1. AP CHEMISTRY 2010-2011 Ms. Paskowski
THE PERIODIC TABLE
AP CHEMISTRY
Ms. Paskowski

2. History of the Periodic Table
Many early scientists organized the elements in a table.
Mendeleev
organized the elements according to physical and chemical properties.
Correctly predicted the existence of previously undiscovered elements and their properties

3. Figure 7.24 Mendeleev's Early Periodic Table, Published in 1872

4. Periodic Trends Ionization Energy Electron Affinity Atomic Radius
Ionic Size
Valence Electrons
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5. Ionization Energy
Energy required to remove an electron from a gaseous atom or ion.
In general, as we go across a period from left to right, the first ionization energy increases.
In general, as we go down a group from top to bottom, the first ionization energy decreases.
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6. Ionization Energy Why does the IE decrease down a group?
Why does the IE increase across a period?
The effective nuclear charge increases across a period and decreases down a column.
What is effective nuclear charge?

7. The Values of First Ionization Energy for the Elements in the First Six Periods
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8. Concept Check
Explain why the graph of ionization energy versus atomic number (across a row) is not linear.
Where are the exceptions?
The graph is not linear due to electron repulsions. Some exceptions include from Be to B and N to O.
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9. Concept Check
Which atom would require more energy to remove an electron? Why?
Li Cs
Li would require more energy to remove an electron because the outer electron is on average closer to the nucleus (so more tightly bound).
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10. Concept Check Which has the larger second ionization energy? Why?
Lithium or Beryllium
Lithium has the larger second ionization energy because then a core electron is trying to be removed which will require a lot more energy than a valence electron.
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11. Successive Ionization Energies (KJ per Mole) for the Elements in Period 3
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12. Electron Affinity
Energy change associated with the addition of an electron to a gaseous atom.
In general as we go across a period from left to right, the electron affinities become more negative.
In general electron affinity becomes more positive in going down a group.
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13. Figure The Electron Affinity Values for Atoms Among the First 20 Elements that Form Stable, Isolated X- Ions

14. Atomic Radius
In general as we go across a period from left to right, the atomic radius decreases.
In general atomic radius increases in going down a group.
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15. Figure The Radious of an Atom (r) is Defined as Half the Distance Between the Nuclei in a Molecule Consisting of Identical Atoms

16. Concept Check Which should be the larger atom? Why? Na Cl 7.12
Na should be the larger atom because the electrons are not bound as tightly due to a smaller effective nuclear charge.
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17. Concept Check Which is larger? The hydrogen 1s orbital
The lithium 1s orbital
Which is lower in energy?
The hydrogen 1s orbital is larger because the electrons are not as tightly bound as the lithium 1s orbital (lithium has a higher effective nuclear charge and will thus draw in the inner electrons more closely).
The lithium 1s orbital is lower in energy because the electrons are closer to the nucleus.
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18. Atomic Radius of a Metal
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19. Atomic Radius of a Nonmetal
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20. Atomic Radii for Selected Atoms

21. Exercise
Arrange the elements oxygen, fluorine, and sulfur according to increasing:
Ionization energy
Atomic size
Ionization Energy: S, O, F
Atomic Size: F, O, S
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22. Final Thoughts
It is the number of valence electrons that chemists use to explain an atom’s chemistry.
BUT do not forget the involvement of the nucleus.
Electrostatic interaction of positive and negative charges is the fundamental force that explains chemical interactions.
i.e., the nucleus of one atom MUST attract the electrons of another atom to create a chemical bond . . .
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