Quantum numbers and orbital energies
Each atom’s electron has a unique set of
quantum numbers to define it { n, l, ml, ms }
Energy level
n=1
Sublevel
1s (l = 0)
# of orbitals/sublevel
1 (ml has one value)
n=2
2s (l = 0)
2p (l = 1)
1 (ml has one value)
3 (ml has three values)
n=3
3s (l = 0)
3p (l = 1)
3d (l = 2)
1 (ml has one value)
3 (ml has three values)
5 (ml has five values)
n = principal
quantum
number
(energy)
l = azimuthal
quantum
number
(shape)
ml = magnetic
quantum
number
(orientation)
Concept: Each electron in an atom has a unique
set of quantum numbers to define it
{ n, l, ml, ms }
2
Fig. 7.14
What is the reason that the periodic table
organizes elements according to
similarities in chemical properties?
Arrangement of electrons in
atoms
•
There is a
relationship between
the quantum number
(n) and its the
number of subshells.
Principal quantum number (n) = number of subshells
multi-electron atoms
H atom
All other atoms
6
As n increases, the difference in energy level _______.
Energy levels for multi-electron
atoms
energy
# of
level
electrons
2n2 electrons
Orbital energy ladder
n = 3  18 e’s
n = 2  8 e’s
n = 1  2 e’s
Basic Principle:
electrons occupy
lowest energy
levels available
Aufbau Principle -- “Bottom Up
Rule”
Electron spin
How could an orbital hold two electrons
without electrostatic repulsion?

SternGerlach
Experiment

2 ways to write electron configurations
spdf NOTATION
for H, atomic number = 1
no. of
1
electrons
spdf Notation
1s
sublevel
value of energy level
Orbital Box Notation
ORBITAL BOX NOTATION
for He, atomic number = 2
2
1s

1s
Arrows show
electron spin
(+½ or -½)
Pauli exclusion principle
An orbital can contain a maximum of 2 electrons,
and they must have the opposite “spin.”
• Example:
•
Determine the electron configuration and
orbital notation for the ground state neon atom.
Write the ground state configuration
and the orbital diagram for oxygen
in its ground state
Hund’s Rule -
Basic Principle:
electrons occupy
lowest energy
levels available
Rules for Filling Orbitals
Bottom-up
(Aufbau’s principle)
Fill orbitals singly before doubling up
(Hund’s Rule)
Paired electrons have opposite spin
(Pauli exclusion principle)
Orbital energy ladder
f
d
n=4
p
d
p
Energy
n=3
s
p
s
s
n=1
n=2
s
Phosphorus




  
Symbol: P
Atomic Number: 15

Full Configuration: 1s22s22p63s23p3

Valence Configuration: 3s23p3
Shorthand Configuration: [Ne]3s23p3
Box Notation


1s
2s


2p


3s


3p

Electron spin &
magnetism
For the ground state oxygen atom:
spdf configuration:
orbital box notation:
Paramagnetic: atoms with unpaired electrons
that are weakly attracted to a magnet.
Diamagnetic: atoms with paired electrons
that are not attracted to a magnet.
Apparatus for measuring
magnetic properties
Identify examples of the following principles:
1) Aufbau 2) Hund’s rule 3) Pauli exclusion
Note: Not written according to Aufbau, but grouping according to n
Electron distribution for the argon atom
Never zero electron distribution
Electron configuration for As
Shorthand notation for silicon
1s22s22p6 3s23p2  [Ne] 3s23p2
[Ne]
Silicon’s valence electrons
Shorthand
practice
[Noble
Gas Core] notation
+ higher energy
electrons
• Examples
• ● Aluminum: 1s22s22p63s23p1
[Ne]3s23p1
• ● Calcium: 1s22s22p63s23p64s2
•
[Ar]4s2
• ● Nickel: 1s22s22p63s23p64s23d8
•
[Ar]4s23d8 {or [Ar]3d84s2}
• ● Iodine: [Kr]5s24d105p5 {or [Kr]4d105s25p5}
• ● Astatine (At): [Xe]6s24f145d106p5
Outer electron configuration for the
elements
Electronic configuration of Br
• 1s2 2s22p6 3s23p63d10 4s24p5
[Ar] 3d104s24p5
• [Ar] = “noble gas core”
• [Ar]3d10 = “pseudo noble gas core”
• (electrons that tend not to react)
Atom’s reactivity is determined by valence electrons
valence e’s in Br: 4s24p5
highest n electrons
Valence e’s for “main group” elements
Valence e- shells for
transition metals v. main group elements
d orbitals sometimes
included in valence shell
d orbitals not included
in valence shell
(pseudo noble gas cores)
Rule-of-thumb for valence electrons
Identify all electrons at the highest
principal quantum number (n)
•
•
•
•
•
•
•
•
•
•
Examples
●
●
●
●
Sulfur: 1s22s22p63s23p4 or [Ne]3s23p4
valence electrons: 3s23p4
Strontium: [Kr]5s2
valence electrons: 5s2
Gallium: [Ar]4s23d104p1
valence electrons: 4s24p1
Vanadium: [Ar]4s23d3
valence electrons: 4s2 or 3d34s2
Selenium’s valence electrons
Written for increasing energy:
Pseudo noble gas core includes:
 noble gas electron core
 d electrons (not very reactive)
Core and valence electrons in Germanium
Written for increasing energy:
Pseudo noble gas core includes:
 noble gas core
 d electrons
d-block: some exceptions to the Aufbau
principle
• Mendeleev’s periodic table generally
organized elements by increasing atomic
mass and with similar properties in
columns. In some places, there were
missing elements whose properties he
predicted.
• When gallium, scandium, and germanium
were isolated and characterized, their
properties were almost identical to those
predicted by Mendeleev for eka-aluminum,
eka-boron, and eka-silicon, respectively.
Figure 8.14: Mendeleev’s periodic table.
• Periodic law states that when the
elements are arranged by atomic number,
their physical and chemical properties vary
periodically.
• We will look in more detail at three
periodic properties: atomic radius,
ionization energy, and electron affinity.
Effective Nuclear Charge
Effective nuclear charge is the positive
charge that an electron experiences from
the nucleus. It is equal to the nuclear
charge, but is reduced by shielding or
screening from any intervening electron
distribution (inner shell electrons).
Effective nuclear charge increases across a
period. Because the shell number (n) is
the same across a period, each
successive atom experiences a stronger
nuclear charge. As a result, the atomic
size decreases across a period.
Atomic Radius
While an atom does not have a definite size,
we can define it in terms of covalent radii
(the radius in covalent compounds).
Figure 8.17:
Representation
of atomic radii
(covalent radii)
of the maingroup
elements.
• Atomic radius is plotted against atomic
number in the graph below. Note the regular
(periodic) variation.
A representation of atomic radii is shown below.
• Refer to a periodic table and arrange
the following elements in order of
increasing atomic radius: Br, Se, Te.
34
35
Se
Br
Te is larger than Se.
Se is larger than Br.
52
Te
Br < Se < Te
• First Ionization Energy (first ionization
potential)
• The minimum energy needed to remove
the highest-energy (outermost) electron
from a neutral atom in the gaseous state,
thereby forming a positive ion
Periodicity of First Ionization
Energy (IE1)
Like Figure 8-18
Fig. 8.15
• Left of the line, valence shell electrons are
being removed. Right of the line, noblegas core electrons are being removed.
Identifying Elements by Its Successive
Ionization Energies
Problem: Given the following series of ionization energies (in kJ/mol)
for an element in period 3, name the element and write its
electron configuration:
IE1
IE2
IE3
IE4
580
1,815
2,740
11,600
Plan: Examine the values to find the largest jump in ionization energy,
which occurs after all valence electrons have been removed.
Use the periodic table!
Solution:
Identifying Elements by Its Successive
Ionization Energies
Problem: Given the following series of ionization energies (in kJ/mol)
for an element in period 3, name the element and write its
electron configuration:
IE1
IE2
IE3
IE4
580
1,815
2,740
11,600
Plan: Examine the values to find the largest jump in ionization energy,
which occurs after all valence electrons have been removed.
Use the periodic table!
Solution:
The largest jump in IE occurs after IE3 so the element has
3 valence electrons thus it is Aluminum ( Al, Z=13), its
electron configuration is :
1s2 2s2 2p6 3s2 3p1
Fig. 8.16
Ranking Elements by First Ionization Energy
Problem: Using the Periodic table only, rank the following elements in
each of the following sets in order of increasing IE!
a) Ar, Ne, Rn
b) At, Bi, Po
c) Be, Na, Mg
d) Cl, K, Ar
Plan: Find their relative positions in the periodic table and apply trends!
Solution:
• Trends
• Going down a group, first ionization
energy decreases.
• This trend is explained by understanding
that the smaller an atom, the harder it is to
remove an electron, so the larger the
ionization energy.
• Generally, ionization energy increases with
atomic number.
• Ionization energy is proportional to the
effective nuclear charge divided by the
average distance between the electron
and the nucleus. Because the distance
between the electron and the nucleus is
inversely proportional to the effective
nuclear charge, ionization energy is
inversely proportional to the square of the
effective nuclear charge.
• Small deviations occur between Groups
IIA and IIIA and between Groups VA and
VIA.
• Examining the valence configurations for
these groups helps us to understand these
deviations:
• IIA ns2
• IIIA ns2np1
• VA ns2np3
• VIA ns2np4
It takes less energy to remove the
np1 electron than the ns2 electron.
It takes less energy to remove the
np4 electron than the np3 electron.
• Electrons can be successively removed
from an atom. Each successive ionization
energy increases, because the electron is
removed from a positive ion of increasing
charge.
• A dramatic increase occurs when the first
electron from the noble-gas core is
removed.
• Refer to a periodic table and arrange
the following elements in order of
increasing ionization energy: As, Br,
Sb.
33
35
As
Br
Sb is larger than As.
As is larger than Br.
51
Sb
Ionization energies:
Sb < As < Br
Overall periodic trends
Note: Electronegativity has similar trend as electron affinity
Reactivity of the Alkali Metals
2Li(s) + 2H2O(l)  2LiOH(aq) + H2(g)
Lithium video
2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g)
Sodium video
2K(s) + 2H2O(l)  2KOH(aq) + H2(g)
Trend?
Potassium video
61
More Sodium Reaction Videos
2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g)
http://www.theodoregray.com/PeriodicTable/
Preppin
g Na
100 g Na
in one
piece
150 g Na
in small
pieces
62
Electronic Configuration Ions
• Na 1s 2 2s 2 2p 6 3s 1
Na+
• Mg 1s 2 2s 2 2p 6 3s 2
Mg+2 1s 2 2s 2 2p6
• Al 1s 2 2s 2 2p 6 3s 2 3p 1
Al+3 1s 2 2s 2 2p 6
• O 1s 2 2s 2 2p 4
O- 2
1s 2 2s 2 2p 6
• F
1s 2 2s 2 2p 5
F- 1
1s 2 2s 2 2p 6
• N
1s 2 2s 2 2p 3
N- 3
1s 2 2s 2 2p 6
1s 2 2s 2 2p 6
Isoelectronic Atoms and Ions
•
H- 1 { He } Li+ Be+2
•
N- 3
•
P- 3 S- 2 Cl- { Ar }
•
As- 3 Se- 2 Br- { Kr } Rb+ Sr+2
•
Sb- 3 Te- 2
O- 2 F- { Ne } Na+ Mg+2 Al+3
K+ Ca+2
Sc+3 Ti+4
Y+3
Zr+4
I- { Xe } Cs+ Ba+2 La+3 Hf+4
Trends when atoms form chemical
bonds
Empirical Observation
“when forming ionic compounds, elements tend to lose
or gain electrons to be more like the nearest noble gas”
Metals tend to lose e-’s
Nonmetals tend to gain e-’s
Are ions bigger or smaller than atoms?
Representative cation
Na → Na+ + e-
Representative anion
F + e- → F-