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Chapter 11 Modern Atomic Theory. Electromagnetic Radiation. Light is a form of electromagnetic (EM) radiation All forms of EM radiation are types of kinetic energy See page 306. EM Radiation. EM radiation can be described as traveling in waves or as packets of energy called photons.
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Electromagnetic Radiation • Light is a form of electromagnetic (EM) radiation • All forms of EM radiation are types of kinetic energy • See page 306
EM Radiation • EM radiation can be described as traveling in waves or as packets of energy called photons
EM Radiation • Describe each form of EM radiation by its: • Wavelength • Frequency • Energy of its photon
All forms of EM radiation travel at the speed of light (c) Wavelength x frequency = speed of light
EM Radiation • The longer the wavelength the: • lower the frequency and the lower the energy of the EM radiation • The shorter the wavelength the: • higher the frequency and the higher the energy of the EM radiation
EM Radiation • Compare the wavelength, frequency, and energy of: • Ultraviolet light and infrared light
EM Radiation • The higher the energy of the EM radiation the more damaging it is to living tissue. • CONSIDER GAMMA RAYS, X-RAYS, AND UV LIGHT
Emission of Energy by Atoms • Thanks to the work of Bohr we know that: • When atoms are energized by an input of energy their electrons are excited (energized) • When excited electrons return to lower energy states they emits energy in the form of light. • Emits photons of energy • Energy of the photons emitted depends upon how excited the electron was.
Bohr Experiment (~1911) • Bohr excited hydrogen atoms by running electricity through a tube of hydrogen gas. • The gas gave off a pink light.
Bohr Experiment (~1911) • Bohr aimed a beam of the pink light at a “prism” • Found the pink light generated a line spectrum not a continuous spectrum • Line spectrum – specific colors of light observed • Continuous spectrum – all colors of light present
Bohr Experiment (~1911) • He observed 4 bands of color: (pg 310) • Purple (410 nm) • Blue (434 nm) • Green (486 nm) • Red (656 nm)
Bohr Experiment (~1911) • He then calculated the energy of each color of light • _____________ was the highest energy and _____________ was the lowest energy.
Bohr’s Interpretation of the Data • The electrons circle the nucleus in orbits of specific energies. • Larger orbits are of higher energy than smaller orbits
Bohr’s Interpretation of the Data • Electrons are always in one of the circular orbits. • The electricity excites electrons and allows them to move to higher energy orbits.
Bohr’s Interpretation of the Data • When the excited electrons return to lower energy orbits they emit energy in the form of light. • The difference in energy levels between the orbits determines the energy of the light given off.
Bohr’s Interpretation of the Data • Because only certain wavelengths of light are emitted only certain energy changes are occurring • 4 bands of light, therefore 4 possible energy changes are possible • Say energy levels are quantized, meaning only specific energy levels are available to electrons.
Bohr’s Model • When Bohr’s mathematical approach was applied to other elements it didn’t work. • Bohr’s model of the atom has been revised to replace the circular orbits with “wave mechanical model” of the atom
Modern Atomic Structure • Still picture electrons to be at specific energy levels, but no longer picture them as traveling in circular orbits. • The current model of the atom locates electrons in orbitals.
Modern Atomic Structure • Orbitals are very different than Bohr’s orbits • Read the firefly analogy on page 313
Orbitals • Each orbital is of a specific energy, size, and shape • Each orbital can hold a maximum of 2 electrons of opposite spin (Pauli exclusion principle)
Orbitals • The exact path of an electron in an orbital is not known. • Heisenberg uncertainty principle states that it is impossible to determine the location and path of an electron at the same time
Orbitals • Orbital shapes describe the region in space where an electron will be found 90% of the time.
Modern Atomic Theory • Atoms have specific energy levels in which electrons may be found. • Called Principal Energy Levels (PEL) • PEL farther from the nucleus are larger and of higher energy. • Assign a number (n) to each PEL • See board
Modern Atomic Theory • Within each PEL are sublevels • Sublevels are named: s, p, d, and f • The larger the PEL the more sublevels it contains
Modern Atomic Theory • Sublevels contain orbitals.
Describing Orbitals • See pages 314/315 for diagrams of the orbitals • S orbitals are spherical • The 3 p orbitals are shaped
Putting it All Together • Unless they are excited, electrons always occupy the lowest energy orbital with room. • Electrons enter orbitals of a given sublevel one at a time before pairing up (Hund’s rule) • Consider 2 electrons in a p sublevel:
The fun part! • Our goal is to write the following for atoms and ions: • Electron configuration • Box/energy diagram • Lewis dot symbol • Our goal is also to: • Identify core and valence electrons
Terms • Electron configuration – shows the number of electrons in each sublevel • Box/energy diagram – shows the number of electrons in each orbital • Orbitals are shown as boxes • electrons are shown as arrows
Terms • Lewis Dot Symbol – shows the valence electrons as dots around the symbol for the element • Maximum of 2 electrons per side of the symbol
Terms • Valence electrons – all the electrons in the highest occupied PEL • Valence electrons are the ones involved in bonding • Core electrons – all electrons not considered valence electrons
You will be given this. • Sublevels listed from lowest energy to higher energy: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p……….