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Chapters 4 and 5

Chapters 4 and 5. The Structure of the Atom And Electrons in Atoms. Early Theories of Matter. Democritus (460-370 B.C.) Named atom ( atomos ). Early Theories of Matter. Aristotle (384-322 B.C.). Early Theories of Matter. John Dalton (1766-1844) First Atomic Theory. Defining an Atom.

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Chapters 4 and 5

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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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