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HC History of the atom.notebook October 21, 2011 History & Development of Atomic Theory Aug 2810:34 AM History of Atomic Theory Democritus ~ 460 to 370 BC Experiment: None. He had beliefs that were disregarded. Beliefs: • Named the atom "Atomos" meaning indivisible. • Matter is composed of atoms too small to be seen. • Believed atoms could not be created, destroyed or further divided. • Atoms are solid and homogeneous with empty space between them. • Different atoms have different sizes and shapes which determine the properties of matter. Democritus Atomic Model: a sphere Aug 201:50 PM 1 HC History of the atom.notebook October 21, 2011 History of Atomic Theory The Alchemists ~ 300 BC to 300 AD Experiment: Pseudoscience influenced by Aristotle, concerned with... Using chemistry to change cheap metal to gold. Finding an eternal life elixir. Aristotle Beliefs: • All matter was made up of a combination of four elements; fire, wind, earth, water Aug 201:35 PM History of Atomic Theory Dalton ~ early 1803 English School Teacher Developed the first "official" atomic theory. Experiment: Studied chemical reactions, making observations and measurements. His experimental results demonstrated the following... John Dalton The Law of Conservation of Matter total mass of reactants = total mass of products Ex. 2 H2 + O2 --> 2 H2O mass of hydrogen + mass of oxygen reacted = mass of water produced The Law of Definite Proportions regardless of starting amounts, a compound will always form with the same mass ratio of elements Ex. The formula for table sugar will always be C12H22O11, whether you start with 5.0 grams or 20.0 grams of carbon, hydrogen & oxygen. The Law of Multiple Proportions different compounds can form from the same elements but with different mass ratios Ex. The formula for water is H2O (2:1 ratio of hydrogen to oxygen) The formula for peroxide is H2O2 (1:1 ratio of hydrogen to oxygen) Aug 201:58 PM 2 HC History of the atom.notebook October 21, 2011 Five Principles-Dalton's Atomic Theory 1. All matter is made of indestructible atoms. 2. Atoms of the same element are identical in their physical and chemical properties. 3. Atoms of different elements have different physical and chemical properties. 4. Atoms of different elements combine in simple whole number ratios to form chemical compounds. 5. In chemical reactions, atoms cannot be subdivided, created, or destroyed. They are combined, separated, or rearranged. Atomic Model: Sphere-Billiard Balls Which principles remain true today? Aug 231:23 PM History of Atomic Theory Experiment: Cathode Ray Tube Electricity is passed through a glass tube filled with gas. The gas beam can be bent with a magnet. Joseph John Thomson J.J. Thomson ~ 1897 English Physicist Discoveries: • Atoms consist of charged particles • The negatively charged particles are called electrons (1897) • The positively charged particles are called protons (1920) Model: Plum Pudding or Chocolate Chip Cookies The chocolate chips represent negatively charged electrons stuck in positively charged dough. Aug 203:07 PM 3 HC History of the atom.notebook October 21, 2011 History of Atomic Theory Robert Millikan ~ 1909 Experiment: Oil Drop Discoveries: measured the charge on an electron determined the mass of an electron Aug 231:56 PM History of Atomic Theory Rutherford ~ 1911 English Physicist (a student of Thomson) Experiment: Gold Foil Experiment Positively charged alpha particles are shot at a piece of thin gold foil. Most alpha particles had little deflection. Some were deflected at large angles. Ernest Rutherford Discoveries: • A positively charged core of an atom called the nucleus • Electrons surround the nucleus • The rest of the atom is empty space Model: Nuclear atom Quote: "It was about as believable as if you had fired a 15 inch shell at a piece of tissue paper, and it came back and hit you."- Rutherford Aug 203:19 PM 4 HC History of the atom.notebook October 21, 2011 History of Atomic Theory James Chadwick ~ 1932 English Physicist (Student of Rutherford) Discoveries: • The nucleus contains another particle with no charge (neutral) and a mass equal to that of a proton. • Called this particle the neutron. Model: Sir James Chadwick Same as Rutherford Aug 203:30 PM Rutherford's Model Revisted The Nuclear Atom: • Atoms have a positively charged nucleus containing protons (+) and neutrons (0). Electrons (-) are located outside of the nucleus. The Problem: According to the laws of physics, with nothing holding the electrons in place, the negatively charged particles would spiral into the positively charged nucleus (- and + charges attract) and Rutherford's atom would collapse. Sep 167:33 AM 5 HC History of the atom.notebook October 21, 2011 Niels Bohr ~ 1913 Danish Physicist • student of both Thomson & Rutherford The Experiment: He studied the light spectrum of hydrogen gas in an attempt to more accurately explain the role of electrons in the atom. Sep 167:45 AM Electromagnetic Radiation: Radiation that is produced by electrically charged particles (like electrons!) All electromagnetic radiation exhibits wave like behavior. Wavelength - the distance from one peak to another Amplitude - the height of a peak Frequency - the number of waves that pass a particular point per second Sep 108:14 PM 6 HC History of the atom.notebook October 21, 2011 The Electromagnetic Spectrum Low Energy Light = longer wavelength & lower frequency red light = low energy light High Energy Light = shorter wavelength & higher frequency blue light = high energy light Sep 108:21 PM Light Spectrums Hydrogen Spectrum Helium Spectrum • Bohr observed when energy (light or heat) was added to a sample of hydrogen gas a distinctive color was given off. • When he looked at the light through a prism, only certain color lines could be seen, instead of the whole rainbow like he expected. • Concluded that the reason for why hydrogen's spectrum only contained a red, a blue, and a purple line involved the behavior of the electrons in the hydrogen atoms. Sep 167:59 AM 7 HC History of the atom.notebook October 21, 2011 Bohr's Experimental Conclusions • Thought the atom contained energy levels which kept the electrons from falling into the nucleus. • Individual elements each have a unique arrangement of electrons in those energy levels. • When an atom absorbs energy the electrons enter into an "excited state" by jumping up to higher energy levels. • The excited state is unstable, so the electrons release the excess energy in the form of colored light. • The color of the light observed is determined by the amount of energy released by the electrons. • Electrons in lower positions (closer to the nucleus) have lower energy (red light) • Electrons in higher positions (farther from the nucleus) have higher energy (blue light) Sep 167:27 AM Important Vocabulary: ground state- all electrons in their lowest possible energy levels excited state- electrons absorbed energy & jump to a higher energy level Sep 169:41 AM 8 HC History of the atom.notebook October 21, 2011 Bohr's Model: Solar System Electrons "orbit" around the nucleus in certain paths (energy levels). An electron must have a certain amount of energy to be "allowed" in a particular energy level. Similar to the way the planets in our solar system orbit around the sun. The energy levels prevent the electrons from falling into the nucleus. Sep 167:59 AM Excited vs. Ground State Cont... The Problem with Bohr's Model??? It worked well for hydrogen, but not for elements containing more than 1 electron. Sep 169:50 AM 9 HC History of the atom.notebook October 21, 2011 History of Atomic Theory Quantum Mechanics Erwin Schrodinger ~ 1926 Austrian Physicist Experiment: NONE! It is a mathematical model. It cannot be represented by anything that exists in the real world. Mathematical Equation: Problem: Erwin Schrodinger Heisenberg's Uncertainty Principle: you can never know how fast an electron is moving and where an electron is at the same time. In other words, you can find out where the electron started and you can see where the electron ended up but how it got there WE DON'T KNOW! Discoveries: • Mathematical model that deals with the probability of finding an electron within a given space around the nucleus • The probability is 90% • The given space are called orbitals (or electron clouds) • There are four orbitals with different shapes: s p d f • These orbitals can be related to the periodic table • Electrons behave like waves. Model: Aug 204:21 PM The Quantum Mechanical Model Quantum Numbers: Numerical values that mathematically describe the energy and location of electrons in atoms. From these values you get orbitals - probability maps that show the possible areas around the nucleus where electrons may be located. Four Quantum Numbers: 1. Principal Quantum Number (n)- Refers to the energy level (the distance of the orbital from the nucleus) where an electron is located. • When n=1 the electron is closest to the nucleus and has the least energy. • n = 1,2,3,4,5, etc. 2. Angular Quantum Number (l)- Refers to the shape of the orbital; s, p, d or f • l = 0, 1, 2 ... (n-1) • l = 0 (s-orbital), l = 1 (p-orbital), l = 2 (d-orbital)... 3. Magnetic Quantum Number (m)- Orientation of orbital(s) • m = any integer from -l to +l 4. Spin Quantum Number (s)- electron spin • electrons can have either "up spin" (+) or "down spin" (-) • s = +1/2 or -1/2 Aug 247:18 PM 10 HC History of the atom.notebook October 21, 2011 Orbital Shapes p - orbitals are dumbbell shaped s - orbitals are spherical Aug 258:00 PM Orbital Shapes d-orbitals are clover shaped f-orbitals Sep 76:48 PM 11 HC History of the atom.notebook October 21, 2011 Electron Configurations: Represent the arrangement of electrons in an atom. The Rules: 1. Aufbau Principle: An electron will occupy the lowest possible energy level. 2. Pauli Exculsion Principle: Each orbital cannot contain more than two electrons, and those electrons must have opposite spin. 3. Hund's Rule: Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all single electrons must have the same spin Sep 107:54 PM Sep 108:32 PM 12 HC History of the atom.notebook October 21, 2011 Using the Periodic Table to Write Configurations. • Each row on the periodic table represents an energy level. • The s, p, d and f-orbitals are represented by different "blocks" on the periodic table. • The d-orbitals are located 1 energy level lower than the s & p orbitals in the same row. • The electrons configuration for a particular element is determined by the location of that element on the table. Aug 234:09 PM The Orbital Blocks Note: When writing configurations, the "d" orbitals are 1 energy level lower than the "s" & the "p" around them. Aug 234:15 PM 13 HC History of the atom.notebook October 21, 2011 Examples... Hydrogen: coefficient = energy level (row on periodic table) H 1s1 letter = orbital type (periodic table block) superscript = number of electrons in that orbital (# of spaces over on table) He C Mg Fe Br Ag Configurations of Ions: Add or remove the number of electrons gained or lost to the end of the configuration. Examples: Mg: Mg+2: N: N-3: Aug 234:16 PM Final Few Notes on Configurations: • Most ions have electron configurations identical to a noble gas. • Noble gases are stable. Therefore when ions form they lose or gain electrons in order to have stable electron configurations as well. • Valence electrons - electrons in the highest energy level of an atom. Example: How many valence electrons does phosphorus have? • f-orbitals... 1 energy level lower than the "d" • Example: Au Aug 234:42 PM 14 Attachments Pictures of atom Chadwick apparatus Time line of scientists Schrodinger equation Conversation with science orbitals