Atomic Structure and the Periodic Table

Transcription

1 Atomic Structure and the Periodic Table In the few years following the announcement of the Bohr theory, a series of revisions to this model occurred. Bohr s single quantum number (n) was expanded to a total of four quantum numbers (n, l, m l, m s ). These quantum numbers were necessary to explain a variety of evidence associated with spectral lines and magnetism. In addition, these same quantum numbers also greatly improved the understanding of the periodic table and chemical bonding. You will recall that the atomic theory you used previously allowed only a limited description of electrons in atoms up to atomic number 20, calcium. In this section, you will see that the four quantum numbers improve the theoretical description to include all atoms on the periodic table and they improve the explanation of chemical properties. Section 3.5 provided the empirical and theoretical background to quantum numbers. The main thing that you need to understand is that there are four quantized values that describe an electron in an atom. Quantized means that the values are restricted to certain discrete values the values are not on a continuum like distance during a trip. There are quantum leaps between the values. Table 4 in Section 3.5 illustrates the values to which the quantum numbers are restricted. The advantage of quantized values is that they add some order to our description of the electrons in an atom. In this section, the picture of the atom is based upon the evidence and concepts from Section 3.5, but the picture is presented much more qualitatively. For example, as you shall see, the secondary quantum number values of 0, 1, 2, and 3 are presented as s, p, d, and f designations to represent the shape of the orbitals (Table 1). What is truly amazing about the picture of the atom that is coming in this section is that the energy description in Section 3.5 fits perfectly with both the arrangement of electrons and the structure of the periodic table. The unity of these concepts is a triumph of scientific achievement that is unparalleled in the past or present. Table 1 Values and Letters for the Secondary Quantum Number value of l letter designation* s p d f name designation sharp principal diffuse fundamental * This is the primary method of communicating values of l later in this section. Electron Orbitals Although the Bohr theory and subsequent revisions were based on the idea of an electron travelling in some kind of orbit or path, a more modern view is that of an electron orbital. A simple description of an electron orbital is that it defines a region (volume) of space where an electron may be found. Figure 1 and Table 2 present some of the differences between the concepts of orbit and orbital. At this stage in your chemistry education, the four quantum numbers apply equally well to electron orbits (paths) or electron orbitals (clouds). A summary of what is coming is presented in Table Figure 1 Orbitals are like electron clouds. This computer-generated image shows a 3d orbital of the hydrogen atom, which has four symmetrical lobes (in this image, two blue and two red-orange), with the nucleus at the centre. The bands in the lobes show different probability levels: the probability of finding an electron decreases while moving away from the nucleus. This is quite a different image from the Bohr electron orbits. orbital a region of space around the nucleus where an electron is likely to be found Table 2 Orbits and Orbitals Orbits Orbitals 2-D path 3-D region in space fixed distance variable from nucleus distance from nucleus circular or no path; elliptical path varied shape of region 2n 2 electrons 2 electrons per orbit per orbital NEL Atomic Theories 185

2 Table 3 Energy Levels, Orbitals, and Shells Principal energy Energy Energy in Additional energy level sublevel magnetic field differences n l m l m s shell orbital shape orbital orientation electron spin subshell shell main energy level; the shell number is given by the principal quantum number, n; for the representative elements the shell number also corresponds to the period number on the periodic table for the s and p subshells subshell orbitals of different shapes and energies, as given by the secondary quantum number, l; the subshells are most often referred to as s, p, d, and f The first two quantum numbers (n and l) describe electrons that have different energies under normal circumstances in multi-electron atoms. The last two quantum numbers (m l, m s ) describe electrons that have different energies only under special conditions, such as the presence of a strong magnetic field. In this text, we will consider only the first two quantum numbers, which deal with energy differences for normal circumstances. As we move from focusing on the energy of the electrons to focusing on their position in space, the language will change from using main (principal) energy level to shell, and from energy sublevel to subshell. The terms can be taken as being equivalent, although the contexts of energy and space can be used to decide when they are primarily used. Rather than a complete mathematical description of energy levels using quantum numbers, it is common for chemists to use the number for the main energy level and a letter designation for the energy sublevel (Table 4). For example, a orbital, a orbital, a 3d, or a 4f orbital in that energy sub-level can be specified. This symbol is simpler than communicating n 1, l 0, and is simpler than n 2, l 1. Notice that this orbital description includes both the principal quantum number and the secondary quantum number; e.g., 5s,,3d,or 4f. Including the third quantum number, m l, requires another designation, for example x, y, and z. Table 4 Classification of Energy Sublevels (Subshells) Value of l Sublevel symbol Number of orbitals 0 s 1 1 p 3 2 d 5 3 f 7 Although the s-p-d-f designation for orbitals is introduced here, the shape of these orbitals is not presented until Section 3.7. The emphasis in this section is on more precise energy-level diagrams and their relationship to the periodic table and the properties of the elements. Creating Energy-Level Diagrams Our interpretation of atomic spectra is that electrons in an atom have different energies. The fine structure of the atomic spectra indicates energy sublevels. The designation of these energy levels has been by quantum number. Now we are going to use energy-level diagrams to indicate which orbital energy levels are occupied by electrons for a particular atom or ion. These energy-level diagrams show the relative energies of electrons in various orbitals under normal conditions. Note that the previous energy-level diagrams that you have drawn included only the principal quantum number, n. Now you are going to extend these diagrams to include all four quantum numbers. 186 Chapter 3 NEL

3 Section e 6s 6p 5d 4f 18e 5s 5p 4d 18e 4s 4p 3d 3p 8e 3s 8e 2e In Figure 2, you see that as the atoms become larger and the main energy levels become closer together, some sublevels start to overlap in energy. This figure summarizes the experimental information from many sources to produce the correct order of energies. A circle is used to represent an electron orbital within an energy sublevel. Notice that the energy of an electron increases with an increasing value of the principal quantum number, n. For a given value of n, the sublevels increase in energy, in order, s<p<d<f. Note also that the restrictions on the quantum numbers (Table 4) require that there can be only one s orbital, three p orbitals, five d orbitals, and seven f orbitals. Completing this diagram for a particular atom provides important clues about chemical properties and patterns in the periodic table. We will now look at some rules for completing an orbital energy-level diagram and then later use these diagrams to explain some properties of the elements and the arrangement of the periodic table. Figure 2 Diagram of relative energies of electrons in various orbitals. Each orbital (circle) can potentially contain up to two electrons. NEL Atomic Theories 187

4 H (a) He (b) Figure 3 Energy-level diagrams for (a) hydrogen and (b) helium atoms Figure 4 Energy-level diagrams for lithium, carbon, and fluorine atoms. Notice that all of the orbitals at the same energy are shown, even though some are empty. Pauli exclusion principle no two electrons in an atom can have the same four quantum numbers; no two electrons in the same atomic orbital can have the same spin; only two electrons with opposite spins can occupy any one orbital aufbau principle aufbau is German for building up; each electron is added to the lowest energy orbital available in an atom or ion Hund s rule one electron occupies each of several orbitals at the same energy before a second electron can occupy the same orbital Li (a) C (b) F (c) In order to show the energy distribution of electrons in an atom, the procedure will be restricted to atoms in their lowest or ground state, assuming an isolated gaseous atom. You show an electron in an orbital by drawing an arrow, pointed up or down to represent the electron spin (Figure 3a). It does not matter if you point the arrow up or down in any particular circle, but two arrows in a circle must be in opposite directions (Figure 3b). This is really a statement of the Pauli exclusion principle, which requires that no two electrons in an atom have the same four quantum numbers. Electrons (arrows) are placed into the orbitals (circles) by filling the lowest energy orbitals first. An energy sublevel must be filled before moving onto the next higher sublevel. This is called the aufbau principle. If you have several orbitals at the same energy (e.g., p, d,or f orbitals), one electron is placed into each of the orbitals before a second electron is added. In other words, spread out the electrons as much as possible horizontally before doubling up any pair of electrons. This rule is called Hund s rule.you follow this procedure until the number of electrons placed in the energy-level diagram for the atom is equal to the atomic number for the element (Figure 4). According to these rules, when electrons are added to the second (n 2) energy level, there are s and p sublevels to fill with electrons (Figure 2). The lower energy s sublevel is filled before the p sublevel is filled. According to Hund s rule, one electron must go into each of the p orbitals before a second electron is used for pairing (Figure 4). There are several ways of memorizing and understanding the order in which the energy levels are filled without having the complete chart shown in Figure 2. One method is to use a pattern like the one shown in Figure 5. In this aufbau diagram, all of the orbitals with the same principal quantum number are listed horizontally. You can follow the diagonal arrows starting with the ls orbital to add the required number of electrons. An alternate procedure for determining the order in which energy levels are filled comes from the arrangement of elements in the periodic table. As you move across the periodic table, each atom has one more electron (and proton) than the previous atom. Because the electrons are added sequentially to the lowest energy orbital available (aufbau principle), the elements can be classified by the sublevel currently being filled (Figure 6). To obtain the correct order of orbitals for any atom, start at hydrogen and move from left to right across the periodic table, filling the orbitals as shown in Figure 6.Check to see that this gives exactly the same order as shown in Figure 5. 7s 6s 5s 4s 7p 6p 5p 4p 7d 6d 5d 4d 7f 6f 5f 4f 3s 4s 5s 6s 7s 3d 4d 5d 6d 3p 4p 5p 6p Figure 5 In this aufbau diagram, start at the bottom () and add electrons in the order shown by the diagonal arrows. You work your way from the bottom left corner to the top right corner. 3s 3p 3d 4f 5f Figure 6 Classification of elements by the sublevels that are being filled 188 Chapter 3 NEL

5 Section 3.6 Drawing Energy-Level Diagrams for Atoms SAMPLE problem Draw the electron energy-level diagram for an oxygen atom. Since oxygen (O) has an atomic number of 8, there are 8 electrons to be placed in energy levels. As the element is in period 2, there are electrons in the first two main energy levels. Using either the aufbau diagram (Figure 5) or the periodic table (Figure 6), we can see that the first two electrons will occupy the orbital. O The next two electrons will occupy the orbital. O The next three electrons are placed singly in each of the orbitals. The last (eighth) electron must be paired with one of the electrons in the orbitals. It does not matter into which of the three p orbitals this last electron is placed. The final diagram is drawn as shown. Note that this energy-level diagram is not drawn to scale. The actual gap in energy is much larger between the and levels than between the and levels. Example Draw the energy-level diagram for an iron atom. Solution 3d 4s O 3p 3s Fe Creating Energy-Level Diagrams for Anions The energy-level diagrams for anions, or negatively charged ions, are done using the same method as for atoms. The only difference is that you need to add the extra electrons corresponding to the ion charge to the total number of electrons before proceeding to distribute the electrons into orbitals. This is shown in the following sample problem. NEL Atomic Theories 189

6 SAMPLE problem Drawing Energy-Level Diagrams for Anions Draw the energy-level diagram for the sulfide ion. Sulfur has an atomic number of 16 and is in period 3. A sulfide ion has a charge of 2, which means that it has two more electrons than a neutral atom. Therefore, we have 18 electrons to distribute in three principal energy levels. Using either the aufbau diagram (Figure 5) or the periodic table (Figure 6), we can see that the first two electrons will occupy the orbital. The next two electrons will occupy the orbital, and six more electrons will complete the orbitals. The next two electrons fill the 3s orbital, which leaves the final six electrons to completely fill the 3p orbitals. Notice that all orbitals are now completely filled with the 18 electrons. 3p 3s S 2 Creating Energy-Level Diagrams for Cations For cations, positively charged ions, the procedure for constructing energy-level diagrams is slightly different than for anions. You must draw the energy-level diagram for the corresponding neutral atom first, and then remove the number of electrons (corresponding to the ion charge) from the orbitals with the highest principal quantum number, n. The electrons removed might not be the highest-energy electrons. However, in general, this produces the correct arrangement of energy levels based on experimental evidence. SAMPLE problem Drawing Energy-Level Diagrams for Cations Draw the energy-level diagram for the zinc ion. First, we need to draw the diagram for the zinc atom (atomic number 30). Using either the aufbau diagram (Figure 5) or the periodic table (Figure 6), we can see that the 30 electrons are distributed as follows: The first two electrons will occupy the orbital. The next two electrons will occupy the orbital, and six more electrons complete the orbitals. The next two electrons fill the 3s orbital, and six more electrons complete the 3p orbitals. The next two electrons fill the 4s orbital and the final 10 electrons fill the 3d orbitals. The zinc ion, Zn 2, has a two positive charge, and therefore has two fewer electrons than the zinc atom. Remove the two electrons from the orbital with the highest n the 4s orbital in this example. 3d 4s 3p 3s Zn Chapter 3 NEL

7 Section 3.6 SUMMARY Electron Energy-Level Diagrams Electrons are added into energy levels and sublevels for an atom or ion by the following set of rules. Remembering the names for the rules is not nearly as important as being able to apply the rules. These rules were created to explain the spectral and periodic-table evidence for the elements. Start adding electrons into the lowest energy level () and build up from the bottom until the limit on the number of electrons for the particle is reached the aufbau principle. For anions, add extra electrons to the number for the atom. For cations, do the neutral atom first, then subtract the required number of electrons from the orbitals with the highest principal quantum number, n. No two electrons can have the same four quantum numbers; if an electron is in the same orbital with another electron, it must have opposite spin the Pauli exclusion principle. No two electrons can be put into the same orbital of equal energy until one electron has been put into each of the equal-energy orbitals Hund s rule. This process is made simpler by labelling the sections of the periodic table and then creating the energy levels and electron configurations in the order dictated by the periodic table. DID YOU KNOW? Gerhard Herzberg Spectroscopy was an essential tool in developing quantum numbers and electron energy levels. A key figure in the development of modern spectroscopy was Gerhard Herzberg ( ). His research in spectroscopy at the University of Saskatchewan and the National Research Council in Ottawa earned him an international reputation and the Nobel Prize in chemistry (1971). Practice Understanding Concepts 1. State the names of the three main rules/principles used to construct an energy-level diagram. Briefly describe each of these in your own words. 2. How can the periodic table be used to help complete energy-level diagrams? 3. Complete electron energy-level diagrams for the (a) phosphorus atom (b) potassium atom (c) manganese atom (d) nitride ion (e) bromide ion (f) cadmium ion 4. (a) Complete electron energy-level diagrams for a potassium ion and a chloride ion. (b) Which noble gas atom has the same electron energy-level diagram as these ions? Extension 5. If the historical letter designations were not used for the sublevels, what would be the label for the following orbitals, using only quantum numbers:,,, 3d? NEL Atomic Theories 191

8 electron configuration a method for communicating the location and number of electrons in electron energy levels; e.g., Mg: s 2 principal quantum number 3p 5 orbital number of electrons in orbital(s) Figure 7 Example of electron configuration SAMPLE problem Electron Configuration Electron energy-level diagrams are a better way of visualizing the energies of the electrons in an atom than quantum numbers, but they are rather cumbersome to draw. We are now going to look at a third way to convey this information. Electron configurations provide the same information as the energy-level diagrams, but in a more concise format. An electron configuration is a listing of the number and kinds of electrons in order of increasing energy, written in a single line; e.g., Li: 2 1. The order, from left to right, is the order of increasing energy of the orbitals. The symbol includes both the type of orbital and the number of electrons (Figure 7). For example, if you were to look back at the energy-level diagrams shown previously for the oxygen atom, the sulfide ion, and the iron atom, then you could write the electron configuration from the diagram by listing the orbitals from lowest to highest energy. oxygen atom, O: sulfide ion, S 2 : s 2 3p 6 iron atom, Fe: s 2 3p 6 4s 2 3d 6 Note that some of the information is lost when going from the energy-level diagram to the electron configuration, but the efficiency of the communication is much improved by using an electron configuration. Fortunately, there is a method for writing electron configurations that does not require drawing an energy-level diagram first. Let us look at this procedure. Writing Electron Configurations 1. Write the electron configuration for the chlorine atom. First, locate chlorine on the periodic table. Starting at the top left of the table, follow with your finger through the sections of the periodic table (in order of atomic number), listing off the filled orbitals and then the final orbital p 5 Figure 8 Polonium is a very rare, radioactive natural element found in small quantities in uranium ores. Polonium is also synthesized in gram quantities by bombarding Bi-209 with neutrons. The energy released from the radioactive decay of Po-210, the most common isotope, is very large (100 W/g) a half gram of the isotope will spontaneously heat up to 500 C. Surprisingly, Po-210 has several uses, including as a thermoelectric power source for satellites. You now have the electron configuration for chlorine: 2, 2, 6, 3s 2, and 3p Identify the element whose atoms have the following electron configuration: s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 4 Notice that the highest n is 6, so you can go quickly to the higher periods in the table to identify the element. The highest s and p orbitals always tell you the period number, and in this case the electron configuration finishes with 6s 2 4f 14 5d 10 6p 4 : the element must be in period 6. Going across period 6 through the two s-elements, the 14 f -elements, and the 10 d-elements, you come to the fourth element in the p-section of the periodic table. The fourth element in the 6p region of the periodic table is polonium, Po (Figure 8). Example 1 Write the electron configuration for the tin atom and the tin(ii) ion. Solution Sn: s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 2 Sn 2 : s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d Chapter 3 NEL

9 Section 3.6 Example 2 Identify the atoms that have the following electron configurations: (a) s 2 3p 6 4s 2 3d 10 4p 5 (b) s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 5 Solution (a) bromine atom, Br (b) technetium atom, Tc Shorthand Form of Electron Configurations There is an internationally accepted shortcut for writing electron configurations. The core electrons of an atom are expressed by using a symbol to represent all of the electrons of the preceding noble gas. Just the remaining electrons beyond the noble gas are shown in the electron configuration. This reflects the stability of the noble gases and the theory that only the electrons beyond the noble gas (the outer shell electrons) are chemically important for explaining chemical properties. Let s rewrite the full electron configurations for the chlorine and tin atoms into this shorthand format. Cl: s 2 3p 5 becomes Cl: [Ne] 3s 2 3p 5 Sn: s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 2 becomes Sn: [Kr] 5s 2 4d 10 5p 2 Writing Shorthand Electron Configurations Write the shorthand electron configuration for the strontium atom. Follow the same procedure as before, but start with the noble gas immediately preceding the strontium atom, which is krypton. Then continue adding orbitals and electrons until you obtain the required number of electrons for a strontium atom (two beyond krypton). Sr: [Kr] 5s 2 Example Write the shorthand electron configuration for the lead atom and the lead(ii) ion. Solution Pb: [Xe] 6s 2 4f 14 5d 10 6p 2 Pb 2 : [Xe] 6s 2 4f 14 5d 10 SAMPLE problem LEARNING TIP Electron Configurations for Cations Recall that energy-level diagrams for cations are done by first doing the energy level for the neutral atom, and then subtracting electrons from the highest principal quantum number, n. Notice in this Example that electrons are removed from the n 5 orbital. The 5p electrons are removed before the 5s electrons. SUMMARY Procedure for Writing an Electron Configuration Step 1 Step 2 Determine the position of the element in the periodic table and the total number of electrons in the atom or simple ion. Start assigning electrons in increasing order of main energy levels and sublevels (using the aufbau diagram, Figure 5, or the periodic table, Figure 6). Step 3 Continue assigning electrons by filling each sublevel before going to the next sublevel, until all of the electrons are assigned. For anions, add the extra electrons to the total number in the atom. For cations, write the electron configuration for the neutral atom first and then remove the required number of electrons from the highest principal quantum number, n. 3s 4s 5s 6s 7s 3d 4d 5d 6d 3p 4p 5p 6p 4f 5f NEL Atomic Theories 193

10 Practice Understanding Concepts 6. Identify the elements whose atoms have the following electron configurations: (a) 2 2 (b) (c) s 1 (d) s 2 3p 4 7. Write full electron configurations for each of the Period 3 elements. 8. (a) Write shorthand electron configurations for each of the halogens. (b) Describe how the halogen configurations are similar. Does this general pattern apply to other families? 9. Write the full electron configurations for a fluoride ion and a sodium ion. 10. A fluoride ion, neon atom, and sodium ion are theoretically described as isoelectronic. State the meaning of this term. 11. Write the shorthand electron configurations for the common ion of the first three members of Group 12. representative elements the metals and nonmetals in the main blocks, Groups 1-2, 13-18, in the periodic table; in other words, the s and p blocks transition elements the metals in Groups 3-12; elements filling d orbitals with electrons LEARNING TIP The lanthanides are also called the rare earths, and the elements after uranium (the highest-atomic-number naturally occurring element) are called transuranium elements. Explaining the Periodic Table The modern view of the atom based on the four quantum numbers was developed using experimental studies of atomic spectra and the experimentally determined arrangement of elements in the periodic table. It is no coincidence that the maximum number of electrons in the s, p, d, and f orbitals (Table 5) corresponds exactly to the number of columns of elements in the s, p, d, and f blocks in the periodic table (Figure 6). This by itself is a significant accomplishment that the original Bohr model could not adequately explain. Table 5 Electron Subshells and the Periodic Table Electron distribution # of groups: Period elements orbitals: s p d f Period 1 22 Period Period Period Period Groups or families in the periodic table were originally created by Mendeleev to reflect the similar properties of elements in a particular group. The noble gas family, Group 18, is a group of gases that are generally nonreactive. The electron configurations for noble gas atoms show that each of them has a filled ns 2 np 6 outer shell of electrons (Table 6). The original idea from the Bohr theory filled energy levels as stable (nonreactive) arrangements still holds, but is more precisely defined. Similar outer shell or valence electron configurations also apply to most families, in particular, the representative elements. Similarly, the transition elements can now be explained by our new theory as elements that are filling the d energy sublevel with electrons. The transition elements are sometimes referred to as the d block of elements. The 5 d orbitals can accommodate 10 electrons, and there are 10 elements in each transition-metal period (Table 6). 194 Chapter 3 NEL

11 Section 3.6 Table 6 Explaining the Periodic Table Sublevel Elements Orbitals Electrons Series of elements s and p representative d transition f lanthanides and actinides Using the same test of the theory on the lanthanides and the actinides, we can explain these series of elements as filling an f energy level. The f block of elements is 14 elements wide, as expected by filling 7 f orbitals with 14 electrons. The success of the quantum-number and s-p-d-f theories in explaining the long-established periodic table led to these approaches being widely accepted in the scientific community. Explaining Ion Charges Previously, we could not explain transition-metal ions and multiple ions formed by heavy representative metals. Now many of these can be explained, although some require a more detailed theory beyond this textbook. For example, you know that zinc forms a 2 ion. The electron configuration for a zinc atom: Zn: [Ar] 4s 2 3d 10 shows 12 outer electrons. If another atom or ion removes the two 4s electrons (the ones with the highest n) this would leave zinc with filled 3d orbitals a relatively stable state, like those of atoms with filled sub-shells: Zn 2 : [Ar] 3d 10 (Note that it is unlikely that zinc would give up 10 electrons to leave filled 4s orbitals.) Another example that illustrates the explanatory power of this approach is the formation of either 2 or 4 ions by lead. The electron configuration for a lead atom: Pb: [Xe] 6s 2 4f 14 5d 10 6p 2 shows filled 4f,5d,6s orbitals, and a partially filled 6p orbital. The lead atom could lose the two 6p electrons to form a 2 ion or lose four electrons from the 6s and 6p orbitals to form a 4 ion. (From the energy-level diagram (Figure 2), you can see that all of these outer electrons are very similar in energy and it is easier to remove fewer electrons than large numbers such as 10 and 14.) Again, our new theory passes the test of being able to initially explain what we could not explain previously. Let s put it to another test. lanthanides and actinides the 14 metals in each of periods 6 and 7 that range in atomic number from and , respectively; the elements filling the f block Explaining Magnetism To create an explanation for magnetism, let s start with the evidence of ferromagnetic (strongly magnetic) elements and write their electron configurations (Table 7). Table 7 Ferromagnetic Elements and Their Electron Configurations Ferromagnetic Electron d-orbital Pairing of d electrons element configuration filling iron [Ar] 4s 2 3d 6 1 pair; 4 unpaired cobalt [Ar] 4s 2 3d 7 airs; 3 unpaired nickel [Ar] 4s 2 3d 8 3 pairs; 2unpaired Based on the magnetism associated with electron spin and the presence of several unpaired electrons, an initial explanation is that the unpaired electrons cause the magnetism. However, ruthenium, rhodium, and palladium, immediately below iron, cobalt, and nickel in the NEL Atomic Theories 195

12 (a) (b) unmagnetized periodic table (i.e., in the same groups) are only paramagnetic (weakly magnetic) and are not ferromagnetic. The presence of several unpaired electrons may account for some magnetism, but not for the strong ferromagnetism. The eventual explanation for this anomaly is that iron, cobalt, and nickel (as smaller, closely packed atoms) are able to orient themselves in a magnetic field. The theory is that each atom acts like a little magnet. These atoms influence each other to form groups (called domains) in which all of the atoms are oriented with their north poles in the same direction. If most of the domains are then oriented in the same direction by an external magnetic field of, for example, a strong bar magnet, the ferromagnetic metal becomes a permanent magnet. However, the magnet is only permanent until dropped or heated or subjected to some other procedure that allows the domains to become randomly oriented again. (Figure 9). Ferromagnetism is a based on the properties of a collection of atoms, rather than just one atom. Paramagnetism is also explained as being due to unpaired electrons within substances where domains do not form. In other words, paramagnetism is based on the magnetism of individual atoms. Again, the theory of electron configurations is able to at least partially explain an important property of some chemicals. In this case, a full description of each electron, including its spin, is involved in the explanation. magnetized Figure 9 The theory explaining ferromagnetism in iron is that in unmagnetized iron (a) the domains of atomic magnets are randomly oriented. In magnetized iron (b) the domains are lined up to form a permanent magnet. LAB EXERCISE Quantitative Paramagnetism (p. 215) Paramagnetism is believed to be related to unpaired electrons. This lab exercise explores this relation. Anomalous Electron Configurations Electron configurations can be determined experimentally from a variety of sophisticated experimental designs. Using the rules created above, let s test our ability to accurately predict the electron configurations of the atoms in the 3d block of elements. First, the predictions: Sc: [Ar] 4s 2 3d 1 Fe: [Ar] 4s 2 3d 6 Ti: [Ar] 4s 2 3d 2 Co: [Ar] 4s 2 3d 7 V: [Ar] 4s 2 3d 3 Ni: [Ar] 4s 2 3d 8 Cr: [Ar] 4s 2 3d 4 Cu: [Ar] 4s 2 3d 9 Mn: [Ar] 4s 2 3d 5 Zn: [Ar] 4s 2 3d 10 Then the evidence: Sc: [Ar] 4s 2 3d 1 Fe: [Ar] 4s 2 3d 6 Ti: [Ar] 4s 2 3d 2 Co: [Ar] 4s 2 3d 7 V: [Ar] 4s 2 3d 3 Ni: [Ar] 4s 2 3d 8 Cr: [Ar] 4s 1 3d 5 Cu: [Ar] 4s 1 3d 10 Mn: [Ar] 4s 2 3d 5 Zn: [Ar] 4s 2 3d 10 Overall, the configurations based on experimental evidence agree with the predictions, with two exceptions chromium and copper. A slight revision of the rules for writing electron configurations seems to be required. The evidence suggests that half-filled and filled subshells are more stable (lower energy) than unfilled subshells. This appears to be more important for d orbitals compared to s orbitals. In the case of chromium, an s electron is promoted to the d subshell to create two half-filled subshells; i.e., [Ar] 4s 2 3d 4 becomes [Ar] 4s 1 3d 5. In the case of copper, an s electron is promoted to the d subshell to create a half-filled s subshell and a filled d subshell; i.e., [Ar] 4s 2 3d 9 becomes [Ar] 4s 1 3d 10 (Figure 10). The justification is that the overall energy state of the atom is lower after the promotion of the electron. Apparently, this is the lowest possible energy state for chromium and copper atoms. Figure 10 The stability of half-filled and filled subshells is used to explain the anomalous electron configurations of chromium and copper. Cr: [Ar] Cu: [Ar] Predicted Actual Cr: [Ar] Cu: [Ar] 4s 3d 4s 3d 196 Chapter 3 NEL

13 Section 3.6 Section 3.6 Questions Understanding Concepts 1. Determine the maximum number of electrons with a principal quantum number (a) 1 (c) 3 (b) 2(d) 4 2. Copy and complete Table 8. Table 8 Orbitals and Electrons in s, p, d, and f Sublevels Sublevel Value of l Number of Max. # of symbol orbitals electrons (a) s 0 (b) p 1 (c) d 2 (d) f 3 3. State the aufbau principle and describe two methods that can be used to employ this principle. 4. If four electrons are to be placed into a p subshell, describe the procedure, including the appropriate rules. 5. (a) Draw electron energy-level diagrams for beryllium, magnesium, and calcium atoms. (b) What is the similarity in these diagrams? 6. The last electron represented in an electron configuration is related to the position of the element in the periodic table. For each of the following sections of the periodic table, indicate the sublevel (s,p,d,f) of the last electron: (a) Groups 1 and 2 (b) Groups 3 to 12(transition metals) (c) Groups 13 to 18 (d) lanthanides and actinides 7. (a) When the halogens form ionic compounds, what is the ion charge of the halide ions? (b) Explain this similarity, using electron configurations. 8. The sodium ion and the neon atom are isoelectronic; i.e., have the same electron configuration. (a) Write the electron configurations for the sodium ion and the neon atom. (b) Describe and explain the similarities and differences in properties of these two chemical entities. 9. Use electron configurations to explain the common ion charges for antimony; i.e., Sb 3 and Sb Predict the electron configuration for the gallium ion, Ga 3. Provide your reasoning. 11. Evidence indicates that copper is paramagnetic, but zinc is not. Explain the evidence. 12. Predict the electron configuration of a gold atom. Provide your reasoning. 13. Use electron configurations to explain the (a) 3 charge on the scandium ion (b) 1 charge on a silver ion (c) 3 and 2 charges on iron(iii) and iron(ii) ions (d) 1 and 3 charges on the Tl 1 and Tl 3 ions 14. Carbon, silicon, and germanium all form four bonds. Explain this property, using electron configurations. Applying Inquiry Skills 15. The ingenious Stern-Gerlach experiment of 1921 is famous for providing early evidence of quantized electron spin. The experimental design called for a beam of gaseous silver atoms from an oven to be sent through a nonuniform magnetic field. There were two possible results: one predicted by classical and one by quantum theory (Figure 11). (a) Which of the expected results is likely the classical prediction and which is likely the quantum theory prediction? Explain your choice. (b) Use quantum theory, the Pauli exclusion principle, and the electron configuration of silver to explain the results of the Stern-Gerlach experiment. beam of silver atoms photographic plate magnet pole N S magnet pole field off two possible results for field on oven non-uniform magnetic field slit to focus beam Figure 11 The Stern Gerlach experiment expected two possible results: one predicted from classical magnetic theory in which all orientations of electron spin are possible and the other predicted from quantum theory in which only two orientations are possible. Making Connections 16. Prior to 1968 Canadian dimes were made from silver rather than nickel. A change was made because the value of the silver in the dime had become greater than ten cents and the dimes were being shipped out of the country to be melted down. (a) Why were the dimes shipped out of Canada before being melted? (b) If you had a box full of Canadian dimes and you wanted to efficiently separate the silver from the nickel ones, what empirical properties of silver and nickel learned in this section could assist you in completing your task? NEL Atomic Theories 197

14 (c) Use theoretical concepts learned in this section to explain your separation technique. 17. Electron spin resonance (ESR) is an analytical technique that is based on the spin of an electron. State some examples of the uses of ESR in at least two different areas. GO Magnetic Resonance Imaging (MRI) is increasingly in demand for medical diagnosis (Figure 12). (a) How is this technique similar to and different from electron spin techniques? (b) Provide some examples of the usefulness of MRI results. (c) What political issue is associated with MRI use? GO Figure 12 MRI was developed using the quantum mechanical model of the atom. 198 Chapter 3 NEL