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AP CHEMISTRY Atomic Structure and Electrons

AP CHEMISTRY Atomic Structure and Electrons. Ch. 7 sec 1-9. Light and Quantized Energy. Chemists found Rutherford’s nuclear model lacking because it did not begin to account for the differences in chemical behavior among the various elements.

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AP CHEMISTRY Atomic Structure and Electrons

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  1. AP CHEMISTRYAtomic Structure and Electrons Ch. 7 sec 1-9

  2. Light and Quantized Energy • Chemists found Rutherford’s nuclear model lacking because it did not begin to account for the differences in chemical behavior among the various elements. • In the early 1900s, scientists observed that certain elements emitted visible light when heated in a flame.

  3. Light and Quantized Energy • Analysis of the emitted light revealed that an element’s chemical behavior is related to the arrangement of the electrons in its atoms. • In order to better understand this relationship and the nature of atomic structure, it will be helpful to first understand the nature of light.

  4. The Electromagnetic Spectrum • Electromagnetic radiation includes radio waves that carry broadcasts to your radio and TV, microwave radiation used to heat food in a microwave oven, radiant heat used to toast bread, and the most familiar form, visible light. • All of these forms of radiant energy are parts of a whole range of electromagnetic radiation called the electromagnetic spectrum.

  5. The Electromagnetic Spectrum

  6. Particle Nature of Light • While considering light as a wave does explain much of its everyday behavior, it fails to adequately describe important aspects of light’s interactions with matter.

  7. The Quantum Concept • In 1900, the German physicist Max Planck (1858–1947) began searching for an explanation as he studied the light emitted from heated objects.

  8. The Quantum Concept • His study of the phenomenon led him to a startling conclusion: matter can gain or lose energy only in small, specific amounts called quanta. • That is, a quantum is the minimum amount of energy that can be gained or lost by an atom.

  9. The Quantum Concept • While a beam of light has many wavelike characteristics, it also can be thought of as a stream of tiny particles, or bundles of energy, called photons • Thus, a photon is a particle of electromagnetic radiation with no mass that carries a quantum of energy.

  10. Atomic Emission Spectra • The atomic emission spectrum of an element is the set of frequencies of the electromagnetic waves emitted by atoms of the element.

  11. Atomic Emission Spectra • Hydrogen’s atomic emission spectrum consists of several individual lines of color, not a continuous range of colors as seen in the visible spectrum. • Each element’s atomic emission spectrum is unique and can be used to determine if that element is part of an unknown compound.

  12. Atomic Emission Spectra

  13. Atomic Emission Spectra • An atomic emission spectrum is characteristic of the element being examined and can be used to identify that element. • The fact that only certain colors appear in an element’s atomic emission spectrum means that only certain specific frequencies of light are emitted.

  14. Bohr Model of the Atom • Why are elements’ atomic emission spectra discontinuous rather than continuous? • Niels Bohr, a young Danish physicist working in Rutherford’s laboratory in 1913, proposed a quantum model for the hydrogen atom that seemed to answer this question. • Impressively, Bohr’s model also correctly predicted the frequencies of the lines in hydrogen’s atomic emission spectrum.

  15. Energy States of Hydrogen • Building on Planck’s and Einstein’s concepts of quantized energy (quantized means that only certain values are allowed), Bohr proposed that the hydrogen atom has only certain allowable energy states. • The lowest allowable energy state of an atom is called its ground state.

  16. Energy States of Hydrogen • When an atom gains energy, it is said to be in an excited state. • And although a hydrogen atom contains only a single electron, it is capable of having many different excited states. A02220MV.mpg

  17. Energy States of Hydrogen • Bohr went even further with his atomic model by relating the hydrogen atom’s energy states to the motion of the electron within the atom. • Bohr suggested that the single electron in a hydrogen atom moves around the nucleus in only certain allowed circular orbits.

  18. Energy States of Hydrogen

  19. Hydrogen’s Line Spectrum • Bohr suggested that the hydrogen atom is in the ground state, also called the first energy level, when the electron is in the n = 1 orbit.

  20. Hydrogen’s Line Spectrum • When energy is added from an outside source, the electron moves to a higher-energy orbit such as the n = 2 orbit shown.

  21. Hydrogen’s Line Spectrum • Such an electron transition raises the atom to an excited state. • When the atom is in an excited state, the electron can drop from the higher-energy orbit to a lower-energy orbit. • As a result of this transition, the atom emits a photon corresponding to the difference between the energy levels associated with the two orbits.

  22. Hydrogen’s Line Spectrum • The four electron transitions that account for visible lines in hydrogen’s atomic emission spectrum are shown.

  23. The Heisenberg Uncertainty Principle • Heisenberg concluded that it is impossible to make any measurement on an object without disturbing the object—at least a little. • The act of observing the electron produces a significant, unavoidable uncertainty in the position and motion of the electron.

  24. The Heisenberg Uncertainty Principle • Heisenberg’s analysis of interactions such as those between photons and electrons led him to his historic conclusion. • The Heisenberg uncertainty principlestates that it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time.

  25. Mathematic Equations c=ln c= speed of light (3.0 x 108m/s) l= wavelength (m) n= frequency (s-1 or Hz) E=hn E=energy (J or kg∙m2/s2) h=Planck’s constant (6.63x10-34 J∙s or kg∙m2/s) n=frequency (s-1) Combining them : E=hc/l

  26. Mathematic Equations (cont’d) deBroglie equation l=h/mv l=wavelength (m) h=6.63x10-34 kg∙m2/s m=mass (kg) v=velocity (m/s) p=mv p=momentum (kg∙m/s) m=mass (kg) v=velocity (m/s)

  27. Energy levels, sublevels, & orbitals • Energy levels=clouds or shells around nucleus (n=1,2,3…) • Sublevels=found inside energy levels (s,p,d,f) • Atomic orbitals=found within sublevels: • s = 1 orbital (sphere) • p = 3 orbitals (dumbell) • d = 5 orbitals (p. 313) • f = 7 orbitals • 2 e- max per orbital

  28. Rules Governing e- Configurations • Aufbau Principle = e- fill orbitals with lowest energy first • Pauli Exclusion Principle = e- in the same orbital have opposite spins → no 2 e- in a single atom will have the same set of quantum numbers • Hund’s Rule = e- occupy one orbital in each sublevel before pairing up (p,d,f)

  29. Electron Configurations • Use Periodic Table to find e- configurations

  30. Electrons • Diamagnetism = all of e- are paired; not strongly affected by magnetic fields • Paramagnetism = has unpaired e-; strongly affected by magnetic fields • Valence e- = e- in outermost energy level • For Representative Elements 1A-8A, groups number=number of valence e- • Period number=energy level of valence e-

  31. Quantum Numbers

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