Transcript Effective Nuclear Charge (Z eff )

Periodic Relationships Among
the Elements
Chapter 8
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ns2np6
ns2np5
ns2np4
ns2np3
ns2np2
ns2np1
d10
d5
d1
ns2
ns1
Ground State Electron Configurations of the Elements
4f
5f
8.2
Electron Configurations of Cations and Anions
Of Representative Elements
Na [Ne]3s1
Na+ [Ne]
Ca [Ar]4s2
Ca2+ [Ar]
Al [Ne]3s23p1
Al3+ [Ne]
Atoms gain electrons
so that anion has a
noble-gas outer
electron configuration.
Atoms lose electrons so that
cation has a noble-gas outer
electron configuration.
H 1s1
H- 1s2 or [He]
F 1s22s22p5
F- 1s22s22p6 or [Ne]
O 1s22s22p4
O2- 1s22s22p6 or [Ne]
N 1s22s22p3
N3- 1s22s22p6 or [Ne]
8.2
-1
-2
-3
+3
+2
+1
Cations and Anions Of Representative Elements
8.2
Na+: [Ne]
Al3+: [Ne]
O2-: 1s22s22p6 or [Ne]
F-: 1s22s22p6 or [Ne]
N3-: 1s22s22p6 or [Ne]
Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne
What neutral atom is isoelectronic with H- ?
H-: 1s2
same electron configuration as He
8.2
Electron Configurations of Cations of Transition Metals
When a cation is formed from an atom of a transition metal,
electrons are always removed first from the ns orbital and
then from the (n – 1)d orbitals.
Fe:
[Ar]4s23d6
Fe2+: [Ar]4s03d6 or [Ar]3d6
Mn:
[Ar]4s23d5
Mn2+: [Ar]4s03d5 or [Ar]3d5
Fe3+: [Ar]4s03d5 or [Ar]3d5
8.2
Effective nuclear charge (Zeff) is the “positive charge” felt
by an electron.
Zeff = Z - s
0 < s < Z (s = shielding constant)
Zeff  Z – number of inner or core electrons
Z
Core
Zeff
Radius (pm)
Na
11
10
1
186
Mg
12
10
2
160
Al
13
10
3
143
Si
14
10
4
132
8.3
Effective Nuclear Charge (Zeff)
increasing Zeff
increasing Zeff
8.3
8.3
8.3
Atomic Radii
8.3
Comparison of Atomic Radii with Ionic Radii
8.3
Cation is always smaller than atom from
which it is formed.
Anion is always larger than atom from
which it is formed.
8.3
The Radii (in pm) of Ions of Familiar Elements
8.3
Ionization energy is the minimum energy (kJ/mol) required
to remove an electron from a gaseous atom in its ground
state.
I1 + X (g)
X+(g) + e-
I1 first ionization energy
I2 + X+(g)
X2+(g) + e-
I2 second ionization energy
I3 + X2+(g)
X3+(g) + e-
I3 third ionization energy
I1 < I2 < I3
8.4
8.4
Variation of the First Ionization Energy with Atomic Number
Filled n=1 shell
Filled n=2 shell
Filled n=3 shell
Filled n=4 shell
Filled n=5 shell
8.4
General Trend in First Ionization Energies
Increasing First Ionization Energy
Increasing First Ionization Energy
8.4
Ionization Energy Trends
DOWN a Group
Ionization Energy decreases as you move down a
group because the electrons that are being lost are
in successively higher energy levels – further from
the nucleus – requiring less energy to remove
them.
(Note – although the nuclear charge is greater
(more protons), there is more shielding because of
the additional electrons… so the overall effective
nuclear charge doesn’t really change)
Ionization Energy Trends
DOWN a Group
Which has a higher ionization energy, Na or Rb? Explain why.
Which has a higher ionization energy, Br or Cl? Explain why.
Ionization Energy Trends
ACROSS a Period
Ionization Energy increases as you move across a period because there
is an increase in the charge on the nucleus, which increases the
effective nuclear charge – requiring more energy to remove electrons
from the same energy level.
Note: The increase is not totally smooth. Between group 2 and 13
there is a ‘dip’ (decrease) in ionization energy because the electron
removed from the element in group 13 is in a p-subshell which is
slightly higher in energy than the s-subshells of group 2 and 1. This
electron is slightly higher in energy already, so it takes less energy to
remove it than it does to remove an electron from the s-subshell.
A similar glitch occurs between groups 15 and 16. In group 15 the
element has only 1 electron in each p-subshell. In group 16, one of the
p-subshells has a pair of electrons. The resulting greater electronelectron repulsion counteracts the effect of the increased nuclear
charge... which lowers the activation energy.
Ionization Energy Trends
ACROSS a Period
Which has a lower ionization energy, Na or Cl? Explain why.
Which has a higher ionization energy, B or Be? Explain why.
Which has a higher ionization energy, O or N? Explain why.
Electronegativity Trends – Why?
Down a Group – EN decreases
• Valence electrons further from nucleus
• Effective nuclear charge isn’t changing
• Atom’s ability to attract an electron further away from the
nucleus is diminishing
Across a Period – EN increases
• Charge on nucleus increases (more protons)
• Electrons in valence shell are more attracted to nucleus –
closer
• Atom’s ability to attract an electron is increased due to
closer distance to nucleus and increase effective nuclear
charge.
Electronegativity Trends – Why?
Which has a higher electronegativity, Na or Rb? Explain why.
Which has a higher electronegativity, Br or Cl? Explain why.
Melting Point Trends – Why?
Down a Metal Group – MPt decreases slightly
• Valence electrons further from nucleus
• Effective nuclear charge isn’t changing
• Atom’s ability to attract an electron further away from the nucleus is
diminishing
Across a Period – MPt increases… then decreases (MPt reflects strength of forces
between particles in solid vs liquid states.
• Gps 1&2 - Strength of metallic bond increases with inc nuclear charge, inc #
mobile valence electrons, atomic radius decreases
• Gps 14 – giant covalent structures… v. strong covalent bonding in all
directions
• Gps 15-16-17 – molecular structures… nonpolar… only weak van der Waals
forces… mpt depends on mass and size of molecules
• Gp 18 – single atoms, v wk vdWaals… very low mpt
Melting Point Trends – Why?
Which has a higher melting point, Na or Rb? Explain why.
Which has a higher melting point, Na or Cl2? Explain why.
Which has a higher melting point, C or O2? Explain why.
Periodic Trends
Ionization Energy, Electronegativity,
Atomic Radius, Ionic Radius, Melting Point