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Chapter 8
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  1. Chapter 8 Electron Configuration and Periodicity

  2. Development of Periodic Table Dmitri Mendeleev and Lothar Meyer independently came to the same conclusion about how elements should be grouped.

  3. Meyer vs Mendeleev • Meyer group the elements according to their chemical and physical properties based on their volumes whereas Mendeleev based it on their masses.

  4. Development of Periodic Table Because Mendeleev was able to predict the existence of certain elements. His table of elements was adopted. Mendeleev predicted the discovery of germanium (which he called eka-silicon) as an element with an atomic weight between that of zinc and arsenic, but with chemical properties similar to those of silicon.

  5. Periodic Law • Original: The chemical and physical properties of the elements are periodic functions of their atomic masses • Current: The chemical and physical properties of the elements are periodic functions of their atomic numbers (THINK! PROTONS = ELECTRONS)

  6. Atomic Properties and Periodic Trends • There are a number of important physical and chemical properties of elements that the periodic table can put in perspective: • metallic character (includes luster, conductivity, malleability, and ductility) • atomic radius — half the distance between the nuclei of two like atoms • ionic radius — the radius of the nuclei after they have formed positive or negative ions

  7. Atomic Properties and Periodic Trends • Ionization energies — the energy required to remove electrons • electron affinity — the energy change that occurs when an atom gains an electron • electronegativity — the ability of an atom to attract a pair of electrons when bonded to another atom

  8. Atomic Size • Atomic size generally increases as you move down a group of the periodic table. Electrons are added to higher principal energy levels, and nuclear charge increases. The outermost orbitals are larger, and the shielding by electrons at lower energy levels increases.

  9. As you proceed across a period, the size of successive atoms decreases. How is this explained? • With each successive element, the positive nuclear charge increases (as does the electron charge). The electrons are being added to the same principal energy level, so the force of attraction between the nucleus and the electron increases, leading to a decrease in size.

  10. 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. • In any period, the number of electrons between the nucleus and the principal energy level is the same for all elements. The shielding effect of these electrons is a constant within a period.

  11. 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.

  12. Example 1 • Choose the larger atom in each pair: • Na or Si • P or Sb • Al or Cl • Al or In

  13. Ionic Radius • If positively charged the radius decreases • If negatively charged 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.

  14. Example 2 • Choose the larger particle in each pair • Na or Na+ • Co3+ or Co2+ • Al3+ or Al

  15. Example 3 Predict which of the following substances has the largest radius: P3, S2, Cl, Ar, K+, Ca2+.

  16. 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.

  17. 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).

  18. Ionization Energy: Periodic table Ionization Energy vs atomic #

  19. Example 4 • Choose the atom with te larger ionization energy in each pair • B or C • O or S • Cl or As

  20. 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).

  21. Example 5 • Use the periodic table to select the element in the following list for which there is the largest difference between the second and third ionization energies.

  22. Electronegativity • Electronegativity is the tendency of atoms of an element to attract electrons when they are chemically combined with another element. They are expressed in a qualitative measurement called the Pauling electronegativity scale. Ionization energy and electron affinity are used to calculate electronegativity.

  23. Electronegativity Trends • Going across a period, electronegativity increases. Metals have low electronegativities, non-metals high electronegativities. • Electronegativity decreases going down a group. The ability of an atom to attract electrons decreases with increasing nuclear charge. • What is electronegativity good for? Predicting the type of bonding that exists between atoms in compound. The term was first coined by Linus Pauling in his landmark 1932 paper, "The Nature of the Chemical Bond."

  24. Example 6 • Predict the order of increasing electronegativity in each of the following groups of elements. • C, N, O • S, Se, Cl • Si, Ge, Sn

  25. 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.

  26. 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.

  27. Example 7 • Order the atoms in each of the following sets from the least exothermic electron affinity to the most • N, O, F • F, Cl, Br, I

  28. 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.

  29. 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.

  30. Example 8 • Indicate whether the following are diamagnetic or paramagnetic • Fe • Ca • Sb