Inorganic Chemistry Lesson 16 Classification of elements. Part 2. Periodic Law

Transcription

1 Inorganic Chemistry Lesson 16 Classification of elements. Part 2. Periodic Law March 2, Atomic mass and the properties of elements. If we arrange halogens according to their nonmetallicity (which we provisionally define as the element s ability to act as an oxidizer) we will see that their nonmetallicity decreases it this order: F >Cl >Br >I In other words, fluorine is the most active nonmetal, whereas iodine is the least active one. Interestingly, the masses of halogens increase in the same order: Halogen Atomic mass F 19 Cl 35.5 Br 79 I 128 As we learned from the Problem number 1 (HW15), the activity of alkaline metals increases in that order: Li <Na <K <Rb <Cs That order corresponds to the increase of the element s mass Alkaline metal Atomic mass Li 7 Na 23 K 39 Rb 85 Cs 133 In both cases, we observe some unambiguous dependence between the properties of elements and their masses. One can conjecture such a dependence is universal, and it can be observed in other groups of elements: chalcogens, alkaline earth metals, and others. Later, we will check if that is the case, and, for a while, let s agree that the lightest element of the group can be used as a representative of the group as whole. In connection to that, let s take the lightest element 1

2 from each group, and arrange all of them according to their chemical properties. Since the ability to oxidize other element, or the ability to be oxidized are hard to characterize quantitatively, let s choose some other, more quantitative criteria. Such criteria are the element s valence in their oxides and hydrides. Before we started, let s look at one more group of elements, the boron group. 2 Boron group. This group is interesting because it is composed by both metals and nonmetals. The lightest element of this group, boron (chemical symbol B), is a nonmetal (sometimes it is called metalloid because of its metallic luster and hardness). Boron, as well as other representatives of this group (aluminium (Al), gallium (Ga), indium (In), and Thallium (Tl)) form oxides in reaction with oxygen, and their valence in the oxides in three: B 2 O 3, Al 2 O 3, Ga 2 O 3, etc. Interestingly, whereas boron is nonmetal, other elements of the boron group are metals, although majority of them are amphoteric. The boron group elements, as well as other elements are capable of forming hydrides: BH 3, AlH 3, etc, where they are also trivalent. 3 Periodic law. Now we know seven groups of elements (halogens, chalcogens, pniktogens, carbon group, boron group, alkaline metals and alkaline earth metals), and we are ready to start to arrange all of them together. Let s take the lightest element from each group and arrange them according to their mass. In parallel, we will write the formula of the element s hydride and oxide. Element Li Be B C N O F Hydride LiH BeH 2 BH 3 CH 4 NH 3 OH 2 FH Oxide Li 2 O BeO B 2 O 3 CO 2 N 2 O Mass We write the formula of water as OH 2 (not H 2 O), and the formula of hydrogen fluoride as FH, not HF. We do that for consistency only, that is not a correct style (although not a mistake either). In future, we will follow a standard style (H 2 O, HF). From this table, it is clear that, when the mass of the listed elements increases the valence of the element in its hydride increases from one (in LiH) to four (in CH 4 ), and then decreases back to one (in HF). In contrast, the valence of the elements in their oxides increases from one (in Li 2 O) to five (in N 2 O 5 ), and never decreases back. 1 Interestingly, if we take the second element of each group, the valence of oxides will monotonously increase from left to right: 1 Obviously, oxygen cannot form oxides. Fluorine is more active than oxygen, so it is incorrect to speak about fluorine oxides. 2

3 Element Na Mg Al Si P S Cl Hydride NaH MgH 2 AlH 3 SiH 4 PH 3 H 2 S HCl Oxide Na 2 O MgO AL 2 O 3 SiO 2 P 2 O 5 SO 3 (Cl 2 O 7 ) Mass Chlorine (VII) oxide (Cl 2 O 7 ) is shown in parentheses because this compound is not stable a pure form. However, its corresponding acid, perchloric (HClO 4 ) is a stable and well characterized compound. By combining the above rows into a single table, and by adding heavier elements from each group, we get the following table (we limit ourselves with four elements from each group, and we draw a generic formulas of hydrides and oxides): Table 1. Group I II III IV V VI VII Hydride XH XH 2 XH 3 XH 4 XH 3 XH 2 XH Oxide X 2 O XO X 2 O 3 XO 2 X 2 O 5 XO 3 X 2 O 7 I Li 7 Be 9 B 10 C 12 N 14 O 16 F 19 II Na 23 Mg 24 Al 27 Si 28 P 31 S 32 Cl 35.5 III K 39 Ca 40 Ga 70 Ge 72.5 As 75 Se 89 Br 80 IV Rb 85 Sr 88 In 115 Sn 119 Sb 122 Te 128 I 127 As we can see, all seven groups fit nicely into this table, and one important thing can be derived from it: atomic masses of elements increase when we move from the top to the bottom, and from the left side to the right. That means the periodic dependence exists between the masses of elements and their chemical properties. Indeed, if we unwrap the above table, we get: Li Be B C N O F Na Mg Al Si P S Cl... In this sequence, when we are going from left to right, the mass of the elements increase monotonously, but the properties change periodically: element s metallicity gradually decreases from lithium to fluorine, then we jump to sodium (an alkaline metal), and then the metallicity decreases again until we arrive to chlorine (a nonmetal). This periodicity was first formulated as a law by a Russian chemist Dmitry Mendeleev. Figure 1: Dmitry Mendeleev ( ), a discoverer of the Periodic Law. Properties of elements are periodic functions of their atomic masses. Accordingly, each raw in the above table is called a period, and the table itself is called Periodic Table. Each table s column is called a group, 3

4 and the number of each group is equal to the highest possible valence of the oxides of the elements from that group. 4 How many groups? As you can see, the Table 1 is incomplete: there is a gap between each halogen and the next alkaline metal. Indeed, whereas the mass of adjacent elements differ by 2 or 3 Daltons, the difference between the mass of fluorine (19) and sodium (23), or between bromine (80) and rubidium (85) is unusually high. By the moment Mendeleev made his discovery (in 1869), scientist knew no good candidates to fill this gap. Only in 1894, British physicist Lord Raleigh and chemist Sir William Ramsay separated a strange gas form air. This gas, called argon, was present in atmosphere in small amount (ca 1%). This gas was totally inert, it did not react with any known acid, base, or oxidizer. Later Ramsay discovered other gases with same properties: Neon (Ne), Krypton (Kr), Xenon (Xe), Neon (Ne). Due to their inertness they, as well as independently discovered helium (He) and radon (Ra) are called inert, or noble gases. Initially they were believed to have maximal possible valence of zero. Accordingly, they occupied the column number zero in the early periodic table (before alkaline metals). However, in second half of XX century chemists managed to oxidize xenon and krypton to tetroxides (XeO 4 and KrO 4, accordingly). Based on that observation, the noble gases now occupy the eighth group of the Periodic table. Currently we know that the maximal possible valence of elements in their oxides is eight, and, accordingly, there are eight columns (groups) in the Periodic table. 2 5 Transition metals Even if we add the noble gas group to the table 1, it still is incomplete. Indeed, there is no space there for such elements as iron, cobalt, zinc, platinum, etc. In addition, there are huge gaps between calcium (atomic weight 40 Da) and gallium (atomic weight 70), between strontium (atomic weight 88) and indium (atomic weight 115 Da). Obviously, the remaining elements are supposed to occupy those gaps, but how concretely should they be arranged? The idea to arrange them was arguably the most brilliant Mendeleev s idea. His considerations were as follows. If we look at the properties of the metals with atomic weight from 40 to 70 Da (scandium, titanium, vanadium, chromium, manganese, iron, cobalt, nickel, copper, zinc) we will see a strange duality. From one hand, all those metals are capable of forming salts with common acids, and they are divalent in those salts. For example, iron or chromium react with hydrochloric acid to form ferrous or chromous chlorides, respectively. Fe + 2HCl FeCl 2 + H 2 (1) 2 Actually, a new style Periodic table, which was obtained by a rearrangement of the old style table, has more groups. However, to understand the Periodic law better, it would be useful to start with the old style Periodic table. 4

5 Cr + 2HCl CrCl 2 + H 2 (2) Accordingly, chromium (II) hydroxide can be prepared, which can react with acids to produce other chromium (II) salts. CrCl 2 + 2NaOH Cr(OH) 2 + 2NaCl (3) Cr(OH) 2 + H 2 SO 4 CrSO 4 + H 2 O (4) In other words, the properties of chromium (II) hydroxide are similar to the properties of the hydroxides of other common divalent metals. However, in addition to a divalent state, other valence state are possible for chromium. The maximal valence of chromium is six, and the typical example of Cr (VI) compounds is chromium trioxide (CrO 3 ). This oxide reacts with water to produce a compounds with the formula H 2 CrO 4. CrO 3 + H 2 O H 2 CrO 4 (5) This equation resembles the reaction we studied during previous lessons, namely the reaction between sulfur trioxide and water: SO 3 + H 2 O H 2 SO 4 (6) This similarity is not just a coincidence. The compound with formula H 2 CrO 4 has strong acidic properties, and, accordingly, is called chromic acid. For example, it reacts with sodium hydroxide to form sodium chromate: H 2 CrO 4 + 2NaOH Na 2 CrO 4 + H 2 O (7) (Compare the reaction 7 with the reaction between sulfuric acid and sodium hydroxide:) H 2 SO 4 + 2NaOH Na 2 SO 4 + H 2 O (8) The similarity between these two acids, sulfuric and chromic, is even deeper. Not only the acid properties of these two acids are similar, but the properties of the salts they form are pretty close. For example, whereas both sodium sulfate and chromate are soluble in water, barium chromate is as insoluble as barium sulfate. BaCl 2 + Na 2 SO 4 BaSO 4 (solid) + 2NaCl (9) BaCl 2 + Na 2 CrO 4 BaCrO 4 (solid) + 2NaCl (10) Due to its dual nature, a special subgroups where allocated for this type elements. 5

6 Homework Read the CW materials. We will continue our discussion of the Periodic Law during the next class. Answer the following questions. 1. What does the Periodic Law says? 2. How do metallic properties of element change when we move from the top of the Table 1 to the bottom? 3. Which element manifests stronger nonmetallic properties, arsenic of antimony? 4. Which sequence of elements is called period? 5. List the elements from the third period that form basic oxides. 6. Using the Table 1, can you predict which acid is stronger, sulfuric acid (H 2 SO 4 ) or perchloric acid (HClO 4 )? Which hydroxide is more basis, magnesium hydroxide, or strontium hydroxide? My is mark.lukin@gmail.com 6