Atomic Structure

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Nice Atomic structure

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Atomic Structure

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All waves have a characteristic wavelength, l, and amplitude, A. Frequency, n, of a wave is the number of cycles which pass a point in one second. Speed of a wave, c, is given by its frequency multiplied by its wavelength: For light, speed = c = 3.00x108 m s-1 . A Brief History of Time The Wave Nature of Light

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The Wave Nature of Light

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The Wave Nature of Light

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The Wave Nature of Light

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Planck: energy can only be absorbed or released from atoms in certain amounts called quanta. The relationship between energy and frequency is where h is Planck’s constant ( 6.626  10-34 J s ) . Quantized Energy and Photons

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The Photoelectric Effect and Photons Einstein assumed that light traveled in energy packets called photons. The energy of one photon is: Quantized Energy and Photons

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Nature of Waves: Quantized Energy and Photons

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Line Spectra Radiation composed of only one wavelength is called monochromatic. Radiation that spans a whole array of different wavelengths is called continuous. White light can be separated into a continuous spectrum of colors. Note that there are no dark spots on the continuous spectrum that would correspond to different lines. Line Spectra and the Bohr Model

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Bohr Model Colors from excited gases arise because electrons move between energy states in the atom. (Electronic Transition) Line Spectra and the Bohr Model

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Bohr Model Since the energy states are quantized, the light emitted from excited atoms must be quantized and appear as line spectra. After lots of math, Bohr showed that where n is the principal quantum number (i.e., n = 1, 2, 3, … and nothing else). Line Spectra and the Bohr Model

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Bohr Model We can show that When ni > nf, energy is emitted. When nf > ni, energy is absorbed Line Spectra and the Bohr Model

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Bohr Model Line Spectra and the Bohr Model CyberChem (Fireworks) video Mathcad (Balmer Series)

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Line Spectra and the Bohr Model: Balmer Series Calculations

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Limitations of the Bohr Model Can only explain the line spectrum of hydrogen adequately. Can only work for (at least) one electron atoms. Cannot explain multi-lines with each color. Electrons are not completely described as small particles. Electrons can have both wave and particle properties. Line Spectra and the Bohr Model

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Knowing that light has a particle nature, it seems reasonable to ask if matter has a wave nature. Using Einstein’s and Planck’s equations, de Broglie showed: The momentum, mv, is a particle property, whereas  is a wave property. de Broglie summarized the concepts of waves and particles, with noticeable effects if the objects are small. The Wave Behavior of Matter

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The Uncertainty Principle Heisenberg’s Uncertainty Principle: on the mass scale of atomic particles, we cannot determine exactly the position, direction of motion, and speed simultaneously. For electrons: we cannot determine their momentum and position simultaneously. If Dx is the uncertainty in position and Dmv is the uncertainty in momentum, then The Wave Behavior of Matter

Energy and Matter : 

Energy and Matter E = m c2

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Schrödinger proposed an equation that contains both wave and particle terms. Solving the equation leads to wave functions. The wave function gives the shape of the electronic orbital. [“Shape” really refers to density of electronic charges.] The square of the wave function, gives the probability of finding the electron ( electron density ). Quantum Mechanics and Atomic Orbitals

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Quantum Mechanics and Atomic Orbitals Solving Schrodinger’s Equation gives rise to ‘Orbitals.’ These orbitals provide the electron density distributed about the nucleus. Orbitals are described by quantum numbers.

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Orbitals and Quantum Numbers Schrödinger’s equation requires 3 quantum numbers: Principal Quantum Number, n. This is the same as Bohr’s n. As n becomes larger, the atom becomes larger and the electron is further from the nucleus. ( n = 1 , 2 , 3 , 4 , …. ) Azimuthal Quantum Number, . This quantum number depends on the value of n. The values of  begin at 0 and increase to (n - 1). We usually use letters for  (s, p, d and f for  = 0, 1, 2, and 3). Usually we refer to the s, p, d and f-orbitals. Magnetic Quantum Number, m. This quantum number depends on  . The magnetic quantum number has integral values between -  and +  . Magnetic quantum numbers give the 3D orientation of each orbital. Quantum Mechanics and Atomic Orbitals

Quantum Numbers of Wavefuntions : 

Quantum Numbers of Wavefuntions

Quantum Mechanics and Atomic Orbitals : 

Quantum Mechanics and Atomic Orbitals

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Orbitals and Quantum Numbers Quantum Mechanics and Atomic Orbitals

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The s-Orbitals Representations of Orbitals

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The p-Orbitals Representations of Orbitals

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d-orbitals

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Many-Electron Atoms Orbitals and Their Energies Orbitals CD

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Electron Spin and the Pauli Exclusion Principle Many-Electron Atoms

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Electron Spin and the Pauli Exclusion Principle Since electron spin is quantized, we define ms = spin quantum number =  ½. Pauli’s Exclusions Principle: no two electrons can have the same set of 4 quantum numbers. Therefore, two electrons in the same orbital must have opposite spins. Many-Electron Atoms

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Figure 6.27 Figure 6.27 Orbitals CD

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Figure 6.28 Orbitals CD

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Many-Electron Atoms Orbitals and Their Energies Orbitals CD

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Electron Configurations

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Metals, Nonmetals, and Metalloids Metals Figure 7.14

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Two Major Factors: principal quantum number, n, and the effective nuclear charge, Zeff. Periodic Trends

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Figure 7.5: Radius video Clip

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Figure 7.6

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Figure 7.10 IE clip

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Figure 7.9

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Electron Affinities Electron affinity is the opposite of ionization energy. Electron affinity: the energy change when a gaseous atom gains an electron to form a gaseous ion: Cl(g) + e-  Cl-(g) Electron affinity can either be exothermic (as the above example) or endothermic: Ar(g) + e-  Ar-(g)

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Figure 7.11: Electron Affinities

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Group Trends for the Active Metals Group 1A: The Alkali Metals

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Group Trends for the Active Metals Group 2A: The Alkaline Earth Metals

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Group Trends for Selected Nonmetals Group 6A: The Oxygen Group

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Group Trends for Selected Nonmetals Group 7A: The Halogens

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Group Trends for the Active Metals Group 1A: The Alkali Metals Alkali metals are all soft. Chemistry dominated by the loss of their single s electron: M  M+ + e- Reactivity increases as we move down the group. Alkali metals react with water to form MOH and hydrogen gas: 2M(s) + 2H2O(l)  2MOH(aq) + H2(g)

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Group Trends for the Active Metals Group 2A: The Alkaline Earth Metals Alkaline earth metals are harder and more dense than the alkali metals. The chemistry is dominated by the loss of two s electrons: M  M2+ + 2e-. Mg(s) + Cl2(g)  MgCl2(s) 2Mg(s) + O2(g)  2MgO(s) Be does not react with water. Mg will only react with steam. Ca onwards: Ca(s) + 2H2O(l)  Ca(OH)2(aq) + H2(g)

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Atomic Structure