Ch1a - Chapter 1. The Electronic Structure of the Atom. The...

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Unformatted text preview: Chapter 1. The Electronic Structure of the Atom. The Schrdinger Wave Equation Shapes of Atomic Orbitals The Polyelectronic Atom 1 Ion Electron Configurations Magnetic Properties of Atoms The Structure of Hydrogen and Hydrogenic Ions As chemists we need to know about what the electrons are doing inside atoms. This helps us to understand a number of very basic atomic properties connected to ionic bonding , such as: Ionization enthalpy (ionization energy or potential), Electron attachment enthalpy (electron affinity) Ionic Radius And various properties of atoms when they are found in covalent compounds or metals , such as: 2 Other types of radii (atomic, covalent and van der Waals) Electronegativity And, of course, we need to understand the formation of covalent bonds between atoms. To make this a little easier, we start with the simplest possible atom, hydrogen, which has only one electron (and other exotic one-electron species, such as He + , Li 2+ , Be 3+ , which also have only one electron, but increasing numbers of protons in their nuclei, the hydrogenic ions .) White light, from the sun or an incandescent filament is split in to a rainbow - a continuous spectrum or continuum Atomic Line Spectra 3 Light coming from a gas discharge tube, for example one where a current is passed through hydrogen gas, produces a line spectrum Bohr Theory The first attempt to devise a model of the atom to explain the line spectra was due to Niels Bohr (1885 1962). He based his theory on the Rutherford planetary model of the atom, and considered only circular orbits. To counter a principle of classical physics, that electrons in a circular path would emit radiation energy and spiral into the nucleus, he said that only certain orbits were allowed; those with radius and energy given by: r = E n = - here n = 1,2,3, = Plancks constant, = the electons mass, e = the electrons charge, hc n 2 n 2 h 2 4 2 me 2 Z 4 where n = 1,2,3, , h = Plancks constant, m = the electons mass, e = the electrons charge, c = speed of light, and = the Rydberg constant (a combination of the others). The energy is always negative, and as required by Coulombs law, becomes more negative the closer the electron is to the nucleus. The value of E n is most negative for n = 1, and becomes less negative as n increases. It goes to zero when n = . The situation where n = 1 is called the ground state , and others are called excited states . An atom will want to be in its ground state, and will generally lose energy as electromagnetic radiation in order to get there. The number n is called the principal quantum number . Bohr Theory (cont.) Bohr theory predicts line spectra for hydrogen: If an electron moves from an orbit n initial to an orbit n final where n initial < n final , energy is absorbed . A photon of the correct energy can make this happen, or a violent collision with an electron or another atom....
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Ch1a - Chapter 1. The Electronic Structure of the Atom. The...

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