Transcript Trends & the Periodic Table

Trends & the Periodic Table
Trends
• see properties change in predictable ways
based location of elements on PT
• some properties can be predicted:
density
melting point/boiling point
* atomic radius
* ionization energy
TABLE S
• electronegativity
anyone know where we can find these numbers
Periodic properties: Graph shows a repetitive pattern
(Note: Doesn’t have to be a straight line)
When you’re
done it will
look like this so
leave room for
writing!
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
Increasing number
of energy levels
Atomic Radius
• atomic radius: defined as ½ distance
between neighboring nuclei in molecule or
crystal
• affected by
1. # energy
levels
2. Proton Pulling
Power (PPP)
TRENDS:
atoms get larger as go
down column:
↑# principal energy levels
atoms get smaller as move
across series:
↑ PPP
“proton pulling power”
Cs has more energy levels, so it’s bigger
Li: group 1 period 2
Cs: group 1 period 6
Increasing Atomic
Radius
Increasing number of energy levels
As we go across, elements gain electrons, but
they are getting smaller! What is happening?
Family
IA or 1
IIA or 2
IIIA or 13
IVA or 14
VA or 15
VIA or 16
VIIA or 17
VIIIA or 18
Element
Li
Be
B
C
N
O
F
Ne
Configuration
2-1
2-2
2-3
2-4
2-5
2-6
2-7
2-8
Increasing number of energy levels
Increasing Atomic Radius
Decreasing
Atomic
Radius
Why does this happen..
• as go from left to right, you gain more
protons (atomic number increases)
• results in greater “proton pulling power”
– remember: nucleus is (+) and electrons are (-)
so e- get pulled towards the nucleus
• more protons you have, the stronger PPP
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as go across row size tends to decrease a bit
because of greater PPP “proton pulling power”
We can “measure” the PPP
by determining the
effective nuclear charge
• this is charge actually felt by valence electrons
• equation to calculate effective nuclear charge:
nuclear charge - # inner shell electrons
(doesn’t include valance e-)
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+7
+1
calculate “effective nuclear charge”
• # protons minus # inner electrons
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H and He:
only elements
whose valence
electrons feel
full nuclear
charge (pull)
NOTHING
TO
SHIELD
THEM
Increasing number of energy levels
Increasing Atomic Radius
Decreasing Atomic Radius
Increased Electron
Shielding
Look at all the shielding Francium's one valance
electron has. It barely feels the proton pull from
the nucleus. No wonder it will lose it’s one
electron the easiest. No wonder it’s the most
reactive metal
Ionization Energy
• definition: amount energy required to
remove farthest valence e- from atom
• 1st ionization energy: energy required to
remove most loosely held valence electron
(valence e- farthest from nucleus)
Trends in Ionization Energy
• What do you think happens to the ionization
energy as go down column of PT?
decreases
• As go across row?
increases
Electronegativity
• ability of atom to attract electrons to itself
so can form bonds with other elements
(to create cmpds)
• noble gases tend not to form bonds, so
don’t have electronegativity values
• Fluorine: most electronegative element
= 4.0 Paulings
• Francium: least electronegative element
= 0.7 Paulings
Increasing number of energy levels
Increasing Atomic Radius
Increasing electron shielding
Decreasing Atomic Radius
Increasing Ionization Energy
Increasing Electronegativity
due to  PPP
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elements in same group: farther away valence electrons
are from nucleus the easier to remove them
easier for Cs (top of column) to lose electrons than
Li (bottom of column) so Cs is a more reactive metal!
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elements in same row: easier to take away valence
electrons when have less protons
Li has less “proton pulling power” so easier to remove
its valence electrons
Reactivity of Metals
• metals are losers!
• judge reactivity of metals by how easily
give up electrons to form (+) ions
• most active metals: Fr (then Cs)
• for metals, reactivity increases as
ionization energy goes down
Trends for Reactivity
(Metallic Character) of Metals
• increases as go down column
–easier to lose electrons!
• decreases as go across row
–more difficult 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 ↑ as electronegativity ↑
Trend for Reactivity of Non-metals:
depends on PPP
• increases as go across row
• decreases as go down column
– (shielded by more inner-shell 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 these than gain extra needed
• non-metals tend to have 5, 6, or 7 valence electrons
– it’s easier to add extra needed than to lose these
• noble gases already have 8 so they don’t form ions
very easily
positive ions (cations)
• formed by loss of electrons
• cations always smaller than parent
atom
2e
8e
8e
2e
Ca
Ca
8e
8e
2e
Ca+2
negative ions or (anions)
• formed by gain of electrons
• anions always larger than parent
atom
Allotropes
• different structural forms of element in
same phase
– different structures and properties
– examples: C and O
Graphite and Diamond: both carbon in solid form
O2 (g) and O3 (g)
O2 (oxygen) - necessary for life
O3 (ozone) - toxic to life