GHW#3. Chapter 3. Louisiana Tech University, Chemistry 481.
POGIL(Process Oriented Guided Inquiry Learning) Exercise on
Chapter 3. Energetics of Ionic Bonding.
Why?
What are the properties of ionic compounds? How periodic table is used to predict
ionic bonding? What is Coulombs law and how it applies to ionic bonding? What is
lattice energy? How lattice energy is calculated from Coulombs model and the Madelung
constant? How Madelung constant is affected by different ionic lattice types: cesium
chloride, rock-salt, fluorite etc. How lattice energy and melting points of ionic
compounds are affected by ionic radii? What are the periodic trends in ionic radii? What
is a charge density of an ion? How are charge density values used to predict more ionic
bonding or more polarizable/less ionic bonding? What factors affect the polarizability on
an ion? How does polarizability of an ion affect the lattice energy and melting points?
How do you calculate lattice energy for an ionic compound from thermodynamic data
using Born-Haber cycle? What is enthalpy of solution, Hsolution, enthalpy of hydration,
Hhydration and solvent-solvent intermolecular attractions, Hsolvent-solvent? How are the
Hsolution, Hhydration, LE and Hsolvent-solvent related for the solution process? What factors
affect the solubility of an ionic compound in a given solvent? How does solvent, LE,
Hhydration, and Hsolvent-solvent affect the solubility of an ionic compound? How does the
size and charge of ions of an ionic compound affect its LE and Hhydration? What controls
the relative change of both LE and Hhydration when ion size is decreased or incresed?
Why ionic compounds with smaller anions- NO3-, ClO4-, ClO3-, and C2H3O2-, and
cations- H+, Na+, K+, and NH4+ form soluble compounds?
Instructional Objectives:
1) Explain properties of ionic compounds: as electrical attraction, electricity of
molten liquid, electrical conduction when dissolved in water, brittleness when
hammered, the high melting and boiling points, solubility in polar solvents.
2) Use periodic table and electronegative to predict ionic bonding.
3) Apply Coulombs law to ionic bonding and qualitatively predict lattice energy.
4) Calculate lattice energy given charge of ions, ionic radii and Madelung constant.
5) Explain Madelung constant values for different ionic lattice types: cesium
chloride, rock-salt, fluorite etc.
6) Predict the trends in lattice energy and melting points of ionic compounds based
on ionic radii.
7) Predcit the trends in ionic radii using periodic table.
8) Calcualte is a charge density of an ion and compare them to predict more ionic
bonding or more polarizable less ionic bonding.
9) Explain the factors that affect the polarizability and its effect the lattice energy
and melting points.
10) Calculate lattice energy from thermodynamic data (IE, EA, BDA, SE, Hf etc.)
and using Born-Haber cycle.
11) Expalin factors affect the solubility of an ionic compound in a solvent, the
enthalpy of solution, Hsolution, the enthalpy of hydration, Hhydration and the
solvent-solvent intermolecular attractions, Hsolvent-solvent and their effect on
solubility.
12) Identify conditions that would make Hsolution, more negatice and ionic compound
more soluble.
13) Expain the effect of size and charge of ions on the LE and Hhydration.
14) Expain the solubility of ionic compounds with smaller anions- NO3-, ClO4-, ClO3-,
and C2H3O2-, and cations- H+, Na+, K+, and NH4+.
RESOURCES
INORGANIC CHEMISTRY By Peter Atkins, Tina Overton, Jon Rourke, Mark Weller,
Fraser Armstrong, 4th Edition 2006.
New Concepts
Properties of ionic compound:
1) An ionic bond is a strong electrical attraction between two oppositely charged
atoms or groups of atoms.
2) Conducts electricity as molten liquid.
3) When dissolved in water produce solutions that conduct an electric current: they
are electrolytes.
4) They break readily when hammered, because of ion repulsions created by the
slippage.
5) They have high melting and boiling points.
6) They are soluble in polar solvents.
Periodic trends in bonding and electronegativity
a) The main Group elements (s and p orbitals) lose or gain electrons to attain a
configuration like a noble gas.
b) Transition elements (d orbitals) lose their s orbital electrons first and then one or
more d orbital electron(s).
c) The ionic character of the bond is proportional to the electronegativity difference of
the elements making the bond.
Electronegativity Difference
0-0.5
0.6-1.6
1.6, or greater
Nonpolar covalent
Polar covalent bond
Ionic
C-H
O-H
Na-Cl
Coulombs law and the ionic bond: The electrostatic model is simply an application of
the charge principles that opposite charges attract and similar charges repel.
k = constant (permeability of the medium)
q+ = cation charge; q- = anion charge
r = distance between two ions
An ionic compound results from the interaction of a positive and negative ion,
such as sodium and chloride in common salt. The ions will cluster together so as to
maximize heir attractions and minimize repulsions. The arrangements assumed by these
ions will be determined by the compound formula and by the sizes of the ions.
Lattice Energy:
In crystalline compounds this net balance of attractive and repulsive forces is
called the lattice energy which is the energy released in the formation of one mole of
ionic solid from the gaseous ions. The lattice energy and melting point are directly
related.
E.g. Mg2+(g) + 2Br-(g)  MgBr2(s) HLE = -2440kJ mol-1
Based on Coulombs law,
as the ionic radii of either the cation or
anion increase, the lattice energies decrease, and the solids consists of di-positive ions
have much larger lattice energies than solids with mono-positive ions. Lattice type also
matters because packing is different. Coulomb law equation, is multiplied a factor called
Madelung constant to get lattice energy.
Madelung Constant: Factor that considers all electrostatic attractions in an ionic lattice
Trends in ionic radii
 Ionic radii increase down a group
 Ionic radii decrease across a period
Lattice
Madelung
Solid
 Cations are smaller than their parent
Type
Constant
atoms - lose electrons
NaCl
Rock salt
1.747558
 Anions are larger than their parent
CsCl
CsCl type
1.747558
atoms - gain electrons
CaF2
Fluorite
2.51939
TiO2
Rutile
2.408
Charge density of ions: Defined as Charge/volume expressed as coulombs/Å3.
Comparing charge density values of series of ions (Appendix 2) one can predict is most
likely to form compounds exhibiting more ionic bonding and ones that will be more
polarizable and show some degree of covalency. As the covalency increases in an ionic
bond the experimental lattice energy will be higher than the predicted by ionic model
alone.
Group 1 chlorides and fluorides: The lithium compound exhibits a lower melting point
than we would anticipate. Remember that lithium is very small and has a slightly higher
electronegativity than the other Group 1 metals, thus it appears that lithium's low melting
points are due to a bit of covalency and thus a slightly reduced polarity of the bonds.
Polarization will be increased by:
1. High charge and small size of the cation. Ionic potential Å Z+/r+ (= polarizing power)
2. High charge and large size of the anion. The polarizability of an anion is related to the
deformability of its electron cloud (i.e. its "softness")
3. An incomplete valence shell electron configuration. Noble gas configuration of the
cation better shielding.
e.g. Hg2+ (r+ = 102 pm) is more polarising than Ca2+ (r+ = 100 pm)
Calculation of lattice energy from thermodynamic data: A Born-Harber cycle could
be drawn for steps in the formation of any ionic compounds.
E.g. Draw a Born-Harber cycle for the formation of BaBr2 from barium metal and
bromine gas. Label each step with the appropriate thermodynamic quantity.
Calculate the enthalpy of formation for
BaBr2.
BaBr2 lattice energy = 1950 kJmol-1
Ba atomization energy= 175
Ba 1st ionization energy = 503
Ba 2nd ionization energy = 965
Br2 bond enthalpy = 193
Br electron affinity= -325
LE = EA + IE2 +IE1 BDE + SE - Hf
LE = [2 x (-325)]+ 965+503+ 193+175 – (-764 )= 1950 LE = 1950 kJmol-1
Enthalpy of solution, Hsolution:
The formation of a solution involves the interaction of solute with solvent molecules.
Many different liquids can be used as solvents for liquid solutions, and water is the most
commonly used solvent. When water is used as the solvent, the dissolving process is
called hydration. The heat change which takes place when one mole of a solute is
completely dissolved in a solvent to form a solution of concentration 1 mol L-1.
Enthalpy of hydration, Hhydration:
The heat evolved when 1 mole of gaseous ions become hydrated (surrounded by water
molecules), measured under standard conditions.
Hhydration(cation): Al3+(g) + aq  Al3+(aq) Hhyd(cation) = -4613 kJ mol-1
Hhydration(anion): Cl 1-(g) + aq  Cl 1-(aq) Hhyd(anion) = -363 kJ mol-1
Hhydration(AlCl3):
Al3+(g) + 3 Cl 1-(g) + aq  Al3+(aq) + 3 Cl 1-(aq) Hhyd = -4613 kJ mol-1+3 x (-363)] kJ mol-1
Hhyd(AlCl3) = -5702 kJ mol-1
Solvent-solvent intermolecular attractions, Hsolvent-solvent:
The energy required to break dipole - dipole interactions between solvent molecules (L)
when they become solvating ligands (L') for the ions.
The enthalpy of solution, Hsolution, enthalpy of hydration, Hhydration, lattice energy and
solvent-solvent intermolecular attractions, Hsolvent-solvent is related in the solution
process by the equation:
Hsol HLE + (Hhyd (anion) + Hhyd(cation)) + Hsolvent-solvent
Factors affecting solubility:
Solvent:
The "like dissolves like" rule is used to predict the solvent needed for solution process.
In other words, if you want to dissolve ionic compound you should use a solvent that is
also highly polar with a high dielectric constant and non-polar compound requires a nonpolar solvent with a low dielectric constant. A small solvent molecular have a smaller
dipole which can approach the ions closely to increase solubility.
Lattice Energy (LE): It takes energy to separate ions from their crystal lattice and from
hydrated ions. Smaller size of the ions increases both the lattice energy and hydration
enthalpy. If the lattice enthalpy increases more than the hydration enthalpy, then heat of
solution become more endothermic and vise versa.
Enthalpy of hydration, Hhydration: The attraction of a dipole solvent (water) dipole to
an oppositely charged ion often released as salvation (hydration) energy. This hydration
energy is used to break the ionic lattice: the lattice energy (LE). Non-polar and weakly
polar solvents do not have sufficiently strong hydration to overcome LE and to dissolve
electrolytes. In a crystal hydrate, the ions are largely hydrated, and consequently the
hydration energy is considerably less than that of the anhydrous solute. The hydrates,
therefore, usually have lower water solubility.
The enthalpy of solution, Hsolution:
Usually substances with a large negative heat of solution (i.e., exothermic reaction) are
more soluble than substances with a smaller negative heat of solution. Compounds that
have a positive heat of solution (endothermic) may also be soluble.
Size and charge of ions on LE and Hhydration:What controls the relative rate of fall
both LE and Hhydration?
Charge factor
Larger the charges on ions increases LE more than the Hhydration, therefore ionic
compounds with di-positve or negative charge tend to make the enthalpy of solution less
negative (less soluble).
Size factor
It turns out that the main factor increasing Hhydration over LE is the size of the negative
and positive ions.
E.g. Small anions- NO3-, ClO4-, ClO3-, and C2H3O2- form soluble compounds.
The lattice energy increase less than the hydration enthalpy of the positive ions. That
means that the enthalpy of solution will become more negative.
E.g. Small cations- H+, Na+, K+, and NH4+ forms soluble compounds.
The lattice energy increase less than the increase in hydration enthalpy of these small
positive ions make enthalpy of solution will become more negative.
E.g. Large anions- PO43-, S2-, CO32-, and SO32- ions are insoluble except those that also
contain alkali metals or Na+, K+, and NH4+. Changes in the size of the positive ion don't
make as great a percentage difference to the inter-ionic distance as they would if the
negative ion was small. The hydration enthalpy of the positive ions decreases more than
the lattice energy. The enthalpy of solution will become less negative.
Success Criteria
Ability to answer the questions and apply concepts related to the topics given in the
instructional objectives.
Resources
DESCRIPTIVE INORGANIC CHEMISTRY by Geoff Rayner-Canhanmi
Prerequisites
Freshman chemistry, chapters 1-4
of DESCRIPTIVE
INORGANIC
CHEMISTRY Rayner-Canhanmi and detailed knowledge of covalent (especially MO
theory) and ionic bonding models.
GHW# 3 Chapter 3. Ionic Bonding Your Name:________________________
Key Questions (relatively simple to answer using the Focus Information)
1. What properties of a compound would lead you to expect that it contains ionic
bonds?
2. Would you expect sodium chloride to dissolve in carbon tetrachloride, CCl4?
Explain your reason.
3. Which would you expect to contain ionic bonds, MgCl2 or SC12? Explain your
reasoning.
4. What is Coulombs law how it applies to ionic bond?
5. What is lattice energy? Take NaCl as an example.
6. Which one of each of the following pairs will be smaller radius? Explain your
reasoning in each case.
a) K or K+:
b) K+ or Ca2+:
c) Br- or Rb+:
d) Se2- or Br-:
e) O2- or S2-:
Compound
Interionic Distance Melting Point Lattice Energy
(Angstroms)
(Centigrade)
(kcal/mol)
NaF
2.31
988
-201
NaCl
2.79
801
-182
NaBr
2.94
790
-173
7. Explain the lattice energy and
melting point trends:
8. Explain the lattice energy
and melting point trends:
NaI
3.18
660
-159
Compound
Cation radius Anion radius Melting Point Lattice Energy
(Angstroms) (Angstroms) (Centigrade)
(kcal/mol)
MgCl2
0.65
1.81
714
2326
CaCl2
0.94
1.81
782
2223
MgO
0.65
1.45
2852
3938
CaO
0.94
1.45
2614
3414
9. Compare the charge density values of the three silver ions: Ag+, Ag+2, and Ag+3
(Appendix 2). Which is most likely to form compounds exhibiting ionic bonding?
10. Compare the charge densities of the fluoride ion and the iodide ion (Appendix 2).
On this basis, which would be the more polarizable?
11. How does polarization and covalency affect lattice energy and melting points?
Compound
Melting Point (oC)
AgF
435
AgCl
AgBr
AgI
455
430
553
12. Calculate the Lattice energy of NaCl from following thermodynamic data:
1.
2.
3.
4.
5.
Steps
Vaporization of sodium: Na(s)
 Na(g)
Decomposition of Cl2: 1/2 Cl2 (g)  Cl(g)
Ionization of sodium: Na(g)  Na+(g)
Electron affinity to chlorine: Cl(g) + e-  Cl-(g)
Formation of NaCl(s):
Na(g)+1/2Cl2 (g)

NaCl(s)
Ho, kJ
+92
+121
+496
-349
-411
13. Define following terms:
a) Enthalpy of solution, Hsolution:
b) Enthalpy of hydration, Hhydration:
c) Solvent-solvent intermolecular attractions, Hsolvent-solvent:
14. How is Enthalpy of solution, Hsolution, Enthalpy of hydration, and Lattice energy
are related?
15. Predict the solubility of following ionic compounds:
Lattice Energy(U)
Hhyd, M+
LiF
1030
-950
LiI
720
-950
CsI
585
-700
MgF2 3100
-2800
a) LiF:
b) LiI:
c) CsI:
d) MgF2:
Hhyd, M-60
-80
-80
-120
16. Give rational explanation to the solubility rules in terms of ioin sizes, lattice
energy(U), Hhyd, and Hsolution.
a) All compounds containing alkali metal cations and the ammonium ion are
soluble.
b) All compounds containing NO3-, ClO4-, ClO3-, and C2H3O2- anions are
soluble.
c) All chlorides, bromides, and iodides are soluble except those containing
Ag+, Pb2+, or Hg22+.
d) All sulfates are soluble except those containing Hg22+, Pb2+, Sr2+, Ca2+, or
Ba2+.
e) All hydroxides are insoluble except compounds of the alkali metals, Ca2+,
Sr2+, and Ba2+.
f) All compounds containing PO43-, S2-, CO32-, and SO32- ions are insoluble
except those that also contain alkali metals or NH4+.