Sec.
5.1
Ch. 5: Electrons in Atoms
I.
Issue: Rutherford’s model could not explain
how/why atoms chemically react
II.
Answer: Niels Bohr’s Solar System Model (1913)
A.
Bohr’s (and others) observations
1. flame test results: metallic elements
burn a unique color
http://www.answers.com/topic/flame-test
2.
spectral lines: gases in glass tubes give
off characteristic, unique combination of
wavelengths of light
a.
set up
http://www.iun.edu/~cpanhd/C101webnotes/modern-atomic-theory/emission-spectrum.html
b.
unique spectral lines results
http://webmineral.com/help/FlameTest.shtml
B.
Bohr’s interpretation: planetary, or solar
system model of atom
1. electrons are arranged in discrete
energy levels around the positive
nucleus
2.
each energy level represents a different
amount of E; E of electrons increase as
distance from nucleus increases
3.
light E has a dual nature: can behave like
both energy and a particle; travels in waves
but absorbed/given off in discrete E
amounts called photons
4.
electrons absorb E and give off photons of E as
they leap from one energy level to the next
a.
ground state
distance in which the electron normally
orbits around the nucleus; represents its
lowest E state
b.
excited state
when heated, electrons endothermically
absorb E to quantum leap to a higher E
level but electrons are unstable at this
higher “excited” energy level
c.
electrons then leap back down to their
ground state and exothermically release
E in the form of visible photons of light E
III. The FINAL (?) Piece of the Atom - 1930
James Chadwick: bombarding Be with radiation
A.
Produced, among other things, a neutral
beam of particles
B.
Called these particles neutrons
1. the mass of a neutron was calculated to be
similar to that of a proton
(defined now as 1 “atomic mass unit” or
1amu)
2. has no charge – neutral particle
IV. Quantum Theory – the Wave Mechanical Model
of the Atom (current thinking)
A.
Erwin Schrodinger (1926) developed a math
equation to help solve the issues that
occurred with the Bohr model of the atom
1.
A.
Schrodinger is credited with discovering
this modern model of the atom
Why needed?
Bohr’s model couldn’t explain the
increasingly complicated spectral line results;
was only good for hydrogen’s
B.
Electrons now depicted as traveling in a lessdefined “cloud” or pathway; unsure as to
where exactly an electron is around the atom
C.
Electron arrangement based on a lot of
math! Don’t need to know the math, just the
end conclusions
D.
Electron location described in four
increasingly more specific areas. They’re
like an address for an electron.
1.
n = principal energy level
number of electrons each n holds: 2n2
n
1
2
3
4
5
6
7
2.
# electrons it can hold
2
8
18
ℓ = sublevels (s, p, d, or f)
a.
General rule to determine the
possible subshells within n energy
level (integers only)
i. ℓ ≤ n - 1
a. When ℓ = 0 (s-sublevel)
ℓ = 1 (p-sublevel)
ℓ = 2 (d-sublevel)
ℓ = 3 (f-sublevel)
b.
n = 1 has one energy level; no
subdivisions, we say that its ℓ = 0
and we know we are talking about
the s-sublevel
c.
sublevels represent multiple
pathways within the energy level
d.
are thought to exist to keep electron
repulsion at a minimum (keep
electrons apart as much as it can in
that tiny space)
e.
have uniquely-shaped pathways
http://en.wikibooks.org/wiki/General_Chemistry/Shells_and_Orbitals
ℓ
s
p
d
f
E level
lowest
higher
next highest
higher still
Shape of ℓ pathway
spherical
dumbbell
dumbbell with o-ring
dumbbell with 2 o-rings
d-orbital
f-orbital
Ex.
When n = 1, what sublevel(s) is/are possible?
Only the s-sublevel because ℓ ≤ n – 1
When n = 3, what sublevel(s) is/are possible?
ℓ = 2 which indicates that the highest sublevel is
the d-sublevel. If the highest sublevel is the d-sublevel
then every sublevel below the d-sublevel is possible.
This includes both an s and p sublevel.
When n = 2, what sublevel(s) is/are possible?
ℓ = 1 which indicates that the highest sublevel is
the p-sublevel. If the highest sublevel is the p-sublevel
then every sublevel below the p-sublevel is possible.
This includes only an s-sublevel.
3.
m = orbitals
a. sublevels after the s sublevel are
further divided into multiple
pathways; regardless, we say s
sublevel is also an orbital
b.
divide orbital into different
planes/axes; each orbital of a
sublevel has same size, distance
from nucleus, and E level
c.
number of orbitals within a sublevel
is always an odd number; starting
with s it is one, and then it goes up
consecutively by odd numbers
d.
orbitals range from - ℓ to + ℓ
i. s = ___
0
a. s-orbitals can only hold a
maximum of 2 electrons
ii. p = ___ ___ ___
-1
0 +1
a. p-subshells can hold a
maximum of 6 electrons
because each orbital can
hold two electrons and there
are three orbitals in a psubshell. Therefore a psubshell can hold a total of 6
electrons.
iii. d = ___ ___ ___ ___ ___
-2 -1
0 +1 +2
a. d-subshells can hold a
maximum of 10 electrons
because each orbital can
hold two electrons and there
are five orbitals in a dsubshell. Therefore a dsubshell can hold a total of
10 electrons.
e.
ℓ
Number of orbitals
possible
1
3
5
7
s
p
d
f
4.
each increasing sublevel
Number of electrons
possible
2
6
10
14
s = spin number
a.
each orbital can hold a maximum of
2 electrons
b.
one electron travels clockwise or
spin-up and has a value of +1/2,
the other counter-clockwise or spindown and has a value of -1/2
c.
further minimizes electron repulsion
Ex.
___
1s
___
2s
___ ___ ___
2p
Identify the four quantum numbers for several
electrons
n Total ℓ types
#
(sublevels)
of ℓ
in each n
m
(orbitals)
in each
sublevel
Total #
of
electrons
(2n2)
1
2
3
4
5
6
7
1
1+3
1+3+5
1+3+5+7
1+3+5+7+9
1+3+5+7+9+11
1+3+5+7+9+11+13
2
8
18
32
50
72
98
0
1
2
3
4
5
6
s
s, p
s, p, d
s, p, d, f
s,p,d,f,g
s,p,d,f,g,h
s,p,d,f,g,h,i
Sec
5.2
V.
Orbital filling diagrams-abbreviation of the
wave-mechanical atom model
A.
Symbols and their meanings:
 squares or circles - represent orbitals
 arrows - represent electrons
VI. To complicate the matter  electrons do not “fill up”
the orbitals as we might expect due to overlapping that
occurs between energy levels. Be sure to follow this
filling order when doing any kind of electron
configuration:
1s
2s
3s
4s
5s
6s
7s
2p
3p
4p
5p
6p
7p
3d
4d
5d
6d
7d
4f
5f 5g
6f 6g 6h
7f 7g 7h 7i
Ex. Phosphorus (Atomic Number = 15)
1s
2s
2p
3s
Fluorine (Atomic Number = 9)
1s
2s
2p
3p
VII. Electron configurations: an even more abbreviated
picture of the wave-mechanical model of the atom
A.
Write the energy level number next to its sublevel,
in the proper order according to the arrow diagram
B.
Write the number of electrons in each sublevel by
making it a superscript of the sublevel letter
C.
Ex: beryllium (atomic number 4)
Be = 1s22s2
D.
Try this for magnesium
VIII. The electron configuration in your reference table is
even more simplified, only showing the number of
electrons in each energy level
A.
Ca = 2 – 8 – 8 – 2
B.
Note: the energy levels are in numerical order and
are not noting the proper order in which the
sublevels were filled
IX. All of the abbreviated diagrams can indicate other
information about the atom as well
A.
Ground states and excited states
1.
2.
O atom (ground state)
a. orbital diagram:
b.
electron config.:
c.
periodic table:
O atom (excited state)(one possibility)
a.
orbital diagram:
b.
electron config.:
c.
periodic table:
3.
What to look for: if a sublevel at a higher
energy level is occupied with electrons
before completely filling up a sublevel
at a lower energy level, then the atom is
in the “excited state.”
4.
Try some!
1. Are the following in a ground state or
excited state?
2. What element is it?
a.
1s22s22p63s2
B.
b.
1s22s22p53s1
c.
2–8–8–1
d.
1s22s22p63s23p5
e.
1s
f.
1s22s22p63p2
g.
1s22s22p63s23p6
h.
1s22s12p2
2s
2p
3s
3p
If an atom is involved in a chemical reaction
or is ionized, its number of electrons will vary
from what the periodic table states it should
have.
1.
2.
Na atom:
1s
2s
2p
1s
2s
2p
Na ion:
C.
3.
Cl atom: 1s22s22p63s23p5
4.
Cl ion: 1s22s22p63s23p6
5.
Try some!
Ques: Ion or atom?
a.
Ca: 1s22s22p63s23p6
b.
Mg: 1s22s22p63s2
c.
O: 2 – 6
d.
Ne: 1s
e.
N:
2s
2p
1s22s22p5
Paired and unpaired electrons
1.
2.
3.
4.
refers to the number of electrons in an
orbital
paired = two electrons in orbital
unpaired = one electron in orbital
unoccupied = zero electrons in orbital
5.
occupied = means there are electrons in
the orbital with no mention of how many
6.
Ex: aluminum
1s
7.
2s
Ex: fluorine
1s
2s
2p
2p
3s
3p