Unit #4
Electron Configuration / Periodic Table
Instructor’s notes
Introduction
1) Remember: Atoms are:
a) Nucleus- dense positively charged center of the atom. Accounts for the mass of
an atom. Contains both protons and neutrons.
i) Proton- positively charged particle (equal to +1) in the nucleus with a mass
of 1 AMU. Approximately equal to 1.67 * 10-24g.
ii) Neutron- an uncharged particle in the nucleus; same mass as a proton.
b) Electron- negatively charged particle (equal to -1) that orbits the nucleus with
an insignificant mass. Approximately equal to 9.11 * 10-28g.
2) Remember: location and types of electrons
a) Valence- electrons in the outer most "shell". The maximum number of these for
any atom is 8. (We will later learn that they fill the "s" & "p" sublevels in the
highest energy levels).
b) Core- any electron not considered a valence electron. In between the nucleus
and the outermost "shell".
c) Electron Cloud- area surrounding the nucleus where the electrons can be
found. Both valence and core electrons are here. This area is negatively charged.
Analogous to clouds surrounding the earth.
Electron Configuration
Now we will learn how the electrons are arranged in an atom.
Remember the various atomic theories we learned about in unit #3.
Considering those ideas, what do you thing the numbers (and sometimes
letters) on the left-hand edge of the periodic chart represent ? Answer:
energy levels.
3) Energy Level- represents the most probable distance of the electron from
the nucleus of the atom. It is always represented as a positive whole
number; 1-7. Electrons in the first energy level have the lowest amount of
energy; electrons in the seventh energy level have the most energy.
(Sometimes: the energy levels are represented by the letters (from lowest
to highest energy) K, L, M, N, O, P, Q.)
a) Ground state- most stable, lowest energy position of an electron.
b) Excited state- any position of electron except the ground state. Less
stable, higher energy.
c) Lighting and the chemistry behind the characteristic colors of some
types of light bulbs:
i) Law of conservation of energy- energy is neither created nor
destroyed, it just changes forms.
(1) Examples: The suns energy is captured by chlorophyll and
other accessory pigments in plants.
(2) The energy is stored in the plant in the form of carbohydrates.
(3) After millions of years, these stored carbohydrates can become
fossil fuels.
(4) Humans can burn these fossil fuels to "create" energy. (Heat
water to make steam.)
(5) The steam spins turbines that in turn spin generators.
(6) The generators make an electric current.
(7) When the electric current is passed through a "gas" this excites
the electrons.
(8) When the electrons drop back down to their ground state they
emit light of a characteristic color.
(a) He
=
yellow light
(b) Ar
=
lavender light
(c) Kr
=
white light
(d) Xe
=
blue light
(e) Ne
=
orange / red
(f) Na
=
yellow /orange
(g) H
=
red
(h) Diagrams of how the above works.
1p
This electron is at the ground state. It has the
lowest energy and is most stable.
Ground state of Hydrogen (note position of lone electron (blue) in its
ground state).
Energy must be absorbed in
order for the electron to
become excited and "jump
up to" a higher energy level.
1p
Excited state of Hydrogen (note how the blue ground state electron is now in
a higher energy level (magenta color). This atom had to absorb energy in
order for this electron to move to a higher energy state.
1p
Energy usually in the form
of light would be emitted as
the electron returns to its
ground state.
Excited state electron returning to its ground state. (As this occurs, the atom
must lose energy. This energy will often be in the form of visible light.)
Explain how this can also account for hot metal glowing red. Answer:
again the electrons become excited and "glow" as they fall to ground state.
4) Electron Configuration- particular distribution of electrons among
available "sublevels". Often represented as: 1s22s22p3 (this as you will
later learn is the electron configuration of Nitrogen.)
a) Sublevel- indicates the shape of the orbital in which the electrons
move. (These are components of a given energy level.) Represented
by the letters: s, p, d, and f.
i) Space orbital- a highly probable location about a nucleus in which
an electron may be found. The number of space orbitals determines
the sublevel.
(1) All "s", have 1 space orbital. Represented by:
(2) All "p", have 3 space orbitals. Represented by:
(3) All "d", have 5 space orbitals. Represented by:
(4) All "f", have 7 space orbitals.
Represented by:
ii) All valence electrons will ALWAYS be located in ONLY the "s"
and "p" sublevels.
iii) Octet Rule- In most cases, chemical bonds form so that each atom
has an octet of electrons in their valence shells.
(What is an Octet? Answer: eight of something.)
5) Pauli Exclusion Principle- rule stating that each space orbital can hold a
maximum of 2 electrons and they must have opposite spins.
a) The first electron in a space orbital is represented by drawing an
upward pointing arrow in the space orbital.
i) Example:
b) The second electron in a space orbital is represented by drawing a
downward pointing arrow in the same space orbital. The opposite
spins are represented by the arrows with opposite directions.
i) Example:
c) If more than space orbital is in a sublevel; then Hund's Rule will
apply.
6) Hund's Rule- when electrons occupy the same sublevel (s, p, d, or f), to
achieve the lowest energy arrangement (most desirable) of the electrons
place the electrons into separate space orbitals, one at a time until all
space orbitals have one electron with all their spins parallel (going the
same direction), before pairing up any of the electrons.
a) Example of how a "p" sublevel would fill with electrons. Any order
other than that below would be incorrect.
i) CORRECT =
The red colored arrow
represents where the "next"
electron should be placed
according to Hund's Rule.
7) Aufbau Principle- scheme used to reproduce the electron configuration
of the ground states of atoms by successfully filling subshells with
electrons in a specific order. Remember: electron configuration:
particular distribution of electrons among available "sublevels".
a) Also known as the
i) The "Order of electron fill".
ii) The "Building up order".
iii) The "Diagonal Rule".
b) Orbital Diagram- notation showing how the orbitals of a subshell are
occupied by electrons.
i) Example:
c) Remember electron configurations are represented as: 1s22s22p3
d) Electron Dot Notation- shows the Element symbol and the valence
electrons. Also called "Lewis Symbols".
i) Symbol- represents the nucleus and the core electrons.
ii) "Dots"- represent the valence electrons and are shown as either
paired or unpaired.
iii) Example (Drawn larger to illustrate electron placement.):
Represents the first valence "p" space orbital
Represents the second valence "p"
space orbital
Represents valence "s" space orbitals
Represents the third valence "p" space orbital
Note: a second electron could be drawn next to the single electrons if
required. (Meaning if two electrons were in one or more of the "p"
space orbitals.)
e) Notice all three above notations (orbital notation, electron
configuration, and electron dot notation) represent Nitrogen.
f) Draw the "Electron Dot Notation for the following atoms:
i) Be
ii) F
iii) S
iv) Na
Now we will draw the above three notations for the first 36 elements.
This should be enough to illustrate the electron "pattern" of the periodic
table and to demonstrate how different elements can chemically react in
similar ways.
8) Series / Period - horizontal rows on the periodic chart. All elements in
the same series or period are in the same energy level.
9) Family / Group - vertical columns on the periodic chart. All elements in
the same family or group have the same number of valence electrons.
10) Electrons fill orbitals in a reasonably definite order starting with the
lowest energy level.
a) You must be able to use (and reproduce) the chart that will follow. It
will make determination of electron placement much easier.
Energy
Level
Sublevels
K
1s
L
2s
2p
M
3s
3p
3d
N
4s
4p
4d
4f
O
5s
5p
5d
5f
P
6s
6p
6d
6f
Q
7s
7p
7d
7f
To use: Simply start at "1s" and fill the sublevel according to Hund's Rule. When a given
sublevel is full, follow the arrow forward. When you reach the head of an arrow, drop
down to the arrow below and continue to fill sublevels as per Hund's Rule.
11) Exceptions to the above "filling pattern" exist. We will focus only on
those within the first 36 elements. As we encounter these exceptions, we
will discuss them since they are often representative of other elements
within the same group.
a) C
1s22sp3 (and several other Group IV elements.)
b) Cr
[Ar] 4s13d5
c) Cu
[Ar] 4s13d10
12) Noble Gas - any member of the gaseous Group VIII elements that has
an octet in their outermost sublevels. They are extremely stable
("satisfied") and therefore do not willingly react with other elements.
a) Also called "inert gases"
b) Remember: What is an "octet" ? Answer: eight of something (in
this case electrons.)
c) Remember:
i) What sublevels are the outermost ? Answer: the "s" and "p"
sublevels.
ii) What do we call electrons in these sublevels ? Answer: valence
electrons.
13) Noble Gas Core - the inner sublevel's electron configuration that
corresponds to a Noble gas and fulfills the Octet Rule.
a) The noble gas core is represented by writing the element symbol of
the noble gas and placing it in brackets.
b) It is a short-cut when writing the electron configuration of the
elements.
Therefore, which elements would have the Noble Gas Core of:
Neon ? Answer: all of the 3rd period elements.
Xenon ? Answer: all of the 7th period elements.
Bromine ? Answer: Trick question. None of them, (Bromine is not a
Noble gas).
Periodic "Short-cuts".
14) Remember: Series / Period - horizontal rows on the periodic chart. All
elements in the same series or period are in the same energy level.
15) Remember: Family / Group - vertical columns on the periodic chart.
All elements in the same family or group have the same number of
valence electrons.
16) Collective group names:
a) Halogens - Have you ever heard of this term before ? Answer:
Automotive headlights and various other forms of lighting.
Definition: very reactive non-metals found in group VII. They have
the general formula "X2", where "X" represents the halogen symbol.
How many valence electrons do the halogens have ? Answer: 7
Why do you think they are so very reactive ? Answer: with 7
valence electrons they are only one electron short of fulfilling the
octet rule and becoming relatively stable. Therefore, they react
willingly with other elements that can "supply" them with that last
electron.
17) Noble gas - any member of the gaseous Group VIII elements that has
an octet in their outermost sublevels. They are extremely stable
("satisfied") and therefore do not willingly react with other elements.
a) Also called "inert gases"
b) While originally thought to be totally un-reactive, compounds have
been formed with Xenon, Krypton, and Radon.
c) Although relatively un-reactive, many uses exist.
i) He - fills weather balloons
ii) He, Ne - mixed with O2 for use in artificial atmospheres like those
required for deep sea diving.
iii) Ar, Kr, Xe - used to produce inert atmospheres for flashbulbs and
aluminum welding (as well as MIG & TIG welding).