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Chapter Eight Electron Configurations and Periodicity Electron Spin In Chapter 7, we saw that electron pairs residing in the same orbital are required to have opposing spins. This causes electrons to behave like tiny bar magnets. (See Figure 8.3) 2 1 An electron configuration of an atom is a particular distribution of electrons among available subshells. The notation for a configuration lists the subshell symbols sequentially with a superscript indicating the number of electrons occupying that subshell. For example, lithium (at. # 3) has two electrons in the 1s subshell and one electron in the 2s sub shell 1s2 2s1. An orbital diagram is used to show how the orbitals of a subshell are occupied by electrons. Each orbital is represented by a circle. Each group of orbitals is labeled by its subshell notation. The Boron atom 1s 2s 2p Electrons are represented by arrows: up for ms = +1/2 and 3 down for ms = -1/2 The Pauli exclusion principle, which summarizes experimental observations, states that no two electrons can have the same four quantum numbers. - In other words, an orbital can hold at most two electrons, and then only if the electrons have opposite spins. The maximum number of electrons and their orbital diagrams are: Maximum Number of Number of Sub shell Orbitals Electrons s (l = 0) 1 2 p (l = 1) 3 6 d (l =2) 5 10 f (l =3) 7 14 4 2 Aufbau Principle Every atom has an infinite number of possible electron configurations. The configuration associated with the lowest energy level of the atom is called the ground state. Other configurations correspond to excited states. Table 8.1 lists the ground state configurations of atoms up to krypton. The Aufbau principle is a scheme used to reproduce the ground state electron configurations of atoms by following the building up order. You need to remember the order of filling the orbitals (next slide) 5 Order for Filling Atomic Subshells 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f 1s 2s 3s 4s 5s 6s 2p 3p 4p 5p 6p 3d 4d 4f 5d 5f 6d 6f 6 3 Aufbau Principle The building up order corresponds for the most part to increasing energy of the subshells. By filling orbitals of the lowest energy first, you usually get the lowest total energy ( ground state ) of the atom. Remember, the number of electrons in the neutral atom equals the atomic number, Z. 7 Table 8.1 8 4 Question Which ground-state electron configuration is INCORRECT? 1) Cr: [Ar] 3d6 2) Ca: [Ar] 4s2 3) Na: 1s2 2s2 2p6 3s1 4) Zn: [Ar] 3d10 4s2 5) Kr: [Ar] 3d10 4s2 4p6 9 Here are a few examples. Using the abbreviation [He] for 1s2, the configurations are: Z=4 Beryllium 1s22s2 or [He]2s2 Z=3 Lithium Z=5 Z=6 Z=7 Z=8 Z=9 Z=10 Boron Carbon Nitrogen Oxygen Fluorine Neon 1s22s1 or [He]2s1 1s22s22p1 1s22s22p2 1s22s22p3 1s22s22p4 1s22s22p5 1s22s22p6 or or or or or or [He]2s22p1 [He]2s22p2 [He]2s22p3 [He]2s22p4 [He]2s22p5 [He]2s22p6 10 5 With sodium (Z = 11), the 3s subshell begins to fill. Z=11 Sodium 1s22s22p63s1 or [Ne]3s1 Z=12 Magnesium 1s22s22p23s2 or [Ne]3s2 Then the 3p subshell begins to fill. Z=13 : : Z=18 Aluminum 1s22s22p63s23p1 or [Ne]3s23p1 Argon 1s22s22p63s23p6 or [Ne]3s23p6 11 PROBLEM Give (1) the expected ground-state electron configurations (full as well as abbrev.) (2) the orbital -filling diagrams for; a. P (Z=15) b. Zn (Z=30) c. Ca (Z=20) 6 Note that elements within a given family have similar configurations. The Group IIA elements are sometimes called the alkaline earth metals. Beryllium 1s22s2 Magnesium 1s22s22p63s2 Calcium 1s22s22p63s23p64s2 13 Configurations and the Periodic Table Electrons that reside in the outermost shell of an atom or in other words, those electrons outside the noble gas core are called valence electrons. These electrons are primarily involved in chemical reactions. Elements within a given group have the same valence shell configuration. This accounts for the similarity of the chemical properties among groups of elements. The total number of valence electrons for an atom equals its group number. 14 7 Another way of writing the electron configurations is the Orbital Diagram Consider carbon (Z = 6) with the ground state configuration 1s22s22p2. Hund s rule states that the lowest energy arrangement (the ground state ) of electrons in a subshell is obtained by putting electrons into separate orbitals of the subshell with the same spin before pairing electrons. Carbon's orbital diagram is 1s 2s 2p 15 Table 8.2 16 8 Question? Which of the following orbital filling diagrams represent : 1) ground state or 2) excited state or 3) forbidden for Boron (Z= 5) 1s ( ) ( 2s ) ( ) ( ) ( ) ( ( ) ( ) 2p ( ) ( ) ( ) ( ) ) ( ) ( ) ( ) ( ) ( ) ( ) ( ) ( ) Magnetic Properties Although an electron behaves like a tiny magnet, two electrons that are opposite in spin cancel each other. Only atoms with unpaired electrons exhibit magnetic susceptibility. A paramagnetic substance is one that is weakly attracted by a magnetic field, usually the result of unpaired electrons. A diamagnetic substance is not attracted by a magnetic field generally because it has only paired electrons. 18 9 Question The ground-state electron configuration of a Fe (s) is 1s2 2s2 2p6 3s2 3p6 4s23d6. Therefore, Fe is 1) diamagnetic. 2) paramagnetic with one unpaired electron. 3) paramagnetic with two unpaired electrons. 4) paramagnetic with four unpaired electrons. 5) paramagnetic with five unpaired electrons. 19 The periodic law states that when the elements are arranged by atomic number, their physical and chemical properties vary periodically. We will look at three periodic properties: Atomic radius, Ionization energy, Electron affinity A. Atomic Radius Trends 1) Within each period (horizontal row), the atomic radius tends to decrease with increasing atomic number (nuclear charge). 2) Within each group (vertical column), the atomic radius tends to increase with the period number. 20 10 Figure 8.10: Representation of atomic radii (covalent radii) of the maingroup elements 21 Two factors determine the size of an atom. 1) One factor is the principal quantum number, n. The larger is n, the larger the size of the orbital. 2) The other factor is the effective nuclear charge, which is the positive charge an electron experiences from the nucleus minus any shielding effects from intervening electrons. 22 11 B. Ionization energy The first ionization energy of an atom is the minimal energy needed to remove the highest energy (outermost) electron from the neutral atom. For a lithium atom, the first ionization energy is illustrated by: Li(1s 2 2s1 ) Li (1s 2 ) e Ionization energy = 520 kJ/mol 23 Ionization energy trends There is a general trend that ionization energies increase with atomic number within a given period. This follows the trend in size, as it is more difficult to remove an electron that is closer to the nucleus. For the same reason, we find that ionization energies, again following the trend in size, decrease as we descend a column of elements. 24 12 Periodic Properties Ionization energy The electrons of an atom can be removed successively. The energies required at each step are known as the first ionization energy, the second ionization energy, and so forth. Table 8.3 lists the successive ionization energies of the first ten elements. 25 C. Electron Affinity The electron affinity is the energy change for the process of adding an electron to a neutral atom in the gaseous state to form a negative ion. For a chlorine atom, the first electron affinity is illustrated by: Cl([Ne]3s 2 3p 5 ) e Cl ([Ne]3s 2 3p 6 ) Electron Affinity = -349 kJ/mol The more negative the electron affinity, the more stable the negative ion that is formed. The general trend goes from lower left to upper right as electron affinities become more negative. Table 8.4 gives the electron affinities of the main-group elements. 26 13 Table 8.4 27 The Main-Group Elements The physical and chemical properties of the main-group elements clearly display periodic behavior. Variations of metallic-nonmetallic character. Basic-acidic behavior of the oxides. Group IA, Alkali Metals Largest atomic radii React violently with water to form H2 Readily ionized to 1+ Oxides dissolved in water form basic solutions 28 14 Group IIA, Alkali Earth Metals Readily ionized to 2+ React with water to form H2 Oxides dissolved in water form basic solutions Group III A Metals (except for boron) Several oxidation states (commonly 3+) Group IV A Form the most covalent compounds Oxidation numbers vary between 4+ and 4- Group V A Form anions generally(1-, 2-, 3-), though positive 29 oxidation states are possible Question Sodium oxide, Na2O, is added to excess water. Which one of the following would occur? 1) reaction to yield an aqueous solution of sodium hydroxide 2) reaction to yield hydrogen and an aqueous solution of sodium hydroxide 3) reaction to yield oxygen and an aqueous solution of sodium hydroxide 4) reaction to yield oxygen, hydrogen, and an aqueous solution of sodium hydroxide 5) no reaction 30 15 Question An unknown element, Y, reacts with oxygen to form an oxide with the general formula Y2O3. The pure element is paramagnetic and the oxide, when added to water, forms an acidic solution. Which of the following elements could be described by these characteristics? 1) Na 2) N 3) Ca 4) S 5) F 31 Group VI A anions generally, though positive oxidation states are possible React vigorously with alkali and alkali earth metals, and nonmetals Halogens Form monoanions High electronegativity (electron affinity) Diatomic gases Most reactive nonmetals (F) Noble Gases Minimal reactivity Monatomic gases; Closed shell 32 16 This document was created with Win2PDF available at http://www.daneprairie.com. The unregistered version of Win2PDF is for evaluation or non-commercial use only.