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Melting and Boiling – Play with the simulations below.
In a solid, atoms or molecules are closely packed together.
When the solid melts, the atoms or molecules typically move further away from each other, forming a liquid. (s) → (l) (Exceptions to this, such as water, are not a feature of this course.)
When a liquid boils, the atoms or molecules move even further away from each other forming a gas. (l) → (g)
The more tightly atoms or molecules are held together, the more energy is needed to move them apart (higher melting point/boiling point).
Consider these three samples: C6H14 (s), C6H14 (l) and C6H14 (g)
a) In which case are the molecules furthest apart?
b) In which case are the molecules closest together?
c) Which sample must be at the highest temperature? How do you know?
Bonding
Ionic bonding
● Electrostatic attraction between oppositely charged ions.
● This occurs
○ between a metal and a nonmetal
○ when polyatomic ions are present
○ very occasionally when two nonmetals have a large enough difference in electronegativity that one nonmetal takes electrons from the other making positive and negative ions (greater than 2.0 difference in electronegativity).
● The attraction between oppositely charged ions forms a very strong bond.
Covalent bonding
● Nonmetal atoms are held together by the shared electrons that fill the outer energy levels of both atoms.
● Covalent bonds are the strongest type of bond.
Which of the following contains covalent bonds, which contains ionic bonds and which contains neither?
a) Na3P b) PF3 c) CH4 d) Cu
Read This
Complex molecules such as NH4NO3 and Al(ClO3)3 contain polyatomic ions (in which the atoms are covalently bonded) attached to either positive or negative ions by electrostatic attraction. They are categorized as ionic compounds, not covalent compounds because it is the ionic bond that hold the substance together and ionic bonds that must be loosened or broken for the substance to melt or react with other chemicals.
Metallic bonding
● In metals, the valence electrons (typically the d-shell electrons) are delocalized. This means they are not attached to any specific metal nucleus. These free-moving electrons form what is called “the electron sea”.
● The metallic bond is the attraction between positive metal ions and the electron sea that surrounds them.
● The metallic bond is the weakest type of bonds which is why metals can be bent and drawn into wires. Some metals have such low melting points that they are liquid at or near room temperature (mercury, gallium etc.)
● https://youtu.be/vOuFTuvf4qk
Read This
From what you have read above, you might think that covalent compounds have the highest melting and boiling temperatures - that it takes the most energy to separate the molecules to form liquids or gases. In fact, the opposite is true - covalent compounds are typically the easiest to melt and boil. How can this be when the covalent bond is the strongest type of bond?
Covalent bonds are not what holds one covalently bonded molecule to another. In order to melt a covalently bonded substance such as water, it is not the covalent bonds between hydrogen and oxygen that must be broken. In order to melt, it is the attraction between two molecules that must be loosened or broken. These attractions are called intermolecular forces (IMF).
● This course will cover a few of the many types of intermolecular forces: London dispersion, dipole-dipole and hydrogen bonding.
3. In each pair, circle the atom or molecule that will probably have the lowest melting point based on the information about ionic, metallic and covalent compounds above:
a) NaI and HI b) Cu and P4 d) NaCl and Sn
Intermolecular Forces
London Dispersion Forces.
● The name can be shortened to London Forces, dispersion forces or LDF
● LDF are caused by the random, slight movement of the electrons in a molecule that create extremely weak, temporary dipoles.
● The more electrons present, the stronger these temporary dipoles are.
● LDF are only noticeable in nonpolar covalent compounds, but are actually present in all covalent compounds.
● These are the weakest of the intermolecular forces between covalent molecules.
4. These three nonpolar covalent compounds are all at room temperature: C25H52, C8H18 and CH4. Which will most likely be a gas, which will most likely be a liquid and which will most likely be a solid at room temperature? How do you know?
Dipole-Dipole Attractions.
● These attractions occur between polar covalent molecules.
● Polar covalent compounds have a permanent molecular dipole (the gold dipole from the PhET simulation, not the black bond dipoles).
● This gives each molecule a permanent, slight positive and slight negative end that will attract the positive and negative ends of other polar molecules.
● Dipole-dipole forces are much stronger IMF than London forces. However, all IMF’s are MUCH weaker than any ionic, metallic or covalent bond. IMF’s are NOT BONDS!!!!! They are the forces between polar covalent molecules.
Stronger dipoles result in stronger dipole forces.
Simulation
Use the above the simulation Comparing Attractive Forces
From the dropdown menu “select a pair of molecules” choose “pull apart F2 and F2.”
Predict how difficult it will be to pull apart the two molecules in the data table.
Using the green star, move one F2 away from the other. Comment on how easy or difficult this was in the data table.
From the dropdown menu, choose “pull apart H2 and H2.”
Predict how difficult it will be to pull apart the two molecules in the data table.
Using the green star, move one H2 away from the other. Comment on how easy or difficult this was in the data table.
From the dropdown menu, choose “pull apart HF and HF.”
Predict how difficult it will be to pull apart the two molecules in the data table.
Using the green star, move one HF away from the other. Comment on how easy or difficult this was in the data table.
In the last two columns, determine whether the molecules are polar or nonpolar and identify the type of intermolecular forces the molecules exhibit.
How would you expect F2’s boiling point to compare to Br2? Explain.
In lab, nonpolar covalent cetyl palmitate melted first, then polar covalent sucrose. Metallic copper and ionic sodium chloride will not melt over a candle flame - the flame does not get hot enough. Explain this data using your knowledge of bonding and intermolecular forces.
Like Dissolves Like.
● Polar substances dissolve in polar solvents while nonpolar substances dissolve in nonpolar solvents.
● Ionic substances dissolve in polar solvents because ionic compounds, like polar compounds have positive and negative parts to them.
Explain why sucrose dissolved in water and not in hexane.
Explain why cetyl palmitate dissolved in hexane and not in water.
Conductivity
● For a substance or mixture to conduct electricity, there must be a free flow of charges.
● Solid metals have the electron sea which consists of electrons that are free to move.
● Dissolved ionic compounds have dissociated (positive ions separate from negative ions), making their positive and negative ions free to move around in the solution.
Explain why solid sodium chloride, sucrose and cetyl palmitate did not conduct electricity as solids.
Explain why copper did not conduct electricity in water when it did conduct electricity as a solid.
Two more intermolecular forces for those who want to know more:
Induced dipole attractions
· Occurs when a polar molecule causes a slight movement in the electrons of a nonpolar molecule. This makes the nonpolar molecule very slightly polar as long as the polar molecule is nearby.
· This is how oxygen (nonpolar molecule) dissolves in water (polar molecule) so that fish can breathe.
Ion-dipole attractions
· These form when an ionic solid dissolves in water. The positive ion attracts the negative (oxygen) part of the water molecule and the negative ion attracts the positive (hydrogen) part of the water molecule.
One more type of solid for those who want to know more:
Network covalent solids
· Network covalent solids are solids in which covalent bonds exist in all directions.
· These solids have the highest melting and boiling points.
· Examples: diamond, glass (SiO2), SiC
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