Electron Configuration & Orbitals 2 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 Continuous spectrum (results when white light is passed through a prism) contains all the wavelengths of visible light Line spectrum each line corresponds to a discrete wavelength: Hydrogen emission spectrum Copyright Cengage Learning. All rights reserved 2 Refraction of White Light The Line Spectrum of Hydrogen To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE Copyright Cengage Learning. All rights reserved 3 To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE Copyright Cengage Learning. All rights reserved 4 Significance Only certain energies are allowed for the electron in the hydrogen atom. Energy of the electron in the hydrogen atom is quantized. CONCEPT CHECK! Why is it significant that the color emitted from the hydrogen emission spectrum is not white? How does the emission spectrum support the idea of quantized energy levels? Copyright Cengage Learning. All rights reserved 5 Copyright Cengage Learning. All rights reserved 6 1
Electron in a hydrogen atom moves around the nucleus only in certain allowed circular orbits. - Bohr Bohr s model gave hydrogen atom energy levels consistent with the hydrogen emission spectrum. Ground state lowest possible energy state (n= 1) Electronic Transitions in the Bohr Model for the Hydrogen Atom a) An Energy-Level Diagram for Electronic Transitions Copyright Cengage Learning. All rights reserved 7 Copyright Cengage Learning. All rights reserved 8 Electronic Transitions in the Bohr Model for the Hydrogen Atom b) An Orbit-Transition Diagram, Which Accounts for the Experimental Spectrum For a single electron transition from one energy level to another: 18 1 1 E = 2.178 10 J 2 2 nfinal ninitial ΔE= change in energy of the atom (energy of the emitted photon) n final = integer; final distance from the nucleus n initial = integer; initial distance from the nucleus Copyright Cengage Learning. All rights reserved 9 Copyright Cengage Learning. All rights reserved 10 The model correctly fits the quantized energy levels of the hydrogen atom and postulates only certain allowed circular orbits for the electron. As the electron becomes more tightly bound, its energy becomes more negative relative to the zero-energy reference state (free electron). As the electron is brought closer to the nucleus, energy is released from the system. Bohr s model is incorrect. This model only works for hydrogen. Electrons do not move around the nucleus in circular orbits. Copyright Cengage Learning. All rights reserved 11 Copyright Cengage Learning. All rights reserved 12 2
Bohr s Model Electrons don t exist in orbits they exist in orbitals! Essentially the model went from Bohr s atomic model of electrons in orbits with specific levels of energy was only partly correct. The idea of electron shells is not quite sufficient to explain all the atomic spectral lines. We need to add the concept of orbitals and quantum numbers for elements with multiple electrons. Bohr Model to Quantum Model http://www.th.physik.uni-frankfurt.de/~jr/gif/phys/nbohr.jpg Orbitals Orbitals Orbitals orbital means small orbit We are interested in: their energy to explain spectral lines their shape to explain configurations of molecules An orbital is not an orbit! It is a probability map ( probability density distribution ) An An orbital orbital is not is a an 3-dimensional orbit! It is a probability region of space map where probability the electron density is likely distribution to be found. n l m 1 m s Schrodinger proposed 4 Quantum Numbers to describe the location of an electron the Principal Quantum Number (describes the SHELL) the Secondary Quantum Number (describes the SUBSHELL) the Magnetic Quantum Number (describes the ORBITAL) the Spin Quantum Number (describes the SPIN) Where are the electrons? Like Bohr said, electrons are found in shells ( energy levels ) but all shells contain subshells all subshells contain orbitals and every orbital contains two electrons. that s four levels of giving an electron s address! 3
Principal Quantum Number, n You already know this as the shell. Shells are identified as K, L, M, N and their given values are n = 1, 2, 3, 4 Describes the size of the orbital (or how far away the electron is from the nucleus). As n increases, so does energy. Secondary Quantum Number, l Each shell has n subshells Subshells are given letter codes: s, p, d, f... : first shell has one orbital type, 2p: second shell has two orbital types 3s, 3p, 3d: third three orbital types 4s, 4p, 4d, 4f: fourth four orbital types Secondary Quantum Number, l Within a given shell, the subshells are in order of increasing energy: Relative Sizes and Principal Quantum Number describes the size. 4s < 4p < 4d < 4f (increasing energy) Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 334 Magnetic Quantum Number, m l Splits the subshells into individual orbitals. Describes orientation in space, thus gives 3D information. Each orbital can hold a maximum of: 2 electrons. Maximum of Electrons In Each Sublevel Maximum Number of Electrons In Each Sublevel Maximum Number Sublevel Number of Orbitals of Electrons s 1 2 p 3 6 d 5 10 f 7 14 LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World, 1996, page 146 4
Shapes of s, p, and d-orbitals p-orbitals Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 335 d-orbitals Spin Magnetic Number, m s Indicates the orientation of the two electrons in each orbital. Values are +1/2 or 1/2 Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 336 Spin Quantum Number, m s North South Principal Energy Levels 1 and 2 - - S The electron behaves as if it were spinning about an axis through its center. This electron spin generates a magnetic field, the direction of which depends on the direction of the spin. N Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 208 5
s, p, and d-orbitals Electron Configurations A s orbitals: Hold 2 electrons (outer orbitals of Groups 1 and 2) B p orbitals: Each of 3 pairs of lobes holds 2 electrons = 6 electrons (outer orbitals of Groups 13 to 18) C d orbitals: Each of 5 sets of lobes holds 2 electrons = 10 electrons (found in elements with atomic no. of 21 and higher) Kelter, Carr, Scott,, Chemistry: A World of Choices 1999, page 82 How do we fill orbitals? Atom is most stable when electrons have the lowest possible energy. That is when they occupy the lowest possible energy orbitals available. Filling Rules for Electron Orbitals Aufbau Principle: Electrons are added one at a time to the lowest energy orbitals available until all the electrons of the atom have been accounted for. Pauli Exclusion Principle: An orbital can hold a maximum of two electrons. To occupy the same orbital, two electrons must spin in opposite directions. Hund s Rule: Electrons occupy equal-energy orbitals so that a maximum number of unpaired electrons results. *Aufbau is German for building up Energy Level Diagram of a Many-Electron Atom Arbitrary Energy Scale 6s 6p 5d 4f 5s 5p 4d 4s 4p 3d 3s 3p 32 18 18 8 Sublevels Energy n = 4 n = 3 4f 4d 4p 3d 4s 3p 3s 2p 8 n = 2 2p 2 n = 1 NUCLEUS O Connor, Davis, MacNab, McClellan, CHEMISTRY Experiments and Principles 1982, page 177 6
Sublevels 2 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 Energy n = 4 n = 3 n = 2 n = 1 4f 4d 4p 3d 4s 3p 3s 2p Writing Electron Configurations: 2p 4 Energy Level n Subshell (s, p, d or f) Number of electrons in that subshell 2 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 etc. Abbreviated Notation Chlorine Longhand is 2 2 2p 6 3s 2 3p 5 You can abbreviate the first 10 electrons with [Ne] (replaces 2 2 2p 6 ) The next energy level after Neon is 3 So you start at level 3 and finish by adding 7 more electrons to bring the total to 17 [Ne] 3s 2 3p 5 http://www.steve.gb.com/images/science/orbital_filling.png Electron Configurations and the Periodic Table 7