Electron Configurations & the Periodic Table

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Chapter 7
Electron Configurations & the
Periodic Table
General Chemistry I
S. Imbriglio
Part A: Electron Configurations
• The arrangement of all of the electrons in an atom
is called the electron configuration
• Electron configurations can be used to explain:
– Reactivity & properties of the elements
– Trends in reactivity & properties (periodic table!)
• The electron configuration of an atom is best
investigated using electromagnetic radiation
A. Electromagnetic (EM) Radiation
1. Electromagnetic (EM) Waves: oscillating
perpendicular magnetic & electric fields that travel
through space at the same rate (the “speed of
light”: c = 3.00x108 m/s)
- Unlike sound waves, electromagnetic waves
require no medium for propagation
eg. This allows the sun’s electromagnetic
radiation to reach the earth as sunlight.
http://ww2.unime.it/dipart/i_fismed/wbt/mirror/ntnujava/emWave/emWave.html
a) Wavelength & Frequency
• All EM waves can be described in terms of
wavelength & frequency.
– Wavelength
( - lambda):
distance between
adjacent crests
(or troughs) in a
wave
a) Wavelength & Frequency
- Frequency ( - nu): the number of complete
waves passing a point in a given period of time
(remember c = 3.00x108 m/s)
Unit of frequency is
the Hertz (Hz):
1 Hz = 1 s-1
= 1 “per second”
a) Wavelength & Frequency
• For EM radiation, frequency is related to wavelength
by  = c.
• If you know one, you know the other.
Calculate the frequency of an X-ray that has a wavelength of 8.21 nm.
 = c
i. Electromagnetic Spectrum
• The type of electromagnetic radiation is defined by its
frequency & wavelength
• Remember, as  increases,  decreases, and vice versa.
b) Amplitude
• The intensity of radiation is related to its
amplitude.
– Amplitude: height
of the wave crest
– In the visible portion
of the spectrum,
brighter light is light
with a greater
amplitude.
c) Refraction
• Classifying EM radiation (light) as a
wave explains many fundamental
properties of light.
• Refraction:
– When white light passes through a
narrow slit & then through a glass prism,
the light separates in a continuous
spectrum.
– The spectrum is continuous because
each color merges into the next without
a break – all wavelengths (or
frequencies) of visible light are observed.
d) Diffraction
• Diffraction:
Waves can add
constructively
or destructively
to amplify or
cancel each
other.
e) Black-Body Radiation
• At high temperatures, matter emits
electromagnetic radiation.
• As the temperature increases, the maximum
intensity of the emitted radiation increases in
frequency.
• The observed
spectrum depends
only on temperature
& not on the particular
elements present.
b) Black-Body Radiation
• THE EXPLANATION: According to classical
physics, as the temperature of a solid increases,
the atoms vibrate more vigorously – some of the
vibrational energy is released as EM radiation.
• THE PROBLEM: Using
the classical picture of
light as a wave,
scientists were unable
to explain the shape of
the observed black-body
radiation spectra.
3. Planck’s Quantum Hypothesis…
THE ANSWER?
• According to classical physics, the energy scale
is continuous – there are no limitations on the
amount of energy a system can gain or lose.
• Planck proposed that variations in energy are
discontinuous – energy changes occur only by
discrete amounts.
eg. The
“Quantization”
of Elevation
Quantized
(1 step = 1 quantum)
Classical
(continuous)
3. Planck’s Quantum Theory
• For electromagnetic radiation of a certain
frequency, the smallest amount of energy, called
a quantum, is defined by the relationship:
E = h
(h = Planck’s constant = 6.626x10-34 Js)
Energy can be absorbed or emitted only as a
quantum, or some whole-number multiple of a
quantum.
3. Planck’s Quantum Theory
• According to Planck’s theory, the energy of one
quantum of EM radiation is dependent on the
frequency (and wavelength) of the radiation:
E = h = hc/
• The energy per quantum increases as the
frequency gets higher & the wavelength gets
shorter.
3. Planck’s Quantum Theory
• The energy per quantum increases as the frequency
gets higher & the wavelength gets shorter.
E = h = hc/
• Which has more energy – a quantum of microwave radiation
( = 1x10-2 m) or a quantum of infrared radiation ( = 1x10-6
m)?
3. Planck’s Quantum Theory
• Planck proposed that vibrating atoms in a heated solid
can absorb and emit EM radiation only in certain discrete
amounts.
• Planck’s quantum theory allowed him to successfully
explain black-body radiation spectra, but his radical
assertion that “energy is quantized” was difficult for the
scientific community to accept.
• Fortunately, five years after its inception, Einstein used
Planck’s Quantum Theory to explain another well-known
phenomenon called the photoelectric effect.
4. The Photoelectric Effect
• Certain metals exhibit a photoelectric effect –
when illuminated by light of certain wavelengths
(photo-), they emit electrons (-electric).
• In order for the photoelectric effect to occur, the
frequency of the light must be higher than a certain
minimum value – called the threshold frequency.
• Each photosensitive metal has a different
threshold frequency.
4. The Photoelectric Effect
• When light of a high enough energy (frequency) is
used, the number of electrons ejected is
proportional to the intensity of the light.
• Light below the threshold frequency will not cause
an electric current to flow – no matter how bright
(intense) the light is.
eg. Light meters use the
photoelectric effect to
measure the intensity
(brightness) of light.
http://jchemed.chem.wisc.edu/JCEDLib/WebWare/collection/open/JCEWWOR006/peeffect5.html
a) Photons
• Classical physics could not explain the existence of
a threshold frequency, so Einstein turned to
Planck’s Quantum Theory.
• Einstein defined a quantum of electromagnetic
radiation as a photon.
• Einstein proposed that light could be thought of as
a stream of photons with particle-like properties as
well as wave properties.
• For light of frequency :
Ephoton = h = hc/
a) Photons
• The photoelectric effect can be explained by
assuming that light has particle-like properties:
– Removing one electron from a photosensitive
metal requires a certain minimum energy
(Emin).
– Each photon has an energy given by E = h.
– Only photons with E > Emin have enough
energy to knock an electron loose.
– Photons of lower frequency (lower energy)
do not have enough energy to knock an
electron loose.
a) Photons
• If the intensity of light is proportional to the
number of photons, then more intense light
means more photons.
• If each photon ejects an electron, then more
photons means more electrons ejected.
• The number of electrons ejected is
proportional to the intensity of light.
The Photoelectric Effect…Explained
b) Wave-Particle Duality of Light
• Depending on the circumstances, light (all EM
radiation) can appear to have either wave-like or
particle-like characteristics.
• Both ideas are needed to fully explain light’s
behavior in different phenomena.
“It not only prohibits the killing of two birds with one
stone, but also the killing of one bird with two
stones.”
- James Jeans, commenting on Einstein’s explanation
of the photoelectric effect
Nobel Prize Winners
• Max Planck won the Nobel Prize for Physics
in 1918 for his quantum theory.
Blackbody radiation spectra explained
• Albert Einstein won the Nobel Prize for
Physics in 1921 for his theory on the quantized
nature of light and how it relates to light’s
interaction with matter (not for his theory of
relativity!).
Photoelectric effect explained
5. Line Emission Spectra
• In the 1920s, another phenomenon was left
unexplained by classical physics – the observance
of atomic line emission spectra.
• When a voltage is applied to a gaseous element at
low pressure, the atoms absorb energy & become
“excited.”
• The “excited” atoms then emit the extra energy as
EM radiation.
5. Line Emission Spectra
• When this radiation is passed through a prism, a
limited number of discrete colored lines are seen –
a discontinuous spectrum.
• This discontinuous spectrum is called a line
spectrum, or a line emission spectrum.
• Unlike black-body radiation, each element has a
unique line emission spectrum.
Why don’t
these atoms
emit
continuous
spectra?
B. Bohr’s Hydrogen Atom: A
Planetary Model
• Classical physics could not explain the presence
of line emission spectra.
• Not long after Einstein used quantum theory to
explain the photoelectric effect, Niels Bohr used
quantum theory to explain the behavior of the
electron in a hydrogen atom.
• Bohr’s model provided the first explanation of
the discontinuous line emission spectrum of
hydrogen.
B. Bohr’s Hydrogen Atom: A
Planetary Model
• Bohr assumed that the single electron in a
hydrogen atom moves around the nucleus in a
circular orbit.
• Bohr applied quantum theory to his model by
proposing that the electron is restricted to circling
the nucleus in orbits of certain radii, each of which
corresponds to a specific energy.
• Thus, the energy of the electron is quantized,
and the electron is restricted to certain energy
levels – orbits.
B. Bohr’s Hydrogen Atom: A
Planetary Model
B. Bohr’s Hydrogen Atom
1. Energy Levels (Orbits):
– Each allowed orbit is assigned a principal
quantum number (n = 1,2,3,…).
– The energy of the electron and the radius
of its orbit increase as the value of n
increases.
– An atom with its electron in the lowest energy
level is said to be in the ground state.
1. Energy Levels (Orbits)
En = _
2.179x10-18 J
n2
(n = 1, 2, 3, …)
• The allowed energies of an electron (orbit) in a
hydrogen atom are restricted by the principal
quantum number (n), according to the equation
above.
• The negative sign is a result of Bohr’s choice to
define En = 0 when n = .
a) Excited State vs. Ground State
• Transitions Between Levels: Electrons can move
from one energy level to another
– An electron must absorb energy to transition from a lower
energy level to a higher energy level
– Energy is emitted when an electron transitions from a
higher energy level to a lower energy level
• When an electron absorbs energy and moves to a
higher energy level, that atom is said to be in an
excited state.
http://www.upscale.utoronto.ca/GeneralInterest/Harrison/BohrModel/Flash/BohrModel.html
a) Excited State vs. Ground State
• Absorb energy
to move to a
higher energy
orbit.
• Emit energy to
move to a lower
energy orbit.
a) Excited State vs. Ground State
• When an “excited”
electron returns to the
ground state, energy
is emitted as a photon
with an energy
corresponding to the
difference in energy between the two levels.
• In the Bohr model, n= is the excited state in which enough
energy has been added to completely separate the electron
from the proton – Bohr arbitrarily assigned this state as
having E = 0 (hence the negative energy values).
2. Explanation of Line Spectra
• Bohr’s model of
the hydrogen
atom can be
used to explain
the line
emission
spectrum of
hydrogen:
E = Efinal - Einitial
|E| = h
2. Explanation of Line Spectra
• Using Bohr’s equation for allowed energies in a hydrogen
atom:
E = Ef - Ei
E = _
2.179x10-18 J
nf2
E = 2.179x10-18 J x
_ _ 2.179x10-18 J
ni2
1 _ 1
ni2
nf2
Only certain energies of light (E ) can be absorbed or emitted
by electrons in a hydrogen atom.
2. Explanation of Line Spectra
• Now, coupling that equation with E = h allows us
to describe the frequencies of light that can be
absorbed or emitted by an electron in a hydrogen
atom.
E = h where E = |E|
• The frequencies () determined by this equation
correlate with the frequencies of light observed
in the line emission spectrum of hydrogen.
• The discrete lines
in the line emission
spectra correspond
to photons of
specific frequencies
that are emitted
when electrons
relax from higher
energy levels to
lower energy levels
Using Bohr’s model, calculate the frequency of the
radiation released by the transition of an electron in a
hydrogen atom from the n = 5 level to the n = 3 level.
Using Bohr’s model, calculate the wavelength of the
radiation absorbed by a hydrogen atom when the electron
undergoes a transition from the n = 4 to n = 5 level.
C. Quantum Mechanical Model of the Atom
• By the early 1920s, the theory of the WaveParticle Duality of light had been accepted, but a
young scientist named Louis De Broglie was ready
to shock the scientific community with another
hypothesis.
• De Broglie proposed that matter can exhibit
wave-like properties.
eg. Electrons exhibit diffraction
similar to that observed with light.
1. De Broglie: Matter as Waves
• De Broglie proposed that a particle of mass m
moving at speed v will have a wave nature
consistent with a wavelength given by the
equation:
 = h/mv
• Large (macroscale) objects have wavelengths too
short to observe.
• Small (nanoscale) objects have longer & more
readily observable wavelengths.
a) Quantum Mechanics
• Current ideas about atomic structure are based
on De Broglie’s theory.
• The treatment of atomic structure using the
wave-like properties of the electron is called
quantum mechanics (or wave mechanics)
• In contrast to Bohr’s precise atomic orbits,
quantum mechanics provides a “less certain”
picture of the hydrogen atom.
b) Wave Equation & Wave Functions
• In 1926, Erwin Schrödinger used De Broglie’s
theory to develop an equation (Schrödinger’s
wave equation) describing the locations &
energies of the electron in a hydrogen atom.
• Acceptable solutions to Schrödinger’s wave
equation are called wave functions ().
• Unlike Bohr’s model, these wave functions do not
describe the exact location of an electron.
b) Wave Equation & Wave Functions
• The square of a wave function (2) gives the
probability of finding an electron in a particular
infinitesimally small volume of space in an atom.
• Because we are treating electrons as waves (not
particles) we cannot pinpoint the specific location
of an electron.
• Instead, mathematical solutions to the wave
functions give 3-dimensional shapes (orbitals)
within which electrons can usually be found.
b) Wave Equation & Wave Functions
• These 3-D orbitals (probability clouds) take the place of
Bohr’s simple well-defined orbits in the modern model of the
atom. We don’t know exactly where the electrons are.
• This “less certain” model is justified by an important principle
of science established in 1927.
2. Heisenberg’s Uncertainty Principle
• It is impossible to determine the exact location
and the exact momentum of a tiny particle like
an electron.
– The very act of measurement would affect the
position and momentum of the electron because
of its very small size and mass.
– The collision of an electron with a high-energy
photon (required to locate the electron) would
change the momentum of the electron.
– The collision of an electron with a low-energy
photon would not provide much information
about the location of the electron.
2. Heisenberg’s Uncertainty Principle
• A macroscale analogy…
High Shutter Speed
Low Shutter Speed
Can judge location,
but not speed.
Can judge speed,
But not location
D. Quantum Numbers & Atomic
Orbitals
• According to quantum mechanics, each electron in
an atom can be described using four quantum
numbers:
–
–
–
–
n
l
ml
ms
Principal Quantum Number
Angular Momentum Quantum Number
Magnetic Quantum Number
Electron Spin Quantum Number
– The first three numbers describe the atomic orbital in
which the electron resides & the fourth differentiates
electrons that are in the same atomic orbital.
1. Principal Quantum Number (n)
• The principal quantum number (n) has only
integer values, starting with 1:
n = 1, 2, 3, 4, . . .
a) The value of n corresponds to the Principal
Electron Shell that the orbital is in.
b) The principal electron shell is the major factor in
determining the energy of the electron(s) in that
orbital – a higher n value means a higher energy.
2. Angular Momentum Quantum
Number (l)
• The angular momentum quantum number (l )
is an integer that ranges from zero to a
maximum of n – 1:
l = 0, 1, 2, 3, . . . (n – 1)
a) The value of l indicates the subshell that the
orbital is in (within the larger energy shell).
n = 1;
l=0
(1 subshell)
n = 2;
l = 0 or 1
(2 subshells)
n = 3;
l = 0, 1 or 2
(3 subshells)
n = 4;
l = 0, 1, 2 or 3 (4 subshells)
2. Angular Momentum Quantum
Number (l)
• Each subshell (l) is designated with a letter:
b) Each letter (s, p, d, f) symbolizes a subshell containing one
specific type of orbital with a unique shape.
eg. All s orbitals are spherical (l = 0) & all p orbitals are
shaped like dumbbells (l = 1) – more on this in a minute.
s orbital
p orbital
2. Angular Momentum Quantum Number
In the third principle shell, there is one s subshell containing one s
orbital, one p subshell containing three p orbitals & one d subshell
containing five d orbitals..
In the second principle shell, there is one
s subshell containing one s orbital & one
p subshell containing three p orbitals.
In the first principle shell, there is one
s subshell containing one s orbital.
Within a p or d subshell, how do you distinguish between the individual orbitals?
3. Magnetic Quantum Number (ml )
•
The magnetic quantum number (ml) can have
any integer value between l and - l, including
zero:
ml = l, . . . , +1, 0, -1, . . . , - l
a) The magnetic quantum number (ml) is related to
the directional orientation of the orbital.
eg. There are three possible p orbitals – each
pointing along a different axis in space.
3. Magnetic Quantum Number (ml )
(s) l = o;
(p) l = 1;
(d) l = 2;
(f) l = 3;
ml
ml
ml
ml
=0
= -1,0,1
= -2,-1,0,1,2
= -3,-2,-1,0,1,2,3
eg. There is only one
type of directional
orientation for any
given s orbital in an
l = 0 subshell because
ml must equal 0.
(1 s orbital)
(3 p orbitals)
(5 d orbitals)
(7 f orbitals)
3. Magnetic Quantum Number (ml )
• There are three different p orbitals in every l =
1 subshell because ml = -1,0,1.
Each of the three
p orbitals is
pointed along a
different axis
(x,y,z).
3. Magnetic Quantum Number (ml )
• There are five different d orbitals in every l = 2
subshell because ml = -2,-1,0,1,2.
Four of the five
d orbitals are
pointed along a
different axis.
The fifth has a
slightly different
shape.
4. Shells (n), Subshells (l ) & Orbitals (ml ):
A Summary
4. Shells (n), Subshells (l ) & Orbitals
(ml ): A Summary
This picture shows all of
the orbitals in the first
three electron shells
(n = 1,2,3).
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
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 = ?
Provide the three quantum numbers describing
each of the three p orbitals in the 2p subshell.
n
2px
2py
2pz
l
ml
5. Electron Spin Quantum Number (ms)
• The first three quantum numbers (n, l, ml) fully
characterize all of the orbitals in an atom.
• But, one more quantum number is necessary
to describe all of the electrons in an atom.
• This is because every orbital can hold two
electrons.
5. Electron Spin Quantum Number (ms)
• The spin quantum number (ms) can have just
one of two values (+1/2 & -1/2).
• Each electron exists in one of two possible spin
states.
- The “spinning”
electron induces an
external magnetic field.
Opposite spins induce
opposing magnetic
fields.
5. Electron Spin Quantum Number (ms)
• When two electrons have the same ms quantum
number, those spins are said to be parallel.
• When two electrons in the same orbital have
different ms quantum numbers, those electrons
are said to be paired.
 

Parallel spins
Paired spins
a) Pauli Exclusion Principle
• The Pauli Exclusion Principle states that no
more than two electrons can be assigned to the
same orbital in an atom & those two electrons
must have opposite spins.
• In other words:
– No two electrons in the same atom can have
the same set of four quantum numbers (n, l,
ml, ms).
– If two electrons occupy the same orbital, their
spins must be paired (+1/2 & -1/2).
Quantum Numbers: A
Macroscale Analogy
• n
•
l
- indicates which train (shell)
- 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).
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
How many orbitals in shell n = 1?
How many electrons possible?
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.
n=2
How many orbitals in shell n = 2?
How many electrons possible?
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.
n=3
# of Orbitals?
# of Electrons?
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.
Provide the four quantum numbers describing
each of the two electrons in the 3s orbital.
n
l
ml
ms
E. Electron Configurations
• The electron configuration of an atom is the
complete description of the orbitals occupied by
all of its electrons
– eg. The electron in a ground state hydrogen
atom occupies the 1s orbital
• There are several ways to represent electron
configurations. . .
1. Representations of Electron
Configuration
• In most cases, it is sufficient to write a list of all
of the occupied subshells and indicate the
number of electrons in each subshell with a
superscript.
H 1s1
C 1s2 2s2 2p2
Ar 1s2 2s2 2p6 3s2 3p6
1. Representations of Electron Configuration
a) Expanded Electron Configuration: In some
cases, it is more informative to write a list of each
occupied orbital and indicate the number of
electrons in each orbital.
N 1s2 2s2 2p3
versus
N 1s2 2s2 2p1 2p1 2p1
The expanded configuration indicates that there is one
electron in each of the three 2p orbitals – the
original configuration doesn’t.
1. Representations of Electron
Configuration
b) An orbital box diagram goes one step further by
also illustrating the spins of the elctrons.
P 1s2 2s2 2p2 2p2 2p2 3s2 3p1 3p1 3p1
P
1s
2s


2p


3s


3p



The orbital box diagram indicates that the three
electrons in the 3p subshell all have parallel
(unpaired) spins.
i) Hund’s Rule
• In the last example, we saw that:
– Atoms can have half-filled orbitals
– the electrons in the half-filled orbitals tend to have
parallel spins
• Hund’s Rule: The most stable arrangement of
electrons in the same subshell has the maximum
number of unpaired electrons, all with the same
spin
• In other words, electrons pair only after each
orbital in a subshell is occupied.
Write the expanded electron configuration and the
box orbital diagram for oxygen (1s2 2s2 2p4).
O
O
O
1s2 2s2 2p4
Write the expanded electron configuration and the
box orbital diagram for boron (1s2 2s2 2p1).
B
B
B
1s2 2s2 2p1
1. Representations of Electron
Configuration
c) When you get deeper into the periodic table,
electron configurations can be abbreviated by
using noble gas notation.
- The noble gases are the elements in group 8A
(He, Ne, Ar, Kr, Xe, Rn)
- Each noble gas has a filled outer subshell
(enough electrons to fill its highest energy
subshell)
c) Noble Gas Notation
• Electron Configurations of Noble Gases
[He]
[Ne]
[Ar]
[Kr]
= 1s2
= 1s2 2s2 2p6
= 1s2 2s2 2p6 3s2 3p6
= 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6
• To use noble gas notation, write the symbol for
the preceding noble gas [in brackets] to
represent all of the electrons in its electron
configuration.
• Add the rest of the electrons at the end.
c) Noble Gas Notation
• Write the following electron configurations using
noble gas notation:
O
1s2 2s2 2p4
Si 1s2 2s2 2p6 3s2 3p2
Now we know how to write electron configurations.
How do we know what the ground state
electron configuration for an element is???
2. Ground State Configuration
• Afbau Principle: Every atom has an infinite
number of possible electron configurations.
• For an atom in its ground state, electrons
are found in the energy shells, subshells &
orbitals that produce the lowest energy for
the atom.
• Other configurations correspond to excited
states.
2. Ground State Configuration
• In other words, when deciding where to “put”
the electrons in the ground state, always start
filling the lowest energy orbitals first.
• In general:
– Orbital energy increases as n increases
– Within the same shell (n), orbital energy
increases as l increases (E: s<p<d<f)
a) Order of Subshell Filling
• The electron configurations of the first ten
elements illustrate this point.
a) Order of Subshell Filling
• In general,
subshells are filled
in order of
increasing n + l
value
• If two orbitals have
the same value for
n + l, fill the
subshell with
lowest n value
first
a) Order of Subshell Filling
i) Using the Periodic Table
• You don’t have to memorize the order of the
subshells, just use the periodic table!
• Start at H & move through the table in order until
the desired element is reached.
Notice: (n – 1)d orbitals are filled after ns and before np orbitals.
a) Order of Subshell Filling
i) Using the Periodic Table
• Write the electron configuration for Al.
Al
Al
Ne
Al
a) Order of Subshell Filling
i) Using the Periodic Table
• Write the electron configuration for As.
As
As
Ar
As
a) Order of Subshell Filling
i) Using the Periodic Table
Write the electron configuration for Sn.
Sn
Sn
Kr
Sn
ii) Transition Metals
• Remember, (n – 1)d orbitals are generally filled
after ns orbitals and before np orbitals.
• There are some exceptions:
– When it is possible to half-fill or fill the (n-1)d
shell, the ns subshell can be left half-filled
– This is an example of Hund’s Rule. The ns
and (n-1)d orbitals are very close in energy,
so the more parallel spins, the better.
ii) Transition Metals
4s
Sc [Ar]3d14s2
- 4s filled before 3d

3d

4s
Ti
[Ar]3d24s2

3d


4s
V
[Ar]3d34s2

3d



ii) Transition Metals
Cr
might expect…
4s
[Ar]3d44s2

3d




Physical properties indicate that this is not the electron
configuration. It is actually…
4s
Cr
[Ar]3d54s1

3d




Notice the 3d subshell is half-filled. This configuration
maximizes unpaired electrons - Hund’s Rule.

ii) Transition Metals
Having a filled subshell is also energetically
favorable, so copper has an unexpected
configuration…
4s
Cu
[Ar]3d104s1

3d
    
The energetic stability gained from having either a
filled or a half-filled subshell has an effect on the
reactivity of different elements.
iii) Magnetic Properties
• The electron configuration of an atom
determines its magnetic properties.
• In atoms (or ions) with completely filled shells,
all of the electron spins are paired, so their
individual magnetic fields effectively cancel
each other out.
• Such substances are called diamagnetic.
iii) Magnetic Properties
• Atoms (or ions) with unpaired
electrons (parallel spins) are
attracted to a magnetic field.
– More unpaired electrons, stronger
attraction.
• Such substances are called
paramagnetic.
eg. Metallic nickel is paramagnetic:
4s
Ni

3d
   

iii) Magnetic Properties
• Ferromagnetic
substances are permanent
magnets.
– Spins of electrons in a
cluster of atoms are
aligned in same
direction, regardless of
external magnetic field
– Metals in the iron, cobalt
& nickel groups exhibit
ferromagnetism
iv) Valence Electrons
• The atomic electron configuration of an element
determines the chemical reactivity of that element,
but it is not the total number of electrons that is
important.
• If that were the case, each element would have
unique reactivity & we would not observe periodicity
in atomic trends and reactivity.
• How do we explain the trends in the periodic table?
Valence Electrons!
iv) Valence Electrons
• When considering the principal electron shells
(n = 1,2,3,…), there are two types of electrons:
– Core Electrons: electrons in the filled
“inner” shell(s) of an atom
– Valence Electrons: electrons in the unfilled
“outer” shell of an atom
• All elements in the same group have similar
chemical properties because they have the
same number of valence electrons in their
outer shell!
iv) Valence Electrons
• For elements in the first three periods:
– The core electrons are those in the preceding
noble gas configuration.
– The additional electrons in the outer shell are the
valence electrons.
eg.
B
1s2 2s2 2p1
B
[He]2s2 2p1
Core: 1s2
(Shell with n = 1)
Valence: 2s2 2p1
(Shell with n = 2)
iv) Valence Electrons
Cl
Cl
Core:
Valence:
iv) Valence Electrons
• For elements in the fourth period and below in groups
3A – 7A, the filled d subshells are also part of the
core, even though they are not included in the noble
gas configuration.
Se
Se
Core:
Valence:
iv) Valence Electrons
• In each A group, the number of valence
electrons is equal to the group number.
1
8
2 # of valence electrons 3 4 5 6 7
p-block
s-block
v) Lewis Dot Symbols
• The number of valence electrons in an atom is
directly related to its reactivity.
• Gilbert Lewis came up with a way to represent an
element & its valence electrons.
One dot equals one valence electron.
3. Ion Electron Configurations
• The number of valence electrons determines the
type of cation (+) or anion (-) an atom will form.
• When s- and p-block elements form ions, electrons
are removed or added such that a noble gas
configuration is achieved.
• The ions are said to be isoelectronic with the
noble gas – more on this later.
• In general:
– Metals lose electrons; form cations
– Non-metals gain electron; form anions
a) Cations
Li
Li+
1s22s1
1s2 = [He]
(loses an electron)
- Group 1 metals form cations with +1 charge.
Mg 1s22s22p63s2 (loses two electrons)
Mg2+ 1s22s22p6 = [Ne]
- Group 2 metals form cations with +2 charges.
b) Anions
O 1s22s22p4
(gains two electrons)
O2- 1s22s22p6 = [Ne]
- Group 6 elements form anions with -2 charge.
F
F-
1s22s22p5
(gains one electron)
1s22s22p6 = [Ne]
- Group 7 elements (halogens) form anions with -1
charge.
Periodic Trends
• When Mendelev created the first periodic table, he
organized the elements based on similarities in
chemical properties & reactivity.
• Now, we can use the electron configurations of
the elements to explain the trends in the periodic
table.
– Atomic & Ionic Radii
– Ionization Energy
– Electron Affinity
F. Atomic Radii
• The atomic radius of an atom is defined as
one-half of the internuclear distance between
two of the same atoms in a simple diatomic
molecule.
- In this simplified
picture, we assume that
each atom is spherical
& the radius is the distance
from the center to the edge
of the sphere.
1. Trends: Atomic Radii
• The size (radius) of an atom is determined by
two main factors:
a) Principal Quantum Number (n): the larger the
principal quantum number (n), the larger the
orbitals
- As you move down a group in the periodic table,
the atomic radii of the atoms increase because n
increases.
1. Trends: Atomic Radii
b) Effective Nuclear Charge (Z*): the nuclear
positive charge experienced by outer-shell
electrons in a many-electron atom
– Outer-shell electrons are shielded from the full
nuclear positive charge (Z) by the inner-shell
electrons (electron-electron repulsion)
– The effective nuclear charge (Z*) felt by an
outer-shell electron is less than the actual
charge of the nucleus (Z).
1. Trends: Atomic Radii
b) Effective Nuclear Charge:
• As Z* increases, the outer electrons are pulled
closer to the nucleus & the atomic radius
decreases.
• Z* increases across a period in the periodic
table (additional attraction to nucleus stronger than
electron-electron repulsion/shielding).
• Atomic Radius decreases across a period in the
periodic table.
1. Trends: Atomic Radii
Increasing Atomic Radius
Main
Group
Elements
Decreasing Atomic Radius
G. Ionic Radii
Periodic trends in ionic radii parallel the trends in
atomic radii within the same group.
Ionic radii increase as you move down a group in the
periodic table (n increases).
G. Ionic Radii
1. Cations: The radius of a
cation is always smaller
than that of the atom from
which it is derived.
– Nuclear charge (Z)
remains the same.
– Electron-electron
repulsion (shielding)
decreases.
– Effective nuclear
charge (Z*) increases.
G. Ionic Radii
2. Anions: The radius of an
anion is always larger than
that of the atom from which it
is derived.
– Nuclear charge (Z)
remains the same.
– Electron-electron
repulsion (shielding)
increases.
– Effective nuclear charge
(Z*) decreases.
a) Isoelectronic Ions
• Atoms or ions with identical electron
configurations are said to be isoelectronic.
• In general:
– Anions in a given period are isoelectronic
with the noble gas in the same period.
– Cations in a given period are isoelectronic
with the noble gas in the preceding period.
a) Isoelectronic Ions
a) Isoelectronic Ions
• When comparing
isoelectronic ions, the
radii depend on the
number of protons in the
nucleus (Z)
• The more protons in the
nucleus, the smaller the
radius.
a) Isoelectronic Ions
Isoelectronic Ions
O2-
F-
Na+
Mg2+
Ionic Radius (pm)
140
133
102
66
Number of Protons
8
9
11
12
Number of
Electrons
10
10
10
10
Rank the following series in order of increasing
atomic or ionic radii. (1 = smallest, 3 = largest)
Series 1:
Al
O
P
Series 2:
Be+2
Mg+2
Sr+2
Series 3:
O-2
F-
Na+
H. Ionization Energies
• The first ionization energy of an atom is the
minimum energy required to remove the highest
energy (outermost) electron from the neutral atom
in the gas phase.
• The larger the I.E., the harder it is to remove the
electron.
eg.
The first ionization energy of lithium is illustrated
by the following equation
Li (g)  Li+ (g) + e1s22s1
1s2
E = 520 kJ/mol
1. Trends: Ionization Energies
• Ionization energies tend to decrease as you move
down a group in the periodic table.
– Size - it is easier to remove an electron that is
further from the nucleus.
• Ionization energies tend to increase as you move
across a period in the periodic table.
– Effective Nuclear Charge – As Z* increases, it
becomes harder to remove an electron.
1. Trends: Ionization Energies
1. Trends: Ionization Energies
• Ionization energies do not increase smoothly
across the periods in the periodic table.
In general:
- It is easier to
remove an
electron if it results
in the formation of
a filled or half-filled
subshell.
- It is harder to
remove an
electron from a
filled or half-filled
subshell.
1. Trends: Ionization Energies
eg. In contrast to periodic trends, the ionization
energy of nitrogen is higher than the ionization
energy of oxygen.
N
N+
O
O+
1. Trends: Ionization Energies
In contrast to periodic trends, the ionization energy of
beryllium is higher than that of boron. Why?
Be
Be+
B
B+
2. Subsequent Ionization Energies
• The first, second & third
ionization energies are the
energies associated with
removing the first, second
& third highest energy
electrons in an atom.
• Ionization energies
increase with each
successive electron
removed because Z*
increases (same number of
protons, less electron
repulsion).
2. Subsequent Ionization Energies
• Ionization energies can be used to explain why
Li forms Li+ cations and Be forms Be2+ cations.
I. Electron Affinities
• The electron affinity of an element is the energy
change resulting from an electron being added to
an atom to form a 1- anion.
– Electron affinity is the measure of the attraction an
atom has for an additional electron.
eg. The electron affinity of fluorine is illustrated by:
F (g) + e-  F- (g) E = EA = -328 kJ/mol
1s22s22p5
1s22s22p5
Fluorine readily accepts an electron to gain a stable
noble gas configuration (filled n = 2 shell).
I. Electron Affinities
• Most electron affinities are <0 (favorable).
• Electron affinities are generally  0 (unfavorable) for
atoms with filled subshells (Groups 2A & 8A).
Knowing the Trends…
• You should be able to explain the trends & use the
periodic table to predict relative magnitudes for the
following properties:
– Atomic Radius
– Ionic Radius
– Ionization Energy
• The trends in electron affinities are less regular, but
you should be able to explain differences in EAs
based on filled/unfilled subshells.
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