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Early Chemistry
 Early Chemists only
believed in 1
element: Dirt
 Later Chemists
believed in 4
elements:
 Air
 Earth
 Fire
 Water
 Various combinations
of these produced
various compounds
Atomic Structure
All matter is composed of atoms.
Understanding the structure of atoms is
critical to understanding the properties
of matter.
properties of solid materials depend on
the geometrical atomic arrangements, and
the interactions between constituent
atoms.
history of the atom
460 BC Democritus developed the idea of atoms!
he pounded up materials in his
pestle and mortar until he had
reduced them to smaller and
smaller particles which he
called
ATOMA
(greek for indivisible)
It took ~2400 years from when it was
conceived to the time experimental
evidence prove of the atom existence.
history of the atom
1808
John Dalton
suggested that all matter was
made up of tiny spheres that
were able to bounce around
with perfect elasticity and
called them
ATOMS
history of the atom
1898
Joseph John Thompson / Cambridge
found that atoms could
sometimes eject a far smaller
negative particle which he
called
ELECTRON
1906 Nobel prize in Physics
History of the atom
1910 Ernest Rutherford / Cambridge
student of Thompson
proposed a more detailed model
with a central nucleus:
positive charge was all in a
central nucleus. With this
holding the electrons in place by
electrical attraction
1908 Nobel prize in Chemistry
History of the atom
1913
Niels Bohr / Danish / a football fanatic
studied under Rutherford at the
Victoria University in Manchester.
Bohr refined Rutherford's idea by
adding that the electrons were
in orbits. Rather like planets
orbiting the sun. With each orbit
only able to contain a set
number of electrons.
1922 Nobel prize in physics
Bohr’s atom
Rutherford’s model predicted a rainbow of colors rather than
discrete lines obtained from an atomic line spectra.
To explain the line spectra, Bohr proposed that electrons of
specific energy moved in circular orbits around the nucleus
and could not exist between these orbits.
Atomic Structure
Atoms are composed of
protons – positively charged particles
neutrons – neutral particles
nucleus
electrons – negatively charged particles
in orbitals surrounding the
nucleus.
Atomic Structure
● Every different atom has a characteristic
number of protons in the nucleus.
● atomic number (Z) = number of protons
● For an electrically neutral atom,
atomic number = number of electrons.
● Atoms with the same atomic number have
the same chemical properties and belong
to the same element.
Atomic Structure
 Z ranges from 1 for hydrogen to 92 for
uranium (the highest for the naturally
occuring elements).
Atomic Structure
The atomic mass (A) of a specific atom:
the sum of the number of protons and
neutrons within the nucleus.
mass number: A = Z + N
number of protons is the same for all atoms
of a given element,
number of neutrons (N) may be variable.
Atomic Structure
The number of protons in the nucleus of the
atom is equal to the atomic number (Z).
The number of electrons in a neutral atom is
equal to the number of protons.
The mass number of the atom (M) is equal to the
sum of the number of protons and neutrons in
the nucleus.
The number of neutrons is equal to the
difference between the mass number of the
atom (M) and the atomic number (Z).
Atomic Structure/isotopes
atoms of some elements have two or more
different atomic masses, called isotopes.
Atomic Structure
Atomic weight: Weighted average of the atomic
masses of the atom’s naturally occurring isotopes.
Boron consists of the isotopes:
19.7% B-10 (5p+5n) and 80.3% B-11 (5p+6n).
atomic weight for Boron,
B = (19.7 x 10)+(80.3 x 11)]/100= 10.8 amu
Bromine’ isotopes:
50.5% Br-79 and 49.5% Br-81.
atomic weight for Bromine,
Br = [(50.5 x 79)+(49.5 x 81)]/100= 80.0 amu
Atomic Structure
In one mole of a substance there are 6.022 x1023
(Avogadro’s number) atoms or molecules.
Atomic weight = weight of 6.023 x 1023 atoms
For example, the atomic weight of iron is 55.85
amu/atom, or 55.85 g/mol.
1 amu/atom = 1g/mol = 1 dalton
Atomic mass unit (amu): 1⁄12 of the atomic mass
of carbon
atomic mass of C:12.011 amu / of H:1.008 amu
subatomic particles
particle
Electron (e-)
Proton (p)
Neutron (n)
Mass (g)
9.11x10-28
1.67x10-24
1.67x10-24
Charge (C/eV)
-1.6x10-19
-1
+1.6x10-19
+1
0
0
Proton is 1837 times heavier than an electron.
Neutron is 1842 times heavier than an electron.
Electron is much lighter with respect to the
protons and neutrons
Atomic Structure
HELIUM
ATOM
Shell
proton
nucleus
# electrons = # protons
++
N
-
electron
N
neutron
ATOMIC MASS NUMBER = number
of protons + number of neutrons
ATOMIC NUMBER = number of
protons
Lithium
Protons
Neutrons
three electrons
three protons
four neutrons.
Electrons
Beryllium
Protons
Neutrons
four electrons
four protons
five neutrons.
Electrons
Boron
Protons
Neutrons
five electrons
five protons
six neutrons.
Electrons
Carbon
Protons
Neutrons
six electrons
six protons
six neutrons.
Electrons
Nitrogen
Protons
Neutrons
seven electrons
seven protons
seven neutrons.
Electrons
Oxygen
Protons
Neutrons
eight electrons
eight protons
eight neutrons
Electrons
Fluorine
Protons
Neutrons
nine electrons
nine protons
ten neutrons.
Electrons
Neon
Protons
Neutrons
ten electrons
ten protons
ten neutrons
Electrons
Sodium
Protons
Neutrons
eleven electrons
eleven protons
twelve neutrons
Electrons
Atomic structure
How many protons, neutrons and electrons?
Atomic structure
How many protons, neutrons and electrons?
Atomic structure
How many protons, neutrons and electrons?
Atomic structure
How many protons, neutrons and electrons?
Atomic structure
How many protons, neutrons and electrons?
Atomic structure
The charge and mass number of an electron are:
a) charge = 0, Mass number = 1
b) charge = -1, Mass number = 0
c) charge = +1, Mass number = 1
d) charge = +1, Mass number = 0
The charge and mass number of a neutron are?
a) charge = +1, Mass number = 1
b) charge = 0, Mass number = 1
c) charge = +1, Mass number = 0
d) charge = -1, Mass number = 0
Atomic structure
Which of the following has 25 protons and 31
neutrons?
a) 56Mn
b)56Ga
c) 25Ga
d)31Mn
e)56Ba
Atomic structure
Why does chlorine have an atomic mass of 35.5,
which is not a whole number?
a) Chlorine contains an extra electron which
makes it weigh more than 35.
b) Chlorine contains 17 protons and 18.5 neutrons
c) Chlorine normally exists in an excited state, and
so it weighs more than 35.
d) The chlorine was not pure when its atomic mass
was measured.
e) Chlorine, as found in nature, contains a mixture
of the isotopes 35Cl and 37Cl, in such proportions
as to give an average atomic mass of 35.5
Atomic structure
The two main parts of an atom are?
a) nucleus and electron energy levels
b) nucleons and protons
c) oxidation number and valence
d) protons and neutrons
e) protons and electrons
Atomic structure
The nucleus of the element having atomic
number 25 and atomic weight 55 will
contain?
a)
b)
c)
d)
25 protons and 30 neutrons
30 protons and 25 neutrons
55 protons
55 neutrons
Atomic structure
A beryllium atom has 4 protons, 5
neutrons, and 4 electrons. What is the
mass number of this atom?
a)
b)
c)
d)
e)
4
5
8
9
13
Atomic structure
The smallest particle into which an element can be
divided and still have the properties of that
element
a) nucleus
b) electron
c) atom
d) neutron
How would you describe the nucleus?
a) dense, positively charged
b) mostly empty space, positively charged
c) tiny, negatively charged
d) dense, negatively charged
Atomic structure
Where are electrons likely to be found?
a) in the nucleus
b) in electron clouds
c) mixed throughout an atom
d) in definite paths
Every atom of a given element has the same
number of
a) protons
b) neutrons
c) electrons
d) isotopes
Atomic structure
What is the meaning of the word atom?
a) dividable
b) invisible
c) hard particles
d) not able to be divided
Which statement is true about isotopes of the
same element?
a) They have the same number of protons
b) They have the same number of neutrons
c) They have a different atomic number
d) They have the same mass
Atomic structure
Which has the least mass in an atom?
a) nucleus
b) proton
c) neutron
d) electron
If an isotope of uranium, uranium-235, has 92
protons, how many protons does the isotope
uranium-238 have?
a) 92
b) 95
c) 143
d) 146
Atomic structure
What is the atomic mass number of a Ba atom?
a) 56
b) 81
c) 137
d) 25
Atomic structure
A carbon atom with 6 protons, 6 electrons, and 6
neutrons would have a mass number of
a) 6
b) 12
c) 15
d) 18
The number at the top is the
a) atomic number
b) element name
c) atomic mass
d) chemical symbol
Atomic structure
How many electrons does a neutral Cl atom
contain?
a)16
b)17
c) 18
d)19
What is the difference between atomic mass and
atomic weight?
Atomic mass is the mass of a single atom or an individual
isotope. The atomic weight is the average mass of all
naturally occurring isotopes of an element.
Atomic Structure
Neutral atoms have the same number
of protons and electrons.
Ions are charged atoms.
●cations – have more protons than
electrons and are positively charged
●anions – have more electrons than
protons and are negatively charged
Atomic Structure
If a neutral atom looses one or more electrons
it becomes a cation.
Na
11 protons
11 electrons
e- +
11 protons
+
Na
10 electrons
If a neutral atom gains one or more electrons
it becomes an anion.
Cl
17 protons
17 electrons
+ e-
Cl-
17 protons
18 electrons
Bohr Atomic model
electrons are assumed to
revolve around the atomic
nucleus in discrete orbitals,
and the position of any
particular electron is more
or less well defined in terms
of its orbital.
Electrons are permitted to
have only specific values of
energy.
Bohr Atomic model
excitation vs relaxation
An electron may change
energy by making a quantum
jump either to an higher
energy (with absorption of
energy) or to a lower energy
(with emission of energy).
relaxation
excitation
Quantum Mechanics
Unfortunately, extremely small particles
(electrons) do not follow the laws of classical
(Newtonian) physics.
The new physics that mathematically treats small
particles is called
Quantum Mechanics.
electron distribution
wave-mechanical model
an electron is no longer
treated as a particle moving
in a discrete orbital;
electron is considered to
exhibit both wavelike and
particle-like characteristics.
The position of an electron is
described by a probability
distribution // electron
cloud.
Quantum Mechanics
Wave behavior is described with the wave function
ψ, incorporating the wave and particle features of
electrons (Erwin Schrödinger)
The probability of finding an electron in a certain
area of space is
proportional to ψ2
electron density.
Austrian; 1933 Nobel
prize in physics
Quantum Mechanics
Heisenberg’s uncertainty principle
more precisely the position of some particle is determined,
the less precisely its momentum can be known
A macroscale analogy…
High Shutter Speed
Can judge location,
but not speed.
Low Shutter Speed
Can judge speed,
But not location
Heisenberg’s uncertainty principle
we cannot precisely measure the momentum and
the position of an electron at the same time.
As the momentum of the electron is more and more
certain, the position of the electron becomes less
and less certain, and vice versa.
n = 2.5 cannot exist as a principal quantum number.
There must be an integral number of wavelengths
(n) in order for an electron to maintain a
standing wave. If there were to be partial waves,
the whole and partial waves would cancel each
other out and the particle would not move.
Quantum Mechanics
The Schrödinger equation
specifies possible energy states
an electron can occupy.
The energy states and wave
functions are characterized by
a set of quantum numbers.
Instead of orbits in the Bohr
model, quantum numbers and
wave functions describe atomic
orbitals in quantum mechanics.
quantum numbers
every electron in an atom is characterized by four
quantum numbers.
There are three quantum numbers necessary to
describe an atomic orbital.
 The principal quantum number (n)
designates size
 The angular moment quantum number (l)
describes shape
 The magnetic quantum number (ml)
specifies orientation
Principal Quantum Number (n)
n designates the size of the orbital.
Larger values of n correspond to larger orbitals.
The allowed values of n are integers: 1, 2, 3 and so
forth.
A collection of orbitals with the same value of n is
frequently called a shell.
n
K
L
M
N
O
P
.......
1
2
3
4
5
6
……
Angular moment Quantum Number (l)
l signifies the subshell
l describes the shape of the orbital.
l values range from 0 to n – 1
Example: If n = 2, l can be 0 or 1.
n
l
subshell
energy state
1
0
2
0,1
s
1
s,p
3
3
4
5
6
0,1,2 0,1,2,3 0  4 0  5
s,p,d s,p,d,f s,p g s,p h
5
7
9
11
Magnetic Quantum Number (ml)
describes the orientation of the orbital in space.
ml are integers that depend on l: – l,…0,…+l
ml identifies # of energy states for each subshell
For an s subshell: a single energy state
For p, d, and f subshells: 3, 5, and 7 energy states
Number of available electron states for
initial shells and subshells
Principal
Quantum No:
Subshell
No. of energy States:
Shell
l
1
2
K
L
s /0
ml
1/0
2
3
M
4
N
5
O
6
P
s/0
p/1
s/0
p/1
d/2
s/0
p/1
d/2
f/3
s/0
p/1
d/2
f/3
g/4
s/0
p/1
d/2
f/3
g/4
h/5
1/0
3 / -1,0,+1
1/0
3 / -1,0,+1
5 / -2,-1,0,+1,+2
1/0
3 / -1,0,+1
5 / -2,-1,0,+1,+2
7 / -3,-2,-1,0,+1,+2,+3
1/0
3 / -1,0,+1
5 / -2,-1,0,+1,+2
7 / -3,-2,-1,0,1,2,3
9 / -4,-3,-2,-1,0,+1,+2,+3,+4
1/0
3 / -1,0,+1
5 / -2,-1,0,+1,+2
7 / -3,-2,-1,0,1,2,3
9 / -4,-3,-2,-1,0,1,2,3,4
11 / -5,-4,-3,-2,-1,0,+1,+2,+3,+4,+5
2
6
2
6
10
2
6
10
14
2
6
10
14
18
2
6
10
14
18
22
n
Number of Electrons
Per Subshell Per Shell
2
8
18
32
50
72
Atomic orbitals
An s subshell has one orbital which is spherically
shaped.
If you were to measure where the electron was
within an s subshell many, many times and plot
the results on a graph you would get something
like this.
Atomic Orbitals
p orbitals: a dumbbell shape with electrons on
either side of the nucleus in tear drop shaped lobes
Three orientations:
l = 1 (as required for a p orbital)
ml = –1, 0, +1
Atomic Orbitals
The d orbitals:
Five orientations:
l = 2 (as required for a d orbital)
ml = –2, –1, 0, +1, +2
Atomic orbitals
d-orbitals are followed by the seven f-orbitals.
7 orientations:
l = 3 (as required for a d orbital)
ml = -3, –2, –1, 0, +1, +2, +3
Quantum Numbers
To summarize quantum numbers:
principal (n) – size
angular (l) – shape
magnetic (ml) – orientation
Required to
describe an
atomic orbital
principal (n = 2)
2px
related to the magnetic
quantum number (ml )
angular momentum (l = 1)
electron spin (ms) direction of spin
Required to describe an electron in an atomic orbital
Electron Spin Quantum Number-ms
used to specify an electron’s spin.
There are two possible
directions of spin.
Allowed values of ms
are +½ and −½.
Pauli exclusion principle
No Two Electrons in an Atom Can Have the Same
Four Quantum Numbers; the same values for n, l,
ml, and ms.
Although the first three quantum numbers identify
a specific orbital and may have the same values,
the fourth is significant and must have opposite
spins.
a set of quantum numbers is specific to a certain
electron.
Quantum numbers / Q
An electron with
n = 2, ℓ = 1, ml = −1, and ms = +1/2
is found in the same atom as a second electron with
n = 2, ℓ = 1, ml = −1.
What is the spin quantum number for the second
electron?
Since the first three quantum numbers are identical for
these two electrons, we know that they are in the same
orbital. As a result, the spin quantum number for the second
electron cannot be the same as the spin quantum number
for the first electron. This means that the spin quantum
number for the second electron must be ms = −1/2.
Quantum numbers / Q
An electron with
n = 5, ℓ = 4, ml = 3, and ms = −1/2
is found in the same atom as a second electron with
n = 5, ℓ = 4, ml = 3.
ms = ?
Since the first three quantum numbers are identical
for these two electrons, we know that they are in
the same orbital. As a result, the spin quantum
number for the second electron cannot be the same
as the spin quantum number for the first electron.
This means that the spin quantum number for the
second electron must be ms = +1/2.
Quantum numbers / Q
Can an electron with
n = 1, ℓ = 0, ml = 0, and ms = +1/2
exist in the same atom with a 2nd electron with
n = 2, ℓ = 0, ml = 0, and ms = +1/2?
Since these two electrons are in different
orbitals, they occupy different regions of space
within the atom.
As a result, their spin quantum numbers can be
the same, and thus these two electrons can
exist in the same atom.
Atomic structure
Maximum number of electrons in a subshell with
l = 3 and n = 4 is
a)
b)
c)
d)
e)
10
12
14
16
18
Principal
Quantum No:
4
Subshell
n
l
s/0
p/1
d/2
f/3
No. of energy States:
ml
1/0
3 / -1,0,+1
5 / -2,-1,0,+1,+2
7 / -3,-2,-1,0,+1,+2,+3
Number of
Electrons
Per Subshell
2
6
10
14
Atomic structure
The lowest principal quantum number for
an electron is?
a)
b)
c)
d)
e)
0
1
2
3
4
Atomic structure
Which sublevel can by occupied by a
maximum of 10 electrons?
a)
b)
c)
d)
s
p
d
f
Subshell
l
s/0
p/1
d/2
f/3
No. of energy States:
ml
1/0
3 / -1,0,+1
5 / -2,-1,0,+1,+2
7 / -3,-2,-1,0,+1,+2,+3
Number of
Electrons
Per Subshell
2
6
10
14
Atomic structure
The K, L and M shells of an atom are full.
Its atomic number is_______.
a)
b)
c)
d)
18
20
10
12
Principal
Quantum No:
Subshell
No. of energy States:
Shell
l
1
2
K
L
s /0
ml
1/0
2
3
M
s/0
p/1
s/0
p/1
d/2
1/0
3 / -1,0,+1
1/0
3 / -1,0,+1
5 / -2,-1,0,+1,+2
2
6
2
6
10
n
Number of Electrons
Per Subshell Per Shell
2
8
18
Atomic structure
If n=3, and l=2, then what are the possible
values of ml ?
Since ml must range from –l to +l, then ml can
be: -2, -1, 0, 1, or 2.
Atomic structure
State whether an electron can be described by
each of the following sets of quantum number.
If a set is not possible, state why not.
a) n = 2, l = 1, ml = -1
b) n = 1, l = 1, ml = +1
c) n = 4, l = 3, ml = +3
d) n = 3, l = 1, ml = -3
Atomic structure
Replace the question marks by suitable
responses in the following quantum number
assignments.
a) n = 3, l = 1, ml = ?
b) n = 4, l = ?, ml = -2
c) n = ?, l = 3, ml = ?
Atomic structure
Replace the question marks by suitable responses
in the following quantum number assignments.
a) n = 3, l = 1, ml = -1,0,1
b) n = 4, l = 2, ml = -2
c) n = 4, l = 3, ml = -3,-2,-1,0,1,2,3
Principal Quantum No:
1
2
3
4
n
Subshell
l
No. of energy States:
s /0
0
s/0
p/1
s/0
p/1
d/2
f/3
0
-1,0,+1
0
-1,0,+1
-2,-1,0,+1,+2
-3,-2,-1,0,+1,+2,+3
ml
Atomic structure / Q
Provide the three quantum numbers
describing each of the three p orbitals in
the 2p subshell.
n
l
ml
2px
2
1
0
2py
2
1
1
2pz
2
1
-1
Principal Quantum No: n
2
Subshell
l
No. of energy States:
s/0
0
p/1
-1,0,+1
ml
Atomic structure
For n = 1, determine the possible values of l.
For each value of l, assign the appropriate letter
designation & determine the possible values of ml.
n=1
l = 0s
Principal Quantum No: n
1
ml= 0
Subshell
s /0
l
No. of energy states:
0
ml
Atomic structure
How many orbitals in shell n = 1?
1st Shell has only the s orbital!
How many electrons possible?
S orbital can hold only 2 electrons!
Atomic structure
For n = 2, determine the possible values of l.
For each value of l, assign the appropriate letter
designation & determine the possible values of ml.
Principal
Quantum No:
n
2
Subshell
No. of energy States:
Shell
l
L
s/0
ml
0
-1, 0, +1
p/1
How many orbitals in shell n = 2 ?
How many electrons possible?
Principal Quantum
No:
n
2
Subshell
Shell
l
L
s/0
p/1
Number of Electrons
Per Subshell
Per Shell
2
6
8
Atomic structure
For n = 3, determine the possible values of l. For
each value of l, assign the appropriate letter
designation & determine the possible values of ml.
Shell
n
l
ml
3
M
s/0
p/1
d/2
1/0
3 / -1,0,+1
5 / -2,-1,0,+1,+2
How many orbitals in shell n = 3?
How many electrons possible?
Shell
Number of Electrons
n
l
3
M
s/0
p/1
d/2
Per Subshell
Per Shell
2
6
10
18
Atomic structure
For n = 4, determine the possible values of l.
For each value of l, assign the appropriate
letter designation & determine the possible
values of ml.
Principal
Quantum No: Shell
n
4
N
Subshell
No. of energy States:
l
ml
s/0
p/1
d/2
f/3
1/0
3 / -1,0,+1
5 / -2,-1,0,+1,+2
7 / -3,-2,-1,0,+1,+2,+3
Atomic structure
Provide the four quantum numbers describing
each of the two electrons in the 3s orbital.
n
3
3
l
0
0
ml
0
0
ms
-1/2
+1/2
Quantum Numbers:
A Macroscale Analogy
n
- indicates which train (shell)
 l
- indicates which car (subshell)
 ml - indicates which row (orbital)
 ms - indicates which seat (spin)
No two people can have exactly the
same ticket (sit in the same seat).
Electron energy states
electrons have discrete energy states
they fill up the lowest possible energy states in the
electron shells
and subshells,
When all the electrons
occupy the lowest
possible energies in
accord with the
foregoing restrictions,
Energy states for
an atom is said to be
a Na atom
in its ground state.
Electron configurations
Most elements: Electron configuration not stable!
Element
Hydrogen
Helium
Lithium
Beryllium
Boron
Carbon
...
Neon
Sodium
Magnesium
Aluminum
...
Argon
...
Krypton
Atomic #
1
2
3
4
5
6
Electron configuration
1s 1
1s 2
(stable)
1s 2 2s 1
1s 2 2s 2
1s 2 2s 2 2p 1
1s 2 2s 2 2p 2
...
10
11
12
13
1s 2 2s 2 2p 6
(stable)
1s 2 2s 2 2p 6 3s 1
1s 2 2s 2 2p 6 3s 2
1s 2 2s 2 2p 6 3s 2 3p 1
...
18
...
36
1s 2 2s 2 2p 6 3s 2 3p 6
(stable)
...
1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 (stable)
Electron Configurations
 Valence electrons – those in unfilled shells
 Filled shells more stable
 Valence electrons are most available for
bonding and tend to control the chemical,
electrical, thermal, optical properties
 example:
C (atomic number = 6)
1s2 2s2 2p2
valence electrons
Electron Configurations
Q. the full electronic configuration of an element is
1s22s22p5.
How many electrons does it have in its outer
shell?
A. # of outer shell-valence electrons: 7
Q: the full electronic configuration of an element.
1s22s22p5.
What is its atomic number?
A. Atomic number: 9
Electronic Configurations
Fe-atomic # = 26
4d
4p
Energy
3d
4s
3p
3s
1s2 2s2 2p6 3s2 3p6 3d
6
4s2
N-shell n = 4
6 electrons left to be located
total # e-s
20 electrons
M-shell n = 3; 18 electrons
2p
2s
L-shell n = 2; 10 electrons
1s
K-shell n = 1; 2 electrons
Valence electrons
Element Atomic esymbol number configuration
# of valence
electrons
H
1
1s1
1
He
Li
Be
B
C
N
O
F
Ne
2
3
4
5
6
7
8
9
10
1s2
1s22s1
1s22s2
1s22s22p1
1s22s22p2
1s22s22p3
1s22s22p4
1s22s22p5
1s22s22p6
2
1
2
3
4
5
6
7
8
Order of Subshell Filling
The electron configurations of the first ten
elements illustrate this point.
Electron configurations for common elements
Shells and subshells
In multi-electron atoms, the energies of the atomic
o
orbitals are split.
Splitting of energy
levels refers to the
splitting of a shell (n=3)
into subshells of
different energies (3s,
3p, 3d)
Splitting of Shells into subshells
3s subshell
3;
l = 0)
3rd shell (n
(n ==3p
3)subshell
(n3d
= subshell
3; l = 1) (n = 3; l = 2)
2s
2ndsubshell
shell (n = 2)2p subshell (n = 2; l = 1)
(n = 2; l = 0)
Electron Configurations
rules for electron configurations:
● Electrons will reside in the
lowest possible energy orbitals
● Each orbital can accommodate
a maximum of two electrons.
● Electrons will not pair in
degenerate orbitals if an
empty orbital is available.
● Orbitals will fill in the order
..3p6/4s2/3d10/4p6/5s2/4d10/
5p6/6s2/4f14/5d10/6p6/7s2
Energy Level Diagram of a multi-electron atom
6s
6p
5d
4f
32
5s
5p
4d
18
4s
Arbitrary
Energy Scale
4p
3d
18
3s
3p
8
2s
2p
8
1s
2
NUCLEUS
Electron Configurations
The electron configuration describes how the electrons are
distributed in the various atomic orbitals.
In a ground state hydrogen atom, the electron is found in the
1s orbital.
Ground state electron
configuration of hydrogen
Energy
principal (n = 1)
2s
1s
2p
2p
2p
1
1s
number of electrons in
the orbital or subshell
angular momentum (l = 0)
The use of an up arrow indicates an electron
with ms = + ½
Electron Configurations
Energy
If hydrogen’s electron is found in a higher energy
orbital, the atom is in an excited state.
2s
2p
2p
2p
A possible excited state electron
configuration of hydrogen
1s
2s1
Electron Configurations
Pauli exclusion principle: no two electrons in an
atom can have the same four quantum numbers.
The ground state electron
configuration of helium
Energy
2p
2p
2p
1s
2s
2
Quantum number
Principal (n)
1s
describes the 1s orbital
Angular moment (l)
Magnetic (ml)
describes the electrons in the 1s orbital
Electron spin (ms)
1
0
0
+½
1
0
0
‒½
Electron Configurations
The Aufbau principle states that electrons are
added to the lowest energy orbitals first before
moving to higher energy orbitals.
Li has a total of 3 electrons
The ground state electron
configuration of Li
Energy
2p
2s
1s
2p
2p
The third electron must go in
the next available orbital with
the lowest possible energy.
1s22s1
The 1s orbital can only accommodate 2 electrons
(Pauli exclusion principle)
Electron Configurations
Be has a total of 4 electrons
Energy
2p
2p
2p
2s
1s
The ground state electron
configuration of Be
1s22s2
Electron Configurations
B has a total of 5 electrons
Energy
2p
2s
1s
2p
2p
The ground state electron
configuration of B
1s 2s 2p
2
2
1
Electron Configurations
Hund’s rule, the most stable arrangement of
electrons is the one in which the number of electrons
with the same spin is maximized.
C has a total of 6 electrons The ground state electron
configuration of C
1s 2s 2p
2
Energy
2p
2p
2p
2
2
2s The 2p orbitals are of equal energy, or degenerate.
1s
Put 1 electron in each before pairing (Hund’s rule).
Electron Configurations
N has a total of 7 electrons
Energy
2p
2p
2p
2s The 2p orbitals are of equal energy, or degenerate.
1s
Put 1 electron in each before pairing (Hund’s rule).
The ground state electron
configuration of N
1s 2s 2p
2
2
3
Electron Configurations
O has a total of 8 electrons
Energy
2p
2s
1s
2p
2p
Once all the 2p orbitals are singly occupied,
additional electrons will have to pair with
those already in the orbitals.
The ground state electron
configuration of O
1s22s22p4
Electron Configurations
F has a total of 9 electrons
Energy
2p
2s
1s
2p
2p
When there are one or more unpaired
electrons, as in the case of oxygen and
fluorine, the atom is called paramagnetic.
The ground state electron
configuration of F
1s22s22p5
Electron Configurations
Ne has a total of 10 electrons
Energy
2p
2s
1s
2p
2p
When all of the electrons in an atom are
paired, as in neon, it is called diamagnetic.
The ground state electron
configuration of Ne
1s 2s 2p
2
2
6
learning check
Write the electron configuration and give the orbital
diagram of a calcium (Ca) atom (Z = 20).
Z = 20, Ca has 20 electrons.
Each s subshell can contain a maximum of
two electrons, whereas each p subshell can
contain a maximum of six electrons.
Solution
Ca
1s22s22p63s23p64s2
1s2 2s2
2p6
3s2
3p6
Remember that the 4s orbital fills
before the 3d orbitals.
4s2
learning check
electron configuration for an arsenic atom (Z = 33) in the
ground state.
Z = 18 for Ar.
The order of filling beyond the noble gas
core is 4s, 3d, and 4p. Fifteen electrons go
into these subshells because there are 33 –
18 = 15 electrons in As beyond its noble gas
core.
Solution
As
[Ar]4s23d104p3
Arsenic is a p-block element; therefore,
we should expect its outermost electrons
to reside in a p subshell.
2
2
6
2
6
2
3
10
electron configuration?
Number of Energy Levels: 3
First Energy Level: 2
Second Energy Level: 8
1s22s22p63s23p1
Third Energy Level: 3
electron configuration?
Number of Energy Levels: 4
First Energy Level: 2
Second Energy Level: 8
22s22p63s23p64s1
1s
Third Energy Level: 8
Fourth Energy Level: 1
electron configuration?
Number of Energy Levels: 4
First Energy Level: 2
Second Energy Level: 8
1s22s22p63s23p64s23d2
Third Energy Level: 10
Fourth Energy Level: 2
electron configuration?
Number of Energy Levels: 4
First Energy Level: 2
Second Energy Level: 8
22s22p63s23p64s13d5
1s
Third Energy Level: 13
Fourth Energy Level: 1
electron configuration?
Number of Energy Levels: 4
First Energy Level: 2
Second Energy Level: 8
22s22p63s23p64s23d5
1s
Third Energy Level: 13
Fourth Energy Level: 2
electron configuration?
Number of Energy Levels: 4
First Energy Level: 2
Second Energy Level: 8
1s22s22p63s23p64s23d6
Third Energy Level: 14
Fourth Energy Level: 2
electron configuration?
Number of Energy Levels: 4
First Energy Level: 2
Second Energy Level: 8
22s22p63s23p64s13d10
1s
Third Energy Level: 18
Fourth Energy Level: 1
electron configuration?
Number of Energy Levels: 4
First Energy Level: 2
Second Energy Level: 8
22s22p63s23p64s23d10
1s
Third Energy Level: 18
Fourth Energy Level: 2
Valence electrons
They occupy the outermost shell.
They participate in the bonding between atoms
They dictate the physical and chemical properties
if the outermost or valence electron shell are
completely filled: stable electron configurations
occupation of the s and p states for the outermost
shell by a total of eight electrons, in neon (Ne),
argon (Ar), and krypton (Kr); inert, or noble, gases,
which are virtually unreactive chemically.
Valence electrons
unfilled valence shells assume stable electron
configurations by gaining or losing electrons to
form charged ions, or by sharing electrons with
other atoms.
This is the basis for some chemical reactions, and
also for atomic bonding in solids.
Valence electrons
Under special circumstances, the s and p orbitals
combine to form hybrid spn orbitals, where n
indicates the number of p orbitals involved, which
may have a value of 1, 2, or 3.
The IIIA, IVA, and VA group elements of the
periodic table often form these hybrids. The
driving force for the formation of hybrid orbitals is
a lower energy state for the valence electrons.
For carbon the sp3 hybrid is of primary
importance in organic and polymer chemistries.
s- and p-orbitals
‘Aufbau’ Principle: filling orbitals
1s
2s
2p
H: 1s1
n=1
l=0
ml = 0
n=2
l=0
ml = 0
n=2
l=0
ml = -1 ml = 0
ml = 1
s- and p-orbitals
‘Aufbau’ Principle: filling orbitals
1s
2s
2p
He: 1s2
n=1
l=0
ml = 0
n=2
l=0
ml = 0
n=2
l=0
ml = -1 ml = 0
ml = 1
s- and p-orbitals
‘Aufbau’ Principle: filling orbitals
1s
2s
2p
Li: 1s2 2s1
n=1
l=0
ml = 0
n=2
l=0
ml = 0
n=2
l=0
ml = -1 ml = 0
ml = 1
s- and p-orbitals
‘Aufbau’ Principle: filling orbitals
1s
2s
2p
Be: 1s2 2s2
n=1
l=0
ml = 0
n=2
l=0
ml = 0
n=2
l=0
ml = -1 ml = 0
ml = 1
s- and p-orbitals
‘Aufbau’ Principle: filling orbitals
1s
2s
2p
B: 1s2 2s22p1
‘core’
closed
shell
open shell: valence
electrons
s- and p-orbitals
‘Aufbau’ Principle: filling orbitals
1s
2s
2p
C: 1s2 2s22p2
Hund’s rule: maximum number of unpaired
electrons is the lowest energy arrangement.
Hund’s rule
electrons fill orbitals one at a time.
we must fill each shell with one electron each
before starting to pair them up.
the charge of an electron is negative and electrons
repel each other. An electron will try to create
distance between itself and other electrons by
staying unpaired. This further explains why the
spins of electrons in an orbital are opposite (i.e.
+1/2 and -1/2).
s- and p-orbitals
‘Aufbau’ Principle: filling orbitals
1s
N: 1s2 2s22p3
O: 1s2 2s22p4
2s
2p
s- and p-orbitals
‘Aufbau’ Principle: filling orbitals
1s
F: 1s2 2s22p5
Ne: 1s2 2s22p6
2s
2p
s- and p-orbitals
‘Aufbau’ Principle: filling orbitals
Na: 1s22s22p63s1
Mg:
1s22s22p63s2
or [Ne]3s1
or
[Ne]3s2
P: [Ne]3s23p3
Ar: [Ne]3s23p6
electron configuration?
Which one of the following is a proper orbital
configuration?
electron configuration?
Which one of the following is a proper orbital
configuration?
electron configuration?
Which one of the following is
a proper orbital configuration?
beyond the d-orbitals
‘s’-groups
group
‘p’-groups
period
d-transition elements
1s2
2s2/2p6
3s2/3p6/
4s2/3d10/4p6
5s2/4d10/5p6/
6s2/4f14/5d10/6p6
lanthanides
actinides
f-transition elements
Organisation of the periodic table
Organisation of the periodic table
Organisation of the periodic table
Organisation of the periodic table
Organisation of the periodic table
electron configuration?
Give electron configurations for the Fe3+and
S2- ions.
The Fe3+ ion is an iron atom that has lost three
electrons. Since the electron configuration of the
Fe atom is 1s22s22p63s23p63d64s2,
the configuration for Fe3+ is 1s22s22p63s23p63d5.
The S2- ion a sulfur atom that has gained two
electrons. Since the electron configuration of the
S atom is 1s22s22p63s23p4,
the configuration for S2- is 1s22s22p63s23p6.
electron configuration?
Give the e- configurations for the following ions?
Fe2+
Al3+
Cu+
Ba2+
BrO2-






[Ar] 3d64s2
[Ne] 3s23p1
[Ar] 3d104s1
[Xe] 6s2
[Ar] 3d104s24p5
[He] 2s22p4
-2 e-: [Ar] 3d6
-3 e-: [Ne]
-1 e-: [Ar] 3d10
-2 e-: [Xe]
+1 e-: [Ar] 3d104s24p6
+2 e-: [He] 2s22p6
electron configuration?
Which of the following electron configurations is
an inert gas, a halogen, an alkali metal, an
alkaline earth metal, a transition metal?
a)
b)
c)
d)
e)
f)
1s22s22p63s23p63d74s2
1s22s22p63s23p6
1s22s22p5
1s22s22p63s2
1s22s22p63s23p63d24s2
1s22s22p63s23p64s1
transition metal
inert gas
halogen
alkaline earth metal
transition metal
alkali metal
electron configuration?
The halogens (group 7A, or group 17, of the periodic
table) all have similar chemical properties (for example,
forming singly charged negative ions). What aspect of
their electron configurations leads to these elements
having such similarities?
a) They all have a complete 1s2 shell at the lowest
energy level
b) They all have an identical Ns2Np5 configuration for
their valence electrons (N is any whole number).
c) They all have p electrons in their outermost shell.
d) They all have an odd number of protons.
e) They all have an even number of neutrons.