| 
>>
View the other Key Equations and Concepts in this
chapter
Identifying Intermolecular Forces
>> Parts of this equation/concept include:
The type of intermolecular force depends on the type of molecule.
Therefore it is important to classify the molecule.
The first classification is whether the molecule is covalent or ionic.
Recall that ionic molecules are a combination of a cation and an
anion. The common cations are either a metal or ammonium ion (NH4+).
The common anions are nonmetals or combinations of nonmetals. On
the other hand, covalent compounds are combinations of nonmetals
only. Remember that hydrogen (H), regardless of its position on
the periodic table, is always a nonmetal. Although strong acids
form ions in water (like ionic compounds), this is due to a reaction
with the water. The ions were not present in the original compound,
as they are in ionic compounds.
The covalent compounds are then classified as either polar or nonpolar.
Determination of polarity often requires drawing the Lewis structure
and determining the molecular shape. You may wish to review that
section of Chapter 7.
Ionic compounds have ion-ion molecular forces. Hydrogen bonding occurs
in molecules where hydrogen is covalently bonded to an oxygen, nitrogen,
or fluorine atom. This can be determined from the Lewis structure.
Determination of polarity is not necessary. Polar covalent molecules
have dipole-dipole forces. Nonpolar covalent molecules have London
or dispersion forces. (London and dispersion are two different names
for the same force. You should recognize both names, but either
answer is correct for nonpolar covalent molecules.)
>> Example 1
What is the force of attraction between molecules of the following substances?
- PCl3
- NiCl2
- I2
- HF
- HCl
- CH2O
- H3PO4
- BF3
Solution:
- This is a covalent compound. (So it does not have ion-ion
forces. No hydrogens, so hydrogen bonding is not a possibility.)
In the Lewis structure, phosphorus is the central atom. It is
bonded to the three chlorine atoms. The phosphorus atom also
has a lone pair. Therefore the molecular shape is a trigonal
pyramid, which makes it a polar compound. It therefore has dipole-dipole
forces.
- This is an ionic compound of the metal ion Ni2+
and the nonmetal Cl–. Therefore it has ion-ion
forces.
- This is a covalent compound. All atoms are the same type,
so the bonds are all nonpolar and the molecule is nonpolar.
Therefore this has London (or dispersion) forces.
- This is a covalent compound. Hydrogen must be bound to fluorine.
(There are no other atoms!) Since fluorine is one of the "big
3" electronegative atoms (FON), this compound will hydrogen
bond.
- This is a covalent compound with hydrogen bonded to chlorine.
This is a linear, polar molecule, so the force of attraction
is dipole-dipole.
- This is a covalent compound with both hydrogens and the oxygen
bonded to the carbon. Since the hydrogens are not bonded to
the oxygen, this does NOT hydrogen bond. It is, however, polar,
so the force of attraction is dipole-dipole.
- This is a covalent molecule. In fact, it is an oxoacid. Recall
that the structure of oxoacids is that hydrogen is bonded to
oxygen. Since hydrogen is directly bonded to electronegative
oxygen, this compound will hydrogen bond
- This is a covalent compound. Boron is satisfied with a deficient
octet. Therefore the structure of this compound is that all
three fluorines are bonded to the boron, with no other bonds
or lone pairs. The molecular geometry is trigonal planar. Since
the molecule is the same on all sides, the molecule is nonpolar.
Therefore this compound has dispersion (or London) forces of
attraction.
| A. Ranking the Strength of Intermolecular Forces |
In general, ion-ion forces are stronger than hydrogen bonds, which
are stronger than dipole-dipole forces, which are stronger than
London (or dispersion) forces. Exceptions can be found for two specific
compounds in nearly every category. However, that is not something
that can be determined from the formula, so your answer should be
based on that general ranking.
>> Example 2
Rank the strength of the following compounds from strongest to
weakest. (Note: These compounds are also used in the previous
example.)
PCl3, NiCl2, I2, HF
Solution:
From the previous example we know that PCl3 has dipole-dipole
forces, NiCl2 has ion-ion forces, I2 has
London forces, and HF has hydrogen bonds. Ranking according to
ion-ion > hydrogen bond > dipole-dipole > dispersion, the ranking
of the four substances is
NiCl2 > HF > PCl3 > I2
Compounds can also be ranked within the category of ion-ion forces.
Compounds made of ions with higher charges will be stronger than
ions with lower charges. Smaller ions are stronger than larger ions.
Ionic size can be found in Figure 9.3. You can also estimate the
size from the periodic table. Large ions will be lower (in higher
periods). There is also an effect of arrangement of ions. This makes
the distance factor harder to estimate. Therefore it is generally
a good principle to rank the strength of ionic compounds based first
on charge, then on size.
>> Example 3
Rank the following ionic compounds according to the strength
of their ion-ion forces of attraction
NiCl2, Fe2O3, MgO, KI, BeF2,
CrCl3
Solution:
First, look at the charges on the ions that make up each compound.
NiCl2 = Ni2+ and Cl–
Fe2O3 = Fe3+ and O2
MgO = Mg2+ and O2
KI = K+ and I
BeF2 = Be2+ and F
CrCl3 = Cr3+ and Cl–
Ranking according to the ions with the highest charge. (Multiply
charges together if the combination makes it difficult to decide.)
Ranking according to charge, the order is
Fe2O3 > MgO > CrCl3 > NiCl2
= BeF2 > KI
Note that MgO is greater than CrCl3, even though +3
> +2. If you multiply the charges, 2(2) = 4, which is greater
than 3(1) = 3. The subscripts do not contribute to this decision.
Note also that, based on charge, NiCl2 and BeF2
are the same. Therefore you need to go to the second criteria
(that of size) to rank these compounds. Beryllium and fluorine
are both in the second period of the periodic table. That makes
them smaller than the nickel and chlorine (fourth and third, respectively).
Smaller ions can be closer together, so they will have stronger
ion-ion forces of attraction.
The final order is
Fe2O3 > MgO > CrCl3 > BeF2
> NiCl2 > KI
Compounds can also be ranked within the category of London forces.
Because these forces are due to a temporary dipole formed by the
perturbation of the electron cloud, molecules with bigger electron
clouds will be more easily and more greatly perturbed. Since the
number of electrons is related to the number of protons, which is
related to molar mass, the easiest way to rank compounds with dispersion
forces is by molar mass. Nonpolar compounds with higher molar masses
will have stronger London forces.
>> Example 4
Rank the following from strongest to weakest intermolecular forces.
Cl2, CH4, BF3, SCl2,
CO2
Solution:
All these compounds, except SCl2, are nonpolar. SCl2
has two lone pairs on the central atom of sulfur. Its molecular
geometry is bent. Consequently, this compound has dipole-dipole
forces of attraction and will be stronger than the other choices.
The other choices are ranked based on molar mass. The molar mass
of Cl2 = 70.91 g/mol, CH4 = 16.05 g/mol,
BF3 = 67.81 g/mol, CO2 = 44.01 g/mol. The
higher molar mass will be stronger, so
SCl2 > Cl2 > BF3 > CO2
> CH4
>> back
to the Top of the Page
| B. Ranking Physical Properties |
Predictions about various physical properties can be made based
on the strength of the intermolecular forces.
>> Solubility
There are two aspects to solubility. The forces between the solvent
and the solute must overcome the solute-solute and solvent-solvent
forces. This is most effective if these forces are similar. Therefore
polar molecules dissolve in polar solvents and nonpolar molecules
dissolve in nonpolar solvents. A common chemist's rule of thumb
expressing this idea is, "Like dissolves like." The most common
solvent is water. It is a polar solvent.
Ionic compounds sometimes dissolve in polar solvents (particularly
water). This occurs when the ion-ion forces are relatively weak.
The solubility rules (Table 5.4 in the textbook) are still the best
guide to determine solubility of ionic compounds. However, these
empirical rules can be partially explained by ion-ion forces. For
example, all the ions that are "always soluble" have a charge of
1. So do the mostly soluble compounds, except sulfate. However,
even sulfate can be somewhat explained by its large size.
Nearly every material dissolves slightly in any other material.
The question, "Is it soluble?" implies that a significant quantity
of solute dissolves.
>> Example 5
Which of the following compounds are likely to dissolve in water?
- SCl2
- O2
- NaCl
- CO2
- PH3
Solution:
- SCl2 is a polar molecule with a bent geometry.
Therefore it will dissolve in water.
- O2 is a nonpolar molecule with a linear geometry
and a nonpolar bond between the oxygen atoms. Therefore it will
not be soluble in water.
- NaCl is an ionic compound. The charges are low (Na+
and Cl–), and according to the solubility rules,
"group I cations are always soluble." Therefore NaCl is soluble
in water.
- CO2 is a linear, nonpolar molecule. Therefore it
is not soluble in water.
- PH3 is a trigonal pyramid, polar molecules. It
will dissolve in water.
>> Example 6
Which of the following will dissolve in C6H14 (hexane)?
- SCl2
- O2
- NaCl
- CO2
- PH3
Solution:
Hexane is a nonpolar solvent. Hydrocarbons, molecules made solely
of hydrogen and carbon, are nearly always nonpolar. Therefore
nonpolar molecules will dissolve in this solvent. In the list
only O2 and CO2 are nonpolar, so only these
two compounds are soluble in hexane.
>> Melting Point
For a substance to melt, kinetic forces (increased by increasing
temperature) must overcome the intermolecular forces keeping the
particles in a fixed position. Therefore higher temperatures are
needed to overcome stronger molecular forces. In other words, the
stronger the intermolecular force, the higher the melting point.
Remember that a freezing point has the same value as a melting
point and can be determined exactly the same way.
>> Example 7
Rank the following substances from highest melting point to lowest
melting point.
KNO3, H2O, N2O, NaCl, F2,
Cl2
Solution:
First, determine the type of intermolecular force involved in
each molecule. Sometimes, it pays to pick out the obvious ones
first, rather than try to draw Lewis structures and determine
molecular geometry for each compound.
Ion-ion forces are the strongest and are found in salts. KNO3
and NaCl are both salts, so these need to be ranked. Since all
ions in both compounds have a +1 charge, size will determine the
ranking. K is lower in the periodic table than Na, and so is larger.
Nitrate ion, NO3–, is a polyatomic
ion and therefore larger than Cl–. Since smaller
ions have stronger forces of attraction, NaCl has a higher melting
point than KNO3, and both have higher melting points
than any of the other choices.
Dispersion forces are sometime obvious. Since F2 and
Cl2 are homoatomic (made of all the same kinds of atom),
these will have dispersion forces and have the lowest melting
points. Since dispersion forces are stronger with larger molecules,
Cl2 has a higher melting point than F2.
H2O must hydrogen bond. The only atom the hydrogen
can bond to is oxygen, since hydrogen itself cannot be a central
atom. That makes it weaker than the salts but stronger than the
halogens.
That leaves only N2O. In drawing Lewis structures,
the least electronegative atom is normally the central atom. Since
nitrogen is less electronegative than oxygen, one of the nitrogen
atoms is the central atom. This skeletal structure requires oxygen
on one side and oxygen on the other. There is no way for the pull
of electrons to cancel with two different atoms. Therefore this
must be a polar molecule, with dipole-dipole forces.
Using this information to arrange the substances by melting point,
the compound with the highest melting is NaCl and the lowest is
F2,
NaCl > KNO3 > H2O > N2O >
Cl2 > F2
>> Boiling Point
Boiling points are always higher than melting points for any one
substance. The method for ranking different substances is the same
as that for melting points. Substances with stronger intermolecular
forces require higher temperatures to escape that attraction and
become a gas (boil). The ranking assumes a constant atmospheric
pressure. Lower atmospheric pressures decrease the boiling point
of any substance.
>> Example 8
Rank the following substances from highest boiling point to lowest
boiling point.
CaS, HNO3, KBr, SO2, XeF4
Solution:
Determining the "obvious" intermolecular forces first, CaS and
KBr are ionic compounds, with ion-ion intermolecular forces. The
forces will be stronger for CaS, since it is made from ions with
a charge of 2, whereas KBr is made from ions with a charge of
1. HNO3 will hydrogen bond, since it is an oxoacid
and hydrogen is bonded to the oxygens rather than the central
atom. (Even if it was bonded to the central atom, nitrogen is
electronegative enough to induce the hydrogen bond.)
The other two atoms require Lewis structures. The Lewis structure
of SO2 has a double bond to one of the oxygens, single
bond to the other, and a lone pair on the sulfur. Since the double
bond could be attributed equally to either oxygen, this is a resonance
compound and the true position of the double bond is equally distributed
between the two oxygens, so both sulfur-oxygen bonds are exactly
the same. However, the lone pair on the sulfur makes the geometry
bent, and the overall molecule is polar. Therefore it has dipole-dipole
forces of attraction.
The XeF4 expands the octet of the xenon. The xenon
not only bonds to the four fluorine atoms but also has two lone
pairs. The overall molecular geometry is square planar because
the lone pairs orient themselves on opposite sides of the octahedral.
Therefore the partial negative charge due to the lone pairs cancel,
as do the polar bonds of each xenon-fluorine bond. Overall, the
molecule is nonpolar and has dispersion forces of attraction.
Based on this, the ranking of from highest boiling point to lowest
boiling point is
CaS > KBr > HNO3 > SO2 > XeF4
>> Vapor Pressure of Pure Substances
Vapor pressure is determined by the number of molecules that can
escape from the surface of a liquid. More molecules can escape if
the force of attraction, which holds the molecules together, is
weak or if the force pulling them away (kinetic energy measured
by temperature) is high. Problems asking you to rank vapor pressure
assume all substances are at the same temperature (in the same way
that ranking in terms of boiling point assumes that all substances
are at the same atmospheric pressure).
>> Example 9
Rank the following substances from highest vapor pressure to
lowest vapor pressure.
Br2, NH3, Ar, PCl3, PCl5
Solution:
NH3 must hydrogen bond, since the hydrogen must be
bonded to the nitrogen. Br2 and Ar must have dispersion
forces, since Br2 has only nonpolar bonds and Ar has
no bonds. You will probably need to draw the Lewis structures
for PCl3 and PCl5. PCl3 has three
P-Cl bonds and one lone pair. Therefore it has a trigonal pyramid
geometry and is polar. PCl5 has an expanded octet with
five P-Cl bonds and no lone pairs. It is a nonpolar molecule.
The nonpolar molecules are ranked by molar mass. The molar masses
are Ar = 39.95 g/mol, PCl5 = 208.24 g/mol, Br2
= 159.80 g/mol.
Since the highest vapor pressure has the weakest intermolecular
forces, the nonpolar molecule (dispersion force) with the lowest
molar mass will have the highest vapor pressure. The substance
with the strongest intermolecular forces will have the lowest
vapor pressure, so
Ar > Br2 > PCl5 > PCl3 > NH3
>> Viscosity and Surface Tension
Like vapor pressure, viscosity and surface tension depend on both
the temperature and the intermolecular forces of the substance.
Increasing temperature decreases both the viscosity and the surface
tension. At any given temperature, the substance with stronger intermolecular
forces has greater viscosity and higher surface tension.
>> Example 10
Rank the following substances from highest viscosity to lowest
viscosity.
Br2, H2O, CCl4, PH3
Solution:
Liquids with stronger intermolecular forces have higher viscosity.
H2O will hydrogen bond. PH3 is polar (phosphorus
has a lone pair). Br2 and CCl4 are both
nonpolar. The molar mass of bromine is 159.8 g/mol and that for
carbon tetrachloride is 153.82 g/mol. Therefore highest viscosity
is for water and the lowest is for carbon tetrachloride. The ranking
is
H2O > PH3 > Br2 > CCl4
>> Example 11
Rank the substances from the previous example from highest surface
tension to lowest surface tension.
Solution:
Liquids with stronger intermolecular forces have higher surface
tensions. Therefore the order of ranking is the same as it is
for viscosity.
H2O > PH3 > Br2 > CCl4
>> View
the other Key Equations and Concepts in this chapter
|
