The Periodic Table and Electron Configurations

3 Which orbital diagram is correct for the ground state S atom  

The placement of the outermost electron in the 45 orbital (rather than in the 2d orbital) of potas-sirtm is strongly supported by experimental evidence. The physical and chemical properties of potassirrm are very similar to those of lithirrm and sodirtm, the first two alkali metals. In both lithirrm and soditrm, the outermost electron is in an 5 orbital (there is no dorrbt that their ontermost electrons occupy 5 orbitals because there is no d or 2d subshell). Based on its similarities to the other alkah metals, we erqrect potassirrm to have an analogous electron configuration that is, we expect the last electron in potassirrm to occupy the 45 rather than the 2d orbital. 

CHAPTER 6 Quantum Theory and the Eleetronie Structure of Atoms 

Student Anrtotation Electron configurations such as these may also be written with the d subshell first For example, [Ar]4s 3d iican also be written as [Ar]3d i 4s. Either way is acceptable. 

For elements Zn (Z = 30) through Kr (Z = 36), the 3d, As, and Ap subshells fill in a straightforward manner With rubidium (Z = 37), electrons begin to enter the n = 5 energy level. 

Which of the following electron configurations correctly represents the Ti atom  

Although zinc and the other elements in Group 2B sometimes are induded under the heading transition metals, they neither have nor readily acquire partially filled f subshells. Strictly speaking, they are not transition metals. 

The first step in using the periodic table to figure out electron configurations is to note that the periods are numbered 1 through 7 from top to bottom along the left side of the table. These numbers correspond to the principal quantum numbers (values of n) for file orbitals that become filled across the period for both s and 

Each set of five d orbitals becomes filled for the transition metals in the three horizontal periods beginning with atomic numbers 21, 39, and 57 and ending with, successively, atomic numbers 30,48, and 80. For each of these orbital s, the value ofn is 1 less than the period number in which the orbitals become filled. The first d orbitals to become filled are the lowest-lying ones possible, the 3d orbitals, which become filled in the fourth period. Across the 5th period, where the 5s and 5p orbitals become filled for elements 37 and 38 and for elements 49-54, respectively, the 4d 

Atomic Number Symbol Electron Configuration Atomic Number Symbol Electron Configuration 

The chapter summary below is presented in a programmed format to review the main points covered in this chapter. It is used most effectively by filling in the blanks, referring back to the chapter as necessary. The correct answers are given at the end of the summary. 

Which of your predicted spectra from CTQ 6 of ChemActivity 10 provides the better match to the experimental spectrum, Model 1 Explain. 

Based on the analysis we have used to assign peaks in photoelectron spectra to shells and subshells in atoms, why is the peak at 0.42 MJ/mole in the K spectrum assigned to the = 4 shell (as opposed to being another subshell of = 3) Refer to the data in Table 1 of ChemActivity 10 Electron Configurations. 

In the photoelectron spectrum of Sc, the peak at 0.63 MJ/mole is assigned to the 45 subshell. Why is the peak at 0.77 MJ/mole in the Sc spectrum assigned as a third subshell of = 3 (named 3d) as opposed to being a second subshell of n = 4 (that is, 4p)  

Note that the periodic table has an unusual form. The elements are arranged in blocks of columns—a block of two columns on the left, six columns on the right, and ten columns in the middle. 

What is the relationship between the form of the periodic table and the electron configurations of the elements  

AIM To learn about the electron configurations of atoms with z greater than 18. 

The next element is calcium, with an additional electron that also occupies the 4s orbital. 

Partial electron configurations for the elements potassium through krypton. The transition metals shown in green (scandium through zinc) have the general configuration [Ar]4s 3d , except for chromium and copper. 

One of the best uses of the periodic tabie is to predict the properties of newiy discovered eiements. For exampie, the artificiaiiy synthesized eiement bohrium (Z= 107) is found in the same famiiy as manganese, technecium, and rhenium and is expected to show chemistry simiiar to these eiements. The probiem, of course, is that oniy a few atoms of bohrium can be made at a time and the atoms exist for oniy a very short time (about 1 7 seconds), it s a reai chaiienge to study the chem- 

A periodic table is almost always available to you. If you understand the relationship between the electron configuration of an element and its position on the periodic table, you can figure out the expected electron configuration of any atom. 

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Because the energies of the 4/and 5d orbitals are very close to each other, the electron configurations of some of the lanthanides involve 5d electrons. For example, the elements lanthanum (La), cerium (Ce), and praseodymium (Pr) have the following electron configurations  

Because La has a single 5d electron, it is sometimes placed below yttrium (Y) as the first member of the third series of transition elements Ce is then placed as the first member of the lanthanides. Based on its chemical properties, however. La can be considered the first element in the lanthanide series. Arranged this way, there are fewer apparent exceptions to the regular filling of the 4/orbitals among the subsequent members of the series. 

After the lanthanide series, the third transition element series is completed by the filling of the 5d orbitals, followed by the filling of the 6p orbitals. This brings us to radon (Rn), heaviest of the known noble-gas elements. 

The final row of the periodic table begins by filling the 7s orbitals. The actinide elements, of which uranium (U, element 92) and plutonium (Pu, element 94) are the best known, are then built up by completing the 5/orbitals. All of the actinide elements are radioactive, and most of them are not found in nature. 

We just saw that the electron configurations of the elements correspond to their locations in the periodic table. Thus, elements in the same column of the table have related outer-shell (valence) electron configurations. As Table 6.4 shows, for example, all 2A elements have an ns outer configuration, and all 3A elements have an n np outer configuration, with the value of n increasing as we move down each column. 

FIGURE 5.10 Electron configurations follow the order of occupied sublevels on the periodic table. 

Using the Periodic Table to Write Electron Configurations 

The s block includes hydrogen and helium as well as the elements in Group lA (1) and Group 2A (2). This means that the final one or two electrons in the elements of the s block are located in an s orbital. The period number indicates the particular s orbital that is filling Is, 2s, and so on. 

Theblock consists of the elements in Group 3A (13) to Group 8A (18). There are six p block elements in each period because three p orbitals can hold up to six electrons. The period number indicates the particular p sublevel that is filling 2p, 3p, and so on. 

The d block, containing the transition elements, first appears after calcium (atomic number 20). There are 10 elements in each period of the d block because five d orbitals can hold np to 10 electrons. The particular d sublevel is one less (n — 1) than the period nnmber. For example, in Period 4, the d block is the 3d snblevel. In Period 5, the d block is the 4d sublevel. 

To construct Table 8.3, we have taken three groups of elements from the periodic table and written their electron configurations. The similarity in electron configuration within each group is readily apparent. If the shell of the highest principal quantum number—the outermost, or valence, shell—is labeled n, then 

For the p-block elements in groups 13 to 18, the number of valence electrons is from 1 to 6. For example, aluminum is in period 3 and group 13, its valence-shell electron configuration is 3s 3p. We use n = 3 since A1 is in the third period and we have to accommodate three electrons after the neon core, which contains 10 electrons. Thus the electron configuration of A1 is 

TABLE 8.3 Electron Configurations of Some Groups of Elements 

To use this figure as a guide to the aufbau process, locate the position of an element in the table. Subshells listed ahead of this position are filled. For example, germanium (Z = 32) is located in group 14 of the blue 4p row. The filled subshells are 1s, 2s,  

and 3d °. At (Z = 32), a second electron has entered the 4p subshell. The electron configuration of Ge is [Ar]3d °4s 4p. Exceptions to the orderly filling of subshells suggested here are found among a few of the d-block and some of the f-block elements. 

We have seen that we can describe the atoms beyond hydrogen by simply filling the atomic orbitals starting with level n = and working outward in order. This works fine until we reach the element potassium (Z = 19), which is the next element after argon. Because the 3p orbitals are fully occupied in argon, we might expect the next electron to go into a 3d orbital (recall that for u = 3 the sublevels are 3s, 3p, and 3d). However, experiments show that the chemical properties of potassium are very similar to those of lithium and sodium. 

Because we have learned to associate similar chemical properties with similar valence-electron arrangements, we predict that the valence-electron configuration for potassium is 4s, resembling sodium (3s ) and lithium (2s ). That is, we expect the last electron in potassium to occupy the 4s orbital instead of one of the 3d orbitals. This means that the principal energy level 4 begins to fill before level 3 has been completed. This conclusion is confirmed by many types of experiments. So the electron configuration of potassium is 

The orbitals being filled for elements in various parts of the periodic table. Note that in going along a horizontal row (a period), the (n + 1)s orbital fills before the nd orbital. The group label indicates the number of valence electrons (the number of s plus the number of p electrons in the highest occupied principal energy level) for the elements in each group. 

In a principal energy level that has d orbitals, the s orbital from the next level fills before the d orbitals in the current level. That is, the (n + l)s orbitals always fill before the nd orbitals. For example, the 5s orbitals fill for rubidium and strontium before the 4d orbitals fill for the second row of transition metals (yttrium through cadmium). 

Valence electrons are the electrons in the outermost principal shell (the principal shell with the highest principal quantum number, n). These electrons are important because, as we will see in the next chapter, they are involved in chemical bonding. Electrons that are not in the outermost principal shell are called core electrons. For example, silicon, with the electron configuration of ls 2s 2p s 3p, has 4 valence electrons (those in the n = 3 principal shell) and 10 core electrons. 

Write an electron configuraton for selenium and identify the valence electrons and the core electrons. 

Write the electron configuration for selenium by determining the total number of electrons from selenium s atomic number (34) and then distributing them into the appropriate orbitals. 

The valence electrons are those in the outermost principal shell. For selenium, the outermost principal shell is the n = 4 shell, which contains 6 electrons (2 in the 4s orbital and 4 in the three 4p orbitals). All other electrons, including those in the 3d orbitals, are core electrons. 

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Florida State University. Electron Configurations and the Periodic Table. Available online. URE http //winel.sb.fsu.edu/ chml045/notes/Struct/EPeriod/Struct09.htm. 

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If you missed 31, go to Electronic Configuration and the Periodic Table, page 239. 

In this chapter we will learn to use electron configurations and the periodic table to predict the type of bond atoms will form, as well as the number of bonds an atom of a particular element can form and the stability of the product. 

Electron Configuration and the Periodic Table Strontium, which is used to produce red fireworks, has an electron configuration of [Kr]5s. Without using the periodic table, determine the group, period, and block of strontium. 

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