Week
Period
Chapter
Chemical Principles / Atkins, Jones
Note
1st
3/4-3/7
2nd
3/11-3/14
3rd
3/18-3/21
4th
3/25-3/28
2.1-2.16
Ionic Bonds, Lewis Structures and Covalent Bonds
SS: Chap 2 (3/25)
5th
4/1-4/4
3.1-3.7
VSEPR and Valence-Bond Theory
Q: Chap 1 (4/1)
6th
4/8-4/11
3.8-3.12
Molecular Orbital Theory
SS: Chap 3 (4/8)
7th
4/15-4/18
8th
4/22-4/25
Mid-term Exam
(Chapters 1-3)
9th
4/29-5/2
4.1-4.11
The Gas Laws, Molecular Motion
10th
5/6-5/9
4.12-5.6
Real Gases, Intermolecular Forces
SS: Chap 4 (5/6)
11th
5/13-5/16
5.7-5.16
Liquids and Solids
SS: Chap 5 (5/13)
12th
5/20-5/23
7.1-7.12
Enthalpy
Q: Chap 4 and 5 (5/20)
13th
5/27-5/30
7.13-8.8
Chemical Change and Entropy
SS: Chap 7 and 8 (5/27)
14th
6/3-6/5
8.9-8.16
Gibbs Free Energy
Q: Chap 7 (6/3)
15th
6/10-6/13
16th
6/17-6/20
Q: Chap 2 (4/15)
Additional Class / Review
Final Exam
* The coverage of the Quiz (Q)
* The coverage of the Study Summary (SS)
2013 General Chemistry I
(Chapters 4, 5, 7, 8)
Chapter 2.
CHEMICAL BONDS
IONIC BONDS
2.1
2.2
2.3
2.4
The Ions That Elements Form
Lewis Symbols
The Energetics of Ionic Bond Formation
Interactions Between Ions
COVALENT BONDS
2.5
2.6
2.7
2.8
Lewis Structures
Lewis Structures of Polyatomic Species
Resonance
Formal Charge
2013 General Chemistry I
Chapter 2. CHEMICAL BONDS
(화학결합)
Lowering of energy by rearranging valence electrons
Understanding the bond formation between atoms ⇒⇒⇒
understanding properties and reactivity of materials
Designing new materials
Ionic bond(이온결합); electron transfer + electrostatic attraction, NaCl
Covalent bond(공유결합); sharing electrons, NH3
Metallic bond(금속결합); cations held by a sea of electrons, copper
Lewis structure (루이스구조)
Octet rule (옥텟 규칙) ---- ---- coordinate covalent bond (배위결합)
Resonance (공명)
Formal charge (형식전하)
Oxidation number (산화수)
2013 General Chemistry I
Chapter 2. CHEMICAL BONDS
 Chemical bond is the link between atoms.
Key Ideas
Goal
Bond formation by lowering of energy
How are we going to apply Q.M. knowledge ?
Understanding the bond formation between atoms ⇒⇒⇒ Designing new materials
Lowering of energy by rearranging valence electrons
Ionic bond; electron transfer + electrostatic attraction, NaCl
Covalent bond; sharing electrons, NH3
Metallic bond; cations held by a sea of electrons, copper
2013 General Chemistry I
IONIC BONDS (Sections 2.1-2.4)
 ionic model: the description of bonding in terms of ions
ionic solid: three-dimensional crystalline solid
an assembly of cations and anions stacked together in a regular pattern
Metal + nonmetal
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2.1 The Ions That Elements Form
 Cations: Remove outermost electrons in the order np, ns, (n-1)d
Metallic elements in the s block and on the left of the p block in Period 3 (Al)
⇒ lose electrons down to their noble-gas cores; 1s2 (duplet) or ns2np6 (octet)
Li+, Be2+, Na+, Mg2+, Al3+, ···
Metallic elements of the p block in Periods 4 and later
⇒ lose electrons down to their noble-gas cores surrounded by d10; (n – 1)d10ns2np6
The delectrons, in most cases, cannot be lost.
Ga3+([Ar]3d10), ···
Elements in the d block
⇒ ns-electrons first, then variable number of (n – 1)d-electrons
Fe2+([Ar]3d6), Fe3+([Ar]3d5), ···
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Many elements in the p and d blocks; variable valence
⇒ Inert-pair effect; p-electrons alone or all their valence p- and s-electrons
In+, In3+, Sn2+, Sn4+, ···
The order of losing electrons:
 Anions: Add electrons until the next noble-gas configuration is reached
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2.2 Lewis Symbols
- valence electrons – depicted as dots; a pair of dots for paired electrons
(1916)
- cations and anions
2013 General Chemistry I
Formulation of ionic compounds
1) Cations by removing all dots from metal atoms:
2) Anions by adding dots to complete the valence shell (octet or duplet):
3) Adjust the number of each element to conserve the total number of dots.
4) Write the charge of each ion:
CaCl2; empirical formula
2013 General Chemistry I
2.3 The Energetics of Ionic Bond Formation
NaCl ionic crystal formation;
electron transfer from Na to Cl and then Coulomb stabilization
The energy required for the formation of ionic bonds is supplied largely by the attraction
between oppositely charged ions.
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2.4 Interactions Between Ions
-In an ionic solid, each cation is attracted to all the anions to a greater or lesser extent. →
a “global” characteristic of the entire crystal
i.e. ionic bond is not a bond between two ions !
 Lattice energy: the difference in energy between the ions packed together in a solid
and the ions widely separated as a gas
- strong electrostatic interactions in ionic solids
→ high melting points and brittleness
2013 General Chemistry I
2.4 Interactions Between Ions
- Coulomb potential energy of the interactions of two individual ions
e is the fundamental charge; z1 and z2 are the charge numbers of the
two ions; r12 is the distance between the centers of the ions; e0 is the
vacuum permittivity.
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- In a one-dimensional crystal in which cations and anions alternate along a line,
toward right hand side
toward left hand side
; A = 2 ln2 or 1.386
- molar potential energy of a three-dimensional crystal
 The factor A is the Madelung constant,
dependent on how the ions are arranged about one another
2013 General Chemistry I
- real potential energy of an ionic solid
→ attractive and repulsive interionic interactions in close range.
Short-range repulsions between ions
Total potential energy = EP + EP*
2013 General Chemistry I
Born-Meyer equation;
correction for repulsions to the Madelung constant
repulsive effect
The coulombic interaction between ions in a solid is large
when the ions are small and highly charged.
2013 General Chemistry I
COVALENT BONDS
"With brilliant insight, and before anyone knew about quantum mechanics or orbitals,
Lewis proposed"
Quantum mechanical view of covalent bond
Chemical bonding in H2+ by sharing an electron between two protons
An electron between the two nuclei exerts an attractive force on the nuclei.
2013 General Chemistry I
2.5 Lewis Structures
Covalent bonding; octet (or duplet) by sharing (Lewis, 1916)
- octet rule: atoms go as far as possible toward completing their octets
Nonmetal atoms share electrons to complete their octet;
lines (bonding pairs), dots (lone pairs)
- A line (-) represents a shared pair of electrons : a bond.
2013 General Chemistry I
2.6 Lewis Structures of Polyatomic Species
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- Each atom completes its octet by sharing pairs of electrons.
Methane; CH4 (why not CH3 nor CH5?)
- Lewis structure does not portray the 3D shape of a molecule or ion,
but simply displays which atoms are bonded together.
- bond order: the number of bonds that link a specific pair of atoms.
2013 General Chemistry I
 Writing a Lewis structure
- terminal atom: bonded to only one other atom
central atom: bonded to at least two other atoms
- The element with the lowest ionization energy (less greedy) as the central atom
electronegativity is a better indicator
Ex. HCN
- Atoms symmetrically around the central atom; SOS for S2O (Exception; NNO)
- OH is attached to the central atom in oxoacids; HO–Cl for HClO
i.e. H2SO4 ----- (HO)2SO2
- Polyatomic ions; total number of electrons should be adjusted to represent the overall
charge
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Toolbox 2.1: How to write the Lewis structure of a polyatomic species
Step 1: # of electron pairs = total # of valence electrons / 2
Step 2: Write down the most likely arrangements of atoms.
Step 3: One electron pair between each pair of bonded atoms
Step 4: Complete the octet (duplet for H) of each atom
using the remaining electron pairs.
Form multiple bonds if in short of electron pairs.
Step 5: A line for each bonding pair
2013 General Chemistry I
Take care of charges.
2013 General Chemistry I
Ex 2.4
Lewis structures for CH3COOH (multiple central atoms)
# of electron pairs = (4 + 3 + 4 + 2 6 + 1) / 2 = 12
CH3COOH
Lewis structures for C2O2H4 (multiple central atoms)
How many structures can you draw?
2013 General Chemistry I
Lewis structures for NO3-
3
# of electron pairs = (5 + 3
12
6 + 1) / 2 = 12
Identical N–O bond lengths of 124 pm
(> 120 pm for N=O, < 140 pm for N–O)
- double-headed arrows (↔), indicating a blend of the contributing structures
- delocalization: a shared electron pair is distributed over several pairs of atoms
and cannot be identified with just one pair of atoms.
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 Benzene, C6H6
- All the carbon-carbon bonds with the same length
- Only one 1,2-dichlorobenzen exists.
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2.8 Formal Charge
70
 Formal charge – the charge it would have if the bonding were perfectly covalent in the
sense that the atom had exactly a half-share in the bonding electrons
V = the number of valence electrons in the free atom
L = the number of electrons present on the bonded atom as lone pairs
B = the number of bonding electrons on the atom
Formal charge indicates the degree of redistribution of electrons relative to free atoms
not the real charge of an atom
- A Lewis structure in which the formal charges of the individual atoms are closest to zero
typically represents the lowest energy arrangement of the atoms and electrons.
2013 General Chemistry I
OCO with lower formal charges is more likely for CO2 than COO.
HCN has lower formal charges than HNC.
NNO with lower formal charges is more likely for ON2 than NON.
- Formal charge exaggerates the covalent character of bonds by
assuming that the electrons are shared equally.
- Oxidation number exaggerates the ionic character of bonds.
It represents the atoms as ions, and all the electrons in a bond are assigned to the
atom with the lower ionization energy (higher electronegativity (2.12)).
formal charge
oxidation state
2013 General Chemistry I
Formal charges differ from oxidation numbers!
Neither of them is the true charge.
Quantum mechanically, there is no true localized charge on an atom!
Ca2+ is an oxidation state of calcium with the oxidation number of “+2”.
Oxidation number is important in following the oxidation-reduction reaction.
- octet rule: In covalent bond formation, atoms go as far as possible
toward completing their octets by sharing electron pairs.
There are many exceptions to the octet rule
2013 General Chemistry I
2.9 Radicals and Biradicals
Odd number of electrons
72
# of electron pairs = (5 + 2 x 6) / 2 = 8.5
 Radicals: species with an unpaired electron, highly reactive
Biradicals: molecules with two unpaired electrons
2.10 Expanded Valence Shells
Expanded valence shell (expanded octet); large atoms with empty d-orbitals
(Period 3 or later) may accommodate more than 8 electrons.
PCl5 vs. NCl5
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- hypervalent compound: a compound that contains an atom with
more atoms attached to it than is permitted by the octet rule
ST 2.10B
Linear I3– ion
3x7 + 1 = 22 electrons, 11 electron pairs
2 bonds, 2 3 + 2 = 8 lone pairs
Remaining one pair into the central I
2013 General Chemistry I
Lewis structures for C2O2H4 (multiple central atoms)
# of electron pairs = (4 + 3 + 4 + 2
2
3
2013 General Chemistry I
6 + 1) / 2 = 12
3
3
3
3
72
- variable covalence: the ability to form different numbers of covalent bonds
Ex 2.8
5
Dominant resonance Lewis structure of SO42–
6 + 2 = 32 valence electrons, 16 electron pairs
most preferred structure
number of resonance structures
2013 General Chemistry I
THE PERIODICITY OF ATOMIC PROPERTIES
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2.11 The Unusual Structures of Some Group 13/III Compounds
- boron and aluminum
- incomplete octet: fewer than eight valence electrons
- completing octets by a coordinate covalent bond, in which both electrons come from
one of the atoms
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2.11 The Unusual Structures of Some Group 13/III Compounds
Possible due to the atomic radius of Al in AlCl3 larger than that of B in BCl3.
borane
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Chapter 2. CHEMICAL BONDS
Ionic bond; electron transfer + electrostatic attraction, NaCl
 Oxidation number– exaggerates the ionic character of bonds.
It represents the atoms as ions, and all the electrons in a bond are assigned to the atom
with the lower ionization energy (higher electronegativity ).
Covalent bond; sharing electrons, NH3
 Formal charge – the charge it would have if the bonding were perfectly covalent in the
sense that the atom had exactly a half-share in the bonding electrons
Nonpolar covalent bond; the average charge on each atom is zero.
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lower energy
structure
the average charge on each atom is not zero.
- partial charges: the charges on the atoms
- polar covalent bond: a bond in which ionic contributions to the resonance result in
partial charges
- electric dipole: a partial positive charge next to an equal but opposite partial negative charge
size of an electric dipole ---- electric dipole moment (m)
Unit: Debye (D)
definition: a dipole between electron and proton
separated by 100 pm is 4.80D
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Cl-H bond: m = ~1.1 D : Cl has ~23% of an electron’s charge
2.12 Correcting the Covalent Model: Electronegativity
 Electronegativity (c) – Electron-pulling power of an atom when it is
part of a molecule (by Linus Pauling in 1932)
Rough guide to the charge separation in a bond between two atoms
Average based on measured bond energies from a large range of compounds; can be revised
Measure of extra stability due to ionic contributions
2013 General Chemistry I
2.12 Correcting the Covalent Model: Electronegativity
Mulliken’s electronegativity scale (1934); properties of an isolated atom
Exactly defined
- Mulliken scale: c = ½(I + Eea)
average of the ionization energy and electron affinity
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- rough rules of thumb
ionic
polar covalent
covalent
2013 General Chemistry I
i.e. NaCl or KF : ionic
C-O
Ca-O
: polar covalent
: ionic
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2.13 Correcting the Ionic Model: Polarizability
- All ionic bonds have some covalent character.
- highly polarizable atoms and ions: readily undergo a large distortion
of their electron cloud
i.e. large anions and atoms such as I-, Br-, and Cl- polarizing power: property of ions (and atoms) that cause large
distortions of electron clouds
- increases with decreasing size and increasing
charge of a cation
i.e. the small and/or highly charged cations
Li+, Be2+, Mg2+, and Al3+
2013 General Chemistry I
THE STRENGTHS AND LENGTHS OF COVALENT BONDS
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2.14 Bond Strength
measured by
 Dissociation energy (D): energy required to separate the bonded atoms
- The bond breaking is homolytic, which means that each atom retains one of the electrons from the bond.
- average dissociation energy for one type of
bond found in different molecules
i.e. C-H single bond:
average strength of bonds in a selection
of organic molecules, such as methane (CH4),
ethane (C2H6), and ethene (C2H4)
2013 General Chemistry I
2.15 Variation in Bond Strength
Strongest bond between two nonmetal atoms; CO (1062 kJ/mol)
Lone pair-lone pair repulsion
due to the short F–F distance
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80
2013 General Chemistry I
ADP(aq) + H2PO4–(aq) + 174 kJ/mol
ATP(aq) + H2O
bond stiffness (
bond strength); resistance to stretching and compressing
will be discussed in the Major Technique 1: Infrared (IR) Spectroscopy
2013 General Chemistry I
2.16 Bond Length
 Bond length: the distance between the centers of two atoms joined
by a covalent bond
- corresponding to the internuclear distance at the potential energy minimum
for the two atoms
- affecting the overall size and shape of a molecule
evaluated by using spectroscopy or x-ray diffraction methods
- Factors influencing bond length
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Box 2.2 Bond length measurements
quantum mechanical rotational energy
with a rotational quantum number J
microwave spectroscopy
Reproducibility of bond lengths;
constant within a few percent in similar arrangements
2013 General Chemistry I
 Covalent radius: contribution an atom to the length of a covalent bond
- Decreases from left to right
(increasing Zeff )
- Increases in going down a group
(size of valence shells and better
shielding by inner core electrons)
Bond length: Approximately the sum of the covalent radii of the two atoms
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INFRARED SPECTROSCOPY
 Infrared radiation: electromagnetic radiation with longer wavelengths
(lower frequencies) than red light
~ 1000 nm or ~ 3×1014 Hz
- Molecules by infrared radiation become vibrationally excited.
bond stiffness (
bond strength); resistance to stretching and compressing
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- “stretching” mode: the atoms moving closer and away again.
“bending” mode: bond angles periodically increase and decrease.
 Vibrational frequencies
- The stiffness of a bond measured by its force constant, k
Force = -k × displacement
by Hooke’s law
- Vibrational frequency, n, of a bond between two atoms A and B of mass mA and mB
n=
1
2π
k
µ
m=
2013 General Chemistry I
𝑚𝐴𝑚𝐵
𝑚𝐴+𝑚𝐵
m = effective mass (or reduced mass)
 Normal modes of vibration of polyatomic molecules
A nonlinear molecule consisting of N atoms
→ 3N-6 normal modes
i.e. H2O, n = 3 → 3 normal modes
A linear molecule → 3N-5 normal modes
CO2, n = 3 → 4 normal modes
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 Actual spectrum
 Characteristic frequencies of functional groups detectable in a spectrum
- fingerprint region: a complex series of absorptions
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