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Chapter 8
Chemical Bonding
• Types of Bonds
• Ionic Bonding
• Covalent Bonding
• Shapes of Molecules
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Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
2
Table 8.1 Two Carbon Compounds
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H2O, CCl4AcidsDissolves in
Very lowHighElectrical Conductivity
as a Liquid
Decomposes
1339 (at high P)
2.71
100.1
White solid
CaCO3
Calcium Carbonate
Sublimes at 78.6
-56.6 (at 5.11 atm)
0.00198
44.01
Colorless gas
CO2
Carbon Dioxide
Boiling Point (°C)
Melting Point (°C)
Density (g/mL)
Molar Mass (g/mol)
Physical State
Formula
Compound
3
Table 8.2 General Properties of Ionic and
Covalent Compounds
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Often soluble in carbon tetrachloride but not in water
Poor conductor of electricity and heat
Low boiling point
Low melting point
Weak, brittle solids or soft and waxy solids
Gases, liquids, or solids
Covalent
Often soluble in water but not in carbon tetrachloride
Good electrical conductor when molten or in solution
Very high boiling point
Very high melting point
Hard, brittle solids
Crystalline solids
Ionic
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4
Chemical Bonds• Chemical bond
– A force that holds atoms together in a
molecule or compound
• Two types of
chemical bonds
– Ionic Bonds
– Covalent Bonds
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Figure 8.2
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Ionic Bond• A bond created by
electrostatic attraction
between oppositely charged ions
• Occurs between a metal and a nonmetal
• Electrons transferred between the cation (positively charged ion)
and the anion (negatively charged ion)
• Extremely strong bonds8 -
6
Covalent Bonds• A bond created by the
sharing of electrons between atoms
• Occurs between two nonmetals (resulting in a neutral overall charge)
• Electrons not transferred in this case
• Electrons shared in pairs typically
• Weaker bonds than ionic bonds
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Figure 8.2 or another
molecule picture
3
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Practice – Identifying Types of Bonding
• Identify the type of bonding in each
of the following substances:
1. NaF
2. ClO2
3. FeSO4
4. SO2
5. Ca(ClO2)2
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Practice Solutions – Identifying Types of Bonding
• Identify the type of bonding in each of the following substances:1. NaF – Ionic bonding (metal + nonmetal)
2. ClO2 – Covalent bonding (2 nonmetals)
3. FeSO4 – Ionic bonding between the metal and nonmetal; Covalent bonding between the nonmetals in the polyatomic ion
4. SO2 – Covalent bonding
5. Ca(ClO2)2 - Ionic bonding between the metal and nonmetal; Covalent bonding between the nonmetals in the polyatomic ion
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Polar vs. Nonpolar• Two general types of covalent bonds:
– Polar covalent• Unequal sharing (or a partial transfer) of
electrons
• Occurs when different elements are covalently bonded to one another
–Why different elements?»Because different elements have
different electronegativities
– Nonpolar covalent• Equal sharing (no transfer) of electrons
• Occurs only when all of the atoms in a molecule belong to the same element
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4
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Polar vs. Nonpolar (Figure 8.4)
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Bonds
Ionic Covalent
Polar Nonpolar
Complete Transfer
of Electrons
No Transfer
of Electrons
No Sharing of Electrons
Equal Sharing of Electrons
Increasing electron
transfer
Increasing equality of
sharing
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Polar vs. Nonpolar• Polar covalent bonds are:
– Typically shorter bonds – Stronger bonds due to their increased ionic
character
• Nonpolar covalent bonds are:– Typically longer bonds – Weaker bonds
• Polarity– Occurs in polar covalent molecules
– Polarity is the degree of transfer of electrons in a covalently bonded molecule composed of different element’s atoms.
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Electronegativity• Ability of an atom to attract bonding electrons
• Proposed by Linus Pauling in the early 1930’s
• A difference in electronegativity between the atoms in a covalent bond results in:– A polar covalent bond
– Increased ionic character
• The greater the difference in electronegativity, the greater the ionic character and the more polar the bond that joins the atoms.
– Decreased bond length and increased bond strength
• No difference in electronegativity between atoms in a covalent bond results in a nonpolar covalent bond.
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Electronegativity
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Figure 8.5
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Trends in Electronegativity
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Figure 8.6
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Practice – Polar Bonds
• Which of the following molecules
have polar bonds? If a bond is
polar, which atom has a partial
negative charge?
1. SO2
2. N2
3. PH3
4. CCl45. O3
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Practice Solutions – Polar Bonds
• Which of the following molecules have polar bonds? If a bond is polar, which atom has partial negative charge?1. SO2 – Polar covalent bonds ���� O is more
electronegative and has a partial negative charge
2. N2 – Nonpolar covalent bonds3. PH3 – Polar covalent bonds ���� P is more
electronegative and has a partial negative charge
4. CCl4 – Polar covalent bonds ���� Cl is more electronegative and has a partial negative charge
5. O3 – Nonpolar covalent bonds
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Ionic Bonding• Formation of ions and ionic bonds relates to an
element’s electron configuration.
• Each element immediately following a noble gas is a metal.– Metals lose electrons, forming a positive
charge, to become cations.• Each element immediately preceding a noble
gas is a nonmetal.
– Nonmetals gain electrons, forming a negative charge, to become anions.
• Therefore, elements (main-group) either lose or gain electrons to become isoelectronic with a noble gas (i.e. have the same electron configuration).
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18
Ionic Bonding
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Lewis Dot Symbols• Lewis Dot symbol
– Electron dot symbol
– Dots placed around an element’s symbol represent valence electrons
– Pair electrons as needed
– Octet rule
• Tendency of an atom to achieve an electron configuration having 8 valence electrons
– Same as the electron configuration of a noble gas
– The 8 electrons exist in 4 pairs
– Ions achieve 8 electrons by losing or gaining electrons
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20
Practice – Lewis Symbols for Ions
• Write the Lewis symbols for the beryllium and nitrogen ions. Then write a formula for the
compound that would form between them, using their Lewis symbols.
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Practice Solutions – Lewis Symbols for Ions
• Write the Lewis symbols for the beryllium and nitrogen ions. Then write a formula for the compound that would form between them, using their Lewis symbols.
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Be N
Be3N2
BeN
N
NBe
Be
N
Lewis Symbols for beryllium and nitrogen.
Lewis Structures for beryllium and nitrogen ions.
Compoundformula2
3
Be2
N3
N3
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Structures of Ionic Crystals• Crystal lattice
– The pattern obtained when an ion, represented as a charged sphere, exerts a force equally in all directions.
• Thus, ions of equal and opposite charge surround it.
– Cations and anions must come into contact for a crystal lattice to form.
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Figure 8.10
23
Structures of Ionic Crystals• Ionic crystal
– Ions are arranged in a regular geometric pattern that maximizes the attractive forces and minimizes the repulsive forces.
– Hard and brittle
– Can shatter if struck forcefully
• The charges and sizes of ions largely determine the characteristic patterns of ionic crystals
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Structures of Ionic Crystals
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Octet Rule• Octet rule
– Tendency of an atom to achieve an electron configuration having 8 valence electrons
• Same as the electron configuration of a noble gas
• Covalently bonded atoms achieve 8 valence electrons by sharing electrons
• The 8 electrons exist in 4 pairs
– H reacts to obtain a total of 2 electrons
like He.
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Covalent Bonding• Single covalent bond
– A covalent bond that consists
of a pair of electrons shared
by two atoms
– Each atom contributes one
electron to the bond
• The orbitals overlap to
allow the electron pair to be located around both
atoms
– Lewis formula
• The atoms are shown
separately and the valence
electrons are represented
by dots
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Figure 8.14
27
Covalent Bonding• Multiple covalent bonds
– Covalent bonds that consist of more than one pair of electrons shared by two atoms
– Double bond• Sharing of two pairs of
electrons (4 electrons total)• In Lewis Dot structures, a
double bond is represented by 4 dots or 2 parallel lines.
– Triple bond• Sharing of three pairs of
electrons (6 electrons total)
• In Lewis Dot structures, a triple bond is represented by 6 dots or 3 parallel lines.
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Figure 8.17
10
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Practice – Lewis Formulas
• Determine the formula of a simple compound
that follows the octet rule and is formed
from nitrogen and fluorine atoms. Use
electron dot structures to describe the
bonding in this compound.
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Figure 8.18
29
Practice Solutions – Lewis Formulas• Determine the formula of a simple compound that
follows the octet rule and is formed from nitrogen and fluorine atoms. Use electron dot structures to describe the bonding in this compound.
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N
Lewis Symbols for nitrogen and fluorine.
F
F
F
F
N N FF
FN FF
F
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Steps for Writing Lewis Dot Structures
1. Write an atomic skeleton.• The arrangement of atoms is usually symmetrical.• In a molecule of two different elements, the one with
the greater number of atoms usually surrounds the one with the lesser number of atoms.
• The central atom, the one surrounded by the other atoms, tends to be the one that is less electronegative and is present in the least quantity. This atom usually forms the greater number of bonds and is found further toward the bottom left side of the periodic table.
• Hydrogen atoms are generally on the outside of the molecule.
• The chemical formula may give clues about the arrangement of atoms.
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Steps for Writing Lewis Dot Structures (Cont’d)
2. Sum the valence electrons from each atom to get the total number of valence electrons.
3. Place two electrons, a single bond, between each pair of bonded atoms.
4. If you have not placed all the valence electrons in the formula, add any remaining electrons as unshared electron pairs, consistent with the octet rule.• Add pairs of electrons first to complete the octet of
atoms surrounding the central atom. Then add any remaining electrons in pairs to the central atom.
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32
Steps for Writing Lewis Dot Structures (Cont’d)
5. If necessary to satisfy the octet rule,
shift unshared electrons from non-
bonded position on atoms with
completed octets to positions
between atoms to make double or
triple bonds.
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33
Writing Lewis Dot Structures• Write a Lewis formula for the
formaldehyde, CH2O, molecule.
CH2O
• Write a Lewis formula for cyanic
acid.
HCN
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12
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Writing Lewis Dot Structures1. Write an atomic skeleton for CH2O:
2. Sum the valence electrons from each atom to get the total number of valence electrons.
Carbon is in Group IVA (14), so it has 4 valence electrons. Each hydrogen contributes 1 valence electron (H is in Group IA (1)). Oxygen contributes 6 valence electrons because it is in Group VIA (16).
Total number of valence electrons = 4 + (1 x 2) + 6 = 12
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CH H
O
35
Writing Lewis Dot Structures3. Next, bond the electrons around each atom in a
single bond first, then use double bonds as necessary.
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CHH
O
C
O
H HC
O
H H
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Writing Lewis Dot Structures1. Write an atomic skeleton for HCN:
H C N
2. Sum the valence electrons from each atom to get the total number of valence electrons.
Carbon is in Group IVA (14), so it has 4 valence electrons. Hydrogen contributes 1 valence electron (H is in Group IA (1)). Nitrogen contributes 5 valence electrons because it is in Group VA (15).
Total number of valence electrons = 4 + 1 + 5 = 10
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Writing Lewis Dot Structures3. Next, bond the electrons around each
atom in a single bond first, then use double and triple bonds as necessary.
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CH
C
N
NHC NH C NH
38
Resonance Structures• When there are several equally valid
arrangements of bonding (i.e. Lewis Dot structures), then the concept of resonance helps explain why.
• Resonance– The electron arrangement in molecules with
several equally valid Lewis Dot structures is represented by them all, each showing a different arrangement of the true arrangement of electrons.
• Resonance hybrid– Representation of the actual molecule– A composite of the formulas drawn
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Practice - Resonance Structures1. Write an atomic skeleton for N2O:
N N O
2. Sum the valence electrons from each atom to get the total number of valence electrons.
Nitrogen is in Group VA (15), so it has 5 valence electrons. Oxygen contributes 6 valence electrons because it is in Group VIA (16).
Total number of valence electrons = (5 x 2) + 6 = 16
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Practice - Resonance Structures3. Next, bond the electrons around each atom in a
single bond first, then use double bonds as necessary.
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N N O
N N ON N O
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Practice - Resonance Structures4. To draw resonance structures, rearrange
the electrons (and bonds) in the structures.
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N N O N N O
42
Exceptions to the Octet Rule• Incomplete octets
– The central atom has less than eight electrons around it.
– Ex. BH3
• Expanded octets– The central atom has greater than eight electrons
around it.
– Ex. PH5, SF6
• Odd-numbered Lewis Dot structures– The total number of electrons is odd.
– The central atom has an odd number of electrons around it.
– Ex. NO, NO2, ClO28 -
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Carbon Compounds
• Carbon has: – Four valence electrons
– The ability to form four bonds
– The ability to catenate, i.e. bond to itself
– Very strong bonds when bonded to itself
• Carbon molecules are ubiquitous in nature.
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44
Hydrocarbons
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Figure 8.20
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Hydrocarbons• Compounds containing hydrogen and carbon
• Aliphatic hydrocarbons
– A class in which the bonds are all localized single, double, and triple bonds
– Alkanes
• Hydrocarbons which contain only carbon-carbon single bonds
– Alkenes
• Hydrocarbons which contain at least one carbon-carbon double bond
– Alkynes
• Hydrocarbons which contain at least one carbon-carbon triple bond
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Hydrocarbons• Aromatic hydrocarbons
– A class of hydrocarbons which has carbon
atoms arranged in a six-atom ring with alternating single and double bonds
– Delocalized structures
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Figure 8.22
47
Functional Groups• Functional group
– A group that is introduced into or substituted in a hydrocarbon chain
– Gives the hydrocarbon its characteristic properties
– The group has a heteroatom, an atom other C and H• Typically O, S, and N
– Alcohol• A hydroxyl group (-OH)
replaces a hydrogen atom in the formula for a hydrocarbon
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Figure 8.23
48
Table 8.4 Functional Groups in Hydrocarbons
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H3C-NH2Methyl amine-N-Amine
Ethyl acetate-C-O-Ester
Acetic acid-C-OHCarboxylic
Acid
Acetone-C-Ketone
Acetaldehyde-C-HAldehyde
H5C2-O-C2H5Diethyl ether-O-Ether
C2H5-OHEthyl alcohol-OHAlcohol
FormulaExampleFunctional
GroupClass
17
49
Odors and Carbon Compounds
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50
Shapes of Molecules
• The relative locations of electron pairs around a central atom play a large role in determining a molecule’s 3-D shape.
• Negatively charged electrons repel one another, so electron pairs in different orbitals stay as far apart as possible.
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51
Shapes of Molecules• Valence shell electron pair repulsion
(VSEPR) theory– The tendency of electron pairs to
adjust the orientation of their orbitals to maximize the distance between them
– The bonded atoms and unshared pairs are arranged around the central atom as far apart as possible
• Bond angle– A shape is characterized by a bond
angle between the central atom and the atoms bonded to it
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52
VSEPR Parent Structures
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Table 8.5
53
VSEPR Derivative Structures
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54
Steps for VSEPR Structures1. Draw a Lewis formula.
2. Count the number of atoms bonded to the central atom, and count unshared pairs on the central atom.
3. Add the number of atoms and the number of unshared electron pairs around the central atom. The total indicates the parent structure.
4. The molecular shape is derived from the parent shape by considering only the positions in the structure occupied by bonded atoms.
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Practice – VSEPR Structures
• What is the shape of the nitrite
ion (NO2-1)? What is the O-N-O
bond angle?
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Practice Solutions – VSEPR Structures
1. Draw the Lewis Dot Structure.
2. Count the number of atoms bonded to the central atom and count unshared pairs on the central atom.
N, which is the central atom in this case, has 2 atoms bonded to it, and 1 unshared pair on it.
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NO
ONO O
-1
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Practice Solutions – VSEPR Structures
3. If we add the number of atoms and unshared pairs around the central atom, we get the number 3. This indicates that the parent structure is trigonal planar.
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A
B
BB
120
20
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Practice Solutions – VSEPR Structures
4. The derived structure, which only considers bonded atoms, is called bent
(or angular) and has a bond angle equal to 118°, since we take 2°off the parent
structure’s bond angle to obtain the
derived structure’s bond angle.
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ON
O118
59
Natural Applications of VSEPR Theory
• Molecular shapes are important in living systems.
• Glycine molecules are typically found in proteins
or gelatins.
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Figure 8.29
60
Natural Applications of VSEPR Theory
• Heme molecule
– Oxygen is carried throughout the body via red blood cells containing heme molecules.
– Histidine, an amino acid in the heme molecule, just fits into the space next to the oxygen molecule.
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Polarity of Molecules• Diatomic molecules
– Polarity lies along the plane of the bond
• Polyatomic molecules
– A nonpolar molecule is one that has all nonpolar bonds or one that has polar bonds that cancel out
• Bonds that cancel out have equal polarities in opposite directions
– This happens when:
» A central atom has no unshared electrons
» The atoms around the central atom all have the same electronegativity
– A polar molecule is one that has polar bonds that DO NOT cancel out
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62
Polarity of Molecules
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Figure 8.32
63
Practice – Polarity of Molecules
• Predict whether C2H6, NO2, CO2, SO2, and SO3 are polar or nonpolar molecules.
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Practice Solutions – Polarity of Molecules
• Predict whether C2H6, NO2, CO2, SO2, and SO3 are polar or nonpolar molecules.– C2H6 ���� tetrahedral at each C ����
nonpolar
– NO2 ���� bent ���� polar
– CO2 ���� linear ���� nonpolar
– SO2 ���� bent ���� polar
– SO3 ���� trigonal planar ���� nonpolar
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“Like” Dissolves “Like”• Ionic salts and
polar liquids dissolve better in
polar liquids than in nonpolar liquids
• Nonpolar liquids dissolve better in
other nonpolar liquids than in polar
liquids
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Figure 8.34