Chemistry: A Molecular Approach, 1st Ed.
Nivaldo Tro
Roy Kennedy
Massachusetts Bay Community College
Wellesley Hills, MA
2008, Prentice Hall
Why
Do
Atoms
Bond?
 processes are spontaneous if they result in a system with
lower potential energy
 chemical bonds form because they lower the potential
energy between the charged particles that compose atoms
 the potential energy between charged particles is directly
proportional to the product of the charges
 the potential energy between charged particles is inversely
proportional to the distance between the charges
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Potential Energy Between
Charged Particles
1  q1  q2 
E potential 


4 0  r 
 0 is a constant
 = 8.85 x 10-12 C2/J∙m
 for charges with the same sign, Epotential is + and the
magnitude gets less positive as the particles get farther
apart
 for charges with the opposite signs, Epotential is  and the
magnitude gets more negative as the particles get closer
together
 remember: the more negative the potential energy, the
more stable the system becomes
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Potential Energy Between
Charged Particles
The attraction
repulsion between likecharged particles increases
opposite-charged
particles as
the particles
increases
as the
get particles
closer together.
get
To bring
closer
together.
them closer
Bringing
requires
them
the addition
closer
lowersofthemore
potential
energy.
energy of the system.
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Bonding
 a chemical bond forms when the potential energy of the
bonded atoms is less than the potential energy of the separate
atoms
 have to consider following interactions:
 nucleus-to-nucleus repulsion
 electron-to-electron repulsion
 nucleus-to-electron attraction
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Types of Bonding
6
Ionic Bonds
 when metals bond to nonmetals, some electrons from the
metal atoms are transferred to the nonmetal atoms
 metals have low ionization energy, relatively easy to remove an
electron from
 nonmetals have high electron affinities, relatively good to add
electrons to
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Covalent
Bonds
 nonmetals have relatively high ionization energies, so it is
difficult to remove electrons from them
 when nonmetals bond together, it is better in terms of
potential energy for the atoms to share valence electrons
 potential energy lowest when the electrons are between the
nuclei
 shared electrons hold the atoms together by attracting
nuclei of both atoms
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Conductivity of NaCl
in NaCl(s), the ions are
stuck in position and not
allowed to move to the
charged rods
in NaCl(aq), the ions
are separated and
allowed to move to the
charged rods
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Lewis Theory and Ionic Bonding
 Lewis symbols can be used to represent the transfer of
electrons from metal atom to nonmetal atom, resulting in
ions that are attracted to each other and therefore bond

Li

 + 
F


 F 



  




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Li +
1
10
Predicting Ionic Formulas
Using Lewis Symbols
 electrons are transferred until the metal loses all its
valence electrons and the nonmetal has an octet
 numbers of atoms are adjusted so the electron transfer
comes out even
Li



Li



 O 



  

O


2 Li +

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2
Li2O
11
Energetics of Ionic Bond
Formation
 the ionization energy of the metal is endothermic
 Na(s) → Na+(g) + 1 e ─ DH° = +603 kJ/mol
 the electron affinity of the nonmetal is exothermic
 ½Cl2(g) + 1 e ─ → Cl─(g) DH° = ─ 227 kJ/mol
 generally, the ionization energy of the metal is larger than
the electron affinity of the nonmetal, therefore the
formation of the ionic compound should be endothermic
 but the heat of formation of most ionic compounds is
exothermic and generally large; Why?
 Na(s) + ½Cl2(g) → NaCl(s)
DH°f = -410 kJ/mol
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Ionic
Bonds
 electrostatic attraction is nondirectional!!
 no direct anion-cation pair
 no ionic molecule
 chemical formula is an empirical formula, simply giving the ratio
of ions based on charge balance
 ions arranged in a pattern called a crystal lattice
 every cation surrounded by anions; and every anion surrounded
by cations
 maximizes attractions between + and - ions
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Lattice Energy
 the lattice energy is the energy released when the solid
crystal forms from separate ions in the gas state
 always exothermic
 hard to measure directly, but can be calculated from
knowledge of other processes
 lattice energy depends directly on size of charges and
inversely on distance between ions
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Born-Haber Cycle
 method for determining the lattice energy of an ionic substance
by using other reactions
 use Hess’s Law to add up heats of other processes
 DH°f(salt) = DH°f(metal atoms, g) + DH°f(nonmetal atoms, g) +
DH°f(cations, g) + DH°f(anions, g) + DH°f(crystal lattice)
 DH°f(crystal lattice) = Lattice Energy
 metal atoms (g)  cations (g), DH°f = ionization energy

don’t forget to add together all the ionization energies to get to the
desired cation

M2+ = 1st IE + 2nd IE
 nonmetal atoms (g)  anions (g), DH°f = electron affinity
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Born-Haber Cycle for NaCl
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Trends in Lattice Energy
Ion Size
 the force of attraction between charged particles is
inversely proportional to the distance between them
 larger ions mean the center of positive charge (nucleus of
the cation) is farther away from negative charge (electrons
of the anion)
 larger ion = weaker attraction = smaller lattice energy
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Lattice Energy vs.
Ion Size
Lattice Energy
Metal Chloride
(kJ/mol)
LiCl
-834
NaCl
-787
KCl
-701
CsCl
-657
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Trends in Lattice Energy
Ion Charge
 the force of attraction between
oppositely charged particles is directly
proportional to the product of the
charges
Lattice Energy =
-910 kJ/mol
 larger charge means the ions are more
strongly attracted
 larger charge = stronger attraction =
larger lattice energy
 of the two factors, ion charge
generally more important
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Lattice Energy =
-3414 kJ/mol
19
Ionic Bonding
Model vs. Reality
 ionic compounds have high melting points and boiling
points
 MP generally > 300°C
 all ionic compounds are solids at room temperature
 because the attractions between ions are strong,
breaking down the crystal requires a lot of energy
 the stronger the attraction (larger the lattice energy), the
higher the melting point
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Ionic Bonding
Model vs. Reality
 ionic solids are brittle and hard
 the position of the ion in the crystal is critical to establishing
maximum attractive forces – displacing the ions from their
positions results in like charges close to each other and the
repulsive forces take over
+
-
+
-+
+
-
-
+
-+ +- -+
+
+
+
+
- +
- + -
+
-+
+
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-+ +- -+ +
+
+
+
- +
- +
- + -
21
Ionic Bonding
Model vs. Reality
 ionic compounds conduct electricity in the liquid state or
when dissolved in water, but not in the solid state
 to conduct electricity, a material must have charged
particles that are able to flow through the material
 in the ionic solid, the charged particles are locked in
position and cannot move around to conduct
 in the liquid state, or when dissolved in water, the ions
have the ability to move through the structure and
therefore conduct electricity
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Covalent Bonding
Model vs. Reality
 molecular compounds have low melting points and
boiling points
 MP generally < 300°C
 molecular compounds are found in all 3 states at room
temperature
 melting and boiling involve breaking the attractions
between the molecules, but not the bonds between the
atoms
 the covalent bonds are strong
 the attractions between the molecules are generally weak
 the polarity of the covalent bonds influences the strength of the
intermolecular attractions
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Intermolecular Attractions vs. Bonding
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Ionic Bonding
Model vs. Reality
 some molecular solids are brittle and hard, but many are soft
and waxy
 the kind and strength of the intermolecular attractions varies
based on many factors
 the covalent bonds are not broken, however, the polarity of
the bonds has influence on these attractive forces
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Ionic Bonding
Model vs. Reality
 molecular compounds do not conduct electricity in the
liquid state
 molecular acids conduct electricity when dissolved in
water, but not in the solid state
 in molecular solids, there are no charged particles around
to allow the material to conduct
 when dissolved in water, molecular acids are ionized, and
have the ability to move through the structure and
therefore conduct electricity
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Bond
Energies
 chemical reactions involve breaking bonds in reactant
molecules and making new bond to create the products
 the DH°reaction can be calculated by comparing the cost of
breaking old bonds to the profit from making new bonds
 the amount of energy it takes to break one mole of a
bond in a compound is called the bond energy
 in the gas state
 homolytically – each atom gets ½ bonding electrons
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Trends in Bond Energies
 the more electrons two atoms share, the stronger the covalent
bond
 C≡C (837 kJ) > C=C (611 kJ) > C−C (347 kJ)
 C≡N (891 kJ) > C=N (615 kJ) > C−N (305 kJ)
 the shorter the covalent bond, the stronger the bond
 Br−F (237 kJ) > Br−Cl (218 kJ) > Br−Br (193 kJ)
 bonds get weaker down the column
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Using Bond Energies to Estimate DH°rxn
 the actual bond energy depends on the surrounding atoms
and other factors
 we often use average bond energies to estimate the
DHrxn
 works best when all reactants and products in gas state
 bond breaking is endothermic, DH(breaking) = +
 bond making is exothermic, DH(making) = −
DHrxn = ∑ (DH(bonds broken)) + ∑ (DH(bonds formed))
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31
Estimate the Enthalpy of the Following
Reaction
H H
O O
H O O
+
H
32
Estimate the Enthalpy of the Following
Reaction
H2(g) + O2(g)  H2O2(g)
reaction involves breaking 1mol H-H and 1 mol O=O
and making 2 mol H-O and 1 mol O-O
bonds broken (energy cost)
(+436 kJ) + (+498 kJ) = +934 kJ
bonds made (energy release)
2(464 kJ) + (142 kJ) = -1070
DHrxn = (+934 kJ) + (-1070. kJ) = -136 kJ
(Appendix DH°f = -136.3 kJ/mol)
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Bond Lengths
 the distance between the nuclei of bonded
atoms is called the bond length
 because the actual bond length depends on
the other atoms around the bond we often
use the average bond length
 averaged for similar bonds from many
compounds
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Trends in Bond Lengths
 the more electrons two atoms share, the shorter the
covalent bond
 C≡C (120 pm) < C=C (134 pm) < C−C (154 pm)
 C≡N (116 pm) < C=N (128 pm) < C−N (147 pm)
 decreases from left to right across period
 C−C (154 pm) > C−N (147 pm) > C−O (143 pm)
 increases down the column
 F−F (144 pm) > Cl−Cl (198 pm) > Br−Br (228 pm)
 in general, as bonds get longer, they also get weaker
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Bond Lengths
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Metallic
Bonds
 low ionization energy of metals allows them to lose electrons
easily
 the simplest theory of metallic bonding involves the
metals atoms releasing their valence electrons to be
shared by all to atoms/ions in the metal
 an organization of metal cation islands in a sea of electrons
 electrons delocalized throughout the metal structure
 bonding results from attraction of cation for the delocalized
electrons
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Metallic Bonding
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Metallic Bonding
Model vs. Reality
 metallic solids conduct electricity
 because the free electrons are mobile, it allows the
electrons to move through the metallic crystal and
conduct electricity
 as temperature increases, electrical conductivity decreases
 heating causes the metal ions to vibrate faster, making it
harder for electrons to make their way through the crystal
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Metallic Bonding
Model vs. Reality
 metallic solids conduct heat
 the movement of the small, light electrons through the
solid can transfer kinetic energy quicker than larger
particles
 metallic solids reflect light
 the mobile electrons on the surface absorb the outside
light and then emit it at the same frequency
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Metallic Bonding
Model vs. Reality
 metallic solids are malleable and ductile
 because the free electrons are mobile, the direction of the
attractive force between the metal cation and free
electrons is adjustable
 this allows the position of the metal cation islands to
move around in the sea of electrons without breaking the
attractions and the crystal structure
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Metallic Bonding
Model vs. Reality
 metals generally have high melting points and boiling
points
 all but Hg are solids at room temperature
 the attractions of the metal cations for the free electrons
is strong and hard to overcome
 melting points generally increase to right across period
 the charge on the metal cation increases across the period,
causing stronger attractions
 melting points generally decrease down column
 the cations get larger down the column, resulting in a
larger distance from the nucleus to the free electrons
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