ELECTRON CONFIGURATIONS AND PERIODICITY

1
CHAPTER 8
ELECTRON CONFIGURATIONS AND PERIODICITY
2
I. Why is this is very important?
A) Chemistry is explained through sharing or
transferring electrons.
B) All that follows this semester and next will
use the vocabulary of orbitals, occupancy, etc.
II. How are the electrons arranged in an atom?
What orbitals are preferentially filled by
electrons?
A) What is an electron configuration?
3
B) THREE RULES GOVERN THE ARRANGEMENT OF
ELECTRONS IN ATOMS.
1) The Pauli Exclusion Principle - no two
electrons in the same atom can have the same set
of all 4 quantum numbers (no 2 can have the same
zip code).
2) The Aufbau Principle - (means building-up in
German) in the ground state, the electrons will
fill the atomic orbital of lowest energy.
4
Lowest energy is determined by the n ? rule.
That is, the sum of those two quantum numbers
determines the lowest energy, the lower the sum
the lower the energy. If the sums are equal, the
lowest value of n determines the lowest energy.
n ?
1 0 1 1s
2 0 2 2s
2 1 3 2p
3 0 3 3s
5
there is a tie, the 2p is lower than the 3s
because n2 is less than n3
3 1 4 3p
3 2 5 3d
4 0 4 4s
4 1 5 4p
there are 2 ties, the 4s is higher than the 3p
because n4 is GREATER than n3 and the 4s (sum
4) comes before the 3d (sum 5)
The second tie is 5 3d comes before 4p BECAUSE
______.
6
The order is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s,
4d etc.
7
3) Hund's Rule - Electrons occupy all the
orbitals of a given sublevel singly before
pairing begins. Spins of electrons in different
incomplete orbitals are parallel in the ground
state.The most stable arrangement of electrons in
the subshells is the one with the greatest number
of parallel spins.
8
II. ELECTRON CONFIGURATIONS - THE GROUND STATE
Element Electron configuration
1H 1s1
2He 1s2
3Li 1s22s1
4Be 1s22s2
5B 1s22s22p1
6C 1s22s22p2
7N 1s22s22p3
9
8O 1s22s22p4
9F 1s22s22p5
10Ne 1s22s22p6
11Na 1s22s22p63s1
12Mg 1s22s22p63s2
13Al 1s22s22p63s23p1
14Si 1s22s22p63s23p2
15P 1s22s22p63s23p3
16S 1s22s22p63s23p4
17Cl 1s22s22p63s23p5
18Ar 1s22s22p63s23p6
10
19K 1s22s22p63s23p64s1
20Ca 1s22s22p63s23p64s2
21Sc 1s22s22p63s23p64s23d1
22Ti 1s22s22p63s23p64s23d2
23V 1s22s22p63s23p64s23d3
24Cr 1s22s22p63s23p64s13d5
There is a tendency toward half-filled and
completely filled d subshells. This is a
consequence of the closeness of the 3d and the 4s
orbital energies.
11
The 3d level becomes more stable as we move from
left to right on the periodic chart. Remember
there is an increase in the number of
__________consequently, an increase in the number
of electrons as we move from left to right on the
chart.
25Mn 1s22s22p63s23p64s23d5
26Fe 1s22s22p63s23p64s23d6
27Co 1s22s22p63s23p64s23d7
28Ni 1s22s22p63s23p64s23d8
29Cu 1s22s22p63s23p64s13d10
12
Additional exceptions are Mo 5s14d5 Ag
5s14d10 Au 6s15d10 That is reasonable
considering their position on the periodic chart.
30Zn 1s22s22p63s23p64s23d10
31Ga 1s22s22p63s23p64s23d104p1
32Ge 1s2 2s22p63s23p64s23d104p2
33As 1s2 2s22p63s23p64s23d104p3
34Se 1s22s22p63s23p64s23d104p4
35Br 1s22s22p63s23p64s23d104p5
36Kr 1s22s22p63s23p64s23d104p6
13
You should be able to do the electron
configuration for any element on the chart which
follows the basic rules and the d5 and d10
exceptions.
Orbital Diagrams
14
1s 2s 2p
15
(No Transcript)
16
(No Transcript)
17
ELECTRON CONFIGURATIONS OF IONS
Electrons do not come out the same way as we put
them in according to the Aufbau Principle.
Electrons leave the outer most shell first.
Let's look at V vs V2 23V 1s22s22p63s23p64s23d
3
23V2 1s22s22p63s23p63d3
18
III. Relationship to the Periodic Chart I
want you to be able to use the Periodic Chart to
do electron configurations.
19
At the top left of the periodic table which you
have been given, write the letter n. At the
beginning of each row write a number beginning
with 1 and ending with 7. Above H put s1 above
He put s2
above B put np1 above C put np2 etc across the
chart.
2 lines above the transition metals put (n- 1)d
20
Above Sc put d1 above Ti put d2above V put d3
above Cr put d5 above Mn put d5 above Fe put
d6 above Co put d7 above Ni put d8 above Cu
put d10 above Zn put d10.
Outside the lanthanides and actinides put (n - 2)f
IV. Magnetic Properties of Atoms
A) Unpaired electrons give a paramagnetic
substance.
21
A paramagnetic substance is one that is attracted
into a magnetic field. This is generally the
result of unpaired electrons.
The experimental evidence is that in a magnetic
field a paramagnetic element weighs more than in
a nonmagnetic field environment. The change in
weight between the two environments can be used
to calculate the number of unpaired electrons in
a particular atom (element).
22
An element in which all the electrons are paired
is a diamagnetic substance. It is repelled in a
magnetic field and weighs __________ between the
two environments.
V. TRENDS IN THE PERIODIC TABLE - USE OF WHAT WE
HAVE JUST LEARNED
A)Atomic radius - a number of physical properties
of elements are related to the size of an atom,
but with our probability picture where does an
atom end?
23
Experimental Measurement - the covalent radius is
the center to center distance between 2 adjacent
atoms. The atomic radius is ½ the distance
between the 2 nuclei of the adjacent atoms. There
are many difficulties associated with obtaining
these numbers, but in spite of the difficulties,
scientists have assembled a set of approximate
atomic radii from interatomic distances.
24
(No Transcript)
25
C) Atomic radius, in general, __________ as we
move from left to right in a row of the periodic
table. (Many students get this wrong for some
reason. DON'T BE ONE OF THEM!!!)
D) Atomic radius increases from top to bottom in
a family or group.
E) These 2 trends are the result of 3 influences
on size.
26
1) The number of the principal energy levels
which are occupied, or the highest principal
energy level which is occupied. As n increases,
the size ________ , shells with a larger value of
n extend further from the nucleus and the
covalent radius_________. This accounts for
the_________ from top to bottom in a group.
2) The nuclear charge is very important. As the
number of protons increases across a row, the
positive charge on the nucleus increases.
27
THERE IS A STRONGER ATTRACTION FOR ELECTRONS IN
THE SAME SHELL. The size _____________especially
considering the s and p block elements.
3) The shielding effect is another important
consideration in predicting the size of atoms.
The attraction for electrons in the outermost
shell by the nucleus is shielded by electrons in
lower energy levels.
a) The smaller size of atoms going across a row
can be attributed to minimum shielding.
28
Electrons in the same shell are attracted more
strongly as the nuclear charge increases, because
the shielding effect remains the same. If the
shielding effect remains the same, the Effective
Nuclear Charge increases. The ENC is the positive
charge that an electron experiences from the
nucleus and is equal to the nuclear charge minus
the number of shielding electrons.
29
In the second row for example Li has 3protons
in the nucleus, 2e in the 1s orbital (shielding)
and 1e in the 2s orbital. ENC 3 - 2 1. The
outermost electron "feels" a net attraction by
the inside of 1.
B has 5 in the nucleus, 2e in the 1s orbital
(shielding) and 2e in the 2s orbital, and 1e in
the 2p orbital. ENC _______ The outermost
electron "feels" a net attraction by the inside
of _____ .
30
Carbon will have an ENC of 4, N one of 5 etc.
As the ENC _________the size ________.
b) Larger size - valence shell - the outermost
shell is further from the nucleus. Shielding and
n are increasing, but the outermost electrons
"don't know" that there is an increase in the
number of positive charges in the nucleus. The
ENC remains the same down a column and the
valence electrons are further out, allowing the
size to_____________ down a column.
31
c) What should you be able to do as a result of
this??? I should be able to give you a list of
elements and you should be able to put them in
order of size from smallest to largest by just
looking at their positions on the chart.
EXAMPLE Arrange the following in the order of
increasing size Na Be Mg .
THINK... Be is the highest and farthest to the
right on the chart, Na is to the left of Mg and
in the same row. Therefore the order should be Be
lt Mglt Na (largest).
32
F) IONIZATION ENERGY - when not modified as
first, second, third, etc., it usually means the
first ionization energy. THIS IS ANOTHER
IMPORTANT TREND IN THE PERIODIC TABLE.
1) The first ionization energy is the minimum
amount of energy required to remove the most
loosely held electron from an isolated gaseous
atom in the ground state to form a 1 charged ion.
33
A(g) energy ? A 1e
A(g) ? A 1e ?H kJ/mol
2) THIS IS A VERY IMPORTANT CONCEPT because the
chemical properties of any atom are determined by
the configuration of an atom's valence electrons,
those electrons in the outermost shell. Those are
the electrons of importance in bonding, the
subject of the next two chapters.
34
3) The ionization energies of electrons in atoms
are a confirmation of the electron configurations.
4) The way in which ionization energies vary with
atomic number provides another good example of
the periodic law and electron configurations.
5) In general the first ionization energy
increases from left to right across a row in the
Periodic Chart with some notable exceptions.
35
(No Transcript)
36
(No Transcript)
37
6) The trend across from left to right is
accounted for by a) the increasing nuclear charge.
38
The electrons in the outermost shell are more
strongly bound to the nucleus due to the
increasing effective nuclear charge.
b) the decrease in atomic radius. This also tends
to reinforce the concept of the electrons being
more strongly bound to the nucleus. Notice the
high FIE of the Noble Gases. Eight electrons in
the valence shell is an extremely stable electron
arrangement.
c) Irregularities with which you should be
familiar
39
(No Transcript)
40
Boron's FIE is less than Be's
B 1s22s22p1 Be 1s22s2
The 2 p electron extends a little further from
the nucleus than the 2s, therefore it is easier
to remove. It is shielded by the 2s electron
less tightly held.
Oxygen's FIE is less than N's and sulfur's is
less than phosphorus. This is due to oxygen
having a shared pair of electrons in one of the p
orbitals.
41
With an electron already in the orbital there is
repulsion between the two in the same orbital and
it comes out with less energy input.
7) The trend from top to bottom of a column shows
a decrease in the FIE which corresponds to an
increase in the atomic radius.
8) The 2nd, 3rd, and 4th ionization energies are
those required to remove the 2nd, 3rd, and 4th
electrons.
42
Each ion formed is more highly positively charged
and the energy required to remove each electron
is greater than that required for the previous
one. When an electron is removed from a lower
shell, there is a big jump in energy.
43
(No Transcript)
44
Element First
Second
Third
Fourth

Fifth
Sixth

Seventh H
1,312 He
2,372
5,250 Li
520
7,298
11,815 Be
899
1,757
14,848
21,006 B
801
2,427
3,660

9) Put in order of increasing first ionization
energy the following elements C N O F
G) ELECTRON AFFINITY
1) Electron affinity is the energy change which
occurs when an electron is accepted by an atom in
the gaseous state. A(g) e
?A-(g)
2) In contrast to ionization energy, what do we
observe on the following graph of EA's?
45
(No Transcript)
46
a) Ionization Energy is always endothermic. It
always takes energy to remove an electron. ?H is
always
b) Electron Affinity can be either endothermic or
exothermic depending on the element.
c) An exothermic (?H -) value for the electron
affinity indicates that energy is released upon
the addition of an electron to a gaseous atom.
47
The greater the negative value of the electron
affinity, the greater the tendency of an atom to
accept an electron.
d) A ?H indicates that energy must be absorbed
for an atom to gain an electron.
e) As we go from left to right on the periodic
chart, the elements have, in general, an
increasing tendency to form negative ions.
However, there are more exceptions than with
Ionization Energy.
48
(No Transcript)
49
The End