Download periodic trends

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PERIODIC
TRENDS
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Electron
Filling
Order
Figure 8.5
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Electron Configurations
and the Periodic Table
Figure 8.7
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Phosphorus
Group 5A
Atomic number = 15
1s2 2s2 2p6 3s2 3p3
[Ne] 3s2 3p3
All Group 5A elements
have [core ] ns2 np3
configurations where n
is the period number.
3p
3s
2p
2s
1s
Lithium
3p
3s
2p
2s
1s
Group 1A
Atomic number = 3
1s22s1 ---> 3 total
electrons
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Electron Properties
Diamagnetic: NOT attracted to a
magnetic field
Paramagnetic: substance is attracted
to a magnetic field. Substance has
unpaired electrons.
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7
Neon
3p
3s
2p
2s
1s
Group 8A
Atomic number = 10
1s2 2s2 2p6 --->
Diamagnetic
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Beryllium
3p
3s
2p
2s
1s
Group 2A
Atomic number = 4
1s22s2
Diamagnetic
Boron
3p
3s
2p
2s
1s
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Group 3A
Atomic number = 5
1s2 2s2 2p1
Paramagnetic
Carbon
3p
Group 4A
Atomic number = 6
1s2 2s2 2p2
3s
2p
2s
1s
Paramagnetic
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11
Fluorine
3p
3s
2p
2s
1s
Group 7A
Atomic number = 9
1s2 2s2 2p5 --->
Paramagnetic
Ion Configurations
How do we know the configurations
of ions?
Determine the magnetic properties
of ions.
Ions with UNPAIRED
ELECTRONS are
PARAMAGNETIC.
Without unpaired electrons
DIAMAGNETIC.
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transition metal ions
Fe
[Ar] 4s2 3d6
loses 2 electrons ---> Fe2+ [Ar] 4s0 3d6
Fe2+
Fe
4s
3d
4s
3d
• loses 3 electrons ---> Fe3+ [Ar] 4s0 3d5
Fe2+
Fe
4s
3d
4s
3d
Transition Metals
How do they fill? How can
we determine?
Chromium
Iron
Copper
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Ion Configurations Mn
Mn
[Ar] 4s2 3d5
---> Mn5+ [Ar] 4s03d2
loses 5 electrons ---> Mn5+ [Ar] 4s2 3d0
2+
P
Fe
D
Fe
4s
3d
4s
3d
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PERIODIC
TRENDS
PERIODICITY
Period Law-physical and chemical
properties of elements are a
periodic function of their
atomic numbers
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General Periodic Trends
• Atomic and ionic size
• Ionization energy
• Electron affinity, electronegativity
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Effective Nuclear Charge
Z*
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Effective Nuclear Charge, Z*
• Z* is the nuclear charge experienced by
the outermost electrons. See p. 295 and Screen 8.6.
• Explains why E(2s) < E(2p)
• Z* increases across a period owing to
incomplete shielding by inner electrons.
• Estimate Z* by --> [ Z - (no. inner electrons) ]
• Charge felt by 2s e- in Li
Z* = 3 - 2 = 1
• Be
Z* = 4 - 2 = 2
• B
Z* = 5 - 2 = 3
and so on!
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Effective Nuclear Charge, Z*
• Atom
•
•
•
•
•
•
•
Li
Be
B
C
N
O
F
Z* Experienced by Electrons in
Valence Orbitals
+1.28
------+2.58
Increase in
+3.22
Z* across a
+3.85
period
+4.49
+5.13
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General Periodic Trends
Higher effective nuclear charge
Electrons held more tightly
Larger orbitals.
Electrons held less
tightly.
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Lithium
Periodic Trend in the
Reactivity of Metals
Sodium
Potassium
MOST
2.
 As
Reactivity for Metals
you go down a group for metals
the number of energy levels
increase.
 Because of this, reactivity increases
because the atom is more willing to
give away its electron (react).
3.Nonmetalic Trends: Gain electrons

Nonmetals on right side, form anions

Going right elements are more
nonmetallic (better gainers of electrons)

Going UP elements become more
nonmetallic (want to gain)
8. Reactivity nonmetals: Gain e

The reason Across = fill the energy level

Going UP a group, nonmetals have
same valence but fewer total electrons

Flourine is the most reactive nonmetal.
Atomic Radii
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Figure 8.9
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Atomic Size
• Size increases, down a group.
• Because electrons are added into
additional energy levels, there is less
attraction.
• Size decreases across a period.
• Because, increased effective nuclear
charge.
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Atomic Size
Size decreases across a period owing
to increase in Z*. Each added electron
feels a greater and greater + charge.
Large
Small
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Trends in Atomic Size
See Figures 8.9 & 8.10
Radius (pm)
250
K
1st transition
series
3rd period
200
Na
2nd period
Li
150
Kr
100
Ar
Ne
50
He
0
0
5
10
15
20
25
Atomic Number
30
35
40
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Ion Sizes
+
Li,152 pm
3e and 3p
Li + , 78 pm
2e and 3 p
Forming
a cation.
• CATIONS are SMALLER than the
atoms from which they come.
• The electron/proton attraction
has gone UP and so size
DECREASES.
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Ion Sizes
F, 71 pm
9e and 9p
F- , 133 pm
10 e and 9 p
Forming
an anion.
• ANIONS are LARGER than the atoms
from which they come.
• The electron/proton attraction has
gone DOWN and so size INCREASES.
• Trends in ion sizes are the same as
atom sizes.
Trends in Ion Sizes
Figure 8.13
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Ionization Energy
IE = energy required to remove an electron
from an atom in the gas phase.
Mg (g) + 738 kJ ---> Mg+ (g) + e-
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Ionization Energy
IE = energy required to remove an electron
from an atom in the gas phase.
Mg (g) + 738 kJ ---> Mg+ (g) + e-
Mg+ (g) + 1451 kJ ---> Mg2+ (g) +
eMg+ has 12 protons
and only 11
electrons. Therefore, IE for Mg+ > Mg.
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Ionization Energy
1st IE: Mg (g) + 735 kJ ---> Mg+ (g) + e2nd IE: Mg+ (g) + 1451 kJ ---> Mg2+ (g) + e-
3rd IE: Mg2+ (g) + 7733 kJ ---> Mg3+ (g) + eEnergy cost is very high to dip into a
shell of lower n (core electrons).
This is why ox. no. = Group no.
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Trends in Ionization Energy
1st Ionization energy (kJ/mol)
2500
He
Ne
2000
Ar
1500
Kr
1000
500
0
1
H
3
Li
5
7
9
11
Na
13
15
17
19
K
21
23
25
27
29
31
Atomic Number
33
35
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Trends in Ionization Energy
• IE decreases down a
group
• Because size increases.
• IE increases across a period
• Because effective nuclear
charge increases
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Electron Affinity
A few elements GAIN
electrons to form anions.
Electron affinity is the energy
involved when an atom
gains an electron to form an
anion.
X(g) + e- ---> X-(g)
E.A. = ∆E
Trends in Electron Affinity
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Trends in Electron Affinity
• Affinity for electron
increases across a
period (EA becomes
more negative).
• Affinity decreases
down a group (EA
becomes less
negative).
Atom EA
F
-328 kJ
Cl -349 kJ
Br -325 kJ
I
-295 kJ
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Electron Affinity of Oxygen
O atom [He] 
 

+ electron
O- ion [He] 
 
EA = - 141 kJ

∆E is
EXOthermic
because O
has an affinity
for an e-.
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Electron Affinity of Nitrogen
N atom [He] 
 

+ electron
N- ion
[He] 


EA = 0 kJ

∆E is zero for Ndue to
electronelectron
repulsions.
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Electronegativity
• So how is this different from
electron affinity?
• Electron Affinity – is rating of how
well an atom wants to gain an
electron
• Electronegativity – is rating of how
well an atom keeps the electron
once it is bonded to another atom
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Electronegativity
Electron Configurations
and the Periodic Trends
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• Periodic table