Download effective nuclear charge

Trend in Atomic Radius – Main Group
 Different methods for measuring the radius of an atom,
and they give slightly different trends
 van der Waals radius = nonbonding
 covalent radius = bonding radius
 atomic radius is an average radius of an atom based on
measuring large numbers of elements and compounds
 Atomic Radius Increases down group
 valence shell farther from nucleus
 effective nuclear charge fairly close
 Atomic Radius Decreases across period (left to right)
 adding electrons to same valence shell
 effective nuclear charge increases
 valence shell held closer
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Effective Nuclear Charge
 in a multi-electron system, electrons are simultaneously
attracted to the nucleus and repelled by each other
 outer electrons are shielded from full strength of nucleus
 screening effect
 effective nuclear charge is net positive charge that is
attracting a particular electron
 Z is nuclear charge, S is electrons in lower energy levels
 electrons in same energy level contribute to screening, but very
little
 effective nuclear charge on sublevels trend, s > p > d > f
Zeffective = Z - S
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Screening & Effective Nuclear Charge
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Trends in Atomic Radius
Transition Metals
 increase in size down the Group
 atomic radii of transition metals roughly the same size
across the d block
 must less difference than across main group elements
 valence shell ns2, not the d electrons
 effective nuclear charge on the ns2 electrons approximately the
same
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Example 8.5 – Choose the
Larger Atom in Each Pair
1)
2)
3)
4)
N or F
F,
C or Ge
Ge,
N or Al,
Al
Al or Ge?
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Electron Configuration of
Cations in their Ground State
 cations form when the atom loses electrons from the valence
shell
 for transition metals electrons, may be removed from the
sublevel closest to the valence shell
Al atom = 1s22s22p63s23p1
Al+3 ion = 1s22s22p6
Fe atom = 1s22s22p63s23p64s23d6
Fe+2 ion = 1s22s22p63s23p63d6
Fe+3 ion = 1s22s22p63s23p63d5
Cu atom = 1s22s22p63s23p64s13d10
Cu+1 ion = 1s22s22p63s23p63d10
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Magnetic Properties of
Transition Metal Atoms & Ions
 both Zn atoms and Zn2+ ions are diamagnetic, showing
that the two 4s electrons are lost before the 3d
 Zn atoms [Ar]4s23d10
 Zn2+ ions [Ar]4s03d10
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 Write the Electron Configuration and Determine
whether the following atoms their ions are Paramagnetic
or Diamagnetic
 Fe and Fe3+
 Cu and Cu1+
 Ag and Ag+
Trends in Ionic Radius
 Ions in same group have same charge
 Ion size increases down the group
 higher valence shell, larger
 Cations smaller than neutral atom; Anions bigger
than neutral atom
 Cations smaller than anions
 except Rb+1 & Cs+1 bigger or same size as F-1 and O-2
 Larger positive charge = smaller cation
 for isoelectronic species
 isoelectronic = same electron configuration
 Larger negative charge = larger anion
 for isoelectronic series
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1A
Periodic
Pattern
Ionic
Radius
(Å)
+1
-1
2A
H
Li
+1
+2
0.68 Be
0.31
+1
Na
0.97 Mg 0.66
+2
1.33 Ca 0.99
+1
Rb
1.47
+1
Cs
+2
Sr
1.13
+2
1.69 Ba 1.35
4A
5A
+3
B 0.23
+2
+1
K
3A
-4
C
+3
Al
0.51
0.62
+3
+1
Ga
0.81
In
0.95
Tl
6A
N
-4
Si
P
-4
Ge
0.71
+3
+1
0.84 +4
+2
Pb
Sn
+4
+2
-3
-2
1.71 O
1.40
-3
-2
-1
1.84 Cl
1.81
-2
-1
2.12
S
-3
As
+3
+1
7A
-1
F
2.22 Se 1.98 Br 1.96
-2
Sb
Bi
1.33
Te
2.21
-1
I
2.20
14
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Ionization Energy
 minimum energy needed to remove an electron from an atom
 gas state
 endothermic process
 valence electron easiest to remove
 M(g) + IE1  M1+(g) + 1 e M+1(g) + IE2  M2+(g) + 1 e
first ionization energy = energy to remove electron from neutral atom;
2nd IE = energy to remove from +1 ion; etc.
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General Trends in 1st Ionization Energy
 larger the effective nuclear charge on the electron, the more
energy it takes to remove it
 the farther the most probable distance the electron is from the
nucleus, the less energy it takes to remove it
 1st IE decreases down the group
 valence electron farther from nucleus
 1st IE generally increases across the period
 effective nuclear charge increases
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Example 8.8 – Choose the Atom in Each Pair
with the Higher First Ionization Energy
1) Al or SS,
2) As or Sb
Sb,
3) N or Si,
Si
4) O or Cl?
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Irregularities in the Trend
 Ionization Energy generally increases from left to right across
a Period
 except from 2A to 3A, 5A to 6A
Be
B


1s
2s


1s
2s
N
2p

O
2p


1s
2s


1s
2s



2p
 

2p
Which
Which isis easier
easier to
to remove
remove an
an electron
electron from
from B
Nor
orBe?
O? Why?
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Irregularities in the
First Ionization Energy Trends
Be


1s
2s
Be+
2p

1s

2s
2p
To ionize Be you must break up a full sublevel, cost extra energy
B


1s
2s

B+
2p


1s
2s
2p
When you ionize B you get a full sublevel, costs less energy
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Irregularities in the
First Ionization Energy Trends
N


1s
2s



N+
2p


1s
2s


2p
To ionize N you must break up a half-full sublevel, cost extra energy
O


1s
2s
 

2p
O+


1s
2s



2p
When you ionize O you get a half-full sublevel, costs less energy
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Trends in Successive
Ionization Energies
 removal of each successive electron
costs more energy
 shrinkage in size due to having more
protons than electrons
 outer electrons closer to the nucleus,
therefore harder to remove
 regular increase in energy for each
successive valence electron
 large increase in energy when start
removing core electrons
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Trends
in
Electron
Affinity
 energy released when an neutral atom gains an electron
 gas state
 M(g) + 1e-  M-1(g) + EA
 defined as exothermic (-), but may actually be
endothermic (+)
 alkali earth metals & noble gases endothermic, WHY?
 more energy released (more -); the larger the EA
 generally increases across period
 becomes more negative from left to right
 not absolute
 lowest EA in period = alkali earth metal or noble gas
 highest EA in period = halogen
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Metallic
Character
 Metals






malleable & ductile
shiny, lusterous, reflect light
conduct heat and electricity
most oxides basic and ionic
form cations in solution
lose electrons in reactions - oxidized






brittle in solid state
dull
electrical and thermal insulators
most oxides are acidic and molecular
form anions and polyatomic anions
gain electrons in reactions - reduced
 Nonmetals
 metallic character increases left
 metallic character increase down
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Example 8.9 – Choose the
More Metallic Element in Each Pair
1)
2)
3)
4)
Sn or Te,
Te
P or Sb,
Sb
Ge or In,
In
S or Br?
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Trends
in
the
Alkali
Metals
 atomic radius increases down the column
 ionization energy decreases down the column
 very low ionization energies




good reducing agents, easy to oxidize
very reactive, not found uncombined in nature
react with nonmetals to form salts
compounds generally soluble in water  found in seawater
 electron affinity decreases down the column
 melting point decreases down the column
 all very low MP for metals
 density increases down the column
 except K
 in general, the increase in mass is greater than the increase in
volume
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2 Na(s) + 2 H2O(l)  2 NaOH(aq) + H2(g)
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Trends
in
the
Halogens
 atomic radius increases down the column
 ionization energy decreases down the column
 very high electron affinities




good oxidizing agents, easy to reduce
very reactive, not found uncombined in nature
react with metals to form salts
compounds generally soluble in water  found in seawater
 reactivity increases down the column
 react with hydrogen to form HX, acids
 melting point and boiling point increases down the
column
 density increases down the column
 in general, the increase in mass is greater than the increase in
volume
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Example 8.10 – Write a balanced chemical
reaction for the following.
 reaction between potassium metal and bromine gas
K(s) + Br2(g) 
(ionic compounds are all solids at room temperature)
 reaction between rubidium metal and liquid water
Rb(s) + H2O(l) 
 reaction between chlorine gas and solid iodine
Cl2(g) + I2(s) 
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Trends
in
the
Noble
Gases
 atomic radius increases down the column
 ionization energy decreases down the column
 very high IE
 very unreactive
 only found uncombined in nature
 used as “inert” atmosphere when reactions with other gases
would be undersirable
 melting point and boiling point increases down the
column
 all gases at room temperature
 very low boiling points
 density increases down the column
 in general, the increase in mass is greater than the increase in
volume
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