1 Chapter 6 Atomic Structure 6.1 The Wave Nature of Light _ ________________ - progressive, repeating disturbance that spreads through a medium from a point of origin to more distant points. - there is little movement of the material in the direction of the disturbance ____________________________ WAVES - oscillations in electric and magnetic fields that are propagated over distances __________________ - distance between any two identical points in consecutive cycles ____________________ - height of a wave _________________ - # of wavelengths that pass through a point in a unit of time "number of cycles of the wave" (wavelength)(frequency) = (speed of the wave) c =____________________ m/s wavelength is ___________________ ___ to frequency2 amplitude is intensity – it does not affect wavelength or frequency EXAMPLE: Calculate the frequency of a radio wave that has a wavelength of 4.00 meters. Calculate the wavelength of light that has the frequency of 1 x 1016 hertz. Calculate the frequency of blue light that has a wavelength of 450 nm. Calculate the wavelength (in nm) of an x-ray with a frequency of 5.25 x 1018 s-1. The Electromagnetic Spectrum gamma rays x rays ultraviolet visible infrared microwave TV-radio3 6.2 Photons: Energy by the Quantum • 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 (or chunks) that are called quanta. * atoms can absorb or emit electromagnetic energy only in _____________________ * smallest amount of energy is a _____________________ * energy of quantum (photon) E = ________ * energy can be absorbed or emitted only as a quantum or as exact multiples of a quantum. * variations in energy are discontinuous • Einstein used this assumption to explain the photoelectric effect. 1. electrons are energized by photons 2. electrons escape from atom (using E from photon) 3. excess photon E = KE of the ejected electron He concluded that energy is proportional to frequency: E = h (where h = Planck’s constant = 6.626  10 J-s) Therefore, if one knows the wavelength of light, one can calculate the energy in one photon, or packet, of that light: c = 4 Practice Problems: 1. Calculate the E (J) of a photon of visible light that has a frequency of 2.50 x 1014 hertz. 2. Calculate the E, in joules per photon, of blue light at 475 nm. 3. What is the wavelength of a photon with an energy of 4.95 X 10-19 J/photon? 4. A laser produces green light of wavelength 535.0 nm. (a) Calculate the frequency of this wavelength. (b) Calculate the energy, in kilojoules, of ONE MOLE of photons of this wavelength.5 6.3 Line Spectra and the Bohr Model • Most radiation sources emit light of many different wavelengths, that when separated produce a spectrum. • We do not observe a continuous spectrum for atoms and molecules, as one gets from a white light source. • Only a line spectrum of discrete wavelengths is observed. prism separates visible light in to a ______________________ white light - continuous spectrum light from a hydrogen lamp - _____________ spectrum __________________spectrum - refers to an analysis of the light emitted when an element is energized In order to explain why an electron does not continuously lose energy and end up spiraling into the nucleus of the H atom, Niels Bohr based his atomic model on the following three postulates: 1. Electrons in an atom can only occupy certain orbits (corresponding to certain energies). 2. Electrons in permitted orbits have specific, “allowed” energies; these energies will not be radiated from the atom. 3. Energy is only absorbed or emitted in such a way as to move an electron from one “allowed” energy state6 to another; the energy is defined by E = h The energy absorbed or emitted from the process of electron promotion or demotion can be calculated by the equation E  RH ( 1/nf2 – 1/ni2) where RH is the Rydberg constant, 2.18  10 J, and ni and nf are the initial and final energy levels of the electron. Example: Calculate the energy change that occurs when an electron falls from the n = 2 to the n = 1 level in an H atom. What is the frequency and wavelength of the emitted energy? 6.4 The Wave Behavior of Matter • Louis de Broglie proposed that if light can have material properties, matter should exhibit wave properties. • He demonstrated that the relationship between mass and wavelength was7 DeBroglie proposed that a particle with a mass(m) moving at a speed (v) will have a wave nature consistent with wavelength = h/mv 1000kg car moving at 100km/h has a wavelength of 2.39 x 10-38 m – undetectable, but subatomic particles, are readily observable Heisenberg’s Uncertainty Principle • Heisenberg showed that the more precisely the momentum of a particle is known, the less precisely is its position known: • In many cases, our uncertainty of the whereabouts of an electron is greater than the size of the atom itself! Quantum Mechanics and Atomic Orbitals • Erwin Schrödinger (1962) developed a mathematical treatment into which both the wave and particle nature of matter could be incorporated. • It is known as quantum mechanics. • The wave equation is designated with a lower case Greek psi (). • The square of the wave equation, 2, gives a probability density map of where an electron has a certain statistical likelihood of being at any given instant in time. (electron density distribution, the probability of finding the electron in a given area) Quantum Numbers • Solving the wave equation gives a set of wave functions, or orbitals, and their corresponding energies.8 • Each orbital describes a spatial distribution of electron density. • An orbital is described by a set of three quantum numbers and they give the “address” of the electron. 1. Principle Quantum Number (n) *The principal quantum number, n, describes the energy level on which the orbital resides. *The values of n     (as n increases, the orbital becomes larger, and the electrons are higher in energy) (It is equal to the Period Number) 2. Angular Momentum Quantum Number (l) *This quantum number defines the