U.S. National Chemistry Olympiad

1988 National Test

Notes

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1. A verbal statement or a mathematical equation that summarizes a broad variety of observations and experience is called a(n)

(A) hypothesis
(B) law
(C) theory
(D) evolution

2. The term that is related to the reproducibility or repeatability of a measurement is:

(A) accuracy
(B) systematic
(C) precision
(D) qualitative

3. A sample of carbon dioxide that undergoes a transformation from solid to liquid to gas would undergo

(A) a change in mass
(B) a change in density
(C) a change in composition
(D) no change in physical properties

4. A solution is a(n)

(A) compound
(B) pure substance
(C) homogeneous mixture of substances
(D) heterogeneous mixture of substances

5. Dalton's theory of atoms consisted of all of the following postulates except

(A) elements are composed of indivisible particles called atoms.
(B) atoms of different elements have different properties.
(C) atoms combine in fixed ratios of whole numbers.
(D) the volumes of gases consumed and produced are in ratios of small whole numbers.

6. Cathode rays are

(A) a beam of positively charged particles.
(B) nuclei of helium atoms
(C) fast-moving neutrons
(D) streams of electrons

7. There are two stable isotopes of carbon. They differ with respect to

(A) atomic mass
(B) number of protons
(C) atomic number
(D) electron configuration

8. One ⁴⁰Ca²⁺ ion contains

(A) 2 protons
(B) 18 electrons
(C) 21 neutrons
(D) 2 electrons

9. The average atomic weight of elemental copper is reported as 63.5. Copper consists of two stable isotopes, ⁶³Cu and ⁶⁵Cu. Approximately what percent of naturally occurring copper is the ⁶³Cu isotope?

(A) 30
(B) 50
(C) 70
(D) 90

10. The best molecular formula for sulfur in its naturally occurring state is

(A) S
(B) S₂
(C) S₆
(D) S₈

11. The correct formula for iron(III) sulfide is

(A) Fe₃S₂
(B) FeS
(C) Fe₂S₃
(D) Fe₃S

12. What is the percentage by weight of potassium in the compound KNO₃?

(A) 38.7
(B) 39.1
(C) 56.6
(D) 63.1

13. If 18.5 moles of the liquid compound C₂Cl₄ are required for a particular chemical reaction, what volume should be taken? The density of C₂Cl₄ is 1.63 g/ml.

(A) 11.3 mL
(B) 30.2 mL
(C) 1.88 L
(D) 5.01 L

14. The principle of DuLong and Petit states that the product of the atomic weight of a solid element and its specific heat, measured at room temperature, is approximately 6.2 calories per mole per degree (25.9 Joules per mole per degree). A prospector has asked you to identify the most abundant element in a mineral sample he has found. You chemically separate the most abundant element and determine that its specific heat is 5.24 x 10¯² cal per gram per degree. The element most likely is

(A) Au
(B) Fe
(C) Ni
(D) Sn

15. In the complete combustion of octane with oxygen, represented by the unbalanced chemical equation,

C₈H₁₈ + O₂ ----> CO₂ + H₂O

one gram of octane will yield what mass of water?

(A) 0.079 g
(B) 1.4 g
(C) 18 g
(D) 162 g

16. The first chemical compound of a rare gas element was prepared in 1962. Since then several such compounds have been prepared and characterized. What is the empirical formula of a compound of Xe which is 67.2% Xe and 32.8% O by mass?

(A) XeO₂
(B) XeO₃
(C) XeO₄
(D) XeO₅

17. The commercial production of phosphoric acid from phosphate ores can be represented by the equation

Ca₃(PO₄)₂ + 3 SiO₂ + 5 C + 5 O₂ + 3H₂O ---> 3 CaSiO₃ + 5 CO₂ + 2 H₃PO₄.

If 1.0 kg. each of calcium phosphate and silica are used with sufficient excess of carbon, oxygen, and water, what quantity of phosphoric acid can be produced?

(A) 0.31 kg
(B) 0.63 kg
(C) 1.0 kg
(D) 1.1 kg

18. A power company burns approximately 474 tons of coal per day to produce electricity. If the sulfur content of the coal is 1.30% by weight, how many tons of SO₂ are dumped into the atmosphere each day?

(A) 12.3
(B) 6.16
(C) 3.08
(D) 0.19

19. A 50.0 mL aliquot of a sulfuric acid solution was treated with barium chloride and the resulting BaSO₄ was isolated and weighted. If 0.667 gram of BaSO₄ was obtained, what was the morality of the H₂SO₄?

(A) 0.00700 M
(B) 0.0286 M
(C) 0.0572 M
(D) 1.43 M

20. Which statement about the electron configuration of electrons in the Cs atom is correct?

(A) The outermost two electrons are paired in the same atomic orbital.
(B) The 4f shell is completely full.
(C) Only one of the 55 electrons is involved in most interactions of Cs with other atoms.
(D) The 4f shell is only partially filled.

21. Those elements in which unpaired valence electrons are in atomic orbitals designated by l = 2 are described as

(A) representative elements.
(B) lanthanide elements.
(C) transition elements.
(D) halogens.

22. Which of the following species is smallest is size?

(A) Cl
(B) I
(C) Cl¯
(D) I¯

23. Which atom has the lowest second ionization energy?

(A) Mg
(B) Na
(C) K
(D) Ar

24. What formula would be expected for a binary compound formed between strontium and nitrogen?

(A) SrN
(B) Sr₂N
(C) SrN₃
(D) Sr₃N₂

25. The formation of an ionic compound from its elements can be understood in terms of steps, each involving a certain energy input or output. Which energy step usually dominates all others in the formation of a stable ionic compound?

(A) ionization energy
(B) electron affinity
(C) lattice energy
(D) dissociation energy

26. Which compound violates the simple octet rule for electron distribution around the central atom?

(A) CO₂
(B) NF₃
(C) OF₂
(D) PF₅

27. Predict the Cl-Sn-Cl bond angles in SnCl₃¯.

(A) 90°
(B) 109°
(C) 180°
(D) between 90° and 109°

28. The ethylene molecule, C₂H₄, is

(A) linear.
(B) planar.
(C) pyramidal.
(D) shaped like two tetahedra joined at the points.

29. Which diatomic molecular species should have the greatest bond dissociation energy?

(A) H₂
(B) F₂
(C) NO
(D) N₂

30. If a sample of silicon is "doped" with a small amount of boron, what sort of material results?

(A) insulator
(B) metallic conductor
(C) n-type semiconductor
(D) p-type semiconductor

31. Tin(II) chloride is a solid with a melting point of 246 °C; tin(IV) chloride is a liquid with a freezing point of - 33 °C. These properties can be explained by

(A) the greater covalence of SnCl₄ compared with that of SnCl₂.
(B) the greater formula weight of SnCl₄ compared with that of SnCl₂.
(C) the greater degree of ionic character of SnCl₄ compared with that of SnCl₂.
(D) the greater number of ionic and molecular particles in SnCl₄ with SnCl₂.

32. The Haber process for preparing ammonia involves the direct conversion of hydrogen and nitrogen gases at high temperature and pressure using a catalyst:

N₂(g) + 3 H₂(g) ----> 2 NH₃(g)

How many liters of ammonia could be prepared from a mixture of 19.0 liters of nitrogen and 34.7 liters of hydrogen, assuming complete conversion and with identical conditions of temperature and pressure?

(A) 19.0
(B) 23.1
(C) 34.7
(D) 38.0

33. A sample of gas confined in a bulb at 1 atm pressure and 25 °C is measured to have a density of 1.309 grams per liter. The gas is

(A) O₂
(B) N₂
(C) Ne
(D) CH₄

34. The diagram below represents a 2-dimensional model of a sample of water at a particular temperature.

Which diagram best represents the same system following the addition of a sample of crystalline urea (NH₂CONH₂)?

35. A solution is prepared by dissolving 1.20 g of sodium bisulfate, NaHSO₄, in 1.00 kg of distilled water. The osmotic pressure of the solution is determined to be 364 mm Hg or 0.48 atm. Which equation best describes how sodium bisulfate behaves in water?

(A) NaHSO₄(s) + H₂O(l) ---> will not dissolve
(B) NaHSO₄(s) + H₂O(l) ---> NaHSO₄(aq)
(C) NaHSO₄(s) + H₂O(l) ---> Na⁺(aq) + HSO₄¯(aq)
(D) NaHSO₄(s) + H₂O(l) ---> Na⁺(aq) + H⁺(aq) + SO₄²¯(aq)

36. Given the following standard enthalpies of formation: CO₂(g), -394 kJ/mol; H₂O(l), -286 kJ/mol; C₄H₃(g), 16.0 kJ/mol. Calculate the heat of combustion of one mole of C₄H₃ if the equation describing the process is

C₄H₃(g) + 6 O₂ ---> 4 CO₂(g) + 4 H₂O(l).

(A) -2736 kJ
(B) -696 kJ
(C) 696 kJ
(D) 2736 kJ

37. For which process is DS° negative?

(A) the rusting of iron
(B) the mixing of two ideal gases
(C) the sublimation of carbon dioxide
(D) a partially soluble salt dissolving in water

38. Which experiment will yield DG°?

(A) measuring the temperature change when a known mass of sodium hydroxide solid is added to a known mass of distilled water.
(B) continually measuring the pH during the titration of a strong base of known concentration with a known volume of a weak acid of unknown concentration.
(C) measuring the volume of a known mass of oxygen at STP
(D) measuring the height of a plug of mercury above a sample of air in a capillary tube, at several temperatures.

39. A 20.0-liter vessel initially contains 0.50 mole each of H₂ and I₂ gases. These substances react and finally reach an equilibrium condition. Calculate the equilibrium condition of HI if K_eq = 49.

(A) 0.78 M
(B) 0.039 M
(C) 0.033 M
(D) 0.021 M

40. In the laboratory the equilibrium constant for a particular reaction can be measured at different temperatures. Plotting the data yields the graph shown below: Which of the following statements is false? (Note: The notation 4.40e-4 is equivalent to 4.40 x 10¯⁴.)

(A) DS° can be obtained from the y-intercept.
(B) The slope of the line is equal to -(DH°/R)
(C) The reaction is endothermic.
(D) The free energy of the reaction is positive.

41. The half-life of radioactive ⁵⁵Cr is 1.8 hours. the delivery of a sample of this isotope from the reactor to your laboratory requires about 10.8 hours. What is the minimum amount of such material that should be shipped in order that you receive 1.0 milligram of ⁵⁵Cr?

(A) 128 mg
(B) 64 mg
(C) 32 mg
(D) 11 mg

42. The following initial rate data were collected for the reaction:

aA(g) + bB(g) ---> cC(g) + dD(g);

The rate law that best fits this data is

(A) rate = k[A]¹[B]¹
(B) rate = k[A]¹[B]²
(C) rate = k[A]²[B]¹
(D) rate = k[A]²[B]⁰

43. If Experiment #1, in problem 42, is allowed to proceed further, it is found that after 5 days the concentrations of all the species are;

Which is the correct balanced chemical equation for the reaction?

(A) A(g) + B(g) ---> C(g) + D(g)
(B) 2 A(g) + B(g) ---> C(g) + 2 D(g)
(C) 2 A(g) + 2 B(g) ---> C(g) + 3 D(g)
(D) A(g) + 2 B(g) ---> C(g) + 3 D(g)

44. Which statement concerning the stratospheric ozone levels is false?

(A) The concentration of stratospheric ozone over Antarctica decreases in the summer and returns to normal levels during the winter.
(B) Chlorine radicals derived from chlorofluorocarbons have been implicated as causing the drop in ozone levels.
(C) The ozone layer is important because it absorbs high energy solar radiation which could cause damage to plants and animals.
(D) New research suggests that heterogeneous reactions may play a crucial role in changing the levels of ozone in the stratosphere.

45. A solution forms a precipitate upon addition of NaOH solution. When the precipitate is dissolved in HCl and treated with KCNS solution, a deep-red color is observed. The original solution probably contained the cation

(A) Fe²⁺
(B) Cr³⁺
(C) Fe³⁺
(D) Mn²⁺

46. A precipitate forms when HCl is added to an aqueous solution, but the precipitate dissolves when a small amount of hot water is added. The original solution probably contained the cation

(A) Ag⁺
(B) Pb²⁺
(C) Hg₂²⁺
(D) Hg²⁺

47. Al(OH)₃ is an insoluble solid whose K_sp = 1.9 x 10¯³³. What is the maximum concentration of OH¯ which can exist in 0.10 M AlCl₃ solution without causing Al(OH)₃ to precipitate?

(A) 2.7 x 10¯¹¹
(B) 1.4 x 10¯¹⁰
(C) 8.7 x 10¯⁸
(D) 3.1 x 10¯⁷

48. The compound CaO is

(A) an acid anhydride.
(B) a basic anhydride.
(C) an amphoteric oxide.
(D) an alkaline earth metal.

49. When H₂PO₄¯ behaves as a Bronsted base it becomes

(A) H₃PO₄
(B) HPO₄²¯
(C) HPO₃
(D) PO₄³¯

50. What is the solubility of MgF₂ in water at 25 °C if its K_sp is 6.4 x 10¯⁹?

(A) 6.4 x 10¯⁹
(B) 5.7 x 10¯⁵
(C) 1.9 x 10¯³
(D) 1.2 x 10¯³

51. Which salt will produce an acid solution when dissolved in water?

(A) NH₄Cl
(B) NaCN
(C) NH₄CN
(D) NaCl

52. What is the pH of a 1.0 x 10¯⁹ molar HCl solution?

(A) 5
(B) 6
(C) 7
(D) 9

53. Forty milliliters (40.00 mL) of 0.0900 M NaOH is diluted to 100.00 mL with distilled water and 30.00 mL of 0.1000 M HCl are added. The pH of the resulting solution is

(A) 9.57
(B) 11.66
(C) 12.18
(D) 12.38

54. Five acids are listed below in order of decreasing acid strength:

HCl > HC₂H₃O₂ > HCN > H₂O > NH₃.

For which of the following reactions must the numerical value of the equilibrium constant be less than one?

(A) HCl + H₂O <==> H₃O⁺ + Cl¯
(B) H₂O + NH₂¯ <==> NH₃ + OH¯
(C) HC₂H₃O₂ + OH¯ <==> C₂H₃O₂¯ + H₂O
(D) HCN + C₂H₃O₂¯ <==> HC₂H₃O₂ + CN¯

55. What is the pH at the equivalence point in a titration of 0.020 M NH₃(aq) with 0.020 M HBr(aq)? For ammonia, K_b = 1.8 x 10¯⁵.

(A) 5.5
(B) 5.6
(C) 7.0
(D) 8.5

56. A solution of a weak monoprotic acid is titrated with a 0.10 M strong base solution. Using a pH meter fitted with a glass and calomel electrode, a titration curve is constructed from pH values versus volume of base added. What information could NOT be obtained from the titration curve alone?

(A) the pK_a of the acid
(B) the molecular weight of the acid
(C) the millimoles of acid in the solution
(D) the best buffering region for the system

57. Which would make the best aqueous buffer solution?

(A) 0.10 M NH₄OH / 0.10 M NH₄Cl
(B) 0.10 M NH₄OH / 0.40 M NH₄Cl
(C) 0.10 M NH₄OH / 0.10 M HCl
(D) 0.10 M KOH / 0.10 M HCl

58. What must be the ratio of acetate ion concentration to acetic acid concentration in an aqueous solution in order to provide a solution of pH 5? The K_a of acetic acid is 1.8 x 10¯⁵.

(A) 0.056
(B) 1.0
(C) 1.8
(D) 5.0

59. The compound often used to provide a mild oxidizing condition in water for bleaching and for disinfecting is

(A) NaCl
(B) Ca(ClO)₂
(C) HF
(D) PbCr₂O₇

60. The highest oxidation number commonly observed for any ionic or molecular transition metal species is

(A) 2
(B) 3
(C) 6
(D) 7

61. In the electrolytic refining of copper, an electric current is passed through a cell containing a pure copper electrode, an electrode of impure copper, and an aqueous solution containing copper cations. As metallic copper is removed from the impure electrode and redeposited on the pure copper electrode other metallic impurities are left behind if the potential of the cell is adjusted properly. In this process, the impure copper electrode is called the

(A) anode.
(B) cathode.
(C) reference electrode.
(D) indicating electrode.

62. Given the standard electrode (reduction) potentials:

Which pair of substances will react spontaneously?

(A) Ni²⁺ with Cr³⁺
(B) Ni with Cr³⁺
(C) Ni²⁺ with Cr
(D) Ni with Cr

63. A certain solid-state ion-selective electrode responds according to the equation

E = constant + 0.0591 log(a_M²⁺)^1/2

where a_M²⁺ is the activity of a divalent cation M²⁺. The potential of the electrode was measured to be 0.2500 V when the electrode was immersed in a solution of Cu²⁺ with an activity, a_Cu²⁺, of 2.0 x 10¯⁴ M. The potential of the electrode in an unknown solution of Cu²⁺ was 0.2796 V. The activity of the Cu²⁺ in the unknown solution is

(A) 2.0 x 10¯³ M.
(B) 6.3 x 10¯⁴ M.
(C) 6.3 x 10¯⁴ M.
(D) 2.0 x 10¯⁴ M.

64. A mixed precipitate of NaCl and KCl weighing 0.2076 gram was dissolved and titrated with silver nitrate. The titration required 28.50 mL of 0.1055 M AgNO₃. What was the weight percent of NaCl in the mixed precipitate?

(A) 78.40%
(B) 71.00%
(C) 43.90%
(D) 29.00%

65. The most precise meausurement of an atomic or molecular weight on the ¹³C mass scale is by

(A) gravimetric carbon-hydrogen analysis.
(B) equivalent weight titration.
(C) gas density measurements.
(D) mass spectrometry.

66. Indifferent electrolytic solutions, that is solutions containing ions other than those involved in the analysis, are used to wash precipitates rather than distilled water because

(A) the wash solution must be able to conduct electrical current.
(B) precipitates are too insoluble in distilled water.
(C) the indifferent electrolytes prevent peptization (dispersion) of the precipitated salt.
(D) the precipitated solid must be kept in the colloidal state.

67. A sample of iron ore, weighing 0.700 gram, is dissolved in nitric acid. The solution is then diluted with water, following which sufficient concentrated aqueous ammonia is added to quantitatively precipitate the iron as Fe(OH)₃. The precipitate is filtered then ignited and weighed as Fe₂O₃. If the weight of the ignited and dried precipitate is 0.541 gram, what is the weight percent of iron in the original iron ore sample?

(A) 27.0%
(B) 48.1%
(C) 54.1%
(D) 81.1%

68. Any saturated hydrocarbon that is classified as an octane must have the molecular formula

(A) C₈H₈
(B) C₈H₁₄
(C) C₈H₁₆
(D) C₈H₁₈

69. The presence of a carbon-carbon double bond in an orbanic molecule

(A) results in a linkage between carbon atoms that has a much smaller bond energy that that of single carbon to carbon bonds.
(B) results in greater reactivity, often with addition of two atoms to the two carbon atoms of the bond.
(C) makes the molecule much more polar and therefore much more soluble in water.
(D) allows internal molecular rotations so that there are no isomers with the same formula possible.

70. Proteins that catalyze a metabolic reaction are

(A) metabolites.
(B) micelles.
(C) substrates.
(D) enzymes.

Part II

Problems (52 percent)

1. (12 pts.) A 50.0 milliliter sample of 0.133 M vanadate, VO₃¯, is reacted with excess solid zinc metal which reduces the vanadium and produces 0.0040 mole of H₂(g) in a side reaction. A total of 0.915 grams of zinc is consumed in both reaction:

(a) Calculate the number of moles of zinc consumed.
(b) Calculate the number of zinc used to reduce the vanadium.
(c) Determine the identity of the vanadium species produced, V, V²⁺, V³⁺, or VO²⁺, and outline your reasoning.
(d) Write a balanced equation for the reaction of this vanadium species with permanganate ion in acid solution if the vanadium species is converted to its original form during this reaction.
(e) Calculate the volume of 0.0100 M potassium permanganate required for the reaction in (d).

2. (15 pts.) Nitrogen dioxide, NO₂, is a red-brown paramagnetic gas which dimerizes under certain conditions to diamagnetic dinitrogen tetroxide, N₂O₄. DH for the dimerization process at 25°C is - 57 kJ mol¯¹ and K_p is 5.85 x 10¯³ at this temperature.

(a) Calculate DG and DS at 25°C
(b) Write the equilibrium expression for this process.
(c) If 0.050 moles of NO₂ are placed in a 1.00 liter flask:

(1) calculate the pressure before any reaction occurs.
(2) calculate the pressures of NO₂ and N₂O₄ after equilibrium has been established.

(d) Discuss the effect which the following changes would have on the [NO₂] / [N₂O₄] ratio and explain your reasoning.

(1) The flask is heated to 50°C
(2) The pressure in the flask is doubled by adding He.
(3) The contents of the 1.00 liter flask are transferred to a 2.00 liter flask.

3. (15 pts.) The peroxodisulfate ion (S₂O₈²¯) reacts with iodide ions to form sulfate ions and a product which contains iodine.

(a) If 21.52 mL of a 0.0820 M solution of K₂S₂O₈ are consumed in the reaction with 0.586 gram of KI (FW = 166.0) write a balanced equation for the reaction between S₂O₈²¯ and I¯.

(b) A kinetic study of the initial stages of this reaction yields the following rate data:

[I¯] mol liter¯1	[S2O82¯] mol liter¯1	- d[S2O82¯]/dt mol liter¯1 sec¯1
2.50 (10¯3)	1.25 (10¯4)	3.94 (10¯10)
5.00 (10¯3)	1.25 (10¯4)	7.98 (10¯10)
5.00 (10¯3)	2.50 (10¯4)	15.96 (10¯10)

(1) Determine the kinetic order of each reactant and explain your reasoning.
(2) Calculate the rate constant and specify its units.
(3) Calculate the rate when the molarities of the iodide and peroxodisulfate ions are 1.20 (10¯³) and 7.50 (10¯⁴), respectively.

(c) Select the reaction scheme(s) which is (are) consistent with the above rate and stoichiometric data. Explain why each of the others is not consistent.

4. (10 pts.) A 0.853 gram sample of liquid ethanol (MW = 46.07) is added to a bomb (constant volume) calorimeter with a heat capacity of 654 Joules/°C. The calorimeter is charged with oxygen and placed in a bath containing 1.100 kilograms of H₂O at 25.12 °C. The combustion of the ethanol is initiated with an electric spark and, after the reaction is complete, the temperature is 29.49 °C.

(a) Write a balanced equation for the complete combustion of ethanol.
(b) Calculate the heat of combustion of this ethanol sample in kJ gram¯¹.
(c) Identify the thermodynamic quantity represented by this heat and calculate its value in kJ mol¯¹
(d) Identify a second thermodynamic quantity that can be obtained from these data alone and write and equation to calculate it.
(e) Calculate the enthalpy of formation of ethanol from the information above and the enthalpies of formation of carbon dioxide (-395.5 kJ mol¯¹) and gaseous water (-241.8 kJ mol¯¹).

EQUATIONS (16 percent)

5. Write net equations for all the following reactions. Use appropriate ionic and molecular formulas for the reactants and products and omit formulas for any species which do not take part in a reaction. Write structural formulas for any organic species. You need not balance the equations.

(a) Solid iron(II) sulfide is added to a solution of HCl.
(b) A solution of tin(II) nitrate is added to an acidified solution of potassium dichromate.
(c) A mixture of ethanol and methanoic (formic) acid is refluxed.
(d) Solid calcium carbonate is heated.
(e) Carbon-14 undergoes radioactive decay.
(f) Solid magnesium oxide is added to water.
(g) Gaseous chlorine is bubbled through a solution of sodium hydroxide.
(h) Aqueous ammonia is added to solid zinc(II) hydroxide.

ESSAY QUESTIONS (32 percent)

6. (12 pts.) Many neutral covalent compounds are known with the general formula EF₄.

(a) Identify (by name and symbol) non-transition elements from three different families of the periodic table which form neutral compounds with this formula.
(b) Sketch Lewis electron-dot structures for the compounds formed by each of the species identified in (a).
(c) Identify which (if any) of the compounds sketched in (b) is (are) polar. Explain your choice(s).
(d) Write equations for the reactions (if any) of the species sketched in (b) in water.

7. (12 pts.) One way of determining the formula mass of an unknown compound is to measure its effect on the freezing point of an appropriate solvent.

(a) Write an equation that could be used to calculate the formula mass from freezing point measurements and define each term.
(b) List criteria which could be used to select an "appropriate" solvent and explain why each is significant.
(c) Predict how the calculated formula mass would compare with the true value (higher, lower or the same) for each of the following factors and explain your reasoning.

(1) Some of the solute remains undissolved during the freezing point determination.
(2) The freezing points of the solvent and solution are obtained with a thermometer which reads two degrees too high.
(3) Some of the solvent evaporates between the time it is first weighed and the time that the solute is added.
(4) The solute dissociates partially in the solvent.

(d) Two students made triplicate formula mass determinations by this method. Give the value that each student should report and explain your reasoning.

(1) Jessica obtained values of 145, 168, and 180.
(2) Jim obtained values of 248, 260, and 322.

8. (8 pts.) The first eight ionization energies (in eV) for a certain element are: 10.49. 19.72, 30.18, 51.37, 65.02, 220.43, 263.21, 309.41.

(a) Use atomic energy level concepts to account for:

(1) the general increase in these values.
(2) any discontinuities in the general trend.

(b) Identify the family to which this element most probably belongs and explain your reasoning.
(c) Identify by name and symbol one of the elements in the middle of this family.
(d) Assume that the ionization energies above were found for the element in (c) and predict the ionization energies for the elements around it (i.e., above, below, to the right, and to the left). Give reasons for your predictions.