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History of Atomic Theory Alchemy ~ Before 400 B.C. Experiment: Pseudoscience concerned with: • Changing metal to gold • Finding an eternal life elixir Aristotle Beliefs: • All matter was made up of a combination of the four elements • Four elements: fire, wind, earth, water 1 History of Atomic Theory Democritus ~ 400 B.C. Experiment: None. He had beliefs that were disregarded. Beliefs: • Named the atom "Atomos." It means indivisible • Matter is composed of atoms too small to be seen. • Empty space between atoms • Atoms are solid and homogeneous • Different atoms have different sizes and shapes Democritus Model: Sphere 2 History of Atomic Theory Dalton ~ early 1803 English School Teacher John Dalton Experiment: Studied chemical reactions, making observations and measurements Beliefs: 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. Model: Sphere-Billiard balls Which principles remain true today? 3 History of Atomic Theory Experiment: Cathode Ray Tube Electricity is passed through a gas tube. 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 are electrons stuck in positive dough 4 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 5 History of Atomic Theory James Chadwick ~ 1932 English Physicist (Student of Rutherford) Experiment: • Beryllium foil was bombarded with alpha particles Sir James Chadwick • A neutral radiation was emitted • Emitted radiation would then knock protons out of the nuclei of other substances • The radiation was a stream of neutral particles having the same mass as a proton Discoveries: • A neutrally charged subatomic particle • The particle was called a neutron Model: Same as Rutherford 6 Review History of Atom http://glencoe.mcgrawhill.com/sites/dl/free/0078759864/164155/00044672.html 7 History of Atomic Theory Bohr ~ 1913 Danish Physicist (a student of Rutherford) Problem: According to the laws of physics, charged particles will radiate energy when orbiting and spiral into the nucleus. Rutherford's atoms would all collapse. Niels Bohr Experiment: Light Spectrums Electrons give off energy in the form of colored light by falling from an excited state to a ground state. Discoveries: • Electrons can only be found at certain energy levels • Each energy level requires a certain amount of energy • Lower (closer) levels have lower energy • Higher (farther) levels have higher energy Model: Photo courtesy NASA Hydrogen Spectrum Helium Spectrum Ladder Rungs-Cannot stand in between rungs Vocabulary: ground state- all electrons in their lowest possible energy levels excited state-electrons absorbed energy & jump to a higher energy level 8 9 10 History of Atomic Theory Quantum Mechanics Erwin Schrodinger ~ 1926 Austrian Physicist Problem: Bohr's model works well for Hydrogen, but fails for every other element Mathematical Equation: Experiment: Erwin Schrodinger NONE! It is a mathematical model. It cannot be represented by anything that exists in the real world. Discoveries: • Mathematical model that deals with the probability of finding an electron within a given space • 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 have wave properties Model: 11 Electromagnetic Radiation: waves that are produced by electrically charged particles Ex: sunlight, X rays, microwaves Electromagnetic Spectrum 12 13 All electromagnetic radiation exhibits wave like behavior (wavelength, frequency, and speed) 14 Quantum Numbers: Specify the properties of atomic orbitals and the properties of electrons in those orbitals. Each e- has a set of four numbers; no two electrons in the same atom can have the same four numbers. Four Quantum Numbers: Principal Angular Magnetic Spin 1. Principal Quantum Number (n)- Refers to the distance of the orbital from the nucleus (says which of the main energy levels an e- is in) • When n=1 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 (also associated a. Possible Shapes with angular momentum) • s, p, d, f (in order of increasing energy) • l=integer values from 0 to n-1 b. Number of possible shapes is limited by the principal quantum number 15 3. Magnetic Quantum Number (m)- Orientation of orbital(s) (direction of angular momentum) - m=any integer from -l to +l 4. Spin Quantum Number (s)- State of the electron that occupies an orbital -Electrons are assigned one of the two possible directions it can be spinning - + spin or - spin 16 17 18 19 20 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! 21 Electron Configuration: The arrangement of electrons in an atom is known as the atoms electron configuration. Rules: 1. Aufbau Principle: An electron occupies the lowest energy level that can receive it. 2. Pauli Exculsion Principle: No two electrons in the same atom can have the same set of four quantum numbers. • If two electrons have the same n, m, and l values they have to have different spins • Two electrons fit into each orbital 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 electrons in singly occupied orbitals must have the same spin 22 23 24 4p 3d 4th energy level 4s 3p 3rd energy level 3s 2nd energy level 2p 2s 1st energy level 1s 25 26 27 28 29 30 Attachments Pictures of atom Chadwick apparatus Time line of scientists Schrodinger equation Conversation with science orbitals