Deer Valley Unified School District



Lesson 7.5 Electron Configurations I Suggested Reading Zumdahl Chapter 7 Sections 7.9-7.11Essential QuestionWhat is the electronic structure of atoms?Learning ObjectivesExplain the Pauli Exclusion Principle.Define the Aufbau Principle.Deduce the electron arrangement for atoms and ions up to Z = 20.Apply the relationship between the electron arrangement of elements and their position in the periodic table.Apply the relationship between the number of elements in the highest occupied energy level for an element and its position in the periodic table.Differentiate between the valence and core electrons.Identify the number and location of the valence electron in an atoms.Apply Hund's rule.IntroductionIn the previous lesson you learned that an electron in an atom has four quantum numbers, n, l, ml, and ms. The first three quantum numbers describe the region of space where the electron is most likely to be found; we say the electron occupies the orbital. The spin quantum number,?ms, describes the spin orientation of an electron. In this lesson, we will look at electron spin and then we will discuss how electrons are distributed among the possible orbitals in electron configurations.Electron Spin?When a beam of hydrogen atoms is exposed to a magnet, it is split into two. Half of the atoms are bent in one direction and half are bent in the other direction. The beam of hydrogen atoms is split in two because the electron in each atom behaves like a tiny magnet. The electron acts like a ball of spinning charge, and like a circulating electric charge it creates a magnetic field. Electron spin however, is subject to a quantum restriction on the possible direction of the spin axis. The resulting direction of spin magnetism corresponds to spin quantum number ms = +1/2 and ms = -1/2. As you will see shortly, electron spin must be taken into account when writing electron configurations.Electron Configurations and Orbital DiagramsAn electron configuration of an atom is the particular distribution of electrons among the subshells of the orbitals. We can use the quantum mechanical model of the atom to show how the electrons are arranged in the orbitals of various atoms. In order to write the electron configurations for atoms we assume that as protons are added to the nucleus?one by one?to "build up" the elements, one electron is added fro each proton. This is called the building up (aufbau principle).?When we write electron configurations, they usually correspond to the ground state of the atom. Recall that the ground state is the lowest energy level of the atom. Other configurations correspond to excited states, which are associated with energy levels other than the lowest.For example, hydrogen has one electron, which occupies the 1s orbital in its ground state. The configuration for hydrogen is written as 1s1, which can be represented by the following orbital diagram.?The arrow represents the electron spinning in a particular direction.?The?Pauli Exclusion Principle?states that no two electrons in an atom can have the same four quantum numbers. If one electron in an atom has the quantum numbers?n?= 1,?l?= 0,?ml?= 0, and?ms?= +1/2, no other electron can have these same quantum numbers. Since electrons is the same orbital have the same values of?n,?l,?and?ml, this postulate says they must differ in the values of ms.?Thus, an orbital can hold at most two electron, and then only if the electrons have opposite spins.?Electrons with the same spin repel each other like magnets do when like poles are moved towards one another.?The next element is helium (Z = 2) has 2 electrons, so its orbital diagram illustrates the Pauli Exclusion Principle.?Lithium has three electrons, two of which can go into the 1s orbital. Since there is only one orbital in the n = 1 level, the other electron must go into the next lowest energy level, n = 2. The allowable values for the angular momentum quantum number (l) when n = 2 are 0 and 1, so there are both s and p orbitals at the n = 2 level. Since there is one s orbital (ml = 0) and three p orbitals (ml = -1, 0, +1) there are a total of 4 orbitals at the n = 2 level. We fill these orbitals as we build up electrons across the periodic table according to the building up principle. Lets illustrate this by looking at the electron configurations for the next several elements.?The next element is carbon. Before we can show the electron configuration for carbon, we must introduce another concepts called Hund's Rule. Hund's rule states that the lowest energy configuration of a subshell is obtained by putting electrons into separate orbitals of the subshell with the same spin before pairing electrons. Lets illustrate this by showing the electron configurations for C, N, and O. Watch.The electron configurations of the remaining elements are built up in a similar fashion. Note that orbitals hold a maximum of twice as many electrons as the number of orbitals in the subshell. Thus, a 2p subshell, which has three orbitals (ml?= -1, 0, and +1), can hold a maximum of six electrons.The following video will help illustrate these concepts.?YouTube VideoMnemonic Diagram for the?Building Up Principle?Following the principle you obtain the electron configuration of an atom by successively filling subshells in the following order:1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5fUse the following mnemonic diagram to help you remember the order.?? ?You obtain this diagram by writing the subshells in rows, each row having subshell(s) of given n. Within each row you arrange the subshells in order of increasing l.?The building up order corresponds, for the most part, to the increasing energy of the orbitals. Thus, the lowest energy orbitals (ground state) are filled first.??Electron Configurations and the Periodic TableBy now you can see a pattern develop among the ground-state electron configurations of the atoms. This pattern explains the period table. Consider the Group VIIIA elements helium, neon, argon, and krypton. These elements are members of a group called the noble gases and are characterized by their relative unreactivity, which results from having an outer shell that is completely filled with electrons.He 1s2Ne 1s22s22p6Ar ?1s22s22p63s23p6Kr ?1s22s22p63s23p63d104s24p6??Now look at the electron configurations for the Group IIA alkaline earth metals, which are similar, moderately reactive elements. ?Be ?1s22s2?? ??? ??? ??? ??? ??? ?? ?? ??? ??? ??? ??? ??? ?or [He]2s2 ???? ??? ?? ??Mg ?1s22s22p63s2?? ??? ??? ??? ??? ???? or [Ne]3s2??? ??? ?? ??Ca ?1s22s22p63s23p63d104s2?? ?? ?? ?or [Ar]4s2?The symbolism on the rights is referred to as the condensed electron configuration. Each of these configurations consists of a noble-gas core, that is, an inner-shell configuration corresponding to one of the noble gases, plus two outer electrons win an ns2?configuration.?The Group IIIA elements boron, aluminum, and gallium also have similarities. There configurations are?B? ??1s22s22p1? ?? ??? ??? ??? ??? ??? ??? ??? ??? ??? ??? ??? ??? ?? ???? ?or [He]2s22p1????? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ??Al? ?1s22s22p63s23p1?? ??? ??? ??? ??? ??? ??? ?? ??or [Ne]3s23p1 ???Ga ?1s22s22p63s23p63d104s24p1?? ?? ??or [Ar]3d104s24p1?Boron and aluminum have noble-gas cores plus three electrons with the configuration ns2np1. Gallium has an additional filled 3d subshell. The noble gas core together with the (n-1)d10 electrons are sometimes referred to as the pseudo-noble-gas core electrons because these electrons usually are note involved in chemical reactions.An electron in an atom outside of the noble-gas of?pseudo-noble-gas core is called a valence electron. Valence electrons are primarily involved in chemical reactions, and similarities among the configurations or valence electrons (the valence shell configuration) account for the similarities of the chemical properties among a group of elements. This is an important concept that you must be aware of as it is important in much of chemistry.?Look at the periodic table with the valence shell electron configurations. Note the similarity in electron configuration within any group (column) of the elements. This similarity explains why the properties of elements in any group are similar. There is a better picture on page 307 of your textbook.?Homework Problems:Practice exercises 6.7-6.9Book questions pg. 323 questions 67,69,71,73 ................
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