University:
Jackson College
Course:
CEM-141 | General Chemistry I
Academic year:

2024
Views:

159

Pages:

14
Author:

Henry Cherry

Electrons in the outermost s and p orbitals. These are the electrons most often involved in bonding. The organization of electrons in an atom or ion, from the lowest energy orbital (“s”) to the highest energy orbital (“f”). s p d f FULL ELECTRON CONFIGURATION: • ALL electrons are shown, according to the energy level and orbital type. • Start at n = 1 • Fill each energy level before moving on to the next. EXAMPLES: 1s1 1s2 1s2 2s1 1s2 2s2 1s2 2s2 2p1 1s2 2s2 2p5 1s2 2s2 2p6 Noble gases have a FULL VALENCE SHELL, that is, a STABLE OCTET 1s2 2s2 2p6 3s2 3p6 4s2 1s2 2s2 2p6 3s2 3p6 4s2 3d1 Note the change in the numbering as you move from 4s to 3d. Note that all of the exponents add up to the atomic number of the element. 1s2 2s2 2p6 3s2 3p6 4s2 3d8 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d5 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f2 Note the change in the numbering as you move from 6s element # 56, Ba to element #57, La which starts to fill its 4f orbitals. 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d6 Note the change in the numbering as you move from 6s element # 56, Ba to elements #57 through 70, the Lanthanides, which fill all 14 of the 4f orbitals, and then back up to the 5d orbitals that begins with Lu (element # 71) and ends with Hg. (Ac, # 89 to No, # 102 fill their 4f orbitals) (Lr, # 103 to Cn , #112 fill their 6 d orbitals) CORE ELECTRON CONFIGURATION: • Show the electrons in the energy levels past the PREVIOUS noble gas. EXAMPLES: 2s1 2s2 2p5 4s2 3d8 4s2 3d10 4p2 5s2 4d5 5s2 4d10 6s2 4f14 5d4 6s2 4f14 5d10 ORBITAL REPRESENTATION DIAGRAMS: Show the pairs of electrons in each orbital. EXAMPLES: 1s2 2s2 2p5 ↑↓ ↑↓ ↑↓ 1s2 2s2 2p6 ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ 4s2 [Ar] 3s2 ↑↓ ↑↓ 3p6 x 4s2 ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ 4s2 3d10 [Ar] 6s2 ↑↓ ↑ 3d8 ↑↓ [Xe] ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ 4f14 ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ 4p2 ↑↑_ 5d4 ↑ ↑ ↑ ↑ _ A “drawing” that represents the nucleus and its outermost electrons. Core Electron Config 1s1 ↑ 1s2 ↑↓ [He] 2s1 ↑ [He] 2s2 ↑↓ [He] 2s2 2p1 ↑↓ ↑ _ _ [He] 2s2 2p2 ↑↓ ↑ ↑ _ [He] 2s2 2p3 ↑↓ ↑ ↑ ↑ [He] 2s2 2p4 ↑↓ ↑↓ ↑ ↑ [He] 2s2 2p5 ↑↓ ↑↓ ↑↓ ↑ [He] 2s2 2p6 ↑↓ ↑↓ ↑↓ ↑↓ Lewis Dot Diagram IONS: Atoms usually lose s and p electrons first since these are the outermost electrons. EXAMPLES: ION CORE CONFIGURATION [He] 2s 2p 2 LEWIS DOT 6 Added one electron [Ar] 4s2 Lost 2 electrons [Ar] 4s2 3d10 4p1 Lost 3 valence p and then s electrons [Ar] 4s2 3d10 4p2 [He] 2s2 2p 6 [He] 2s 2p 6 2 Sample Lewis dot diagrams for ions: ELECTRONIC CONFIGURATIONS EXCITED STATE: Atoms absorb energy and exist at higher energy levels. As the atom returns to a lower energy level, it gives off energy. This relates to emission spectra / spectral lines. Example: Carbon atom. [He] [He] 2s2 ↑↓ 2p ↑ 2 2s1 ↑ 2p ↑ 3 ↑ _ ↑ ↑ Electron excited from s orbital to a higher energy level. This makes a more stable arrangement, as having all orbitals half filled is more stable than having some orbitals with pairs, some with singles, and some empty. This means that C can easily combine with 4 atoms such as H to make CH4. More on this in Organic Chemistry. Consider the element GALLIUM, atomic number 31. This neutral (uncharged) element has 31 electrons. We NOW know that each of the electrons exist in complex three dimensional orbitals: Since Gallium is in the 4 th row of the periodic table: n=4 0 < ℓ < n – 1, so ℓ= 0, 1, 2, 3, meaning that orbitals represented by EACH of the l values, s, p, d, f, are possible . (But they may not all be filled by electrons!!!) The mℓ and ms just tell us the orientations in space of the orbitals and the spin of the paired electrons so it is NOT necessary to look at these each time you write an electron configuration. NOTE: The complete understanding of this concept will occur over several classes, and requires you to practice, daily, until you reach proficiency. Reading these pages is not enough to gain a full understanding. This only scratches the surface. In fact, the practice that we will do in class will put this into a more practical approach, rather than the mathematical approach as indicated by the quantum numbers, above. Full electronic configuration Showing ALL of the electrons in every energy level from n = 1, for Ga: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p1 Core electronic configuration Showing electrons of an element, starting at the PREVIOUS noble gas [Ar] 4s2 3d10 4p1 Orbital representation diagram Showing paired electrons with opposite spins, in each orbital: Some textbooks, and even teacher worksheets, show full electron configurations as ENERGY DIAGRAMS, so that the energy differences of the electrons in each level is clearly demonstrated: Practice done in class, with the teacher, on the white board, is essential for the understanding of electron configurations. We will start with basic FULL configurations, understanding the progression of the electron configurations from Hydrogen to Helium to Lithium, to Beryllium, etc. We will learn how to best distribute electrons in p orbitals so that we know when to have paired and unpaired electrons. As we enter the 3rd energy level, and d orbital electrons are possible, it is necessary to understand how to write the MOST STABLE electron configurations. And entering the 4th energy level, which introduces f orbital electron distributions, we must understand how an element would fill it’s orbitals in the proper energy level order. Practice in class will be accompanied by three dimensional images and videos on screen, as well has examples of emission spectra, so that you can understand this work that you are doing, IN CONTEXT. The sample questions will evolve organically, as we develop understanding of each aspect, so providing a full notes package on the practice questions would not be beneficial to your understanding, nor is it possible to predict each question in the order it will be presented. Once we are comfortable, we will also introduce ions, so that we understand WHERE an element would lose or gain electrons from, in order to achieve the most stable configuration. This will lead us into the ionic and covalent (molecular) bonding that we will be studying next. You will learn better by being able to work through this unit in consultation with your lab partners.

Electronic Configurations

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