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I. Electrons and Period Trends Summary Properties of Light A. Wave Description of Light 1. Electromagnetic radiation is a form of energy that exhibits wavelike behavior as it travels through space. 2. Electromagnetic spectrum is all forms of electromagnetic radiation put together. http://lot.astro.utoronto.ca/images/spectrum.png II. III. 3. Wavelength the distance between corresponding points on adjacent waves. Photoelectric Effect A. The emission of electrons from a metal when light shines on the metal. B. Quantum: the minimum quantity of energy that can be lost or gained by an atom. C. Photon: is a particle of electromagnetic radiation having zero mass and carrying a quantum energy. Bright Line Emission A. Ground state: the lowest energy state of a atom. B. Excited state: a state in which the atom has a higher potential energy than it has in its ground state. C. Line‐emission spectrum: Bands of light that correspond to an elements identity. D. Continuous spectrum: the emission of a continuous range of frequencies of electromagnetic radiation. IV. Bohr Model of the Atom A. Niels Bohr proposed a model of the hydrogen atom based on its line‐emission spectrum. (http://www.splung.com/nuclear/images/atoms10.gif) V. Quantum Model of the Atom A. Quantum Theory (Heisenberg and Schrodinger) describes mathematically the wave properties of electrons and other very small particles. (http://farm1.static.flickr.com/59/169563400_cf75d3ccf9.jpg) B. Quantum Number: specifies the properties of atomic orbitals and the properties of electrons in orbitals. C. Principle Quantum Number (n) the main energy level occupied by the electrons (1,2,3,4,5,6,7 ‐‐one for each period on the periodic table) http://image.tutorvista.com/content/atomic-structure/relative-energy-levels-hydrogen-and-hydrogen-like-atoms.gif VI. D. Magnetic Quantum Number: the orientation of the orbital around the nucleus (s, px, py, pz, dx2-y2, dxy, dyz, dxz, dz, f etc…) E. Spin Quantum Number: + or – the two fundamental spin states of an electron in orbit. F. General rules that hold true: IF n = the number of the energy level (1-8) THEN n = the number of orbital types in that energy level n2 = the total number of orbitals in that energy level 2n2 = the total number of electrons that can exist in the energy level. Electron Configurations A. Aufbau Principle: an electron occupies the lowest –energy orbital that can receive it. B. Pauli Exclusion Principle: no two electrons in the orbital can have the same spin (they must have opposite spins). C. Hund’s rule: when filling an energy level all orbitals MUST get one electron before any orbital can have two electrons. D. Electron Dot Notation: Dots are used to denote the outermost shell of electrons. The dots are placed according to which orbital electrons they represent. 3 4 7 X 2 1 5 6 8 E. Electron Configuration: 1. Lists the energy level, the orbital type and the number of electrons. EXAMPLE: 3p3 this means that in the p orbitals of the third energy level there are three electrons. 2. When listing a complete configuration notation, group the energy levels together. Even though the order of filling for an atom may be: 1s2 2s2 2p6 3s2 3p6 4s2 3d6 We group our energy levels and show it as follows: 1s2 2s2 2p6 3s2 3p6 3d6 4s2 The sum of the superscripts will always be the number of electrons in the atom (atomic number). F. Orbital Notation: This is very similar to configuration. We label the orbitals by 1s, 2s, etc, but we show the electrons by slashes or arrows. VII. History of the Periodic Table A. Dmitri Mendeleev: 1. Organized elements based on properties. 2. Left empty spaces for “new” elements B. Henry Moseley: 1. Arranged elements by their increasing atomic charge (number of protons) C. Periodic Law: The physical and chemical properties of the elements are periodic functions of their atomic numbers. 1. Elements with the same properties fall in the same column. VIII. Groups/Families (Columns) A. Noble gases: chemically unreactive, full outer shell (octet rule) B. Alkali metals: Group 1 C. Alkaline Earth Metals: Group 2 D. Transition Metals (Groups 3‐12) E. Halogens: (Group 17) IX. Periods (Rows) A. The Lanthanides (58‐71)—fill the f block B. The Actinides (90‐103)—fill f block X. Trends A. Atomic Radius: as one half the distance between the nuclei of identical atoms that are bonded together. B. Ionization Energy: the energy required to remove one electron from a neutral atom of an element. 1. Second Ionization Energy: the addition energy need to remove a second electron. C. Electron Affinity: the energy CHANGE that occurs when an electron is acquired by a neutral atom. D. Ionic Radii: The size of the charged atom 1. Cation: + charge 2. Anion: ‐ charge E. Valence Electrons: the electrons available to be lost, gained or shared in the formation of chemical bonds. F. Electronegativity: the measure of the ability of an atom/element to attract/steal electrons from another element/atom. http://www.meta‐synthesis.com/webbook/24_complexity/periodic_trends.jpg http://www.chemistry.wustl.edu/~courses/genchem/Tutorials/LED/images/periodic.jpg