electron configurations follow the order of sublevels on ...€¢ for 3d transition metals, 3d has...
TRANSCRIPT
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1
Electron configurations follow the order of sublevels on
the periodic table.
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The periodic table consists of sublevel blocks
arranged in order of increasing energy.
• Groups 1A(1)-2A(2) = s level
• Groups 3A(13)-8A(18) = p level
• Groups 3B(3) to 2B(12) = d level
• Lanthanides/Actinides = f level
2
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To write an electron
configuration using
Sublevel blocks,
• locate the element on the periodic table
• starting with H in 1s,write
each sublevel block in order
going from left to right
across each period
• write the number of electrons
in each block
3
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Solution
Period 1 1s block 1s2
Period 2 2s → 2p blocks 2s2 2p6
Period 3 3s → 3p blocks 3s23p2 (Si)
Writing all the sublevel blocks in order gives
1s22s22p63s23p2
4
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A. 1s22p5
B. 1s22s22p6
C. 1s22s22p3
D. 1s22s22p4
5
1s22p5
1s22s2
2p6
1s22s2
2p3
1s22s2
2p4
25% 25%25%25%
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A. 1s22s22p63s23p64s2
B. 1s22s22p63s23p64s1
C. 1s22s22p63s23p63d2
D. 1s22s22p63s23p54s23d1
6
1s22s2
2p63s23p64s2
1s22s2
2p63s23p64s1
1s22s2
2p63s23p63d2
1s22s2
2p63s23p54s2
3d1
25% 25%25%25%
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• For 3d transition metals, 3d has lower energy after the
electrons is filled. 3d and 4s are close in energy
1s 2s 2p 3s 3p 4s 3d
Ar 1s2 2s2 2p6 3s2 3p6
K 1s2 2s2 2p6 3s2 3p6 4s1
Ca 1s2 2s2 2p6 3s2 3p6 4s2
Sc 1s2 2s2 2p6 3s2 3p6 3d14s2
Ti 1s2 2s2 2p6 3s2 3p6 3d24s2
7
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8
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A. 1s22s2 2p63s2 3p6 4s2
B. 1s22s2 2p63s2 3p6 4s2 3d5
C. 1s22s2 2p63s2 3p6 4s2 3d3
D. 1s22s2 2p63s2 3p6 4s2 3d2
10
1s22s2
2p63s2
3p6 4
s2
1s22s2
2p63s2
3p6 4
s2 3
d5
1s22s2
2p63s2
3p6 4
s2 3
d3
1s22s2
2p63s2
3p6 4
s2 3
d2
25% 25%25%25%
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Solution
Period 1 1s block 1s2
Period 2 2s → 2p blocks 2s2 2p6
Period 3 3s → 3p blocks 3s2 3p6
Period 4 4s → 3d blocks 4s2 3d2 (at Ti)
Writing all the sublevel blocks in order gives
1s2 2s2 2p63s2 3p6 4s2 3d2
11
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A. Na
B. S
C. K
D. Ca
12
Na S K Ca
0% 0%0%0%
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A. 3p64s2
B. 4s24d7
C. 4s23d7
13
3p64s2
4s24d7
4s23d7
100%
0%0%
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The last filled principal energy level is called the
valence level, or valence shell which is the
outermost sublevels that are highest in energy.
The valence electrons
• occupy orbitals in the valence level All the other
electrons are called core electrons, or inner electrons.
• determine the chemical properties of an element
• are related to the group number of the element
Example: Phosphorus has 5 valence electrons
5 valence electrons
P Group 5A(15) 1s22s22p6 3s23p3
14
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All the elements in a group have the same number
of valence electrons.
Example:
Be 1s2 2s2
Mg 1s2 2s2 2p6 3s2
Ca 1s2 2s2 2p6 3s2 3p6 4s2
Sr 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2
Can you see a pattern?
Elements in Group 2A (2) have two (2) valence electrons.
15
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7-16
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State the number of valence electrons for each of following:
A. oxygen
1) 4 2) 6 3) 8
B. aluminum
1) 13 2) 3 3) 1
C. chlorine
1) 2 2) 5 3) 7
Solution:
A. 2) 6 B. 2) 3 C. 3) 7
17
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State the number of valence electrons for each.
A. 1s22s22p63s23p3
B. 1s22s22p63s23p64s23d104p4
C. 1s22s22p5
Solution:
A. 5 B. 6 C. 7
18
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Noble gas-core abbreviated electron configurations are
often used for elements with many electrons.
Notice that iron’s electron configuration starts out with
argon’s electron configuration, but ends differently:
Fe 1s22s22p63s23p64s23d6
Ar 1s22s22p63s23p6
We use the symbol [Ar] to represent argon’s electron
configuration:
Fe [Ar] 4s23d6
argon core + valence electrons
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In atoms, the number of electrons is equal to the
number of protons, which is the atomic number.
In ions, the number of electrons does not equal the
atomic number. We must add or subtract electrons,
depending on whether the ion is an anion or cation.
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An example problem:
Write the electron configuration for Na+:
Na+ has a positive charge of 1; therefore,
we need to subtract 1 electron from the
total number of electrons, 11. Na+ has 10
electrons and is isoelectronic with Ne.
1s22s22p6
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7-22
Write the electron configuration in spdf notation and noble gas-core abbreviated electronic configuration for the following ions.
1. Br-
2. N3-
3. K+
4. Sr2+
5. S2-
6. Ni2+
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Write the electron configuration in spdf noation and
noble –gas core abbreviated for the following ions.
1. Br- 1s22s22p63s23p64s23d104p6 [Kr]
2. N3- 1s22s22p6 [Ne]
3. K+ 1s22s22p63s23p6 [Ar]
4. Sr2+ 1s22s22p63s23p64s23d104p6 [Kr]
5. S2- 1s22s22p63s23p6 [Ar]
6. Ni2+ 1s22s22p63s23p6 3d8 [Ar] 3d8
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Electron configurations are related to the
following properties:
Relative reactivity
Atomic radii (atomic size)
Ionization Energy (tendency to lose
electrons)
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From the activity series we see that the most active metals are those in Groups IA and IIA.
The more active the metal, the more easily it loses electrons to form ions.
Figure 7.24
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The alkali metals form +1 ions since they all have 1 valence electron.
Oxides of the alkali metals:
Li2O, Na2O, K2O, Rb2O, Cs2O
The alkaline earth metals form +2 ions since they all have 2 valence electrons.
Oxides of the alkaline earth metals:
BeO, MgO, CaO, SrO, BaO
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Ionization energy
• is the energy it takes to remove a electron from an atom
or ion in the gas state.
27
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the ionization energy
decreases in going
down a group.
28
In general, for Main group elements:
the ionization energy increase in going left to right across
a in period. Except, IE(group 3A) < IE(group 2A) and
IE(group 6A) < IE(group 5A)
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Select the element in each pair with the higher
ionization energy.
A. Li or K
B. K or Br
C. P or Cl
Solution
A. Li
B. Br
C. Cl
29
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What do these values tell us about the stability of
core (inner) electrons?
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31
Atomic size is usually reported atomic radius
covalent radius: half the distance between
identical nuclei of the neighboring atoms of the
same element in a molecules
Metallic radius: is based
on the distance of
closest approach of
adjacent atoms in a
solid metal.
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The atomic radius in
general increases
• going down each group
as the number of energy
levels increases
• Except : Al > Ga
32
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The atomic radius in general decreases
• going from left to right across a period
• because electrons are held closer to the nucleus by the
increasingly greater charge of the nucleus.
33
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Select the element in each pair with the larger
atomic radius.
A. Li or K
B. K or Br
C. P or Cl
34
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Select the element in each pair with the larger
atomic radius.
A. K is larger than Li
B. K is larger than Br
C. P is larger than Cl
35
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A positive ion
• has lost its valence
electrons
• is smaller than the
corresponding neutral atom
• Example: Group 1A and
metals ions about half the
size of its corresponding
neutral atoms
36
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The sodium ion Na+
• forms when the Na atom loses one electron from
the third energy level
• is smaller than a Na atom
37
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A negative ion
• increases the number of
valence electrons
• is larger than the
corresponding neutral atom
• Example: Group 7A and
nonmetal ions about twice the
size of its corresponding
neutral atoms
38
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The fluoride ion F-
• forms when a valence electron is added
• has increased repulsions due to the added valence
electron
• is larger than a F atom
39
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For any isoelectronic (same number of electrona) series,
as the number of protons increases, the ion size
decreases.
Figure from p. 286
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The general trend for ionic size (or radius) is for ionic size to increase from top to bottom in the periodic table.
For cations, as the charge increases, the ionic size decreases.
For anions, as the
charge increases
(becomes more
negative), the ionic
size increases.
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1. Which is larger in each of the following?
A. K or K+
B. Al or Al3+
C. S2- or S
2. Which is smaller in each of the following?
A. N3- or N
B. Cl or Cl-
C. Sr2+ or Sr
42
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1. Which is larger in each of the following?
A. K > K+
B. Al > Al3+
C. S2- > S
2. Which is smaller in each of the following?
A. N < N3-
B. Cl < Cl-
C. Sr2+ < Sr
43
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Ability of an atom to attract bonding electrons
Proposed by Linus Pauling in the early 1930’s
A difference in electronegativity between the atoms in a covalent bond results in:
A polar covalent bond
Increased ionic character
The greater the difference in electronegativity, the greater the ionic character and the more polar the bond that joins the atoms.
Decreased bond length and increased bond strength
No difference in electronegativity between atoms in a covalent bond results in a nonpolar covalent bond.
8-44
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Electronegativity Values
Figure 8.5
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The difference in electronegativity between metals
and nonmetals is so large, that the electrons are
transferred, not shared.
The greater the electronegativity difference, the
more polar the bond.
Si-F > N-F> O-F >F-F
What partial charges go
on each atom?
8-46
Figure 8.6