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Chapter 4 Atomic Theory

Chapter 4 Atomic Theory. Democritus hypothesis of the atom Dalton’s atomic theory Subatomic particles Rutherford’s atomic model Atomic mass of an element Mass number of an element Isotopes. The Atom. In chapters 1-3, we described matter in terms of physical and chemical properties

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Chapter 4 Atomic Theory

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  1. Chapter 4Atomic Theory • Democritus hypothesis of the atom • Dalton’s atomic theory • Subatomic particles • Rutherford’s atomic model • Atomic mass of an element • Mass number of an element • Isotopes

  2. The Atom • In chapters 1-3, we described matter in terms of physical and chemical properties • Specifically, we explained matter with regards to its composition • When we referred to composition, we were actually referring to the ATOM. • Definition of an atom: A particle of matter that uniquely defines a chemical element • Therefore, an atom is the basis for the composition of matter.

  3. Democritus (Greek: ca. 400 B.C.) Matter---composed of tiny indestructible particles (Democritus’s hypothesis) Atoms remain unaltered, but, move about in space to combine in various ways to form all macroscopic objects No experiments to prove existence of atoms Hypothesis did not explain chemical behavior Among the first to claim atoms are real Dalton (English: ca. 1800) Transformed Democritus’s hypothesis into theory Used an experimental approach Linked macroscopic changes to activity at atomic level Developed a system of chemical symbols based on atomic mass Proposed atoms determine the composition of matter Chemical combination of different elements occur in simple numerical ratios by weight Dalton’s atomic theory is the basis for modern atomic theory Early Models of the Atom

  4. Structure of the nuclear atomSubatomic particles Subatomic particles • One major change in Dalton’s theory is that atoms can be broken down into: 1. Electrons (negatively charged) 2. Protons (positively charged) 3. Neutrons (neutral)

  5. Electrons • Discovery of the electron is attributed to J. J. Thomson (English physicist; 1856-1940) • He discovered electrons using a Cathode ray tube • Measured mass-to-charge ratio (m/z) of particles emitted after using a variety of gases and electrodes • He found that the m/z ratio of the particle did not change in all the different gases he used

  6. Cathode Ray Tube (CRT)

  7. Electrons (continued) • Conclusion: electrons are parts of atoms of all elements • Showed electrons were negatively charged • Proposed “Plum pudding” model of the atom: electrons are embedded in a sphere of positive charge

  8. Ernest Rutherford (ca. 1870) • Student of J. J. Thomson • J. J. Thomson: an atom was a solid sphere of positive charge • Electrons, embedded in the sphere like the seeds in a watermelon. • Rutherford disproved the Plum pudding model

  9. Rutherford Gold Foil Experiment

  10. Ernest Rutherford (continued) • Theory: i) Atoms are mostly empty space ii) Most of mass centralized in a very small area called nucleus • Rutherford’s model known as the NUCLEAR ATOM • Nuclear atom: protons and neutrons in nucleus • Electrons are distributed around the nucleus • Electrons occupy almost all the volume of the atom • Based on the model, nucleus is very tiny compared to the entire atom (football stadium vs. a marble)

  11. The Atom (Rutherford Model)

  12. Atomic and Mass number • Atomic number: Number of protons (and electrons) in an atom • Mass number: Total number of particles (protons and neutrons) in an atom's nucleus. Mass # = (# of Protons) + (Number of Neutrons) • For krypton, this equation becomes: 83.80 = 84 = (Number of Protons) + (Number of Neutrons) # of Neutrons = 84 – (# of Protons) # of Neutrons = 84 – 36 = 48

  13. Atomic and Mass number • In the shorthand notation, in Krypton for example, the atomic number is usually listed on the top left corner of the box and the mass number below the element symbol • Note: Mass number is always larger than the atomic number

  14. Isotopes • Isotopes = atoms with the same atomic number but different neutrons • Have different mass numbers • Because isotopes contain the same # of electrons and protons they have similar chemical properties • This is because electrons determine the chemical reactivity of elements To calculate the relative mass (mass #) of an element, e.g. chlorine, based on the abundance of its isotopes: Mass # = (35 * 0.7577) + (37 * 0.2423) = 26.5195 + 8.9651 = 35.4846 = 35.49 (35.5) amu

  15. Atomic, Mass number, Isotopes (cont.) In Summary... • For any element: • Number of Protons = Atomic Number • Number of Electrons = Number of Protons = Atomic Number • Number of Neutrons = Mass Number - Atomic Number • For krypton: • Number of Protons = Atomic Number = 36 • Number of Electrons = Number of Protons = Atomic Number = 36 • Number of Neutrons = Mass Number - Atomic Number = 84 - 36 = 48

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