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Chapter 8

Chapter 8. Electron Configuration and Periodicity. Overview. Electron Structure of Atoms Electron spin and the Pauli Exclusion Principle. Aufbau Principle and the Periodic Table Electron Configuration Orbital Diagram of atoms; Hund’s Rule Periodicity of the Elements

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Chapter 8

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  1. Chapter 8 Electron Configuration and Periodicity

  2. Overview • Electron Structure of Atoms • Electron spin and the Pauli Exclusion Principle. • Aufbau Principle and the Periodic Table • Electron Configuration • Orbital Diagram of atoms; Hund’s Rule • Periodicity of the Elements • Mendelev’s periodic table predicted undiscovered elements. • Periodic Properties • Periodicity and the main group elements.

  3. Orbitals in Multielectron Atoms • Electrons are attracted to the nucleus but also repelled by each other. • Repulsion from other electrons reduces the attraction to the nucleus by a small amount giving rise to an “effective nuclear charge” • Effective nuclear charge: the net nuclear charge felt by an electron after shielding from other electrons in the atom is taken into account. Zeff = Zact Zshield.

  4. 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f Diagonal Rule for Build-up Rule • The periodic table can also be used to determine the electron configuration of an element.

  5. Electron configurations of multielectron atoms (Aufbau principle) • Electron configuration determined since electrons tend to be in lowest energy orbitals. • The Aufbau principle guides us in the filling of orbitals: • Fill lowest energies first. • Maximum of two electrons with opposite spins in each orbital. • Degenerate orbitals (orbitals with same energy) follow Hund’s rule • Hund’s rule: If two or more orbitals have the same energy, fill each orbital with one electron before pairing electrons. E.g. Determine the electron configurations of H and He • H  1s1;  • He  1s2;  E.g. 2 Determine the electron configuration of the second row elements. E.g.3 Determine the electron configuration of the 4th row elements. • Shorthand: electron configuration of arsenic is [Ar]4s23d104p3.

  6. Magnetic Properties • Although an electron behaves like a tiny magnet, two electrons that are opposite in spin cancel each other. Only atoms with unpaired electrons exhibit magnetic susceptibility (see Fig. 8.2). • A paramagnetic substance is one that is weakly attracted by a magnetic field, usually the result of unpaired electrons. • A diamagnetic substance is not attracted by a magnetic field generally because it has only paired electrons.

  7. Periodic Table and Electron Configurations • Build-up order given by position on periodic table; row by row. • Elements in same column will have the same outer shell electron configuration.

  8. Anomalous Electron Configurations • A few exceptions to the Aufbau principles exist. Stable configuration: • half-filled d shell: • Cr has [Ar]4s13d5; • Mo has [Kr] 5s14d5 • filled d subshell: • Cu has [Ar]4s13d10 • Ag has [Kr]5s14d10. • Au has [Xe]6s14f145d10 • Exceptions occur with larger elements where orbital energies are similar.

  9. Electron Configuration of Excited States & Ions • Metals form cations by losing e; nonmetals become anions by gaining e. • Both adopt inert gas electron configuration. E.g. The alkali metals will lose a single electron to become M+. The electron configuration is [He], [Ne], [Ar], [Kr], and [Xe] for Li+, Na+, K+, Rb+ respectively.

  10. ISOELECTRONIC SUBSTANCES and EXCITED STATES • Substances with the same number of electrons are isoelectronic ions. • Isoelectronic ions (or molecules) ions (or molecules) with the same number of valence electrons. • Isoelectronic substances: P3, S2, Cl, Ar, K+, Ca2+. • The electron configation of an element in an excited state will have an electron in a high-energy state E.g. [Ar]4s13d94p1 is an excited-state electron configuration for Cu.

  11. Development of the Periodic Table • Mendeleev developed periodic table to group elements in terms of chemical properties. • Alkali metals develop +1 charge, alkaline earth metals + 2 • Nonmetals usually develop negative charge (1 for halides, 2 for group 6A, etc.) • Blank spots where elements should be were observed. • Discovery of elements with correct properties.

  12. Periodic Properties • Periodic law = elements arranged by atomic number gives physical and chemical properties varying periodically. • We will study the following periodic trends: • Atomic radii • Ionization energy • Electron affinity

  13. Atomic Radius Fig. 8.17 Atomic Radii for Main Group Elements • Atomic radii actually decrease across a row in the periodic table. Due to an increase in the effective nuclear charge. • Within each group (vertical column), the atomic radius tends to increase with the period number.

  14. Atomic Radius 2 • If positively charged the radius decreases while if the charge is negatively the radius increases (relative to the atom). • When substances have the same number of electrons (isoelectronic), the radius will depend upon which has the largest number of protons. E.g. Predict which of the following substances has the largest radius: P3, S2, Cl, Ar, K+, Ca2+.

  15. IONIZATION ENERGY • Ionization energy, Ei: minimum energy required to remove an electron from the ground state of atom (molecule) in the gas phase. M(g) + h M+ + e. • Ei related to electron configuration. Higher energies = stable ground states. • Sign of the ionization energy is always positive, i.e. it requires energy for ionization to occur. • The ionization energy is inversely proportional to the radius and directly related to Zeff. • Exceptions to trend: • B, Al, Ga, etc.: their ionization energies are slightly less than the ionization energy of the element preceding them in their period. • Before ionization ns2np1. • After ionization is ns2. Higher energy  smaller radius. • Group 6A elements. • Before ionization ns2np4. • After ionization ns2np3 where each p electron in different orbital (Hund’s rule). • Electron-electron repulsion by two electrons in same orbital increases the energy (lowers EI).

  16. Ionization Energy: Periodic table Fig. 8.18 Ionization Energy vs atomic #

  17. HIGHER IONIZATION ENERGIES • The energies for the subsequent loss of more electrons are increasingly higher. For the second ionization reaction written as • M+(g) + h M2+ + e Ei2. • Large increases in the ionization energies vary in a zig-zag way across the periodic table. • States with higher ionization energies have: 1s22s22p6 (stable).

  18. ELECTRON AFFINITY • Electron Affinity, Eea, is the energy change that occurs when an isolated atom in the gas phase gains an electron. E.g. Cl + e Cl Eea = 348.6 kJ/mol • Energy is often released during the process. • Magnitude of released energy indicates the tendency of the atom to gain an electron. • From the data in the table the halogens clearly have a strong tendency to become negatively charged • Inert gases and group I & II elements have a very small Eea.

  19. Fig. 8.2 Stern-Gerlach Experiment • Hydrogen atoms split into two beams when passed through magnetic field. Beams correspond to spin on atom. Return to slide 6

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