More Intermolecular Forces, Vapor Pressure, and Lattice Energy

This week we spent more time discussing intermolecular forces in depth. We did Lecture 24 which discussed vaporization, physical and chemical changes, heating curves, vapor pressure, viscosity surface tension, and capillary action.

In this lecture worksheet, there were questions where we had to choose which compound would have the highest vaporization and why this was the case. We learned that the reason there is a different enthalpy of vaporization is because of the difference in the intermolecular forces between each compound.

We also discussed physical and chemical changes and what type of reaction classified each category. We found out that a phase change in molecular compounds is a physical change. A chemical change for molecular compounds would be a decomposition reaction. The interesting thing about ionic compounds is that when they become aqueous, the ions break apart. This type of reaction is both a chemical and physical reaction. It is chemical because the ionic bonds are being broken between the ions due to ion-dipole interactions with the water. The reaction is also physical because the ionic compound is going from a solid to a liquid via the water. It can be argued that it is a physical change because if the water is taken out, the ionic compound returns to a solid.

 

The standard heating curve (seen below) was discussed in detail, including when phase changes were happening, if the change was endothermic or exothermic, and if the substance was gaining energy or losing energy. 

From A-> B the substance is a solid, from C->D the substance is a liquid, and from E->F the substance is a gas. The lines that are parallel with the X-axis are where the phase changes are happening. The reason that they are not increasing in temperature at this point is because all of the extra energy is going into breaking apart the intermolecular forces between the molecules so that the compound can change phases. Going up to the right, the reactions are endothermic meaning that heat is being absorbed during the reaction. Going down to the left the reactions are exothermic meaning that heat is being lost during the reaction.

We discussed vapor pressure, which is the pressure of a vapor inside of a closed system. We defined vapor as the gaseous particles of a substance that is a liquid at room temperature.  We learned that liquids like water don't need extra heat to evaporate and that these substances are continuously evaporating and condensing. Dynamic equilibrium is when particles are condensing and evaporating at the same rate inside a closed system. The boiling point of a liquid is the temperature where the atmospheric pressure is equal to the vapor pressure. The pressure needed to boil is 760 torr. The reason the boiling point of water is lower than 100 degrees celsius at higher elevations is because there is higher atmospheric pressure in the mountains and therefore it doesn't take as much heat energy to reach the equilibrium between atmospheric pressure and vapor pressure. We also learned that as IMFs increase the boiling point also increases because more energy is needed to break the molecules from these forces. 

 

The last thing that we discussed this week was lattice energy. Lattice energy is the energy required to separate a mole of an solid ionic compound into gas ions.  Lattice energy increases with charge, and decreases with size of the ions. Coulombs law can be used to explain this relationship. As lattice energy goes up, the melting point goes up as well because it takes more energy to break apart the ions from a solid lattice than a liquid. 

I would give my understanding of the material this week a 9/10. I felt like the discussions that we had about the various IMFs and the properties that are influenced by them like vapor pressure, and boiling point were very helpful. I now feel that with a little more studying I will be very prepared for the test on Tuesday.

Metals,, Intermonlecular Forces, and Water

This week we learned about metals, intermolecular forces, and water. We learned many things about metals through a POGIL . We learned that metals usually have the following properties:

• shine or luster
• are malleable, which means that they can be hammered into different shapes with out breaking. 
• are ductile, which means that they can be made into wire
• conduct electricity and heat

We were introduced to the concept of metallic bonding, and how it is different than ionic or covalent bonding. This is because the electrons from all of the metal ions are shared between more than 2 atoms. These electrons are delocalized from their original atom and this is called the sea of electrons.

In this picture you can see that there are lots of electrons around all of the ions. This is the sea of electrons.

Next we discussed the intermolecular forces which are also called van der Waals forces.  There are four types of van de Waals forces:

• dipole-dipole
• dipole-induced dipole
• induced dipole-induced dipole
• hydrogen bonding

Dipole-dipole  forces are intermolecular forces that from between polar molecules. Because there is a separation  of charges throughout the molecule, the negative part can attract a positive part of another molecule and visa verse. This is a relatively weak force and it only releases about 5 kJ/mol of energy when broken.

Dipole-induced dipole are forces that act between a polar molecule and a nonpolar molecule. When there different molecules come together the nonpolar molecule can have an induced dipole. This means that temporarily there will be a unequal distribution of charges in the nonpolar molecule. Dipole-induced dipole is the weakest intermolecular force. It only takes about 2 kJ/mol of energy to break theses bonds.

Induced dipol-induced dipole forces are forces that form between two nonpolar molecules. These are caused by the rapidly moving electrons about the molecules that cause a momentary change in the charge of molecules. These forces are also called London Dispersion Forces (LDFs). They have about the same energy as dipole-dipole forces, which is about 5 kL/mol energy to break them apart.

Hydrogen bonding is bonding between hydrogen atom that is covalently bonded to a oxygen, nitrogen, or fluorine atom that bonds with a different O, N, or F in a different molecule. An abundant example of hydrogen bonding would be in water. This compound is formed by many H2O molecules that are hydrogen bonded to each other. This is the strongest of the intermolecular forces, needing about 20 kJ/mol of energy to break.



The last thing that we discussed this week was about water and ice. We learned that when water freezes it expands. This is because it forms ridged hexagonal crystals that take up more volume than the liquid water. We constructed an ice cube using magnetic water molecules to simulate hydrogen bonding.

Notice how there are spaces in between the molecules in the ice and how all of the spaces are filled in the liquid water.


The last thing that we talked about regarding water was  ion-dipole interactions. We mainly focused on how this related to ions/ionic compounds dissolved in water.When an ionic compound is dissolved in water the cations become surrounded by the oxygen atoms in the water molecules and the anions become surrounded by the hydrogen atoms of the water molecules. If there is a stronger cation-anion force of attraction than the force of attraction from the water molecules the substance will not be soluble in water.

This week I think that I came to understand the intermolecular forces better. I now know that LDFs are a result of induced dipole-induced dipole interactions. I also think that this week has helped me fill in some of the gaps that I had regarding metals and ionic compounds dissolving in water. I would give my understanding of the week a 8/10 because I feel as if I understand the main concepts but I still want to study this material more to make sure that I am very comfortable with it.