Chemistry 12 APRIL Course Code = CH. Student Instructions

Transcription

1 MINISTRY USE ONLY MINISTRY USE ONLY Place Personal Education Number (PEN) here. Place Personal Education Number (PEN) here. MINISTRY USE ONLY Chemistry Ministry of Education APRIL 2001 Course Code = CH Student Instructions 1. Place the stickers with your Personal Education Number (PEN) in the allotted spaces above. Under no circumstance is your name or identification, other than your Personal Education Number, to appear on this booklet. 2. Ensure that in addition to this examination booklet, you have a Data Booklet and an Examination Response Form. Follow the directions on the front of the Response Form. 3. Disqualification from the examination will result if you bring books, paper, notes or unauthorized electronic devices into the examination room. 4. When instructed to open this booklet, check the numbering of the pages to ensure that they are numbered in sequence from page one to the last page, which is identified by ENDÊOFÊEXAMINATION. 5. At the end of the examination, place your Response Form inside the front cover of this booklet and return the booklet and your Response Form to the supervisor.

2 Question 1: 1.. (5) Question 6: 6.. (2) Question 2: 2.. (3) Question 7: 7.. (5) Question 3: 3.. (4) Question 8: 8.. (5) Question 4: 4.. (3) Question 9: 9.. (5) Question 5: 5.. (4) Question 10: 10.. (4)

3 CHEMISTRY 12 APRIL 2001 COURSE CODE = CH

4 GENERAL INSTRUCTIONS 1. Aside from an approved calculator, electronic devices, including dictionaries and pagers, are not permitted in the examination room. 2. All multiple-choice answers must be entered on the Response Form using an HBÊpencil. Multiple-choice answers entered in this examination booklet will not be marked. 3. For each of the written-response questions, write your answer in the space provided in this booklet. 4. Ensure that you use language and content appropriate to the purpose and audience of this examination. Failure to comply may result in your paper being awarded a zero. 5. This examination is designed to be completed in two hours. Students may, however, take up toê30 minutes of additional time to finish.

5 CHEMISTRY 12 PROVINCIAL EXAMINATION 1. This examination consists of two parts: Value Suggested Time PART A: 48 multiple-choice questions PART B: 10 written-response questions Total: 100 marks 120 minutes 2. The following tables can be found in the separate Data Booklet. Periodic Table of the Elements Atomic Masses of the Elements Names, Formulae, and Charges of Some Common Ions Solubility of Common Compounds in Water Solubility Product Constants at 25 C Relative Strengths of Br nsted-lowry Acids and Bases Acid-Base Indicators Standard Reduction Potentials of Half-cells No other reference materials or tables are allowed. 3. A calculator is essential for the Chemistry 12 Provincial Examination. The calculator must be a hand-held device designed primarily for mathematical computations involving logarithmic and trigonometric functions and may also include graphing functions. Computers, calculators with a QWERTY keyboard, and electronic writing pads will not be allowed. Students must not bring any external devices to support calculators such as manuals, printed or electronic cards, printers, memory expansion chips or cards, or external keyboards. Students may have more than one calculator available during the examination. Calculators may not be shared and must not have the ability to either transmit or receive electronic signals. In addition to an approved calculator, students will be allowed to use rulers, compasses, and protractors during the examination.

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7 PART A: MULTIPLE CHOICE Value: 60 marks INSTRUCTIONS: Suggested Time: 70 minutes For each question, select the best answer and record your choice on the Response Form provided. Using an HB pencil, completely fill in the circle that has the letter corresponding to your answer. Selected multiple-choice questions are worth 2 marks. 1. Which of the following reactions occurs most rapidly at standard conditions? (1Êmark) A. 2Fe( s) + O2( g) 2FeO( s) B. CaO( s) + 3C( s) CaC2( s) + CO( g) C. SnO2( s) + 2CO( g) Sn( s) + 2CO2( g) D. 2AgNO + Na CrO Ag CrO + 2NaNO 3( aq) 2 4( aq) 2 4( s) 3( aq) 2. Consider the following reaction: (1Êmark) CaO( s) + 2HCl( aq) CaCl2( aq) + H2O( l) Which of the following could be used to measure the rate of this reaction? A. change in acidity B. change in volume C. change in pressure D. change in total mass 3. In order for a collision between reactant particles to be successful (1Êmark) A. a H must be positive. B. the system must be closed. C. there must be sufficient KE. D. the change in KE must be less than the change in PE OVER

8 4. Consider the following PE diagram: (1Êmark) PE (kj) I II III IV Progress of the reaction The activation energy for the forward reaction is represented by A. I B. II C. III D. IV 5. What is the relationship between the activation energy and the rate of a reaction? (1Êmark) A. When the activation energy is high, the rate of reaction is fast. B. When the activation energy is low, the rate of reaction is slow. C. When the activation energy is high, the rate of reaction is slow. D. There is no relationship between activation energy and rate of reaction. 6. Consider the following reaction mechanism: (1Êmark) Step 1 OCl + H O HOCl + OH Step 2 I + HOCl HOI + Cl Step 3 HOI + OH H O + OI 2 2 Which of the following is correct for the overall reaction? A. HOI is a product. B. HO 2 is a reactant. C. HOCl is a catalyst. D. OH is a reaction intermediate

9 7. Consider the following equilibrium reaction: (2Êmarks) ( ) ( ) + ( ) 2ICl g I2 g Cl2 g Some ICl is added to an empty flask. How do the reaction rates change as the system approaches equilibrium? forward rate reverse rate A. increases increases B. increases decreases C. decreases increases D. decreases decreases 8. In an equilibrium system, continuing microscopic changes indicate that the equilibrium is (1Êmark) A. dynamic. B. complete. C. exothermic. D. spontaneous. 9. Consider the following equilibrium: (1Êmark) 4CuO( s) + energy 2Cu2O( s) + O2( g) The equilibrium will shift to the right as a result of A. adding CuO ( s). B. removing O 2( g). C. adding a catalyst. D. decreasing the temperature OVER

10 10. Consider the following equilibrium: (2Êmarks) N2( g) + 3H 2( g) 2NH3 ( g) The volume of the system is decreased. The equilibrium shifts A. left since the reverse rate is greater than the forward rate. B. left since the forward rate is greater than the reverse rate. C. right since the reverse rate is greater than the forward rate. D. right since the forward rate is greater than the reverse rate. 11. Consider the following equilibrium: (2Êmarks) 2SO3 ( g) 2SO2( g) O2( g) H = kj When the temperature is increased, the equilibrium will shift A. left with K eq becoming larger. B. right with K eq becoming larger. C. left with K eq becoming smaller. D. right with K eq becoming smaller. 12. Starting with equal concentrations of reactants, which of the following will be closest to completion at equilibrium? (1Êmark) A. CO( g) + Cl 2( g) COCl2( g) K eq = 22 B. PCl3 ( g) + Cl2( g) PCl5 ( g) K = eq C. CO( g) + Cl 2( g) COCl2( g) K eq = D. CH3O2( g) + NO2( g) CH3O2NO2( g) K = eq - 4 -

11 13. Consider the following equilibrium: (1Êmark) ( ) ( ) + ( ) 2COF2 g CO2 g CF4 g At equilibrium, a 1. 00L container contains mol COF2, mol CO2, andê mol CF 4. What is the value of K eq? A B C D Which of the following dissolves in water to form a molecular solution? (1Êmark) A. KCl B. Na 2 O C. NH4Br D. CHOH A saturated solution is formed by adding g PbI2( s) to ml of water in a beaker. Describe the situation which exists in the beaker. (1Êmark) 2+ A. [ Pb ]= [ I ] 2+ B. moles PbI2( s) = moles Pb ( aq) C. mass of PbI2( s) = mass of PbI2( aq) D. rate of crystalization = rate of dissociation 16. What is the concentration of barium ions in a 1. 00L solution containing 208. g of BaCl 2? (1Êmark) A M B M C M D M OVER

12 17. Which of the following salts has low solubility? (1Êmark) A. MgS B. ZnCl 2 C. SrSO 4 D. AgNO Consider the following solubility equilibrium: (2Êmarks) AgCl + ( s) Ag ( aq) + Cl ( aq) Some NaCl ( s) is added to the equilibrium. When equilibrium is reestablished, how have the ion concentrations changed from the original equilibrium? [ Ag + ] [ Cl ] A. decreased increased B. decreased decreased C. increased decreased D. increased increased - 6 -

13 19. A precipitate forms when a 020. M solution containing an unknown cation is added to SO 2 4, but not when an equal volume is added to S 2. (2Êmarks) 0.20 M unknown cation 0.20 M unknown cation 0.20 M SO M S 2 The unknown cation is A. Na + B. Ca 2+ C. Pb 2+ D. Zn The K sp expression for a saturated solution of Ni3( PO4) 2 is (1Êmark) sp = [ 3 2+ ] [ 4 ] sp = [ 2 2+ ] [ 4 ] A. K Ni PO B. K Ni PO C. K 3Ni 2PO sp = 2+ 3 [ ][ 4 ] sp = 3 2+ [ ] [ 4 ] D. K 3Ni 2PO Which of the following are general properties of bases in aqueous solution? (2Êmarks) + A. feel slippery and increase [ HO 3 ] B. turn litmus red and accept a proton C. conduct electricity and turn litmus blue D. feel slippery and react with Au to produce H 2( g) OVER

14 22. The conjugate base of HPO 2 4 is (1Êmark) A. PO 4 3 B. HPO 4 C. HPO 4 2 D. HPO The electrical conductivities of 010. M solutions of NaCl, HCN and HNO 2 are measured. The order by conductivity from highest to lowest is (2Êmarks) A. NaCl > HNO2 > HCN B. HCN > HNO2 > NaCl C. NaCl > HCN > HNO2 D. HNO2 HCN NaCl 24. Which of the following acids has the weakest conjugate base? (1Êmark) A. HIO 3 B. HNO 2 C. HPO 3 4 D. CH3COOH 25. When ml of M HCl is added to ml of water, the concentration ofê HO + 3 in the final solution is (1Êmark) A M B M C M D M - 8 -

15 26. Which of the following chemical species are amphiprotic in aqueous solution? (2Êmarks) I. F II. NH 4 + III. HPO 4 2 A. I only. B. II only. C. III only. D. II and III only. 27. A solution is prepared by mixing mol HCl with Calculate the moles of OH present after mixing. mol KOH. (1Êmark) A. 0 mol B C D mol mol mol 28. Calculate the ph in a M solution of Sr( OH) 2. (2Êmarks) A B C D The K b value for HPO 4 2 is (1Êmark) A B C D OVER

16 30. Which of the following 1. 0M salt solutions is acidic? (1Êmark) A. BaS B. NH4Cl C. Ca( NO 3 ) 2 D. NaCH3COO 31. Which of the following represents the hydrolysis reaction that occurs in a solution of K CO 2 2 4? (1Êmark) A. KCO 2K + CO B. K + H O KOH + H O C. C O + H O HC O + OH D. K2C2O4 + H O 2 K2CO3 + CO2 + H2 32. When the indicator thymol blue is added to a 010. M solution of an unknown acid, the solution is red. The acid could be (1Êmark) A. HF B. HS 2 C. HCN D. HNO

17 33. The complete neutralization of ml of KOH requires mol H2SO 4. TheÊ[ KOH]Êwas (1Êmark) A M B M C M D M What is the [ HO 3 ] at the equivalence point for the titration between HBr and KOH? (1Êmark) A B C D. 00. M M M M 35. Which of the following would form a buffer solution when equal moles are mixed together? (1Êmark) A. HCl and NaCl B. HCN and NaCN C. KNO3 and KOH D. Na SO and NaOH Which of the following oxides dissolves to form a solution with a ph greater than 7? (1Êmark) A. SO 2 B. CO 2 C. NO 2 D. KO OVER

18 37. The ph of acid rain could be (1Êmark) A. 50. B. 70. C. 90. D Consider the following reaction: (1Êmark) The species being oxidized is + 2+ s aq aq 2 g Zn( ) + 2H ( ) Zn ( ) + H ( ) A. H 2 B. H + C. Zn D. Zn When SO reacts to form S O, the sulphur atoms (2Êmarks) A. lose electrons and are reduced. B. gain electrons and are reduced. C. lose electrons and are oxidized. D. gain electrons and are oxidized. 40. Which of the following is a list of metals in order from strongest to weakest reducing agents? (1Êmark) A. Au > Ni > Rb B. Ni > Au > Rb C. Ni > Rb > Au D. Rb > Ni > Au

19 41. Consider the following spontaneous reaction: Mg( s) + 2HCl( aq) MgCl2( aq) + H2( g) Which of the following statements is correct? (1Êmark) A. Mg is a weaker reducing agent than H 2 B. Mg is a weaker reducing agent than H + C. Mg is a stronger reducing agent than H 2 D. Mg is a stronger reducing agent than H Which of the following will not react spontaneously with H O 2 at standard conditions? (1Êmark) A. F 2 B. Ca C. Na D. Sn 43. When a piece of Cu is placed in 1. 0M AgNO 3, (1Êmark) A. the [ Ag + ] increases. B. the [ Cu 2+ ] increases. C. the [ NO 3 ] decreases. D. no change occurs OVER

20 Use the following diagram to answer questions 44 and 45. Volts Cu 1.0 M KNO3 Ni 1.0 M Cu( NO 3 ) M Ni( NO 3 ) Which of the following diagrams represents the relationship between [ Ni 2+ ] and the mass of the Cu electrode as the cell above is in operation? (1Êmark) A. B. [ Ni 2+ ] [ Ni 2+ ] Mass of Cu Mass of Cu C. D. [ Ni 2+ ] [ Ni 2+ ] Mass of Cu Mass of Cu 45. The E for the above cell is (1Êmark) A volts B volts C volts D volts

21 46. Which of the following describes an electrochemical cell? (2Êmarks) E cell Type of reaction A. positive spontaneous B. positive non-spontaneous C. negative spontaneous D. negative non-spontaneous 47. Which of the following aqueous solutions should not be used as an electrolyte in an electrolytic cell? (1Êmark) A. 10. M KOH B. 10. M H2SO4 C. 10. M CuSO4 D. 10. M C6H12O6 48. When 1. 0M Na2SO 4 is electrolyzed, the solution near the anode becomes (2Êmarks) A. basic and bubbles form. B. acidic and bubbles form. C. basic and no bubbles form. D. acidic and no bubbles form. This is the end of the multiple-choice section. Answer the remaining questions directly in this examination booklet OVER

22 PART B: WRITTEN RESPONSE Value: 40 marks INSTRUCTIONS: Suggested Time: 50 minutes You will be expected to communicate your knowledge and understanding of chemical principles in a clear and logical manner. Your steps and assumptions leading to a solution must be written in the spaces below the questions. Answers must include units where appropriate and be given to the correct number of significant figures. For questions involving calculation, full marks will NOT be given for providing only an answer. 1. An Alka-Seltzer tablet is added to water to produce carbon dioxide gas. The gas was collected using water displacement. The following data is recorded: Time (s) Volume of CO 2 (ml)

23 a) Calculate the average rate of reaction for the formation of CO 2 gas for the times: i) 010s (1Êmark) ii) s (1Êmark) b) Suggest a reason why the rate of reaction from 0 to 10. 0s is slower than the rate from to s? (1Êmark) c) The rate of reaction is not constant during the entire interval from to s. Describe the change in rate and explain a reason for the change. (2Êmarks) OVER

24 2. A flask is initially filled with some HI. At equilibrium, the [ HI]= 080. mol L. What is the H 2 [ ] at equilibrium? (3Êmarks) 2HI( g) H2( g) + I2( g) Keq =

25 3. Consider the following equilibrium system: 2NOCl ( g) 2NO( g) + Cl2( g) Keq = A 1. 00L flask is filled with 020. mol NOCl, 0. 10mol NO and 0. 10mol Cl 2. State and show by calculation the direction in which the reaction proceeds to reach equilibrium. (4Êmarks) Direction: Calculations: OVER

26 4. In a titration, ml of NaCl aq ( ) reacts completely with ml of M AgNO. What is the Cl [ ] in the original solution? (3Êmarks)

27 5. The following data was obtained when ml of a saturated solution of PbI 2 was evaporated to dryness. Mass of evaporating dish g Mass of evaporating dish and residue g Use this information to determine the Ksp of PbI2. (4Êmarks) OVER

28 6. a) Write the equation for the predominant reaction of HC O 2 4 with HSO 3. (1Êmark) HC O acid + HSO + base b) Identify a Br nsted-lowry conjugate acid base pair from the above reaction. (1Êmark) Acid: Base: 7. a) In the space below, sketch the titration curve for the reaction whenê 010. M HCl is added to 100. ml of 010. M NaOH. (3Êmarks) ph Volume of added HCl (ml) b) Describe two changes in the titration curve that would result from using 010. M CH3COOH in place of the HCl. (2Êmarks) i) ii)

29 [ ] 050 3( ) 8. Calculate the OH in. M NH aq. (5Êmarks) OVER

30 9. Balance the following redox reaction in basic solution. (5Êmarks) MnO C O MnO CO basic + + ( )

31 10. a) Draw and label the parts of an operating electrolytic cell during the electrolysis of molten potassium chloride KCl ( l). (3Êmarks) b) Define the term oxidizing agent. (1Êmark) END OF EXAMINATION

32 Data Booklet CHEMISTRY 12 Work done in this booklet will not be marked. Ministry of Education Revised January 2000

33 CONTENTS Page Table 1 Periodic Table of the Elements 2 Atomic Masses of the Elements 3 Names, Formulae, and Charges of Some Common Ions 4 Solubility of Common Compounds in Water 5 Solubility Product Constants at 25 C 6 Relative Strengths of Brønsted-Lowry Acids and Bases 7 Acid-base Indicators 8 Standard Reduction Potentials of Half-cells REFERENCE D.R. Lide, CRC Handbook of Chemistry and Physics, 80 th edition, CRC Press, Boca Raton, 1999.

34 PERIODIC TABLE OF THE ELEMENTS H Hydrogen Li Lithium Be Beryllium Si Silicon 28.1 Atomic Number Symbol Name Atomic Mass 5 B Boron C Carbon N Nitrogen O Oxygen F Fluorine He Helium Ne Neon Na Sodium Mg Magnesium Al Aluminum Si Silicon P Phosphorus S Sulphur Cl Chlorine Ar Argon K Potassium Ca Calcium Sc Scandium Ti Titanium V Vanadium Cr Chromium Mn Manganese Fe Iron Co Cobalt Ni Nickel Cu Copper Zn Zinc Ga Gallium Ge Germanium As Arsenic Se Selenium Br Bromine Kr Krypton Rb Rubidium Sr Strontium Y Yttrium Zr Zirconium Nb Niobium Mo Molybdenum Tc Technetium (98) 44 Ru Ruthenium Rh Rhodium Pd Palladium Ag Silver Cd Cadmium In Indium Sn Tin Sb Antimony Te Tellurium I Iodine Xe Xenon Cs Cesium Fr Francium (223) 56 Ba Barium Ra Radium (226) 57 La Lanthanum Ac Actinium (227) 72 Hf Hafnium Rf Rutherfordium (261) 73 Ta Tantalum Db Dubnium (262) 74 W Tungsten Sg Seaborgium (263) 75 Re Rhenium Bh Bohrium (262) 76 Os Osmium Hs Hassium (265) 77 Ir Iridium Mt Meitnerium (266) 78 Pt Platinum Au Gold Hg Mercury Tl Thallium Pb Lead Bi Bismuth Po Polonium (209) 85 At Astatine (210) 86 Rn Radon (222) Based on mass of C 12 at Values in parentheses are the masses of the most stable or best known isotopes for elements which do not occur naturally. 58 Ce Cerium Th Thorium Pr Praseodymium Pa Protactinium Nd Neodymium U Uranium Pm Promethium (145) 93 Np Neptunium (237) 62 Sm Samarium Pu Plutonium (244) 63 Eu Europium Am Americium (243) 64 Gd Gadolinium Cm Curium (247) 65 Tb Terbium Bk Berkelium (247) 66 Dy Dysprosium Cf Californium (251) 67 Ho Holmium Es Einsteinium (252) 68 Er Erbium Fm Fermium (257) 69 Tm Thulium Md Mendelevium (258) 70 Yb Ytterbium No Nobelium (259) 71 Lu Lutetium Lr Lawrencium (262) 1

35 ATOMIC MASSES OF THE ELEMENTS Based on mass of C 12 at Values in parentheses are the mass number of the most stable or best known isotopes for elements that do not occur naturally. 2 Actinium Aluminum Americium Antimony Argon Arsenic Astatine Barium Berkelium Beryllium Bismuth Boron Bromine Cadmium Calcium Californium Carbon Cerium Cesium Chlorine Chromium Cobalt Copper Curium Dubnium Dysprosium Einsteinium Erbium Europium Fermium Fluorine Francium Gadolinium Gallium Germanium Gold Hafnium Helium Holmium Hydrogen Indium Iodine Iridium Iron Krypton Lanthanum Lawrencium Lead Lithium Lutetium Magnesium Manganese Mendelevium Ac Al Am Sb Ar As At Ba Bk Be Bi B Br Cd Ca Cf C Ce Cs Cl Cr Co Cu Cm Db Dy Es Er Eu Fm F Fr Gd Ga Ge Au Hf He Ho H In I Ir Fe Kr La Lr Pb Li Lu Mg Mn Md (227) 27.0 (243) (210) (247) (251) (247) (262) (252) (257) 19.0 (223) (262) (258) Mercury Molybdenum Neodymium Neon Neptunium Nickel Niobium Nitrogen Nobelium Osmium Oxygen Palladium Phosphorus Platinum Plutonium Polonium Potassium Praseodymium Promethium Protactinium Radium Radon Rhenium Rhodium Rubidium Ruthenium Rutherfordium Samarium Scandium Selenium Silicon Silver Sodium Strontium Sulphur Tantalum Technetium Tellurium Terbium Thallium Thorium Thulium Tin Titanium Tungsten Uranium Vanadium Xenon Ytterbium Yttrium Zinc Zirconium Hg Mo Nd Ne Np Ni Nb N No Os O Pd P Pt Pu Po K Pr Pm Pa Ra Rn Re Rh Rb Ru Rf Sm Sc Se Si Ag Na Sr S Ta Tc Te Tb Tl Th Tm Sn Ti W U V Xe Yb Y Zn Zr (237) (259) (244) (209) (145) (226) (222) (261) (98) Element Symbol Atomic Number Atomic Mass Element Symbol Atomic Number Atomic Mass

36 NAMES, FORMULAE, AND CHARGES OF SOME COMMON IONS * Aqueous solutions are readily oxidized by air. ** Not stable in aqueous solutions. Positive Ions (Cations) Al 3+ Aluminum Pb 4+ Lead(IV), plumbic NH 4 + Ammonium Li + Lithium Ba 2+ Barium Mg 2+ Magnesium Ca 2+ Calcium Mn 2+ Manganese(II), manganous Cr 2+ Chromium(II), chromous Mn 4+ Manganese(IV) Cr 3+ Chromium(III), chromic Hg 2 2+ Mercury(I)*, mercurous Cu + Copper(I)*, cuprous Hg 2+ Mercury(II), mercuric Cu 2+ Copper(II), cupric K + Potassium H + Hydrogen Ag + Silver H 3 O + Hydronium Na + Sodium Fe 2+ Iron(II)*, ferrous Sn 2+ Tin(II)*, stannous Fe 3+ Iron(III), ferric Sn 4+ Tin(IV), stannic Pb 2+ Lead(II), plumbous Zn 2+ Zinc Negative Ions (Anions) Br Bromide OH Hydroxide CO 3 2 Carbonate ClO Hypochlorite ClO 3 Chlorate I Iodide Cl Chloride HPO 4 2 Monohydrogen phosphate ClO 2 Chlorite NO 3 Nitrate CrO 4 2 Chromate NO 2 Nitrite CN Cyanide C 2 O 4 2 Oxalate Cr 2 O 7 2 Dichromate O 2 Oxide** H 2 PO 4 Dihydrogen phosphate ClO 4 Perchlorate CH 3 COO Ethanoate, acetate MnO 4 Permanganate F Fluoride PO 4 3 Phosphate HCO 3 Hydrogen carbonate, bicarbonate SO 4 2 Sulphate HC 2 O 4 Hydrogen oxalate, binoxalate S 2 Sulphide HSO 4 Hydrogen sulphate, bisulphate SO 3 2 Sulphite HS Hydrogen sulphide, bisulphide SCN Thiocyanate HSO 3 Hydrogen sulphite, bisulphite 3

37 SOLUBILITY OF COMMON COMPOUNDS IN WATER The term soluble here means > 0.1 mol/l at 25 C. Negative Ions (Anions) Positive Ions (Cations) Solubility of Compounds All Alkali ions: Li +, Na +, K +, Rb +, Cs +, Fr + Soluble All Hydrogen ion: H + Soluble All Ammonium ion: NH 4 + Soluble Nitrate, NO 3 All Soluble or or Chloride,Cl Bromide, Br Iodide, I All others Ag +, Pb 2+, Cu + Soluble Low Solubility Sulphate, SO 4 2 All others Ag +, Ca 2+, Sr 2+, Ba 2+, Pb 2+ Soluble Low Solubility Sulphide, S 2 Alkali ions, H +, NH 4 +, Be 2+, Mg 2+, Ca 2+, Sr 2+, Ba 2+ All others Soluble Low Solubility Hydroxide, OH Alkali ions, H +, NH 4 +, Sr 2+ All others Soluble Low Solubility or or Phosphate, PO 4 3 Carbonate, CO 3 2 Sulphite, SO 3 2 Alkali ions, H +, NH 4 + All others Soluble Low Solubility 4

38 SOLUBILITY PRODUCT CONSTANTS AT 25 C Name Barium carbonate Barium chromate Barium sulphate Calcium carbonate Calcium oxalate Calcium sulphate Copper(I) iodide Copper(II) iodate Copper(II) sulphide Iron(II) hydroxide Iron(II) sulphide Iron(III) hydroxide Lead(II) bromide Lead(II) chloride Lead(II) iodate Lead(II) iodide Lead(II) sulphate Magnesium carbonate Magnesium hydroxide Silver bromate Silver bromide Silver carbonate Silver chloride Silver chromate Silver iodate Silver iodide Strontium carbonate Strontium fluoride Strontium sulphate Zinc sulphide Formula BaCO 3 BaCrO 4 BaSO 4 CaCO 3 CaC 2 O 4 CaSO 4 CuI Cu IO 3 CuS FeS ( ) AgBr AgCl AgI ZnS ( ) 2 Fe( OH) 2 Fe OH 3 PbBr 2 PbCl 2 ( ) 2 Pb IO 3 PbI 2 PbSO 4 MgCO 3 ( ) 2 Mg OH AgBrO 3 Ag 2 CO 3 Ag 2 CrO 4 AgIO 3 SrCO 3 SrF 2 SrSO 4 K sp

39 RELATIVE STRENGTHS OF BR NSTED-LOWRY ACIDS AND BASES in aqueous solution at room temperature. Name of Acid Acid Base K a STRONG STRENGTH OF ACID WEAK Perchloric HClO H + ClO Hydriodic HI H + I + Hydrobromic HBr H + Br + Hydrochloric HCl H + Cl Nitric HNO H + NO very large very large very large very large very large Sulphuric H2SO4 H + HSO4 very large + Hydronium Ion H O + 3 H + H2O Iodic HIO + 3 H + IO Oxalic H C O + H + HC O ( 2 ) , ( III) 6 + Fe( H2O) ( OH) ( ) + Sulphurous SO H O H SO H HSO Hydrogen sulphate ion HSO H SO Phosphoric H PO H H PO Hexaaquoiron ion iron ion Fe H O H Citric H C H O H H C H O Nitrous HNO H NO Hydrofluoric HF H + F Methanoic, formic HCOOH H + HCOO Hexaaquochromium ion chromium III ion Cr H , ( ) 2O H Cr H 6 2O OH H2C6H5O + H HC H O Al( H2O) H Al H O OH 6 + ( 2 ) ( ) Benzoic C H COOH H C H COO Hydrogen oxalate ion HC O H C O Ethanoic, acetic CH COOH H CH COO. Dihydrogen citrate ion ( ) + ( ) ( ) Hexaaquoaluminum ion, aluminum ion Carbonic ( CO + HO) HCO 2 H + HCO Monohydrogen citrate ion HC H O H C H O Hydrogen sulphite ion HSO H SO Hydrogen sulphide H S H HS Dihydrogen phosphate ion H PO H HPO Boric H BO H H BO Ammonium ion NH H NH Hydrocyanic HCN H + CN Phenol C 6H5OH H + C6H5O Hydrogen carbonate ion HCO3 H + CO Hydrogen peroxide H2O2 H + HO Monohydrogen phosphate ion 2 + HPO4 H + PO Water H O + H + OH Hydroxide ion OH H + O very small Ammonia NH H + NH very small WEAK STRENGTH OF BASE STRONG Ð 6 Ð

40 ACID-BASE INDICATORS Indicator Methyl violet Thymol blue Orange IV Methyl orange Bromcresol green Methyl red Chlorophenol red Bromthymol blue Phenol red Neutral red Thymol blue Phenolphthalein Thymolphthalein Alizarin yellow Indigo carmine ph Range in Which Colour Change Occurs Colour Change as ph Increases yellow to blue red to yellow red to yellow red to yellow yellow to blue red to yellow yellow to red yellow to blue yellow to red red to amber yellow to blue colourless to pink colourless to blue yellow to red blue to yellow 7

41 STANDARD REDUCTION POTENTIALS OF HALF-CELLS Ionic concentrations are at 1M in water at 25 C. Oxidizing Agents Reducing Agents E ( Volts) STRONG STRENGTH OF OXIDIZING AGENT WEAK Overpotential Effect F2( g) + 2e 2F SO e 2SO HO H + 2e 2HO MnO4 + 8H + 5e Mn + 4H2O Au + 3e Au( s) BrO3 + 6H + 5e 1 Br 2 2( l) + 3H2O ClO4 + 8H + 8e Cl + 4H2O Cl2( g) + 2e 2Cl Cr2O H + 6e 2Cr + 7H2O O 2 2( ) + 2H + 2e g H2O MnO2( s) H 2e Mn + 2H2O IO3 6H 5e I ( ) s H2O Br2( l) + 2e 2Br AuCl + 4 3e Au( s) + 4Cl NO H 3e NO( g) + 2H2O Hg + 2e Hg( l) O ( ) + H ( M)+ e g H2O + + 2NO3 + 4H + 2e N2O4 + 2H2O Ag + e Ag( s) Hg e Hg( l) Fe + e 2+ Fe O ( ) + 2H + 2e H O 2 g MnO H2O 3e MnO2( s) + 4OH I2( s) + 2e 2I Cu + e Cu( s) HSO H 4e S( s) + 3H2O Cu + 2e Cu( s) SO4 + 4H + 2e H2SO3 + H2O Cu + e + Cu Sn + 2e 2+ Sn S( ) + 2H + 2e s H2S( g) H + 2e H2( g) Pb + 2e Pb( s) Sn + 2e Sn( s) Ni + 2e Ni( s) HPO H + 2e HPO HO Co + 2e Co( s) Se( ) + 2H + 2e s H2Se Cr + e 2+ Cr H2O + e 7 2 H2 + 2OH ( 10 M) Fe + 2e Fe( s) Ag S + e 2 2 ( s) 2 2Ag() s + S Cr + 3e Cr( s) Zn + 2e Zn( s) Te + H + e ( s) 2 2 H2 Te 079. HO+ e H2( g) + 2OH Mn + 2e Mn( s) Al + 3e Al( s) Mg + 2e Mg( s) Na + e Na( s) Ca + 2e Ca( s) Sr + 2e Sr( s) Ba + 2e Ba( s) K + e K( s) Rb + e Rb( s) Cs + e Cs( s) Li + e Li( s) 304. Ð 8 Ð Overpotential Effect WEAK STRENGTH OF REDUCING AGENT STRONG