Transcript Trends & the Periodic Table

Trends & the Periodic Table
Trends
• more than 20 properties change in predictable
way based location of elements on PT
• some properties:
– Density
– melting point/boiling point
– atomic radius
– ionization energy
– electronegativity
Atomic Radius
• Atomic radius: defined as ½ distance
between neighboring nuclei in molecule or
crystal
• “size” varies a bit from substance to
substance
X-ray
diffraction
pinpoints
nuclei – then
measures
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
previous | index | next
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
size atoms actually get a bit smaller as go across row left to
right - what’s going on?
What do you remember about
charge?
• opposites attract
• like charges repel
• largest influence on atomic size in
order:
• # principal energy levels
• “proton pulling power (PPP)”
Effective nuclear charge
• Charge actually felt by valence electrons
• = Atomic Number minus # inner shell electrons
• Not same as nuclear charge or # protons in
nucleus
• Charge felt by valence electrons is attenuated
(shielded) by inner shell electrons
previous | index | next
H and He:
only elements
whose valence
electrons feel
full nuclear
charge (pull)
previous | index | next
Li’s valence e- feels effective nuclear charge of +1
previous | index | next
Calculate “effective nuclear charge”
• # protons minus # inner electrons
previous | index | next
as go across row size tends to decrease a bit
because of greater PPP “proton pulling power”
previous | index | next
• size  as you go  column
• size  as you go  row
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?
decreases
• As go across row?
increases
previous | index | next
•Valence electrons in atoms: effective nuclear charge of +1
•Cs valence electron lot farther away from nucleus than Li
•electrostatic attraction much weaker so easier to steal
electron away from Cs
previous | index | next
•easier to steal electron from Li than from Ne
•Li: smaller effective nuclear charge - valence electron
farther away from nucleus
•Li has less “proton pulling power” than Ne
Trends in ionization energy
• Ionization energy decreases as go
down a column
– gets 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
don’t have electronegativity values
• 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 most electronegative element!
Reactivity of Metals
• Metals are losers!
• judge reactivity of metals by how easily
give up electrons
• most active metals: Fr (then Cs)
• For metals, reactivity increases as
ionization energy goes down
Trends for Reactivity of Metals
or
Metallic Character
• 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)
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 (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
Allotropes
• Different forms of element in same phase
– different structures and properties
O2 and O3 - both gas phase
O2 (oxygen) - necessary for life
O3 (ozone) - toxic to life
Graphite, diamond:
both carbon in solid form
Graphite and Diamond