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Electron Configuration Patterns WE CANNOT KNOW THE EXACT LOCATIONS OF ELECTRONS WITHIN ENERGY LEVELS, BUT WE CAN DESCRIBE PROBABLE REGIONS OF ELECTRON LOCATION. Our understanding of electron distribution is based on mathematical probabilities that result from knowledge about the behavior of charged particles in an atom: a) b) c) Consequently, there are three “rules” for predicting the most probable locations of electrons: 1) Electrons always enter the region of lowest energy available. This concept is called the Why does it make sense? 2) Orbitals never hold more than 2 electrons at a time. This concept is called the Why does it make sense? 3) Electrons enter separate orbitals in the same sublevel one-at-a time before pairing up. This concept is called the Why does it make sense? Try the interactive simulation in your on-line text: Electron Configuration Simulation (needs Shockwave) The chart below can be used to show the energy levels and sublevels, in order of increasing energy value. You should not memorize it, but look at the patterns it holds. Each energy level is listed (as a “coefficient”), along with the type(s) of sublevels it holds (letters s, p, d, f). The superscripts tell you the maximum number of electrons that can ever be in that sublevel (orbitals combined). With this in mind, what does “1s2” mean? What does “2p6” mean? What is the first principal energy level to have a d sublevel? How many electrons can be found in any p sublevel? How many energy levels have an s sublevel? Determining an element’s ground-state configuration: 1s2 First, determine the number of electrons you are trying 2s2 2p6 continue directional arrows as needed to configure. Start at the first diagonal arrow (it goes through 1s) and write that energy level and sublevel. 3s2 3p6 * 3d10 Add a superscript for each electron present in that sublevel, up to the maximum # allowed. If you have 4s2 4p6 4d10 4f14 more electrons to place, move to the next level and sublevel along the arrow. If you reach the end of an 5s2 5p6 5d10 5f14 … additional sublevels not used arrow, move to the next one and start at the top again. Keep track of the total number of electrons you need to 6s2 6p6 6d10 6f!4 … … show, and keep track of the maximum number of electrons allowed in each sublevel. When you finish, 7s2 7p6 7d10 7f14 … … … double-check the sum of all the superscripts. The total should match the atomic number of the element. Here is * Note: the 4s sublevel accepts electrons immediately after an example for Titanium, atomic number 22: the 3p sublevel. The 3d sublevel fills after 4s. A new principal energy level “opens up” after every 22Ti = 1s22s22p63s23p64s23d2 p-sublevel is completed. This is consistent with observed chemical behaviors of elements. Energy-order takes priority over numeric-order! Practice: Writing electron configurations. (Hint: to type a superscript, press Ctrl, Shift, and the = keys at the same time. Type your superscript. Press the same keys to toggle back to normal font. You can use the Font menu to do this, but it takes more time.) Write the electron configuration of phosphorus, P: How many unpaired electrons does it hold? Write the electron configuration of calcium, Ca: How many unpaired electrons does it hold? Use a periodic table to locate the following elements. Write the electron configuration of each one Then answer the CONCEPT QUESTIONS below each set. (Only one or a few elements have been selected from each Group. These will guide you to making important generalizations about the periodic table and element placement.) I) Alkali Metals, Group IA (write the complete configuration for each element listed) Hydrogen, H: Lithium, Li: Sodium, Na: Potassium, K: Concept Questions: Why doesn’t potassium have any electrons in the 3d sublevel? What Group Number are these elements located in? What do you notice about the configuration in the outermost energy level (called the valence level) for each of them? Do you think this pattern would be consistent for all Alkali Metals? Predict the electron configuration of the valence level only of Cesium, Cs. Write the number for the principal energy level, the letter of the sublevel, and the superscript for the number of electrons. II) Alkaline Earth Metals , Group IIA Beryllium: Magnesium: Concept Questions: What electron configuration feature places these two elements in Group IIA? Predict the electron configuration of the valence level only of Barium, Ba. III) Group IIIA Aluminum: Concept Questions: How many principal energy levels are present in atoms of aluminum? What sublevels are occupied with electrons in aluminum’s valence level? How many valence electrons does this element have altogether? Would this number of valence electrons be consistent for all elements in Group IIIA? IV) Group IVA Silicon: Concept Questions: Which principal energy level is the “valence level” of electrons of silicon? What sublevels are occupied with electrons in silicon’s valence level? How many valence electrons does this element have altogether? Name two other elements that would have the same number of electrons in their valence levels: V) Group VA (sometimes called the “pnictogen” group… but not in your book!) Nitrogen: Concept Questions: How many electrons are in the valence level of this element? How many of the valence electrons are unpaired? (Consider Hund’s Rule) VI) Group VIA (sometimes called the “chalcogen” group, but not in your book!) Sulfur: Concept Questions: How many electrons are in the valence levels of elements in this Group? How many of the valence electrons are unpaired? VII) Halogens, Group VIIA Chlorine: VIII Noble Gases, Group VIII Neon: Argon: Concept Question: What sublevel configuration pattern is common to all noble gases? Summarize: How does an element’s Group and Period placement correspond to its electron configuration? Be specific. Extend: Transition Metals, B-Groups Write the electron configuration for Scandium, Sc: Write the electron configuration for Iron, Fe: Concept Questions: What sublevel is not complete in the electron configuration of these transition metals? How many transition metals are there across any given Period? Name the transition element in the 4th Period that has a complete set of d-electrons. Lanthanides and Actinides (so-named because they follow the elements Lanthanum and Actinium, respectively) What Period are the Lanthanide elements in? What Period are the Actinide elements in? Concept Questions: Why are these elements referred to as “inner” transition metals? In general, what sublevel of electrons is being filled in the inner-transition metals?
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