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Through the work of • • • 1. 2. Einstein, Planck, DeBroglie, Bohr, Schrodinger, Heisenberg we know Depending on the experiment, light has particle like behavior (an example…photoelectric effect) Particles (the electron) has wave like behavior (example: the emission spectrum) Our goals in the first part of this chapter is To describe light and particles in terms of energy and wavelength, apply DeBroglie’s relationship and the photoelectric effect To describe likely positions of where the electron is located to understand chemical behavior. • Light is a form of electromagnetic radiation • Electromagnetic radiation is energy (radiant) that as it travels consists of electric and magnetic fields • Travel in waves that can be described by their wavelength (lamda –l) and frequency (n- nu) • In a vacuum, travel at the speed of light, c = 3 x 108 m/s ln= c White light is made up of light waves of different wavelengths Disperses light into spectrum of colors. There are many lines that blend together producing a continuous spectrum ( 400-700 nm). Hydrogen Neils Bohr Bohr theorized that the lines in the hydrogen spectrum corresponded to certain quantum of energy emitted or absorbed when the electron moved from one “orbit” to another. • Proposed electrons travel in successively larger orbits and when an electron jumps from an outer orbit to an inner one, it emits light. • Proposed the planetary model - suggested the electron orbits the nucleus as planets orbit the sun Spectroscopy • Method of studying substance that are exposed to exciting energy. Can be used to identify substances, since each element emits a unique collection of lines. • The presence of spectral lines of specific frequencies suggests that the energy of the electron in an atom is restricted to a series of discrete values Max Planck • Proposed energy is quantized • Energy can only be lost or gained in packets • Each packet is called a quantum Energy of quantum of electromagnetic radiation is proportional to frequency, DE = hu Indium compounds give a blueviolet flame test. The atomic emission responsible for this blue-violet color has a wavelength of 451 nm. Calculate the energy of a single photon of this wavelength. Albert Einstein • Used Planck’s equation to explain photoelectric effect Light consists of quanta of energy that behave like tiny particles of light(photon). Electrons can absorb energy from photons, but they follow an "all or nothing" principle. All of the energy from one photon must be absorbed and used to liberate one electron. 1/2mv2 = hn - f 1/2mv2: Kinetic energy of electron ejected hn: Energy of incident radiation f: Work function, minimum energy needed to eject electron • The photoelectric work function for Mg is 5.90 x 10-19J. Calculate the minimum frequency of light required to eject electrons from Mg. • Light of wavelength 285 nm shines on a piece of Mg metal. What is the speed of the ejected electron? Albert Einstein • Used Planck’s equation to explain photoelectric effect Light consists of quanta of energy that behave like tiny particles of light(photon). Electrons can absorb energy from photons, but they follow an "all or nothing" principle. All of the energy from one photon must be absorbed and used to liberate one electron. •In related development, proposed special theory of relativity 2 E = mc • What were the contributions of DeBroglie? Matter Waves, l = h / mv • Heisenberg? • Schrodinger? Uncertainty Principle – location and momentum of particle are complimentary; can’t both be known simultaneously with precision; can’t specify precise location of particle if it behaves like a wave Developed an equation that describes the wavelike properties of matter, we use the wave function to express the probability of finding the electron in a specific region in space (orbital), Quantum number used to describe the orbital and the electron 1. Probability of finding electron at different points is calculated • Some points will have higher probability than others 2. If connect all points of high probability, three dimensional shapes are formed 3. The most probable place to find the electron will be some place in that shape Quantum Numbers • Set of four numbers that completely specify each electron 1. Prinicipal quantum number, n, corresponds to the energy level 2. Second quantum number, l, corresponds to the sublevel within the energy level. It provides information about the shape of the electron cloud 3. Third quantum number, m, is the orbital quantum number and describes the orientation in space of orbital 4. Fourth quantum number, s, is the spin quantum number and describes rotation of an electron on its axis. Values are ½ or - ½ Sublevel s p d f Orientations (Orbitals) Region around nucleus where electrons are likely to be found. Each orbital holds two electrons that have opposite spin 6d __ __ __ __ __ 5f __ __ __ __ __ __ __ 7s __ 6p __ __ __ 5d __ __ __ __ __ 4f __ __ __ __ __ __ __ 6s __ 5p __ __ __ 4d __ __ __ __ __ 5s __ 4p __ __ __ 3d __ __ __ __ __ 4s __ 3p __ __ __ 3s __ 2p __ __ __ 2s __ 1s __ Atomic Orbitals • Region around nucleus where electrons are likely to be found Energy Level Sublevels 1 s 2 s,p 3 s, p, d 4 s, p, d, f • Each orbital holds two electrons that have opposite spin Follow 3 rules to configure the electrons 1. Aufbau Principle - electrons fill orbitals starting at the lowest available (possible) energy states before filling higher states 2. Pauli Exclusion Principle - two electrons cannot share the same set of quantum numbers within the same system. Therefore, there is room for only two electrons in each orbital and the electrons have opposite spin. 3. Hund’s Rule – in equal energy orbitals, arrange the electrons to achieve the maximum number of unpaired electrons. • What is a trend? • Atomic and Ionic Radius •Ionization Energy •Electron Affinity •What do these terms mean in relation to the atom and its subatomic particles… effective nuclear charge, shielding, penetration, and electron repulsions?