![Electronic
Structure
of Atoms
Some Anomalies
For instance, the
electron
configuration for
copper is
[Ar] 4s1 3d5
rather than the
expected
[Ar] 4s2 3d4.](https://image.slidesharecdn.com/chapter06au-141020130046-conversion-gate02/85/Ch-6-Electronic-Structure-of-Atoms-40-320.jpg)
Ch. 6 Electronic Structure of Atoms
- 1. Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten Electronic Structure of Atoms Chemistry, The Central Science, 10th edition Chapter 6 Electronic Structure of Atoms John D. Bookstaver St. Charles Community College St. Peters, MO ã 2006, Prentice Hall, Inc.
- 2. Electronic Structure of Atoms Waves • To understand the electronic structure of atoms, one must understand the nature of electromagnetic radiation. • The distance between corresponding points on adjacent waves is the wavelength (l).
- 3. Electronic Structure of Atoms Waves • The number of waves passing a given point per unit of time is the frequency (n). • For waves traveling at the same velocity, the longer the wavelength, the smaller the frequency.
- 4. Electronic Structure of Atoms Electromagnetic Radiation • All electromagnetic radiation travels at the same velocity: the speed of light (c), 3.00 ´ 108 m/s. • Therefore, c = ln
- 5. Electronic Structure of Atoms The Nature of Energy • The wave nature of light does not explain how an object can glow when its temperature increases. • Max Planck explained it by assuming that energy comes in packets called quanta.
- 6. Electronic Structure of Atoms The Nature of Energy • Einstein used this assumption to explain the photoelectric effect. • He concluded that energy is proportional to frequency: E = hn where h is Planck’s constant, 6.63 ´ 10−34 J-s.
- 7. Electronic Structure of Atoms The Nature of Energy • Therefore, if one knows the wavelength of light, one can calculate the energy in one photon, or packet, of that light: c = ln E = hn
- 8. Electronic Structure of Atoms The Nature of Energy Another mystery involved the emission spectra observed from energy emitted by atoms and molecules.
- 9. Electronic Structure of Atoms The Nature of Energy • One does not observe a continuous spectrum, as one gets from a white light source. • Only a line spectrum of discrete wavelengths is observed.
- 10. Electronic Structure of Atoms The Nature of Energy • Niels Bohr adopted Planck’s assumption and explained these phenomena in this way: 1. Electrons in an atom can only occupy certain orbits (corresponding to certain energies).
- 11. Electronic Structure of Atoms The Nature of Energy • Niels Bohr adopted Planck’s assumption and explained these phenomena in this way: 2. Electrons in permitted orbits have specific, “allowed” energies; these energies will not be radiated from the atom.
- 12. Electronic Structure of Atoms The Nature of Energy • Niels Bohr adopted Planck’s assumption and explained these phenomena in this way: 3. Energy is only absorbed or emitted in such a way as to move an electron from one “allowed” energy state to another; the energy is defined by E = hn
- 13. Electronic Structure of Atoms The Nature of Energy The energy absorbed or emitted from the process of electron promotion or demotion can be calculated by the equation: DE = −RH ( 1 ) nf 2 1 ni 2 - where RH is the Rydberg constant, 2.18 ´ 10−18 J, and ni and nf are the initial and final energy levels of the electron.
- 14. The Wave Nature of Matter • Louis de Broglie posited that if light can have material properties, matter should exhibit wave properties. • He demonstrated that the relationship between mass and wavelength was Electronic Structure of Atoms l = h mv
- 15. Electronic Structure of Atoms The Uncertainty Principle • Heisenberg showed that the more precisely the momentum of a particle is known, the less precisely is its position known: (Dx) (Dmv) ³ h 4p • In many cases, our uncertainty of the whereabouts of an electron is greater than the size of the atom itself!
- 16. Electronic Structure of Atoms Quantum Mechanics • Erwin Schrödinger developed a mathematical treatment into which both the wave and particle nature of matter could be incorporated. • It is known as quantum mechanics.
- 17. Electronic Structure of Atoms Quantum Mechanics • The wave equation is designated with a lower case Greek psi (y). • The square of the wave equation, y2, gives a probability density map of where an electron has a certain statistical likelihood of being at any given instant in time.
- 18. Electronic Structure of Atoms Quantum Numbers • Solving the wave equation gives a set of wave functions, or orbitals, and their corresponding energies. • Each orbital describes a spatial distribution of electron density. • An orbital is described by a set of three quantum numbers.
- 19. Principal Quantum Number, n • The principal quantum number, n, describes the energy level on which the orbital resides. • The values of n are integers ≥ 0. Electronic Structure of Atoms
- 20. Azimuthal Quantum Number, l • This quantum number defines the shape of the orbital. • Allowed values of l are integers ranging from 0 to n − 1. • We use letter designations to communicate the different values of l and, therefore, the shapes and types of orbitals. Electronic Structure of Atoms
- 21. Azimuthal Quantum Number, l Electronic Structure of Atoms Value of l 0 1 2 3 Type of orbital s p d f
- 22. Magnetic Quantum Number, ml • Describes the three-dimensional orientation of the orbital. • Values are integers ranging from -l to l: Electronic Structure of Atoms −l ≤ ml ≤ l. • Therefore, on any given energy level, there can be up to 1 s orbital, 3 p orbitals, 5 d orbitals, 7 f orbitals, etc.
- 23. Magnetic Quantum Number, ml • Orbitals with the same value of n form a shell. • Different orbital types within a shell are subshells. Electronic Structure of Atoms
- 24. Electronic Structure of Atoms s Orbitals • Value of l = 0. • Spherical in shape. • Radius of sphere increases with increasing value of n.
- 25. Electronic Structure of Atoms s Orbitals Observing a graph of probabilities of finding an electron versus distance from the nucleus, we see that s orbitals possess n−1 nodes, or regions where there is 0 probability of finding an electron.
- 26. Electronic Structure of Atoms p Orbitals • Value of l = 1. • Have two lobes with a node between them.
- 27. Electronic Structure of Atoms d Orbitals • Value of l is 2. • Four of the five orbitals have 4 lobes; the other resembles a p orbital with a doughnut around the center.
- 28. Electronic Structure of Atoms Energies of Orbitals • For a one-electron hydrogen atom, orbitals on the same energy level have the same energy. • That is, they are degenerate.
- 29. Electronic Structure of Atoms Energies of Orbitals • As the number of electrons increases, though, so does the repulsion between them. • Therefore, in many-electron atoms, orbitals on the same energy level are no longer degenerate.
- 30. Electronic Structure of Atoms Spin Quantum Number, ms • In the 1920s, it was discovered that two electrons in the same orbital do not have exactly the same energy. • The “spin” of an electron describes its magnetic field, which affects its energy.
- 31. Electronic Structure of Atoms Spin Quantum Number, ms • This led to a fourth quantum number, the spin quantum number, ms. • The spin quantum number has only 2 allowed values: +1/2 and −1/2.
- 32. Electronic Structure of Atoms Pauli Exclusion Principle • No two electrons in the same atom can have exactly the same energy. • For example, no two electrons in the same atom can have identical sets of quantum numbers.
- 33. Electronic Structure of Atoms Electron Configurations • Distribution of all electrons in an atom • Consist of Number denoting the energy level
- 34. Electronic Structure of Atoms Electron Configurations • Distribution of all electrons in an atom • Consist of Number denoting the energy level Letter denoting the type of orbital
- 35. Electronic Structure of Atoms Electron Configurations • Distribution of all electrons in an atom. • Consist of Number denoting the energy level. Letter denoting the type of orbital. Superscript denoting the number of electrons in those orbitals.
- 36. Electronic Structure of Atoms Orbital Diagrams • Each box represents one orbital. • Half-arrows represent the electrons. • The direction of the arrow represents the spin of the electron.
- 37. Electronic Structure of Atoms Hund’s Rule “For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized.”
- 38. Electronic Structure of Atoms Periodic Table • We fill orbitals in increasing order of energy. • Different blocks on the periodic table, then correspond to different types of orbitals.
- 39. Electronic Structure of Atoms Some Anomalies Some irregularities occur when there are enough electrons to half-fill s and d orbitals on a given row.
- 40. Electronic Structure of Atoms Some Anomalies For instance, the electron configuration for copper is [Ar] 4s1 3d5 rather than the expected [Ar] 4s2 3d4.
- 41. Electronic Structure of Atoms Some Anomalies • This occurs because the 4s and 3d orbitals are very close in energy. • These anomalies occur in f-block atoms, as well.