Chapter 15 Review
15.1 (Part One)
Lattice Enthalpy = the enthalpy change that occurs when one mole of a solid ionic compound breaks into its ions in the gaseous state under standard conditions
The heat or energy absorbed will be used to break the bonds between the element and the compound (so the process will be endothermic)
So lattice enthalpy is positive
*Lattice enthalpy is a measurement of the stability of an ionic crystal*
The greater the lattice enthalpy, the more stable the ionic crystal (you can also say that the greater the lattice enthalpy, the stronger the ionic bond)
Lattice Enthalpy Values for Ionic Compounds
Lattice enthalpy depends on the strength of the electrostatic force of attraction between the positive and negative ions in the crystal
In an ionic compound, the strength of the EFA (electrostatic force of attraction) depends on two factors
1.
Ion charge
The greater the charge, the stronger the strength of EFA, therefore the greater the lattice enthalpy
2.
Ion size
The bigger the size/ radius, the lower the strength of EFA, therefore the weaker the lattice enthalpy

Example Question #1: Explain the differences in the lattice enthalpy values for the Group 1 chlorides:
Li, Na, and K are all in the same group so they have the same charge (+1) (so you can’t use the factor of ion charge to explain the difference in their lattice enthalpy)
Therefore, you can use the size of their ions
We know that the potassium ion has a greater radius, therefore it would have a lower strength of EFA so it has a weaker lattice enthalpy
Example Question #2: Explain the differences in the lattice enthalpy values for these compounds: a) Be would have a greater lattice enthalpy because it has a greater charge than Li b) Mg would have a greater lattice enthalpy because it has a greater charge than Li c) MgO would have a greater lattice enthalpy because O has a greater charge than Cl

Born Haber Cycles
Lattice enthalpy can’t be determined directly by experiment
Instead, it can be found indirectly from other known enthalpy changes by applying Hess’s Law
We start the cycle from the MX(s) phase where we have our solid ionic compound
(crystal) and based on the enthalpy change definition, it’s when one mole of a solid ionic compound breaks into it’s gaseous ions
So that turns into M + (g) + X (g)
We learned that the enthalpy of formation of that ionic compound, depends on the elements in their standard state
Therefore, we have the solid metal + the gaseous nonmetal
In order to figure out how we get the value of the standard enthalpy of formation, we go back to the cycle and see that on step A, there are “several reactions and several delta H values” that occur as shown by the four arrows
Arrows:
1.
Solid metal to gaseous metal atoms
convert the solid metal atom to a gaseous one (this process is called enthalpy of atomization/ vaporization)
this process is endothermic because you’re moving from a solid to gas so heat is absorbed to break the bonds

2.
Gaseous metal atoms to ions
when you have that gaseous metal that was formed from the solid metal, it looses an electron to form an ion (therefore you have first ionization energy: energy required to remove one electron from an element in the standard state)
the process is also endothermic here because you’re moving from a gaseous atom to a gaseous ion so the bond is being broken
3.
Non-metal molecules to gaseous atoms
this process is the bond enthalpy that we previously learned because we have a diatomic molecule present (X + X is broken to form X)
also endothermic reaction because the bond is being broken
moving from diatomic to monatomic gases
4.
Gaseous non-metal atoms to ions
talking about first electron affinity = the energy released when one mole of a compound/ atom gains an electron
exothermic reaction because energy is being released to form this ion (gaseous atom to negative ion)
Born Haber Cycles – Individual Reactions
In arrow 4…
First electron affinity = the energy released when one mole of gaseous atoms accepts one mole of electrons to form ions with a -1 charge

Second electron affinity = the energy absorbed when one mole of gaseous ions with a -1 charge accepts one mole of electrons to form ions with a -2 charge
In steps B and C…
B) elements in their standard state form 1 mole of ionic compound
C) solid ionic crystal dissociates into gaseous ions

Worked Example #1: Use the following data, plus values from the data booklet, to construct a
Born Haber cycle to find the lattice enthalpy for sodium chloride.
Elements always start in their standard state: Na(s) + 1/2Cl
2
(g)
Solid ionic compound to the right: NaCl(s)
Different gaseous ions at the top of the triangle: Na + (g) + Cl (g)
The element of Na(s) at room temperature moves (called enthalpy of atomization) to form a gaseous atom of Na(g) and this is an endothermic process
Move from the gaseous atom to a gaseous ion in the process of first ionization energy since you’re only losing one electron to form Na + (g)
Now that this cation has formed, you can focus on your other atom of 1/2Cl
2
(g) that goes through the process of average (1/2) bond enthalpy/ enthalpy of atomization in order to form Cl(g) and it’s endothermic because the bond is being broken
This Cl(g) undergoes first electron affinity to form Cl (g) which is exothermic since energy is released when one electron is accepted by the chlorine atom (which works since it’s in group 17)
When you look at the bottom of the pyramid, you see enthalpy of formation of sodium chloride that’s obtained with the value of -411 (definition of lattice enthalpy: when one

mole of a solid ionic compound breaks into it’s gaseous ions) (therefore, it’s an exothermic process)
To calculate the lattice enthalpy, you add all the values together and equal them to the enthalpy of formation + x
15.1 (Part Two)
Enthalpy Change of Solution = the enthalpy change that takes place when 1 mole of a solution dissolves in a solvent to form an ‘infinitely’ dilute solution under standard conditions of temperature and pressure
To Determine Enthalpy Change of Solution:
1.
Lattice Enthalpy
2.
Hydration Enthalpy (only for ions: Cations/ Anions)
Enthalpy Change of Solution – Lattice Enthalpy
lattice enthalpy is the enthalpy change that occurs when 1 mole of a solid ionic compound/ lattice breaks into its ions in the gaseous state under standard conditions
ΔH
LE
is defined endothermically
The less closely ions pack together in the solid lattice, the more endothermic the ΔH
LE
is.
The smaller and more highly charged the ions (the greater their charge density), the closer they pack

Enthalpy Change of Solution – Enthalpy of Hydration
= the enthalpy change that occurs when 1 mole of gaseous ions is dissolved to form an infinitely dilute solution of 1 mole of aqueous ions under standard conditions of temperature and pressure.
If you have a gaseous metal ion, it will dissolve into 1 mole of an aqueous ion (so you’ll have enthalpy of hydration of the cation)
If you have a gaseous nonmetal ion, it will dissolve into an aqueous ion (so you’ll have enthalpy of hydration of the anion)
The smaller and more highly charged the ions (the greater their charge density), the more exothermic the ΔH hyd
Calculation of Enthalpy Changes in Aqueous Solution
*under standard conditions of temperature and pressure:
298K
1 atm.
Worked Example #1: Calculate the enthalpy of solution of NaCl (s).
Method #1 – Algebra
Look up values on data booklet

Method #2 – Enthalpy Diagram
Start from the solid ionic compound
At the top of the diagram, break the compound into gaseous ions (and write their lattice enthalpy)
Turn those gaseous ions into aqueous ions (enthalpy of hydration)
Finally, turn the initial solid ionic compounds into their aqueous ionic forms using enthalpy of solution