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Chapter 6

Chapter 6. The Periodic Table & Periodic Law. Section 6.1. Development of the Modern Periodic Table. John Newlands. In 1864, noticed when the elements were arranged in order of increasing atomic mass, their properties repeated every eight elements. THE LAW OF OCTAVES. Meyer & Mendeleev.

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Chapter 6

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  1. Chapter 6 The Periodic Table &Periodic Law

  2. Section 6.1 Development of the Modern Periodic Table

  3. John Newlands • In 1864, noticed when the elements were arranged in order of increasing atomic mass, their properties repeated every eight elements. • THE LAW OF OCTAVES

  4. Meyer & Mendeleev • In 1869, published almost identical versions with the elements in order of increasing atomic mass and in columns with similar properties.

  5. Mendeleev • Mendeleev is given more credit than Meyer BECAUSE: • He published his table first • He better demonstrated his table • Suggested some of the previously measured masses were incorrect • Left blanks for not yet discovered elements "Father" of the Periodic Table

  6. Table 8.1 Mendeleev’s Predicted Properties of Ge “eka Silicon” and Its Actual Properties Predicted Properties eka Silicon Property Actual Properties Ge atomic mass 72amu 72.61amu appearance gray metal gray metal density 5.5g/cm3 5.32g/cm3 molar volume 13cm3/mol 13.65cm3/mol specific heat capacity 0.31J/g*K 0.32J/g*K oxide formula EO2 GeO2 oxide density 4.7g/cm3 4.23g/cm3 sulfide formula and solubility ES2; insoluble in H2O; soluble in aqueous (NH4)2S GeS2;insoluble H2O; soluble aqu (NH4)2S ECl4; (<1000C) chloride formula (boiling point) GeCl4; (840C) chloride density 1.9g/cm3 1.844g/cm3

  7. Development of Periodic Table Was Mendeleev psychic???? • periodic law: when arranged by atomic # elements exhibit a periodic recurrence of similar properties • Quantum-mechanical model of atom explains organization of table

  8. Mosley • In 1913, using X-rays, he discovered a unique number of protons in the nuclei of atoms for each element. • Today the elements are arranged in order of increasing atomic number • PERIODIC LAW • There is a periodic repetition of chemical and physical properties of the elements when they are arranged in order ofincreasing atomic number

  9. Arrangement of the Periodic Table • Groups/Families • 18 vertical columns (↑↓) • Two Labeling Systems • Number-and-letter system • A through 8A columns (representative elements) • 1B through 8B short columns (transition elements) 2. Number system - 1-18 • Periods • 7 horizontal rows (↔)

  10. GROUPS/FAMILIES PERIODS

  11. Arrangements of the Periodic Table

  12. Metals • Shiny • Good conductors of heat and electricity • Malleable & Ductile • Generally Solid at room temperature Group 1 Alkali Metals Group 2 Alkaline Earth Metals Groups 3-12 Transition Metals Lanthanide & Actinide Groups Inner Transition Metals

  13. Nonmetals & Metalloids B Si Ge As Sb Te Po At • Nonmetals • Dull • Generally gases or brittle solids at room temperature • Poor conductors of heat and electricity • Metalloids • Elements with physical and chemical properties of both metals and nonmetals • Rest on the “stair-step” Nonmetals → ←Metals

  14. Section 6.2 Classification of Elements

  15. Element Placement Why are elements put into groups/families together? Because they have similar chemical properties Why do elements have similar chemical properties? Because they have the same number of valence electrons Group 1 – Alkali Metals Period 2 Lithium 1s22s1 [He]2s1 Period 3 Sodium 1s22s22p63s1 [Ne]3s1 Period 4 Potassium 1s22s22p63s23p64s1 [Ne]4s1 ALL ELEMENTS IN GROUP 1 (ALKALI METALS) HAVE ONE VALENCE ELECTRON

  16. Recurring pattern in e- configuration is basis for periodic behavior. • Main group, group # = valence e- count • Valence e- responsible for chemistry • Elements in same group behave similarly

  17. Dot Diagrams for Representative Elements

  18. Figure 8.12 A periodic table of partial ground-state electron configurations

  19. Representative Elements • s-block elements • Groups 1&2, hydrogen & helium • Valence electrons occupy outermost s sublevels only • p-block elements • Groups 13-18 (except helium) • Valence electrons include a full outermost s sublevel and a filled or partially filled p sublevel Period number is equal to the principle energy level where the valence electrons are located

  20. Transition Elements • d-block elements • Groups 3-12 • Valence electrons include a full outermost s sublevel and a filled or partially filled d sublevel The period number minus 1 equals the principle energy level where the valence electrons are located

  21. Inner transition metals • f-block elements • Lanthanide & Actinide Groups • Full or partially full outermost s sublevel, and full or partially full outermost f sublevel The period number minus 2 equals the principle energy level where the valence electrons are located.

  22. Section 6.3 Periodic Trends

  23. Atomic Radius • Half the distance between two nuclei of identical atoms that are chemically bonded together • Down the group • atomic radius increases • Across the period • atomic radius decreases

  24. Atomic Radius Decreases Atomic Radius Increases

  25. Figure 8.16 Periodicity of atomic radius

  26. Practice Atomic Radius • Which has the larger atomic radii of the following? B or Al Na or Mg F or Cl • Which has the smaller atomic radii of the following? H or He K or Cs N or Ne • Circle the one with the largest atomic radius and underline the one with the smallest. C, Si, Ge V, Cr, W N, Mg, Ca

  27. Ionization Energy • The amount of energy required to remove an electron from the atom (how tightly an atom holds on to its electrons) • Down a group • ionization energy decreases • Across a period • ionization energy increases

  28. Trends in Atomic Properties • Ionization Energy (IE) • Energy required for complete removal of 1 mole of e- from 1 mole of atoms • Atoms w/ low IE form cations (lose e-) • Atoms w/ high IE form anions (gain e-) Na(g)  Na+(g) + e- I1 Na+(g)  Na2+(g) + e- I2 I1 < I2 < I3

  29. 8.4 Trends in Atomic Properties • Greater IE, more difficult to remove e- • Positive values, energy into atom • Larger atoms easier to ionize Figure 8.18

  30. Ionization Energy Increases Ionization Energy Decreases

  31. Figure 8.17 Periodicity of first ionization energy (IE1)

  32. Practice Ionization Energy • Which has the greater ionization energy? Ne or Ar N or O Sc or Ti • Which has the smaller ionization energy? Al, Si, P K, Rb, Sr Be, Mg, Ca

  33. Ionic Radius • Octet Rule • Atoms tend to gain, lose, or share electrons in order to achieve a full outer energy level (typically 8 are needed) • Ion • An atom that has an overall charge due to the gaining or losing of electrons

  34. Main-group ions and noble gas configurations Figure 8.25

  35. Ionic Radius Comparisons • Metals have LOW ionization and electron affinity • They lose electrons to form positively charged ions • Positive charged ions are smaller than the original atom • Nonmetals have HIGH ionization energy and electron affinity • They gain electrons to form negatively charged ions • Negatively charged ions are larger than the original atom

  36. Trends in Properties of Monatomic ions • Cation smaller than parent • e- removed, other e- feel greater Zeff • Anion larger than parent • e- added, e-/e- repulsions occupy more space Figure 8.29

  37. Ionic Radius Increases Ionic Radius Increases FOR IONIC RADIUS… MUST FOLLOW METAL/NON-METAL RULES

  38. Ionic Radius Practice • Which is the smaller of the two? Lithium ion or Lithium atom Chlorine ion or Chlorine atom • Underline the following that will form a positively charged ion and circle the ones that will form a negatively charged ion. Mg F Al Cu Br N S K • How will the radius of each of the above change when an ion is formed? Mg F Al Cu Br N S K

  39. Electronegativity • The ability of an atom to attract electrons in a chemical bond. • Down the group • Electronegativity values decrease • Across the period • Electronegativity values increase *Noble gases are the exception to this rule.

  40. Electronegativity Increases Electronegativity Decreases

  41. Electronegativity Practice • Which has the greater electronegativity value? B or N Si or Sn Cr or W • Which has the smaller electronegativity value? Rb, Sr, Y Ga, In, Sn As, Se, S

  42. CUMULATIVE REVIEW • Which has the smallest atomic radius between Ga, In, & Tl? • Which has the highes ionization energy? • Which is the smallest: an atom of sodium, an ion of sodium, or an atom of potassium? • Which has the greatest electron affinity between zinc, arsenic, or bromine? Which has the lowest ionization energy?

  43. 8.4 Trends in Atomic Properties Figure 8.21

  44. Electronegativity Increases Atomic Radius Decreases Ionic Radius Increases Electron Affinity Increases Ionization Energy Increases Electronegativity Decreases Atomic Radius Increases Ionic Radius Increases Electron Affinity Decreases Ionization Energy Decreases

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