# Electron Configurations

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1 Section 3 Electron Configurations Key Terms electron configuration Pauli exclusion principle noble gas Aufbau principle Hund s rule noble-gas configuration Main Ideas Electrons fill in the lowest-energy orbitals first. There are three ways to indicate electron configuration. No electrons can occupy a higher-energy sublevel until the sublevel below it is filled. The quantum model of the atom improves on the Bohr model because it describes the arrangements of electrons in atoms other than hydrogen. The arrangement of electrons in an atom is known as the atom s electron configuration. Because atoms of different elements have different numbers of electrons, a unique electron configuration exists for the atoms of each element. Like all systems in nature, electrons in atoms tend to assume arrangements that have the lowest possible energies. The lowest- energy arrangement of the electrons for each element is called the element s ground-state electron configuration. A few simple rules, combined with the quantum number relationships discussed in Section 2, allow us to determine these ground-state electron configurations. Main Idea Electrons fill in the lowest-energy orbitals first. To build up electron configurations for the ground state of any particular atom, first the energy levels of the orbitals are determined. Then electrons are added to the orbitals, one by one, according to three basic rules. (Remember that real atoms are not built up by adding protons and electrons one at a time.) The first rule shows the order in which electrons occupy orbitals. According to the Aufbau principle, an electron occupies the lowest-energy orbital that can receive it. Figure 3.1 shows the atomic orbitals in order of increasing energy. The orbital with the lowest energy is the 1s orbital. In a ground-state hydrogen atom, the electron is in this orbital. The 2s orbital is the next highest in energy, then the orbitals. Beginning with the third main energy level, n = 3, the energies of the sublevels in different main energy levels begin to overlap. Note in the figure, for example, that the 4s sublevel is lower in energy than the 3d sublevel. Therefore, the 4s orbital is filled before any electrons enter the 3d orbitals. (Less energy is required for two electrons to pair up in the 4s orbital than for those two electrons to occupy a 3d orbital.) Once the 3d orbitals are fully occupied, which sublevel will be occupied next? Figure 3.1 Energy 6d 5f 7s 6p 5d 4f 6s 5p 4d 5s 4p 3d 4s 3p 3s 2s 7s 6s 5s 4s 3s 2s > Virginia standards CH.2 The student will investigate and understand that the placement of elements on the periodic table is a function of their atomic structure. The periodic table is a tool used for the investigations of: CH.2.e periods. CH.2.g electron configurations, valence electrons, and oxidation numbers. CH.2.EKS-11; CH.2.EKS-12; CH.2.EKS-13 Atomic Sublevels The order of increasing energy for atomic sublevels is shown on the vertical axis. Each individual box represents an orbital. 6p 5p 4p 3p 6d 5d 4d 3d 5f 4f 1s 1s Arrangement of Electrons in Atoms 105

2 Figure 3.2 Pauli Exclusion Principle According to the Pauli exclusion principle, an orbital can hold two electrons of opposite spin states. In this electron configuration of a helium atom, each arrow represents one of the atom s two electrons. The direction of the arrow indicates the 1s orbital electron s spin state. The second rule reflects the importance of the spin quantum number. According to the Pauli exclusion principle, no two electrons in the same atom can have the same set of four quantum numbers. The principal, angular momentum, and magnetic quantum numbers specify the energy, shape, and orientation of an orbital. The two values of the spin quantum number reflect the fact that for two electrons to occupy the same orbital, they must have opposite spin states (see Figure 3.2). The third rule requires placing as many unpaired electrons as possible in separate orbitals in the same sublevel. In this way, electron- electron repulsion is minimized so that the electron arrangements have the lowest energy possible. According to Hund s rule, orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin state. Applying this rule shows, for example, that one electron will enter each of the three p orbitals in a main energy level before a second electron enters any of them. This is illustrated in Figure 3.3. What is the maximum number of unpaired electrons in a d sublevel? Figure 3.3 Hund s Rule The figure shows how (a) two, (b) three, and (c) four electrons fill the p sublevel of a given main energy level according to Hund s rule. (a) (b) (c) Main Idea There are three ways to indicate electron configuration. Three methods, or notations, are used to indicate electron configurations. Two of these notations will be discussed for the first-period elements, hydrogen and helium, on the following page. The third notation is used mostly with elements of the third period and higher and will be discussed later in this section. In a ground-state hydrogen atom, the single electron is in the lowestenergy orbital, the 1s orbital. The electron can be in either one of its two spin states. Helium has two electrons, which are paired in the 1s orbital. 106 Chapter 4

3 Orbital Notation In orbital notation, an unoccupied orbital is represented by a line,, with the orbital s name written underneath the line. An orbital containing one electron is represented as. An orbital containing two electrons is represented as, showing the electrons paired and with opposite spin states. The lines are labeled with the principal quantum number and sublevel letter. For example, the orbital notations for hydrogen and helium are written as follows: H _ 1s He _ 1s Electron-Configuration Notation Electron-configuration notation eliminates the lines and arrows of orbital notation. Instead, the number of electrons in a sublevel is shown by adding a superscript to the sublevel designation. The hydrogen configuration is represented by 1s 1. The superscript indicates that one electron is present in hydrogen s 1s orbital. The helium configuration is represented by 1s 2. Here the superscript indicates that there are two electrons in helium s 1s orbital. Electron Configurations Sample Problem A The electron configuration of boron is 1s 2 2s 2 1. How many electrons are present in an atom of boron? What is the atomic number for boron? Write the orbital notation for boron. solve The number of electrons in a boron atom is equal to the sum of the superscripts in its electron- configuration notation: = 5 electrons. The number of protons equals the number of electrons in a neutral atom. So we know that boron has 5 protons and thus has an atomic number of 5. To write the orbital notation, first draw the lines representing orbitals. _ 1s _ 2s _ Next, add arrows showing the electron locations. The first two electrons occupy n = 1 energy level and fill the 1s orbital. _ 1s _ 2s _ The next three electrons occupy the n = 2 main energy level. Two of these occupy the lower- energy 2s orbital. The third occupies a higher-energy p orbital. _ 1s _ 2s _ Answers in Appendix E 1. The electron configuration of nitrogen is 1s 2 2s 2 3. How many electrons are in a nitrogen atom? What is the atomic number of nitrogen? Write the orbital notation for nitrogen. 2. The electron configuration of fluorine is 1s 2 2s 2 5. What is the atomic number of fluorine? How many of its p orbitals are filled? Arrangement of Electrons in Atoms 107

4 Figure 3.4 The Aufbau Principle Follow the diagonal arrows from the bottom of one column to the top of the next to get the order in which atomic orbitals are filled according to the Aufbau principle. 1s 2s 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f Main Idea No electrons can occupy a higher-energy sublevel until the energy sublevel below it is filled. In the first-period elements, hydrogen and helium, electrons occupy the orbital of the first main energy level. Figure 3.4 provides a pattern to help you remember the order in which orbitals are filled according to the Aufbau principle. The ground-state configurations in Figure 3.5 illustrate how the Aufbau principle, the Pauli exclusion principle, and Hund s rule are applied to atoms of elements in the second period. 6s 6p 6d 6f 7s 7p 7d critical thinking Interpret Review the statement of the Aufbau principle given at the beginning of this section, then study Figure 3.4 carefully. Based on your understanding of the two, explain in your own words why electrons fill atomic orbitals in this seemingly unusual order. According to the Aufbau principle, after the 1s orbital is filled, the next electron occupies the s sublevel in the second main energy level. Thus, lithium, Li, has a configuration of 1s 2 2s 1. The electron occupying the 2s level of a lithium atom is in the atom s highest, or outermost, occupied level. The highest-occupied energy level is the electron-containing main energy level with the highest principal quantum number. The two electrons in the 1s sublevel of lithium are no longer in the outermost main energy level. They have become inner-shell electrons, which are electrons that are not in the highest-occupied energy level. The fourth electron in an atom of beryllium, Be, must complete the pair in the 2s sublevel because this sublevel is of lower energy than the sub level. With the 2s sublevel filled, the sublevel, which has three vacant orbitals of equal energy, can be occupied. One of the three p orbitals is occupied by a single electron in an atom of boron, B. Two of the three p orbitals are occupied by unpaired electrons in an atom of carbon, C. And all three p orbitals are occupied by unpaired electrons in an atom of nitrogen, N. Hund s rule applies here, as is shown in the orbital notations in Figure 3.5. According to the Aufbau principle, the next electron must pair with another electron in one of the orbitals rather than enter the third main energy level. The Pauli exclusion principle allows the electron to pair with one of the electrons occupying the orbitals as long as the spins of the paired electrons are opposite. Thus, atoms of oxygen, O, have the configuration 1s 2 2s 2 4. Oxygen s orbital notation is shown in Figure 3.5. Two orbitals are filled in fluorine, F, and all three are filled in neon, Ne. Atoms such as those of neon, which have the s and p sublevels of their highest occupied level filled with eight electrons, are said to have an octet of electrons. Examine the periodic table inside the back cover of the text. Notice that neon is the last element in the second period. 110 Chapter 4

5 Figure 3.5 Electron Configurations of Atoms of Second-Period Elements Showing Two Notations Name Symbol 1s 2s Orbital notation Electronconfiguration notation Lithium Li _ _ _ _ _ 1s 2 2s 1 Beryllium Be _ _ _ _ _ 1s 2 2s 2 Boron B _ _ _ _ _ 1s 2 2s 2 1 Carbon C _ _ _ Nitrogen N _ _ _ _ Oxygen O _ _ _ _ _ _ 1s 2 2s 2 2 _ 1s 2 2s 2 3 _ 1s 2 2s 2 4 Fluorine F _ _ _ _ _ 1s 2 2s 2 5 Neon Ne _ _ _ _ _ 1s 2 2s 2 6 Elements of the Third Period After the outer octet is filled in neon, the next electron enters the s sublevel in the n = 3 main energy level. Thus, atoms of sodium, Na, have the configuration 1s 2 2s 2 6 3s 1. Compare the configuration of a sodium atom with that of an atom of neon in Figure 3.5. Notice that the first 10 electrons in a sodium atom have the same configuration as a neon atom, 1s 2 2s 2 6. In fact, the first 10 electrons in an atom of each of the third-period elements have the same configuration as neon. This similarity allows us to use a shorthand notation for the electron configurations of the third-period elements. Noble-Gas Notation Neon is a member of the Group 18 elements. The Group 18 elements (helium, neon, argon, krypton, xenon, and radon) are called the noble gases. To simplify sodium s notation, the symbol for neon, enclosed in square brackets, is used to represent the complete neon configuration: [Ne] = 1s 2 2s 2 6. This allows us to write sodium s electron configuration as [Ne]3s 1, which is called sodium s noble-gas notation. Figure 3.6, on the next page, shows the electron configuration of each of the third-period elements using noble-gas notation. The last element in the third period is argon, Ar, which is a noble gas. As in neon, the highest-occupied energy level of argon has an octet of electrons, [Ne]3s 2 3p 6. In fact, each noble gas other than He has an electron octet in its highest energy level. A noble-gas configuration refers to an outer main energy level occupied, in most cases, by eight electrons. Arrangement of Electrons in Atoms 111

6 Figure 3.6 Name Electron Configurations of Atoms of Third-Period Elements Symbol Atomic number Number of electrons in sublevels 1s 2s 3s 3p Noble-gas notation Sodium Na *[Ne]3s 1 Magnesium Mg [Ne]3s 2 Aluminum Al [Ne]3s 2 3p 1 Silicon Si [Ne]3s 2 3p 2 Phosphorus P [Ne]3s 2 3p 3 Sulfur S [Ne]3s 2 3p 4 Chlorine Cl [Ne]3s 2 3p 5 Argon Ar [Ne]3s 2 3p 6 *[Ne] = 1s 2 2s 2 6 Elements of the Fourth Period The electron configurations of atoms in the fourth-period elements are shown in Figure 3.7 (on the next page). The period begins by filling the 4s orbital, the empty orbital of lowest energy. The first element in the fourth period is potassium, K, which has the electron configuration [Ar]4s 1. The next element is calcium, Ca, which has the electron configuration [Ar]4s 2. With the 4s sublevel filled, the 4p and 3d sublevels are the next available vacant orbitals. Figure 3.1 (on the first page of this section) shows that the 3d sublevel is lower in energy than the 4p sublevel. Therefore, the five 3d orbitals are next to be filled. A total of 10 electrons can occupy the 3d orbitals. These are filled successively in the 10 elements from scandium (atomic number 21) to zinc (atomic number 30). Scandium, Sc, has the electron configuration [Ar]3d 1 4s 2. Titanium, Ti, has the configuration [Ar]3d 2 4s 2. And vanadium, V, has the configuration [Ar]3d 3 4s 2. Up to this point, three electrons with the same spin have been added to three separate d orbitals, as required by Hund s rule. Surprisingly, chromium, Cr, has the electron configuration [Ar]3d 5 4s 1. Not only did the added electron go into the fourth 3d orbital, but an electron also moved from the 4s orbital into the fifth 3d orbital, leaving the 4s orbital with a single electron. Chromium s electron configuration is contrary to what is expected according to the Aufbau principle. However, in reality the [Ar]3d 5 4s 1 configuration is of lower energy than a [Ar]3d 4 4s 2 configuration. For chromium, having six orbitals with unpaired electrons is a more stable arrangement than having four unpaired electrons in the 3d orbitals and forcing two electrons to pair up in the 4s orbital. On the other hand, for tungsten, W, which is in the same group as chromium, having four electrons in the 5d orbitals and two electrons paired in the 6s orbital is the most stable arrangement. There is no simple explanation for such deviations from the expected order given in Figure Chapter 4

7 Manganese, Mn, has the electron configuration [Ar]3d 5 4s 2. The added electron goes to the 4s orbital, completely filling this orbital while leaving the 3d orbitals still half-filled. Beginning with the next element, electrons continue to pair in the d orbitals. Thus, iron, Fe, has the configuration [Ar]3d 6 4s 2 ; cobalt, Co, has the configuration [Ar]3d 7 4s 2 ; and nickel, Ni, has the configuration [Ar]3d 8 4s 2. Next is copper, Cu, in which an electron moves from the 4s orbital to pair with the electron in the fifth 3d orbital. The result is an electron configuration of [Ar]3d 10 4s 1 the lowest- energy configuration for Cu. As with Cr, there is no simple explanation for this deviation from the expected order. In atoms of zinc, Zn, the 4s sublevel is filled to give the electron configuration [Ar]3d 10 4s 2. In atoms of the next six elements, electrons add one by one to the three 4p orbitals. According to Hund s rule, one electron is added to each of the three 4p orbitals before electrons are paired in any 4p orbital. Figure 3.7 Name Electron Configuration of Atoms of Elements in the Fourth Period Symbol Atomic number Number of electrons in sublevels above 3s 3p 3d 4s 4p Noble-gas notation Potassium K *[Ar]4s 1 Calcium Ca [Ar]4s 2 Scandium Sc [Ar]3d 1 4s 2 Titanium Ti [Ar]3d 2 4s 2 Vanadium V [Ar]3d 3 4s 2 Chromium Cr [Ar]3d 5 4s 1 Manganese Mn [Ar]3d 5 4s 2 Iron Fe [Ar]3d 6 4s 2 Cobalt Co [Ar]3d 7 4s 2 Nickel Ni [Ar]3d 8 4s 2 Copper Cu [Ar]3d 10 4s 1 Zinc Zn [Ar]3d 10 4s 2 Gallium Ga [Ar]3d 10 4s 2 4p 1 Germanium Ge [Ar]3d 10 4s 2 4p 2 Arsenic As [Ar]3d 10 4s 2 4p 3 Selenium Se [Ar]3d 10 4s 2 4p 4 Bromine Br [Ar]3d 10 4s 2 4p 5 Krypton Kr [Ar]3d 10 4s 2 4p 6 *[Ar] = 1s 2 2s 2 6 3s 2 3p 6 Arrangement of Electrons in Atoms 113

8 Elements of the Fifth Period In the 18 elements of the fifth period, sublevels fill in a similar manner as in elements of the fourth period. However, they start at the 5s orbital instead of at the 4s orbital. Successive electrons are added first to the 5s orbital, then to the 4d orbitals, and finally to the 5p orbitals. This can be seen in Figure 3.8. There are occasional deviations from the predicted configurations here also. The deviations differ from those for fourthperiod elements, but in each case the preferred configuration has the lowest possible energy. Figure 3.8 Electron Configurations of Atoms of Elements in the Fifth Period Name Symbol Atomic number Number of electrons in sublevels above 3d 4s 4p 4d 5s 5p Noble-gas notation Rubidium Rb *[Kr]5s 1 Strontium Sr [Kr]5s 2 Yttrium Y [Kr]4d 1 5s 2 Zirconium Zr [Kr]4d 2 5s 2 Niobium Nb [Kr]4d 4 5s 1 Molybdenum Mo [Kr]4d 5 5s 1 Technetium Tc [Kr]4d 6 5s 1 Ruthenium Ru [Kr]4d 7 5s 1 Rhodium Rh [Kr]4d 8 5s 1 Palladium Pd [Kr]4d 10 Silver Ag [Kr]4d 10 5s 1 Cadmium Cd [Kr]4d 10 5s 2 Indium In [Kr]4d 10 5s 2 5p 1 Tin Sn [Kr]4d 10 5s 2 5p 2 Antimony Sb [Kr]4d 10 5s 2 5p 3 Tellurium Te [Kr]4d 10 5s 2 5p 4 Iodine I [Kr]4d 10 5s 2 5p 5 Xenon Xe [Kr]4d 10 5s 2 5p 6 *[Kr] = 1s 2 2s 2 6 3s 2 3p 6 3d 10 4s 2 4p Chapter 4

9 Premium Content Electron Configurations Sample Problem B a. Write both the complete electron-configuration notation and the noble-gas notation for iron, Fe. Learn It! Video HMDScience.com Solve It! Cards HMDScience.com b. How many electron-containing orbitals are in an atom of iron? How many of these orbitals are completely filled? How many unpaired electrons are there in an atom of iron? In which sublevel are the unpaired electrons located? solve a. The complete electron-configuration notation of iron is 1s 2 2s 2 6 3s 2 3p 6 3d 6 4s 2. The periodic table inside the back cover of the text reveals that 1s 2 2s 2 6 3s 2 3p 6 is the electron configuration of the noble gas argon, Ar. Therefore, as shown in Figure 3.7, iron s noble-gas notation is [Ar]3d 6 4s 2. b. An iron atom has 15 orbitals that contain electrons. They consist of one 1s orbital, one 2s orbital, three orbitals, one 3s orbital, three 3p orbitals, five 3d orbitals, and one 4s orbital. Eleven of these orbitals are filled, and there are four unpaired electrons. They are located in the 3d sublevel. The notation 3d 6 represents: 3d _ _ _ _ _ Answers in Appendix E 1. a. Write both the complete electron- configuration notation and the noble-gas notation for iodine, I. How many inner-shell electrons does an iodine atom contain? b. How many electron-containing orbitals are in an atom of iodine? How many of these orbitals are filled? How many unpaired electrons are there in an atom of iodine? 2. a. Write the noble-gas notation for tin, Sn. How many unpaired electrons are there in an atom of tin? b. How many electron-containing d orbitals are there in an atom of tin? Name the element in the fourth period whose atoms have the same number of electrons in their highest energy levels that tin s atoms do. 3. a. Write the complete electron configuration for the element with atomic number 25. You may use the diagram shown in Figure 3.2. b. Identify the element described in item 3a. 4. a. How many orbitals are completely filled in an atom of the element with atomic number 18? Write the complete electron configuration for this element. b. Identify the element described in item 4a. Elements of the Sixth Period The sixth period consists of 32 elements. It is much longer than the periods that precede it in the periodic table. To build up electron configurations for elements of this period, electrons are added first to the 6s orbital in cesium, Cs, and barium, Ba. Then, in lanthanum, La, an electron is added to the 5d orbital. Arrangement of Electrons in Atoms 115

10 With the next element, cerium, Ce, the 4f orbitals begin to fill, giving cerium atoms a configuration of [Xe]4f 1 5d 1 6s 2. In the next 13 elements, the 4f orbitals are filled. Next the 5d orbitals are filled and the period is completed by filling the 6p orbitals. Because the 4f and the 5d orbitals are very close in energy, numerous deviations from the simple rules occur as these orbitals are filled. The electron configurations of the sixth-period elements can be found in the periodic table inside the back cover of the text. The seventh period is incomplete and consists largely of synthetic elements. Electron Configurations Sample Problem C a. Write both the complete electron-configuration notation and the noblegas notation for a rubidium atom. b. Identify the elements in the second, third, and fourth periods that have the same number of highest-energy-level electrons as rubidium. solution a. 1s 2 2s 2 6 3s 2 3p 6 3d 10 4s 2 4p 6 5s 1, [Kr]5s 1 b. Rubidium has one electron in its highest energy level (the fifth). The elements with the same outermost configuration are, in the second period, lithium, Li; in the third period, sodium, Na; and in the fourth period, potassium, K. Answers in Appendix E 1. a. Write both the complete electron- configuration notation and the noble-gas notation for a barium atom. b. Identify the elements in the second, third, fourth, and fifth periods that have the same number of highest- energy-level electrons as barium. 2. a. Write the noble-gas notation for a gold atom. b. Identify the elements in the sixth period that have one unpaired 6s electron. Section 3 Formative ASSESSMENT Reviewing Main Ideas 1. a. What is an atom s electron configuration? b. What three principles guide the electron configuration of an atom? 2. What three methods are used to represent the arrangement of electrons in atoms? 3. What is an octet of electrons? Which elements contain an octet of electrons? 4. Write the complete electron-configuration notation, the noble-gas notation, and the orbital notation for the following elements: a. carbon b. neon c. sulfur 5. Identify the elements having the following electron configurations: a. 1s 2 2s 2 6 3s 2 3p 3 b. [Ar]4s 1 c. contains four electrons in its third and outer main energy level d. contains one set of paired and three unpaired electrons in its fourth and outer main energy level Critical Thinking 6. RELATING IDEAS Write the electron configuration for the third-period elements Al, Si, P, S, and Cl. Is there a relationship between the group number of each element and the number of electrons in the outermost energy level? 116 Chapter 4

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