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Electronic Structure of
Atom and Periodicity
1. Electromagnetic Radiation
2. The Nature of Photon
3. The Atomic Spectrum of Hydrogen
4. The Bohr Model
5. The Quantum Mechanical Model of the Atom
6. Quantum Numbers
7. Orbital Shapes and Energies
8. Electron Spin and the Pauli Principle
9. Polyelectronic Atoms
10. The History of the Periodic Table
11. The Aufbau Principle and the Periodic Table
12. Periodic Trends in Atomic Properties
13. The Properties of a Group: The Alkali Metals
2
Colors in Fireworks:
Red: Sr, Li
Orange: Fe/Charcoal
Yellow: Na
Blue: Cu
Green: Ba
Purple: Cu+Sr
Why do different chemicals give
us different colors?
3
Electromagnetic (EM) Radiation:
One of the ways that energy travels through space.
Characteristics of EM Waves
• Wavelength ( l ) – distance
between two peaks or troughs
in a wave.
• Speed (c) – speed of trough or
crest passing through the
space. speed of light
(2.9979×108 m/s)
4
Frequency (v) vs. Wavelength (l)
Frequency ( v ) – number of waves
(cycles) per second that pass a
given point in space.
• Unit: sec-1, or Hertz
• For EM wave, frequency is
inversely proportional to
Wavelength: c = v x l
Or v = c/l
5
Practice: What is the frequency of red light from common laser
pointer (wavelength at 650 nm)?
c = 2.9979×108 m/s
• 4.62 x 1014 Hz
6
Classification of Electromagnetic
Radiation
7
Types of Electromagnetic Radiation
• By Wavelength
– Radiowaves
– Microwaves
– Infrared (IR)
– Visible: 400 nm < l < 800nm
• ROYGBIV
– Ultraviolet (UV)
– X-rays
– Gamma rays
• high frequency and energy
8
Pickle Light:
https://www.youtube.com/watch?
v=2vNeMeniYWY
(Dangerous! don’t try
at home without Dr.
Wen :O )
Household 120V
through the pair of
gater clip
Yellow glow from
sodium in the pickled
cucumber
9
Marx Planck’s theory
• Energy is quantized.
• Electromagnetic radiation is a stream of “particles”
called photons. Energy of one photon:
Ephoton
h (Planck’s constant) = 6.625×10-34 J • s
hc
= hν =
l
• Dual nature of light:
 Electromagnetic radiation (and all matter) exhibits
wave properties and particulate properties.
10
Mystery in the Line Spectrum
• Continuous spectrum vs. Line spectrum:
11
Bohr’s Theory on Hydrogen Atom
• Electron in a hydrogen atom
moves around the nucleus only
in certain allowed circular orbits.
• Energy of the electron in the
hydrogen atom is quantized.
• Ground state – lowest possible
energy state (n = 1)
• Only certain energies are
allowed for the electron in the
18
2.178 10 J
hydrogen atom.
En  
2
n
12
Electronic Transitions in the Bohr Model
vs. Line Spectrum for Hydrogen Atom
13
The Bohr explanation of three series of spectral lines emitted by the H atom.
Energy for Electronic Transitions in
the Hydrogen Atom
• For a single electron transition from one energy level
to another:
E =  2.178  10
18
 1
1 
J 2  2 
ninitial 
 nfinal
ΔE = change in energy of the atom (energy of the emitted photon)
nfinal = integer; final energy level (orbital)
ninitial = integer; initial energy level (orbital)
For emission spectrum, the frequency of light can be
calculated through
E = hv
15
Exercise
What color of light is emitted when an excited
electron in the hydrogen atom falls from:
a) n = 5 to n = 2
b) n = 4 to n = 1
 1
1 
E =  2.178  1018 J  2  2 
ninitial 
 nfinal
E = hv
h = 6.625×10-34 J • s
blue, λ = 434 nm
uv, λ = 97.3 nm
16
Bohr’s Theory Explains the Line
Spectrum in Hydrogen Atom
• The model suggests only certain allowed circular orbits
for the electron.
• As the electron becomes more tightly bound (smaller
n), its energy becomes more negative relative to the
zero-energy reference state (free electron).
• As the electron is brought closer to the nucleus, aka, nf
< ni, energy is released from the system.
• When electron receives photon, it transits into higher
energy levels, ni < nf
17
Beyond Bohr’s Model
• Bohr’s model only works for hydrogen, also electrons
do not move around the nucleus in circular orbits.
• Heisenberg uncertainty principle:
 There is a fundamental limitation to just how
precisely we can know both the position and
momentum of a particle at a given time.
x    m 
h

4
Δx = uncertainty in a particle’s position
Δ(mν) = uncertainty in a particle’s momentum
h = Planck’s constant
18
Wave Behavior of Electrons
Louis de Broglie (1892–1987)
• de Broglie proposed that particles could have wavelike character
• de Broglie predicted that the wavelength of a particle
was inversely proportional to its momentum
• Because it is so small, the wave character of electrons
is significant
19
Tro: Chemistry: A Molecular Approach, 2/e
Electron Diffraction
• Proof that the electron had wave nature came a few
years later with the demonstration that a beam of
electrons would produce an interference pattern as
waves do
If electrons behave only
like particles, there
should only be two
bright spots on the
target
However, electrons
actually present an
interference pattern,
demonstrating they
behave like waves
20
Tro: Chemistry: A Molecular Approach, 2/e
Diffraction patterns of aluminum with x-rays (“true wave”) and electrons.
x-ray diffraction of aluminum foil
electron diffraction of aluminum foil
Electron microscope:
High speed electrons behaves like light, used for imaging
Physical Meaning of a Wave Function
• The square of the function (2) indicates the
probability of finding an electron near a particular
point in space.
 Probability distribution – intensity of color is used
to indicate the probability value near a given point
in space.
23
Relative Orbital Size
• Difficult to define precisely.
• Orbital is a wave function.
• Picture an orbital as a three-dimensional electron
density map.
• Hydrogen 1s orbital:
 Radius of the sphere that encloses 90% of the
total electron probability.
24
• Principal quantum number (n) – size and energy of
the orbital.
• Angular momentum quantum number (l) – shape of
atomic orbitals (sometimes called a subshell). For
each n, l = 0, 1, … n-1.
• Subshells we learned: l = 0 (s), 1 (p), 2 (d), 3 (f)
• Magnetic quantum number (ml) – orientation of the
orbital in space relative to the other orbitals in the
atom. For each l, ml = -l, -l+1, …, 0, 1,.., l
25
Quantum Numbers for the First Four Levels of Orbitals in the Hydrogen Atom
26
The Quantum-Mechanical Model of the Atom
The matter-wave of the electron occupies the
space near the nucleus and is continuously
influenced by it.
The Schrödinger wave equation () allows us
to solve for the energy states associated with
a particular atomic orbital.
The square of the wave function (2) gives
the probability density, a measure of the
probability of finding an electron of a
particular energy in a particular region of the
atom.
Sample Wave Functions of Hydrogen Atom
1s orbital: n = 1, l = 0
2s orbital: n = 2, l = 0
2p orbital: n = 2, l = -1, 0, +1
Electron probability density in the ground-state H atom.
Quantum Numbers (n, l, m ) and Atomic Orbitals
l
An atomic orbital is specified by three quantum numbers.
The principal quantum number (n) is a positive integer.
The value of n indicates the relative size of the orbital and therefore its relative distance from the
nucleus. Also known as “Shell number”
The angular momentum quantum number (l) is an integer from 0 to (n –1).
The value of l indicates the shape of the orbital. Aka, “Subshell number”.
The magnetic quantum number (m ) is an integer with values from –l to +l.
l
The value of m indicates the spatial orientation
of the orbital.
l
Table 7.2 The Hierarchy of Quantum Numbers for Atomic Orbitals
Name, Symbol
(Property)
Principal, n
(size, energy)
Angular momentum, l
(shape)
Magnetic, m
l
(orientation)
Allowed Values
Quantum Numbers
Positive integer
(1, 2, 3, ...)
1
0 to n – 1
0
0
0
0
-l,…,0,…,+l
2
3
1
0
1
2
0
-1
0
+1
-1
-2
0
-1
+1
0
+1
+2
Sample Problem 7.6
PROBLEM:
Determining Quantum Numbers for an Energy Level
What values of the angular momentum (l) and magnetic (m ) quantum numbers are
l orbitals are allowed
allowed for a principal quantum number (n) of 3? How many
for n = 3?
Values of l are determined from the value for n, since l can take values from 0 to (n – 1). The
values of m then follow from the values of l.
l
PLAN:
SOLUTION:
For n = 3, allowed values of l are = 0, 1, and 2
For l = 0, m = 0
l
For l = 1, m = –1, 0, or +1
l
For l = 2, m = –2, –1, 0, +1, or +2
l
There are 9 m values and therefore 9 orbitals with n = 3.
l
Practice: Determining Sublevel Names and
Orbital Quantum Numbers
Give the name, magnetic quantum numbers, and number of
orbitals for each sublevel with the following quantum numbers:
(a)
(b)
(c)
(d)
n = 3, l = 2
n = 2, l = 0
n = 5, l = 1
n = 4, l = 3
Exercise
For principal quantum level n = 3, determine the
number of allowed subshells (different values
of l), and give the designation (such as 3s, etc.)
of each.
34
Exercise
For l = 2, determine the magnetic quantum
numbers (ml) and the number of orbitals.
Note: Number of available orbitals depends on
the number of magnetic quantum numbers.
35
Exercise
A.
B.
C.
D.
Which of the following sets of quantum
numbers is for an electron in 5f subshell?
n = 5, l = 3, ml = 3, ms = -1/2
n = 5, l = 3, ml = -4, ms = +1/2
n = 5, l = 3, ml = -2, ms = +1/2
n = 5, l = 2, ml = -3, ms = -1/2
36
The 1s, 2s, and 3s orbitals
Energy levels of the H atom.
The Boundary Surface Representations of All Three 2p
Orbitals
(n = 2, l = 1)
39
The Boundary Surfaces of All of the 3d
Orbitals
(n =
,l=
)
40
Representation of the 4f Orbitals in Terms of Their Boundary Surfaces
(n = , l =
)
41
Can you identify each orbitals?
2s, 3s
2pz, 3pz, 3px
2p, 2dz2, 3d, 3d
4s, 4p, 4p, 4d, 4d
4d, 4f, 4f, 4f, 4f
42
Electron Spin
• Electron spin quantum number (ms) – can be +½ or ½.
• Pauli exclusion principle - in a given atom no two
electrons can have the same set of four quantum
numbers.
• An orbital can hold only two electrons, and they
must have opposite spins.
43
• Atoms with more than one electron.
• Electron correlation problem:
 Since the electron pathways are unknown, the
electron repulsions cannot be calculated exactly.
• When electrons are placed in a particular quantum
level, they “prefer” the orbitals in the order s, p, d,
and then f.
44
Penetration Effect
• A 2s electron penetrates to the nucleus more than
one in the 2p orbital.
• This causes an electron in a 2s orbital to be attracted
to the nucleus more strongly than an electron in a 2p
orbital.
• Thus, the 2s orbital is lower in energy than the 2p
orbitals in a polyelectronic atom.
45
A Comparison of the Radial Probability Distributions of the 2s and 2p Orbitals
46
Penetration Effect for Electron in a 3s
Orbital: Opportunity of electron
getting “access” to nucleus (low
energy state)
47
A Comparison of the Radial Probability
Distributions of the 3s, 3p, and 3d
Orbitals:
Penetration Effect leads to the order of
energy among the subshells, s < p < d <
f
48
• Originally constructed to represent the patterns
observed in the chemical properties of the elements.
• Mendeleev is given the most credit for the current
version of the periodic table. The year 1876 version:
49
Aufbau Principle
• As protons are added one by one to the nucleus to
build up the elements, electrons are similarly added
to hydrogen–like orbitals.
• An oxygen atom as an electron arrangement of two
electrons in the 1s subshell, two electrons in the 2s
subshell, and four electrons in the 2p subshell.
Oxygen: 1s22s22p4
50
Hund’s Rule
• The lowest energy configuration for an atom is the
one having the maximum number of unpaired
electrons allowed by the Pauli principle in a
particular set of degenerate (same energy) orbitals.
51
Orbital Diagram
• A notation that shows how many electrons an atom
has in each of its occupied electron orbitals.
Oxygen: 1s22s22p4
Oxygen:
1s
2s
2p
52
Valence Electrons
• The electrons in the outermost principal quantum
level of an atom.
1s22s22p6 (valence electrons = 8)
1s22s22p63s23p64s23d104p2 (valence electrons= 4)
• The elements in the same group on the periodic table
have the same valence electron configuration.
53
The Orbitals Being Filled for Elements in Various Parts of the Periodic Table
54
Exercise
Determine the expected electron
configurations for each of the following.
a) S
1s22s22p63s23p4 or [Ne]3s23p4
b) Ba
[Xe]6s2
c) Eu
[Xe]6s24f7
55
Order of Subshell Filling
in Ground State Electron
Configurations
1. Diagram
putting each energy shell on
a row and listing the subshells,
(s, p, d, f), for that shell in
order of energy, (left-to-right)
2. draw arrows through
the diagonals, looping back
to the next diagonal
each time
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d
7s
56
Effective Nuclear Charge Zeff
Electron configuration determines chemical properties:
the Periodicity. Why?
• Core electrons form cloud of negative charge,
“shielding” the attraction between nuclear charge Z
and outer electrons.
• Different atoms have different electron
configuration, valence electrons “feel” differently
from nuclear attraction
• Effective Nuclear Charge Zeff: The actual positive
charge from nucleus attracting the outer electrons
(valence electrons).
57
Estimate of Effective Nuclear Charge Zeff
Zeff = full nuclear charge Z - #core electrons
Example:
• Mg atom: 1s22s22p63s2
Zeff  12 – 10 = 2
• Mg2+ ion: 1s22s22p6
Zeff  12 – 2 = 10
• O atom: 1s22s22p4
Zeff  8 – 2 = 6
58
Application of Effective Nuclear Charge
Higher Zeff leads to stronger attraction
towards the valence electrons, causing
• Higher Ionization Energy
• Higher Electron Affinity
• Smaller Atomic Radius
59
9000
8000
Ionization Energy
7000
6000
5000
Energy required to remove an electron from a
gaseous atom or ion.
 X(g) → X+(g) + e–
4000
3000
2000
1000
0
Mg → Mg+ + e– I1 = 735 kJ/mol (1st IE)
Mg+ → Mg2+ + e– I2 = 1445 kJ/mol (2nd IE)
Mg2+ → Mg3+ + e– I3 = 7730 kJ/mol *(3rd IE)
0
1
2
3
*Core electrons are bound much
more tightly than valence electrons.
60
4
The Values of First Ionization Energy for the Elements in the First Six Periods
61
Ionization Energy within Period
• Why In general, as we go across a period from left to
right, the first ionization energy increases?
 Electrons added in the same principal quantum
level do not completely shield the increasing
nuclear charge caused by the added protons. Zeff
consistently increases from left to right across the
period.
 Electrons in the same principal quantum level are
generally more strongly bound from left to right
on the periodic table.
62
Ionization Energy within Group
• In general, as we go down a group from top to
bottom, the first ionization energy decreases.
• Why?
 Zeff maintains at similar level.
 However, the outer electrons are at higher energy
level. The electrons being removed are, on
average, farther from the nucleus, thus easier to
be removed.
63
Successive Ionization Energies (KJ per Mole) for the Elements in Period 3
64
Electron Affinity
• Energy change associated with the addition of an
electron to a gaseous atom.
 X(g) + e– → X–(g)
• In general as we go across a period from left to right,
the electron affinities become more negative.
• In general electron affinity becomes more positive in
going down a group.
65
Figure 8.18
Electron affinities of the main-group elements (in kJ/mol).
Atomic Radius
• In general as we go across a period from left to right,
the atomic radius decreases.
 Effective nuclear charge increases, therefore the
valence electrons are drawn closer to the nucleus,
decreasing the size of the atom.
• In general atomic radius increases in going down a
group.
 Orbital sizes increase in successive principal
quantum levels.
67
Atomic Radii for Selected Atoms
68
Practice: Estimate Zeff for outmost eK, Na+, Al, S, Cl, F
Step 1. Write the electron configuration
Step 2. identify core electrons vs. outmost
electrons
Step 3. Zeff = Z – #core electrons
K = 1, Na = 1, Na+ = 9, Al = 3, S = 6, Cl = 7, F = 7
69
K = 1, Na = 1, Al = 3, S = 6, Cl = 7, F = 7
Application of Zeff
Rank the first ionization energy for atom of Na,
Al, K, S, Cl.
K < Na < Al < S < Cl < F
70
K = 1, Na = 1, Na+ = 9, Al = 3, Al+ = 3, S = 6, Cl = 7, F = 7
Application of Zeff
Rank the atomic radii for atom of Na, Al, K, S, Cl,
and F.
F < Cl < S < Al < Na < K
71
K = 1, Na = 1, Na+ = 9, Al = 3, Al+ = 3, S = 6, Cl = 7, F = 7
*Application of Zeff
Which has the larger second ionization energy?
Why?
Sodium or Aluminum
IE2: loss of second outmost electron
M+  M2+ + e-
Na+ = 9 (1s22s22p6, only two core electrons)
Al+ = 3 (1s22s22p63s2 , ten core electrons)
72
Practice: Estimate Zeff for
Na+, Al3+, S2-, Cl-,
Step 1. Write the electron configuration
Step 2. identify core electrons vs. outmost
electrons
Step 3. Zeff = Z – #core electrons
Na+ = 9, Al3+ = 11, S2- = 6, Cl- = 7
73
Na+ = 9, Al3+ = 11, S2- = 6, Cl- = 7
Application of Zeff
Within a period, higher Zeff leads to smaller
radius. Which of the following pairs has smaller
radiu:
Al3+ vs. Na+
S2- vs. Cl-
74
The Periodic Table – Final Thoughts
1. It is the number and type of valence electrons that
primarily determine an atom’s chemistry.
2. Electron configurations can be determined from the
organization of the periodic table.
3. Certain groups in the periodic table have special
names.
75
Special Names for Groups in the
Periodic Table
76
Metals (including
Lanthanides/Actinides), Nonmetals,
Metalloids
77
The Alkali Metals
• Li, Na, K, Rb, Cs, and Fr
 Most chemically reactive of the metals
 React with nonmetals to form ionic solids
 Going down group:




Ionization energy decreases
Atomic radius increases
Density increases
Melting and boiling points smoothly decrease
78
Self study materials
• Li, Na, K, Rb, Cs, and Fr
 Most chemically reactive of the metals
 React with nonmetals to form ionic solids
 Going down group:




Ionization energy decreases
Atomic radius increases
Density increases
Melting and boiling points smoothly decrease
79
Behavior Patterns for IE and EA
Reactive nonmetals have high IEs and highly negative EAs.
-
These elements attract electrons strongly and tend to form negative ions in ionic compounds.
Reactive metals have low IEs and slightly negative EAs.
-
These elements lose electrons easily and tend to form positive ions in ionic compounds.
Noble gases have very high IEs and slightly positive EAs.
-
These elements tend to neither lose nor gain electrons.
Figure 8.19
Trends in three atomic properties.
` Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Metallic Behavior
• Metals are typically shiny solids with moderate to high
melting points.
• Metals are good conductors of heat and electricity, and
can easily be shaped.
• Metals tend to lose electrons and form cations, i.e., they
are easily oxidized.
• Metals are generally strong reducing agents.
• Most metals form ionic oxides, which are basic in
aqueous solution.
Figure 8.20
Trends in metallic behavior.
Figure 8.21
Metallic behavior in Group 5A(15) and Period 3.
` Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Metallic behavior decreases across the period
Metallic behavior increases down the group
The McGraw-Hill Companies, Inc./Stephen Frisch Photographer
Figure 8.22
Highest and lowest O.N.s of reactive main-group elements.
Figure 8.23
Oxide acidity.
CaO, the oxide of a main-group metal, is strongly basic.
P O , the oxide of a main-group nonmetal, is acidic.
4 10
Acid-Base Behavior of Oxides
Main-group metals form ionic oxides, which are basic in aqueous solution.
Main-group nonmetals form covalent oxides, which are acidic in aqueous
solution.
Some metals and metalloids from amphoteric oxides, which can act as acids or
bases in water:
Al O (s) + 6HCl (aq) → 2AlCl (aq) + 3H O (l)
2 3
3
2
Al O (s) + 2NaOH (aq) + 3H O (l) → 2NaAl(OH) (aq)
2 3
2
4
Figure 8.24
Acid-base behavior of some element oxides.
Oxides become more basic down a group and more acidic across a period.
Electron configurations of Monatomic Ions
Elements at either end of a period gain or lose electrons to attain a filled outer
level. The resulting ion will have a noble gas electron configuration and is said
to be isoelectronic with that noble gas.
Na(1s22s22p63s1) → e– + Na+([He]2s22p6)
[isoelectronic with Ne]
Br([Ar]4s23d104p5) + e– → Br- ([Ar]4s23d104p6)
[isoelectronic with Kr]
Figure 8.25
Main-group elements whose ions have noble gas electron configurations.
Electron configurations of Monatomic Ions
A pseudo-noble gas configuration is attained when a metal atom empties its
highest energy level.
The ion attains the stability of empty ns and np sublevels and a filled (n – 1)d sublevel.
Sn ([Kr]5s24d105p2) → 4e– + Sn4+ ([Kr]4d10)
A metal may lose only the np electrons to attain an inert pair configuration.
The ion attains the stability of a filled ns and (n – 1)d sublevels.
Sn([Kr]5s24d105p2) → 2e– + Sn2+ ([Kr]5s24d10)
Writing Electron Configurations of Main-Group Ions
Using condensed electron configurations, write reactions for the formation of the
common ions of the following elements:
PROBLEM:
(a) Iodine (Z = 53)
PLAN:
(b) Potassium (Z = 19)
(c) Indium (Z = 49)
Identify the position of each element on the periodic table and recall that:
•
•
Ions of elements in Groups 1A(1), 2A(2), 6A(16), and 7A(17) are usually isoelectronic
with the nearest noble gas.
Metals in Groups 3A(13) to 5A(15) can lose the ns and np electrons or just the np
electrons.
Sample Problem 8.6
SOLUTION:
(a) Iodine (Z = 53) is in Group 7A(17) and will gain one electron to be isoelectronic with Xe: I ([Kr]
5s24d105p5) + e– → I– ([Kr] 5s24d105p6)
(b) Potassium (Z = 19) is in Group 1A(1) and will lose one electron to be isoelectronic with Ar: K
([Ar] 4s1) → K+ ([Ar]) + e–
(c) Indium (Z = 49) is in Group 3A(13) and can lose either one electron or three electrons: In ([Kr]
5s24d105p1) → In+ ([Kr] 5s24d10) + e–
In ([Kr] 5s24d105p1) →
In3+([Kr] 4d10) + 3e–
Figure 8.26
The crossover of sublevel energies in Period 4.
Magnetic Properties of Transition Metal ions
A species with one or more unpaired electrons exhibits paramagnetism – it is
attracted by a magnetic field.
Ag (Z = 47)
↑
5s
↑↓
↑↓
↑↓
4d
↑↓
↑↓
5p
A species with all its electrons paired exhibits diamagnetism – it is not attracted
(and is slightly repelled) by a magnetic field.
Figure 8.27
Measuring the magnetic behavior of a sample.
The apparent mass of a diamagnetic
substance is unaffected by the
magnetic
field.
The apparent mass of a
paramagnetic substance
increases as it is attracted
by the magnetic field.
Magnetic Properties of Transition Metal Ions
Magnetic behavior can provide evidence for the electron configuration of a
given ion.
Ti (Z = 22)
↑↓
↑
↑
4s
↑
Ti2+
4s
3d
4p
3d
4p
↑
Ti2+ has 2 unpaired electrons and is paramagnetic, providing evidence that
the 4s electrons are lost before the 3d electrons.
Sample Problem 8.7
PROBLEM:
Writing Electron Configurations and Predicting Magnetic
Behavior of Transition Metal Ions
Use condensed electron configurations to write the reaction for the formation of
each transition metal ion, and predict whether the ion is paramagnetic or
diamagnetic.
(a) Mn2+(Z = 25)
PLAN:
(b) Cr3+(Z = 24)
(c) Hg2+(Z = 80)
Write the condensed electron configuration for each atom, recalling the irregularity for
Cr. Remove electrons, beginning with the ns electrons, and determine if there are any
unpaired electrons.
Sample Problem 8.7
SOLUTION:
(a) Mn2+(Z = 25) Mn ([Ar] 4s23d5) → Mn2+ ([Ar] 3d5) + 2e−
There are 5 3d electrons and they are all unpaired; Mn2+ is paramagnetic.
(b) Cr3+(Z = 24) Cr ([Ar] 4s13d5) → Cr3+ ([Ar] 3d3) + 3e−
There are 3 3d electrons and they are all unpaired; Cr3+ is paramagnetic.
(c) Hg2+(Z = 80) Hg ([Xe] 6s24f145d10) → Hg2+ ([Xe] 4f145d10) + 2e−
The 4f and the 5s sublevels are filled, so there are no unpaired electrons. Hg2+ is diamagnetic.
Ionic Size vs. Atomic Size
Cations are smaller than their parent atoms while anions are larger.
Ionic radius increases down a group as n increases.
Cation size decreases as charge increases.
An isoelectronic series is a series of ions that have the same electron
configuration. Within the series, ion size decreases with increasing nuclear
charge.
3– > 2– > 1– > 1+ > 2+ > 3+
Figure 8.28
Ionic radius.
Ionic vs. atomic radii.
Sample Problem 8.8
PROBLEM:
Ranking Ions by Size
Rank each set of ions in order of decreasing size, and explain your ranking:
(a) Ca2+, Sr2+, Mg2+
PLAN:
(b) K+, S2−, Cl−
Find the position of each element on the periodic table and apply the trends for ionic size.
SOLUTION:
(a)
(c) Au+, Au3+
Sr2+ > Ca2+ > Mg2+
All these ions are from Group 2A, so size increases down the group.
Sample Problem 8.8
SOLUTION:
(b) S2− > Cl− > K+
These ions are isoelectronic, so size decreases as nuclear charge increases.
(c) Au+ > Au3+
Cation size decreases as charge increases.