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Understanding Modern Atomic Theory in Chemistry

Delve into Rutherford's atomic theory, electromagnetic radiation, atomic emission spectra, Bohr's model, wave mechanical model, orbitals, Pauli exclusion principle, electron configurations, valence electrons, ionization energy, and the periodic table in chemistry. Explore the fundamental principles shaping our understanding of atomic structure.

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Understanding Modern Atomic Theory in Chemistry

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  1. Chapter 10: Modern atomic theory Chemistry 1020: Interpretive chemistry Andy Aspaas, Instructor

  2. Rutherford’s atom • Recall Rutherford’s atomic theory • Positively charged nucleus • Surrounded by negatively charged electrons • Unanswered questions • How are electrons arranged • How do they move?

  3. Electromagnetic radiation • Electromagnetic radiation: energy transmitted by waves, “radiant energy” • Wavelength: distance between peaks of these waves • Different forms of electromagnetic radiation have different wavelengths

  4. Electromagnetic spectrum

  5. Types of electromagnetic radiation • Radio waves: low frequency and energy • Microwaves • Infrared • Visible • Ultraviolet • X-rays • Gamma rays: high frequency and energy

  6. Energy and electromagnetic radiation • The shorter the wavelength, the higher the energy transmitted • Blue light: shorter wavelength: higher frequency: higher energy • Red light: longer wavelength: lower frequency: lower energy

  7. Wave calculations • Velocity = c = speed of light • 2.997925 x 103 m/s • All types of light energy travel at same speed • Amplitude = A = height of wave, brightness of light • Wavelength = = distance between peaks • Frequency =  = number of waves that pass a point in a given amount of time • Generally measured in Hertz (Hz) • 1 Hz = 1 wave/sec = 1 sec-1 • c = x

  8. Planck’s nuclear theory • Light energy behaves as particles in certain situations • Each particle of light (a photon) has a certain fixed amount of energy • Energy of photon is directly proportional to frequency of the light • Higher frequency = more energy in photon

  9. Atomic emission spectra • Atoms that gain extra energy will release that energy in the form of light • Light is given off in very specific wavelengths • Different atoms give off different characteristic wavelengths of light when excited • Line spectrum: shows wavelengths of light that are emitted • Only certain wavelengths are given off, so only specific amounts of energy can be absorbed or given off for any one type of atom • Atoms are “quantized” - only specific energy levels

  10. Bohr’s model of the atom • Explains line spectrum of hydrogen • Energy of atom is related to distance of electron from nucleus • Electrons can “jump” to different possible orbits around nucleus • Gain in energy: electron jump to higher quantum level - “excited state” • Lines in spectrum correspond to difference in energy levels • Ground state: minimum energy level • But, only explains hydrogen atom behavior • Plus, electrons do not have simple circular orbits

  11. Wave mechanical model of the atom • Electrons can be treated as waves (in the same way that light can also be treated as particles) • Mathematics can calculate the probability densities of finding an electron in a particular region of the atom • Schrödinger equation - cannot predict location of any one particle, only probability of it being a certain place

  12. Orbitals • Solutions to wave equations give regions of high probability for finding electrons • Called orbitals • 90% probability of finding an electron • 3-dimensional shape

  13. Orbitals and energy levels • Principal energy level (n) = how much energy the electrons in the orbital have • Higher values mean higher energy and farther average distance from nucleus • Each principal energy level has n sublevels • Different shape and energy • Named s, p, d, f • Each sublevel has 1 or more orbitals • s = 1 orbital, p = 3, d = 5, f = 7

  14. Pauli exclusion principle • No orbital may have more than 2 electrons • Electrons in same orbital must have opposite spin • s holds 2 electrons • p holds 6 electrons • d holds 10 electrons • f holds 14 electrons

  15. Electron configurations • Hydrogen electron configuration: 1s1 • Superscript indicates number of electrons in orbital • Helium: 1s2 • Follow the periodic table: row number = principal energy level (first number in electron configuration) • Column and section determine which sublevel (s, p, d, f) is filled

  16. Valence electrons • Valence electrons: only those in outermost energy level - determine most of an atom’s reactivity properties • Can indicate Na electron configuration as • 1s22s22p63s1 or [Ne]3s1 (using nearest Noble gas with smaller atomic number than the atom)

  17. Atomic properties and the periodic table • Ionization energy: energy required to remove an electron from an atom • Decreases down a group (less energy required to remove electron) • Increases across a period (more energy required to remove an electron) • Atomic size • Increases down group to account for greater mass • But decreases across period because more electrons mean more attraction to the nucleus

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