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Chapter 5

Chapter 5. Periodicity and Atomic Structure. Development of the Periodic Table. The periodic table is the most important organizing principle in chemistry. Periodic table powerpoint – elements of a group have similar properties Chapter 2 – elements in a group form similar formulas

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Chapter 5

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  1. Chapter 5 Periodicity and Atomic Structure

  2. Development of the Periodic Table • The periodic table is the most important organizing principle in chemistry. • Periodic table powerpoint – elements of a group have similar properties • Chapter 2 – elements in a group form similar formulas • Predict the properties of an element by knowing the properties of other elements in the group

  3. Light and the Electromagnetic Spectrum • Radiation (light) composed waves of energy • Waves were continuous and spanned the electromagnetic spectrum

  4. Light and the Electromagnetic Spectrum

  5. Light and the Electromagnetic Spectrum

  6. Light and the Electromagnetic Spectrum • Speedof a wave is the wavelength (in meters) multiplied by its frequency in reciprocal seconds. Wavelength x Frequency = Speed  (m) x  (s–1) = c (m/s–1) • C – speed of light - 2.9979 x 108 m/s–1

  7. Electromagnetic Radiation and Atomic Spectra • Classical Physics does not explain • Black-body radiation • Photoelectric effect • Atomic Line Spectra

  8. Particlelike Properties of Electromagnetic Radiation: The Plank Equation • Blackbody radiation is the visible glow that solid objects emit when heated. • Max Planck (1858–1947): Developed a formula to fit the observations. He proposed that energy is only emitted in discrete packets called quanta. • The amount of energy depends on the frequency:

  9. Particlelike Properties of Electromagnetic Radiation: The Plank Equation • A photon’s energy must exceed a minimum threshold for electrons to be ejected. • Energy of a photon depends only on the frequency.

  10. Electromagnetic Radiation and Atomic Spectra • Atomic spectra: Result from excited atoms emitting light. • Line spectra:Result from electron transitions between specific energy levels.

  11. Electromagnetic Radiation and Atomic Spectra 1/λ = R [1/m2 – 1/n2]

  12. Quantum Mechanics and the Heisenburg Uncertainty Principle • Niels Bohr (1885–1962): Described atom as electrons circling around a nucleus and concluded that electrons have specific energy levels. • Erwin Schrödinger (1887–1961): Proposed quantum mechanical model of atom, which focuses on wavelike properties of electrons.

  13. Quantum Mechanics and the Heisenburg Uncertainty Principle • Werner Heisenberg (1901–1976): Showed that it is impossible to know (or measure) precisely both the position and velocity (or the momentum) at the same time. • The simple act of “seeing” an electron would change its energy and therefore its position.

  14. Wave Functions and Quantum Mechanics • Erwin Schrödinger (1887–1961): Developed a compromise which calculates both the energy of an electron and the probability of finding an electron at any point in the molecule. • This is accomplished by solving the Schrödinger equation, resulting in the wave function, .

  15. Wave Functions and Quantum Mechanics • Wave functions describe the behavior of electrons. • Each wave function contains three variables called quantum numbers: • • Principal Quantum Number (n) • • Angular-Momentum Quantum Number (l) • • Magnetic Quantum Number (ml)

  16. Wave Functions and Quantum Mechanics • Principal Quantum Number (n):Defines the size and energy level of the orbital. n = 1, 2, 3,  • As n increases, the electrons get farther from the nucleus. • As n increases, the electrons’ energy increases. • Each value of n is generally called a shell.

  17. Wave Functions and Quantum Mechanics • Angular-Momentum Quantum Number (l): Defines the three-dimensional shape of the orbital. • For an orbital of principal quantum number n, the value of l can have an integer value from 0 to n – 1. • This gives the subshell notation:l = 0= s orbital l = 1 = p orbital l = 2= d orbitall = 3= f orbital l = 4= g orbital

  18. Wave Functions and Quantum Mechanics • Magnetic Quantum Number (ml):Defines the spatial orientation of the orbital. • For orbital of angular-momentum quantum number, l, the value of ml has integer values from –l to +l. • This gives a spatial orientation of:l = 0 giving ml = 0 l = 1 giving ml = –1, 0, +1l = 2 giving ml = –2, –1, 0, 1, 2, and so on…...

  19. Wave Functions and Quantum Mechanics

  20. Problem • Why can’t an electron have the following quantum numbers? • (a) n = 2, l = 2, ml = 1 (b) n = 3, l = 0, ml = 3 • (c) n = 5, l = –2, ml= 1 • Give orbital notations for electrons with the following quantum numbers: • (a) n = 2, l = 1, ml = 1 (b) n = 4, l = 3, ml = –2 • (c) n = 3, l = 2, ml= –1

  21. The Shapes of Orbitals • s Orbital Shapes:

  22. The Shapes of Orbitals • p Orbital Shapes:

  23. The Shape of Orbitals • d and f Orbital Shapes:

  24. Orbital Energy Levels in Multielectron Atoms

  25. Orbital Energy Levels in Multielectron Atoms • Zeff is lower than actual nuclear charge. • Zeff increases toward nucleus ns > np > nd > nf • This explains certain periodic changes observed.

  26. Orbital Energy Levels in Multielectron Atoms • Electron shielding leads to energy differences among orbitals within a shell. • Net nuclear charge felt by an electron is called the effective nuclear charge (Zeff).

  27. Wave Functions and Quantum Mechanics • Spin Quantum Number: • The Pauli Exclusion Principle states that no two electrons can have the same four quantum numbers.x

  28. Electron Configurations of Multielectron Atoms • Pauli Exclusion Principle: No two electrons in an atom can have the same quantum numbers (n, l, ml, ms). • Hund’s Rule: When filling orbitals in the same subshell, maximize the number of parallel spins.

  29. Electron Configurations of Multielectron Atoms • Rules of Aufbau Principle: • Lower n orbitals fill first. • Each orbital holds two electrons; each with different ms. • Half-fill degenerate orbitals before pairingelectrons.

  30. Electron Configurations and Multielectron Atoms Li 1s2 2s1 1s 2s Be1s2 2s2 1s 2s B 1s2 2s2 2p1 1s 2s 2px 2py 2pz C  1s2 2s2 2p2 1s 2s 2px 2py 2pz

  31. Electron Configurations and Multielectron Atoms N   1s2 2s2 2p3 1s 2s 2px 2py 2pz O  1s2 2s2 2p4 1s 2s 2px 2py 2pz Ne1s2 2s2 2p5 1s 2s 2px 2py 2pz S[Ne]  [Ne] 3s2 3p4 3s 3px 3py 3pz

  32. Problems • Give the ground-state electron configurations for: • Ne (Z = 10) Mn (Z = 25) Zn (Z = 30) • Eu (Z = 63) W (Z = 74) Lw (Z = 103) • Identify elements with ground-state configurations: • 1s2 2s2 2p4 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 5s2 4d6 • 1s2 2s2 2p6 [Ar] 4s2 3d1 [Xe] 6s2 4f14 5d10 6p5

  33. Electron Configurations and the Periodic Table

  34. Some Anomalous Electron Configurations • Anomalous Electron Configurations: Result from unusual stability of half-filled & full-filled subshells. • Chromium should be [Ar] 4s2 3d4, but is [Ar] 4s1 3d5 • Copper should be [Ar] 4s2 3d9, but is [Ar] 4s1 3d10 • In the second transition series this is even more pronounced, with Nb, Mo, Ru, Rh, Pd, and Ag having anomalous configurations (Figure 5.20).

  35. Electron Configurations and Periodic Properties: Atomic Radii

  36. Optional Homework • Text – 5.24, 5.26, 5.28, 5.30, 5.32, 5.34, 5.44, 5.56, 5.58, 5.66, 5.68, 5.70, 5.72, 5.76, 5.78, 5.82, 5.84, 5.94, 5.98, 5.108 • Chapter 5 Homework online

  37. Required Homework • Assignment 5

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