AP CHEMISTRY Atomic Structure and Electrons

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AP CHEMISTRYAtomic Structure and Electrons

• Ch. 7 sec 1-9

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Light and Quantized Energy

• Chemists found Rutherfords 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.

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Light and Quantized Energy

• Analysis of the emitted light revealed that an
  elements 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.

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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.

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The Electromagnetic Spectrum
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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 lights
  interactions with matter.

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The Quantum Concept

• In 1900, the German physicist Max Planck
  (18581947) began searching for an explanation as
  he studied the light emitted from heated objects.

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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.

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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.

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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.

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Atomic Emission Spectra

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

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Atomic Emission Spectra
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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
  elements atomic emission spectrum means that
  only certain specific frequencies of light are
  emitted.

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Bohr Model of the Atom

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

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Energy States of Hydrogen

• Building on Plancks and Einsteins 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.

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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
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Energy States of Hydrogen

• Bohr went even further with his atomic model by
  relating the hydrogen atoms 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.

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Energy States of Hydrogen
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Hydrogens 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.

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Hydrogens 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.

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Hydrogens 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.

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Hydrogens Line Spectrum

• The four electron transitions that account for
  visible lines in hydrogens atomic emission
  spectrum are shown.

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The Heisenberg Uncertainty Principle

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

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The Heisenberg Uncertainty Principle

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

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Mathematic Equations
cln c speed of light (3.0 x 108m/s) l
wavelength (m) n frequency (s-1 or Hz)
Ehn Eenergy (J or kgm2/s2) hPlancks constant
(6.63x10-34 Js or kgm2/s) nfrequency (s-1)
Combining them Ehc/l
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Mathematic Equations (contd)
deBroglie equation lh/mv lwavelength
(m) h6.63x10-34 kgm2/s mmass (kg) vvelocity
(m/s)
pmv pmomentum (kgm/s) mmass (kg) vvelocity
(m/s)
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Energy levels, sublevels, orbitals

• Energy levelsclouds or shells around nucleus
  (n1,2,3)
• Sublevelsfound inside energy levels (s,p,d,f)
• Atomic orbitalsfound within sublevels
• s 1 orbital (sphere)
• p 3 orbitals (dumbell)
• d 5 orbitals (p. 313)
• f 7 orbitals
• 2 e- max per orbital

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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
• Hunds Rule e- occupy one orbital in each
  sublevel before pairing up (p,d,f)

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Electron Configurations

• Use Periodic Table to find e- configurations

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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
  numbernumber of valence e-
• Period numberenergy level of valence e-

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Quantum Numbers