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To understand how the principal energy levels fill with electrons in atoms beyond hydrogen

Objectives. To understand how the principal energy levels fill with electrons in atoms beyond hydrogen To learn about valence electrons and core electrons To learn about the electron configurations of atoms with Z < 18 To understand the general trends in properties in the periodic table.

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To understand how the principal energy levels fill with electrons in atoms beyond hydrogen

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  1. Objectives • To understand how the principal energy levels fill with electrons in atoms beyond hydrogen • To learn about valence electrons and core electrons • To learn about the electron configurations of atoms with Z < 18 • To understand the general trends in properties in the periodic table

  2. A. Electron Arrangements • Orbital Notation for Carbon  __ __ 1s 2s 2p • Electron Configuration for carbon 1s22s22p2

  3. “Orbital Filling Rules” • Aufbau Principle Electrons fill the lowest energy orbitals available • Pauli Exclusion Principle Orbitals can hold a maximum of two electrons; two electrons in one orbital must have opposite spins. • Hund’s Rule Electrons fill equal energy orbitals one at a time until each is occupied by one electron; Electrons in singly occupied orbitals have the same spin.

  4. A. Electron Arrangements in the First 18 Atoms on the Periodic Table • Valence electrons – electrons in the outermost (highest) principal energy level of an atom • Core electrons – inner electrons • Elements with the same valence electron arrangement show very similar chemical behavior. • Classifying Electrons

  5. B. Electron Configurations and the Periodic Table • Orbital filling and the periodic table

  6. A. Electron Arrangements in the First 18 Atoms on the Periodic Table

  7. B. Electron Configurations and the Periodic Table • Look at electron configurations for K through Kr

  8. Ne Na Mg F O 1s22s22p6  filled outer energy level 1s22s22p63s1 can lose 1 electron to have a filled energy level 1s22s22p63s2 loses 2 electrons 1s22s22p5 gains 1 electron 1s22s22p4 gains 2 electrons C. Atomic Properties and the Periodic Table

  9. C. Atomic Properties and the Periodic Table Metals and Nonmetals

  10. Metals Conduct electricity Left side of the periodic table Form (+) ions; lose electrons Nonmetals Do not conduct electricity Right side of the periodic table Form (-) ions; gain electrons C. Atomic Properties and the Periodic Table

  11. Atomic Size C. Atomic Properties and the Periodic Table • Measured as the atomic radius (distance from the center of an atom to its outermost electrons • Remember: • The locations of electrons are not exact  atomic radius is an approximation.

  12. C. Atomic Properties and the Periodic Table Atomic Size

  13. C. Atomic Properties and the Periodic Table Atomic Radius Trends: 1. Atoms increase in size as you move down a group on the periodic table Why? • Outermost electrons are in higher energy levels. Orbital size increases with energy level.

  14. C. Atomic Structure and the Periodic Table • Trends (cont.) - Atomic Radius • Atoms decrease in size as you move from left to right across the periodic table. • Why? • It is the result of increasing nuclear charge going left to right. As the number of protons increases, there is a stronger attraction for the outermost electrons. (which are located in the same energy level. The electrons will be held closer to the nucleus.

  15. C. Atomic Properties and the Periodic Table Ionization Energies • Ionization Energy – energy required to remove an electron from an individual atom (gas)

  16. C. Atomic Properties and the Periodic Table Ionization Energy

  17. C. Atomic Properties and the Periodic Table Ionization Energies Trends: • I.E. increases as you go from left to right on the periodic table Why? • Increasing nuclear charge as you move to the right; the nucleus has a stronger attraction for the outer electrons • I.E increases as you move up a group on the periodic table. Why? • Electrons are closer to the nucleus as you move up a group; there is a stronger attraction to the nucleus

  18. Atomic Properties and the Periodic Table Ionization Energies I.E. can also be explained in terms of atomic radius. • The smaller an atom is…the greater the attraction between the nucleus and outer electrons… and the greater the I.E.

  19. Practice • Which element (Cs, Hf, Au) has the smallest atomic radius? Au 2. Arrange the following elements in order of decreasing ionization energy: Li, O, C, K, Ne, F Ne, F, O, C, Li, K • For each of the following pairs of elements select the largest atom. C, O Sr, Mg Si, N Li, Mg

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