* END OF EXAMINATION*. 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 1 JANUARY Ministry of Education Course Code = CH 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.. 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. Student Instructions 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.. (5) Question :.. (3) Question 3: 3.. (3) Question 4: 4.. (4) Question 5: 5.. (5) Question 6: 6.. (3) Question 7: 7.. (4) Question 8: 8.. (3)

3 Chemistry 1 JANUARY 004 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.. 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 1 PROVINCIAL EXAMINATION 1. This examination consists of two parts: Value Suggested Time PART A: 60 multiple-choice questions PART B: 8 written-response questions Total: 90 marks 10 minutes. 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 5 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 1 Provincial Examination. The calculator must be a hand-held device designed primarily for mathematical computations involving logarithmic and trigonometric functions and may be capable of performing graphing functions. 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. Students may have more than one calculator available during the examination, of which one may be a scientific calculator. 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. Calculators must not have any information programmed into memory that would not be acceptable in paper form. Specifically, calculators must not have any built-in notes, definitions, or libraries. There is no requirement to clear memories at the beginning of the examination but the use of calculators with built-in notes is equivalent to the use of notes in paper form. Any student deemed to have cheated on a provincial examination will receive a 0 on that examination and will be permanently disqualified from the Provincial Examination Scholarship Program.

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7 PART A: MULTIPLE CHOICE Value: 60 marks INSTRUCTIONS: Suggested Time: 80 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. 1. Which of the following factors only affects the rate of heterogeneous reactions? A. nature of reactants B. presence of a catalyst C. temperature of reactants D. surface area of reactants. Consider the following reactions in open systems: I. H( g) + O( g) Æ HO( g) II. CaCO3( s) Æ CaO( s) + CO( g) III. CaO( s) + SiO( s) Æ CaSiO3( s) IV. AgNO3( aq) + NaCl( aq) Æ NaNO3( aq) + AgCl( s) In which of the above could reaction rate be determined by D A. I B. II C. III D. IV bit tricky but consider that it is an OPI system mass of system D time? 3. Which of the following best describes activation energy? A. PE of activated complex B. PE of products PE of ( )- ( reactants) ( )- ( ) ( )- ( reactants) C. PE of reactants PE of activated complex D. PE of activated complex PE of OVER

8 4. Consider the following PE diagram: 300 PE (kj) Progress of the reaction Which of the following is true for the forward reaction? Reaction PE of Activated Complex kj ( ) DH kj ( ) A. catalyzed B. uncatalyzed C. catalyzed D. uncatalyzed Consider the following reaction: NO( g) + O( g) Æ NO( g) Why would this reaction probably involve more than one step? A. There is insufficient activation energy. B. This reaction has high activation energy. C. Reactions between gases are typically slow. D. A successful collision between more than two molecules is unlikely. - -

9 6. Consider the following reaction mechanism: Step 1 O3 Æ O + O Step O3 + O Æ O Which of the following could represent the activated complex for Step? A. O B. O C. O 3 D. O 4 7. In a certain reaction DH =-136 kj and Ea = 96 kj. Which of the following is true of its reverse reaction? I draw A. The reverse reaction is exothermic and Ea =-40 kj. B. The reverse reaction is exothermic and Ea kj C. The reverse reaction is endothermic and Ea kj D. The reverse reaction is endothermic and Ea = 3 kj. out diagram 1 8. Two experiments were performed involving the following equilibrium. The temperature was the same in both experiments. H( g) + I( g) Æ HI( g) In experiment A, 1. 0M H and 10. M I were initially added to a flask and equilibrium was established. In experiment B, 0. M HI was initially added to a second flask and equilibrium was established. Which of the following statements is always true about the equilibrium concentrations? A. [ H ] equals [ HI] in experiment A. B. HI [ ] equals [ H] in experiment A. C. [ HI] in experiment A equals [ HI] in experiment B. D. HI [ ] in experiment A equals 1 I [ ] in experiment B OVER

10 9. Which of the following reactions is accompanied by an increase in enthalpy? A. NO( g) + O( g) Æ NO( g) + 113kJ B. H( g) + O( g) -484 kj Æ HO( g) C. SO3( g) Æ SO( g) + O( g) DH = kj D. 4HCl( g) + O( g) Æ HO( g) + Cl( g) D H = kJ 10. Two substances are mixed and no reaction occurs. With respect to enthalpy and entropy, which of the following could explain why no reaction occurs? Enthalpy Entropy A. increases increases B. increases decreases C. decreases increases D. decreases decreases - 4 -

11 11. Consider the following reaction: tricks N( g) + 3H( g) Æ NH3( g) Which of the following diagrams represents what happens to the forward and reverse reaction rates when the catalyst Fe O 3 4 is added? A. B. forward rate forward rate reverse rate reverse rate C. D. forward rate forward rate reverse rate reverse rate 1. Temperature is gradually decreased then held constant in an exothermic equilibrium. Which of the following represents the change in the reverse reaction rate? A. B. reverse reaction rate reverse reaction rate time time C. D. reverse reaction rate reverse reaction rate time time OVER

12 13. Consider the following equilibrium: CaCO3( s) + HF( g) Æ CaF( s) + HO( g) + CO( g) Which of the following represents the equilibrium [ HO ]? A. [ HO ]= [ ] HF K CO eq[ ] [ ] Keq HF B. [ HO ]= CO [ ] HF CaCO3 C. [ [ ] [ ] HO ]= K CO CaF eq[ ][ ] [ ] [ 3] [ ][ ] Keq HF CaCO D. [ HO ]= CO CaF 14. Which of the following reactions will proceed furthest toward completion? A. Si( s) + O( g) Æ SiO( s) Keq = B. - HBr( g) Æ H( g) + Br( g) Keq = C. HO ( g) Æ H( g) + O( g) Keq = D. CH 4( g) + H( g) Æ CH 6( g) Keq = Consider the following equilibrium: CH4( g) + HO( g) Æ CO( g) + 3H( g) Keq = An equilibrium mixture of this system was found to contain the following concentrations: [ CH4]= M, [ HO ]= 0. 63M, [ CO]= 0. 5 M. What was the equilibrium H [ ]? A M B M C. 0. M D. 84. M - 6 -

13 Use the following information to answer questions 16 and 17. H( g) + I( g) Æ HI( g) Keq = at 440 C 16. If 50. M HI is initially placed into a container, what will be the equilibrium [ HI]? A M B. 39. M C. 44. M D. 48. M 17. If M H, 0. 10M I and M HI are placed into a container at 440 C, which of the following is true as equilibrium is approached? A. I B. HI C. H D. H [ ] decreases significantly. [ ] decreases significantly. [ ] decreases significantly. [ ] remains the same. 18. Which of the following solutes will produce a molecular solution? A. HCl B. Fe 3 S 3 C. HNO 3 D. CH3OH 19. Which of the following would best describe the solubility of a solute? A. litres per gram B. moles per litre C. grams per mole D. moles per second OVER

14 0. Which compound will have the lowest solubility? A. FeS B. CaSO 4 C. AgBrO 3 ( ) D. Fe NO Which of the following precipitates may form when equal volumes of 03. M AgNO, 03. M SrCl and 03. M Na CO are mixed together? 3 3 A. SrCO3 and AgCl B. AgCO3 and AgCl C. SrCO3 and AgCO3 D. SrCO3, AgCO3 and AgCl. An experiment is conducted to identify an unknown cation that is present in each of four beakers. Na CO 3 MnSO 4 ( NH 4 ) S RbNO 3 precipitate no precipitate precipitate no precipitate Which of the following could be the unknown cation? A. Ag + B. Fe +3 C. Ba + D. Be

15 3. Given the equilibrium reaction: NaNO3( aq) + AgS( s) Æ AgNO3( aq) + NaS( aq) Which K sp expression best describes the net ionic reaction? + - sp = [ ] [ ] A. K Ag S B. Ksp = 1 [ Ag + ] [ S - ] [ + Ag ] [ - S ] C. Ksp = Ag S [ ] [ AgNO ] [ Na S ] 3 D. Ksp = NaNO [ ] 3 4. For the salt PbCl, what will be the value for its K sp when a solution [ ] with Pb + solution at 0. M? A B C D of 0. M is mixed with an equal volume of a chloride ion 5. Two salt solutions were mixed and a Trial K sp was calculated to be The K sp value is From this information, which of the following is a true statement? K sp comparison Outcome A. Trial K < K precipitate forms sp sp B. Trial K > K precipitate forms sp sp C. Trial K < K no precipitate forms sp sp D. Trial K > K no precipitate forms sp sp OVER

16 - 6. A saturated solution of SrSO 4 has a SO 4 What is the [ Sr + ]? A M B M C M D M [ ] of M. 7. Which of the following is a common property of acid solutions? A. They have a ph > 7. B. They turn red litmus blue. C. They have a slippery feeling. D. They turn pink phenolphthalein colourless. 8. What is a general characteristic of all Brønsted-Lowry bases? A. They all accept H +. B. They all accept OH -. C. They will turn litmus a pink colour. D. They will react with acids to produce H gas. 9. Select the equation that best represents the reaction of CH3NH acting as a base with water ( aq) ( l) 3 3 ( aq) aq A. CH NH + H O Æ CH NH + OH ( ) ( aq) ( l) 3 ( aq) 3 aq B. CH NH + H O Æ CH NH + H O ( ) ( aq) ( l) 3 ( aq) aq C. CH NH + H O Æ CH NH OH + H ( ) + ( ) ( ) ( ) D. CH3NH aq + HO l Æ CH3 aq + NH3 ( aq) + OH ( aq)

17 30. Which of the following solutions will show the greatest electrical conductivity? A. 01. M HCl B. 05. M HCO3 C. 05. M H3BO3 D. 01. M HCO4 31. When comparing 0. 10M HPO - 4 and 0. 10M HCO - 4 as acids, which of the following is true? A. HC O 4 - is weaker and its ph is larger. B. HPO 4 - is stronger and its ph is larger. C. HPO 4 - is weaker and its ph is smaller. D. HC O 4 - is stronger and its ph is smaller. 3. Which of the following will have the smallest K b value? A. - IO 3 B. NH 3 C. CN - D. HPO Which of the following equations can be used to calculate poh? A. poh =-log K w B. poh = pk + ph w C. poh = pkw -ph + =- [ ] D. poh log H3O OVER

18 34. Which of the following solutions would have a ph =. 00? A M HCl B M HCN C M HSO4 D M NaOH 35. Consider the following acid equilibrium: + - ( aq) ( l) 3 ( aq) aq HCN + H O Æ H O + CN ( ) When writing the K a expression for HCN, why is HO ( l) not included in the expression? A. The concentration of HO ( l) is too large. B. The concentration of HO ( l) is too small. C. The concentration of HO ( l) does not exist. D. The concentration of HO ( l) is relatively constant. 36. What is the K b value for HC H O ? A B C D Which of the following describes the dissociation of calcium chloride? + - s aq aq A. CaCl( ) Æ Ca ( ) + Cl ( ) + - s aq aq B. CaCl( ) Æ Ca ( ) + Cl ( ) + - s aq aq C. CaCl( ) Æ Ca ( ) + Cl ( ) + - s aq aq D. CaCl( ) Æ Ca ( ) + Cl ( ) - 1 -

19 38. Which of the following properties is true for a solution of KNO 3? A. It is neutral. B. It is very basic. C. It is slightly basic. D. It is slightly acidic. 39. Which term does the following statement best describe? A mixture of a weak acid and its conjugate base, each with distinguishing colours. A. buffer B. titration C. indicator D. primary standard 40. A weak acid is titrated with a strong base using the indicator phenolphthalein to detect the end point. What is the approximate ph at the transition point? A. 70. B. 80. C. 90. D Which of the following titrations always results in ph = 70. at the equivalence point? A. A weak acid is titrated with a weak base. B. A weak acid is titrated with a strong base. C. A strong acid is titrated with a weak base. D. A strong acid is titrated with a strong base. 4. What volume of M NaOH is required to neutralize 5. 0 ml of M HBr? A ml B ml C ml D ml OVER

20 43. Which of the following graphs best describes the effect on the ph of a buffer solution when a small amount of acid is added at t 1? A. B. ph ph t 1 t 1 C. D. ph ph t 1 t A buffer solution is prepared using sufficient amounts of H S and NaHS. What limits this buffer s effectiveness when NaOH is added? [ ] A. HS [ ] [ ] [ ] B. HS - C. OH - + D. HO

21 45. What is produced when MgO is added to water? A. the metal Mg B. the acid HMgO ( ) C. the base Mg OH D. the amphiprotic species H MgO 46. Which of the following is a major source of NO ( g), which contributes to the problem of acid rain? A. a fuel cell B. an air conditioner C. a nuclear power plant D. the automobile engine 47. Identify the oxidizing agent in the following equation: A. H + B. Pb C. PbO D. - SO Pb + PbO + 4H + SO Æ PbSO + H O Which of the following is a redox equation? A. H + O Æ H O Æ + 4 ( 3) Æ B. Ag CrO Ag CrO C. Ag NH H Cl AgCl NH ( ) + Æ D. Mn OH HC H O Mn H O C H O OVER

22 49. Which of the following contains molybdenum with its highest oxidation number? A. MoCl 5 B. Mo S 3 C. - MoO 4 D. Mo Cl Which of the following skeletal half-reactions are not oxidations? - - I. ClO Æ ClO3 II. CH5OH Æ CH4O III. NO Æ NO4 A. I B. II C. III D. I and II 51. Consider the following half-reactions under standard conditions: - - I. ClO + e Æ ClO II. PbSO + e Æ Pb + SO 3+ - III. Fe + 3e Æ Fe In an experiment when ClO and Fe were combined, they reacted. In a second experiment when PbSO 4 and Fe were combined, there was no observable change. Which of the following shows the reduction half-reactions I, II and III in order of decreasing E? A. I, II, III B. I, III, II C. II, III, I D. III, II, I

23 5. Which of the following combinations will react spontaneously? A. I + Cu B. Pb Ag C. Zn Mg D. Sn + Ni 53. Consider the following skeletal redox equation for a reaction in basic solution: Æ + ( ) Zn NO ZnO NH basic 3 Which of the following best represents the reduction half-reaction occurring in this solution? Æ Æ Æ - + A. 9H + NO + 8e Æ NH + 3H O B. 3H O NO 5e NH 6OH C. 6H O NO 8e NH 9OH D. 4OH Zn e ZnO H O OVER

24 Use the following diagram to answer questions 54 to 56. Volts Pt (inert) 1.0 M KNO3 Ag 1.0 M K Cr O 7 (acidified) 1.0 M AgNO Which of the following represents the overall cell reaction? Æ Æ Æ A. Cr O + H + Ag Æ Ag + Cr + H O B. Cr O H Ag Ag Cr H O C. Cr O H Ag Ag Cr H O D. Cr O H Ag Ag Cr H O 55. What happens to the ph at each electrode? ph at Anode ph at Cathode A. increases decreases B. increases increases C. stays the same decreases D. stays the same increases 56. What is the cell voltage at equilibrium? A V B V C V D V

25 57. Consider the following diagram: Volts Zn 1.0 M KNO3 Co 1.0 M Zn(NO 3 ) 1.0 M Co(NO 3 ) Which of the following best describes the Co + ion movement and the mass of the zinc electrode as the cell operates? Co + movement Mass of zinc electrode A. toward the Co electrode increases B. toward the Co electrode decreases C. toward the Zn electrode increases D. toward the Zn electrode decreases OVER

26 58. Which of the following would protect an iron pipeline from rusting? A. connecting it to a solution of silver nitrate B. connecting it to the positive terminal of a direct current power supply C. connecting it to the negative terminal of a direct current power supply D. connecting it to electrodes made of copper which are buried beside the pipeline 59. Which of the following best describes a car battery as it is being recharged? A. It is an electrolytic cell. B. It is an electrochemical cell. C. It is an example of a short circuit. D. It is a system moving to a state of lower potential energy. 60. What are the most likely products of the electrolysis of 1. 0M MgI using inert electrodes? A. H and I B. Mg and I C. H and O D. Mg and O This is the end of the multiple-choice section. Answer the remaining questions directly in this examination booklet

27 PART B: WRITTEN RESPONSE Value: 30 marks INSTRUCTIONS: Suggested Time: 40 minutes You are 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. The release of O ( g ) resulting from the decomposition of bleach was measured in two different experiments. Data was collected and the following graph was drawn: Volume O ( ml) Experiment 1 Experiment Time (min) a) Calculate the average rate of reaction for each experiment. ( marks) Experiment 1: Experiment : b) Identify a variable from Experiment 1 and how it was changed to produce the Xdifferent reaction rate for Experiment. Explain using collision theory. (3 marks) OVER

28 . Consider the following equilibrium system: C( s) + H ( g) Æ CH4( g) DH =- 75 kj State three different ways to make more C ( s) react. (3 marks) i) ii) iii) 3. Sufficient NaSO4( s) is added to M Ba NO3 ( ) to cause a precipitate to form. a) Write the net ionic equation for the precipitate formation. (1 mark) - b) Calculate the [ SO 4 ] at the moment the precipitate starts to form. ( marks) 4. a) Write the equation to represent the reaction that results when NH 4 + ions are mixed with HCO 3 - ions. ( marks) b) Identify the two bases in the reaction in part a). (1 mark) c) Predict whether the reaction will favour the reactants or products. Justify your answer. (1 mark) Prediction: Justification: - -

29 5. Calculate the ph of M NH4I. Start by writing the equation for the predominant equilibrium reaction. (5 marks) OVER

30 ( 4( s) ) 6. A solution of NaOH ( aq) was standardized by titration using oxalic acid HCO as the primary standard. The following data was collected: Mass of HCO4( s) used = 1. 0g Volume of NaOH( aq) used = 40. 6mL Calculate the concentration of the NaOH ( aq). (3 marks) 7. Balance the following skeletal redox equation in acidic solution: (4 marks) MnO + As O Æ Mn + AsO (acidic) - 4 -

31 8. Draw an electrolytic cell that could be used to plate an iron ring with gold. Be sure to include all of the necessary parts. In addition, label the anode, solution used and composition of the electrodes. (3 marks) END OF EXAMINATION - 5 -

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33 Data Booklet CHEMISTRY 1 Work done in this booklet will not be marked. Ministry of Education Revised January 000

34 CONTENTS Page Table 1 Periodic Table of the Elements 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 5 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.

35 PERIODIC TABLE OF THE ELEMENTS Be Beryllium B Boron C Carbon N Nitrogen O Oxygen F Fluorine 19.0 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 (09) (10) () Fr Ra La Hf Ac Rf Ta Francium Radium Actinium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium (3) (6) (7) (61) (6) (63) (6) (65) (66) 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 1 at Th Thorium Pr Praseodymium Pa Protactinium Nd Neodymium U Uranium Pm Promethium (145) 93 Np Neptunium (37) 6 Sm Samarium Pu Plutonium (44) 63 Eu Europium Am Americium (43) 64 Gd Gadolinium Cm Curium (47) 65 Tb Terbium Bk Berkelium (47) 66 Dy Dysprosium Cf Californium (51) 67 Ho Holmium Es Einsteinium (5) 68 Er Erbium Fm Fermium (57) 69 Tm Thulium Md Mendelevium (58) 70 Yb Ytterbium No Nobelium (59) 71 Lu Lutetium Lr Lawrencium (6) Values in parentheses are the masses of the most stable or best known isotopes for elements which do not occur naturally. 1

36 ATOMIC MASSES OF THE ELEMENTS Based on mass of C 1 at Values in parentheses are the mass number of the most stable or best known isotopes for elements that do not occur naturally. 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 (7) 7.0 (43) (10) (47) (51) (47) (6) 16.5 (5) (57) 19.0 (3) (6) (58) 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 (37) (59) (44) (09) (145) 31.0 (6) () (61) (98) Element Symbol Atomic Number Atomic Mass Element Symbol Atomic Number Atomic Mass

37 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 + Barium Mg + Magnesium Ca + Calcium Mn + Manganese(II), manganous Cr + Chromium(II), chromous Mn 4+ Manganese(IV) Cr 3+ Chromium(III), chromic Hg + Mercury(I)*, mercurous Cu + Copper(I)*, cuprous Hg + Mercury(II), mercuric Cu + Copper(II), cupric K + Potassium H + Hydrogen Ag + Silver H 3 O + Hydronium Na + Sodium Fe + Iron(II)*, ferrous Sn + Tin(II)*, stannous Fe 3+ Iron(III), ferric Sn 4+ Tin(IV), stannic Pb + Lead(II), plumbous Zn + Zinc Negative Ions (Anions) Br Bromide OH Hydroxide CO 3 Carbonate ClO Hypochlorite ClO 3 Chlorate I Iodide Cl Chloride HPO 4 Monohydrogen phosphate ClO Chlorite NO 3 Nitrate CrO 4 Chromate NO Nitrite CN Cyanide C O 4 Oxalate Cr O 7 Dichromate O Oxide** H 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 Sulphate HC O 4 Hydrogen oxalate, binoxalate S Sulphide HSO 4 Hydrogen sulphate, bisulphate SO 3 Sulphite HS Hydrogen sulphide, bisulphide SCN Thiocyanate HSO 3 Hydrogen sulphite, bisulphite 3

38 SOLUBILITY OF COMMON COMPOUNDS IN WATER The term soluble here means > 0.1 mol/l at 5 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 +,Cu + Soluble Low Solubility Sulphate, SO 4 All others Ag +,Ca +, Sr +,Ba +, Pb + Soluble Low Solubility Sulphide, S Alkali ions, H +, NH 4 +,Be +,Mg +,Ca +, Sr +,Ba + All others Soluble Low Solubility Hydroxide, OH Alkali ions, H +, NH 4 +, Sr + All others Soluble Low Solubility or or Phosphate, PO 4 3 Carbonate, CO 3 Sulphite, SO 3 Alkali ions, H +, NH 4 + All others Soluble Low Solubility 4

39 SOLUBILITY PRODUCT CONSTANTS AT 5 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 O 4 CaSO 4 CuI Cu IO 3 CuS FeS ( ) AgBr AgCl AgI ZnS ( ) Fe( OH) Fe OH 3 PbBr PbCl ( ) Pb IO 3 PbI PbSO 4 MgCO 3 ( ) Mg OH AgBrO 3 Ag CO 3 Ag CrO 4 AgIO 3 SrCO 3 SrF SrSO 4 K sp

40 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 HSO4 H + HSO4 very large + Hydronium Ion H O + 3 H + HO Iodic HIO + 3 H + IO Oxalic H C O + H + HC O ( ) + + 3, ( III) 6 + Fe( HO) ( 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 HO H Cr H 6 O OH Benzoic C6H5COOH H + C6H5COO Hydrogen oxalate ion HCO 4 H + CO Ethanoic, acetic CH3COOH H + CH3COO Dihydrogen citrate ion H C H O H + HC H O ( ) + ( ) ( ) Hexaaquoaluminum ion, aluminum ion Al HO H Al H O OH Carbonic ( CO + HO) HCO 3 H + HCO Monohydrogen citrate ion + 3 HC6H5O 7 H + C6H5O Hydrogen sulphite ion HSO3 H + SO Hydrogen sulphide HS + H + HS Dihydrogen phosphate ion HPO4 H + HPO Boric H3BO3 H + HBO Ammonium ion + + NH4 H + NH Hydrocyanic HCN + H + CN Phenol C 6H5OH H + C6H5O Hydrogen carbonate ion HCO3 H + CO Hydrogen peroxide HO H + HO Monohydrogen phosphate ion + HPO4 H + PO Water H O + H + OH Hydroxide ion OH H + O very small + 3 Ammonia NH H + NH very small WEAK STRENGTH OF BASE STRONG 6

41 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

42 STANDARD REDUCTION POTENTIALS OF HALF-CELLS Ionic concentrations are at 1M in water at 5 C. Oxidizing Agents Reducing Agents E ( Volts ) STRONG STRENGTH OF OXIDIZING AGENT WEAK Overpotential Effect F( g) + e F SO + e SO HO + H + e HO MnO + 8H + 5e Mn + 4H O Au + 3e Au( s) + BrO3 + 6H + 5e 1 Br ( l) + 3HO ClO + + 8H + 8e Cl + 4H O Cl( g) e Cl O ( g) H e HO + + MnO( s) ( s) Br( l) e Br AuCl4 e Au( s) Cl + 3 ( g) + Hg e Hg( l) O ( g) H M e HO Ag e Ag( s) 1 + Hg e Hg( l) O ( ) H e g HO 4 ( s) I( s) e I + Cu e Cu( s) ( s) + 3 O + Cu e Cu( s) SO + 4H + e H SO + H O Cu + e Cu Sn + e Sn S( ) + H + e s HS( g) + H + e H( g) + Pb + e Pb( s) + Sn + e Sn( s) Ni + e Ni( s) 06. HPO + + H + e HPO + HO Co e Co( s) + Se( s) H e HSe Cr + e Cr H + Fe e Fe( s) AgS( s) e Ag() s S 3+ Cr e Cr( s) + Zn e Zn( s) ( s) Te HO e H( g) OH + Mn e Mn( s) 3+ Al e Al( s) + Mg e Mg( s) + Na e Na( s) + Ca e Ca( s) + Sr + e Sr( s) + Ba e Ba( s) + K e K( s) + Rb e Rb( s) + Cs e Cs( s) + Li e Li( s) O + e H + OH ( 10 M) Te H e H Cr O + 14H + 6e Cr + 7H O H + e Mn + H O + 1. IO + 6H + 5e I + 3H O NO + 4H + 3e NO + H O ( 10 ) NO + 4H + e N O + H O Fe e Fe MnO + H O + 3e MnO + 4OH HSO H e S H Overpotential Effect WEAK STRENGTH OF REDUCING AGENT STRONG