AP Chemistry Chapter 6
Electronic Structure and the
Periodic Table
Nature of Light
 Wavelengths
and frequencies
 Wavelength – distance between “troughs”


Measured in meters or nanometer
1nm = 10-9m
– number of of wave cycles
that pass a given point in unit time
 Frequency


Hertz – represents cycles per second
Ex. If 1010 cycles pass a particular point in
one second v = 1010 Hz
 λv
=c
 c = speed of light in vacuum
 2.998 x 108 m/s
 λ expressed in meters
 v in hertz
Photon energies
E
= hc
λ
Use to determine energy in joules of a
photon emitted by an excited atom
Also use to determine energy, in joules, of a
mole of photons
multiple by 6.02 x 1023
 Remember
to convert to kilojoules if
necessary!!!
 103J
= 1 kJ
Bohr model
 Based
on the hydrogen atom
 Why the electrons that were circling the
nucleus did not release their energy and
spiral into the nucleus
 He calculated energies associated with
each allowed orbit
 En
= -2.178 x 10-18 J
n2
En = energy of the electron
n = principal energy level
This formula is included on AP constants
sheet
Quantum Mechanical Model of
electron placement
First
quantum number, n
designates Principal
Quantum Level
 must be an integer
 Important for determining
energy of the electron

l = Second Quantum Number
 Determines
shape of electron cloud
 l = 0, 1,2,3……(n-1)
=1
n = 2
n= 3
n = 4
n
l=0
l = 0, 1
l = 0,1,2
l = 0,1,2,3
ml = Third Quantum Number
 Determines
electrons orientation in space
Corresponds to number of orbitals allowed in
that sublevel (l)
l = 0 ml = 0
(1 orbital)
l = 1 m l = 1,0,-1
(3 orbitals)
l = 2 m l = 2,1,0,-1,-2
(5 orbitals)
l = 3 m l = 3,2,1,0,-1,-2,-3
(7 orbitals)

Ms = Fourth Quantum Number;
electron spin
 Electron
has magnetic properties like that
of charged particles spinning on an axis
 Either of two spins is possible – clockwise
or counterclockwise
 +1/2 -1/2
 Electrons
with different ms values
(one +1/2 and the other -1/2)

Said to have “opposed” spins
 Electrons
with same value for ms
(both +1/2 or -1/2)

Said to have “parallel” spins
Pauli Exclusion Principle
 No
two electrons in an atom can
have the same set of four
quantum numbers
 If they occupy the same orbital,
must have opposing spins
 Pg. 140 example 6.4, 6.5
 Review
shape of sublevels
s - sphere, p - figure-8
 Pg.
412 shape of d sublevel orbitals
Hund’s Rule
When
several orbitals of equal
energy are available, as in a
given sublevel, electrons enter
singly with parallel spins
– is possible to determine number
of unpaired electrons in an atom by their
behavior in a magnetic field
 If unpaired electrons are present, the solid
will be attracted into the field
 Solids

That substance called
“paramagnetic”
 If
the atoms in the solid contain only
paired electrons, is slightly repelled
by the field

Called “diamagnetic”
Review electron configurations and
orbital notation
 Aufbau


Principle
Sublevels are filled in order of
increasing energy
Stability exceptions: Cr, Mo, W, Cu, Ag,
Au
• break from strict Aufbau Principle
 Review
“blocks” on Periodic Table
 Lanthanides – filling 4f
 Actinides – filling 5f (all of these elements
are radioactive, only thorium and uranium are
found in nature

Stability decreases with increasing atomic number
Monatomic ions
 Electrons
are added to or removed
from sublevels in the highest principal
energy level
 Want to achieve configuration like a
noble gas – more stable
 Na+1 (1s22s22p6) + e-
Species
(whether ion or not)
with same electron
configuration called
“isoelectronic”
Transition Metal Cations
 When
transition metals from positive ions,
the outer s electrons are lost first
 25Mn
[Ar] 4s23d5
 Mn+2
[Ar]3d5
to
 After
the outer s electrons are lost, then
the d can be lost
 26Fe
[Ar]4s23d6
forms Fe+3
 [Ar]3d5
 “first
in, first out” rule
Periodic Trends
Radius – one half the distance
of closest approach between atoms in
an elemental substance (pg. 151
drawing)
 Decreases across a period from left
to right
 Increases down a group as atomic
number increases
 Atomic
 In
a group, are increasing one whole
pel as you go down
 The inner electrons “shield” the
outer electrons from the positive
nucleus
 In periods, inner electrons are a poor
“shield” because they are at about the
same distance from the nucleus
Period trend continued
 Effective
nuclear charge (charge felt
by an outer electron) increases
steadily with atomic number
 As effective nuclear charge
increases, the outermost electrons
are pulled in more tightly, and atomic
radius decreases
Ionic Radius
 Increases
moving down a group
 Both cations and anions decrease
from left to right across a period
 Positive ions smaller that their atoms
 Negative ions larger than their atoms
 Pg. 152 Figure 6.13
 Cation – excess of protons draws the outer
electrons closer
 Anion – extra electron adds to the repulsion
between outer electrons (makes the negative ion
larger that the corresponding atom)

Pg. 153 example 6.10
Ionization Energy
 Measure
of how difficult it is to remove an
electron from a gaseous atom
 Energy must always be absorbed to
remove an electron, so always a positive
quantity
 First ionization energy – removal of
outermost electron
 X(g)  X+ + eΔE1 = first ionization energy
Trends
 Increases
across a period (left to right)
 Decreases down a group (increasing a.n.)
 Indirect relationship between atomic radius
and ionization energy


Large atom, electron far from the nucleus,
easier to remove
Smaller atom, electrons closer to the nucleus,
held tighter, so harder to remove
 Pg.
153 Figure 6.15
 First ionization energies in kJ/mol
 Pg.
154 Example 6.11
Electronegativity
 Measure
the ability of an atom in a
molecule to attract electrons to itself
 The greater the electronegativity, the
greater is its ability to attract electrons to
itself
 Dependent on ionization energy and
electron affinity
affinity – tendency to form anions
 EA = energy required to add an electron
 Z + e- + energy  Z EA = energy released on removing e- from
anion
 Z-  Z + e- + energy
 Electron
 If
EA is large and negative, the atom
“wants” to add an electron and form an
anion
 Atom with a very negative electron affinity
and a high ionization energy will attract eand resist any e- being removed from it.
 Is highly electronegative
EN Trends
– generally a steady increase
(metal to nonmetal)
 Period
– decrease within a group (are
some exceptions)
 Scale of electronegativities pg. 154
 Important scale when we get to bonding!
 Group