1. Periodic Table
GLY 4200 – Lecture 6 – Fall, 2017

2. Dmitri Ivanovich Mendeleev
Russian Chemist
"The properties of the elements are a periodic function of their atomic masses"
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Dmitri Ivanovich Mendeleev, born in Siberia in 1834, originally constructed the periodic table through a careful study of the properties of the sixty-three elements known in 1869, when he first published his results.
He stated, “The properties of the elements are a periodic function of their atomic masses."

3. Henry G.J. Moseley British chemist 1887-1915
“Similar properties recur periodically when elements are arranged by increasing atomic number”
British chemist Henry G.J. Moseley used X-ray study to more accurately position elements in the periodic table in a paper published in Moseley’s work was cut short when he was killed in action at Gallipoli in 1915, at the age of 27.
He proposed a modification of Mendeleev’s rule, “Similar properties recur periodically when elements are arranged by increasing atomic number.”
This is really the correct way to formulate the periodic relationship among elements.

4. Periodic Table
To see how the periodic arrangement works, we can look at the filling order of electrons, using the n ℓm s scheme.
Hydrogen (100½) 1s1
Helium (100 -½) 1s2
This completes the n = 1 shell. The valance electrons are denoted using the notation 1sx.

5. Second Row

6. N = 2 Shell This completes the n = 2 shell
The valance electrons are denoted using the notation 2sx 2py
The complete electronic configuration is 1s2 2sx 2py, which may also be expressed as [He] 2sx 2py
It also introduces the p block

7. Periodic Table Blocks
The complete electronic configuration is 1s2 2sx 2py, which may also be expressed as [He] 2sx 2py.

8. Third Row

9. Periodic Table Blocks
This completes the n = 3 s and p blocks. However, since ℓ can now = 2, 3d electrons are possible.
Argon acts like a filled shell, so we show it as the end of row 3.

10. Electronic Energy Levels
Note that 4s electrons are lower energy than 3d
We need to consider the filling order of orbits, which depends on their energy levels.
We see that the 4s orbitals are lower energy than the 3d, so they fill first.

11. Beginning of Fourth Row
Potassium [Ar] 4s1
Calcium [Ar] 4s2

12. Electronic Energy Levels
Note that 3d electrons are lower energy than 4p

13. First Transition Row
Note the behavior of Cr and Cu, where the 3d half shell and full shell fill at the expense of a 4s electron.

14. Periodic Table Blocks
This completes the filling of the 3d orbital, which generates a new block on the periodic table. We then fill the 4p orbital.

15. 4p Block

16. Periodic Table Blocks

17. Electronic Energy Levels
Note that 5s electrons are lower energy than 4d
The energy level diagram shows that the 5s electrons fill next.
Rubidium [Kr] 5s1
Strontium [Kr] 5s2
As before, the 4d orbitals fill before the 5p. This generates the second transition row.

18. Second Transition Row Yttrium [Kr] 4d1 5s2 Zirconium [Kr] 4d2 5s2
Niobium [Kr] 4d4 5s1
Molybdenum [Kr] 4d5 5s1
(Technetium) [Kr] 4d5 5s2
Ruthenium [Kr] 4d7 5s1
Rhodium [Kr] 4d8 5s1
Palladium [Kr] 4d10 5s0
Silver [Kr] 4d9 5s2
Cadmium [Kr] 4d10 5s2
Note the filling order of Nb, Mo, Ru, Pd, and Ag, which borrow 1 or 2 5s electrons in order to add additional 4d electrons.

19. Periodic Table Blocks
This completes the second transition row. However, the n = 4 shell is not filled, since ℓ can = 3, allowing 4f orbitals to exist.

20. Electronic Energy Levels
Note that 5p and 6s electrons are lower energy than 4f
The energy level chart shows that both the 5p and 6s orbitals are lower energy than the 4f, however.

21. 5p Block

22. Periodic Table Blocks
As before Xe acts like a filled shell, and is an inert gas.

23. Electronic Energy Levels
Note that 6s electrons are lower energy than 4f or 5d
Following xenon, the energy level diagram shows that the 6s electrons are the next lowest energy.

24. Row 6 Beginning
Next we have lanthanum, which begins by putting one electron into the 5d orbital.

25. Electronic Energy Levels
Note that 4f electrons are lower energy than 5d
Following lanthanum, a look at the energy table shows that the 4f orbitals are of slightly lower energy than the 5d, so the next electron goes into 4f.
In addition, we need to realize that the 4f orbitals are closer to the nucleus than the 6s or 5d, and thus play little role in the valance state of their ions.

26. 4f Block

27. Periodic Table Blocks
This completes the filling of the 4f orbitals, and completes the first “f” block on the periodic table.
These elements, usually beginning with lanthanum, are known either as the lanthanides or the rare earth elements, very commonly abbreviated as REE.
The elements from cerium through lutetium are known as the lanthanide series.
Next, the remainder of the 5d orbitals are filled in order, with no irregularities.

28. 5d Block
Hafnium [Xe] 4f14 5d2 6s2 Tantalum [Xe] 4f14 5d3 6s2 Tungsten [Xe] 4f14 5d4 6s2 Rhenium [Xe] 4f14 5d5 6s2 Osmium [Xe] 4f14 5d6 6s2 Iridium [Xe] 4f14 5d7 6s2 Platinum [Xe] 4f14 5d1 6s1 Gold [Xe] 4f14 5d10 6s1 Mercury [Xe] 4f14 5d10 6s2

29. 6p Block

30. Periodic Table Blocks
Radon is at the end of the row, so it is again an inert gas. Radon is followed by two elements filling the 7f orbitals.
Francium [Rn] 7s1
Radium [Rn] 7s2
Immediately after radium is actinium, which begins the filling of the 5f orbital. Thorium follows, and the 5f electron jumps to the 6d orbitals.

31. 5f Block

32. Periodic Table Blocks
This completes the 5f block, commonly known as the actinides. Naturally occurring elements end at uranium, although atomic reactors and bombs have produced a considerable amount of plutonium.
Mendeleev noted that the columns of the periodic table contain elements with similar properties, and predicted the existence of some elements not known in 1869.
Today, we have names for some of the columns.
He corrected the atomic weights of some elements. One of such elements was indium (In) known to occur in zinc (Zn) ores and at the time assumed to form the oxide InO (similar to that of zinc, ZnO). From the percent composition of the oxide (82.5% In) and its assumed formula, indium had been assigned an atomic weight of 76. This atomic weight would have placed Indium, a metal, between arsenic and selenium, a metalloid and a nonmetal. Mendeleev proposed that indium formed an oxide In2O3 and that its true atomic weight was 113. Mendeleev then placed indium between cadmium and tin, both metals. Mendeleev also corrected the atomic weights of beryllium (Be) from 13.5 to 9 and uranium (from 120 to 240).
Column 1 is known as the alkali metals. All of the elements are strongly electropositive, and have a valance of +1 in compounds. Column 2 is known as the alkaline earths. These metals have a valance of +2 in compounds.
Columns 3 through 12 are the transition elements. The first transition row contains many elements which are important in mineralogy. They are characterized by multiple oxidation states. Examples:

33. Oxidation States of First Transition Row ElementsSc Ti V Cr Mn Fe Co Ni Cu Zn +3 +4 +5 +6 +7 +2 +1
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Why are multiple oxidation states possible? When the transition metals become ions, they lose their 4s electrons first. The low ionization potentials mean that the elements show variable valency states by loss of electrons from the 4s and 3d orbitals.
We need to look at the electronic configurations of the ions. Chromium is 3d5 4s1. It loses the 4s electron, and either 2, or all 5 3d electrons. Manganese is 3d5 4s2, It loses the two 4s electrons, and either two or five 3d electrons.
Iron is 3d6 4s2. It also loses the two 4s electrons, and sometimes the paired 3d electron. Copper is 3d10 4s1. It can lose only the 4s electron, but sometimes loses a 3d electron as well. . The transition elements readily form alloys with themselves and with other elements (e.g. a copper-tin alloy is used for mirrors, brass is a copper-zinc alloy). Tungsten, because of its high melting point, is used to make tools and filaments in incandescent light bulbs. The atomic size is fairly constant since the electrons in the outer most shells have similar environments.
The rare earth elements are commonly found together. The principal source of rare earths is the mineral monazite. The lanthanides usually exist as trivalent cations, in which case their electronic configuration is (Xe) 4fn, with n varying from 1 (Ce3+) to 14 (Lu3+). Because the 4f electrons are shielded by the 5s and 5p electrons, the outer configuration of the ions is very similar, and they behave almost as a single element. Many of these elements were discovered late, because they are separated with difficulty.
As the atomic number increases across the lanthanide row, size decreases. The increasing nuclear charge pulls the outer electrons in. Since the filling 4f orbitals are inside the outermost electrons, they play little role in size or in bonding. The decrease of ionic radius increasing atomic number is called the lanthanide contraction. According to Goldschmidt’s rules, among ions with identical charge and differing radii, the smaller ions will be incorporated into crystal structures first. Garnets do show preferential concentration of heavy REE’s.
States highlighted in red are the most common

34. REE Abundance
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The rare earth elements are somewhat misnamed, since three of them (La, Ce & Nd) are more common than lead. The abundance pattern shows the typical odd/even sawtooth, and a general decrease in abundance as Z increases. Only promethium, which has no stable isotopes, is truly rare.

35. Europium Anomaly
Figure 9-5. REE diagram for 10% batch melting of a hypothetical lherzolite with 20% plagioclase, resulting in a pronounced negative Europium anomaly
From Winter, (2001) An Introduction to Igneous and Metamorphic Petrology, Prentice Hall
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Europium, [Xe] 4f7 5d0 6s2 , is often found in the +2 state if the oxygen fugacity is low. The two 6s electrons are removed, but the 4f subshell is half full, and therefore somewhat more stable. Eu+2 can readily substitute for Ca+2, and is often found in plagioclase. Since europium in the 2+ state decreases the amount of Eu3+ present, there is often a decrease in a plot of REE concentration versus atomic number at the europium position. This is known as the europium anomaly.

36. Uranium and Thorium Decay
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Among the actinide elements only actinium, thorium, protactinium and uranium occur naturally (i.e. Z ≤ 92). The trans-uranium elements have all been produced synthetically, although there is evidence that plutonium existed on earth during its early history. All isotopes of all the actinides are radioactive, and most decay initially by alpha decay. Helium gas is sometimes associated with deposits of actinide containing minerals. The most common elements are uranium (99.3% 238U, 0.7% 235U) and thorium (232Th). The decay schemes of all three result in different isotopes of lead.
The electronic configurations of Actinides are not always easy to confirm, since the atomic spectra of heavy elements are very difficult to interpret in terms of configuration. There is competition between 5fn7s2 and 5fn-16d17s2 configurations. For light actinides, promotion 5f to 6d occurs to provide more bonding electrons, and is much easier than corresponding 4f to 5d promotion in lanthanides. The second half of actinide series resemble lanthanides more closely.

37. Periodic Table Blocks
Group 16 are the oxides, with configuration ns2 np4. All members of this group, with the exception of oxygen itself, are capable of being in the +6 or -2 state.
Oxygen is almost always -2.
Group 17 are the halogens, with configuration ns2 np5.
They usually have a charge of -1 (always in mineralogy).
Group 18 are the Nobel or inert gases. Since they are ns2 np6, they are not reactive, and do not form compounds in nature.

38. Ionization Potential Element Electronic Configuration (valance)
Size, nm
Ionization Energy,
Cal/mol Be
2s2
0.112
215 Mg
3s2
0.160
176 Ca
4s2
0.197
141 Sr
5s2
0.215
125 Ba
6s2
0.222
119
The periodic table can be used to explain many observed trends.
For example, the ionization energy, designated I, is defined as energy an atom in a gaseous form must absorb in order to lose an electron and become an ion.
It is observed that the larger the atom, the smaller the ionization energy.
This is because the larger atoms have electrons further from the nucleus, where they are less strongly held.