Ionic Bonding

Lets look at the formation of the prototypical ionic compound, NaCl.

We could first write the reaction using Lewis dot symbols, as shown in equation 8.2 (page 257) in your text.

Na(1) + Cl(7) Na(0)+ + Cl(8)-

[Here I have used a text friendly way to write Lewis dot symbols, where I put the number of valence electrons in parentheses.]

So we see that the formation of sodium chloride allows both sodium and chlorine to satisfy the octet rule.

Born-Haber Cycle

Now lets take a look at the thermodynamics of the formation of NaCl:

Na(s) + Cl2(g) NaCl(s) DHf = -410.9 kJ/mol

By comparison the standard enthalpy of formation of water is much smaller:

H2(s) + O2(g) H2O(l) DHf = -285.8 kJ/mol

Yet we saw in class how much energy is released when hydrogen and oxygen come together to form water, so we see that formation of NaCl releases a great deal of energy. Put another way, NaCl(s) is much more stable than the isolated elements.

If we consider formation of NaCl from the elements in a stepwise process we can determine the energy gain associated with bringing the cations and anions together.

Reaction

DHf

Na(s) Na(g)

108 kJ

Cl2(g) Cl(g)

122 kJ

Na(g) Na+(g) + e-

496 kJ (1st IE of Na)

Cl(g) + e- Cl-(g)

-349 kJ (EA of Cl)

Na+(g) + Cl-(g) NaCl(s)

x

If we sum up the reactions we get

Na(s) + Cl2(g) NaCl(s)

So according to Hess’ Law

108 kJ + 122 kJ + 496 kJ – 349 kJ + x = -410.9 kJ

x = -788 kJ

This energy is highly exothermic. It arises because of the ionic interactions (bonding) between cations (Na+) and anions (Cl-) in the NaCl lattice. The reverse process is called the lattice energy.

 

Lattice Energy

Lattice Energy The energy required to completely separate 1 mole of an ionic compound into its gaseous ions.

NaCl(s) Na+(g) + Cl-(g) DH = +788 kJ/mol

Lattice energy is a direct measure of the strength of ionic bonding. The larger the lattice energy the stronger the ionic bonding.

As stated previously ionic bonding arises from the Coulombic attraction between cations and anions. The potential energy associated with two interacting charges is:

E = (k Q1Q2)/d

Where k is a constant, Q1 & Q2 represent the magnitude of the charges (+1, +2, -1, etc.) and d is the distance between the charges. From this relationship we can draw the following conclusions about the magnitude of lattice energy/ionic bonding in real compounds:

To illustrate this consider the following compounds, all of which adopt the same crystal structure (rock salt):

Compound

Q1

Q2

d ()

Lattice Energy

KBr

+1

-1

3.30

-671 kJ

NaCl

+1

-1

2.82

-788 kJ

LiF

+1

-1

2.01

-1030 kJ

MgO

+2

-2

2.11

-3795 kJ

ScN

+3

-3

2.22

-7547 kJ

We see that either increasing the charge or decreasing the distance between ions leads to an increase in the strength of the ionic bonding (although increasing the charge has a much larger effect). What is it then that stops the cations and anions from becoming very highly charged and moving very close together?

The energy required to remove or add electrons beyond the nearest noble gas configuration of an element (i.e. to form Na2+ and Cl2-) limits the charges on the ions.

Repulsions between core electrons on neighboring ions limits the distance between ions.

 

Ionic Radii

By measuring the distances between nuclei in ionic compounds we can determine the radii of cations and anions, just as we did in chapter 7 for neutral atoms.

Ionic radii for several common ions are shown in Figure 8.5 of your text. Some properties of ionic radii that you should know are as follows:

Ionic radii are very useful for predicting the distances between ions in ionic solids. For example we could have used ionic radii to predict the distances in the above table.

 

Compound

Observed distance ()

Cation radius ()

Anion radius ()

Calculated distance ()

KBr

3.30

1.33

1.96

2.29

NaCl

2.82

0.97

1.81

2.78

LiF

2.01

0.68

1.33

2.01

MgO

2.11

0.66

1.40

2.06