You
really need to read pages 100 - 120 in your book.
Electrons
orbit the nucleus in seven distinct regions called energy levels.
Energy
level one (abbreviated n =1) is closest to the nucleus, with each subsequent
energy level being further from the nucleus and energy level 7 (n = 7) being
the furthest from the nucleus.
To
keep matters simple, let’s do what Bohr did and consider only the element
Hydrogen
(only one electron)
When
the atom is in the ground state, the
electron is as close to the nucleus as it can get (the first energy level, n =
1) and electron has the least amount of energy possible for an electron in a
hydrogen atom.
When
the atom absorbs energy the electron can jump up to a higher energy level (n =
2, or n = 3)
When this happens we say the atom is in an excited
state.
ground state excited state ground state
(electron (blue) in lowest
level n = 1) (electron has jumped
to n=3) (electron
emits photon (p) and falls back to n = 1)
When
an excited electron drops back down to a lower energy level, it must get rid of
the energy that it absorbed when it jumped up.
In order to do this, the electron emits a photon (light particle) of a
specific energy (color) equal to the amount it absorbed. By measuring the energy of the photons that
excited atoms emit, scientists have been able to determine the structure of
electron energy levels.
As energy levels get farther away from the
nucleus they get bigger and can hold more electrons.
The
number of electrons that each level can hold is 2n2, where n is the energy level
number.
So
the first level can hold 2(12) = 2 electrons, the second can hold 2(22) = 8, the third can hold
18, the fourth can hold 32, etc.
Not
all electrons in a given energy level have equal amounts of energy. Each energy level has between 1 and 4 sublevels. From lowest energy to highest energy, these
levels are s, p, d, and f
(these letters come from the line spectra each produces —
sharp, principal, diffuse, and fringe)
In
any given energy level, s electrons have the least energy and f electrons have
the most.
Example in the fourth energy level the energy in each
sublevel would be 4s < 4p < 4d
< 4f
Each
sublevel can hold only a certain number of electrons, regardless of the energy
level it is in.
An
s sublevel can hold 2 electrons
A p
sublevel can hold 6 electrons
A d
sublevel can hold 10 electrons
An
f sublevel can hold 14 electrons
The
lowest energy levels don’t have all four sublevels. Below is a chart showing which sublevels are
present in each of the seven energy levels
Total
electrons in energy level
n =
1 1s 2
n =
2 2s
2p 2 + 6 = 8
n =
3 3s
3p 3d
2 + 6 + 10 = 18 even though levels 5 –7
n =
4 4s
4p 4d 4f 2
+ 6 + 10 + 14 = 32 could hold more
electrons,
n =
5 5s
5p 5d 5f 2
+ 6 + 10 + 14 = 32 they never need
to. If you add
n =
6 6s
6p 6d 6f 2
+ 6 + 10 + 14 = 32 up all the
electrons shown here
n =
7 7s
7p 7d 7f 2
+ 6 + 10 + 14 = 32 you could get to
atom # 156
An
electron configuration can be written as follows
3rd
energy level
sublevel p
Sublevels
are written from left to right in order of increasing energy.
Q:
what is wrong with the following electron configurations? A) 1s2 2s3 2p5 b) 1s2 2p5 2s2
A:
a) s3 ,
s sublevel only holds 2 electrons b) 2s2 should come
before 2p5,
it has less energy
The
electron configurations for elements 1 – 18 all follow the above rule exactly,
here are some examples
H 1s
He 1s2
Li 1s2 2s
F 1s2 2s2 2p5
P 1s2 2s2 2p6 3s2 3p3
But
when you get to Potassium (K) a weird thing happens: K 1s2 2s2 2p6 3s2 3p6 4s1
Notice
that the 4s sublevel fills before the 3d sublevel. This is because energy level 3 overlaps
energy level 4, so the highest energy electrons in level 3 (3f) are higher in
energy than the lowest energy electrons in level 4 (4s). In fact, the 6s electrons have lower energy
(are closer to the nucleus) than the 4f electrons, and therefore fill up first.
4f
5p
4d
5s
4p
3d
4s
3s
4s is closer
to the nucleus (lower energy) than 3d
2p
2s
1s
To
determine the order in which the sublevels fill, use the following chart and
draw diagonal arrows (from northeast to southwest).
1s
2s 2p
3s 3p
3d
4s 4p
4d 4f
5s 5p
5d 5f
6s 6p
6d 6f
7s 7p 7d 7f
With
larger atoms this form of notation becomes a bit unwieldy. Take a look at Zinc
1s2 2s2 2p6 3s2 3p6 4s2 3d10 instead we can write it this way: [Ar] 4s2
3d10
Since
the first 18 electrons are configured exactly like those of Argon
, we can consolidate the first 5 sublevels as being equivalent to Argon
and simply write [Ar]. This is only done with the
noble gases: He, Ne, Ar,
Kr, Xe, and Rn
Each
sublevel is made up of one or more orbitals that can
each hold no more than two electrons
Number of orbitals electrons per orbital total electrons in sublevel
Each
s sublevel has one orbital 1 2
2
Each
p sublevel has three orbitals 3 2 6
Each
d sublevel has five orbitals 5 2 10
Each
f sublevel has seven orbitals 7 2 14
The two electrons that make up an orbital each spin in the opposite way
(clockwise or counter clockwise). The magnetic
properties of some elements result from the electron’s spin.
When
electrons “park” around a nucleus, they always fill the lowest energy positions
first. Also, remember that electrons
repel each other, so they will “park” in an empty orbital if it is available
before pairing up (remember that each orbital holds two electrons).
The
best way to picture the arrangement of electrons in an atom is with orbital
notation.
In
this notation each orbital is represented with a dash and each electron is
shown as an arrow that points either up or down (opposite spins)
1s
2s 2p 2p 2p
Because
electrons repel each other, they will not pair up in one p orbital if there is
an empty one available. Orbital pairing
will not begin until each orbital already has one electron.
1s
2s 2p 2p 2p the occupy the empty sublevel
before
pairing up in the others
Every
electron in any given atom can be described in terms of four quantum numbers:
Energy level, sublevel, orbital, and spin.