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PERIODICITY and ATOMIC STRUCTURE

PERIODICITY and ATOMIC STRUCTURE. Electromagnetic Radiation (1-3) Quantum Mechanics (4-8) Electronic Configurations (9-15). STRUCTURE OF ATOM. Electrons and Protons observed and characterized Models of the Atom - pre-20th Century “Raisin Pudding”: J.J. Thomson

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PERIODICITY and ATOMIC STRUCTURE

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  1. PERIODICITY and ATOMIC STRUCTURE Electromagnetic Radiation (1-3) Quantum Mechanics (4-8) Electronic Configurations (9-15)

  2. STRUCTURE OF ATOM • Electrons and Protons observed and characterized • Models of the Atom - pre-20th Century • “Raisin Pudding”: J.J. Thomson • Small positive nucleus surrounded by a lot of empty space through which the electrons are dispersed: Rutherford (1911) • Atom nucleus also includes neutrons: Chadwick (1932)

  3. LIGHT or EM RADIATION: WAVE • Classical science (CS) considered electromagnetic radiation (EM) or light as a wave based on observations of diffraction, reflection, interference, refraction. • Form of energy • Wavelength, λ = c/ν nm • Frequency, ν Hz = 1/s • Speed, c = λν 3.00E+08 m/s

  4. ELECTROMAGNETIC SPECTRUM • Light or electromagnetic radiation spans many orders of magnitude in E, ν, and λ. • Figure 5.3 • Visible: ROY G. BIV 400-800 nm • At lower E and ν , λ increases: Infrared, microwave, radiowave • At higher E and ν , λ decreases: Ultraviolet, X-rays, gamma-rays

  5. ELECTRONS, PROTONS, NEUTRONS: PARTICLES • Classical science considered these subatomic particles to be particles with mass (m), velocity (v) and momentum (mv) • It was assumed that an object was either a wave (light) or a particle (electron).

  6. FROM CLASSICAL TO QUANTUM THEORY • From the late 1800’s to the 1920’s, many experimental observations that could not be explained by Classical Science/Theories were recorded. These led to the development of Quantum Mechanics • atomic line spectra (Balmer,1885) • properties of radiation from heated solid or blackbody (Planck, 1900)

  7. TRANSITION (2) • photoelectric effect (Einstein, 1905) • heat capacity of solids • electron diffraction • As scientists worked to understand these exptal results, several conclusions emerged: • Electrons have WAVE and particle properties. • Light has PARTICLE and wave properties. • deBroglie Eqn expresses this: λ = h/mv

  8. ATOMIC LINE SPECTRA • CS: Rutherford model of the atom. • Expt: When atoms are excited, they return to their stable states by emitting light. This light can be recorded to produce an atomic spectrum. Early experiments showed that the spectra consists of lines and that atoms from different elements gave different line spectra. Fig 5.6 • What do these spectra tell us about the structure of the atom?

  9. ATOMIC LINE SPECTRA (2) • Balmer measured the emission spectrum of H and fit the observed wavelengths of the emitted light to an equation: • ν = Rc (1/22 – 1/n2) where R = Rydberg constant = 1.097E-2 1/nm • The emission lines of the H atom in other regions of the EM spectrum fit the Balmer-Rydberg Eqn: ν = Rc (1/m2 – 1/n2) for n > m; n and m are integers. This is an empirical eqn.

  10. BOHR ATOM (Fig 5.14) • The Rutherford model could not explain these results, but Bohr’s “planetary” model of the atom could (1914) . • This model led to quantized electronic energy levels and to an eqn consistent with the Balmer-Rydberg Eqn. • The energy of an electron in the nth energy level is quantized and equals • En = - hcRZ2/n2 where n = 1, 2, 3...

  11. BOHR ATOM (2) • Then when an electron goes from one quantized level (n) to another (m), light is emitted or absorbed with a wavelength defined by 1/λ = R(1/m2 - 1/n2 ) • The Bohr atom is the basis for the modern theory of the atom but it has limitations. For example, it is only accurate for 1-electron atoms and ions.

  12. QUANTUM MECHANICS (Schrodinger, 1926) • The QM model of the atom replaced the Bohr model. This model is based on electron’s wave properties. • The electron in an atom was viewed as a standing wave around the nucleus. • These standing waves (Ψ) are called wave functions and are interpreted as the allowed atomic orbitals for electrons in an atom.

  13. QUANTUM MECHANICS (2) • The goal of QM is to solve the Schrodinger Eqn and find Ψ plus its associated (quantized) energy. • Ψ2 is related to the probability of finding an electron at a particular (x,y,z) location. • Heisenberg Uncertainty Principle (1927) states that we cannot know the position and momentum of an electron (considered a wave) exactly: • Δx Δ(mv) ≥ h/4 π

  14. QUANTUM MECHANICS (3) • When the Ψs are found, each one is defined by three quantum Numbers. • A set of Ψs lead to atomic electronic configurations. • QM is the basis for understanding chemical bonding and molecular shapes (Chap.7), chemical reactions, physical and chemical properties (Chap. 6).

  15. ATOMIC ORBITALS AO) and QUANTUM NUMBERS (QN) • AOs are wavefunctions, Ψ, and are characterized by QNs which are related to each other. • These relationships determine the identity and number of AOs in an atom. • Principal QN, n = 1, 2, 3…(K, L, M...shell); determines energy (quantized) and size of atomic orbital.

  16. AOs and QNs (2) • Angular momentum QN, ℓ = 0, 1, 2…n-1 (s, p, d…) subshell; determines shape of atomic orbital. For each n value, there are n ℓvalues. • Magnetic, mℓ= - ℓ, …-2, -1, 0, +1, +2, …+ ℓ; determines spatial orientation of orbital. For each ℓ value, there are 2ℓ + 1 mℓvalues. • Spin, ms = +1/2, -1/2; determines orientation of electron spin axis.

  17. AOs and QNs (3) • There are relationships (limitations) between four quantum numbers (Table 5.2) • For the H atom and other one-electron atoms, all AOs with the same n value have the same energy. This is called energy degeneracy. (Fig 5.9)

  18. ATOMS WITH ≥2 ELECTRONS • We will apply the 1-electron results to the many-electron atom. • For the 1-electron atom, AO energy depends only on n and as n increases, energy increases (becomes less positive). So AOs can be ordered from low to high energy: 1s < [2s, 2p] < [3s, 3p, 3d]...

  19. MULTI-ELECTRON ATOMS • For the many-electron atom, energy depends on n and ℓ: 1s < 2s < 2p < 3s < 3p, etc. See Fig 5.9 • This is due to electron-electron repulsions. • Also, because of the difference in AO shape for different ℓ values, electrons with the same n but in different subshells experience different attractive forces to the nucleus. • Zeff = Zactual – electron shielding

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