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Periodic Properties of the Elements

Periodic Properties of the Elements. Paramagnetic vs. Diamagnetic. Diamagnetic Elements. Have unpaired electrons Are weakly repelled by a magnetic field. Paramagnetic elements. Have unpaired electrons Attracted to a magnetic field Some can even become charged by a magnetic field.

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Periodic Properties of the Elements

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  1. Periodic Properties of the Elements

  2. Paramagnetic vs. Diamagnetic Diamagnetic Elements Have unpaired electrons Are weakly repelled by a magnetic field Paramagnetic elements • Have unpaired electrons • Attracted to a magnetic field • Some can even become charged by a magnetic field Examples: copper, neon, mercury Examples: iron, sodium, oxygen

  3. Effective Nuclear Charge Q = magnitude of charge on a particle d = the distance between their nuclei k = constant Conclusion: As the distance between the nucleus and an electron increases, their attractive force decreases Coulomb’s law: the magnitude of the electric force between two charged particles is given as F = kQ1Q2 d2

  4. Effective Nuclear Charge Shielding of inner electrons stays constant Nucleus increases + charge • We can estimate the net attraction between each electron and the nucleus • Electrons repel each other BUT • They are attracted to the protons in the nucleus • Always smaller than the actual nuclear charge due to ELECTRON SHIELDING • Zeffincreases as we move across a period • Why?

  5. Effective Nuclear Charge A proton is added to the nucleus as atomic number increases No new electrons are added to the core – shielding remains constant The distance between the nucleus and valence electrons increases significantly Increases as we move across a period WHY? Changes far less down a group WHY?

  6. Atomic Radius Measured by the distance between two nuclei of nonbonding atoms or bonding atoms

  7. General trends for Atomic Radius As n increases, the outer electrons have a greater probability of being further from the nucleus Zeff increases as we move across the period Electrons are drawn closer to the nucleus • Down a group: INCREASES WHY? • Across a period: DECREASES WHY?

  8. General Trends for Ionic Radii The most spatially extended orbital is vacated Amount of e- repulsions decreases More electron/electron repulsions occur Electrons spread out Cations: SMALLER than parent atoms WHY? Anions: LARGER than parent atoms WHY?

  9. General Trends for Ionization Energy The amount of energy required to remove an electron from the gas phase Of an atom or ion (in Joules) - With each added energy level, valence electrons are further from the nucleus - Less energy is needed to overcome the weaker force between the nucleus and the valence electrons - Zeff increases, pulling outer electrons closer to the nucleus - More energy is required to overcome that force • Down a group: DECREASES WHY? • Across a period: INCREASES WHY?

  10. Electron removal in atoms 4s Electrons in a higher principle energy level have a greater probability of being further from the nucleus Electrons in higher energy levels leave first These experience the least effective nuclear charge, and require the least amount of energy to remove s electrons leave before d electrons WHY? 3d

  11. Electron configurations in ions Write the configuration for vanadium ion, V+3 Write the configuration for the Se2- ion.

  12. Exceptions to trends: Ionization Energy Which has higher ionization energy: Mg or Al?

  13. Another exception example: Which has lower ionization energy: P or S?

  14. General exceptions to ionization energy trend In Group 13, the electron is removed from a p orbital, which is farther from the nucleus In Group 2, the electron is removed from an s orbital, which is closer To the nucleus • In Group 16, the electron is easier to remove because one of the p orbitals has • two electrons, giving more electron repulsions and increasing the distance • between the nucleus and the outer electrons • - In Group 15, the electron harder to remove, because the electrons are closer to • the nucleus • Group 2 to Group 13  DECREASE (expect increase!) • WHY? • Group 15 to Group 16  DECREASE (expect increase!) • WHY?

  15. Electron Affinity The attraction or the atom for an added electron Measured in Joules The ease with which an atom GAINS an electron

  16. General trends in Electron Affinity Outer energy level is farther away from nucleus Nucleus is less attracted to outer electrons Zeff increases Nucleus is more attracted to outer electrons • Down a group: becomes more POSITIVE (endothermic) WHY? • Across a period: becomes more NEGATIVE (exothermic) WHY?

  17. Exceptions to trend: Electron affinity • Noble gases: Have positive electron affinities • WHY? • Adding an electron requires a new principal energy level • All other electrons are in the core and greatly shield the nucleus from the valence electrons • Other exceptions? • Group 2 is more positive than Group 1 • Group 15 is more positive than the Group 14

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