Transcript Trends & the Periodic Table

Trends & the Periodic Table
Trends
• More than 20 properties change in predictable
way based location of elements on PT
• Includes:
density, melting point, boiling point, atomic
radius, ionization energy, electronegativity
Atomic Radius
• Atomic radius: defined as ½ the distance
between neighboring nuclei in molecule or
crystal
• “size” varies a bit from substance to
substance
X-ray
diffraction
pinpoints
nuclei measure
distance
between them
Cannot measure
electron cloud
Trends:
Atoms get larger as go
down column:
↑# principal energy levels
Atoms get smaller as
move across series:
↑ PPP
“proton pulling power”
Going down column 1:
Period
Element
Configuration
1
H
1
2
Li
2-1
3
Na
2-8-1
4
K
2-8-8-1
5
Rb
2-8-18-8-1
6
Cs
2-8-18-18-8-1
7
Fr
2-8-18-32-18-8-1
increasing # energy levels as go down - makes sense that
atoms get larger in size
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Source: Conceptual Chemistry by John Suchocki
Li: Group 1 Period 2
Cs: Group 1 Period 6
Going across row 2:
Family
IA or 1
IIA or 2
IIIA or 13
Element
Li
Be
B
Configuration
2-1
2-2
2-3
IVA or 14
C
2-4
VA or 15
N
2-5
VIA or 16
O
2-6
VIIA or 17
F
2-7
VIIIA or 18
Ne
2-8
Atoms sizes actually get a bit smaller as you go across a
row What’s going on?
What do you remember about
charge?
• opposites attract
• like charges repel
• After # principal energy levels, next
largest influence on atomic size is
“proton pulling power (PPP)”
Effective nuclear charge
• Charge actually felt by valence electrons
• = Atomic Number minus # of inner shell
electrons
• Not same as nuclear charge or # protons in
nucleus
• Charge felt by valence electrons is attenuated
(shielded) by inner shell electrons
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H and He:
only elements
whose valence
electrons feel
full nuclear
charge (pull)
Source: Conceptual Chemistry by John Suchocki
previous | index | next
Source: Conceptual Chemistry by John Suchocki
Li’s valence e- feels effective nuclear charge of +1
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Source: Conceptual Chemistry by John Suchocki
Calculating “effective nuclear charge”
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as go across row size tends to decrease a bit because
of greater PPP “proton pulling power”
Source: Conceptual Chemistry by John Suchocki
previous | index | next
Source: Conceptual Chemistry by John Suchocki
size  as you go  and size  as you go 
Ionization Energy
• Amount energy required to remove
valence electron from atom in gas phase
• 1st ionization energy = energy required to
remove most loosely held valence electron
(e- farthest from nucleus)
Trends in ionization energy
• What do you think happens to the ionization
energy as go down column of PT?
• As go across row?
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
•Valence electrons in atoms feel effective nuclear charge of
+1
•Cs valence electron is lot farther away from nucleus
•electrostatic attraction much weaker so easier to steal
electron away from Cs
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
•easier to steal electron from Li than from Ne
•Li has smaller effective nuclear charge and valence
electron is farther away from nucleus
•Li has less “proton pulling power” than Ne
Trends in ionization energy
• Ionization energy decreases as go
down a column
– easier & easier to remove valence electron
• Ionization energy increases as go
across a row
– more difficult to remove valence electron
Periodic properties: Graph shows a repetitive pattern
(Note: Doesn’t have to be a straight line)
Electronegativity
• ability of atom to attract electrons in bond
• noble gases tend not to form bonds, so
they don’t have electronegativities
• Unit = Pauling
• Fluorine: most electronegative element
= 4.0 Paulings
Trends in electronegativity
• Related to PPP
• Increases as go across row
• Decreases as go down column
• Remember: F is most electronegative
element!
Reactivity of Metals
• Metals are losers!
• judge reactivity of metals by how easily
give up electrons
• most active metals are Fr (then Cs)
• For metals, reactivity increases as
ionization energy goes down
Trends for Reactivity of Metals
or
Metallic Character
• Increases as go down
–Easier to lose electrons!
• Decreases as go across
–Harder to lose electrons!
Reactivity of Non-metals
• Non-metals are winners!
• judge reactivity of non-metals by how
easily gain electrons
• F:most active non-metal
• For non-metals, reactivity increases
as electronegativity increases
Trend for Reactivity of Non-metals:
Depends on PPP
• Increases as go across
• Decreases as go down
– (Shielded by more inner-shell electrons)
Ionic Size Relative to Parent Atom
• Depends if (+) ion or (-) ion
• How do you make a positive ion?
Remove electrons
• How do you make a negative ion?
Add electrons
How do you know if an atom gains
or loses electrons?
• Think back to the Lewis structures of ions
• Atoms form ions to get a valence of 8
(or 2 for H)
• Metals tend to have 1, 2, or 3 valence electrons
– It’s easier to lose them
• Nonmetals tend to have 5, 6, or 7 valence electrons
– It’s easier to add some
• Noble gases already have 8 so they don’t form ions
very easily
Positive ions or cations
• Formed by loss of electrons
• Cations always smaller than parent
atom
Negative ions or anions
• Formed by gain of electrons
• Anions always larger than parent
atom
Allotrope
• Different forms of an element in the
same phase
– Have different structures and properties
• O2 and O3 - both gas phase
• Graphite, diamond – both solid
phase carbon
Graphite and Diamond