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Atomic Structure Part 1: The History of the Atomic Model Democritus • About 400 BC: • Democritus thought the atom existed, but had no proof • The word atom comes from “atomos” which means indivisible • He said that matter must be made of these indivisible particles Aristotle • • Aristotle taught that matter was continuous – that it was not made of particles, such as Democritus suggested. Aristotle was a Greek philosopher who’s ideas were the standard at that time. He was highly respected and thought to be a very wise source of information. John Dalton • • 1800’s John Dalton has proof that atoms exist and writes the first “Atomic Theory” 1. All elements are made up of atoms 2. Atoms of the same element are identical. Different elements have different kinds of atoms. 3. Atoms can form mixtures or chemically combine in whole number ratios. 4. Atoms cannot be destroyed or changed into other elements. Dalton’s model • He assumed the atoms were small, hard spheres with different sizes and different masses for different elements. J.J. Thomson • 1897: • J.J Thomson discovered electrons by using a cathode ray tube. • By using magnets, he bent the cathode rays and determined that the charge of the particles in the ray are negative. Animation - Cathode Ray Video Thomson’s model • Theory at the time said that the whole atom was neutral (had no charge), so if the electrons had a negative charge, there also must be a positively charged part of the atom, thus the idea of the proton. • Thomson suggested a model that had a positive base, with negative parts scattered throughout – “plum pudding” model. Ernest Rutherford • 1911 • Rutherford and his associates discover the nucleus. • Gold Foil Experiment – positively charged alpha particles are deflected by a dense positively charged nucleus in gold atoms. animation Gold Foil Experiment Gold Foil 2 Rutherford’s model • The nucleus must be very tiny. • The nucleus has a positive charge. • The nucleus contains most of the atom’s mass • The electrons are very lightweight and are found in the area surrounding the nucleus. • The atom is mostly empty space. One question remained: Why don’t the electrons collapse into the nucleus? Niels Bohr • 1913 • Suggested that the electrons are held in “fixed” positions around the nucleus. • Electrons have a fixed amount of energy – the energy is “quantized.” • Planetary model – electrons orbit the nucleus. • In Search of a Giant w/ Brian Cox Robert Millikan Watch the Experiment • 1916 • Oil Drop Experiment - Using charge ratios, Millikan found the mass of the electron to be 1/1837 of the mass of the atom. • He also stated that the electron had one unit of negative charge. Therefore suggesting the idea of a second particle (a proton) with one unit of positive charge. Eugen Goldstien • canal rays – opposite cathode rays • Goldstein's work with anode rays of H+ was apparently the first observation of the proton James Chadwick • 1932 • Confirmed the existence of a third particlethe neutron. • The neutron has no charge. • The neutron has the same mass as a proton. Irwin Schrödinger • 1926 • Quantum Mechanical Model of the AtomSchrödinger used mathematical equations to formulate the probability of finding electrons located at fixed energy levels around the nucleus. • Electrons are not contained to orbits, but are inside an “electron cloud.” http://en.wikipedia.org/wiki/Schr%C3%B6di nger_equation The Modern Quantum Model • An electron cloud shape is called an “orbital.” • The orbital shapes are probable locations of the electrons based on the amount of energy of the electron. http://www.orbitals.com/orb/orbtable.htm Atomic Structure Part 2: The Parts of the Atom Major Subatomic Particles • Proton – positively charged + found inside the nucleus • Electron – negatively charged – found outside of the nucleus • Neutron – neutral charge 0 found inside the nucleus Atomic Mass • The total mass of the atom is found inside the nucleus. This includes the protons and the neutrons. Atomic Mass = Number of Protons + the number of neutrons Atomic Number • The number of protons in the nucleus of an atom. • Also equal to the number of electrons in a neutral atom. Atomic number = number of protons = number of electrons Isotopes • Atoms of the same element which have a different number of nuetrons, and therefore a different atomic mass. • Example: Hydrogen has three isotopes. • http://education-portal.com/academy/lesson/isotopes-and-calculation-of-aweighted-average.html Representation • What do these mean? (Examples of isotopes) Average Atomic Mass • Masses recorded on the periodic table of elements are the “weighted average” of the various isotope masses that exist for that element. • Therefore the element’s mass is generally not a whole number. Calculating average atomic mass Quantum Mechanical model http://www.youtube.com/watch?v=45KGS1 Ro-sc • http://www.youtube.com/watch?v=B7pACq_xWyw http://www.youtube.com/watch?v=ARuyAX mjjVI&list=PL69CA6EEFAEF15D1D Chapter 5 Quantum Mechanical Model of the Atom • This model suggests that electrons are located at definite energy levels inside an electron cloud • According to this theory, the electrons do not travel in paths around the nucleus, but can be found anywhere within a defined region of space around the nucleus. We call this area the electron cloud. Electrons are “quantized” • The lowest level is called ground state. Here, the electrons have enough energy to maintain their position and not collapse into the nucleus. • Electrons can move from one level to another by gaining or losing energy or “quanta” (tiny bundles of energy) • The more energy the electron has, the farther away from the nucleus it will be found. Schrödinger's Probability • The exact location of an electron is unknown. • Mathematically, we can predict the most probable areas where the electrons can be found. https://images=_&imgdii=_&images Electron Density Diagram The real picture of an orbital/ The famous equation Orbitals • The most probable areas where the electrons can be found are called orbitals. s orbital p orbital d orbital Quantum Numbers • Quantum numbers are assigned to describe the location and energy of electrons. • The lowest energy level of the electrons is called ground state. Quantum Numbers • Principal quantum numbers – these designate the major energy levels, and thus the distance from the nucleus. • n = 1 - 1st and lowest energy level (closest to the nucleus) n=3 n=2 n = 2 - next closest n=1 n = 3, n = 4, n = 5, etc…. • Orbital quantum numbers - describe the sublevels of energy and show the shape of the orbitals. • s – spherical (lowest sublevel) • p – dumbbell shaped (three kinds) py px pz • d – clover shaped ( five kinds) dxy dxz dx2 – y2 dxy dz • f - complicated patterns (seven kinds) http://www.uky.edu/~holler/html/orbitals_2.html http://www.youtube.com/watch?v=A6DiVsp oZ1E Electron Configurations • Electron Configuration – the way that electrons are arranged in their most stable pattern around the nucleus. • Three rules apply: Aufbau principle Pauli Exclusion Principle Hund’s Rule Aufbau Principle • Aufbau principle – electrons enter the orbitals of lowest energy first. • Follow this pattern: p 133 1s 2s 2p 3s 3p 3d 4s 4p 5s 4d 5p 4f 5d 5f Pauli Exclusion Principle • Pauli exclusion principle – each orbital can hold a maximum of two electrons. These electrons must have opposite spins. Electrons are represented by up and down arrows.) 1s Hund’s Rule • Hund’s Rule – When electrons enter orbitals of equal energy, one electron enters each orbital until all the orbitals contain one electron with spins in parallel. Then the second electrons are entered with opposite spins. px py pz Writing Orbital Notations Another Example • Cr has 24 electrons Writing Electron Configurations Shortcut Method