Name ___________________________________ Date ______________________ Class ______________________
Worksheet #2
Chapter 13 – States of Matter
Source of States
When atoms share valence electrons they do not always share them equally. Frequently one atom has a stronger
attraction for the electrons than the other atom does. This uneven attraction causes the electrons to be held closer to
one end of the bond than the other; we say this makes one end of the bond slightly positive and the other end of the
bond slightly negative. A covalent bond with uneven sharing of the electrons is called a polar covalent bond. A bond
in which the electrons are shared equally is called a nonpolar covalent bond.
Define the following terms:
 Polar covalent –

Nonpolar covalent –
Electronegativity is a measure of the ability of an atom of an element to attract electrons to its self. Put another way,
electronegativity is a measure of the attraction that exists between an atom and a shared pair of electrons in a covalent
bond. Linus Pauling developed a scale of electronegativities that run from a low of 0.7 for several metals in Group 1 to a
high of 4.0 for fluorine.
The table below gives Pauling Values for Electronegativity
We use electronegativity values when we discuss bond polarity. If two atoms sharing a pair of electrons have equal
values for electronegativity the bond is clearly nonpolar. As the difference in electronegativity increases the polarity
of the bond increases, and if the difference in electronegativity is very large the bond is ionic.
1. What is electronegativity?
2. Sodium chloride (NaCl) is an example of an ionic bond. What is the difference in electronegativity between
sodium and chlorine?
3. Nitrogen dioxide (NO2) is an example of a covalent bond. What is the difference in electronegativity between
nitrogen and oxygen?
Kinetic Molecular Theory describes the states of matter in terms of the movement of the molecules in each state. The
word kinetic means “to move”. Objects in motion have energy called kinetic energy. Temperature is a measurement
of the average kinetic energy of all the molecules in a sample of matter. For example, molecules of H2O in water vapor
at 100°C are moving much faster than molecules of H2O in a block of ice that is 0°C.
Considering Kinetic Molecular Theory, one might ask “If all the molecules at room temperature have the same average
kinetic energy, why are some materials gases and others liquids or solids?” To understand why this is, we must consider
intramolecular and intermolecular forces. We defined the following terms in chapter 8:
 Intramolecular forces –

Intermolecular forces –

Which are stronger – intramolecular forces or intermolecular forces?
In this worksheet, we will focus on the three intermolecular forces: Dispersion/London forces, dipole-dipole forces and
hydrogen bonds.
Dispersion (London) Forces
At room temperature oxygen molecules (O2) act as a gas. Under the right conditions; however, oxygen molecules can be
compressed (squished) into a liquid. For oxygen to be compressed there must be some force of attraction between its
molecules. Because oxygen molecules are non-polar, they do not have a positive or negative end. The weak force
between non-polar molecules is called a dispersion (London) force. Draw a picture of this below:
Dipole-Dipole Forces
When carbon dioxide (CO2) gas molecules get close to each other, the partially positive carbon atom in one molecule is
attracted to the partially negative oxygen atoms in another molecule. The strong fore between positive and negative
ends of polar molecules is called dipole-dipole force. Draw a picture of this below.
Hydrogen Bonds
In a water molecule (H2O), the hydrogen atoms have a large partial positive charge and the oxygen atom has a large
partial negative charge. When water molecules get close to each other, a hydrogen atom on one molecule is attracted
to the oxygen atom on another molecule. The very strong force between polar molecules containing hydrogen is
called a hydrogen bond. Draw a picture of this below.
Now that we know about intermolecular forces, we can connect the ideas of electronegativity and bonding forces:
Difference in
electronegativity
(BIG – small)
Intramolecular Forces
(within ONE molecule)
Intermolecular Forces
(between MULTIPLE molecules)
Examples
0 – 0.49
Nonpolar Covalent
(electrons are SHARED EQUALLY)
Dispersion (London) Forces
O2
0.5 – 1.69
1.7 or greater
Polar Covalent
(electrons are SHARED UNEQUALLY)
Ionic (electrons are TRANSFERRED)
1. Hydrogen Bonding (any polar
molecules WITH HYDROGEN)
H2O
2. Dipole-Dipole (any polar molecule
(WITHOUT HYDROGEN)
Ionic (positive ions are attracted to
negative ions)
NaCl
Use your electronegativity table and the chart above to answer the following questions:
1. Determine the INTRAmolecular force for the following compounds: (nonpolar covalent, polar covalent, ionic)
CH4 = ___________________
CF4 = ___________________
HI = ___________________
CO2 = ___________________
NH3 = ___________________
NaCl = ___________________
2. Determine the INTERmolecular force for the following compounds: (dispersion forces, dipole-dipole, H bonding,
ionic)
CH4 = ___________________
CF4 = ___________________
HI = ___________________
CO2 = ___________________
NH3 = ___________________
NaCl = ___________________
Melting or boiling a substance breaks down (weakens) the forces between molecules in the substance.
Therefore, the stronger the intermolecular and intramolecular forces in a substance, more energy is required to melt the
solid or boil the liquid. Below is a list of all the forces we have discussed in order of strength.
1. Why is water a liquid at room temperature when compounds of similar mass are gases? (HINT: think about the
strength of the intermolecular forces)
2. For the following, state AND EXPLAIN which one of the two has a higher melting/boiling point:
a. Sodium Chloride (NaCl) or Ammonia (NH3)?
b. Iron (Fe) or Methane (CH4)?
c. Carbon Dioxide (CO2) or Sulfur Dioxide (SO2)?
Complete the table below: Room temperature 22°C
Example: since the melting point of ethanol is -114°C, above that temperature it is a liquid. Since the boiling point of
ethanol is 79°C, above that temperature it is a gas. Since room temperature is 22°C, ethanol is a gas at room
temperature.
Chemical Name
Melting Point
(°C)
Boiling Point
(°C)
Hydrogen
-259
-252
Oxygen
-218
-183
Nitrogen
-210
-195
Ethanol (alcohol)
-114
79
Water
0
100
Common salt
804
808
Iron
1535
2750
State at Room Temperature
LIQUID (see explanation above)