1 / 41

Periodic Trends

Learn about the interpretation of group trends in atomic radii, ionic radii, ionization energies, melting point, boiling point, electronegativity, and chemical properties. Understand the factors influencing atomic size and ionization energy, and how these trends change across periods and groups.

karenblack
Download Presentation

Periodic Trends

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Periodic Trends • OBJECTIVES: • Interpret group trends in atomic radii, ionic radii, ionization energies, m.p., b.p., electronegativity and chemical properties

  2. Trends in Atomic Size • First problem: Where do you start measuring from? • The electron cloud doesn’t have a definite edge. • They get around this by measuring more than 1 atom at a time.

  3. Atomic Size } • Atomic Radius = half the distance between two nuclei of a diatomic molecule. Radius

  4. Trends in Atomic Size • Influenced by three factors: 1. Energy Level • Higher energy level is further away. 2. Charge on nucleus • More charge pulls electrons in closer. • 3. Shielding effect e <-> e repulsion

  5. Group trends H • As we go down a group... • each atom has another energy level, • so the atoms get bigger. Li Na K Rb

  6. Periodic Trends • As you go across a period, the radius gets smaller. • Electrons are in same energy level. • More nuclear charge. • Outermost electrons are closer. Na Mg Al Si P S Cl Ar

  7. Rb Overall K Na Li Atomic Radius (nm) Kr Ar Ne H Atomic Number 10

  8. Trends in Ionization Energy • The amount of energy required to completely remove a mole of electrons from a mole of gaseous atoms. • Removing an electron makes a +1 ion. • The energy required to remove (1 mole of) the first electron is called the first ionization energy.

  9. Ionization Energy • The second ionization energy is the energy required to remove (1 mole of) the second electron(s). • Always greater than first IE. • The third IE is the energy required to remove a third electron. • Greater than 1st or 2nd IE.

  10. Symbol First Second Third 5247 7297 1757 2430 2352 2857 3391 3375 3963 1312 2731 520 900 800 1086 1402 1314 1681 2080 HHeLiBeBCNO F Ne 11810 14840 3569 4619 4577 5301 6045 6276

  11. Symbol First Second Third 11810 14840 3569 4619 4577 5301 6045 6276 5247 7297 1757 2430 2352 2857 3391 3375 3963 1312 2731 520 900 800 1086 1402 1314 1681 2080 HHeLiBeBCNO F Ne

  12. What determines IE • The greater the nuclear charge, the greater IE. • Greater distance from nucleus decreases IE • Filled and half-filled orbitals have lower energy, so achieving them is easier, lower IE. • Shielding effect

  13. Shielding • The electron in the outermost energy level experiences more inter-electron repulsion (shielding). • Second electron has same shielding, if it is in the same period

  14. Group trends • As you go down a group, first IE decreases because... • The electron is further away. • More shielding.

  15. Periodic trends • All the atoms in the same period have the same energy level. • Same shielding. • But, increasing nuclear charge • So IE generally increases from left to right. • Exceptions at full and 1/2 full orbitals.

  16. He • He has a greater IE than H. • same shielding • greater nuclear charge H First Ionization energy Atomic number

  17. He • Li has lower IE than H • Outer electron further away • outweighs greater nuclear charge H First Ionization energy Li Atomic number

  18. He • Be has higher IE than Li • same shielding • greater nuclear charge H First Ionization energy Be Li Atomic number

  19. He • B has lower IE than Be • same shielding • greater nuclear charge • p orbital is slightly more diffuse and its electron easier to remove H First Ionization energy Be B Li Atomic number

  20. He C H First Ionization energy Be B Li Atomic number

  21. He N C H First Ionization energy Be B Li Atomic number

  22. He • Breaks the pattern, because the outer electron is paired in a p orbital and experiences inter-electron repulsion. N O C H First Ionization energy Be B Li Atomic number

  23. He F N O C H First Ionization energy Be B Li Atomic number

  24. Ne He F • Ne has a lower IE than He • Both are full, • Ne has more shielding • Greater distance N O C H First Ionization energy Be B Li Atomic number

  25. Ne He • Na has a lower IE than Li • Both are s1 • Na has more shielding • Greater distance F N O C H First Ionization energy Be B Li Na Atomic number

  26. First Ionization energy Atomic number

  27. Driving Force • Full Energy Levels require lots of energy to remove their electrons. • Noble Gases have full orbitals. • Atoms behave in ways to achieve noble gas configuration.

  28. Trends in Electron Affinity • The energy change associated with adding an electron (mole of electrons) to a (mole of) gaseous atom(s). • Easiest to add to group 7A. • Gets them to full energy level. • Increase from left to right: atoms become smaller, with greater nuclear charge. • Decrease as we go down a group.

  29. Trends in Ionic Size • Cations form by losing electrons. • Cations are smaller that the atom they come from. • Metals form cations. • Cations of representative elements have noble gas configuration.

  30. Ionic size • Anions form by gaining electrons. • Anions are bigger that the atom they come from. • Nonmetals form anions. • Anions of ‘A’ groups elements have noble gas configuration.

  31. Configuration of Ions • Ions have noble gas configurations (not transition metals). • Na is: 1s22s22p63s1 • Forms a 1+ ion: 1s22s22p6 • Same configuration as neon. • Metals form ions with the configuration of the noble gas before them - they lose electrons.

  32. Configuration of Ions • Non-metals form ions by gaining electrons to achieve noble gas configuration. • They end up with the configuration of the noble gas after them.

  33. Group trends • Adding energy level • Ions get bigger as you go down. Li1+ Na1+ K1+ Rb1+ Cs1+

  34. Periodic Trends • Across the period, nuclear charge increases so they get smaller. • Energy level changes between anions and cations. N3- O2- F1- B3+ Li1+ C4+ Be2+

  35. Size of Isoelectronic ions • Iso- means the same • Iso electronic ions have the same # of electrons • Al3+ Mg2+ Na1+ Ne F1- O2- and N3- • all have 10 electrons • all have the configuration: 1s22s22p6

  36. Size of Isoelectronic ions • Positive ions that have more protons would be smaller. N3- O2- F1- Ne Na1+ Al3+ Mg2+

  37. Electronegativity • The tendency for an atom to attract electrons to itself when it is chemically combined with another element. • High electronegativity means it pulls the electron toward it. • Atoms with large negative electron affinity have larger electronegativity.

  38. Group Trend • The further down a group, the farther the electron is away, and the more electrons an atom has. • More willing to share. • Low electronegativity.

  39. Periodic Trend • Metals are at the left of the table. • They let their electrons go easily • Low electronegativity • At the right end are the nonmetals. • They want more electrons. • Try to take them away from others • High electronegativity.

  40. Ionization energy, Electronegativity, and Electron Affinity INCREASE

  41. Atomic size increases, shielding constant Ionic size increases

More Related