Chapter 9
Chemical Bonding
Section 9.1: Why does bonding occur in the first place?
Bonding lowers the potential energy between positive and negative particles (p341).
What is potential energy?
1 type of Potential Energy: Gravitational P.E.
m
Ball On Top of a Hill
P.E. = mgh
h
Energy changes forms:
P.E.  Kinetic Energy (K.E.)
Section 9.1: Why does bonding occur in the first place?
Bonding lowers the potential energy between positive and negative particles (p341).
Energy changes forms
Mechanical
Energy
Friction
Engines
Motor
Generator
Electrical
Energy
Heat (Thermal)
Energy
FIre
Solar
Heater
Light (Radiant)
Energy
Battery
Battery
Charger
Chemiluminescence
Photosynthesis
Chemical
Energy
Section 9.1: Why does bonding occur in the first place?
Bonding lowers the potential energy between positive and negative particles (p341).
When chemical bonds form: Chemical P.E. changes to Heat Energy & Light Energy
Mechanical
Energy
Heat (Thermal)
Energy
Light (Radiant)
Energy
Electrical
Energy
Chemical
Energy
Section 9.1: Why does bonding occur in the first place?
Bonding lowers the potential energy between positive and negative particles (p341).
Energy changes forms: Chemical P.E.  Heat & Light Energy
http://chemsite.lsrhs.net/chemKinetics/PotentialEnergy.html
Section 9.1: Three Type of Bonds
Ionic bonding: Metal + Nonmetal (Valence e- transferred)
Covalent bonding: Nonmetal + Nonmetal (Valence e- shared)
Metallic bonding: Metal + Metal (“Sea” of e-)
http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch10/non.php
Concept Check
Review: Valence Electrons – e- involved in forming compounds (Ch 8, p315)
Boron (B)
How many valence e-?
How many needed
for full outer shell?
Total valance e-:
Magnesium (Mg)
Hydrogen (H)
Section 9.1: Two Bond Types With Localized Electrons
Ionic & Covalent Bonding
Representative Elements
For REPRESENTATIVE elements:
• period (row) = shell # (n = 1, 2, 3, 4….n)
• group (column) = # of e- in outer shell
Group #
# of valence e-
Transition Elements
IA
1
IIA
IIIA
2
3
IVA
4
VA
5
VIA
6
VIIA
7
VIIA
8
Shells
of an
atom
Section 9.1: Two Bond Types With Localized Electrons
Ionic & Covalent Bonding:
Why do ionic bonds form instead of covalent bonds, and vice versa?
“Bonding Continuum”
Covalent Bonds
Ionic Bonds
nonmetals + nonmetal
metal + nonmetal
Nonpolar Covalent Bond
Polar Covalent Bond
Electrons are shared unequally.
Ionic Bond
Electrons are transferred.
Extent of electron sharing in Covalent Bonds
e-’s shared between atoms
of the same element:
Equal Sharing
e-’s shared between atoms
of different elements:
Unequal Sharing
Unequal sharing – occurs because one of the atoms in a bond has a stronger attraction
for the pair of e-’s than does the other atom
Why does one atom have a stronger attraction for e-?
Electronegativity
Definition: electronegativity (E.N) – the ability of an atom to attract the shared electrons
Decreasing E.N.
Increasing E.N.
Rule for Bond Formation
The atom with the greater E.N. pulls the shared electrons closer to its nucleus
resulting in (1) – charge on high E.N. atom
(2) + charge on low E.N. atom
More later: Section 9.5
Why do ionic bonds form instead of covalent bonds, and vice versa?
“Bonding Continuum”
Covalent Bonds
Ionic Bonds
e- sharing
2 nonmetals
e- transfer
metal + nonmetal
Nonpolar Colvalent
Polar Colvalent
0.4 < E.N.
0.4 < E.N. < 1.7
1.7
E.N. difference > 1.7
Answer: Electronegativity Differences
Example:
Oxygen (O) bonds with
Magnesium (Mg): MgO
E.N. of O = 3.5
E.N. of Mg = 1.2
E.N. difference = 2.3
Section 9.1: “The Other” Bond Type With Delocalized Electrons
Metallic Bonding
Metallic Bonding - Delocalized
A messy “sea” of electrons
Covalent Bonding, Ionic Bonding
- Delocalized
Electrons fit neatly into shells.
Section 9.1: “The Other” Bond Type With Delocalized Electrons
Metallic Bonding
Metallic Bonding - Delocalized
A messy “sea” of electrons
Outer e-
Inner e-
Lewis Electron-Dot Symbols
Two parts:
(1) Element symbol – nucleus + inner electrons
Ex: The element lithium has an element symbol Li
(2) Surrounding dots – valence electrons (outer most shell)
Different elements can have the same number of dots
Be
Mg
Same Group
(Column)
Li
Review: Ions
Ion – charged particles that form when an atom gains or loses one or more electrons
(Ch2, p60)
Element
Ion
Ion Type
Mg
Mg2+
Cation
Cl
Cl-
Anion
Review: Electron Configuration and Orbital Diagrams (Ch8, p304-317)
Example:
Mg
Concept Check
• End of Chapter Problems in-class (for now):
9.7, 9.9, 9.13, 9.15
Write the ion for the following elements: K, Br, Sr, Ar, O
For example, the ion for Mg is Mg2+.
• Suggested Optional Practice Problems (for outside of class):
9.6, 9.8, 9.10, 9.12, 9.14 (Answers in back of book or online)
Section 9.2: Ionic Bonding
Central idea: Electrons are transferred from metal atoms to nonmetal atoms to form
ions that come together in a solid ionic compound.
Solid Ionic compound
Na – metal
Cl - nonmetal
Sodium chloride (NaCl)
Contrast with molecules formed during covalent bonding (more later).
Examples: Water (H2O)
O
H
Carbon Dioxide (CO2)
O
H
C
O
Section 9.2: Ionic Bonding
Rule: The total number of e- lost by the metal atom equals the total number
gained by the nonmetal atom.
Cl-
Na+
lost
gained
Behavior of Ionic Compounds
Why is the melting point of MgO higher than the melting point of KCl?
Lattice Energy
(∆Hºlattice)
Section 9.2: Lattice Energy
Definition – The enthalphy change that occurs when 1 mol of ionic solid
separates into gaseous ions.
For Review of Enthalpy: Ch6, p243
Lattice Energy denoted as: ∆Hºlattice
∆Hºlattice cannot be measured directly, BUT it can be calculate using the:
Born-Haber cycle
Section 9.2: Born-Haber Cycle
Uses Hess’s Law: Total enthalpy of an overall reaction is the sum of the enthalpy
changes of individual reactions. (∆Htotal = ∆Hrxn1 + ∆Hrxn2 +……….)
*Not actual
steps.
Section 9.2: Trends in Lattice Energy
Coulomb’s Law (Ch2)
Section 9.2: Trends in Lattice Energy
Behavior of Ionic Compounds
So, why is the melting point of MgO higher than the melting point of KCl?
Concept Check
• End of Chapter Problems in-class (for now):
9.27, 9.30
• Suggested Optional Practice Problems (for outside of class):
9.26, 9.28 (Answers in back of book)
Problem 9.30
Section 9.3: Covalent Bonding
e- sharing – primary way that atoms interact
Nonmetal + Nonmetal
Examples: Water (H2O)
O
H
Carbon Dioxide (CO2)
O
H
C
Organic Compounds
O
H
H
H
H
C
C
C
H
H
H
H
Contrast with ionic solids formed during ionic bonding (discussed previously).
Na – metal
Cl - nonmetal
Sodium chloride (NaCl)
Section 9.3: Covalent Bonding
Why do covalent bonds form?
Lower P.E. = More stable
Section 9.3: Covalent Bonding
How are the electrons distributed?
Electron
Density
In order for each atom to have a full outer shell (2 e- for H, He; 8 e- for others),
the electrons arrange themselves in certain configurations:
• Bonding Pairs & Lone Pairs
• Bond Type – double, single, triple
Section 9.3: Covalent Bonding
Bond Energy (B.E.) – aka Bond Enthalpy or Bond Strength
Covalent Bond Strength – depends on strength of attraction between nuclei and
shared electrons
Bond Energy – energy needed
to overcome attraction and
break the bond
Section 9.3: Covalent Bonding
Bond Energy (B.E.)
Bond formation is exothermic:
∆Hº always +
Bond breakage is endothermic: ∆Hº always Absolute value of B.E. – Each bond has its own unique B.E. due to variations in:
(1) e- density
(2) charge
(3) atomic radii
Section 9.3: Covalent Bonding
Strength of Bond different than E required to pull atoms apart (B.E.)
Less E needed
to break.
Lower B.E.
Weaker Bonds =
Higher Energy
“Shallow Energy Well”
Stronger Bonds =
Lower Energy
“Deeper Energy Well”
More E needed
to break.
Higher B.E.
Section 9.3: Covalent Bonding
Bond Energy (B.E.) and Bond Length
Bond Length – sum of the radii of the bonded atoms
(analogous to distance in Coloumb’s Law)
At minimum E point.
Section 9.3: Covalent Bonding
Bond Energy (B.E.) and Bond Length
300
R2 = 0.3155
Bond Length
250
200
150
This relationship
holds, in general,
ONLY for single
bonds.
100
50
0
150
250
350
Bond Energy
450
550
Section 9.3: Covalent Bonding
Bond Type (Single, Double, Triple) also matters
Same two
elements,
different B.E.
Nuclei more attracted to 2 shared pairs of e- than one shared pair of e-.
Higher bond order = Shorter bond length = Higher Bond Energy
Section 9.3: Covalent Bonding
Periodic Table Trends Without Detailed Bond Lengths
The closer the
atoms, the
stronger the bond.
Bond Energy:
C—F > C—Cl > C—Br
Section 9.3: Covalent Bonding
Covalent Bonds are stronger than Ionic Bonds
So why, then, do covalent compounds have lower melting points
than ionic compounds?
Example: CCl4 m.p. = -23 ºC
Strong covalent bonding forces
Hold atoms together
NaCl m.p. = 800 ºC
Weak intermolecular forces
Hold molecules together
(More in Chapter 12)
O
+
Chemical Reaction
H
H
+
solid  liquid 
Phase Change
gas
Section 9.4: Bond Energy and Chemical Change
Where does the heat that is released come from?
http://chemsite.lsrhs.net/chemKinetics/PotentialEnergy.html
Section 9.4: Bond Energy and Chemical Change
Total energy of a chemical system = K.E. + P.E.
Example of a chemical system
A container filled with molecules.
Kinetic Energy (K.E.)
Three types:
(1) Vibrational
(2) Rotational
(3) Translational
• Does not change during
chemical reaction (depends on T).
Changes during a Phase Change
(Chapter 12).
http://www.landfood.ubc.ca/courses/fnh/301/water/motion.gif
solid  liquid 
gas
Section 9.4: Bond Energy and Chemical Change
This leaves us with changes in P.E. during chemical reactions.
P.E. contributions can from electrostatic forces between:
Separate Vibrating Atoms
Nucleus & Electrons in Atoms
Protons & Neutrons in Nucleus
Nuclei and Shared Electron Pair in Each Bond = Bond Energy
Where does the heat that is released come from?
The energy released or absorbed during a chemical change is due to the
differences between the reactant bond energies and the product bond energies.
B.E.reactants - B.E.products = Heat
Section 9.4: Bond Energy and Chemical Change
Heat of reaction, ∆Hºrxn
Exothermic reaction:
- ∆Hºrxn
Endothermic reaction: + ∆Hºrxn
∆Hºrxn = ∆Hºreactant bonds broken + ∆Hºproduct bonds formed
∆Hºrxn = ∆BEreactant bonds broken – ∆BEproduct bonds formed
Analogous to ionic compound formation:
Lattice Energy, ∆Hºlattice
Born-Haber cycle
(∆Hºtotal = ∆Hºrxn1 + ∆Hºrxn2 +……+ ∆Hºlattice)
Section 9.4: Bond Energy and Chemical Change
Example: H2 + F2  2 HF
Weaker Bonds
Less Stable, More Reactive
H2 and F2
Stronger Bond
More Stable, Less Reactive
HF
Section 9.4: Bond Energy and Chemical Change
Another way to looks at this reaction:
H2 + F2  2 HF
Heat of reaction, ∆Hºrxn
2H +2F
H2 + F2
HF
∆Hºrxn = ∆Hºreactant bonds broken + ∆Hºproduct bonds formed
Section 9.4: Bond Energy and Chemical Change
Use bond energies to calculate ∆Hºrxn (Table 9.2)
H2 + F2  2 HF
9.39, 9.47, 9.49
Optional Homework Problems: 9.38, 9.46, 9.48, 9.50
Section 9.4: Bond Energy and Chemical Change
Application: Energy Released From Combustion of Fuel
∆Hºrxn = ∆BEreactant bonds broken – ∆BEproduct bonds formed
Energy Released = B.E.(fuel + O2) – B.E.(CO2 + H2O)
Fuels with more weak bonds yield more energy than fuels with fewer weak bonds.
Carbs:
Food fuels the body: More
O-H
C-O
Fats:
More
C-H
C-C
Section 9.5: Between the Extremes
Scientific models are idealized descriptions of reality.
“Bonding Continuum”
Covalent Bonds
Ionic Bonds
e- sharing
2 nonmetals
e- transfer
metal + nonmetal
Nonpolar Colvalent
Polar Colvalent
0.4 < E.N.
0.4 < E.N. < 1.7
1.7
E.N. difference > 1.7
Electronegativity – the relative ability of a bonded atom to attract the shared e-
Section 9.5: Between the Extremes
Electronegativity – inversely related to atomic size (radius)
atomic size
E.N.
WHY?
Section 9.5: Between the Extremes
Nonmetals are more electronegative than metals.
Section 9.5: Between the Extremes
Electronegativity and Oxidation Number (O.N.)
(Review of O.N.: Section 4.5)
Oxidation-reduction (redox) reactions: The net movement of electrons from one
reactant to the other.
Oxidation – the loss of e- (LEO), Reduction – the gain of e- (GER)
“LEO the lions says GER!”
Oxidizing agent – becomes reduced; Reducing agent – becomes oxidized
Which element is oxidized? Reduced? Which is the oxidizing agent? Reducing agent?
Oxidation Number and Electronegativity
When dead organisms (such as plankton) fall to the bottom of the sea, their
dead bodies are eaten (respiration) by bacteria living in the ocean sediments:
CH2O + O2  CO2 + H2O
What might be a problem for bacteria trying to eat CH2O deep in sediments?
In addition to O2: SO42- and NO32- are present in the sediments.
Which might they use?
Section 9.5: Between the Extremes
Electronegativity and Oxidation Number (O.N.)
E.N. is used to determine an atom’s O.N. in a given bond.
(1) The more E.N. atom in a bond is assigned ALL the SHARED e-; The less
E.N. atoms is assigned NONE
Example: HCl
Cl: 8
H: 0
(2) O.N. = # valence e- - # shared eExample:
O.N.Cl = 7 – 8 = -1
O.N.H = 1 – 0 = +1
Section 9.5: Between the Extremes
Polar Covalent Bonds
“Bonding Continuum”
Covalent Bonds
Ionic Bonds
e- sharing
2 nonmetals
e- transfer
metal + nonmetal
Nonpolar Colvalent
Polar Colvalent
0.4 < E.N.
0.4 < E.N. < 1.7
1.7
E.N. difference > 1.7
This bond type is indicated by:
(1) polar arrow (
) pointing toward negative pole H–F
(2) delta symbol ()
O
+
H
H
+
Section 9.5: Between the Extremes
Polar Covalent vs. Nonpolar Covalent
“Bonding Continuum”
Covalent Bonds
Ionic Bonds
e- sharing
2 nonmetals
e- transfer
metal + nonmetal
Nonpolar Colvalent
Polar Colvalent
0.4 < E.N.
0.4 < E.N. < 1.7
1.7
E.N. difference > 1.7
Section 9.5: Between the Extremes
Partial Ionic Character – related directly to the electronegativity difference (∆EN)
Why?
A greater ∆EN results in larger partial charges () and a higher partial ionic character.
Example: HCl, LiCl, Cl2
Arrange these compounds in order of least to most partial ionic character.
Section 9.5: Between the Extremes
Two approaches for getting a sense of a compound’s ionic character:
#1: Arbitrary cutoffs used in bonding continuum.
“Bonding Continuum”
Covalent Bonds
Ionic Bonds
e- sharing
2 nonmetals
e- transfer
metal + nonmetal
Nonpolar Colvalent
Polar Colvalent
0.4 < E.N.
0.4 < E.N. < 1.7
1.7
E.N. difference > 1.7
Section 9.5: Between the Extremes
Two approaches for getting a sense of a compound’s ionic character:
#2: Calculate the percent ionic character (increases with ∆EN)
Compare actual behavior of a polar molecule in an electric field with the
behavior it would show if the e- were completely transferred (pure ionic).
50 % is dividing line.
Notice: Cl2 is 0% ionic, but no molecule has 100 % ionic character (e- sharing
occurs to some extent in every bond.
Section 9.5: Between the Extremes
Notice, now: Why metal that bond with nonmetals form ionic bonds.
Why nonmetals that bond with other nonmetals form covalent bonds.
Section 9.5: Between the Extremes
Properties of substances are indicative of their ionic or covalent character.
Section 9.6: Metallic Bonding (More in Chap 12)
Electron Sea Model
In reactions with nonmetals, metals (Na) transfer their outer e- to form ionic solids (NaCl).
What holds together bonded metals (Na)? All metal atoms contribute their valence e-,
which are shared among all the atoms in a sample.
Metallic Bonding - Delocalized
A messy “sea” of electrons
Covalent Bonding, Ionic Bonding
- Localized
Electrons fit neatly into shells.
Alloys - more than one metal element involved in a metallic “sea”
Section 9.6: Metallic Bonding (More in Chap 12)
Properties of metal substances are explained by the electron sea model.
Most metals are solids.
High m.p. = attractions b/w cations and anions need not be broken
Much higher b.p. = attractions b/w cations and anions broken
m.p. depends on # of valence e-:
Problems for today
9.62, 9.64, 9.66
What would you expect the B.E. of a H–F bond to be given that:
H–H = 432 kJ/mol
F–F = 159 kJ/mol
?