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## Unit 3 at a Glance

Unit 3 covers intermolecular forces and their properties. It asks you to relate the physical properties of solids, liquids, and gases to the forces occurring within them. At 18-22% of the AP Exam, this unit is a doozy. However, it also contains a lot of different concepts – from intermolecular forces to molarity to ideal gas laws.

Success in Unit 3 requires a strong grasp of the previous unit (which is pretty unfair, to be honest.) To be specific, you should know your molecular geometry and polarity rules as well (link TBA).

## Important Formulas You Should Know

All bolded formulas are on the equation sheet.

• PV = nRT
• PA = Ptotal × XA, where XA = moles A/total moles;
• Ptotal = PA + PB + PC + …
• KE = ½ mv2
• 0° Celsius = 273° Kelvin
• M = Mol of solute/L of solution.
• (E = ℎν) and (c = λν)
• A = εbc

## Unit 3 Outline

The following outline is derived from the College Board’s Course and Exam Description found here.

Study Tip #1: Diagrams are your saving grace in this Unit. Paul Anderson of Bozeman Science has a great YouTube playlist with lots of them! You can also find some great ones online. Let us know in the comments if you can think of any other good ones.

Study Tip #2: Units. Units are key! As my Physics C teacher, Mr. Murray, always says, “Units will set you free.” If you keep track of them, they’re a great way to verify that your calculations are correct. They can also cause you a lot of trouble if you have difficulty with them. I missed three questions on my IMF test because I couldn’t convert milliliters to liters. Other places where units might be a concern are when you’re doing Celsius to Kelvin conversions and PV = nRT calculations.

## 3.1 | Intermolecular Forces

• An Intermolecular force is a force between molecules. For example, the force between two H2O molecules. Breaking or forming them causes a physical change.
• London Dispersion Forces, hydrogen bonds, and ion-dipole forces are all examples of Intermolecular forces.
• The forces discussed in this unit are mostly intermolecular forces.
• An Intramolecular force is a force within a molecule. For example, the force between two H atoms in H2. Breaking or forming them causes a chemical change.
• Covalent and Ionic bonds are examples of intramolecular forces.

### Types of IMF

• London Dispersion Forces are caused by the reactions between temporary dipoles.
• They are present in all molecules.
• They are generally one of the weaker bonds, compared to dipole-dipole forces and hydrogen bonds.
• Dispersion forces increase in strength as the contact area between molecules and the polarizability of the molecules increases.
• Polarizability simply means the ability to become polar and gain a temporary dipole.
• The more electrons there are, the more the molecule can gain a temporary dipole and thus the stronger the London Dispersion Force.
• Dipole-Induced Dipole, Dipole-Dipole, and Ion-Dipole forces all only occur when polar molecules are present.
• Dipole-induced dipole forces happen between polar and non-polar molecules.
• They’re called that because the polar molecule induces a temporary dipole in the non-polar molecule.
• The greater the dipole is in the polar molecule and the higher the polarizability of the non-polar molecule (which correlates with a higher number of electrons), the higher the dipole-induced dipole forces will be.
• Dipole-dipole forces happen between polar molecules.
• The stronger the dipole, the greater the force.
• Ion-dipole forces happen between ions and polar molecules. (Ex: salt in water)
• For an ion, polarization increases the higher the charge and the lower the size (Columb’s Law)
• Hydrogen Bonding occurs when H is bonded to the highly electronegative atoms of F, O, and N and this hydrogen atom is attracted to the negative end of a
• Hydrogen Bonding is FON!
• It can be thought of as an extreme/stronger dipole.
• Not every H bonded to an F, O or N is necessarily hydrogen bonded. Be careful!

### Strength Order

• When comparing molecules of the same size, the order (by strength of bond) for forces typically goes as follows:
• Network Covalent Bonding (Covalent Crystals)
• Ionic Bonding (Link TBA) (Ionic Solids)
• Metallic Bonding (Metals)
• Hydrogen Bonding
• Dipole-Dipole Forces
• Ion-Dipole Forces
• Dipole-Induced Dipole Forces
• London Dispersion Forces

## 3.2 | Properties of Solids

• Many properties of solids and liquids can be determined based on the strength and types of IMF present.
• The stronger the IMF…
• …the higher the melting point
• …the higher the boiling point
• …the higher the viscosity
• …the greater the surface tension
• …the lower the vapor pressure
• Covalent Network Solids (quartz, diamond): Generally rigid solids with high melting points
• Metallic Solids (Iron, Sodium): Good conductors of heat and electricity, malleable; the valence electrons create a sea of electrons
• Ionic Solids: The measure of ionic bond strength is called Lattice Energy. Ionic solids tend to have low vapor pressures and high melting and boiling points. They tend to be brittle and only conduct electricity when the ions are mobile.
• Molecular Solids: low melting point, doesn’t conduct electricity

## 3.3 | Solids, Liquids, and Gases

• The molecules of solids are packed more tightly than those of liquids, which are packed more tightly than gases.
• The arrangement of liquid particles are dependant on IMF.
• The solid and liquid phases of a substance typically have similar molar volumes.
• Gas molecules are constantly in motion, so they have no definite shape.

## 3.4 | Ideal Gas Law

• Ideal gas molecules have no volume and experience no intermolecular forces.
• Real gases do take up volume and experience IMF.
• The ideal gas law is PV = nRT. The equation and its constants are on your equation sheet. Remember to use the right constant when you’re doing PV = nRT calculations.
• For a sample containing a mix of gases, the equations for pressure (P) are
• PA = Ptotal × XA, where XA = moles A/total moles;
• Ptotal = PA + PB + PC + …
• these are on your formula chart as well.
• STP stands for Standard Temperature and Pressure. It is defined as 273 Kelvin and 1 atm. In these conditions, one mol of gas occupies 22.4 liters of volume.

## 3.5 | Kinetic Molecular Theory

• KE = ½ mv2
• Hi, Physics! Nice to see you again.
• The temperature (in Kelvin) is proportional to the kinetic energy
• hotter = more kinetic energy
• The Kinetic Molecular Theory is represented by Maxwell-Boltzmann graphs (link TBA).

## 3.6 | Deviation from Ideal Gas Law

• Real gases don’t behave exactly like ideal gases.
• A gas behaves most ideally when it is at a high temperature and low pressure.
• A gas with smaller molecules and weaker IMF will behave more ideally than one with large molecules and a stronger IMF.

## 3.7 | Solutions and Mixtures

• A solute is what is being dissolved.
• A solvent is what is doing the dissolving.
• Homogeneous means that something is the same throughout.
• A solution is a homogeneous mixture. It does not scatter light and it cannot be separated by gravity filtration.
• The equation for the concentration, or molarity, of a solution is:
• M = Mol of solute/L of solution.
• This is on your equation sheet!

## 3.8 | Representations of Solutions

• This concept tests if you can read diagrams and illustrations of solutions, such as the ones found in the 2019 FRQ’s questions 1B and 3B.

## 3.9 | Separation of Solutions and Mixtures Chromatography

• The methods used to separate solutions take advantage of the IMF between molecules in the solution.
• Chromatography comes in several forms. Paper chromatography takes advantage of the differing polarities of the mixture’s components to separate the solution; a significant sample of the components cannot be collected. Column chromatography also takes advantage of different polarities and attraction, but a significant sample can be collected.
• Distillation works by using the differing boiling and vapor pressures of the components of mixtures. It is possible to collect a significant sample of the separated components with this method.

## 3.10 | Solubility

• Remember like attracts like: Substances with similar IMF tend to be the most soluble in one other.
• For example, a polar solute will dissolve more readily in a polar solvent than a non-polar one.

## 3.11 | Spectroscopy and the Electromagnetic Spectrum

• There are three major types of radiation on the Electromagnetic Spectrum that the College Board wants you to know: microwave, infrared, and ultraviolet.
• Microwave radiation increases the rotation of the molecule.
• Like how microwaves turn your food around!
• Infrared radiation increases the vibration of the molecule.
• Ultraviolet radiation causes transitions of electrons in the energy levels in the molecule.

## 3.12 | Photoelectric Effect

• The energy of a photon equals the frequency of the electromagnetic wave times Planck’s constant.
• (E = ℎν)
• The speed of light equals the wavelength of the electromagnetic wave times the frequency of the wave.
• (c = λν)
• These constants and equations are all on your formula sheet.

## 3.13 | Beer-Lambert Law

• The equation for the Beer-Lambert Law is:
• A = εbc
• A = Absorbance
• ε = molar absorptivity constant
• b = length of the path that the light travels through
• c = concentration or molarity.
• In an experiment where ε and b are kept constant, A is directly proportional to c.
• The process undertaken to measure the concentration of a solution is called Spectrophotometry.
• Occasionally, there are errors in the measuring and the College Board will want you to know why the errors happen.
• If the measured absorbance or concentration is lower than expected:
• The solution could have been diluted.
• If the measured absorbance or concentration is higher than expected:
• There could have been fingerprints on the vials used in the experiment.
• There could be contamination from a previously-measured higher concentration solution.

Happy studying! 😀

[Sources: Mrs. Smith’s Chemistry Notes, College Board Course and Exam Description]