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Chemical Periodicity? What?
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Periodicity When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities. Elements in the same group have similar chemical properties.
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Periodicity When elements are organized in order of atomic mass, and
grouped by similarities of chemical properties, a certain “pattern” or periodicity of properties becomes evident Draft for first version of Mendeleev's periodic table (17 February 1869).
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Review of the Periodic Table
Group IA – alkali metals Review of the Periodic Table
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Group IA – alkali metals
H s1 Li s22s1 Na11 - 1s22s22p63s1 K s22s22p63s23p64s1 Rb37 - 1s22s22p63s23p63d104s24p65s1 Cs55 - 1s22s22p63s23p63d104s24p64d105s25p66s1 Fr s22s22p63s23p63d104s24p64d104f145s25p65 d106s26p67s1
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Review of the Periodic Table
Group IIA – alkaline earth metals Review of the Periodic Table
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Group IIA – alkaline earth metals
Be s2 2s2 Mg12 - 1s22s22p63s2 Ca20 - 1s22s22p63s23p64s2 Rb38 - 1s22s22p63s23p63d104s24p65s2 Cs55 - 1s22s22p63s23p63d104s24p64d105s25p66s2 Fr s22s22p63s23p63d104s24p64d104f145s25p65 d106s26p67s2
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Review of the Periodic Table
Group VIII – Noble Gases Review of the Periodic Table
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Group VIII – Noble Gases
He s2 Ne10 - 1s22s22p6 Ar18 - 1s22s22p63s23p6 Kr36 - 1s22s22p63s23p63d104s24p6 Cs54 - 1s22s22p63s23p63d104s24p64d105s25p6 Rn s22s22p63s23p63d104s24p64d104f145s25p65 d106s26p6
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Review of the Periodic Table
Group VII – Halogens Review of the Periodic Table
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Group VII – Halogens F9 - 1s22s22p5 Cl17 - 1s22s22p63s23p5
Br35 - 1s22s22p63s23p63d104s24p5 I s22s22p63s23p63d104s24p64d105s25p5 At85 - 1s22s22p63s23p63d104s24p64d104f145s25p65 d106s26p5
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Review of the Periodic Table
Group VI
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Group VI O8 - 1s22s22p4 S16 - 1s22s22p63s23p4 Se34 - 1s22s22p63s23p63d104s24p4 Te52 - 1s22s22p63s23p63d104s24p64d105s25p4 Po84 - 1s22s22p63s23p63d104s24p64d104f145s25p65 d106s26p4
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Atomic Radii
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Group trends As you go down a group, first IE decreases because...
The electron is further away. More shielding. The electron in the outermost energy level experiences more inter-electron repulsion (shielding).
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Group trends As we go down a group...
H As we go down a group... each atom has another energy level, so the atoms get bigger. Li Na K Rb
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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.
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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
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Ionization Energy Energy needed to remove an electron from a gaseous atom: M(g) + I M+(g) + e-; (I = ionization energy)
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The Periodic Trend of Ionization Energy
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He has a greater IE than H. same shielding greater nuclear charge
First Ionization energy Atomic number
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Outer electron further away outweighs greater nuclear charge
Li has lower IE than H Outer electron further away outweighs greater nuclear charge H First Ionization energy Li Atomic number
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greater nuclear charge
He Be has higher IE than Li same shielding greater nuclear charge First Ionization energy H Be Li Atomic number
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greater nuclear charge
He B has lower IE than Be same shielding greater nuclear charge p orbital is slightly more energy and its electron easier to remove First Ionization energy H Be B Li Atomic number
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First Ionization energy
He First Ionization energy H C Be B Li Atomic number
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First Ionization energy
He N First Ionization energy H C Be B Li Atomic number
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First Ionization energy
He Breaks the pattern, because the outer electron is paired in a p orbital and experiences inter-electron repulsion. N First Ionization energy H C O Be B Li Atomic number
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First Ionization energy
He F N First Ionization energy H C O Be B Li Atomic number
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Ne has a lower IE than He Both are full, Ne has more shielding
Greater distance F N First Ionization energy H C O Be B Li Atomic number
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Na has a lower IE than Li Both are s1 Na has more shielding
He Ne Na has a lower IE than Li Both are s1 Na has more shielding Greater distance F N First Ionization energy H C O Be B Li Na Atomic number
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Periodic Patterns and Ionization Energy
In group: ionization energy as you move down the column Across period: Ionization energy as you move from left to right Rationale: group –outermost shell electrons farther away from nucleus row – outermost electron more tightly held
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Ions When atoms lose or gain electrons, they become ions.
Cations are positive and are formed by elements on the left side of the periodic chart. Anions are negative and are formed by elements on the right side of the periodic chart.
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Ionic Radius Cations form by losing electrons.
190pm vs pm 97pm vs pm Cations form by losing electrons. Cations are smaller that the atom they come from. Metals form cations. Anions form by gaining electrons. Anions are bigger that the atom they come from. Nonmetals form anions
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The Periodic Table Always remove electrons from the highest n value?
Na is: 1s22s22p63s1 Forms a 1+ ion: 1s22s22p6 Same configuration as neon (isoelectronic with NEON. Ions have noble gas configurations (not transition metals)
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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
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Size of Isoelectronic ions
Positive ions that have more protons would be smaller. N3- O2- F1- Ne Na1+ Al3+ Mg2+
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Periodicity and Electronegativity
Electronegativity ≡ a measure of how tightly held an electron is.
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Periodicity and Electronegativity
In group: electronegativity as you move down the column Across period: Electronegativity as you move from left to right Rationale: group –atoms more willing to accept another electrons due to nuclear charge
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Periodic Trends – Summary
Decreasing Atomic Radii Decreasing Ionic Size Increasing Ionization Energy Increasing Electronegativity Decreasing Ionization Energy Decreasing Electronegativity Increasing Atomic Radii Increasing Ionic Radius
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