1. Section 14.2 Periodic Trends
OBJECTIVES:
Interpret group trends in atomic radii, ionic radii, ionization energies, and electronegativities.
Interpret period trends in atomic radii, ionic radii, ionization energies, and electronegativities.
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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.
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3. Atomic Size }
Radius
Atomic Radius = half the distance between two nuclei of atoms
in the solid state (by X-ray diffraction)
or of a diatomic molecule.
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4. Atomic Size Influenced by three factors: Energy Level
Higher energy level is further away.
Charge on nucleus
More positive charge pulls electrons in closer.
Shielding effect
The inner electrons shield the outer electrons from the nuclear charge/attraction.
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5. Shielding
The electron on the outermost energy level has to look through all the other energy levels to see the nucleus.
Second electron has same shielding, if it is in the same period
Shielding Increases down a Group, and is Constant across a Period.
Shielding Across a Group is Constant, but the EFFECTIVE NUCLEAR CHARGE INCREASES.
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6. Group Trends in Atomic Size
As we go down a group...
each atom has another energy level,
so the atoms get bigger.
Shielding increases as well, so the nucleus has less of a hold on e-… distance is longer.
The Increased size of the Energy Levels down a group outweighs the increased nuclear charge H Li Na K Rb
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7. Periodic Trends in Atomic Size
As you go across a period, the radius gets smaller.
Electrons are in same energy level.
More nuclear charge.
Shielding is constant … not an issue.
Outermost electrons are closer. Na Mg Al Si P S Cl Ar
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8. Atomic Radius Overall
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9. Overall Periodic Trend for Atomic RadiiRb K
Overall Periodic Trend for Atomic Radii Na Xe Li
Atomic Radius (nm) Kr Ar Ne H He 10
Atomic Number
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10. Trends in Ionization Energy
The amount of energy required to completely remove an electron from a gaseous atom.
Removing one electron makes a 1+ ion.
The energy required to remove the first electron is called the first ionization energy.
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11. Ionization Energy
The second IE is the energy required to remove the second electron.
Always greater than first IE.
The third IE is the energy required to remove a third electron.
Greater than 1st or 2nd IE.
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12. Ionization Energy Table Where are the Group Effects?
Symbol First Second Third H
HeLi
BeB C N O
F Ne
900
4619
4577
5301
6045
6276 1A 2A 3A
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13. Can we See the Effect of Nuclear Charge in this Table?
Symbol First Second Third H
HeLi
BeB C
NO F Ne
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14. What Affects the IE
The greater the nuclear charge, the greater the IE.
Larger positive nucleus has a greater attraction for the electrons, so the IE increases.
Greater distance from nucleus decreases IE
Electrons are further away from the attractive nucleus, and are easier to remove.
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15. Let’s look at Group & Period Trends for IE
What Affects the IE
Filled and half-filled orbitals have lower energy, so the removal of an electron to achieve this ½ filled orbital requires unusually low IE.
Shielding effect
As Shielding increases, it is easier to “pluck” the outer electron, so the IE would decrease.
Let’s look at Group & Period Trends for IE
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16. Group trends on IE H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87
As you go down a group, first IE decreases because...
The electron is further away.
More shielding.
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17. Periodic trends on IE
Across the representative elements, the atoms are in the same period & have the same energy level.
Same shielding.
But, increasing nuclear charge holds e-’s tighter.
So IE generally increases from left to right.
Exceptions at full and 1/2 full orbitals.
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18. Summarizing 1st Ionization Energy
1st Ionization Energy Increases
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19. He has a greater IE than H. same shielding greater nuclear charge
First Ionization Energy He
He has a greater IE than H.
same shielding
greater nuclear charge H
First Ionization energy
Atomic number
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20. these outweigh greater nuclear charge
First Ionization Energy He
Li has lower IE than H
more shielding
further away
these outweigh greater nuclear charge H
First Ionization energy Li
Atomic number
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21. greater nuclear charge
First Ionization Energy He
Be has higher IE than Li
same shielding
greater nuclear charge
First Ionization energy H Be Li
Atomic number
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22. First Ionization EnergyHe
B has lower IE than Be
same shielding
greater nuclear charge
By removing an electron we make the s orbital half-filled
First Ionization energy H Be B Li
Atomic number
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23. First Ionization EnergyHe
First Ionization energy H C Be B Li
Atomic number
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24. First Ionization EnergyHe N
First Ionization energy H C Be B Li
Atomic number
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25. First Ionization EnergyHe
Breaks the pattern, because removing an electron leaves 1/2 filled p orbital N
First Ionization energy H C O Be B Li
Atomic number
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26. First Ionization EnergyHe F N
First Ionization energy H C O Be B Li
Atomic number
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27. Ne has a lower IE than He Both are full, Ne has more shielding
First Ionization Energy He Ne
Ne has a lower IE than He
Both are full,
Ne has more shielding
Greater distance F N O
First Ionization energy H C Be B Li
Atomic number
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28. Na has a lower IE than Li Both are s1 Na has more shielding
First Ionization Energy He Ne
Na has a lower IE than Li
Both are s1
Na has more shielding
Greater distance F N O
First Ionization energy H C Be B Li Na
Atomic number
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29. First Ionization EnergyHe Ne
First Ionization Energy Ar Kr
First Ionization energy Li Na K
Atomic number
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30. What’s the 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.
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31. 2nd Ionization Energy
For elements that reach a filled or half-filled orbital by removing 2 electrons, 2nd IE is lower than expected.
True for s2
Alkaline earth metals form 2+ ions.
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32. 3rd Ionization Energy
Using the same logic s2p1 atoms have a low 3rd IE.
Atoms in the aluminum family form 3+ ions.
2nd IE and 3rd IE are always higher than 1st IE!!!
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33. Trends in Electron Affinity
What is Electron Affinity?
It’s the energy change associated with adding an electron to a gaseous atom.
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34. Trends in Electron Affinity
It’s easiest to add an electron to Group 7A.
It gets them to a full energy level, or completes the OCTET.
Increase from left to right: atoms become smaller, with greater nuclear charge.
Decrease as we go down a group.
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35. Electron Affinity in the Periodic Table
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36. Electron Affinity in 3-D
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37. 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.
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38. Trends in Ionic Size Anions form by gaining electrons.
Anions are bigger that the atom they come from.
Nonmetals form anions.
Anions of representative elements have noble gas configuration.
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39. Revisiting Configuration of Ions
Ions always have noble gas configuration.
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.
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40. Revisiting 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.
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41. Group Trends Li1+ Na1+ K1+ Rb1+ Cs1+
Going down a Group, you are adding energy levels
Ions get bigger as you go down.
Li1+
Na1+
K1+
Rb1+
Cs1+
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42. 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+
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43. 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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44. Size of Isoelectronic ions
Positive ions that have more protons would be smaller.
Increase in size from most positive to most negative
N3-
O2-
F1- Ne
Na1+
Al3+
Mg2+
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45. Electronegativity
The tendency for an atom to attract electrons to itself when it is chemically combined with another element.
How fair is the sharing?
Big electronegativity means it pulls the electron towards it.
Atoms with large negative electron affinity have larger electronegativity.
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46. Electronegativity Group Trend
The further down a group, the farther the electron is away, and the more electrons an atom has.
Pull/Attraction of the positive nucleus is lessened due to increased distance and Shielding.
Electronegativity decreases.
More willing to share.
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47. Electronegativity Period Trend
As you move across a Period, there are the same number of energy levels, the same shielding, however …
Pull/Attraction of the positive nucleus on other’s electrons increases as the nucleus gets larger
Electronegativity Increases.
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48. Electronegativity 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.
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49. Electronegativity in 3-D
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50. Can We Possibly Summarize all of this Stuff???
Ionization Energy, Electronegativity, and Electron Affinity Increases
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51. Can We Possibly Summarize all of this Stuff???
Atomic size & Ionic Size Increases
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52. Can We Possibly Summarize all of this Stuff???
Shielding is constant
Shielding Increases
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