Quantum Mechanical Atom Part III: Trends in the Periodic Table Chapter 7 Section 8 of Brady

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Quantum Mechanical AtomPart III Trends in the
Periodic TableChapter 7 Section 8of Brady
Senese 5TH ed)

• C. Yau
• Spring 2009

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Table of Elements

• The necessity of a table of elements arose from
  the need to find a pattern in the various
  properties of the elements.
• Science involves looking for patterns.
• Learning about patterns allow us to place
  elements in categories. We would need to learn
  only the general trends and not try to memorize
  the properties of every single element.

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Mendeleev's Periodic Table

• There were many other attempts at writing a table
  of elements.
• Mendeleev's was the most successful in helping us
  predict phys. chem. properties.
• (1) He placed the elements in order of
  increasing mass (protons were not discovered
  yet).
• (2) He then placed elements of similar properties
  in the same column.
• He found that the properties start repeating in
  cycles of 8. (Every 8th element has very similar
  properties.)
• These cycles are what we refer to as "periods"
  and hence, the PERIODIC TABLE.

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The Modern Day Periodic Table

• There were discrepancies in Mendeleev's table
  based on mass.
• Our current periodic table is based on increasing
  atomic number rather than mass. Why does it
  work?
• Atomic electrons
• Reactions involve addition, removal or sharing of
  electrons.
• electrons determines a lot of the chemistry.
• Mass determines HOW MUCH is made, not WHAT is
  made.

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• Group valence electrons
• Period n of the outershell

Valence electrons of S 3s2 3p4 (6 electrons in
n 3) Valence electrons of iodine
5s2 5p5 (7 electrons in n 5)
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• Electron Configuration of Valence
  Electrons
• Group IA ns1 Group VA ns2 np3
• Group IIA ns2 Group VIA
• Group IIIA ns2 np1 Group VIIA
• Group IVA ns2np2 Group VIIIA

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Electron Dot Symbols

• Gp IA ns1
• Gp IIA ns2
• Gp IIIA ns2 np1
• Gp IVA ns2 np2

Gp VA ns2 np3 Gp VIA ns2 np4 Gp VIIA ns2
np5 Gp VIIIA ns2 np6
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• Halogens
• F 2s2 2p5
• Cl 3s2 3p5
• Br 4s2 4p5
• I 5s2 5p5

Halides F- 2s2 2p6 Cl- 3s2 3p6 Br- 4s2
4p6 I- 5s2 5p6
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Fig. 11.22 Atomic radii in picometers (1 m
1012 pm)
Size Incr
Size Increase
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We summarize the periodic trend of atomic size by
drawing a diagonal arrow.
smallest atoms
largest atoms
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Why are there such trends in size? To explain why
atomic size increases going down any column 1)
As n increases, outershell e- are further from
the nucleus. 2) As n increases, there are also
more innershells of electrons shielding the
protons from the electrons.
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• Consider N and P
• radii 74 pm 110 pm

innershell shielding of nuclear charge
1s2 1s2
2s2 2p6
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To explain why atomic size increases going from
right to left in the periodic table, As you move
right to left of any row, there is a decrease in
the of protons, which leads to a decrease in
charge in the nucleus (the nuclear charge)
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• Consider N and F
• radii 74 pm 72 pm

Outershell is the same (n 2) Innershell is the
same (1s2) What is different? protons
nuclear charge
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• Consider N and F
• radii 74 pm 72 pm

nuclearcharge 7 p 9 p Higher
nuclear charge pulls electrons closer to the
nucleus.
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Ion vs. Atom Radii

• Positive ions are always smaller than the atoms
  from which they are formed
• due to decreased shielding effects
• Negative ions always larger than the atoms from
  which they are formed
• due to increased electron repulsion

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Ionization energy (IE)

• IE is the energy required to remove an electron
  from an isolated, gaseous atom
• Successive ionizations are possible until no
  electrons remain
• Trends in IE are the opposite of the trends in
  atomic size
• Why?
• Valence electrons are closer to the nucleus for
  the smaller atoms and are held tighter by the
  nuclear charge.
• They are harder to pull off HIGHER IE

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Trends in IE
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Successive Ionization Energy (IE) Mg (g)
energy Mg (g) e- Mg (g)
energy Mg2 (g) e- Mg2 (g)
energy Mg3 (g) e- 1st IE
737 kJ/mol 2nd IE 1450 kJ/mol 3rd IE 7731
kJ/mol 2nd IE gt 1st IE Why? 3rd IE gtgtgtgt 2nd IE
Why?
1st IE
2nd IE
3rd IE
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Successive IE
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Why the sudden jump in IE's?
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Irregularities in I.E.
Easier to pull e- off B than from Be b/c it's
from 2p not 2s.
Easier to pull e- off O than from N b/c its from
orbital with 2 e- instead of 1 e-
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Periodic Trend in the Ionization Energy
highest IE
lowest IE
Nobel gases are excluded.
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Electron Affinity (EA) Definition Electron
affinity is the energy released when an electron
is added to a neutral gaseous atom. EA of the 1st
electron is almost always negative (exothermic)
F e- F- energy (EA of F
is -328 kJ/mole.) Addition of subsequent
electrons always requires energy (EA's are
positive). We will focus on only the 1st EA.
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We talk about how negative the electron affinity
is. The more negative the EA, the more the
element likes electrons. It gets a bit confusing
with the negative sign, so we sometimes refer to
the absolute value of the EA. This gives us the
"magnitude" of the EA.
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• For example, EA of O is -141 kJ/mol
• EA of F is -328 kJ/mol
• EA of F is more negative than O which means it
  likes electrons more.
• Technically, the EA of F is smaller because it is
  a smaller number.
• It is misleading to say F has a lower electron
  affinity, when it actually has a higher
  "affinity" for electrons.
• To avoid the confusion, we talk in terms of the
  magnitude of O is 141 kJ/mol,
• and the magnitude of F is 328 kJ/mol.

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For example, EA of F is -328 kJ/mol
EA of Cl is -348 kJ/mol We say that
the magnitude of the EA of chlorine is larger
than that of fluorine. (absolute value is
larger) Which do you predict to have a higher
magnitude of EA? the smaller atom or the larger
atom? why? Draw a diagram to show the shells of
electrons in F and Cl.
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• Periodic Trend of the Electron Affinity

F has the largest EA in magnitude.
lowest EA in magnitude
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Summary of the periodic trends Draw in the arrows
to show the trends for 1) atomic size 2) IE
3) magnitude of EA
What generalization can you make about the IE of
metals and nonmetals?
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Compare the electron affinity for metals vs.
nonmetals
What generalizations can you make of the
metalloids concerning IE and EA?
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Metals vs. Nonmetals

• Metals have low ionization energies.
• It is easier to strip off electrons from metals.
• They tend to form CATIONS.
• Nonmetals have high magnitude of electron
  affinities.
• It is easier to add electrons to nonmetals.
• They tend to form ANIONS.
• If there is no source of electrons to be added,
  they tend to SHARE electrons and form COVALENT
  BONDS (as molecules or polyatomic ions).