Understanding Atom Structure & Electromagnetic Radiation: Wavelength, Quantization, & Spec, Study notes of Chemistry

Download Understanding Atom Structure & Electromagnetic Radiation: Wavelength, Quantization, & Spec and more Study notes Chemistry in PDF only on Docsity! The Structure of Atoms Electrons Electromagnetic Radiation Radiation that consists of wave-like electrons and magnetic fields, including light, microwaves and x-rays. Wavelength (l): The distance between successive crests (or troughs) of a wave. The units of length are typically meters or nanometers. Frequency (n): The number of cycles the wave undergoes in one second. The unit is typically s-1 or 1/s and is called a hertz (Hz). The Ultraviolet Catastrophe When a metal is heated, EM radiation is emitted in wavelengths that depend on the temperature. As T increases, the red color brightens, and eventually white light is emitted. At the end of the 19th century, scientists could not explain the relationship between intensity and wavelength for radiation given off by a heated object. Our eyes detect the radiation that occurs in the visible region of the EM spectrum, but other wavelengths are also given off by the heated object. They predicted the intensity should increase continuously with decreasing wavelength, never reaching a maxima. The theories failed in the ultraviolet region (short l). Quantization and Planck’s Constant Max Planck described that emitted EM radiation was caused by vibrating atoms (oscillators) and each oscillator had a fundamental frequency of oscillation. Energy is quantized, such that only certain energies are allowed. E = nhn Planck’s constant: h = 6.626 x 10-34 J.s If the oscillator changes from a higher energy to a lower one, energy is emitted as EM radiation. DE = Dnhn If Dn = 1 (lower E level) E = hn Planck’s Equation Einstein and the Photoelectric Effect Electrons are ejected when light strikes the surface of a metal ONLY if the frequency of the light is high enough (threshold freq). If the frequency is too low, no electrons are ejected. If the frequency is high enough, increasing the intensity of the light causes more e- to be ejected (b/c more photons of light). Light has particle-like properties called photons, which are packets of energy. Photons are massless and have E proportional to the frequency of radiation. Metal only loses electrons if photons have enough energy. E = hn = hc  l Atomic Line Spectra Every element has a unique spectrum, which can be used to identify an element or determine how much is present. Why do excited gaseous atoms emit light of only certain frequencies? Balmer Equation John Balmer found an equation that could calculate the wavelength of red, green and blue lines in the visible emission spectrum of hydrogen. = 1  n1 2 1  l R - 1  n2 2 Rydberg constant: R = 1.0974 x 107 m-1 UV series: n1 = 1, n2 = 2, 3, 4… Vis series: n1 = 2, n2 = 3, 4, 5… IR series: n1 = 3, n2 = 4, 5, 6… For hydrogen, n1 > 2 Bohr Model of the Atom Niels Bohr proposed a planetary structure for the atom where electrons move in a circular orbit around the nucleus, much like a planet revolves around the sun. Classic physics in the early 20th century described that electrons moving in the positive electric field of the nucleus eventually lose energy and crash into the nucleus. (If this was true, matter would self-destruct!) Change in Energy Levels If an electron moves from one energy level to another, energy must be either absorbed or released. If an electron is excited from n = 1 to n = 2: DE = Ef – Ei = (-NARhc/2 2) – (-NARhc/1 2) = (0.75)NARhc = 984 kJ/mol Therefore moving an electron from the first to the second energy state requires an input of 984 kJ/mol of atoms. The electron can only be excited at this precise amount of energy. (QUANTIZATION!) Moving an electron from a higher energy state to lower energy state leads to the release or “emission” of energy. ie n = 2 n = 1 DE = -984 kJ/mol NA: Avogadro’s Number Change in Energy Levels: Qualitative Classify each of the hydrogen transitions: A B C D Absorption: (A,C) Electron goes from a lower to higher energy level. Emission: (B, D) Electron goes from a higher to lower energy level. Ionization: (C) Electron goes from the lowest energy level to infinity. The electron is moved completely away from the nucleus. Bohr Model of Hydrogen Electrons are excited to higher energy levels and absorb energy. The electrons can then return to any lower energy level (either directly or in a series of steps) and release energy (as photons of EM radiation). DE = 1  nfinal 2 -NARhc 1  ninitial 2 - The change in energy for the emission lines in excited hydrogen atoms can be calculated as follows: Heisenberg’s Uncertainty Principle For an electron in an atom, it is impossible to determine accurately both its position and energy. If we know the energy of the electron with a small uncertainty, there is a large uncertainty on its position. Chemists predict the approximate location of an electron. Schrodinger’s Wave Function Schrodinger developed a model for electrons in atoms using the idea that electrons could behave as waves. The model uses mathematical equations of wave motion to generate wave functions (Y). Each describes an allowed energy state of an electron. Schrodinger’s Wave Function • An electron can be described as a standing wave. • Only certain vibrations are possible; n(l/2) • Vibrations are quantized where n is a quantum number • Nodes occur at points of zero amplitude Schrodinger’s Wave Function • An electron in 3-D space requires three quantum numbers • n, l and ml (all are integers) • Each Y is associated with an allowed energy level • Y2 is proportional to the probability of finding an electron at a given point • Orbitals describe the region of space where an electron of a given energy is most likely to be located Quantum Numbers ml, Magnetic Quantum Number: ml = 0, ±1, ±2…, ±l • Orientation in space of orbitals in a subshell • l determines ml • Number of orbitals in a subshell = 2l + 1 Any atomic orbital is specified by three quantum numbers. Value of l Subshell Label Value of ml 0 s 0 1 p -1, 0 +1 2 d -2, -1, 0, +1, +2 3 f -3, -2, -1, 0, +1, +2, +3 Shells and Subshells n2 = number of orbitals in the shell The First Electron Shell, n = 1 • n = 1, l = 0, ml = 0 (one s orbital denoted 1s) • In the shell closest to the nucleus, only a single orbital exists The Second Electron Shell, n = 2 • n = 2, l = 0, ml = 0 (one s orbital denoted 2s) • n = 2, l = 1, ml = -1, 0, +1 (three p orbitals denoted 2p) • The p orbitals have the same shape but different orientation The Third Electron Shell, n = 3 • n = 3, l = 0, ml = 0 (one s orbital denoted 3s) • n = 3, l = 1, ml = -1, 0, +1 (three p orbitals denoted 3p) • n = 3, l = 2, ml = -2, -1, 0, +1, +2 (five d orbitals denoted 3d) n Possible value of l Subshell Designation Possible value of ml # of orbitals in subshell Total # of orbitals in a shell 1 0 1s 0 1 1 2 0 2s 0 1 1 2p -1, 0, 1 3 4 3 0 3s 0 1 1 3p -1, 0, 1 3 2 3d -2, -1, 0, 1, 2 5 9 4 0 4s 0 1 1 4p -1, 0, 1 3 2 4d -2, -1, 0, 1, 2 5 3 4f -3, -2, -1, 0, 1, 2, 3 7 16 Relationship Among Quantum Numbers The Shape of Atomic Orbitals No more than 2 electrons can be held in an orbital. p orbitals: (l = 1) • Shape resembles a dumbbell (e- spends time in both lobes) • A nodal surface passes through the nucleus (no probability of finding an electron) • Three p orbitals are in a subshell, all have the same shape but lie along a different axis (px, py, pz) The Shape of Atomic Orbitals No more than 2 electrons can be held in an orbital. d orbitals: (l = 2) • Four orbitals shaped like clover-leafs, one shaped differently • For the clover-leaf orbitals, two nodal surfaces pass through the nucleus so there are four regions of e- density • The orbitals lie along two planes The Shape of Atomic Orbitals No more than 2 electrons can be held in an orbital. f orbitals: (l = 3) • Seven total orbitals • Three nodal planes • Electron density lies in eight regions of space Magnetism Diamagnetism: • All electrons in an element or compound are paired • When placed in a magnetic field, the substance experiences slight repulsion Paramagnetism: • At least one electron in an element or compound is unpaired • In the absence of a magnetic field, the electrons are randomly oriented • When placed in a magnetic field, the substance experiences attraction Magnetism • A form of paramagnetism where the magnetic effect is greatly enhanced • In the absence of a magnetic field, the electrons align themselves in the same direction • When placed in a magnetic field, the substance experiences a strong attraction • Very few metals (such as iron, cobalt and nickel) exhibit these properties Ferromagnetism: