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The Structure of the Atom and Electrons in Atoms: Early Theories, Discovering the Electron, and Nuclear Model

This chapter explores the early theories of matter, including Democritus and Aristotle, as well as the contributions of scientists like John Dalton, J.J. Thomson, Robert Millikan, and Ernest Rutherford. It also covers the discovery of subatomic particles, how atoms differ, atomic masses, and the basics of radioactivity and nuclear reactions. Additionally, it discusses electromagnetic radiation, the particle nature of light, the photoelectric effect, atomic emission spectra, the Bohr model of the atom, and the wave properties of electrons.

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The Structure of the Atom and Electrons in Atoms: Early Theories, Discovering the Electron, and Nuclear Model

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  1. Chapters 4 and 5 The Structure of the Atom And Electrons in Atoms

  2. Early Theories of Matter • Democritus (460-370 B.C.) • Named atom (atomos)

  3. Early Theories of Matter • Aristotle (384-322 B.C.)

  4. Early Theories of Matter • John Dalton (1766-1844) • First Atomic Theory

  5. Defining an Atom • The smallest particle of an element that retains the properties of the element. • About 1 X 10-10 m in diameter. • Can be seen with a scanning tunneling microscope.

  6. Discovering the Electron • William Crookes (1800’s)

  7. Discovering the Electron • J.J. Thomson (late 1890’s) • Determined the charge-to-mass ratio • Mass must be less than a hydrogen atom • Plum Pudding Model of atom

  8. Discovering the Electron • Robert Millikan (1909) • Determined charge of electron • 1/1840 mass of a hydrogen atom

  9. The Nuclear Atom • Ernest Rutherford (1911)

  10. The Nuclear Atom • Atom contains: • Mostly empty space • Tiny, dense nucleus which is positively charged • Creates nuclear model of atom

  11. Other Subatomic Particles • Rutherford (1920) • Concluded nucleus contains proton • Proton as equal but opposite charge of electron • James Chadwick (1932) • Discovered neutron • Neutron has no charge

  12. Subatomic Particles

  13. How Atoms Differ • Moseley (shortly after Gold Foil) • Atoms of each element contain a unique number of protons • Atomic Number= #protons • Identifies the atom

  14. Isotopes • Isotopes – atoms that contain the same number of protons but different number of neutrons. • Most elements contain a mixture of isotopes. • The relative abundance of each isotope is constant.

  15. Isotopes • Mass Number = #protons + #neutrons

  16. Simple Practice 12 13 12 65 30 30 4 5 4 80 200 80

  17. Mass of Atoms • Atomic mass unit – 1/12 of a carbon-12 atom. • Atomic Mass – weighted average mass of the isotopes of that element.

  18. Calculating Atomic Masses • 6X has mass of 6.015 amu and abundance of 7.50%. 7X has mass of 7.016 amu and abundance of 92.5%. • (6.015)(.0750) + (7.016)(.925) = 6.94 amu

  19. More Challenging Problems! • Cu-63 has a mass of 62.940 amu and an abundance of 69.17%. Find the mass and abundance of the other isotope. • Boron has two isotopes with the masses of 10.013 amu and 11.009 amu. Find the abundance of each isotope.

  20. Radioactivity • Nuclear Reactions – changes an atom’s nucleus. • Atom changes into a new element • Due to unstable nuclei • Radiation contains rays and particles emitted from a radioactive material. • Radioactive decay is the spontaneous emission of radiation.

  21. Types of Radiation

  22. Types of Radiation

  23. Nuclear Reactions • Mass numbers and Atomic numbers on both sides of the reaction must be equal • Practice Problem:

  24. Chapter 5 Electrons in Atoms

  25. Electromagnetic Radiation • Electromagnetic Radiation is a form of energy that has wave-like behavior. • 4 properties of waves: wavelength, amplitude, speed and frequency.

  26. Properties of Waves • Frequency()- number of waves that pass a given point per second. (hertz or 1/s or s-1) • Speed (c)- is constant for all waves. 3 x 108 m/s

  27. Calculating Properties of Waves • c= • What is the frequency of light with a wavelength of 5.80 x 10-7 m? • A radio station broadcasts with a frequency of 104.3 MHz. What is the wavelength of the broadcast?

  28. Particle Nature of Light • Max Planck (1900) discovered that matter can gain or lose energy in small, specific amounts called quanta. • Equantum= h • Planck’s Constant (h)=6.626 x 10-34J·s

  29. Practice Problems • What is the energy of a wave with a frequency of 6.25 x 1019Hz? • What is the frequency of a wave that contains 8.64 x 10-18J of energy? • A wave contains 4.62 x 10-15J of energy. Determine its wavelength.

  30. Photoelectric Effect • Photoelectric effect – electrons are emitted from a metal’s surface when light of a certain frequency shines on it. • Frequency (color) of light, not brightness of light determines if electrons are emitted. • Einstein (1905)- light has wave-like properties but is also a stream of tiny particles or bundles of energy called photons. • Photon – a piece of EM with no mass and carries a quantum of energy.

  31. Atomic Emission Spectrum • When atoms absorb energy they become excited. • Atomic Emission Spectrum- unique set of frequencies emitted by excited atoms.

  32. Bohr Model of the Atom • Bohr (1913) proposed why the emission spectrum of hydrogen is not continuous. • Electrons can have only certain “energy states” • Ground State - the lowest allowable energy state. • Excited State – energy state of an electron when it gains energy

  33. Bohr Model of the Atom

  34. Electrons as Waves • Louis de Broglie (1924) thought Bohr’s model had electrons having similar properties to waves. • de Broglie equation: • Predicts that all moving particles have wave properties.

  35. Heisenberg Uncertainty Principle • When viewing an electron, a photon of light hits it and changes the velocity and position of the electron. • It is impossible to know precisely both the velocity and position of a particle at the same time.

  36. Quantum Mechanical Model of the Atom • Schrödinger (1926) derived an equation that treated hydrogen’s electron as a wave. • Allows electron to have only certain energy but does not give path of electron. • Atomic orbital – a 3-D region around the nucleus in which the electron can be found 90% of the time.

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