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Chemistry Chapter 5

Chemistry Chapter 5. Electrons in Atoms. 5.3 Electron Configuration. Objectives: 1. Apply the Pauli exclusion principle, the Aufbau principle and Hund’s rule to write electron configurations using orbital diagrams and electron configuration notation

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Chemistry Chapter 5

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  1. Chemistry Chapter 5 Electrons in Atoms

  2. 5.3 Electron Configuration • Objectives: • 1. Apply the Pauli exclusion principle, the Aufbau principle and Hund’s rule to write electron configurations using orbital diagrams and electron configuration notation • 2. Define valence electrons and draw electron-dot structures representing an atom’s valence electrons

  3. Background Information • The rules for electron configuration help explain atomic arrangement of protons & neutrons as well as electrons • The energy level of an electron is the region around the nucleus where the electron is likely to be moving • The energy levels in which electrons exist are fixed and run from the lowest level to higher levels • Electrons can jump from one level to another moving just the right distance using a quantum of energy • A quantum of energy is the amount of energy required to move an electron from energy level to the next highest level

  4. The energy levels are not equidistant from each other • The levels become more closely spaced as you move father away from the nucleus • Electron energy levels are designated by principal quantum numbers, with each level increasing in energy • n= 1, 2, 3, 4 and so on • The average electron distance from the nucleus increases with increasing values of n • The electron position is not confined to fixed circular path, so they are called atomic orbitals (not orbits)

  5. The number of energy sublevels equals the principal quantum number • The orbitals are named by the letters s, p, d and f • s orbitals are spherical • p orbitals are dumbbell shaped • d orbitals are clover-leaf shaped • f orbitals are complex • The p and d orbitals have regions close to the nucleus where the probability of finding the electron is low

  6. The regions where the probability of finding the electron is low are called nodes • The lowest energy level has one sublevel (1s) & one orbital • The second principal level has two sublevels & 4 orbitals (2s, 2px, 2py, & 2pz) • The third principal energy level has 3 sublevels & 9 orbitals (one 3s,three 3p, and five 3d oribitals)

  7. The fourth principal energy level has 4 sublevels & 16 orbitals (one 4s, three 4p, five 4d & seven 4f orbitals) • The maximum number of electrons that can occupy a principal energy level is given by the formula 2n2 (where n is the principal quantum number) • High energy systems are unstable • Electrons and neutrons interact to make the most stable arrangement possible • The ways electrons are arranged around the nucleus are called electron configurations

  8. Aufbau, Pauli & Hund • There are three rules regarding the electron configuration of atoms • The first is the Aufbau principle which states that electrons enter the orbitals of lowest energy first • The level s is always the lowest energy sublevel • If energy levels overlap, electrons enter the lowest level first

  9. The Pauli exclusion principle states that an atomic orbital may describe at most two electrons • One or two electrons can occupy each orbital • To occupy the same orbital the electrons must spin in the opposite direction • Spin is a quantum property of electrons • Hund’s rule states that when electrons occupy orbitals of equal energy, one electron enters each orbital until all the orbitals contain one electron with parallel spins

  10. Writing Configurations • Notice that the electrons fill out the orbitals with one electron first, then go and backfill if there are enough electrons • Half-filled orbitals are more energy efficient and more stable than having some full and some empty • The longhand method for showing electron configurations involves writing the energy level (s,p,d or f) followed by a superscript with the number of electrons (1 or 2)

  11. For oxygen with its eight electrons, the shorthand method is 1s22s22p4 • The sum of the superscript equals the number of electrons in the atoms • Electron configurations are correct using this method up to element 23 (vanadium) • Cr and Cu would be incorrect using this method (Cr:1s22s22p63s23p64s13d5 & Cu:1s22s22p63s23p64s13d10) • The half-filled & completely filled d sublevels are more stable than other configurations

  12. What you have learned so far is longhand notation for electron configurations • Noble gas notation is the shorthand way to write the configurations • To use the shorthand, go to the nearest noble gas that occurs before the element, use brackets for that noble, then continue using longhand • For oxygen: [He]2s22p4, Al:[Ne]3s23p1 & so on

  13. Valence Electrons & dot structures • Recall that valence electrons are electrons in an atom’s outermost orbitals • Def: electron dot structures consist of an element’s symbol(s), surrounded by dots representing valence electrons • The dots are placed in a box like arrangement around the symbol

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