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Electron Configuration

Electron Configuration. Chemistry. Chapter 4 – Section 2. The Quantum Model of the Atom. I. Introduction. Bohr’s Model did not make sense to most scientists, who felt that the electron should exist at any distance from the nucleus, depending on energy , rather than in certain levels.

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Electron Configuration

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  1. Electron Configuration Chemistry

  2. Chapter 4 – Section 2 • The Quantum Model of the Atom

  3. I. Introduction • Bohr’s Model did not make sense to most scientists, who felt that the electron should exist at any distance from the nucleus, depending on energy, rather than in certain levels

  4. II. Electrons as Waves • A. De Broglie proposed that an electron, like light, might exist as a particle and a wave • 1. A wave confined to a certain space can only have certain frequencies • a. This corresponded to Bohr’s orbits • b. It also corresponded to the specific frequencies produced in line-emissionspectrums

  5. 2. Beams of electrons can also be bent or refracted, or experience interference, just as light does • a. Diffraction – the bending of a wave around the edge of an object or opening • b. Interference – the overlapping of waves, which either increases or decreases each the waves’ energy

  6. III. Heisenberg Uncertainty Principle • A. The dual wave-particle nature of an electron made many scientists uneasy, because they could not pin point where the electron would be located in the atom • B. Heisenberg proposed that electrons and photons have about the sameenergy • 1. Since photons are used to knock electrons off metal (and thus, detect the electrons), the exact location of an electron cannot be determined with certainty

  7. 2. The Heisenberg Uncertainty Principle – it is impossible to determine the location and speed of the electron or photon at the same time

  8. IV. The Schrodinger Wave Equation • A. Schrodinger developed an equation to show that electrons behave as waves (only certain frequencies would solved the equation) • B. Together, Heisenberg and Schrodinger laid the foundation for our modern quantum theory – the mathematical description of the wave properties of electrons and other small particles

  9. 1.Solutions to the equations are called wavefunctions • 2.They do not locate the electron, but merely give the probability of locating it • 3.This showed that electrons do not travel in neat orbits around the nucleus as Bohr described, but exist in certain regions called orbitals • 4.Orbital– a three-dimensional region around the nucleus that indicates the probable location of the electron • 5.Orbitals can have • different shapes and sizes

  10. V. Atomic Orbitals and Quantum Numbers • A. Quantum numbers are used to describe orbitals and the properties of electrons in them • 1.The first three quantum numbers indicate: • a.Main energy level of the orbital • b.Shape of orbital • c.Orientation of orbital

  11. 2.The fourth quantum number is the spin quantum # and indicates the fundamental spin state of the electron • B.Principal Quantum Number (n) • 1.This is the main energy level of the electron • 2.As (n) increases, electron energy and distance from the nucleus increases • 3.Electrons with the same (n) value will be in the same energy level • 4.(n2) is the total number of • orbitals in each energy level

  12. C.Angular Momentum Quantum Number • 1.These sublevels are orbitals of different shapes • 2.(l) indicates the shape of the orbital and is called the (l) value • 3.If the (l) value is s, then the shape is spherical • 4.If the (l) value is p, then the shape is a dumbbell • 5.If the (l) value is d, then the shape is a clover leaf • 6.If the (l) value is f, then the shape is a flower petal s p d f

  13. D. Magnetic Quantum Number • 1.Orbitals can be the same shape, but have different orientations (axes) • 2.Magnetic Quantum Number (m) – indicates the orientation (or axis) of the orbital • 3.The (m) value = +lto -l

  14. E. Spin Quantum Number • 1.These values can be +1/2 or –1/2 • 2.This indicates the spin state of the electron • 3.Any orbital can hold 2 electrons, but they must have opposite spin states

  15. Chapter 4 – Section III • Electron Configuration

  16. I. Electron Configuration • A. Bohr’s model of the atom only described electron arrangement in the Hydrogen atom • B.The quantum model of the atom describes electron arrangement for allatoms • 1. Electron Configuration – the arrangement of electrons in an atom • 2. Electrons always assume the lowest possible energy state

  17. II. Rules Governing Electron Configurations • A. Aufbau principle – an electron occupies the lowest energy orbital that has a space for it • B. Pauli exclusion principle – no two electrons in the same atom can have the same 4 quantum numbers • 1. No two electrons can be in the same place at the same time • 2. If two electrons occupy the same orbital, then they must have opposite spin states

  18. C. Hund’s rule– orbitals of equal energy will have one electron each until all orbitals have one before a second electron enters an occupied orbital • 1. This minimizes repulsion between electrons • 2. This allows the electron to have the lowestenergy possible

  19. III. Representing Electron Configurations • A. There are 3 methods of electron configuration notation • 1. Two are used for elements in the 1st two periods • 2. The third is used for elements in the 3rd period and higher 1s2 2s2 2p6 3s2 3p5 OR [Ne]3s2 3p5 1s2 2s2 2p6 3s2 3p6 3d10 4s1 OR [Ar]3d10 4s1

  20. Carbon is paramagnetic • B. Orbital Notation • 1. A blank line (___) represents an orbital with no electrons • 2. An orbital with 1 electron is represented with___ • An orbital with its maximum number of two electrons is represented with _____(shows the opposite spins) • 4. Each line is labeled with the (n) and (l) value Neon is diamagnetic

  21. C. Electron Configuration Notation – each orbital letter has a superscript to show how many electronsare in the level • D. Elements of the 2nd Period • 1. Highest occupied energy level – the energy level farthest from the nucleus that contains electrons • 2. Inner shell electrons – those electrons not in the highest energy level • 3. Octet of electrons – highest energy level is filled with 8 electrons

  22. E. Elements of the 3rd Period • 1. 1st 10 electrons of any atom in the period are the same as Neon • 2. Noble Gases – the Group VIII (or 18) elements • 3. Noble Gas Notation – also called abbreviatednotation • a.Use the symbol for the noble gas in brackets [Ne] that is previous to the element • b.Then continue with the electron configuration notation • 4. Noble Gas Configuration – the outermost energy level containing 8 electrons (Helium is the exception)

  23. F. Elements of the 4th Period • 1. 4s fills 1st • 2. 3d fills 2nd • 3. 4p fills 3rd • G. 5th Period Elements • 1. 5s fills 1st • 2. 4d fills 2nd • 3. 5p fills 3rd

  24. H. Elements of the 6th and 7th Periods • 1. 6s fills 1st • 2. 5dfills 2nd in La • 3. 4ffills 3rd • 4. The rest of 5d fills 4th • 5. 6p fills 5th • 6. 7s fills 6th • 7. 6d fills 7th in Ac • 8. 5ffills 8th • 9. Rest of 6d fills 9th

  25. 1s 2s 2px 2py 2pz 3py 3dxy 3dxz 3dyz 3dz2 3dx2y2 3s 3px 3pz 4px 4pz 4dxy 4dxz 4s 4py 4dyz 4dx2y2 4dz2 4fz3 4fxz2 4yz2 4fy(3x2-y2) 4fxyz 4fx(x2-3y2) 4fz(x2-y2)

  26. 1s2 2s2 2p6 3s2 3p5 • OR • [Ne]3s2 3p5

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