AP Chemistry Unit 3: Intermolecular Forces and Properties
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3.1 Intermolecular Forces
Intramolecular Forces
Bonds that exist within a molecule are always more potent than intermolecular forces.
Different from Intermolecular Forces
Intermolecular Forces (IMFs)
Attractive or repulsive forces between molecules.
Force
Strength
LDFs
Weakest*
Dipole-Dipole Forces
Intermediate
Hydrogen Bonding
Strong
Ion-Dipole Forces
Strongest
London Dispersion Forces (LDFs)
Also rarely known as Van der Waals Forces, LDFs are the weakest type of IMFs that exist in all molecules.
LDFs are Instantaneous/temporary dipoles, where, for an instant, a molecule would have a particle negative and positive side and become polar. This is because electrons move around.
Nonpolar Molecules only have LDFs
The strength of LDFs can increase depending on the size of a molecule and can be even stronger than Dipole-Dipole Forces, sometimes even Hydrogen Bonding! This is because, with more electrons, there will be a stronger instantaneous dipole because of a polarizable electron cloud.
Polarizability, or the ability to distort a molecule's electron cloud that creates temporary dipoles that increase the IMFs. Essentially, the ability of an atom to form a dipole.
Dipole-Dipole Forces
Dipole-dipole forces exist from permanent (induced) dipoles. The more polar the molecule, the stronger the Dipole-Dipole force and the decreased distance between the dipoles.
All Polar Covalent Molecules have Dipole-Dipole forces.
Dipole-Dipole Attraction
The positive dipole end of a molecule is attracted to the negative end and vice versa.
Hydrogen Bonds
A very strong type of Dipole-Dipole force that, contrary to its name, is not a bond.
A dipole-dipole attraction in polar molecules where hydrogen is bonded to a highly electronegative atom, Oxygen (O), Nitrogen (N), and Florine (F).
Hydrogen MUST be chemically (directly) bonded to N, O, or F for it to count as a Hydrogen bond.
Ion-Dipole Forces
Ionic compounds with water have Ion-Dipole forces.
In solutions, oxygen (in the water) surrounds the positively charged ion, and hydrogen surrounds the negatively charged ion.
The strength increases as the charge on the ion increases or if the magnitude of the dipole of the polar molecule increases.
Substances with weak LDFs are found in the gas phase at STP (standard temperature and pressure)
Weakest to Strongest IMFs Gas → Liquid → Solid
Boiling/Melting Points
A higher boiling/melting point means the molecule has stronger IMFs.
Volatility
Volatility, the tendency of a substance to vaporize (transition to the gas phase), is higher (more likely) with weaker IMFs. Essentially, quick evaporation means the substance's IMFs are weak.
Viscosity
Viscosity, the measure of a fluid's resistance to flow, is lower with weak IMF forces.
Surface Tension
Liquids with high surface tension have strong IMFs.
Vapor Pressure
Vapor Pressure, the pressure exerted by a substance's vapor, is lower with strong IMFs.
3.2 Properties of Solids
Amorphous Solids
Solids with considerable disorder in their structures.
Crystalline solids are arranged in a geometrical pattern as a crystal lattice.
Unit Cell
The smallest repeating unit that makes up a crystal lattice
Types of Crystalline Solids
Ionic Solids
Has ions at the lattice points, which are held strongly together by cations and anions, resulting in high melting and boiling points.
The ions repel each other, making ionic solids hard and brittle. This allows one layer to slide off another.
Molecular Solids
Held together by IMFs with covalently bonded molecules at each lattice point.
Most covalent compounds without dipoles are gases at room temperature, but they can exist in solid form if they are larger molecules (polarizability).
Metallic Solids
Composed of metal atoms with delocalized (free to move around) electrons that are metallically bonded.
Covalent-Network Solids
Atoms are held together very strongly in large networks by covalent bonds, giving them the properties of being harder and having higher melting points.
Most common forms of carbon:
Diamond
Has strong sigma bonds in an interlocked tetrahedral shape of sp^3 hybridized, making it the hardest naturally occurring substance
Excellent insulators because the space between the orbitals makes it very hard for electrons to move around.
Graphite
The carbons are connected to three others and are sp^2 hybridized. The molecules are bonded with pi bonds, causing the layers to slide by each other. The electrons are free to move, making graphite an excellent conductor.
3.3 Solids, Liquids, and Gases
Compressibility
The degree to which a material's volume can be changed under pressure
Solids
Solids have definite shape and volume. They have limited particle moments with strong intermolecular forces.
Liquids
Low compressibility, very unrigid, and highly dense, which allows liquids to fill the space provided by the container they’re in.
Gases
Gases do not have a definite shape or volume, so they fill the container they're in, and the particles are constantly moving with enough energy not to be affected by intermolecular forces.
Temperature is constant. Pressure and volume are inversely related. Source: bismarckschools
Charles's Law
Pressure is constant. Volume and temperature are directly proportional. Source: bismarckschools
Gay-Lussac’s Law
Pressure is constant. Pressure and temperature are directly proportional.
Avogadro’s Law
= Moles
Volume and # of moles are directly proportional.
Dalton's Law of Partial Pressures
The sum of all partial pressures of each gas in a mixture of gasses is equal to the total pressure.
= Total Pressure
= Partial pressure of the specific gas
Calculating Partial Pressure
= Partial Pressure of the Specific Gas
= Mole Fraction ()
= Total Pressure
3.5 Kinetic Molecular Theory
Kinetic Molecular Theory (KMT)
There are four main principles (postulates):
No forces between particles
Gas particles move in random, constant, straight lines
Particles are separated by great distances relative to their volume of ≈ 0
Collisions of gas particles transfer energy with no energy lost
Average Kinetic Energy of Gaseous Particles
KE= Kinetic Energy (in joules)
m = Mass of the molecule (in kg)
v= Speed of the Molecule (m/s)
Maxwell-Boltzmann Distribution
Also known as Boltzmann Distribution
Shows the relationship between temperature and average kinetic energy Source: janosh.dev
X-axis: Particle Speed
Y-axis: Number of particles
A higher peak means that a higher number of particles have that speed
The area under each curve has an equal amount of gas particles. Only the particle speed changes.
3.6 Deviation from Ideal Gas Law
Real gases are not ideal and do have forces between them
Reasons Gas deviates from the Ideal Gas Law
High Intermolecular Force(s)
The greater the IMF force, the more it deviates.
Highly polar molecules and larger molecules deviate more.
Low Temperatures
A lower temperature causes the gas to deviate.
A higher temperature causes the gas to be the most ideal.
High Pressure
A lower pressure causes the gas to be the most ideal.
A higher pressure causes the gas to deviate.
3.7 Solutions and Mixtures
Mixtures: Physically made up of 2 or more substances
Homogeneous Mixtures (A Solution)
Composed of 2(+) substances and is Uniform in Composition (made with the same thing throughout)
Usually inseparable
Heterogeneous Mixtures
Composed of 2(+) elements or compounds and is Ununiform in Composition (not made with the same thing throughout, like air)
Usually separable
Solutions (A Homogeneous Mixture)
Composed of 2(+) substances and is Uniform in Composition (made with the same thing throughout)
Usually inseparable
Molarity
Molarity (M) =
Numerator (top) = Moles of solute, referring to the amount of substance dissolved.
Denominator (bottom) = Liters of solvent, referring to the volume of the solution
Defined as the number of moles of a solute dissolved in one liter of solvent.
3.8 Representations of Solutions
Water Molecule (H2O)
A polar molecule with a positive partial charge on the oxygen side and a negative partial charge on the hydrogen side.
Solution Formation Visually
Example: NaCl
The water molecules are oriented properly, with the charges facing the appropriate charge.
The atom & molecule sizes are drawn appropriately, with the anion larger than the cation (usually true).
Drawn with the correct ratios: 1 Na means 1 Cl In the example, there are 2 Na and 2 Cl, which is correct. However, you could draw 20 Na and 20 Cl as long as the ratio is constant.
3.9 Separation of Solutions and Mixtures Chromatography
Separating Mixtures (Explained in depth in 3.9)
Physical changes that can be made to separate mixtures into pure substances:
Distillation
The process of separating two or more liquids in a solution through vaporization of the liquid with the lowest boiling point, then condensing it back to its pure liquid state. This would separate the liquid with the lowest boiling point from the rest of the component(s)
Used for mixtures that contain undissolved (insoluble) solids in the liquid by filtering the mixture through a filtering agent (i.e., mesh), leaving the solids on the filtering agent.
Chromatography
Method of separation is usually used to separate dyes/pigments where the components are distributed between mobile and stationary phases.
A more attracted (more similar) sample to the mobile phase causes it to travel further.
Paper Chromatography
Stationary Phase: Paper
Mobile Phase: Solvent that moves through the paper
Uses: Separating pigments, inks, and amino acids Source: wikipedia
Thin-Layer Chromatography
Stationary Phase: Thin layer of absorbent material coated onto a plate
Mobile Phase: Solvent that moves through the paper
Uses: Testing purity, monitoring reactions, and identifying compounds Source: microbenotes
Column Chromatography
Stationary Phase: Packed column of solid material
Mobile Phase: Solvent that moves through the paper
Substances with similar intermolecular interactions tend to be miscible (the ability of one liquid to dissolve in another entirely) or soluble (the ability to which a substance dissolves in a solvent) in one another.
“Like dissolves like”
It is a phrase you will commonly hear but is only used to understand solubility conceptually. Do notuse it to explain something on a test.
Polarity’s effect on Solubility
Ionic solutes usually dissolve in polar solvents.
Polar molecular solutes usually dissolve in polar solvents.
Nonpolar molecular solutes usually dissolve in nonpolar solvents.
3.11 Spectroscopy and the Electromagnetic Spectrum
Amplitude
Verticle height of a wave from its midline, determines the light wave's intensity (brightness).
Causes transitions in molecular vibrational levels
Infrared Spectroscopy
X-axis: Wavelength (µm or cm^-1)
Y-axis: Intensity of Transmitted Radiation
Molecules absorb IR radiation, which causes their bonds to vibrate. These vibrations can cause stretching (change of bond lengths) or bending (change of bond angles).
Dips in the graphs, called Absorption Peaks, indicate the wavelength where the molecule absorbs IR radiation.
Stronger bonds and lighter atoms absorb IR radiation at higher frequencies.
Microwave Radiation
Causes transitions in molecular rotational levels
Ultraviolet Rays (UV)
Visible light/radiation that causes electrons to go into their excited state, which transitions electronic energy levels (when an electron moves from one energy level to another)
3.12 Photoelectric Effect
Photoelectric Effect (Properties of Photons)
Light waves hitting a metal surface will cause photons to be ejected at specific energies.
When an atom or molecule absorbs a photon, the energy of that atom or molecule increases by the same amount of energy as the photon. Consequently, when a photon is emitted (released), the atom or molecule decreases by the same amount of energy as the photon.
c = λ ν
The wavelength of the electromagnetic wave is related to its frequency and the speed of light shown in this equation.
c = Speed of Light (m/s)
λ = Wavelength (m)
Often, the questions will give you the wavelength in nanometers. Remember to convert it to meters.
ν= Frequency ()
It is not actually a v, but instead the Greek symbol for “new.”
Planck’s Equation
The energy of a photon is related to the frequency of the electromagnetic wave through this equation.
E = Change in Energy (J)
h = Planck’s Constant (6.626 Js)
ν = Frequency in
c = Speed of Light ( 2.998 m/s)
λ = Wavelength (m)
Often, the questions will give you the wavelength in nanometers. Remember to convert it to meters.
3.13 Beer-Lambert Law
The Beer-Lambert Formula
A = εbc
Relates the amount of light a solution absorbs to the solution's concentration
A = Absorbance (Amount of Light absorbed)
Ε = Molar absorptivity (How intensely light of a specific wavelength is absorbed in L/molcm)
b = Path Length (Distanced the light traveled in cm)
c = Concentration (mol/L)
Spectrometer
A tool used to measure and analyze the spectrum of light scattered by materials allows for the identification of the substance.
Cuvette
A small transparent rectangular container that holds the sample/solution.
Absorbance
The amount of light a sample absorbs at different wavelengths.
Fingerprints scatter light, causing less light to pass through the solution and reach the detector.
= Moles 
= Total Pressure
= Partial pressure of the specific gas
= Partial Pressure of the Specific Gas
= Mole Fraction (
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), which is given on your equations sheet
m/s)
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Js)
m/s)
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