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Theory of Atomic Structure

Theory of Atomic Structure. Greeks – Democritus, Leucippus. Over 2000 years ago All matter is composed of tiny particles These particles are so small that they cannot be broken down (indestructible) Named these particles atoms (from the Greek work for indestructible)

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Theory of Atomic Structure

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  1. Theory of Atomic Structure

  2. Greeks – Democritus, Leucippus • Over 2000 years ago • All matter is composed of tiny particles • These particles are so small that they cannot be broken down (indestructible) • Named these particles atoms (from the Greek work for indestructible) • Common Greek theory was that all matter consisted of four "elements" - earth, air, fire, and water

  3. John Dalton • 1803 - Proposed an "atomic theory" with spherical solid atoms based upon measurable properties of mass • All elements are composed of atoms, which are indivisible and indestructible particles • All atoms of the same element are exactly alike; in particular they all have the same mass • Joining atoms of two or more elements forms compounds. In any compound, the atoms are joined in a definite whole number ratio (H2O, 2 to 1 ratio)

  4. Dalton • Did account for the following laws: • Law of conservation of mass (a chemical change is simply a rearrangement of atoms) • Law of definite proportions (atoms are combined in definite ratios) • Law of multiple proportions (elements can combine in different ratios to form new compounds – H2O, H2O2 • Did not account for isotopes

  5. J.J Thompson • Through the use of cathode rays, Thompson was able to discover a negatively charged particle with almost no mass • Discovered electrons • See Cathode Ray Tube

  6. “Plum Pudding” Model (1890) • Negative charges were scattered throughout the positive atom

  7. Ernest Rutherford • Discovered that atoms have a positive, dense nucleus surrounded by negative “empty” space (the electrons are in the “empty” space) • Gold Foil Experiment – 1909 (chem ASAP)

  8. Gold Foil Experiment • What was used • Alpha (α) particles – positively charged • Same as Helium nuclei (42He2+) • Gold Foil – only a few atoms thick • What was done • Shot a beam of α particles through the gold foil

  9. Gold Foil Experiment • What happened? • The majority of the particles went straight through the foil and exposed the film behind it • A few particles were deflected back from the foil and scattered

  10. Gold Foil Experiment • What does this mean? • Since the α particles are (+) charged, they must have been pulled through the foil by a strong (-) charge • The atom was mostly negatively charged “empty” space • Since some α particles bounced back, they must have hit a positive part of the foil • The atom must have a small positive part located in the center

  11. Gold Foil Experiment • Atomic Theory (Conclusion) • An atom consists of a very small, positively charged nucleus and the rest is negatively charged empty space. This empty space must contain the electrons. • Lets see the experiment again!

  12. Bohr Model – Neils Bohr 1919 • Electrons resemble our solar system • Electrons revolve (orbit) around the nucleus • Principle energy levels (PEL) = rings around the nucleus

  13. Principle Energy Levels • The only regions in which electrons can be found • Each level can only contain a specific number of electrons (2 in the 1st, 8 in the 2nd) • Electrons farther away from the nucleus have higher energy • Number of PEL’s is given by the period number (found on the periodic table)

  14. James Chadwick • 1932 – discovered neutrons • Using alpha particles he discovered a neutral atomic particle with a mass close to a proton.

  15. Valence Electrons – outer electrons • Outermost electrons give an atom its properties • These are the electrons that are involved in reactions • Last number in the electron configuration • Kernel electrons – inner electrons

  16. Electron Dot Symbols / Lewis Dot Diagrams • Dots represent valence electrons • Examples: • K • Ca • S • Ar • Ca2+ • S2-

  17. Electron Configurations • Distribution of electrons • Given in the lower left hand corner of the periodic table (below the atomic number)

  18. Rules • 1st PEL can hold 2 electrons • 2nd PEL can hold 8 electrons • 3rd PEL can hold 18 electrons • 4th PEL can hold 32 electrons • No more than 8 electrons are in the outermost principle energy level • This equal 8 valence electrons, the maximum amount of valence electrons possible • The electrons are added one at a time to the unfilled principle energy levels

  19. Examples • Ca: 2-8-8-2 What does this mean? • 4 principle energy levels contain electrons • 1st PEL contains 2e-, 2nd and 3rd PEL contain 8e-, and the 4th PEL contains 2e- • 1st and 2nd are filled, 3rd and 4th are not completely full • Ca has 2 valence electrons, and 18 kernel electrons

  20. Examples: • C • Li • F

  21. Detailed Electron Configurations • Within each principle energy levels there are sublevels • Each s sublevel contains 2 electrons • Each p sublevel contains 6 electrons • Each d sublevel contains 10 electrons • Each f sublevel contains 14 electrons

  22. Example • Calcium: • 1s22s22p63s23p64s2 • What does this mean??? • The number in front of the sublevel indicates the PEL (4s = 4th PEL, s sublevel) • The number of electrons is given by the superscript (4s2 = 4th PEL, s sublevel, 2 e-)

  23. Rules • Fill the lowest energy sublevels first • See energy chart and diagonal method • No more than 8 electrons can be in the outermost principle energy level

  24. Examples • C • Al • Ca • Ni

  25. Ground State • Electrons are in the lowest available energy levels • This is how the electron configuration is written on the periodic table • Normal configuration • Stable

  26. Excited State • When an electron gains energy, the electron is at a higher energy state (excited state) • Electrons absorb energy and move to new higher energy levels • UNSTABLE

  27. Excited State • When an electron returns from the excited state to the ground state, energy is released • The energy is often emitted in the form of light • The wavelength (color) of the light can be used to identify the element because each element absorbs and releases a specific amount of energy

  28. Quantum Theory • Electrons can only absorb (or release) specific amounts of energy called quanta • This energy corresponds to the differences between energy levels • Energy is always absorbed (or released) in definite amounts rather than in a continuous flow • A quantum of radiation is called a photon

  29. Flame Tests • Every element absorbs and hence emits different amounts of energy (different colors) • Can be used for identification

  30. Spectra/Spectroscope • Spectra • The photons of light that are given off can be broken down into lines (spectral lines) • Spectroscope • An instrument that breaks light into colored bands • Every element has a unique band length (arrangement)

  31. Modern Orbital Theory (Wave Mechanical/Electron Cloud Model) • Electrons do not have a specific path or location – they have probabilities of being located there • Regions of high probabilities are called orbitals • Electrons are located in orbitals

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