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Chapter 10: Modern atomic theory Chemistry 1020: Interpretive chemistry Andy Aspaas, Instructor 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? 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 Electromagnetic spectrum Types of electromagnetic radiation • • • • • • • Radio waves: low frequency and energy Microwaves Infrared Visible Ultraviolet X-rays Gamma rays: high frequency and energy 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 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 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 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 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 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 Orbitals • Solutions to wave equations give regions of high probability for finding electrons – Called orbitals – 90% probability of finding an electron – 3-dimensional shape 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 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 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 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) 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