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Atomic Theory

Atomic Theory. A Brief History. Atoms are made up of subatomic particles called protons, neutrons and electrons. How do we know that?. Vocabulary.

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Atomic Theory

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  1. AtomicTheory A Brief History

  2. Atoms are made up of subatomic particles called protons, neutrons and electrons • How do we know that?

  3. Vocabulary • Atom: The smallest unit of an element, having all the characteristics of that element and consisting of a dense, central, positively charged nucleus surrounded by a system of electrons. • Molecule: The smallest particle of a substance that retains the chemical and physical properties of the substance and is composed of two or more atoms.

  4. Compound: A compound is a substance made up of atoms representing more than one element bonded together and exhibiting distinct physical and chemical characteristics • Example: H2O, H2SO4

  5. Background • Law of Conservation of Mass (Lavoisier, 1789) • During a chemical reaction, the total mass of the reactants is equal to the total mass of the products. • Law of Definite Proportions (Proust, 1799) • When atoms combine to form compounds, they always combine in the same simple, small whole number proportions. • Example: Water is always H2O • Example: Sulfuric Acid is always H2SO4

  6. Aristotle(circa. 400 B.C.) • Matter is not made of particles, but rather is continuous. • The continuous matter is called “hyle.” • There were only four elements • Earth, Air, Fire, Water

  7. Democritus(circa. 400 B.C.) • Matter is made of empty space and tiny particles called “atoms.” • Atoms are indivisible. • There are different types of atoms for each material in the world.

  8. Why was Democritus Ignored? Because the early Greek philosophers did not experiment and because Aristotle was an established teacher and because the church was opposed to “soul atoms”, the views of Democritus were not accepted until the 19th century.

  9. Pre-Atomic Theory Postulates • Law of Conservation of Mass • During a chemical reaction, the total mass of the reactants is equal to the total mass of the products. • Law of Definite Proportions • When atoms combine to form compounds, they always combine in the same simple, small whole number proportions. • Example: Water is always H2O • Example: Sulfuric Acid is always H2SO4

  10. John Dalton(early 1803) • Matter consists of tiny particles called atoms which are indivisible and indestructible. • All atoms of a particular element are identical. • Atoms of different elements differ in mass and properties. • Atoms combine in whole number ratios to form compound atoms. • In chemical reactions, atoms are combined, separated, or rearranged but are never created, destroyed, or changed

  11. Why were Dalton’s views accepted? • The scientific method is now the proper way to “do science.” • Dalton’s theory was based on experimental observations: the law of Conservation of Mass and the law of Definite Proportions. • Dalton’s theory correctly predicted the outcome of future experiments. These predictions became the law of Multiple Proportions.

  12. The Dalton Atom • John Dalton examined the empirical proportions of elements that made up chemical compounds. • At this stage, the atom was still seen as an indivisible object, with no internal structure.

  13. Amedo Avogadro • Avogadro, among other achievements, was able to explain the existence of diatomic molecules. • Avogadro’s Law: Equal volumes of any gas at the same temperature and pressure, have the same number of particles. • 1 mole = 22.4 Liters

  14. J.J. Thomson set up a crookes tube with a anodic and cathodic ends

  15. When electricity was applied to the tube, a beam was emitted from the cathodic (-) plate • Thomson then assumed the particles emitted were negative • To test this theory, he applied a magnetic field to the tube and “bent” the beam • What happens with like charges?

  16. He tested the tube further by applying an electrical field to the tube using paddles • The tube turned around • Thomson determined that the tube turned as tiny particles hit the paddles

  17. Demonstration • Molecular Expressions: Electricity and Magnetism - Interactive Java Tutorials: Crookes Tube

  18. He concluded that the particles in the tube were negatively charged and had mass • mass = 9.109 x 10-31kg

  19. Since these particles are negatively charged, but the atoms are neutral, there must be other particles in an atom • Problem: This requires too many electrons!

  20. Thomson Model • The discovery of the electron by J. J. Thomson showed that atoms did have some kind of internal structure. • The Thomson model of the atom described the atom as a "pudding" of positive charge, with negatively charged electrons embedded

  21. J.J. Thomson’s Plum Pudding Model Positively charged “pudding” Negatively charged particles later named electrons

  22. Thomson movie

  23. Milliken and the Oil Droplet • In 1909, Robert Milliken performed an experiment using droplets of oil to determine the charge of an electron. • electrons, e, e-, -1.602 x 10-19C

  24. Ernest Rutherford conducted experiments to test the Thomson model • He directed alpha particles through a thin gold foil and measured them with a film • Most particles went through the foil • But, some were deflected, Why?

  25. Rutherford’s Hypothesis England, 1911 • Rutherford hypothesized that the particles were travelling through a void and occasionally bouncing off a concentrated positive charge. a

  26. Conclusion • There must be a dense region with positive charges surrounded by the electrons • An atom is mostly empty space with a dense region in the middle. • This dense region is called the “nucleus” • He measured the number of particles deflected and the angles and calculated that the radius of the nucleus was 1/10,000 of the whole atom Problem: Electrons should spiral into the nucleus.

  27. Let there be protons! • The discovery was made and protons were recognized • The mass of a proton is 2000x the mass of an electron • 1.673 x 10-27 kg

  28. We’re not done yet ... • 30 years later, Irene Curie, the daughter of the great Madame Curie, produced a beam of particles that could go through almost anything • And James Chadwick determined this beam was not affected by a magnetic field (no charge!) • Neutrons were given credit

  29. Coulomb’s Law • Since like charges repel, how can the nucleus be stable with protons (+) and neutrons (0)? • Coulomb’s Law: the closer two charges are, the greater the force between them • As the distance between like charges decreases, the force between them increases. • Try it!

  30. Problems with Rutherford’s model • According to classical physics, an electron in orbit around an atomic nucleus should emit photons continuously as they are accelerating in a curved path. • The loss of energy should cause the electron to collide with the nucleus and collapse the atom.

  31. Elemental Quandary • The Rutherford model was unable to explain the difference in the visible spectrum for each element.

  32. Visible-line Spectrum • When an elemental gas is excited by electricity, it emits a distinct visible light pattern. • The color of each spectral line is identified by the wavelength ()

  33. Electromagnetic Spectrum • All of the frequencies or wavelengths of electromagnetic radiation.

  34. Wavelength • The wavelength is the distance between repeating units of a wave pattern (λ) and measured in nm

  35. Frequency • Frequency is the measurement of the number of times that a repeated event occurs per unit of time (Hz) • The blue wave has the greatest frequency.

  36. Hydrogen

  37. Carbon

  38. Oxygen

  39. Xenon

  40. Compare these spectrum • Hydrogen, Carbon, Oxygen and Xenon

  41. In comes Niels Bohr Denmark, 1913 • In 1913, Bohr proposed that electrons were restricted to certain fixed circular orbits. • Orbits are energy levels • Electrons can jump from ground state to an excited state by absorbing energy or a photon with the precise wavelength.

  42. Neils Bohr(early 1900’s) • Electrons travel around the nucleus in specific energy levels. • Electrons have a ground state and an excited state • Electrons do not radiate energy in their normal energy level called the ground state. • Electrons absorb energy and move to energy levels further from the nucleus called excited states. • Electrons lose energy (light) as they return to lower energy levels.

  43. The Bohr Atom Light Excited States - + Ground State Nucleus

  44. The Bohr Planetary Atomic Model

  45. In the Bohr Model the neutrons and protons occupy a dense central region called the nucleus, and the electrons orbit the nucleus much like planets orbiting the Sun The Bohr Atom

  46. The Modern Atom • The modern atom is further defined by the works of these scientists: • de Broglie • Max Plank • Albert Einstein • Heisenberg • Erwin Schrodinger

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