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Inorganic Chemistry Chem315L

Inorganic Chemistry Chem315L. Inorganic Chemistry. For a long time “inorganic chemistry” was synonymous with “general chemistry,” but since the 1950’s it has moved away from that focus.

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Inorganic Chemistry Chem315L

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  1. Inorganic ChemistryChem315L

  2. Inorganic Chemistry • For a long time “inorganic chemistry” was synonymous with “general chemistry,” but since the 1950’s it has moved away from that focus. • The current emphases in the field of Inorganic Chemistry involve transition metals and coordination complexes. • Solid State • Organometallic • Bioinorganic • also Main Group

  3. Inorganic vs Organic • One of the major differences in the study of Inorganic vs Organic Chemistry is in the relative emphasis placed on structure and reactivity. • The basic structural elements in organic chemistry are relatively simple. • Because of this, Organic Chemistry tends to emphasize the various mechanisms of reactions as these are the more exciting aspects of the subject. • In contrast, Inorganic Chemistry has a wide variety of structural types, and even for a given element there are many factors to consider. • Thus, Inorganic Chemistry has been more concerned with the “static” structures of reactants and products than the way they interconvert. • Another difference: COLOR! C C=C ― C≡C―

  4. Spectroscopy • When white light is passed through a prism, split into a continuous spectrum (ROYGBIV). • Electromagnetic radiation behaves according to: = 3.0 x 108 m/s

  5. Spectroscopy • In the mid-1800’s it was observed that when Hydrogen is excited and examined through a prism, the emitted light consists of a small number of narrow bands of color, separated by dark. • λ = 656, 486, 434, 410 & 397 nm • No one could explain why these particular wavelengths were observed until 1885, when Balmer developed a formula: • Balmer had no physical explanation for this formula; it just worked. Formula was modified when lines were found in other areas of the electromagnetic spectrum. • Note: Your book uses a different version of this equation along with a different value for RH. E=hν=hc/λ “our” RH x h x c = “book” RH (2.179 x 10-18 J) • 1 = RH1 1 • λ ns2 nl2 Lyman ns = 1 UV Balmer ns = 2 Vis Paschen ns = 3 IR ( ) - • RH = Rydberg constant = 1.097 x 107 m-1 • nl = small # > ns

  6. Development of Atomic Theory • Thomson: Raisin Pudding Model • Rutherford: Nuclear Atom • Newtonian Physics says the electron in the Rutherford Model should radiate energy, and as it does it should lose energy and spiral into the nucleus. • it should thus radiate a continuous spectrum of energy. • neither of the above is observed. • Bohr: Quantized Model • Neils Bohr developed a new model of H-atom by making the crucial assumption that not all of the classical physical laws apply to the atomic scale.

  7. Development of Atomic Theory • Bohr’s Postulates: • 1. Energy of H-atom is not a continuous function; it can only exist in a limited number of discreet energy states---the energy is quantized! • allowed states = stationary states • lowest = ground state n = 1 • others = excited states n = 2, 3, 4... • 2. The atom does not radiate energy as long as it remains in a stationary state • absorbs E = goes from lower to higher state • emits E = goes from higher to lower state • 3. The different states correspond to different orbits of the electron. • 4. Energy emitted or absorbed is quantized because the energy states are quantized: E = h

  8. Bohr Model: Energy States for Hydrogen Bohr (Quantized) model can explain line spectra for hydrogen.

  9. Development of Atomic Theory • Meanwhile, deBroglie proposed his hypothesis of the wave-particle duality of the electron, that the electron has properties of both. • Heisenberg proposed the indeterminancy principle, which stated that even if the electron were strictly a particle, it could not be described in as precise a manner as the Bohr model (i.e. orbits). • In 1926 Schrödinger proposed a wave equation which describes the behavior of a subatomic particle the same way macroscopic particles are described by classical mechanics. δ2Ψδ2Ψ δ2Ψ8π2m δx2δy2δz2 h2 + + + (E – V)Ψ = 0 where Ψ = wave function x, y, & z = spacial coordinates m = mass h = Plack’s constant E = total evergy V = potential energy

  10. Development of Atomic Theory δ2Ψδ2Ψ δ2Ψ8π2m δx2δy2δz2 h2 • The solution to the wave equation is the wavefunction, Ψ, which corresponds to the amplitude of a wave. • The probability of finding a particle at a particular point in space is proportional to the square of its wavefunction at that point, Ψ2. This is similar to the intensity of a light wave, which is given by the amplitude squared. • The solutions describe orbitals, which are 3-dimensional plots of all points in space where the probability of finding the electron is some constant (normally 90%). • We won’t solve the Schrödinger equation here, but we will use the solutions. • It is found, for the hydrogen atom, that the set of solutions can be described by three quantum numbers, n, l and ml. + + + (E – V)Ψ = 0

  11. Allowed values for Quantum Numbers • Principle Quantum Number, n • can have any integer value greater than zero;n = 1, 2, 3, 4 … • as n increases, E increases • also, the relative size of the orbital increases with increasing n (i.e. the electron is further away from nucleus; can’t say anything about radius). • Orbital Angular Momentum Quantum Number, l • l can any have integral value from 0 to n-1 • determines the shape of an orbital. • the energy of the orbital depends on l only in a multi-electron case; for electrons with same n, energy of l = 1< l = 2 < l = 3 ... • Magnetic Quantum Number, ml • allowed values range from –l to +1 • determines the orientation of the orbital in space. l = 0 1 2 3 s p d f

  12. The Schrödinger Equation • Each solution to the Schrödinger equation (eigenfunction) has a unique set of n, l, and ml values and represent different orbitals of the H-atom. • If the complete wavefunction were plotted, would need 4-D graph paper with coordinates for the 3 spatial dimensions and one for the value of the wavefunction. • It is common to break the wavefunction down into 3 parts: Ψ(rθφ) = R(r) Θ(θ) Φ(φ) • where R(r) gives the dependence ofΨ upon the distance from the nucleus, and Θ(θ) Φ(φ) give the angular dependence (polar coordinates).

  13. The Radial Wave Function • Solutions for R(r) in Fig 2.3. Solutions squared in Fig 2.5. • These show the probability of finding the electron at certain distances away from the nucleus and also the most probable (but not only) location for finding the electron in an orbital.

  14. Angular Wave Functions • If we solve for ΘΦ we get the angular dependence, or shape, of the orbital. • But interested in probability, so need to examine Θ2Φ2. For s-orbital shape remains the same, but for p-orbital the plot becomes elongated.

  15. The Wave Function Ψ(rθφ) = R(r) Θ(θ) Φ(φ) • It is the presence of different values of the various quantum numbers in each of the functions that makes one different and distinct from another. • So, we have a situation where the electron is “smeared” about the nucleus in a way that varies with distance (as governed by the radial prortion of the wave function) and in different angular patterns (as governed by the angular portion of the wave function).

  16. Energy Level Diagram for Hydrogen l 0 1 2 3 ml 0 -1 0 +1 -2 -1 0 +1 +2 -3 -2 -1 0 +1 +2 +3 4 4s 4p 4d 4f 3 3s 3p 3d ---same shell = same value of n --same sub-shell = same value of l --same energy = degenerate 2 2s 2p 1 n 1s

  17. The Polyelectronic Atom • The quantum mechanics we’ve discussed so far have dealt with only the H-atom, the only atom for which the Schrödinger equation has been solved exactly. • The next simplest atom, He, consists of a nucleaus and two electrons, so there are now 3 interactions: • attraction of electron #1 to the nucleus • attraction of electron #2 to the nucleus • repulsion of two electrons for each other • Can’t solve Schrödinger equation directly, but can make approximations. These solutions indicate that the major difference in shape is that the polyelectronic orbitals tend to be a bit contracted due to increased Z. The major difference in energy is that the energy of the orbitals within a shell are no longer degenerate with the added result that the 4s orbital is lower in energy than the 3d.

  18. Energy Levels for Multielectron Atoms l 0 1 2 ml 0 -1 0 +1 -2 -1 0 +1 +2 4f’s 4d 4p 3d 4s 3p 3s 2p 2s 1s

  19. Electron Configurations and the Periodic Table

  20. Electron Spin • To completely describe an electron in an atom, a fourth quantum number, electron spin (ms = ±½), must be specified. • Each electron has a magnetic moment which is quantized in one of two possible orientations: parallel or opposed to an applied magnetic field. • Pauli Exclusion Principle: in a given atom, no two electrons may have the same four quantum numbers. So, for a given orbital, described by n, l, and ml, a maximum of two electrons may be present. • Aufbau Principle: electron configurations for multielectron atoms are built up and utilize the lowest available energy level. • Hund’s Rule: for a set of degenerate orbitals, the electrons will tend to utilize the largest number of available energy levels (maximum multiplicity). Spins will tend to align within a set of degenerate orbitals. Finally, nature loves filled and half-filled subshells.

  21. Shielding Effects • The energy of an electron in a 1 electron system is proportional to the atomic number (Z) and the shell the electron is in (n). • So, as Z increases, the energy becomes more negative; as n increases, the energy becomes less negative. Negative = more stable. • If we extend this to a multielectron case we find, for example, 55Cs; Z = 55 and n =6. • Compare to hydrogen: • This would seem to imply that all of the electrons in Cs are of lower energy than the electron in the hydrogen atom. NOT SO! Z2 E α - n2 (55)2 E α - = -84 (6)2 (1)2 E α - = -1 (1)2

  22. Shielding Effects • How can we explain/remedy thisseemingly incorrect result. • Look, for example, at Li (3 electrons). Remember that as n increases, the distance from the nucleus increases. • So, the 1s electrons tend to lie between the nucleus and the 2s electron. • Therefore the negatively charged 1s electrons are said to shield the 2s electron from the attractions of the nucleus. • This -2 charge essentially reduces the nuclear charge felt outside the 1s sphere from +3 to +1. So, the 2s electron “feels” an effective nuclear charge, Zeff, of approximately +1. • So, for Cs, the valence electron has a Zeff ≈ 1. (1)2 E α - = -0.027 (6)2

  23. Shielding Effects • The effective nuclear charge, Zeff, is not a whole number due to penetration by orbitals. • The penetration decreases with type of orbital, s > p > d. • This, then, explains why the energy of an atom in a polyelectronic case depends on both n and l. • The Zeff “felt” by the valence electrons varies with l. s p d

  24. Slater’s Rules For Calculating Zeff e.g. Zn (4s,4p) 1 x 0.35 (3d) 10 x 0.85 (3s,3p) 8 x 0.85 (2s,2p) 8 x 1.00 (1s) 2 x 1.00 S = 25.65 Zeff= 30- 25.65 = 4.35 • Zeff = Z – S • Determining S • Put electronic structure in groups as follows (NOT order of filling): (1s) (2s,2p) (3s,3p) (3d) (4s,4p) (4d) (4f) (5s,5p) etc. • Ignore electrons to the right (or above) and leave out electron of interest. • For ns and np valence electrons: • electrons in same ns,npgoup contribute 0.35. • electrons in n-1 group contribute 0.85. • electrons in n-2+ groups contribute 1.00. • For nd or nf electrons: • electrons in same ndor nf group (not both) contribute 0.35 • electrons in groups to left contribute 1.00 e.g. Zn+2 (4s,4p) (3d) 9 x 0.35 (3s,3p) 8 x 1.00 (2s,2p) 8 x 1.00 (1s) 2 x 1.00 S = 21.15 Zeff= 30- 21.15 = 8.85

  25. Trends in Atomic Sizes • Have already seen that as n increases, r increases. • Counteracting this is effect of increasing Zeff. • Trends: • R increases slowly down Periodic Table (Zeff increases slowly; n fast). • R increases rapidly across Periodic Table (Zeff increases considerably while n is constant). • For a given element, M+2 < M+ < M0 < M- < M-2

  26. Trends in Ionization Energies. • Energy difference between highest occupied energy level and n = ∞. • Trends: • Ionization Energy decreases down Periodic Table (due to increased size and decreased Zeff). • Ionization Energy tends to increase across Periodic Table (due to increasing Zeff). • however, B<Be due to change from 2s to 2p orbitals; full 2s more stable • and, O<N due to change from stable half-filled subshell. • With transition metals, generally the ns will ionize before the (n-1)d.

  27. Trends in Electron Affinities. • Electron Affinity now defined as energy required to remove an electron from a negative ion: A-(g) → A(g) + e- • Old definition: energy released when an electron is added to an atom’s outer shell: M(g) + e- → M- ΔH = -EA • i.e. value positive for an exothermic reaction • new definition gets rid of confusion with sign. • Trends parallel to Ionization Energies, though smaller values (easier to remove an extra electron than an electron from a neutral atom). • Mostly endothermic, except for noble gases and alkaline earths which want to get rid of the extra electron.

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