Chapter 9 (Hill/Petrucci/McCreary/Perry
Chemical Bonds: Ionic and Covalent
This chapter deals with the forces that chemists call chemical bonds …chemical bonds hold ions
together in ionic compounds like salts and hold atoms together in molecular compounds. These
intramolecular forces are in contrast to the bonds between molecules i.e. those that hold
different molecules together … here, stronger intermolecular forces result in higher melting
points and boiling points, for example …
“Chemical bonds are electrical forces; they reflect a balance in the forces of attraction and
repulsion between electrically charged particles.” (Hill, p.339)
Chemical Bonds: A Preview
] Read Sections 9.1 and 9.2, pp. 339-341 (Hill)
chemical bond: a force that holds atoms together in molecules or ions together in crystals
covalent bond: a bond formed by a pair of electrons, shared more or less equally between two
atoms
ionic bond: the sum of the attractive forces between the positive and negative ions (cations and
anions) that holds them together in a solid, crystal lattice
Lewis Theory (G. N. Lewis)
1. The valence electrons of atoms are the only electrons involved in chemical bonds.
2. Reactions between a metal and a nonmetal: the metal loses one or more valence electrons to
the nonmetal… “ionic bonds”
3. Reactions between two nonmetals: the metal loses one or more pairs of valence electrons are
shared between the atoms…. “covalent bonds”
4. Atoms gain or lose electrons in such a way that they acquire a noble gas electron
configuration.
The Octet Rule
In molecular structures, many main (A-) group elements have arrangements of valence electrons
around them that contain eight electrons (an octet) whenever possible … ns 2 np6 … with eight
electrons in their valence shell.
Lewis symbol: chemical symbol for an element with its valence electrons represented as “dots”
or pairs of dots.
Lewis dot notation for a neon atom
a Lewis symbol "
Ne
Generalizations About Lewis Dot Symbols for Atoms
Lewis dot symbols are used for A-group elements, i.e. Groups IA through VIIIA
Recall that the number of valence electrons for A-group elements is equal to the group number,
e.g. C, Group IVA, 4 valence electrons. See and work through Exercise 9.1A and 9.1B on
your own …
Lewis Dot Symbols for Other Selected A-Group Elements
See Mg, P and Si in Hill, top of p. 342
Reaction with Na with Cl2 to Form NaCl
The violent reaction of a green gas with a shiny soft metal produces common table salt …
2 Na(s) + Cl2 (g) " 2 NaCl(s) + heat + light!!
Ionic Bonding and Formation of Ionic Crystals …
Stepwise schematic formation of a NaCl crystal … see Fig. 9.4, p. 343, Hill
• two sodium atoms each “give” an electron to a Cl atom in a Cl2 molecule .. result 2 Na+ and 2
Cl 1‚ two Na+ combine with two Cl1- to form 2 NaCl “units”
Stepwise Formation of NaCl: Lewis Approach (see bottom of p. 343, Hill)
Formation of Na+ : Na0 " Na+ + e1Formation of Cl1-: ½ Cl2 + e1- " Cl1Look at example 9.2 and Exercises 9.2A and 9.2B, p. 344, Hill
Skip Section 4.5: “Energy Changes in Ionic Compound Formation,” pp. 344-347, Hill.
Reaction of Two H· (Atoms) to Form H2 (see graphic, bottom of p. 347, Hill)
The electron pair between the nuclei of the two H-atoms constitutes a “single” covalent bond in
which the two positive nuc lei share the electron density of the two electrons equally.
H2 Bond Energy as a Function of Internuclear Distance (Figure 9.2, p. 340, Hill)
• two separated H-atoms approach each other
‚ the bottom of the potential energy well at 74 pm separation … an energy minimum at
equilibrium bond distance
ƒ repulsion increases exponentially as H-atoms get too close
Covalent Bonding
In covalent bonds, one, two or three pairs of electrons are shared between two atomic nuclei
a Cl-Cl single bond
bonding pair
bonding e- pair (shared, (between nuclei)
lone e- pair (“local,” on only 1 atom)
Cl
Cl
lone pair
Lewis Dot Structure for CO2
O
C
O
CO2 has 4 bonding pairs and 4 lone pairs of electrons
Lewis Dot Structure for N2
N
N
N2 has 3 bonding pairs and 2 lone pairs of electrons
Lewis Dot Structure for O2
O2 has 3 bonding pairs and 2 lone pairs of electrons
O
O
=
O
O
Oxygen is Paramagnetic … What is the significance? (see Figure 9.7, p. 349, Hill)
What does paramagnetic behavior suggest about unpaired electrons in a molecular species? One
or more unpaired electrons! Does this support the Lewis structure we found for O2 ? Nope!!
We can draw plausible Lewis structures, but they don’t necessarily conform to the “way”
molecules really are!!
Molecular Structures for Some Selected Small Molecules
Note: Lewis structure gives no info about molecular shape!!
Electronegativity Values for the Elements
electronegativity (Χ): a quantitative measure of the ability of a bonded atom to draw electron
density to itself and away from the other atom in the bond
Χ = (I.E. + E.A.)/2
Robert Mulliken (1934)
∆Χ is a measure of bond polarity
Trends in Electronegativity
The same as for both the ionization energy and the electron affinity …. Χ increases up and to the
right on the periodic table! Also, inversely proportional to atomic size
∆Χ for Selected Bonds
Example. CO
∆Χ = Χ(O) – Χ(C) = 3.5 – 2.5 = 1.0
When ∆Χ ≥ 1.7, a bond achieves about 50% “ionic character.”
HCl: The H-Cl Bond Is a Polar Covalent Bond
(see “Depicting Polar Covalent Bonds, p. 353, Hill)
direction of shift for negative charge in the bond …from H toward Cl
" partial + and partial – charges are shown on H and Cl, respectively, where δ = “partial”
Distribution of Charge: Nonpolar and Polar Covalent Bonds (Fig. 9.11, p. 353)
Note symmetrical distribution of electron charge in diatomic H2, where both atoms in the bond
are identical ….In HCl, the more electronegative Cl atom “shifts” electron density in the bond
toward itself …
Structure of NH3 : Central and Terminal Atoms (see NH3 , p. 354, Hill)
Terminal atoms are attached to a central atom…e.g. H atoms are terminal atoms…N is the
“central atom” …bonded to two or more atoms …
Molecular Structure and Electronegativities (p. 354)
COCl2 … here, ΧCl > ΧC …the central atom usually has the lowest electronegativity, with
terminal atoms tending to have higher electronegativities.
Strategy for Writing Lewis Structures for Molecules
1. Count total number of valence electrons for all atoms – use group number
2. Write or begin with a given skeletal structure
3. Give all terminal atoms enough electron pairs to satisfy their octets (exceptions are Groups IAIIIA …H atoms get one pair
4. Assign remaining pairs to central atom as lone pairs until all electron pairs are added.
5. Form multiple bonds on central atom as necessary
See Examples 9.6, 9.7 and 9.8, pp. 355-357, Hill
See Exercises 9.6, 9.7 and 9.8, pp. 355-357, Hill
Other examples: CCl2 F2 , CO2 , H3 O+, ClO 2 1Formal Charge
If we can draw more than one correct Lewis structure, how can we say one structure is “better”
than another structure? Formal charge provides a rational basis for us to favor one structure over
another. We determine the formal charge on each atom in a molecular structure …
formal cha rge (FC) = # valence e 1- - lone pair e 1- - ½*(number of bonded e 1-)
C
Example. CS2
S
S
1
2
Formal Charges for CS 2
C
FC = 4 – 4 – ½(4) = - 2
S1
FC = 6 – 0 – ½(8) = +2
S2
FC = 6 – 4 – ½(4) = 0
Application of Formal Charge to Choice of the Best Lewis Structure
1. Good structures have formal charges of 0 on all atoms.
2. If all formal charges are not 0, then they should be as small as possible with negative charges
only on electronegative atoms.
3. Two neighboring atoms should not have formal charges with the same sign (+ or -).
4. The sum of formal charges must equal the net charge on the structure.
Example 9.9, p. 359: “best” structure is structure (a)
Example. H3 NO
H
O
N
H H
3 + 5 + 6 = 14 e 1-
N
H
H
I
II
O
H
To answer, first find and assign formal charge to each
atom in each structure…
Assign formal charges to each atom … the winner!! Structure I
Formal Charges for Two Possible O3 Structures (p. 360, Hill)
Note that the formal charges are the same for the two structures. The double- headed arrow ( 1 )
means “is equivalent to.”
Delocalized Bonding: “Resonance”
We say that resonance exists whenever we can write two (or more) plausible structures for a
given molecule…the only difference is in the arrangement of the electrons in the structure.
The “best” representation for the molecule is a composite of the contributing resonance
structures.
Example. O3 see Structure I and Structure II (discussed just above) for O3
Two Resonance Forms for O3
What would the “composite” structure look like?”
O
O
O
Note: the ---- represent four(4) electrons shared among the three O atoms
Skeleton Structure for SO3
O
O
S
O
O
O
S
O
O
I
O
S
II
O
O
O
S
O
III
Lewis Structures for “Odd-Electron” Compounds
Normally, all electrons appear as pairs in Lewis structures … the types of compounds shown
here are unusual … no octets for odd-electron atoms. See NO, NO2 and ClO 2 , top of page 362
(Hill).
Molecules with Incomplete Octets: Electron-Deficient Molecules (see p. 362)
Occurs when central atom is Be, B, or Al.
BF3 : Four resonance structures …but the last three have unfavorable formal charges… why?
Structures with “Expanded Octets”
Cl
Cl
Cl
P
Cl
Cl
Phosphorus pentachloride - the central P atom has 5 bonding pairs of electrons around it.
Elements in Groups V, VI and VII tend to have more than an octet of electrons for elements
below the 2nd Period.
Structures with “Expanded Octets”
S
S
S
S
S
S
S
Sulfur hexafluoride - the central S atom has 6 bonding pairs of electrons around it. Sulfur is in
the 3rd Period in Group VI.
See Examples 9.11 and 9.12, pp. 365-366, Hill
Work through Exercises 9.11 and 9.12 (pp. 365-366, Hill) on your own and check answers in the
back of the text.
Bond Lengths and Bond Energies
bond length: the distance between the nuclei of two atoms that are chemically bonded together
bond energy: the energy required to break a covalent bond that joins two atoms together… bond
energy is a measure of bond strength
] Bond-breaking is an endothermic process!
bond dissociation energy: energy required to break exactly one mol of a given type of covalent
bond
Bond length ∝ 1/bond energy
Weaker bonds are longer bonds and vice versa.
Bond Order, Bond Length and Bond Energies
bond order: the number of bonds between two atoms
Bond Order for C to C Bonds
C
C
B.O. = 1.00
C
C
B.O. = 2.00
C
C
B.O. = 3.00
Bond Energies and Bond Lengths for Selected Covalent Bonds (Table 9.1, p.367)
Molecular Bond Energies
Bond breaking = endothermic
Bond formation = exothermic
To form a molecule, we first have to break bonds of the reactant molecules and then reform
bonds of the product molecule …
See Example 9.14, p. 370, Hill