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History of the atom

History of the atom. Dalton J.J. Thompson Rutherford Bohr. Dalton’s Atomic Theory. All elements are composed of atoms which are indivisible Atoms of the same element are identical Atoms of different elements can mix together in simple whole number ratios to form compounds

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History of the atom

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  1. History of the atom • Dalton • J.J. Thompson • Rutherford • Bohr

  2. Dalton’s Atomic Theory • All elements are composed of atoms which are indivisible • Atoms of the same element are identical • Atoms of different elements can mix together in simple whole number ratios to form compounds • Chemical reactions occur when atoms are separated, joined or rearranged. Atoms of one element can’t be changed into atoms of another element Law of Definite Proportions and Law of Multiple Proportions Law of Conservation of Matter

  3. Early Atomic Models • J.J. Thompson’s plum-pudding model • Rutherford’s Model of the atom

  4. Problems with Rutherford Model • Couldn’t account for the chemical properties of the elements (why do elements react in the way that they do?) Rutherford

  5. The Wave Nature of Light • Electromagnetic Radiation Forms of energy that exhibit wave like characteristics • Wavelength • How far apart the waves are • Frequency • The # of waves to pass a point in a certain timeframe • Wavelength and Frequency are Inversely Proportional

  6. Wavelength and Frequency • Not Working very hard • Long Wavelength • Low Frequency

  7. Wavelength and Frequency • Working VERY hard • Short Wavelength • High Frequency What does this tell you about the relationship between wavelength, frequency and energy?

  8. The Photo Electric Effect • In the early 1900s an experiment was done that COULD NOT be explained by light being a wave. • Different colors were shined on to a metal plate • Electrons would come off the metal plate for ONLY CERTAIN FREQUENCIES (colors) OF LIGHT If light only behaved like a “wave” ANY frequency of light would cause the electrons to be released.

  9. Brick Wall Analogy • Different balls thrown at wall • Each one with different mass • All at same speed of 90 mile/hour • Each ball is like a color or light - each has its own energy No effect Dislodges Brick and send it flying No effect Super Dense Steel Ball Ping Pong Ball Softball JB

  10. Einstein and the Photo Electric Effect • This observation led Einstein to believe that light acted like a particle and a wave • This is called the “dual nature” of light • Light carried packets of energy called quanta, or photons

  11. Neils Bohr and the New Model of the Atom • Bohr hypothesized that electrons could only be at certain energy “levels” around the nucleus • Electrons could “jump” from lower to higher energy states by absorbing a quantum of energy • When an electron releases the energy it gained, specific colors or wavelengths of light are emitted Mercury Line Spectra

  12. Release Energy in the form of colored light Add Energy e- e- e- electron at “Excited State” Quantum Leap Just like a jumper has potential energy at the top of the jump, the electron has stored potential energy in the higher orbit. Electron at “Ground State” Electron at “Ground State” Electrons and the Atom Electrons disappear from one orbit and reappear at another without visiting the space in between! Nucleus JB

  13. Energy and the Electron • Ground State – lowest energy state for an electron • Excited state – high energy state for e- • Quantum – exact amount of energy to move an electron from one energy level to another

  14. Heisenberg Uncertainty Principle • It is impossible to pinpoint the exact location and velocity of an electron at any point in time • You can estimate where an electron will be 90% of the time • An electron cloud shows where an electron spends most of its time

  15. Quantum Mechanical Atomic Model • Similar to Bohr model except that e- cannot be found in distinct orbits around the nucleus • Determines how likely it is for an electron to be found in various regions around the nucleus.

  16. Atomic Orbitals • Region around the nucleus where an electron of a given energy is likely to be found • Each orbital has a characteristic size, shape, and energy • There are four different orbitals: s, p, d, f Different types of atomic orbitals

  17. Principle Energy Levels • Symbolized by n = 1,2,3,4 etc. • Each energy level contains sublevels denoted by a number and a letter (ex. 1s) • Each sublevel contains a certain # of orbitals • Every orbital can hold 2 e-

  18. Electron Configurations • Aufbau principle: e- occupy the lowest energy sublevels 1st (see pg. 133 figure 5.7) • Pauli Exclusion Principle: every orbital can hold a maximum of 2 e- (paired e- spin in the opposite direction) • Hund’s Rule: e- fill all empty orbitals in a sublevel BEFORE they pair up Aufbau Diagram

  19. Valence Electrons • Valence electrons – electrons found in the highest energy level of an atom • Valence e- determine the chemical properties of the element • A filled outer energy level (stable octet) of 8 e- makes an atom stable (or for He 2e- fills its outer energy level)

  20. Electron Dot Diagrams • Shows only number of valence e- • Electrons shown as dots Carbon has 4 valence e-

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