Chemistry 12 JANUARY 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 JANUARY 2003 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 END OF EXAMINATION. 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.. (3) Question 7: 7.. (2) Question 2: 2.. (3) Question 8: 8.. (2) Question 3: 3.. (3) Question 9: 9.. (5) Question 4: 4.. (3) Question 10: 10.. (3) Question 5: 5.. (4) Question 11: 11.. (4) Question 6: 6.. (4) Question 12: 12.. (4)

3 Chemistry 12 JANUARY 2003 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 HB pencil. 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 to 30 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: 12 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 Brønsted-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. The calculator must not be programmable. Computers, calculators with a QWERTY keyboard or symbolic manipulation abilities, and electronic writing pads will not be allowed. Students must not bring any external devices (peripherals) to support calculators such as manuals, printed or electronic cards, printers, memory expansion chips or cards, CD-ROMS, libraries or external keyboards. 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. Note that some multiple-choice questions are worth 2 marks. 1. Given the reaction: CaCO3( s) + 2HCl( aq) Æ CaCl2( aq) + H2O( l) + CO2( g) + heat Which of the following will cause the reaction rate to increase? A. increasing pressure B. decreasing pressure C. increasing temperature D. decreasing temperature 2. Consider the following reaction: Fe O s + 2Al s Æ Al O s + 2Fe l 2 3( ) ( ) 2 3( ) ( ) If mol of Fe is produced in sec, what is the rate of consumption of Fe 2 O 3 in mol s? (2 marks) A mol s B mol s C mol s D. 50. mol s 3. Which of the following would result in a successful collision between reactant particles? A. particles have sufficient KE B. particles convert all their PE into KE C. particles are in an excited state and are catalyzed D. particles have sufficient KE and proper molecular orientation OVER

8 4. Which of the following best describes the E a of a fast reaction and the stability of its activated complex? E a Activated Complex A. small unstable B. small stable C. large unstable D. large stable 5. How does the addition of a catalyst increase the reaction rate of an endothermic reaction? A. It reduces the DH of the reaction. B. It increases the DH of the reaction. C. It reduces the required activation energy. D. It causes the reaction to become exothermic. 6. Consider the following PE diagram: PE (kj) Progress of the reaction What is the activation energy for the reverse reaction? A. -60 kj B. -20 kj C. +40 kj D. +60 kj - 2 -

9 7. Which of the factors below is not a condition necessary for equilibrium? A. a closed system B. a constant temperature C. equal forward and reverse reaction rates D. equal concentrations of reactants and products 8. In order for a chemical reaction to go to completion, how must the entropy and enthalpy change? Entropy Enthalpy A. increases increases B. increases decreases C. decreases increases D. decreases decreases 9. Consider the following equilibrium system: 2SO2( g) + O2( g) Æ 2SO3 ( g) Keq = If additional SO 2 is added to the system, what happens to the equilibrium and the value of K eq? Equilibrium K eq A. shifts left decreases B. shifts right increases C. shifts right no change D. no change no change OVER

10 10. Consider the following equilibrium system: N2( g) + 3H2( g) Æ 2NH3 ( g) Determine the changes in reaction rates as a catalyst is added. Forward Rate Reverse Rate A. increases increases B. increases decreases C. decreases increases D. decreases decreases 11. Consider the following equilibrium system: 2KClO3( s) Æ 2KCl( s) + 3O2( g) Which of the following is the equilibrium constant expression? A. K O eq = [ 2] 3 B. Keq = 1 [ O2] [ KClO ] C. K eq = 3 2 KCl O 3 2 [ ] [ 2] KCl O D. Keq = [ ] 2 [ ] 2 2 KClO [ 3]

11 12. Consider the following equilibrium: CO2( g) + 2H2O( g) Æ CH4( g) + 2O2( g) Which of the options below indicates that the reactants are favoured? A. K eq is zero. B. K eq is very large. C. K eq is slightly less than 1. D. K eq is slightly greater than Consider the following equilibrium: NO 2 4( g) + energy Æ 2NO2( g) How are K eq and [ N 2 O 4 ] affected by the addition of Ne (an inert gas) into the container at constant volume. (2 marks) K eq [ NO 2 4 ] A. no change no change B. no change increases C. increases decreases D. decreases increases 14. Consider the following equilibrium: Cl2( g) + 2NO( g) Æ 2NOCl( g) Keq = 5. 0 At equilibrium, [ Cl2]= 10. M and [ NO]= 20. M. What is the NOCl equilibrium? [ ] at A M B M C M D. 10 M OVER

12 15. Solid Ba( OH) 2 is added to water to prepare a saturated solution. Which of the following is true for this equilibrium system? A. [ anion]= [ cation] B. trial K sp is less than K sp C. blue litmus paper would turn red D. the rate of dissolving = the rate of crystallization 16. A saturated solution of PbI 2 was subjected to a stress and the following graph was obtained. [I ] Concentration [Pb 2+ ] t 1 Time Which stress was applied at time t 1? A. the addition of PbI 2 B. a temperature change C. an increase in volume D. the evaporation of water 17. Which of the following would be true when equal volumes of 0. 2 M NaBr and 0. 2 M AgNO 3 are combined? (2 marks) A. No precipitate forms. B. A precipitate of AgBr forms. C. A precipitate of NaNO 3 forms. D. Precipitates of both NaNO 3 and AgBr form

13 18. Using the solubility table, determine which of the following ions could not be used to separate S 2- from SO 4 2- by precipitation? A. Be 2+ B. Ca 2+ C. Ba 2+ D. Sr Which of the following is true when solid Na S 2 is added to a saturated solution of CuS and equilibrium is reestablished? [ ] increases. [ ] increases. [ ] does not change. [ ] does not change. A. S 2- B. Cu 2+ C. S 2- D. Cu Which of the following describes the relationship between the solubility product constant K sp () of PbI 2? ( ) and the solubility s A. K sp = s 2 B. K sp = 4s 3 C. s 3 = K sp 4 D. s = Ksp OVER

14 21. Which of the following saturated solutions will have the lowest [ S 2- ]? (2 marks) A. BaS B. CaS C. CuS D. ZnS 22. What is the solubility of SrF 2? (2 marks) A M B M C M D M 23. Which of the following represents the neutralization reaction between Ca( OH) 2( s ) and HCl ( aq)? A. HO 2 ( l) Æ H ( aq) + OH ( aq) B. Ca ( ) + 2Cl ( ) Æ CaCl aq aq 2( s) ( ) + Æ + C. Ca OH 2( s) 2HCl( aq) CaCl2( aq) 2H2O( l) D. Ca ( aq) + 2OH ( aq) + 2H ( aq) + 2Cl ( aq) Æ CaCl2 s + 2H2O l - ( ) ( ) 24. Which of the following solutions will have the lowest electrical conductivity? A M HI B M H2S C M NaOH D M NaNO3-8 -

15 25. Consider the following equilibrium: HCO + H O Æ H CO + H O Which of the following statements is true? (2 marks) A. Products are favoured because H 2 O is a stronger acid than H2CO3 B. Products are favoured because H 3 O + is a stronger acid than H2CO3 C. Reactants are favoured because HCO 3 - is a stronger base than H O 2 D. Reactants are favoured because H O 3 + is a stronger acid than H CO Which of the following factors of an acidic solution would affect its ph? I. the strength of the acid II. III. the concentration of the acid the temperature A. I and II only. B. II and III only. C. I and III only. D. I, II and III. 27. Consider the following equilibrium: + - HO( ) Æ HO ( aq) + OH ( aq) 2 2 l 3 + What changes occur to [ H 3 O ] and ph when NaOH is added? (2 marks) [ ] increases and ph increases. [ ] increases and ph decreases. [ ] decreases and ph increases. [ ] decreases and ph decreases. A. + H 3 O B. + H 3 O C. + H 3 O D. + H 3 O OVER

16 28. The ionization of water is endothermic. How is K w related to the temperature of water? A. K w increases as temperature increases. B. K w decreases as temperature increases. C. K w increases as temperature decreases. D. K w remains constant as temperature decreases. 29. Which of the following represents the dissociation equation of a salt in water? + - s aq aq A. KCl( ) Æ K ( ) + Cl ( ) B. Ca ( aq) + SO4 ( aq) Æ CaSO4( s) C. HCl + KOH Æ KCl + H 2 O l ( aq) ( aq) ( aq) ( ) D. 2Na( s) + 2H2O( l) Æ 2NaOH( aq) + H2( g) 30. Which of the following represents the equilibrium constant expression for the hydrolysis reaction that occurs in NaF ( aq)? A. K B. K C. K [ HF ][ - OH ] b = - [ F ] [ - F ][ + H O ] a = 3 [ HF] [ + ][ - ] eq = Na F NaF [ ] w = + - [ 3 ][ ] D. K H O OH 31. Which of the following salt solutions will be acidic? A. KClO 4 B. NH4Br C. NaHCO 3 D. Na2C2O4-10 -

17 32. The indicator phenol red will be red in which of the following solutions? A M HF B M HBr C M NH4Cl D M Na2CO3-33. Which of the following chemical indicators has a K a = ? A. methyl orange B. phenolphthalein C. thymolphthalein D. bromcresol green 34. At a certain point in a strong acid-strong base titration, the moles of H + are equal to the moles of OH -. This is a definition of which of the following? A. the end point B. the titration point C. the transition point D. the equivalence point 35. A ml sample of H2SO4( aq) is titrated with ml of M NaOH ( aq). What is the concentration of the H2SO4( aq)? (2 marks) A M B M C M D M 36. What is the complete ionic equation for the neutralization of 0. 1M Sr( OH) 2 ( aq ) with 0. 1M H2SO4 ( aq)? (2 marks) + A. H ( aq) + OH ( aq) Æ H 2 O l ( ) B. Sr ( aq) + SO4 ( aq ) Æ SrSO4 s () C. Sr ( aq) + 2OH ( aq) + 2H ( aq) + SO4 ( aq) Æ SrSO4 s + 2H2O l ( ) ( ) D. Sr ( aq) + 2OH ( aq) + 2H ( aq) + SO4 ( aq) Æ Sr ( aq) + SO4 ( aq) + 2H2O( l) OVER

18 37. Carbon dioxide gas in the atmosphere dissolves in normal rainwater. This causes normal rainwater to A. be slightly basic. B. have a ph slightly less than C. be unaffected and remain neutral. D. have a ph slightly greater than Consider the following equation: H + 2MnO + 5C O Æ 10 CO + 2Mn + 8H O 2 Identify the chemical species which is reduced. A. H + B. Mn 2+ C. MnO 4-2- D. C 2 O Consider the following equation: H AsO + 4Zn + 8H Æ AsH + 4Zn + 4H O 2 Which of the following is correct? A. Oxygen is reduced. B. Arsenic is oxidized. C. Zinc is the oxidizing agent. D. The reaction is a redox reaction. 40. What is the oxidation number of iron in magnetite, Fe O 3 4? A B. + 2 C D

19 41. Consider the following spontaneous reactions: Cd + 2Np Æ 3Cd + 2Np Cd + Pd Æ Pd + Cd Np + Ce Æ Np + Ce Which is the strongest oxidizing agent? (2 marks) A. Cd 2+ B. Ce 3+ C. Np 3+ D. Pd The volumetric analysis of hydrogen peroxide is often carried out by titrating with aqueous KMnO 4 in an acidic solution. The unbalanced formula equation for the reaction is: HO KMnO4 + H2SO4 Æ K2SO4 + MnSO4 + H2O + O2 Which of the following is the correct set of balancing coefficients? (2 marks) A. 1, 1, 1, 1, 1, 1, 1 B. 1, 2, 3, 1, 2, 4, 3 C. 2, 5, 3, 2, 1, 8, 5 D. 5, 2, 3, 1, 2, 8, Which of the following would not be found in all electrochemical cells? A. ions B. an anode C. a cathode D. two beakers OVER

20 Use the following diagram to answer questions 44 and 45. Volts Zn 1.0 M KNO3 Pb 1.0 M Zn M Pb As this cell operates, the cations move towards the A. Pb electrode and the electrode gains mass. B. Pb electrode and the electrode loses mass. C. Zn electrode and the electrode gains mass. D. Zn electrode and the electrode loses mass. 45. As the cell operates, the electrons flow towards the A. Zn electrode and the cell voltage increases over time. B. Pb electrode and the cell voltage decreases over time. C. Zn electrode and the cell voltage decreases over time. D. Pb electrode and the cell voltage remains constant over time

21 46. Consider the following electrochemical cell: Volts Cl 2(g) Ni 1.0 M KNO3 Pt (Inert) 1.0 M KCl 1.0 M Ni(NO 3 ) 2 What is the E for this cell? A V B V C V D V 47. An aqueous solution of CuSO 4 is electrolyzed using copper electrodes. Which of the following would correctly describe the changes in the mass of each electrode and the Cu 2+ [ ] in solution? (2 marks) Mass of Anode Mass of Cathode [ Cu 2+ ] A. increases increases increases B. increases decreases stays the same C. decreases increases stays the same D. decreases decreases increases OVER

22 48. Which of the following are produced at the anode and cathode in the electrolysis of aqueous potassium sulfate using carbon electrodes? Anode Cathode A. potassium oxygen B. hydrogen oxygen C. oxygen hydrogen D. sulfur potassium This is the end of the multiple-choice section. Answer the remaining questions directly in this examination booklet

23 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 calculations, full marks will NOT be given for providing only an answer. 1. Using the axes below, sketch a PE diagram for the reacting system where: (3 marks) DH =-30 kj mol E a = 50 kj mol PE (kj) Progress of the reaction OVER

24 2. Consider the following reaction mechanism: Step 1 2NO Æ N2O2 Step 2 NO H2 Æ NO 2 + HO 2 Step 3 NO 2 + H2 Æ N2 + HO 2 a) Determine the overall reaction. (2 marks) Overall Reaction: b) Identify a reaction intermediate. 3. Consider the following equilibrium: CH4( g) + H2O( g) Æ CO( g) + 3H2( g) K eq Temperature C C Is the forward reaction in this equilibrium exothermic or endothermic? Explain your answer. (3 marks)

25 4. Consider the following equilibrium: CO( g) + Cl2( g) Æ COCl2( g) At equilibrium, the system contains mol CO, mol Cl2 and mol COCl 2 in a 2. 0 L container. Calculate the value of K eq. (3 marks) 5. Calculate the mass of NaI necessary to begin precipitation of Cu + from a ml sample of M CuNO 3. (4 marks) OVER

26 ( ) ( ) 6. When a solution of Na2CO3 ( aq) is mixed with a solution of Ca NO 3 2 precipitate forms. a) Write the net ionic equation for the precipitation reaction. aq a b) Explain what happens to the precipitate when HCl is added. (3 marks) 7. Write a chemical reaction showing an amphiprotic anion reacting as a base in water. (2 marks)

27 8. Calculate the poh of M Sr( OH) 2. (2 marks) 9. A M diprotic acid has a ph of Calculate its K a value. (5 marks) OVER

28 10. The following two experiments were conducted: Titration A: A strong acid was titrated with a strong base. Titration B: A weak acid was titrated with a strong base. a) How does the ph at the equivalence point of Titration B compare with the ph at the equivalence point of Titration A? b) Explain your answer to a). (2 marks) 11. Balance the following redox reaction. (4 marks) - + Æ + ( ) Sb NO Sb O NO acidic

29 12. A 1. 0 M HCl solution is electrolyzed using a copper anode and an inert carbon cathode. Predict the half-reactions that will occur and describe what you would observe at each electrode. (4 marks) Anode half-reaction: Anode observation: Cathode half-reaction: Cathode observation: END OF EXAMINATION

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

32 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.

33 PERIODIC TABLE OF THE ELEMENTS Be Beryllium B Boron C Carbon N Nitrogen O Oxygen F Fluorine H He Hydrogen Helium Ne Neon Atomic Number Symbol Name Atomic Mass 3 Li Lithium Si Silicon Na Mg Sodium Magnesium Aluminum Silicon Phosphorus Sulphur Chlorine Argon Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Ni Cu Zn Ga Ge As Se Br Kr Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton Nickel Rb Y Zr Nb Mo Tc Ru Rh Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon (98) Sr Pd Ag Cd In Sn Sb Te I Xe Cs Ba Cesium Barium Lanthanum Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury Thallium Lead Bismuth Polonium Astatine Radon (209) (210) (222) Fr Ra La Hf Ac Rf Ta Francium Radium Actinium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium (223) (226) (227) (261) (262) (263) (262) (265) (266) 74 W Db Sg Re Os Bh Hs Ir Mt Pt Au Hg Tl Pb Bi Po At Rn 58 Ce Cerium Based on mass of C 12 at 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) Values in parentheses are the masses of the most stable or best known isotopes for elements which do not occur naturally. 1

34 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

35 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

36 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

37 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

38 RELATIVE STRENGTHS OF BRØNSTED-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 H2O H Cr H 6 2O OH Benzoic C6H5COOH H + C6H5COO Hydrogen oxalate ion HC2O 4 H + C2O Ethanoic, acetic CH3COOH H + CH3COO Dihydrogen citrate ion H C H O H + HC H O ( ) + ( ) ( ) Hexaaquoaluminum ion, aluminum ion Al H2O H Al H O OH Carbonic ( CO + HO 2 ) HCO H + HCO Monohydrogen citrate ion HC6H5O 7 H + C6H5O Hydrogen sulphite ion HSO3 H + SO Hydrogen sulphide H2S + H + HS Dihydrogen phosphate ion H2PO4 H + HPO Boric H3BO3 H + H2BO Ammonium ion + + NH4 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 6

39 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

40 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 + 2e 2SO HO + 2H + 2e 2HO MnO + 8H + 5e Mn + 4H O Au + 3e Au( s) + BrO3 + 6H + 5e 1 Br 2 2( l) + 3H2O ClO + + 8H + 8e Cl + 4H O Cl2( g) e Cl O 2 2( g) H e H2O + 2+ MnO2( s) ( s) 2 Br2( l) e Br AuCl4 e Au( s) Cl + 3 ( g) 2 2+ Hg e Hg( l) O 2 2( g) H M e H2O Ag e Ag( s) 1 2+ Hg 2 2 e Hg( l) O ( ) H e 2 g H2O ( s) I2( s) e I + Cu e Cu( s) ( s) + 3 2O 2+ Cu e Cu( s) SO + 4H + 2e H SO + H O Cu + e Cu Sn + 2e Sn S( ) + 2H + 2e s H2S( g) + 2H + 2e H2( g) 2+ Pb + 2e Pb( s) 2+ Sn + 2e Sn( s) Ni + 2e Ni( s) HPO + + 2H + 2e HPO + HO Co e Co( s) + Se( s) H e H2Se Cr + e Cr H Fe e Fe( s) 2 Ag2S( s) e Ag() s S 3+ Cr e Cr( s) 2+ Zn e Zn( s) ( s) Te HO 2 e H2( g) OH 2+ Mn e Mn( s) 3+ Al e Al( s) 2+ Mg e Mg( s) + Na e Na( s) 2+ Ca e Ca( s) 2+ Sr + e Sr( s) 2+ Ba e Ba( s) + K e K( s) + Rb e Rb( s) + Cs e Cs( s) + Li e Li( s) O + 2e H + 2OH ( 10 M) Te H e H Cr O + 14H + 6e 2Cr + 7H O H + 2e Mn + 2H O IO + 6H + 5e I + 3H O NO + 4H + 3e NO + 2H O ( 10 ) NO + 4H + 2e N O + 2H O Fe e Fe MnO + 2H O + 3e MnO + 4OH HSO H e S H Overpotential Effect WEAK STRENGTH OF REDUCING AGENT STRONG