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513 100 General Chemistry
• Lecture
Wed
13.00 – 13.50
Fri
13.00 – 14.45
หอประชุม
หอประชุม
• Discussion Thu
19.25 – 20.25 5406
• Aj. Nattawan Worawannotai
– Room # 218 Sci 3
– [email protected]
– Office hours: 8.30 – 9.30 am M W R or by appointments
– Website: www.chem.sc.su.ac.th/decharin
• Exam
Midterm
Final
Thu October 9th
Fri December 12th
13.00 – 16.00
9.00 – 12.00
• 513 105 General Chemistry Laboratory
Experiments related to the contents in 513 100 General Chemistry.
• http://www.chem.sc.su.ac.th/decharin/2557-1-513100.html
• http://bit.ly/chem513100
• Quick scan –QR code reader
• QR Droid
• Bring a calculator!
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
2
513 100 General Chemistry
• Grading scheme
– Participation and quizzes (0.3 x 10 classes)
3
– Homework assignment (0.6 x 5 weeks)
3
– Exam (midterm and final)
27
• Homework assignment due on Thursday @ 20.30 h
• Discussion session
•
Reference books
• General Chemistry textbooks in the library
• John C. Koltz et al. Chemistry & Chemical Reactivity, 6th ed., Thomson
Brooks/ Cole: Canada, 2006:
• Ralph H. Petrucci et al. General Chemistry: Principles and Modern
Applications; 10th ed.; Pearson Canada: Toronto, 2011.
• Darrell D. Ebbing et al. General Chemistry; 9th ed.; Houghton Mifflin
Company: Boston, 2007
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
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513 100 General Chemistry
• Atomic theory
• Atomic structures and properties of the elements in the periodic
table
• Chemistry of main group elements, non metals and transition metals
• Chemical bonding
• Properties of gases, liquids, solids and solutions
• Stoichiometry
• Chemical equilibrium and ionic equilibrium
• Chemical kinetics
513 100 General Chemistry
Contents (33% of total grade, 15 hours)
Atom & Atomic Theories
• Early Chemical Discoveries
and the Atomic Theory
• Quantum Theory
• The Bohr Atom
• Particle-Wave Duality
• Quantum Numbers and
Electron Orbitals
• Multielectron Atoms
• Electron Configurations
• Periodic Table and periodic
trends
Bonding
•
•
•
•
•
•
•
•
•
Lewis Theory
Resonance
Exceptions to the Octet Rule
Shapes of Molecules
Bond Properties
Valence-Bond Theory
Multiple Covalent Bonds
Molecular Orbital Theory
Bonding in Metals
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
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513 100 General Chemistry
Atom &
Atomic Theories
Dr. Nattawan Decharin, Department of Chemistry, Faculty of Science, Silpakorn University
7
What do you know about an Atom?
• subatomic particles
– Proton +
– Neutron neutral
– Electron -
• Nucleus contains p+ and n -> + charge and almost all the mass of
an atom
• e- surround the nucleus and occupy most of the volume of the atom
• Atoms have no net charge --> # e = # p+
What is the experimental basis of atomic structure?
What is the experimental basis of atomic structure?
• Electricity
– Objects can bear an electric charge.
– Two types of electric charges, + and – (Benjamin Franklin).
– Like charges repel each other and opposite charges attract.
• Radioactivity
– Uranium ore emits rays that could darken a photographic plate
(Henri Becquerel).
– Marie and Pierre Curie isolated polonium and radium which also
emitted the same kind of rays.
– Radioactive atoms emit these unusual rays when they
disintegrate. => radioactivity.
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
9
What is the experimental basis of atomic structure?
• Law of conservation of mass
• Law of constant composition
John Dalton (1803)
– All matters is made of atoms.
– Atoms of a given element are identical.
– Atoms are indivisible and indestructible.
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
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Cathode-Ray Tubes and Characterization of e-’s
• Cathode ray – flows from cathode (-) to anode (+)
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
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Cathode-Ray Tubes and Characterization of e-’s
• Cathode ray
– Attracted toward
+ charged plate
A beam of – charged particles
=>electron
– Can be deflected by
a magnetic field
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
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Cathode-Ray Tubes and Characterization of e-’s
• Joseph John Thomson (1897)
• Balancing the effect of the electric and magnetic fields
• Charge-to-mass ratio of the electron can be determined by
FElectric = FMagnetic
e/m = 1.76x1011 C/kg
• Same e/m in experiments using different metals as cathode and
several different gases. => Electrons are present in atoms of all
elements.
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
13
Charge and Mass of an Electron
• Robert Andrew Millikan (1906-1914)
• Adjustment of the electric charge on
the plate above and below
• Balance the electrostatic force and the
force of gravity acting upon the droplet
Fgravity = Felectrostatic
• Charge of a droplet = n(1.60x10-19 C)
• e = charge of an electron = 1.60x10-19 C
• From e/m = 1.76x1011 C/kg, me = 9.109x10-31 kg
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
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Thomson’s Atomic Model
• A uniform sphere of positively charged matter
•
• Electrons are embedded in this sphere
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
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Canal Ray
• Eugen Goldstein (1886)
• Canal ray
– a stream of + charged particle travelling in the opposite direction
to that of cathode ray
– can be deflected by electric and magnetic field but much less so
than cathode ray for a given value of the field.
– much smaller charge-to-mass ratio than that of cathode ray
– Charge-to-mass values vary depending on the nature of the gas
in the tube
Particles of higher mass than electron
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
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Discovery of the Proton
• Ernest Rutherford
• Irradiation of various element with alpha particles
• Most alpha particles pass straight through, a few were deflected at
large angles, and some came straight back.
• Atoms contain a small + charged nucleus with most of the mass of
am atom
• From the experiments using gaseous element, it was observed that
the deflection of alpha particles was a function of atomic mass
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
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Neutron
• An atom has no net charge=> # of proton = # of electron
• Most atoms have masses greater than the mass of p+e=>atom must also contain relatively massive particles with no
charge
• James Chadwick (1932)
– Experimental evidence of neutron
– No electric charge, mass 1.675x10-24 g
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
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Rutherford’s Atomic Model
• Atom contains a very tiny core which all the positive charge and
most of the mass is concentrated => nucleus
• Electron occupy the rest of the space in the atom
• Subatomic particles
Particle
Mass (g)
Charge (C)
Electron
9.109x10-28
1.602x10-19
Proton
1.673x10-24
1.602x10-19
Neutron
1.675x10-24
0
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
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Atomic Number
• Atoms of the same element have the same number of protons.
• Atomic number = the number of protons in the nucleus of an
element
Atomic number (Z)
Name
Element
symbol
Atomic
weight
• Relative atomic mass:
– Mass of an atom was given relative to the mass of a standard
atom
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
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Atomic Mass
• A carbon atom as a standard:
– 6 protons + 6 neutrons in the nucleus => assigned a mass value
of exactly 12.000
– 1 atomic mass unit (u) = 1/12 mass of a 12C atom
–  a carbon atom has a mass of 12.000 u.
– 1 u = 1.661x10-24 g
• Mass number (A) = (number of protons) + (number of neutrons)
• Atomic symbol
32
eg. 16 S
Mass
number
Atomic
number
A
Z
X
Element
symbol
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
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Isotopes
• All atoms of the same element have the same number of proton.
• Not all of the atoms of the same element have the same mass
number.
10B
Eg.
Z=5, A=10
#p = 5, #n = 5
11B
Z=5, A=11
#p = 5, #n = 6
1H
(proton)
2H (D, deuterium)
3H (T, tritrium)
• Isotope abundance
percent abundance=
# of atoms of a given isotope
×100
total # of atoms of all isotope of that element
– The mass of isotopes and their % abundance are determined
using a mass spectrometer
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
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Mass Spectrometer
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
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Atomic Weight
• Atomic weight = the average weight of a representative sample of
atoms
% abundance isotope 1
Atomic weight=
100
+
% abundance isotope 2
100
mass of isotope 1
mass of isotope 2 + …
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
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Black Body Radiation
Radiation from heated metal
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
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Planck’s Equation
• Max Planck (1900)
– Vibrating atoms in a heated object give rise to electromagnetic
radiation. These vibrations are quantized (only certain vibration,
with specific frequencies, are allowed).
Energy (J)
𝐸 = ℎ
Planck’s constant
(6.626x10-34 Js)
Frequency (s-1)
𝑐 = 
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
27
Photoelectric Effect
• Electrons are ejected from metal surface only if the frequency of the
light is high enough
– Lower frequency=>no electron ejected, regardless of the light’s
intensity.
– Above minimum frequency=>increasing the light intensity
increases the number of electron ejected.
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
28
Photoelectric Effect
• Albert Einstein
– Light has particle-like properties. These massless “particle”,
called photon, are packets of energy.
– The energy of each photon is proportional to the frequency of
the irradiation.
𝑐
𝐸 = ℎ = ℎ

Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
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Atomic Line Spectra
• Excited atoms in gas phase emit only certain wavelengths of light.
• Rydberg Equation and Balmer Series (visible spectrum lines of
hydrogen atoms)
1
1
1
=R 2− 2
2
n

when n > 2
Rydberg constant (1.0974x107 m-1)
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
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The Bohr Model of Hydrogen Atom
• Niels Bohr
– Connection between line emission spectra and quantum idea
– Electron could occupy only certain orbits or energy level in which
it is stable. => the energy of the electron in the atom is
quantized.
– Equation for the energy of an electron in the nth orbit or energy
level of H atom (En)
R = Rydberg constant = 1.0974x107 m-1
Rhc
En = − 2
n
h = Planck’s constant = 6.626x10-34 Js
c = speed of light
n = principle quantum number
– As n increases, En becomes less negative
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
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The Bohr Theory and the Spectra of Excited Atoms
• An electron in an atom would remain in its lowest
energy level.
• Energy must be absorbed to excite the electron
to a higher energy level.
• Energy is emitted when electron “falls” from
higher energy level to lower energy level.
• Atomic spectra arise from the movement of
electron between quantized energy state
∆𝐸 = 𝐸𝑓𝑖𝑛𝑎𝑙 − 𝐸𝑖𝑛𝑖𝑡𝑖𝑎𝑙
 1
1 

E   Rhc 2  2
n

 final ninitial 
Rhc = 2.179x10-18 J/atom
= 1312 kJ/mol
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
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Bohr Model and Its Flaw
Bohr’s model of the atom only
explained the spectrum of H atom
and other system having one
electron (eg. He+, Li2+)
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
33
The Wave Properties of Electron
• A beam of electrons was diffracted like light wave by the atoms of a
shin sheet of metal foil
• Louis Victor de Broglie
– A free electron of mass m moving with a velocity v should have
an associated wavelength given by
h
 = mv
– Electron can be described as having wave properties under
some circumstances
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
34
Quantum Mechanics
• Werner Heisenberg
Heisenberg uncertainty principle
it is impossible to fix both the position of an
electron in an atom and its energy with any degree of certainty
h
x  p 
4
– If we choose to know the energy of an electron in an atom with
only a small uncertainty, then we must accept a correspondingly
large uncertainty in its position in the space about the atom’s
nucleus.
– => We can assess only the likelihood, or probability, of finding
an electron with a given energy within a given region in space
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
35
Schrödinger’s Model and Wave Equation
• Wave equations
• Wave function = solution to wave equations ()
– Only certain wave functions are allowed for the electron in the
atom
– Each wave function is associated with an allowed energy value
for the electron
– Energy of the electron is quantized (electron can have only
certain values of energy)
– 2 is related to the probability of finding the electron within a
given region space. This probability is referred to as the electron
density
– The region of space in which an electron of a given energy is
most probably located is called its orbital
– Quantum numbers are used to define the energy states and
orbitals available to the electron.
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
36
Principal Quantum Numbers
• The principal quantum number, n
– n = 1, 2, 3, …
– Determine the energy of an electron
– Define the size of an orbital: greater n, greater electron’s
average distance from the nucleus
– Electron shell, electron level
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
37
Angular Momentum Quantum Numbers
•
The angular momentum quantum number,
– l = 0, 1, 2, 3, …, n-1
n values
– subshell
– Define the shape or type of the orbital
l
– The values of l are usually coded by letters
Value of l
Subshell label
0
s
1
p
2
d
3
f
– Eg. When n = 2, l can be either 0 or 1.
The 2nd energy level consists of s and p subshells. Orbitals in the s
and p subshells are called s and p orbitals.
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
38
Magnetic Quantum Numbers
• The magnetic quantum number, ml
– Integer values range from –l to + l
(–l, –l+1, –l+2, …, 0, …, l-2, l-1, l)
2l+1 values
– Define the orientation in space of the orbitals within a subshell,
thus the number of orbitals in a given subshell.
– Orbitals in a given subshell differ only in their orientation in
space, not in their energy.
– Eg. l=0 (s subshell)
1 possible ml value is 0 => only 1 s orbital in s subshell
l=1 (p subshell)
3 possible ml value is -1, 0, 1 => 3 p orbitals in p subshell
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
39
Quantum Numbers
n
l
ml
1
0
0
2
0
1
0
-1, 0, 1
3
0
1
2
0
-1, 0, 1
-2, -1, 0, 1, 2
4
0
1
2
3
0
-1, 0, 1
-2, -1, 0, 1, 2
-3, -2, -1, 0, 1, 2, 3
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
40
s Orbitals
• n = 1, l = 0, ml = 0
• Each dot represent where electron could be. The density of dots is
greater close to the nucleus. => the electron is most often found
near the nucleus, and it is less likely to be found further away.
• Spherical in shape
• Probability of finding electron at given distance from the nucleus
=>surface density plot or radial distribution plot (4r22)
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
41
p Orbitals
• n = 2, l = 1, ml = -1, 0, 1
• The values of l is equal to the number of nodal
surfaces that slices through the nucleus.
• Nodal surface = a surface on which there is zero
probability of finding the electron.
• One imaginary plane (nodal surface) that slice
through the nucleus and divides the region of
electron density in half.
• dumbbell shape
• Three orientations possible (3 ml values)
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
42
d Orbitals
• n = 3, l = 2, ml = -2, -1, 0, 1, 2
• Two nodal surfaces that slice through the nucleus.
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
43
f Orbitals
• n = 4, l = 3, ml = -3, -2, -1, 0, 1, 2, 3
• Three nodal surfaces that slice through the nucleus.
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
44
Same Type of Orbitals with Different n Values
• The difference between the same type of orbitals with different n
values is the size . The size increases as n increases.
• Spherical node
• Number of spherical nodes for any orbital is n-l-1
• Number of total node = n-1
http://winter.group.shef.ac.uk/orbitron/
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
45
Electron Spin Quantum Number
• The electron spin quantum number, ms
• When an electron is placed in a magnetic field, two orientations are
possible for the electron spin, aligned or opposed to the magnetic
field. (ms = +1/2 or ms = -1/2)
• Description of an electron in an atom requires 4 quantum numbers
(n, l, ml, ms)
• Substances that are slightly repelled by a strong magnet are called
diamagnetic.
• Substances that are attracted to a strong magnet are called
paramagnetic.
• Paramagnetism arises from unpaired electron.
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
46
The Pauli Exclusion Principle
• No two electrons in an atom can have the same set of four quantum
numbers (n, l, ml, ms)
=> no atomic orbital can contain more than two electrons and
these electrons must have opposite spins.
• Orbital box diagram
Quantum number set
H atom: 1s orbital
1s
n = 1, l = 0, ml = 0, ms = +1/2
Quantum number set
He atom: 1s orbital
1s
n = 1, l = 0, ml = 0, ms = +1/2
n = 1, l = 0, ml = 0, ms = 1/2
• The maximum number of electrons in any shell is 2n2.
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
47
Effective Nuclear Charge (Z*)
• Penetrating effect
– As more electrons are added to an atom, the outermost electron
will penetrate the region occupied by the inner electrons
The energies of the orbitals alter
• Effective nuclear charge = the nuclear
charge “felt” by a particular electron in a
multielectron atom, as modified by the
presence of the other electrons.
• For H atom, with a single electron, the energy depends on the value
of n (E= Rhc/n2).
• For atoms with more than one electron, the energy depends on the
value of n and l.
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
48
Atomic Subshell Energies and Electron Assignment
• Aufbau principle: electrons are assigned to shell and subshell of
increasingly higher energy.
• Assign electrons to subshell in order of
– increasing “n+l”
– For two subshells with the same “n+l” value, assign electron to
the subshell with lower n first
H atom
Multielectron atom
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
49
Atomic Electron Configuration
• Electron configuration = the arrangement of electrons in the element
that results in the lowest energy for the atom.
• spdf notation
e- shell (n)
H atom:
# of e-
1s1
Orbital type (l)
1s
Orbital box notation
He atom:
spdf notation
1s2
Combine with noble
gas notation
1s
Li atom:
1s22s1
1s
or
[He]2s1
2s
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
50
Hund’s Rule
• The most stable arrangement of electrons is that with the maximum
number of unpaired electrons, all with the same spin direction. This
arrangement makes the total energy of an atom as low as possible.
C atom:
1s22s22p2
1s
2s
or
[He]2s22p2
2p
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
51
Electron Configuration and Periodic Table
• Core electron = electrons included in the noble gas notation
• Valence electron = electrons beyond the core electrons
=> determine the chemical properties of an element
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
52
Electron Configurations of Ions
• Cation = one or more of the valence electrons is removed from a
neutral atom
=> from which shell or subshell?
– Remove electron(s) from the electron shell of highest n first
– If several subshells are present, remove electron(s) of maximum
l value
– Eg.
Na+: 1s22s22p6
Na: 1s22s22p63s1
Ge: [Ar]3d104s24p2
Ge2+: [Ar]3d104s2
Ti: [Ar]3d24s2
Ti2+: [Ar]3d2
Fe: [Ar]3d64s2
Fe3+: [Ar]3d5
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
53
Atomic Properties and Periodic Trends
• Similarities in properties of the elements are the result of similar
valence shell electron configurations
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
54
Atomic Size
• How can we define the size of an atom when an orbital has no sharp
boundary?
• Measure the distance between nuclei of adjacent atoms
(internuclear distance)
• For the main group elements, atomic radii increases going down a
group and decreases going across a period.
– going down a group, the outer electrons are in the orbitals with
higher n values =>electrons are further away from the nucleus
– Going across the period, n is the same, but more protons are
added. => increasing effective nuclear charge (Z*) => attraction
between the nucleus and electrons increases => atomic radius
decreases
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
55
Atomic Radius
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
56
Ionization Energy (IE)
• Ionization energy = energy required to remove an electron in gas
phase
X (g)
X+(g) + e• Energy must be added to separate an electron from an atom (an
endothermic process) => IE is always positive
• For main group elements, first IE generally increases across a
period and decrease down a group
– Across a period => increasing effective nuclear charge =>
stronger attraction between the electron and the nucleus => IE1
increases
– Down a group => the electron is further away from the nucleus
=> weaker nucleus-electron attraction => IE1 decreases
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
57
Ionization Energy (IE)
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
58
Ionization Energy (IE)
Mg(g)
[Ne]3s2
Mg+(g)
[Ne]3s1
Mg2+(g)
[Ne]=1s22s22p6
Mg+(g) + e[Ne]3s1
Mg2+(g) + e[Ne]3s0
Mg3+(g) + e1s22s22p5
IE1 = 738 kJ/mol
IE2 = 1451 kJ/mol
IE3 = 7733 kJ/mol
IE1
IE2
IE3
IE4
IE5
IE6
IE7
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
59
Electron Affinity (EA)
• Electron affinity = the energy of a process in which an electron is
acquired by the atom in the gas phase
A(g) + e (g)
A(g)
• The greater the affinity an atom has for an electron, the more
negative the value of EA.
• Similar trend for EA to IE: an element with a high IE generally has a
high affinity for an electron
– EA generally becomes more negative across the period and less
negative down a group
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
60
Electron Affinity (EA)
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
61
Ion Size
• The radius of a cation is always smaller than that of the atom from
which it is derived due to more effective nuclear charge
• The radius of an anion is always larger than that of the atom from
which it is derived due to electron-electron repulsion and less
effective nuclear charge
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
62
Ion Size
• Isoelectronic ions have the same number of electrons
Eg. O2-, F-, Na+, Mg2+
• For isoelectronic ions, the ions with more protons are smaller than
ions with protons.
Dr. Nattawan Worawannotai, Department of Chemistry, Faculty of Science, Silpakorn University
63