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
to certain energy
levels. So instead of being all around the nucleus the electrons are located
in specific
or energy levels. An orbital is the region of
space surrounding a nucleus in which there is a high probability of finding
up to electrons. The
of these levels increases as the
distance between the electron and the nucleus
. This is a
, 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
and are designated by whole numbers
1 - 7.
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
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
energy.
b) For an electron to go to a lower energy level, it must
energy.
This can be seen as
.
c) The energy emitted is
to the energy difference of the two
levels.
d) The maximum number of electrons in an energy level is
, 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
.
1. These sub-levels are labeled
, and
B. Each orbital has its own shape.
1. s orbital =
2. p orbital =
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=
electrons
p=
electrons
d=
electrons
f=
electrons
 Filling the sublevels
1. Lowest energy sublevel is filled first.
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
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.
exclusion principle. This says that each orbital can have only 2
electrons.
4.
rule. This says that you fill orbitals so that you maximize the
number of
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
charge. This
holds ions together in an ionic compound.
B.
These oppositely charged ions are formed by the
of electrons
from one atom to another.
Covalent bond
Chemical bond formed by
of electrons of the atoms
The smallest unit of a covalent bond is a
compound
Covalent vs. Ionic compounds
Covalent compounds have
melting points than ionic
compounds. O2 is a covalent compound.
Covalent compounds
conduct electric current in a liquid solution.
Ionic compounds do conduct electricity.
All of the
are covalent compounds H2, N2, O2, F2, Cl2, Br2, I2
Bond length: the distance between two
of atoms that are bonded
together.
Energy
Energy is given off in the formation of bonds. You must add energy to break
bonds.
Electronegativity of covalent bonds.
The
the radius of the atom the greater the
attraction of the nucleus to the outer electrons.
Coulomb’s Law:
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
of electrons in
covalent bonds.
Electronegativity differences
nonpolar covalent bond
moderately polar covalent
very polar covalent
ionic bond
C.
Example
HCl
EN
of
H=2.1
EN
of
Cl=3.0
The difference is .9 this is a polar bond
H2O
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
or sea of electrons around the metals and this is
why metals are good conductors of electricity.
The
Covalent bond.
A type of a chemical bond formed when
supplies both electrons of the
electron pair that make a chemical bond and the other atom only has an
empty orbital.
Example NH4+
H
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
is
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
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
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.