (including prefix) - ___ first element ___ second element-ide Example: N2O = CO2 = P4O10 = to go from name to formula: write the symbol for each elelemt, then use the prefixes to determine th subscripts Example: carbon monoxide = carbon tetrachloride = remember the memorizers?????? NH3 = ammonia H2O = water Chemistry 20 Unit 0: Chemistry Review H2S = hydrogen sulphide HF, HCl, HBr, HI = mo prefixes CH4 = methane CH3OH = methanol C2H6 = ethane C2H5OH = ethanol C6H12O6 = glucose C12H22O11 = sucrose O3 = ozone H2O2 = hydrogen perioxide Acids always have aqueous (aq) as the state and always have hydrogen Rules hydrogen _________ide becomes hydro_______ic acid hydrogen _________ate becomes __________ic acid hydrogen _________ite becomes __________ous acid Examples: Change each of the following to the appropriate acid name and give the formula: hydrogen iodide = hydrogen phosphate = hydrogen nitrite = hydrogen sulphite = Chemistry 20 Unit 0: Chemistry Review Lesson #R5: Review of Molecular Nomenclature and Acids Complete the following chart: Formula 1. NO3(g) ammonia 2. 3. IUPAC Name H2S(g) 4. oxygen difluoride 5. methane 6. CH3OH(l) 7. HBr(aq) sulphurous acid 8. 9. CS2(l) hydrosulphuric acid 10. 11. SO2(g) dinitrogen tetraoxide 12. 13. HNO2(aq) carbon monoxide 14. 15. C12H22O11(s) 16. hypochlorous acid 17. diarsenic trioxide 18. ethanol 19. H2CO3(aq) perchloric acid 20. 21. P4O10(s) sulphur trioxide 22. 23. CF4(l) silicon dioxide 24. 25. CH3COOH(aq) Chemistry 20 Unit 0: Chemistry Review Lesson #R6: Review of Nomenclature Complete the following chart: Class 1. Formula IUPAC Name H3PO4(aq) 2. chlorous acid 3. magnesium 4. Al2(SO4)3(s) magnesium chloride 5. 6. NH4NO2(s) phosphorus trihydride 7. 8. KNO3(s) sodium nitrate hexahydrate 9. 10. HNO2(aq) 11. Al(OH)3(s) sodium sulphate octahydrate 12. 13. (NH4)2SO4(s) 14. PbF4(s) hydrogen peroxide 15. 16. PbO(s) hydrofluoric acid 17. 18. KClO(s) bromine 19. 20. N2O3(g) 21. K2CO3 2 H2O(s) nitric acid 22. 23. HF(g) sodium hydroxide 24. 25. NaHSO3(s) Chemistry 20 Unit 0: Chemistry Review Class Formula magnesium sulphate octahydrate 26. 27. Ca(OH)2(s) gold (I) chloride 28. 29. IUPAC Name CaO(s) 30. copper (II) sulphate pentahydrate 31. sulphur 32. Ca(HCO3)2(s) 33. KBr(s) titanium (IV) oxide 34. 35. PCl5(g) sodium chlorate 36. 37. N2H4(l) 38. hydrogen chloride 39. chloric acid 40. lithium thiosulphate 41. B2H6(g) 42. nitrogen trichloride 43. sodium hydrogen sulphite 44. Al(s) 45. HBr(aq) 46. silicon 47. ammonium phosphate 48. xenon 49. SF2(s) 50. Na2SiO3(s) Chemistry 20 Unit 0: Chemistry Review States acids – always (aq) elements – can be (s), (l) or (g) … see periodic table molecular compounds – can be (s), (l) or (g) ionic compounds – if not in a solution always (s) or if in a solution either (s) or (aq)… look up on solubility chart High solubility (aq) all NO3– ClO3– ClO4– all Low solubility (s) none none NH4+ Ions CH3COO– Ag+ Hg+ most Cl– Br– I– most SO42– S2– OH– most group 1 group 2 NH4+ Ag+ Pb2+ Cu+ Hg+ Tl+ Ag+ Pb2+ Ca2+ Ba2+ Sr2+ Ra2+ most group 1 NH4+ Sr2+ Ba2+ Tl+ most Examples: 1. NaCH3COO( 2. BaSO4( 3. KOH( 6. CaCO3( ) 7. FeSO4( ) ) 4. Pb(NO3)4( ) 5. Hg(CH3COO)2( ) endothermic vs. exothermic reaction types: ) 9. Pb(SO4)2( ) 10. Ca3(PO4)2( 1. hydrocarbon combustion C?H? + O2(g) CO2(g) + H2O(g) Example: CH4(g) + 2 O2(g) ) 8. (NH4)2S( Chemical Reactions ) CO2(g) + 2 H2O(g) ) CO32– PO43– SO32– group 1 NH4+ most Chemistry 20 Unit 0: Chemistry Review 2. formation (simple composition) element + element compound Example: 2 Mg(s) + O2(g) 2 MgO(s) 3. simple decomposition compound element + element Example: 2 H2O(l) 2 H2(g) + O2(g) 4. single replacement element + compound element + compound Example: Cu(s) + 2 AgNO3(aq) 2 Ag(s) + Cu(NO3)2(aq) 5. double replacement compound + compound compound + compound Example: Pb(NO3)2(aq) + 2 KI(aq) Balancing Reactions law of conservation of matter says that matter cannot be created or destroyed, it can only change forms we must balance chemical equations to conserve matter 2 KNO3(aq) + PbI2(aq) Chemistry 20 Unit 0: Chemistry Review Examples: ___ CH4(g) + ___ O2(g) ___ CO2(g) + ___ H2O(g) ___ C2H4(g) + ___ O2(g) ___ CO2(g) + ___ H2O(g) Lesson #R7: Review of Chemical Reactions A. For each of the following reactions, identify the reaction type and balance the reaction. ________________1. _____Al(s) + _____O2(g) _____ Al2O3(s) ________________2. _____HCl(aq) + _____Ca(OH)2(s) ________________3. _____CH4(g) + _____ O2(g) _____CaCl2(aq) + _____HOH(l) _____CO2(g) + _____H2O(g) ________________4. _____Zn(s) + _____Pb(CH3COO)2(aq) _____Pb(s) + _____Zn(CH3COO)2(aq) ________________5. _____SO3(g) + _____H2O(g) ________________6. _____HgO(l) _____Hg(l) ________________7. _____CaCO3(s) _____H2SO4(aq) + _____O2(g) _____CaO(s) + _____CO2(g) ________________8. _____NaI(aq) + _____Pb(NO3)2(aq) ________________9. _____Cl2(g) + _____NaI(aq) _____PbI2(s) + _____NaNO3(aq) _____I2(s) ________________10. _____Al2(SO4)3(aq) + _____Ca(OH)2(aq) + _____NaCl(aq) _____Al(OH)3(s) + _____CaSO4(s) ________________11. ___Al2(SO4)3(aq) + ___Ca(HCO3)2(aq) ___Al(OH)3(s) + ___CaSO4(s) + ___CO2(g) ________________12. _____C8H18(l) + _____ O2(g) _____CO2(g) + _____H2O(g) Chemistry 20 Unit 0: Chemistry Review ________________13. _____H2O(l) ________________14. _____Ba(s) _____H2(g) + _____O2(g) + _____HOH(l) _____H2(g) ________________15. _____H2SO4(aq) + _____Ca3(PO4)2(s) + _____Ba(OH)2(aq) _____H3PO4(aq) + _____CaSO4(s) B. For each of the following word equations, write out the balanced chemical reaction including all states and identify the reaction type. ________________1. water hydrogen + oxygen ________________2. nitrogen + hydrogen ammonia gas ________________3. sulphuric acid + sodium hydroxide water + sodium sulphate ________________4. aluminum + copper (II) nitrate copper + aluminum nitrate ________________5. chlorine + potassium bromide bromine + potassium chloride ________________6. sodium hydroxide + aluminum sulphate aluminum hydroxide + sodium sulphate ________________7. phosphorus + oxygen solid tetraphosphorus decaoxide ________________8. lead (II) nitrate + sodium iodide lead (II) iodide + sodium nitrate ________________9. methanol + oxygen carbon dioxide + water vapour Chemistry 20 Unit 0: Chemistry Review ________________10. nitrogen dioxide gas + water nitric acid + nitrogen monoxide gas Predicting Reactions SR and DR reactions always happen in solutions so for ionic compounds check solubility table composition and decomposition do NOT happen in solutions so ionic compounds are (s) Example: 1. potassium iodide solution is added to lead(II) nitrate solution 2. copper metal is added to a solution of silver nitrate 3. chlorine gas is bubbled through a solution of phosphide Lesson #R8: Review of Predicting Chemical Reactions For each of the following reactions: 1. Write the correct equation including states for each element and compound. 2. Balance the equation. 3. State the reaction type. 1. nitrogen triiodide decomposes explosively into its elements. 2. gallium metal reacts with hydrochloric acid. 3. In a charcoal barbeque, some of the carbon undergoes incomplete combustion to produce deadly carbon monoxide gas. Chemistry 20 Unit 0: Chemistry Review 4. Solutions of calcium nitrate and potassium phosphate are mixed. 5. chlorine gas is bubbled through an aluminum iodide solution. 6. iron reacts with silver nitrate. The iron (III) compound is formed. 7. acetylene (C2H2(g)) burns in a welding torch. 8. copper ore (copper (II) oxide) is decomposed to produce copper metal. 9. titanium (IV) chloride solution reacts with a sodium phosphate solution. 10. sulphuric acid is neutralized by sodium hydroxide. Significant Digits any digit from 1-9 is significant trailing zeros are significant Examples: 6.3800 12 000 “sandwich” zeros are significant Examples: 2.04 1005.002 Chemistry 20 Unit 0: Chemistry Review leading zeros are not significant Example: 0.0065 counted objects and constants are not included in sig digs / : multiply or divide then round answer to the lowest number of significant digits +/ : add or subtract then round answer to the lowest number of decimal places Lesson #R9:Review of Significant Digits, Scientific Notation and SI Units State the number of significant digits in each of the following measured values: 1. 18.56 g _________________ 4. 1.00 W 2. 1500C _________________ 5. 0.05730 mol _________________ 3. 0.0062 L _________________ 6. 8.0 × 10-2 mL _________________ 7. 14.08 cm _________________ 9. 0.100 km _________________ 10. 62 km/h _________________ 8. 1.58 × 108 m _________________ _________________ Convert the following numbers into scientific notation. The number in brackets indicates the number of significant digits the answer is to be rounded to. 1. 1000 _________________ (1) 4. 0.00001098 _________________ (3) 2. 492.32 _________________ (3) 5. 6 995 000 _________________ (3) 3. 0.0573 _________________ (2) 6. 62.49 _________________ (2) Using the SI Prefixes table on your data sheet, perform the following conversions. Maintain the same number of significant digits in each conversion. 1. 0.520 km = _________________m 6. 200 ML = _________________L 2. 100 mL = _________________ L 7. 45 g = _________________kg 3. 152.5 cm = _________________m 8. 10.8 mol = _________________ mmol 4. 3300 mg = _________________g 9. 0.450 L = _________________ mL 5. 650 kg = _________________g 10. 1500 m =_________________ km Perform the following calculations. Round your answer to the correct number of significant digits, using scientific notation where necessary. Include units. 1. 16.56 mL – 6.3 mL = _________________ 2. 21.4 g ÷ 0.825 mol = _________________ Chemistry 20 Unit 0: Chemistry Review 3. 480 km + 24.07 km = _________________ 4. 0.550 mol × 40.00 g/mol = _________________ 5. 18.4 g/mL × 5.5 mL = _________________ 6. 22.99 g/mol + 35.45 g/mol = _________________ 7. 18.5C 4.5C = _________________ 8. 6.0 g ÷ 24.30 g/mol = _________________ 9. 19.55 mL 17.55 mL = _________________ 10. 15 600 g ÷ 2000 mol = _________________ The Mole it is a number = 6.02x1023 “molecules” 1. Molar Mass sum of the individual atomic masses for each element in a compound Examples: 1. CO2 = 2. Al(OH)3 = 3. Cu(ClO3)2 = Practice: 1. NaOH 2. Ca3(PO4)2 3. ammonium hydroxide Chemistry 20 Unit 0: Chemistry Review 2. Mole/Mass Calculations n = m M m = nM where: n = number of moles in mol m = mass in g M = molar mass in g mol Examples: 1. How many moles are in 8.06 g of MgO? 2. What is the mass of 0.677 mol of potassium sulphide? Chemistry 20 Unit 0: Chemistry Review Lesson #R10: Review of Molar Mass and Mole Calculations Complete the following chart, showing all calculations, formulas, substitutions, units and significant digits. Name and Formula Molar Mass Mass Moles 1. NaCl(s) 0.20 mol 2. sodium hydroxide 5.48 g 3. (NH4)3PO4(s) 0.600 mol 4. sodium carbonate octahydrate 50 g Chemistry 20 Unit 0: Chemistry Review Name and Formula Molar Mass Mass Moles 5. Ca(NO3)2(s) 8.45 g 6. potassium dichromate 5.65 g 7. Na2CO3(s) 0.850 mol 8. sulphur trioxide 1.45 mol Chemistry 20 Unit 0: Chemistry Review Science 10 Review 1.Define the following terms: a) proton b) neutron c) electron d) atom e) ion f) valence electron g) octet rule h) monovalent i) multivalent j) endothermic k) exothermic l) law of conservation of matter m) mole n) molar mass o) mass number p) isotope 2. Draw the energy level diagrams for the following: a) fluorine atom d) nitride ion b) carbon atom e) argon atom c) lithium ion f) magnesium ion 3. Where on the periodic table would you find nonmetals? What kind of charge do all nonmetals have? 4. Where on the periodic table would you find metals? What kind of charge do all metals have? 5. Perform the following unit conversions: a) 500 kg = _________________ g b) 25.5 mL = _________________ L c) 102.6 mmol = __________________mol d) 58.2 MJ = _____________________ J e) f) g) h) 600 mg = ____________________ g 9.85 GL = ____________________ L 6.85 cm = ____________________ m 680 nm = _____________________ m 6. Calculate the number of moles in 6.55 g of NaHCO3(s). 7. Calculate the mass of 8.98 mol of AgNO3(s). 8. What is the mass of 0.155 mol of potassium phosphate? Chemistry 20 Unit 0: Chemistry Review 9. How many moles are in 0.558 kg of dinitrogen dioxide? 10. Complete the following chart: Class Formula 1. SrCl2 2. H2S(aq) 3. Na2O 4. H2O 5. CaS2O3 7H2O 6. Fe(IO3)3 7. P2O4 8. S8(s) 9. Ni(OH)2 10. H3PO4(aq) 11. NaCl 12. N2(g) 13. Sb2(SO3)5 14. Ca(s) IUPAC Name 15. sodium chloride 16. copper (II) sulphate pentahydrate 17. ammonium sulphide 18. bismuth (III) sulphate 19. sodium sulphate decahydrate 20. water 21. copper (I) oxide 22. calcium 23. hydroiodic acid 24. radon gas 25. ethanol 26. sucrose Chemistry 20 Unit 0: Chemistry Review 27. nitrogen gas 28. carbonic acid 29. dinitrogen monoxide 30. nitrous acid 11. After the chemical formula for each compound, state the solubility with either (aq) for soluble or (s) for low solubility in water. 1. K2S ( ) 2. NH4CH3COO ( ) 3. Fe(OH)3 ( ) 4. HgBr ( ) 5. BaSO4 ( ) 6. CaCl2 ( ) 7. CuI2 ( ) 8. Ca(CH3COO)2 ( ) 9. FeSO4 ( ) 10. Co(NO3)2 ( ) 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. Zn3(PO4)2 ( ) PbI2 ( ) ZnSO4 ( ) Cu(NO3)2 ( ) AgCl ( ) CdSO4 ( ) NH4Cl ( ) CuS ( ) PbCl2 ( ) Na3PO4 ( ) 12. Balance the following reactions and give the reaction type: a) _____Al2S3(s) _____Al(s) + _____S8(s) b) _____N2(g) + _____O2(g) _____ NO2(g) c) _____Na(s) + _____Pb(CH3COO)2(aq) d) _____Ba(s) e) _____CH4(g) + _____HOH(l) + _____ O2(g) _____H2(g) + _____Ba(OH)2(aq) _____CO2(g) + _____H2O(g) f) _____CaSO4(s) + _____AgNO3(aq) g) _____CH3OH(l) _____Pb(s) + _____NaCH3COO(aq) + _____ O2(g) _____Ag2SO4(s) + _____Ca(NO3)2(aq) _____CO2(g) + _____H2O(g) h) _____Na2SO4(aq) + _____FeCl3(aq) _____NaCl(aq) + _____Fe2(SO4)3(aq) Chemistry 20 Unit 0: Chemistry Review i) _____Cr2O3(s) _____Cr(s) + _____O2(g) j) _____V(s) + _____S8(g) _____ V2S5(s) 13. For each of the following word problems, give the reaction type and write out a balanced chemical reaction including all states of matter. a) nitrogen triiodide decomposes explosively into its elements. b) In a charcoal barbeque, some of the carbon undergoes incomplete combustion to produce deadly carbon monoxide gas. c) Solutions of calcium nitrate and potassium phosphate are mixed. d) The main fuel used to propel rockets into outer space is liquid hydrogen combining with liquid oxygen to produce water vapour. e) chlorine gas is bubbled through an aluminum iodide solution.