Periodic Properties of Chapter 8 Periodic Properties of the Elements

For an atom, electrons are in atomic orbitals. Energy of an atomic orbital

Orbital Energy Levels for the Hydrogen Atom H atom: E only depends on n degenerate

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↑

A Picture of the Spinning Electron

Chapter 8, Figure 8.2 The Stern–Gerlach Experiment

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?

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.

1s1 2s1 2p1 H atom electron configuration Lowest energy: ↑ ground state ↑ 1s 2s1 ↓ 2s Excited states 2p1 ↑ 2p orbital diagram

Now we can write the ground state electron configurations and draw orbital diagrams according to Pauli principle. Electron configurations explain many chemical properties.

Chapter 8, Figure 8.6 General Energy Ordering of Orbitals for Multielectron Atoms

Chapter 2, Figure 2.13 The Periodic Table: Main-Group and Transition Elements

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

Chapter 8, Figure 8.7 Outer Electron Configurations of the First 18 Elements in the Periodic Table

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.

Review Problem Set 10

Chapter 2, Figure 2.13 The Periodic Table: Main-Group and Transition Elements

Chapter 8, Figure 8.8 The s, p, d, and f Blocks of the Periodic Table

Periodic trends in atomic properties • Atomic radius

Atomic Radii (in Picometers) for Selected Atoms

Atomic radius In a period: decreases from left to right In a group: increases from top to bottom

(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

Chapter 2, Figure 2.13 The Periodic Table: Main-Group and Transition Elements

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 Ca2+ (c) Br– or Kr

Periodic trends in atomic properties • Atomic radius • Ionization energy

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

Chapter 8, Figure 8.15 First Ionization Energy versus Atomic Number for the Elements through Xenon

Ionization energy In a period: increases from left to right In a group: decreases from top to bottom (general trend)

Periodic trends in atomic properties • Atomic radius • Ionization energy • Electron affinity

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

Chapter 8, Figure 8.17 Electron Affinities of Selected Main-Group Elements

Electron affinity In a period: increases from left to right In a group: no clear trend (very rough trend)

Periodic trends in atomic properties • Atomic radius • Ionization energy • Electron affinity Remember the trends

Chapter 2, Figure 2.13 The Periodic Table: Main-Group and Transition Elements

Problem Set 11