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The Modern Periodic Table

4.2. The Modern Periodic Table. Periodic Trends of the Elements. 4.1 Development of the Periodic Table 4.2 The Modern Periodic Table Classification of Elements 4.3 Effective Nuclear Charge 4.4 Periodic Trends in Properties of Elements Atomic Radius Ionization Energy

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The Modern Periodic Table

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  1. 4.2 The Modern Periodic Table

  2. Periodic Trends of the Elements 4.1 Development of the Periodic Table 4.2 The Modern Periodic Table Classification of Elements 4.3 Effective Nuclear Charge 4.4 Periodic Trends in Properties of Elements Atomic Radius Ionization Energy Electron Affinity Metallic Character 4.5 Electron Configuration of Ions Ions of Main Group Elements Ions of d-Block Elements 4.6 Ionic Radius Comparing Ionic Radius with Atomic Radius Isoelectronic Series 4

  3. Valence electrons determine most properties Groups have the same valence electron makeup Groups have similar properties (including reactivity, ionization, etc.) Early attempts to develop a periodic table organized elements by AW but didn’t account for missing elements. Mendeleev accounted for missing elements, e.g. Ga & Ge, and predicted their properties. He also assumed some AW calculations were wrong, e.g. Te/I and reversed order. He did this because he realized that elements with similar properties belonged in same column.

  4. Periodic Trends describe how specific properties of the elements change as you go … across a period or down a family. Periodic Trends are influenced by the effective nuclear charge (+) felt by the valence electrons In an atom.

  5. Periodic Trends atom size — radius of atom — sphere of ‘repulsive’ influence of electrons 1st ionization energy — energy required to knock e- from atom: e.g. Na(g) → Na+ + e- electron affinity – energy released when -1 anion formed: e.g. Cl + e- → Cl- Electronegativity– relative desire of atom to share e- from other atoms – only has meaning with covalently bonded atoms

  6. 4.4 Periodic Trends in Properties of Elements Atomic radius is the distance between the nucleus of an atom and its valence shell. (a) Atomic radius in metals, or metallic radius, is half the distance between the nuclei of two adjacent, identical metal atoms. (b) Atomic radius in nonmetals, or covalent radius, is half the distance between adjacent, identical nuclei connected by a chemical bond.

  7. Periodic Trends — Atomic Radii What will happen to atomic radii as you go across the periodic table? a. Increase b. Decrease c. stay same

  8. 4.3 Effective Nuclear Charge Effective nuclear charge (Zeff)is the actual magnitude of positive charge that is “experienced” by a valence electron in the atom. Shielding – The attractive effect of the nucleus on valence shell electrons is reduced by the by the influence of the inner shell electrons. Valence shell electrons shield each other to a lesser extent - As a result, the value of Zeff increases steadily from left to right because the inner shell electrons remain the same but Z increases. +3 +4

  9. Atomic Radius Atomic radii (in picometers) The atomic radius increases from top to bottom down a group. • Increasing n, so outermost shell lies farther from the nucleus Atomic radius decreases from left to right across a period. • Increasing Zeffwhich draws the valence shell closer to the nucleus

  10. Atomic Radius Atomic radius decreases left to right across a period due to increased electrostatic attraction between the effective nuclear charge and the charge on the valence shell.

  11. Periodic Trends — Atomic Radii Na Li Radii (pm) Period 3 Ar Ne Period 2 Atomic #

  12. Cs Rb K Radii (pm) Na Li Atomic Radii Periodic Trends - down group/family At I Br Radii (pm) Cl F Atomic #

  13. Ionization Energy Ionization energy (IE)is the minimum energy required to remove an electron from an atom in the gas phase. The result is an ion, a chemical species with a net charge. Sodium has an ionization energy of 495.8 kJ/mol. Specifically, 495.8 kJ/mol is the first ionization energy of sodium, IE1(Na), which corresponds to the removal of the most loosely held electron. Na(g) → Na+(g) + e− What is the ionization energy of one Na atom in J? What general trend do you expect for 1st ionization energy across a period? a) increase b) decrease c) stay the same

  14. 4.3 Effective Nuclear Charge Effective nuclear charge (Zeff)is the actual magnitude of positive charge that is “experienced” by an electron in the atom. In a multi-electron atom, electrons are simultaneously attracted to the nucleus and repelled by one another. • This results in shielding, where an electron is partially shielded from the positive charge of the nucleus by the other electrons. • Although all electrons shield one another to some extent, the most effective are the core electrons. • As a result, the value of Zeff increases steadily from left to right because the core electrons remain the same but Z increases.

  15. 4.4 Periodic Trends in Properties of Elements Atomic radius is the distance between the nucleus of an atom and its valence shell. (a) Atomic radius in metals, or metallic radius, is half the distance between the nuclei of two adjacent, identical metal atoms. (b) Atomic radius in nonmetals, or covalent radius, is half the distance between adjacent, identical nuclei connected by a chemical bond.

  16. Atomic Radius Atomic radii (in picometers) The atomic radius increases from top to bottom down a group. • Increasing n, so outermost shell lies farther from the nucleus Atomic radius decreases from left to right across a period. • Increasing Zeffwhich draws the valence shell closer to the nucleus

  17. Atomic Radius Atomic radius decreases left to right across a period due to increased electrostatic attraction between the effective nuclear charge and the charge on the valence shell.

  18. Ionization Energy Ionization energy (IE)is the minimum energy required to remove an electron from an atom in the gas phase. The result is an ion, a chemical species with a net charge. Sodium has an ionization energy of 495.8 kJ/mol. Specifically, 495.8 kJ/mol is the first ionization energy of sodium, IE1(Na), which corresponds to the removal of the most loosely held electron. Na(g) → Na+(g) + e−

  19. Ionization Energy

  20. Ionization Energy • In general, as Zeff increases, ionization energy also increases. • Thus, IE1 increases from left to right across a period.

  21. Ionization Energy Within a given shell, electrons with a higher value of l are higher in energy and thus, easier to remove.

  22. Ionization Energy Within a given shell, electrons with a higher value of l are higher in energy and thus, easier to remove.

  23. Ionization Energy It is possible to remove additional electrons in subsequent ionizations, giving IE1, IE2, and so on. IE1(Na) = 496 kJ/mol IE2(Na) = 4562 kJ/mol Na(g) → Na+(g) + e− Na+(g) → Na2+(g) + e−

  24. Ionization Energy It takes more energy to remove the 2nd, 3rd, 4th, etc. electrons because it is harder to remove an electron from a cation than an atom. It takes much more energy to remove core electrons than valence. • Core electrons are closer to nucleus. • Core electrons experience greater Zeff because of fewer filled shells shielding them from the nucleus.

  25. Electron Affinity Electron affinity (EA) is the energy released when an atom in the gas phase accepts an electron. Cl(g) + e−→ Cl−(g)

  26. Electron Affinity • Like ionization energy, electron affinity increases from left to right across a period as Zeff increases. • Easier to add an electron as the positive charge of the nucleus increases.

  27. 4.6 Ionic Radius The ionic radiusis the radius of a cation or an anion. When an atom loses an electron to become a cation, its radius decreases due in part to a reduction in electron-electron repulsions in the valence shell. A significant decrease in radius occurs when all of an atom’s valence electrons are removed.

  28. Comparing Ionic Radius with Atomic Radius When an atom gains one or more electrons and becomes an anion, its radius increases due to increased electron-electron repulsions.

  29. Isoelectronic Series An isoelectronic series is a series of two or more species that have identical electron configurations, but different nuclear charges. O2−: 1s22s22p6 F−: 1s22s22p6 Ne: 1s22s22p6 isoelectronic

  30. Isoelectronic Series An isoelectronic series is a series of two or more species that have identical electron configurations, but different nuclear charges. O2−: 1s22s22p6 F−: 1s22s22p6 Ne: 1s22s22p6 isoelectronic

  31. electron affinity Write the equation representing the electron affinity for tin, Sn …. represent energy as either a reactant or product

  32. Electron Affinity Electron affinity (EA) is the energy released when an atom in the gas phase accepts an electron. Cl(g) + e−→ Cl−(g)

  33. Electron Affinity • Like ionization energy, electron affinity increases from left to right across a period as Zeff increases. • Easier to add an electron as the positive charge of the nucleus increases.

  34. H = 2.1 Periodic Trends — electronegativity desire to ‘hog’ shared electrons 3.0 Cl 2.8 Br 2.5 I This is qualitatively the same trend as for electron affinity C = 2.5, N = 3.0, O = 3.5, F = 4.0

  35. H 2.1 He Li Be B C N O F Ne 1.0 1.5 2.0 2.5 3.0 3.5 4.0 Na Mg Al Si P S ClAr 0.9 1.2 1.5 1.8 2.1 2.5 3.0 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn GaGe As Se Br Kr 0.8 1.0 1.3 1.5 1.6 1.6 1.5 1.8 1.9 1.9 1.9 1.6 1.6 1.8 2.0 2.4 2.8 RbSr Y ZrNb Mo TcRuRh Pd Ag Cd In SnSb Te I Xe 0.8 1.0 1.2 1.4 1.6 1.8 1.9 2.2 2.2 2.2 1.9 1.7 1.7 1.8 1.9 2.1 2.5 Cs Ba La Hf Ta W Re Os Ir Pt Au Hg TlPb Bi Po At Rn 0.7 0.9 1.0 1.3 1.5 1.7 1.9 2.2 2.2 2.2 2.4 1.9 1.8 1.9 1.9 2.0 2.2 Fr Ra 0.7 0.9 Electronegativity (EN) Polarity & Polar covalent bonds

  36. 4.6 Ionic Radius The ionic radiusis the radius of a cation or an anion. When an atom loses an electron to become a cation, its radius decreases due in part to a reduction in electron-electron repulsions in the valence shell. A significant decrease in radius occurs when all of an atom’s valence electrons are removed.

  37. Comparing Ionic Radius with Atomic Radius When an atom gains one or more electrons and becomes an anion, its radius increases due to increased electron-electron repulsions.

  38. Isoelectronic Series An isoelectronic series is a series of two or more species that have identical electron configurations, but different nuclear charges. What cations are isoelectronic to Neon? Ne: 1s22s22p6 O2−: 1s22s22p6 F−: 1s22s22p6 isoelectronic

  39. Two ions (or an atom and an ion) are isoelectronic if they have the same electronic configuration. Isoelectronic ions (same electronic configuration) get smaller as the nuclear charge increases. These ions are all isoelectronic to Neon, Ne.

  40. Chapter Summary: Key Points 4

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