2.8 End of Chapter 2: Atoms Molecules Ions

1. In the following drawing, the green spheres represent atoms of a certain element. The purple spheres represent atoms of another element. If the spheres touch, they are part of a single unit of a compound. Does the following chemical change represented by these symbols violate any of the ideas of Dalton’s atomic theory? If so, which one?

1. A sample of compound X (a clear, colorless, combustible liquid with a noticeable odor) is analyzed and found to contain 14.13 g carbon and 2.96 g hydrogen. A sample of compound Y (a clear, colorless, combustible liquid with a noticeable odor that is slightly different from X’s odor) is analyzed and found to contain 19.91 g carbon and 3.34 g hydrogen. Are these data an example of the law of definite proportions, the law of multiple proportions, or neither? What do these data tell you about substances X and Y?

2. The existence of isotopes violates one of the original ideas of Dalton’s atomic theory. Which one?

3. Dalton originally thought that all atoms of a particular element had identical properties, including mass. Thus, the concept of isotopes, in which an element has different masses, was a violation of the original idea. To account for the existence of isotopes, the second postulate of his atomic theory was modified to state that atoms of the same element must have identical chemical properties.

4. How are electrons and protons similar? How are they different?

5. How are protons and neutrons similar? How are they different?

6. Both are subatomic particles that reside in an atom’s nucleus. Both have approximately the same mass. Protons are positively charged, whereas neutrons are uncharged.

7. Predict and test the behavior of α particles fired at a “plum pudding” model atom.

8. (a) Predict the paths taken by α particles that are fired at atoms with a Thomson’s plum pudding model structure. Explain why you expect the α particles to take these paths.

9. (b) If α particles of higher energy than those in (a) are fired at plum pudding atoms, predict how their paths will differ from the lower-energy α particle paths. Explain your reasoning.

10. (c) Now test your predictions from (a) and (b). Open the Rutherford Scattering simulation and select the “Plum Pudding Atom” tab. Set “Alpha Particles Energy” to “min,” and select “show traces.” Click on the gun to start firing α particles. Does this match your prediction from (a)? If not, explain why the actual path would be that shown in the simulation. Hit the pause button, or “Reset All.” Set “Alpha Particles Energy” to “max,” and start firing α particles. Does this match your prediction from (b)? If not, explain the effect of increased energy on the actual paths as shown in the simulation.

11. Predict and test the behavior of α particles fired at a Rutherford atom model.

12. (a) Predict the paths taken by α particles that are fired at atoms with a Rutherford atom model structure. Explain why you expect the α particles to take these paths.

13. (b) If α particles of higher energy than those in (a) are fired at Rutherford atoms, predict how their paths will differ from the lower-energy α particle paths. Explain your reasoning.

14. (c) Predict how the paths taken by the α particles will differ if they are fired at Rutherford atoms of elements other than gold. What factor do you expect to cause this difference in paths, and why?

15. (d) Now test your predictions from (a), (b), and (c). Open the Rutherford Scattering simulation and select the “Rutherford Atom” tab. Due to the scale of the simulation, it is best to start with a small nucleus, so select “20” for both protons and neutrons, “min” for energy, show traces, and then start firing α particles. Does this match your prediction from (a)? If not, explain why the actual path would be that shown in the simulation. Pause or reset, set energy to “max,” and start firing α particles. Does this match your prediction from (b)? If not, explain the effect of increased energy on the actual path as shown in the simulation. Pause or reset, select “40” for both protons and neutrons, “min” for energy, show traces, and fire away. Does this match your prediction from (c)? If not, explain why the actual path would be that shown in the simulation. Repeat this with larger numbers of protons and neutrons. What generalization can you make regarding the type of atom and effect on the path of α particles? Be clear and specific.

16. (a) The Rutherford atom has a small, positively charged nucleus, so most α particles will pass through empty space far from the nucleus and be undeflected. Those α particles that pass near the nucleus will be deflected from their paths due to positive-positive repulsion. The more directly toward the nucleus the α particles are headed, the larger the deflection angle will be. (b) Higher-energy α particles that pass near the nucleus will still undergo deflection, but the faster they travel, the less the expected angle of deflection. (c) If the nucleus is smaller, the positive charge is smaller and the expected deflections are smaller—both in terms of how closely the α particles pass by the nucleus undeflected and the angle of deflection. If the nucleus is larger, the positive charge is larger and the expected deflections are larger—more α particles will be deflected, and the deflection angles will be larger. (d) The paths followed by the α particles match the predictions from (a), (b), and (c).

17. A sample of magnesium is found to contain 78.70% of 24Mg atoms (mass 23.98 amu), 10.13% of 25Mg atoms (mass 24.99 amu), and 11.17% of 26Mg atoms (mass 25.98 amu). Calculate the average mass of a Mg atom.

18. Naturally occurring copper consists of 63Cu (mass 62.9296 amu) and 65Cu (mass 64.9278 amu), with an average mass of 63.546 amu. What is the percent composition of Cu in terms of these two isotopes?

19. In what way are isotopes of a given element always different? In what way(s) are they always the same?

20. Write the symbol for each of the following ions:

(a) the ion with a 1+ charge, atomic number 55, and mass number 133

(b) the ion with 54 electrons, 53 protons, and 74 neutrons

(c) the ion with atomic number 15, mass number 31, and a 3− charge

(d) the ion with 24 electrons, 30 neutrons, and a 3+ charge

(a) 133Cs+; (b) 127I−; (c) 31P3−; (d) 57Co3+

21. Write the symbol for each of the following ions:

(a) the ion with a 3+ charge, 28 electrons, and a mass number of 71

(b) the ion with 36 electrons, 35 protons, and 45 neutrons

(c) the ion with 86 electrons, 142 neutrons, and a 4+ charge

(d) the ion with a 2+ charge, atomic number 38, and mass number 87

22. Determine the number of protons, neutrons, and electrons in the following isotopes that are used in medical diagnoses:

(a) atomic number 9, mass number 18, charge of 1−

(b) atomic number 43, mass number 99, charge of 7+

(c) atomic number 53, atomic mass number 131, charge of 1−

(d) atomic number 81, atomic mass number 201, charge of 1+ 

(e) Name the elements in parts (a), (b), (c), and (d).

23. The following are properties of isotopes of two elements that are essential in our diet. Determine the number of protons, neutrons and electrons in each and name them.

(a) atomic number 26, mass number 58, charge of 2+ 

(b) atomic number 53, mass number 127, charge of 1−

24. An element has the following natural abundances and isotopic masses: 90.92% abundance with 19.99 amu, 0.26% abundance with 20.99 amu, and 8.82% abundance with 21.99 amu. Calculate the average atomic mass of this element.

25. Average atomic masses listed by IUPAC are based on a study of experimental results. Bromine has two isotopes 79Br and 81Br, whose masses (78.9183 and 80.9163 amu) and abundances (50.69% and 49.31%) were determined in earlier experiments. Calculate the average atomic mass of bromine based on these experiments.

26. Variations in average atomic mass may be observed for elements obtained from different sources. Lithium provides an example of this. The isotopic composition of lithium from naturally occurring minerals is 7.5% 6Li and 92.5% 7Li, which have masses of 6.01512 amu and 7.01600 amu, respectively. A commercial source of lithium, recycled from a military source, was 3.75% 6Li (and the rest 7Li). Calculate the average atomic mass values for each of these two sources.

27. The average atomic masses of some elements may vary, depending upon the sources of their ores. Naturally occurring boron consists of two isotopes with accurately known masses (10B, 10.0129 amu and 11B, 11.0931 amu). The actual atomic mass of boron can vary from 10.807 to 10.819, depending on whether the mineral source is from Turkey or the United States. Calculate the percent abundances leading to the two values of the average atomic masses of boron from these two countries.  Turkey source: 26.49% (of 10.0129 amu isotope); US source: 25.37% (of 10.0129 amu isotope)

28. The 18O:16O abundance ratio in some meteorites is greater than that used to calculate the average atomic mass of oxygen on earth. Is the average mass of an oxygen atom in these meteorites greater than, less than, or equal to that of a terrestrial oxygen atom?

29. A molecule of metaldehyde (a pesticide used for snails and slugs) contains 8 carbon atoms, 16 hydrogen atoms, and 4 oxygen atoms. What are the molecular and empirical formulas of metaldehyde?

30. Give the group name for each of the following elements:

    1. krypton

    2. selenium

    3. barium

    4. lithium

31. Write the molecular formula for each compound.

    1. Nitrous oxide, also called “laughing gas,” has 2 nitrogen atoms and 1 oxygen atom per molecule. Nitrous oxide is used as a mild anesthetic for minor surgery and as the propellant in cans of whipped cream.

    2. Sucrose, also known as cane sugar, has 12 carbon atoms, 11 oxygen atoms, and 22 hydrogen atoms.

    3. Sulfur hexafluoride, a gas used to pressurize “unpressurized” tennis balls and as a coolant in nuclear reactors, has 6 fluorine atoms and 1 sulfur atom per molecule.

32. Predict the charge on the most common monatomic ion formed by each element.

    1. calcium, used to prevent osteoporosis

    2. iodine, required for the synthesis of thyroid hormones

    3. zirconium, widely used in nuclear reactors

33. Ionic and covalent compounds are held together by electrostatic attractions between oppositely charged particles. Describe the differences in the nature of the attractions in ionic and covalent compounds. Which class of compounds contains pairs of electrons shared between bonded atoms?

34. Which contains fewer electrons than the neutral atom—the corresponding cation or the anion?

35. What is the difference between an organic compound and an inorganic compound?

36. What is the advantage of writing a structural formula as a condensed formula?

37. The majority of elements that exist as diatomic molecules are found in one group of the periodic table. Identify the group.

38. Discuss the differences between covalent and ionic compounds with regard to

39. the forces that hold the atoms together.

40. melting points.

41. physical states at room temperature and pressure.

42. Why do covalent compounds generally tend to have lower melting points than ionic compounds?

43. What is the total number of electrons present in each ion?

    1. F−

    2. Rb+

    3. Ce3+

    4. Zr4+

    5. Zn2+

    6. Kr2+

    7. B3+

44. What is the total number of electrons present in each ion?

    1. Ca2+

    2. Se2−

    3. In3+

    4. Sr2+

    5. As3+

    6. N3−

    7. Tl+

45. Predict how many electrons are in each ion.

    1. an oxygen ion with a −2 charge

    2. a beryllium ion with a +2 charge

    3. a silver ion with a +1 charge

    4. a selenium ion with a +4 charge

    5. an iron ion with a +2 charge

    6. a chlorine ion with a −1 charge

    7. Predict how many electrons are in each ion.

    8. a copper ion with a +2 charge

    9. a molybdenum ion with a +4 charge

    10. an iodine ion with a −1 charge

    11. a gallium ion with a +3 charge

    12. an ytterbium ion with a +3 charge

    13. a scandium ion with a +3 charge

    14. Predict the charge on the most common monatomic ion formed by each element.

    15. chlorine

    16. phosphorus

    17. scandium

    18. magnesium

    19. arsenic

    20. oxygen

46. Predict the charge on the most common monatomic ion formed by each element.

    1. sodium

    2. selenium

    3. barium

    4. rubidium

    5. nitrogen

    6. aluminum

47. Write the name of each binary covalent compound.

    1. IF7

    2. N2O5

    3. OF2

48. Write the name of each binary covalent compound.

1. IF7

2. N2O5

3. OF2

49. Name each cation: K+, Al3+, NH4+, Mg2+, Li+

50. Name each anion: Br−, CO32−, S2−, NO3−, HCO2−, F−, ClO−, C2O42−

51. Name each anion: PO43−, Cl−, SO32−, CH3CO2−, HSO4−, ClO4−,NO2−, O2−

52. Name each anion: SO42−, CN−, Cr2O72−, N3−, OH−, I−, O22−

53. Name each compound: MgBr2, NH4CN, CaO, KClO3, K3PO4, NH4NO2, NaN3

54. Name each compound: NaNO3, Cu3(PO4)2, NaOH

55. For each ionic compound, name the cation and the anion and give the charge on each ion: BeO, Pb(OH)2, BaS, Na2Cr2O7

56. Write the formula for each compound: magnesium carbonate, aluminum sulfate, potassium phosphate, lead(IV) oxide, silicon nitride, sodium hypochlorite, titanium(IV) chloride, disodium ammonium phosphate

57. Using the periodic table, classify each of the following elements as a metal or a nonmetal, and then further classify each as a main-group element, transition metal, or inner transition metal:

a. uranium
b. bromine
c. strontium
d. neon
e. gold
f. americium
g. rhodium
h. sulfur
i. carbon
j. potassium

58. Using the periodic table, identify the lightest member of each of the following groups:

a. noble gases
b. alkaline earth metals
c. alkali metals
d. chalcogens

59. Use the periodic table to give the name and symbol for each of the following elements:

(a) the noble gas in the same period as germanium
(b) the alkaline earth metal in the same period as selenium
(c) the halogen in the same period as lithium
(d) the chalcogen in the same period as cadmium