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DO NOW: What do sodium, potassium, and rubidium have in common?

DO NOW: What do sodium, potassium, and rubidium have in common?. Chemistry periodicity history of the periodic table, Electron configurations and the periodic table. Organizing the Elements. By 1700, only 13 elements had been identified.

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DO NOW: What do sodium, potassium, and rubidium have in common?

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  1. DO NOW: What do sodium, potassium, and rubidium have in common? Chemistryperiodicityhistory of the periodic table,Electron configurations and the periodic table

  2. Organizing the Elements • By 1700, only 13 elements had been identified. • As more elements were discovered, the need to organize them increased. • Chemists used common properties of elements to group them.

  3. Dimitri Mendeleev • Chemistry professor who was working on a textbook for his students. • Arranged ~60 known elements by atomic mass. • Noticed a repeating pattern.

  4. The Pattern • Every so often, this pattern would repeat: • Reactive gas, nonreactive gas, very reactive metal • Mendeleev was so convinced of these patterns that he predicted that there existed elements that hadn’t been discovered yet, because of “holes” in the patterns. • Later, gallium and germanium were discovered, and they fit in these holes, validating Mendeleev’s periodic table.

  5. Mendeleev or Meyer? • A German scientist, Lothar Meyer, also made a table similar to Mendeleev. • Mendeleev’s predictions of not yet discovered elements proved to be superior.

  6. Periodic Law • Mendeleev arranged the elements by atomic mass. • However, he switched iodine and tellurium, because iodine behaved like fluorine and chlorine. • He assumed that the atomic mass of iodine was incorrect.

  7. Periodic Law • The atomic mass wasn’t incorrect, but rather, the periodic table works best when arranged by atomic number as opposed to atomic mass. • Henry Moseley introduced this idea in 1913.

  8. Periodic Law • The periodic law states that when atoms are arranged by atomic number, properties repeat. • Specifically, atoms in the same column have very similar properties.

  9. Metals, Nonmetals, and Metalloids • Metals are • Nonmetals are on the right side of the stair-step line (with the exception of H)

  10. Metals, Nonmetals, and Metalloids • Metals are • Metalloids border the stair-step line, but Al is considered to be a metal.

  11. Metals, Nonmetals, and Metalloids • Metals are • Metals are on the left side of the stair-step line.

  12. Periodic Table • The periodic table is organized by periods (rows) and groups or families(columns). • The groups are organized by two letters, A and B. • The number before “A” represents the number of valence electrons, or outer electrons. • The number before “B” includes the d-shell electrons • Valence electrons determine nearly everything about chemical bonding, and ALL OF CHEMISTRY. I’m not kidding. This is important.

  13. Periodic Table • All elements in a group have similar chemical properties. The “A” elements have their valence electrons in the s and p orbitals. • Group 1A – alkali metals • Group 2A – alkaline earth metals • Group 3A – boron group • Group 4A – carbon group • Group 5A – nitrogen group • Group 6A – chalcogens • Group 7A – halogens • Group 8A – noble gases

  14. Periodic Table • The metals in the d-block are collectively known as the transition metals. • The metals in the f-block are collectively known as the rare earth metals • The top row elements are the lanthanides • The bottom row elements are the actinides

  15. Review – Aufbau Principle • The order in which orbitals fill up based on their energies. • Lowest energy gets filled up first. • Recall the “diagonal chart” • 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d…

  16. Electron Configurations and the Periodic Table • The fact that elements have similar properties within the same column is no coincidence. • Valence electrons determine chemical properties and bonding trends. • Therefore, elements with the same number of valence electrons will act in similar ways. • Alkali metal reactivity w/ water

  17. Group A Electron Configurations • The valence electrons in group A are the sum of the s and p electrons.

  18. “Stable Octet” • Having a completely full valence shell (s2p6) is a very stable configuration. • Atoms will lose or gain electrons to be like the noble gases, or have 8 valence electrons. • Called a stable octet • Hydrogen/helium will have a stable duet (2 electrons) because there is no p orbital in the first energy level. • This makes each column of the periodic table have predictable ionic charges.

  19. Common Ionic Charges • Groups 1A-3A will lose electrons to be gain a noble gas configuration • Groups 5A-7A will gain electrons to gain a noble gas configuration

  20. Practice • What is the electron configuration for a neutral atom of oxygen? • What is the electron configuration for the oxide anion, O2-? • What is the electron configuration for the noble gas neon?

  21. Isoelectronic Series • Atoms and ions with the same number of electrons are called isoelectronic. • Example: Neon has 10 electrons • Fluoride: F- : 9 + 1 = 10 • Sodium ion: Na+ : 11 – 1 = 10 • F-, Ne, Na+ are isoelectronic because all three species have 10 electrons, and the same electron configuration as a result. • 1s22s22p6

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