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Periodic Properties of
Chapter 8 Periodic Properties of the Elements
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For an atom, electrons are in atomic orbitals.
Energy of an atomic orbital
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Orbital Energy Levels for the Hydrogen Atom
H atom: E only depends on n degenerate
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E depends on n and l same n, l↑ ↔ E↑ Chapter 8, Figure 8.6
General Energy Ordering of Orbitals for Multielectron Atoms E depends on n and l same n, l↑ ↔ E↑
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A Picture of the Spinning Electron
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Chapter 8, Figure 8.2 The Stern–Gerlach Experiment
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Spin quantum number ms ms = +1/2 or −1/2 4 quantum numbers are used to specify an electron. How do electrons fill up atomic orbitals?
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Pauli Exclusion Principle
In a given atom, no two electrons can have the same set of four quantum numbers. An orbital can hold only two electrons, they must have opposite spins.
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1s1 2s1 2p1 H atom electron configuration Lowest energy: ↑
ground state ↑ 1s 2s1 ↓ 2s Excited states 2p1 ↑ 2p orbital diagram
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Now we can write the ground state electron
configurations and draw orbital diagrams according to Pauli principle. Electron configurations explain many chemical properties.
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Chapter 8, Figure 8.6 General Energy Ordering of Orbitals for Multielectron Atoms
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Chapter 2, Figure 2.13 The Periodic Table: Main-Group and Transition Elements
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Hund’s rule For degenerate orbitals, the lowest energy is attained
when the number of electrons with the same spin is maximized. Valence electrons: electrons in the outermost shell for main group elements. They are involved in bonding and reactions. Core electrons: inner electrons
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Chapter 8, Figure 8.7 Outer Electron Configurations of the First 18 Elements in the Periodic Table
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Elements in the same group have similar valence
electron configuration — similar chemical properties. Number of valence electrons = main group number Number of filled shells = period number Noble gases have 8 (He 2) valence electrons. Stable structure. Metals: tend to lose valence electrons to reach 8(2) valence electron. Nonmetals: tend to gain electrons to reach 8(2) valence electrons.
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Review Problem Set 10
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Chapter 2, Figure 2.13 The Periodic Table: Main-Group and Transition Elements
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Chapter 8, Figure 8.8 The s, p, d, and f Blocks of the Periodic Table
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Periodic trends in atomic properties
• Atomic radius
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Atomic Radii (in Picometers) for Selected Atoms
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Atomic radius In a period: decreases from left to right
In a group: increases from top to bottom
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(a) N or F (b) C or Ge (c) N or Al (d) Al or Ge
EXAMPLE 8.5 Atomic Size On the basis of periodic trends, choose the larger atom in each pair (if possible). Explain your choices. (a) N or F (b) C or Ge (c) N or Al (d) Al or Ge
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Chapter 2, Figure 2.13 The Periodic Table: Main-Group and Transition Elements
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Choose the larger atom or ion from each pair.
EXAMPLE 8.7 Ion Size Choose the larger atom or ion from each pair. (a) S or S2– (b) Ca or Ca (c) Br– or Kr
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Periodic trends in atomic properties
• Atomic radius • Ionization energy
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Ionization energy Energy required to remove an electron from
a gaseous atom or ion. X(g) → X+(g) + e− first ionization energy X+(g) → X2+(g) + e− second ionization energy
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Chapter 8, Figure 8.15 First Ionization Energy versus Atomic Number for the Elements through Xenon
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Ionization energy In a period: increases from left to right
In a group: decreases from top to bottom (general trend)
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Periodic trends in atomic properties
• Atomic radius • Ionization energy • Electron affinity
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Electron affinity Energy change associated with the addition
of an electron to a gaseous atom. X(g) + e− → X−(g) X(g) + e− X−(g) E Ei Ef stable X− ∆E = Ef − Ei = EA < 0
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Chapter 8, Figure 8.17 Electron Affinities of Selected Main-Group Elements
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Electron affinity In a period: increases from left to right
In a group: no clear trend (very rough trend)
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Periodic trends in atomic properties
• Atomic radius • Ionization energy • Electron affinity Remember the trends
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Chapter 2, Figure 2.13 The Periodic Table: Main-Group and Transition Elements
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Problem Set 11
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