Electrons in principle energy levels
I.
Arrangement of electrons in principle
energy levels.
A. What keeps an electron from crashing into
the nucleus?
1. Electrons in an atom have their energies
restricted to certain energy levels. So
instead of being all around the nucleus the
electrons are located in specific orbitals or
energy levels. An orbital is the region of
space surrounding a nucleus in which there is
a high probability of finding up to 2
electrons. The energy of these levels
increases as the distance between the
electron and the nucleus increases. This is a
quantized property, which means it is a
property that can have only certain values,
that is not all values are allowed.
2. These energy levels are called electron shells
in this book and in others are called
principle energy levels and are designated
by whole numbers 1 - 7. Electron shell is the
region of space around the nucleus that
contains electrons that have approximately
the same energy and that spend most of their
time approximately the same distance from
the nucleus. You can think of energy levels
like steps on stairs.
a) For an electron to go to a higher energy
level, it must absorb energy.
b) For an electron to go to a lower energy
level, it must emit energy. This can be
seen as color.
c) The energy emitted is equal to the energy
difference of the two levels.
d) The maximum number of electrons in an
energy level is 2n2, where n = the principle
energy level (i.e. 1 – 7).
N=5 ________ Bigger gap equals smaller  (wavelength)
N=4 ________ The bigger the gap the more energy.
N=3 ________ Energy is inversely proportional to wavelength.
N=2 ________
N=1 ________
II. Arrangement of electrons in sub-levels or
sub-shells.
A. Principle energy levels are divided into
sublevels also called sub-shells.
1. These sub-levels are labeled s, p, d, and f.
B. Each orbital has its own shape.
1. s orbital = sphere.
2. p orbital = dumb-bell shape on an axis – px
is on the x axis, ….
px holds 2 e-, py …  6 electrons total for
the p orbital.
3. d orbital = double dumb-bell shape.
4. f orbital = very strange.
C. They have a maximum number of electrons
that they can contain.
s = 2 electrons
p = 6 electrons
d = 10 electrons
f = 14 electrons
Filling the sublevels
1. Lowest energy sublevel is filled first.
Aufbau principle states that electrons
normally occupy electron sub-shells in an
atom in order of increasing energy.
1s<2s<2p<3s<3p<4s<3d<4p<5s<4d<5p<6s<
4f<5d<6p<7s<5f<6d.
The principle number being smaller does not mean that
it fills first.
2. Each sublevel can only hold so many electrons
(identified earlier: s = 2; p = 6; d = 10; f = 14).
3. Pauli exclusion principle. This says that each
orbital can have only 2 electrons.
4. Hund’s rule. This says that you fill orbitals so
that you maximize the number of unpaired
electrons.
5. Fill each sublevel before proceeding to the
next until you run out of electrons.
Bonding
Ionic bond
A. Ionic bond – the force of attraction between
ions of opposite charge. This holds ions
together in an ionic compound.
B. These oppositely charged ions are formed by
the transfer of electrons from one atom to
another.
Covalent bond
Chemical bond formed by sharing of electrons
of the atoms
The smallest unit of a covalent bond is a
covalent compound
Covalent vs. Ionic compounds
Covalent compounds have lower melting
points than ionic compounds. O2 is a
covalent compound.
Covalent compounds do not conduct
electric current in a liquid solution.
Ionic compounds do conduct
electricity.
All of the diatomics are covalent compounds H2,
N2, O2, F2, Cl2, Br2, I2
Energy
Bond length: the distance between two nuclei of
atoms that are bonded together.
Energy is given off in the formation of bonds.
You must add energy to break bonds.
Electronegativity of covalent bonds.
The smaller the radius of the atom the
greater the attraction of the nucleus to
the outer electrons.
Coulomb’s Law: attractive force
between the electron and the proton
increases as the distance decreases
Distance
Two
independent
particles
As the distance decreases the energy
decreases until a bond is formed. Then if
you try to make the distance smaller you
must add a large amount of energy.
Atoms with fewer electrons have less
shielding of the outer electrons.
Shielding would block the charge of
the nucleus from the electron
a. As you get more electrons you get
more distance between the outer
electrons and the nucleus.
When filling the same period the
electronegativity increases as the
nuclear charge increases.
Polar Bonds
Differences in electronegativity cause unequal
sharing of electrons in covalent bonds.
Electronegativity differences
0 - 0.4
bond
0.4 - 1.0
1.0-2.0
2.0 - 4.0
C.
Example
nonpolar covalent
moderately polar covalent
very polar covalent
ionic bond
HCl
EN of H=2.1
EN of Cl=3.0
The difference is .9 this is a polar bond
H 2O
EN of H=2.1
EN of O=3.5
The difference is 1.4 polar bond
N2
EN of N=3.0
EN of N=3.0
The difference is 0 covalent bond
LiF
EN of Li=1.0
EN of F = 4.0
Ionic bond
Metallic Bond
The metallic bond is formed between two or
more metals where the valance electrons are
shared between the metals. This sharing of
electrons makes a conduction band or sea
of electrons around the metals and this is
why metals are good conductors of
electricity.
The coordinate Covalent bond.
A type of a chemical bond formed when one
atom supplies both electrons of the electron
pair that make a chemical bond and the other
atom only has an empty orbital.
H
Example NH4+
N
H
H
+
H+
H
H
N
H
H
EN of N=3.0
EN of H=2.1
.9 Polar
Why do we care about all this bonding?
The internal bonding of the molecules
determines the way the atoms come together
and the structures that are formed.
Why do we do the modeling lab that shows
that the structure of H2O is bent and not
linear? It is because of the bent shape of
water that water has the properties that it
does. The bent shape allows for hydrogen
bonding between the different molecules of
water. Figure 8.26 p.241 in the book. See
the animation
http://www.elmhurst.edu/~chm/vchembook/
161Ahydrogenbond.html
This hydrogen bonding holds water together
and gives it a much higher boiling point than
you would expect if the there was no hydrogen
bonding. If water did not have a bent shape, it
would not be a polar compound, it would not
hydrogen bond and we would not be here.