Atoms, Molecules, and Ions

1. Chapter Two Atoms, Molecules, and Ions

2. Laws of Chemical Combination • Law of Conservation of Mass • The total mass remains constant during a chemical reaction. • Law of Definite Proportions • All samples of a compound have the same composition, or … • All samples have the same proportions, by mass, of the elements present.

3. Example 2.1 A Conceptual Example Jan Baptista van Helmont (1579–1644) first measured the mass of a young willow tree and, separately, the mass of a bucket of soil and then planted the tree in the bucket. After five years, he found that the tree had gained 75 kg in mass even though the soil had lost only 0.057 kg. He had added only water to the bucket, and so he concluded that all the mass gained by the tree had come from the water. Explain and criticize his conclusion.

4. Example 2.1 A Conceptual Example Jan Baptista van Helmont (1579–1644) first measured the mass of a young willow tree and, separately, the mass of a bucket of soil and then planted the tree in the bucket. After five years, he found that the tree had gained 75 kg in mass even though the soil had lost only 0.057 kg. He had added only water to the bucket, and so he concluded that all the mass gained by the tree had come from the water. Explain and criticize his conclusion. Analysis and Conclusions Van Helmont’s explanation anticipated the law of conservation of mass by dismissing the possibility that the increased mass of the tree could have been created from nothing. His main failure, however, was in not identifying all the substances involved. He did not know about the role of carbon dioxide gas in the growth of plants. In applying the law of conservation of mass, we must focus on all the substances involved in a chemical reaction. Exercise 2.1A A sealed photographic flashbulb containing magnesium and oxygen has a mass of 45.07 g. On firing, a brilliant flash of white light is emitted and a white powder is formed inside the bulb. What should the mass of the bulb be after firing? Explain. Exercise 2.1B Is the mass of the burned match pictured in the photograph the same as, less than, or more than the mass of the burning match? Explain your answer.

5. … have the same composition. Three different sources of a compound … The Law of Definite Proportions

6. Example 2.2 The mass ratio of oxygen to magnesium in the compound magnesium oxide is 0.6583:1. What mass of magnesium oxide will form when 2.000 g of magnesium is completely converted to magnesium oxide by burning in pure oxygen gas?

7. Example 2.2 The mass ratio of oxygen to magnesium in the compound magnesium oxide is 0.6583:1. What mass of magnesium oxide will form when 2.000 g of magnesium is completely converted to magnesium oxide by burning in pure oxygen gas? Strategy Solution We usually write mass ratios in a form such as “the ratio of O to Mg is 0.6583:1.” The first number represents the mass of the first element named—in this case, a mass of oxygen, say 0.6583 g oxygen—and the second number represents the mass of the second element named—here a mass of magnesium. Although written only as “1,” we assume that the second number is as precisely known as the first, that is, 1.0000 g magnesium. According to the law of constant composition, the mass of oxygen that combines with the 2.000 g of magnesium must be just the right amount that makes the mass ratio of oxygen to magnesium in the product 0.6583:1. According to the law of conservation of mass, the mass of magnesium oxide must equal the sum of the masses of the magnesium and oxygen that react. We can state the mass ratio in the form of a conversion factor and then determine the required mass of oxygen. The mass of the sole product, magnesium oxide, equals the total of the masses of the substances entering into the reaction.

8. Example 2.2 continued Assessment As a simple check, compare the calculated mass with the starting mass of magnesium. From the law of conservation of mass, the product of the reaction of magnesium with oxygen must have a mass greater than the starting mass of magnesium; and 3.317 g magnesium oxide is indeed greater than 2.000 g magnesium. Exercise 2.2A What mass of magnesium oxide is formed when 1.500 g of oxygen combines with magnesium? Exercise 2.2B When a strip of magnesium metal was burned in pure oxygen gas, 1.554 g of oxygen was consumed and the only product formed was magnesium oxide. What must have been the masses of magnesium metal burned and magnesium oxide formed?

9. Law of Multiple Proportions When two or more different compounds of the same two elements are compared, the masses of one element that combine with a fixed mass of the second element are in the ratio of small whole numbers.

10. Law of Multiple Proportions (cont’d) Ratio of oxygen-to-carbon in CO2 is exactly twice the ratio in CO.

11. Law of Multiple Proportions (cont’d) • Four different oxides of nitrogen can be formed by combining 28 g of nitrogen with: • 16 g oxygen, forming Compound I • 48 g oxygen, forming Compound II • 64 g oxygen, forming Compound III • 80 g oxygen, forming Compound IV What is the ratio 16:48:64:80 expressed as small whole numbers? • Compounds I–IV are N2O, N2O3, N2O4, N2O5

12. Dalton’s Atomic Theory Proposed in 1803 to explain the law of conservation of mass, law of definite proportions, and law of multiple proportions. • Matter is composed of atoms:tiny, indivisible particles. • All atoms of a given element are the same. • Atoms of one element differ from atoms of other elements. • Compounds are formed when atoms of different elements unite in fixed proportions. • A chemical reaction involves rearrangement of atoms. No atoms are created, destroyed, or broken apart.

13. Dalton’s Atomic Theory: Conservation of Mass and Definite Proportions … six fluorine atoms and four hydrogen atoms after reaction. Mass is conserved. Six fluorine atoms and four hydrogen atoms before reaction … HF always has one H atom and one F atom; always has the same proportions (1:19) by mass.

14. Subatomic Particles • Protons and neutrons are located at the center of an atom (at the nucleus). • Electrons are dispersed around the nucleus.

15. Isotopes • Atoms that have the same number of protons but different numbers of neutrons are called isotopes. • The atomic number (Z) is the number of protons in the nucleus of a given atom of a given element. • The mass number (A) is an integral number that is the sum of the numbers of protons and neutrons in an atom. • The number of neutrons = A – Z.

16. Isotopes (cont’d) • How many protons are in chlorine-35? • How many protons are in chlorine-37? • How many neutrons are in chlorine-37? Atoms can be represented using the element’s symbol and the mass number (A) and atomic number (Z): 37 Cl 17 35 Cl 17 A E Z

17. Example 2.3 How many protons, neutrons, and electrons are present in a 81Br atom?

18. Example 2.3 How many protons, neutrons, and electrons are present in a 81Br atom? Strategy Solution We can use the identity of the element and several simple relationships between the subatomic particles described above to determine the numbers of these particles. The atomic number of bromine Z = 35 = number of protons is not given here, but we can obtain it from the list of elements on the inside front cover. In the bromine atom, the number Number of protons = number of electrons = 35 of positively charged protons equals the number of negatively charged electrons. From mass number 81 and Number of neutrons = A – Z = 81 – 35 = 46 atomic number 35, we calculate the number of neutrons, using Equation (2.1).

19. Atomic Mass • Atoms are very tiny, so a tiny unit is needed to express the mass of an atom or molecule. • One atomic mass unit (u) = 1/12 the mass of a C-12 atom. • 1 u = 1.66054 × 10–24 g • The mass of an atom is not exactly the sum of the masses of the protons + neutrons + electrons (we will see why in Chapter 19).

20. Atomic Mass (cont’d) • Question: do all isotopes of an element have the same mass? Why or why not? • The atomic mass given on the periodic table is the weighted average of the masses of the naturally occurring isotopes of that element.

21. Example 2.4 Use the data cited above to determine the weighted average atomic mass of carbon. Example 2.5 An Estimation Example Indium has two naturally occurring isotopes and a weighted average atomic mass of 114.82 u. One of the isotopes has a mass of 112.9043 u. Which is likely to be the second isotope: 111In, 112In, 114In, or 115In?

22. Example 2.4 Use the data cited above to determine the weighted average atomic mass of carbon. Strategy Assessment The contribution each isotope makes to the weighted average atomic mass is given by Equation (2.2). The weighted average atomic mass is the sum of the two contributions. This is the value listed in a table of atomic masses. As expected, the atomic mass of carbon is much closer to 12 u than to 13 u. Solution The contributions are Contribution of carbon-12 = 0.98892 x 12.00000 u = 11.867 u Contribution of carbon-13 = 0.01108 x 13.00335 u = 0.1441 u The weighted average mass is Atomic mass of carbon = 11.867 u + 0.1441 u = 12.011 u Exercise 2.4A There are three naturally occurring isotopes of neon. Their percent abundances and atomic masses are neon-20, 90.51%, 19.99244 u; neon-21, 0.27%, 20.99395 u; neon-22, 9.22%, 21.99138 u. Calculate the weighted average atomic mass of neon. Exercise 2.4B The two naturally occurring isotopes of copper are copper-63, mass 62.9298 u, and copper-65, mass 64.9278 u. What must be the percent natural abundances of the two isotopes if the atomic mass of copper listed in a table of atomic masses is 63.546 u?

23. Example 2.5 An Estimation Example Indium has two naturally occurring isotopes and a weighted average atomic mass of 114.82 u. One of the isotopes has a mass of 112.9043 u. Which is likely to be the second isotope: 111In, 112In, 114In, or 115In? Analysis and Conclusions The masses of isotopes differ only slightly from whole numbers, which tells us that the isotope with mass 112.9043 u is 113In. To account for the observed weighted average atomic mass of 114.82 u, the second isotope must have a mass number greater than 114. It can be only 115In. Exercise 2.5A The masses of the three naturally occurring isotopes of magnesium are 24Mg, 23.98504 u; 25Mg, 24.98584 u; 26Mg, 25.98259 u. Use the atomic mass given inside the front cover to determine which of the three is the most abundant. Can you determine which is the second most abundant? Explain. Exercise 2.5B For the three magnesium isotopes described in Exercise 2.5A, (a) could the percent natural abundance of 24Mg be 60.00%? (b) What is the smallest possible value for the percent natural abundance of 24Mg?

24. Mendeleev’s Periodic Table • Mendeleev arranged the known elements in order of increasing atomic weight from left to right and from top to bottom in groups. • Elements that closely resembled one another were arranged in the same vertical group. • Gaps were left where undiscovered elements should appear. • From the locations of the gaps, he was able to predict properties of some of the undiscovered elements.

25. Germanium:Prediction vs. Observation

26. The Modern Periodic Table Except for H, elements left of the zigzag line are metals. To the right of the line we find nonmetals, including the noble gases. Some elements adjacent to the line are called metalloids.

27. Molecules and Formulas • A molecule is a group of two or more atoms held together by covalent bonds. • A molecular formula gives the number of each kind of atom in a molecule. • An empirical formula simply gives the (whole number) ratio of atoms of elements in a compound.

28. Structural Formulas and Models • Structural formulas and models show how atoms are attached to one another. The condensed structural formula for acetic acid is CH3COOH. C2H4O2: two C atoms, four H atoms, two O atoms. CH3COOH shows how the atoms are arranged. Ball-and-stick model Space-filling model

29. Nomenclature … is the method for naming compounds and writing formulas for compounds. • We could have a specific name for each compound—but we would have to memorize each one! • Can you imagine having to memorize the names of half a million different inorganic compounds? Twenty million organic compounds?? • Instead we have a systematic method— conventions and rules—for naming compounds and writing formulas.

30. Nomenclature ofBinary Molecular Compounds • Binary compounds contain ___ elements. • Molecular compounds exist as ________.

31. Naming Binary Molecular Compounds • The name consists of two words. • First word: name of the element that appears first in the formula. • Second word: stem of the name of the second element, ending with -ide. • Names are further modified by adding prefixes to denote the numbers of atoms of each element in the molecule.

32. Which element is named first? Begin with boron and follow the line to determine the order of naming. Rule of thumb: the element that is farthest down and to the left on the periodic table is usually written first. In a compound consisting of arsenic (As) and sulfur (S), which element is named first?

34. Example 2.6 Write the molecular formula and name of a compound for which each molecule contains six oxygen atoms and four phosphorus atoms. Example 2.7 Write (a) the molecular formula of phosphorus pentachloride and (b) the name of S2F10.

35. Example 2.6 Write the molecular formula and name of a compound for which each molecule contains six oxygen atoms and four phosphorus atoms. Strategy We need to write the chemical symbols of the two elements and use the stated numbers of atoms as subscripts following the symbols. We then determine which element to place first in the molecular formula. Solution We represent the six atoms of oxygen as O6 and the four atoms of phosphorus as P4. According to the scheme in Figure 2.9, the element O is followed only by F. Phosphorus therefore comes first in the formula; we write P4O6. The name of a compound with four (tetra-) P atoms and six (hexa-) oxygen atoms in its molecules is tetraphosphorus hexoxide. Exercise 2.6A Write the molecular formula and name of a compound for which each molecule contains four fluorine atoms and two nitrogen atoms. Exercise 2.6B Write the molecular formula and name of a compound the molecules of which each contain one oxygen atom and eight sulfur atoms.

36. Example 2.7 Write (a) the molecular formula of phosphorus pentachloride and (b) the name of S2F10. Solution (a) Choosing which element symbol goes first. The order of elements shown in the molecular formula must be the same as the order in the name. Writing subscripts. The lack of a prefix on phosphorus signifies one P atom per molecule. The prefix penta- indicates five chlorine atoms. The molecular formula is PCl5. (b) The subscripts indicate two (di-) sulfur atoms and ten (deca-) fluorine atoms. The compound is disulfur decafluoride. Exercise 2.7A Write (a) the molecular formula of tetraphosphorus decoxide and (b) the name of S7O2. Exercise 2.7B Write a plausible molecular formula for a compound that has one sulfur atom, two oxygen atoms, and two fluorine atoms in each of its molecules. Comment on any ambiguity that exists in this case.

37. Ions and Ionic Compounds • An atom that either gains or loses electron(s) is an ion. • There is no change in the number of protons or neutrons in the nucleus of the atom. • Cation – has a positive charge from loss of electron(s). • Anion – has a negative charge from gain of electron(s).

38. Ions and Ionic Compounds (cont’d) In anionic compound, oppositely charged ions are attracted to each other such that the compound has no net charge. There are no distinct molecules of sodium chloride. Sodium chloride simply consists of sodium ions and chloride ions, regularly arranged.

39. Example 2.8 Determine the formula for (a) calcium chloride and (b) magnesium oxide. Example 2.9 What are the names of (a) MgS and (b) CrCl3?

40. Example 2.8 Determine the formula for (a) calcium chloride and (b) magnesium oxide. Solution (a) First we write the symbols for the ions, with the cation first: Ca2+ and Cl–. The simplest combination of these ions that gives an electrically neutral formula unit is one Ca2+ ion for every two Cl– ions. The formula is CaCl2 . Ca2+ + 2 Cl– = CaCl2 (b) Figure 2.10 tells us that the ions are Mg2+ and O2–. The simplest ratio for an electrically neutral formula unit is 1:1. The formula of this binary ionic compound is MgO. Mg2+ + O2– = MgO Exercise 2.8A Give the formula for each of the following ionic compounds: (a) potassium sulfide (b) lithium oxide (c) aluminum fluoride Exercise 2.8B Give the formula for each of the following ionic compounds: (a) chromium(III) oxide (b) iron(II) sulfide (c) lithium nitride

41. Example 2.9 What are the names of (a) MgS and (b) CrCl3? Solution (a) MgS is made up of Mg2+ and S2– ions. Its name is magnesium sulfide. (b) From Figure 2.10 we see that there are two simple ions of chromium, Cr3+ and Cr2+. Because there are three Cl– ions in the formula unit CrCl3, the cation must have a charge of 3+, that is, it must be Cr3+. Because there are two chromium cations, there are two chlorides, CrCl2 and CrCl3, and we must assign a different name to each. Therefore, the name of our compound cannot be simply chromium chloride; instead, to indicate the 3+ cation, we say the name is chromium(III) chloride. Exercise 2.9A Name the following compounds: (a) CaBr2 (b) Li2S (c) FeBr2 (d) CuI Exercise 2.9B Write the name and formula for each of the following compounds: (a) the sulfide of copper(I) (b) the oxide of cobalt(III) (c) the nitride of magnesium

42. Monatomic Ions • Group IA metals form ions of 1+ charge. • Group IIA metals form ions of 2+ charge. • Aluminum, a group IIIA metal, forms ions with a 3+ charge. • Nonmetal ions of groups V, VI, and VII usually have charges of 3–, 2–, and 1–, respectively. • Group B metal ions (transition metal ions) often have more than one possible charge. A Roman numeral is used to indicate the actual charge. • A few transition elements have only one common ion (Ag, Zn, Cd), and a Roman numeral is not often used.

43. Symbols and Periodic Table Locations of Some Monatomic Ions Copper forms either copper(I) or copper(II) ions. Titanium forms both titanium(II) and titanium(IV) ions. What is the charge on a zirconium(IV) ion?

44. Formulas and Names ofBinary Ionic Compounds • Binary ionic compounds are made up of monatomic cations and anions. • These combinations must be electrically neutral. • The formula unit is the simplest collection of cations and anions that represents an electrically neutral unit. • Formula unit is to ion as ________ is to atom. • To write a formula, combine the proper number of each ion to form a neutral unit. • To name a binary ionic compound, name the cation, then the anion. • Monatomic anion names end in -ide.

45. Polyatomic Ions • A polyatomic ion is a charged group of covalently bonded atoms. • There are many more polyatomic anions than there are polyatomic cations. • You should (eventually!) commit to memory much of Table 2.4 • hypo- and per- are sometimes seen as prefixes in oxygen-containing polyatomic ions (oxoanions). • -ite and -ate are commonly found as suffixes in oxygen-containing polyatomic ions.

47. Example 2.10 Write the formula for (a) sodium sulfite and (b) ammonium sulfate. Example 2.11 What is the name of (a) NaCN and (b) Mg(ClO4)2?

48. Example 2.10 Write the formula for (a) sodium sulfite and (b) ammonium sulfate. Solution (a) It is possible to identify the sulfite ion without memorizing all the ions in Table 2.4. If you remember the name and formula of one of the sulfur–oxygen polyatomic anions, you should be able to deduce the names of others. Suppose you remember that sulfate is SO42–. The -ite anion has one fewer oxygen atom, 3 instead of 4, and so it is SO32–. The charges of the two ions in a formula unit must balance, which means the formula unit of sodium sulfite must have Na+ and SO32– in the ratio 2:1. The formula is therefore Na2SO3. (b) The ammonium ion is NH4+, and the sulfate ion is SO42–. A formula unit of ammonium sulfate has two NH4+ ions and one SO42– ion. To represent the two NH4+ ions, we place parentheses around the NH4, followed by a subscript 2, (NH4)2, and thus arrive at the formula (NH4)2SO4. Exercise 2.10A What is the formula for (a) ammonium carbonate, (b) calcium hypochlorite, and (c) chromium(III) sulfate? Exercise 2.10B Write a plausible formula for (a) potassium aluminum sulfate (b) magnesium ammonium phosphate

49. Example 2.11 What is the name of (a) NaCN and (b) Mg(ClO4)2? Solution (a) The ions in this compound are Na+, sodium ion, and CN–, cyanide ion (see Table 2.4). The name of the compound is sodium cyanide. (b) The ions present are Mg2+, magnesium ion, and ClO4–, perchlorate ion. The name of the compound is magnesium perchlorate. Exercise 2.11A Name each of the following compounds: (a) KHCO3 (b) FePO4 (c) Mg(H2PO4)2 Exercise 2.11B Give a plausible name for the following: (a) Na2SeO4 (b) FeAs (c) Na2HPO3

50. Hydrates • A hydrate is an ionic compound in which the formula unit includes a fixed number of water molecules associated with cations and anions. • To name a hydrate, the compound name is followed by “___hydrate” where the blank is a prefix to indicate the number of water molecules. • The number of water molecules associated with each formula unit is written as an appendage to the formula unit name separated by a dot. • Examples: BaCl2.2 H2O; CuSO4. 5 H2O