AP Chemistry Unit 3 Guided Notes: Properties of Substances and Mixtures

Topic 3.1 – Intermolecular and Interparticle Forces

Intermolecular Forces (IMFs) * Definitions: The forces that make molecules stick together. * Properties: Usually affect physical properties, such as melting point, boiling point, etc. * Strength: Normally weaker than covalent and ionic bonds.
Intramolecular Forces * Definitions: The forces that hold a molecule or ionic compound together internally. * Examples: Covalent and Ionic Bonds. * Properties: Usually affect chemical properties. * Strength: Usually much stronger than intermolecular forces.
London Dispersion Forces (LDFs) * Visualization (Neon Atom Example): In a typical electron distribution, electrons are always moving. At some point, the distribution of electrons becomes uneven or "lopsided." * Instantaneous Dipole: The temporary, asymmetrical distribution of electrons in an atom or molecule. * Induced Dipole: A temporary dipole created in a neighboring molecule caused by the instantaneous dipole of a nearby molecule. * Universal Presence: All molecules exhibit London dispersion forces. * Nonpolar Molecules: These exhibit ONLY London dispersion forces. * Polarizability: The more electrons a molecule has, the more polarizable it is, which results in stronger London dispersion forces. * Relationship to Boiling Point: The stronger a molecule’s intermolecular forces, the higher its boiling point (and melting point) will be. * Relative Strength: Of all the IMFs, London dispersion forces are generally considered the weakest. However, some large molecules have so many electrons and are so polarizable that their IMFs end up being stronger than those of some smaller polar molecules. * Examples for practice: * Both $CH_4$ and $C_4H_{10}$ exhibit London dispersion forces because all molecules do. * Ranking nonpolar molecules by boiling point: He < Ar < C_3H_8 (increasing with number of electrons/size).
Dipole-Dipole Forces * Visualization (HCl Molecule): Since Chlorine ( $Cl$ ) is much more electronegative than Hydrogen ( $H$ ), the electron distribution is uneven, making it a polar molecule. * Mechanism: The polarity causes attractions between the positive end of one molecule and the negative end of another. * Occurrence: Only polar molecules exhibit dipole-dipole forces. * Combined Forces: Polar molecules exhibit BOTH London dispersion forces and dipole-dipole forces. * Relative Strength: Generally, dipole-dipole forces are stronger than London dispersion forces for molecules of similar size. Given similar numbers of electrons, the polar molecule should have a higher boiling point than the nonpolar molecule. * Example: $NCl_3$ (polar) has a higher boiling point than $C_4H_{10}$ (nonpolar) if we assume dipole-dipole dominance, though large nonpolar molecules can sometimes override this.
Hydrogen Bonding * Visualization (Water Molecule): Oxygen is a highly electronegative atom, causing the $H_2O$ molecule to be very polar. * Mechanism: This is a special, very strong type of dipole-dipole force. It occurs because Hydrogen is a very small atom, allowing a neighboring molecule’s electronegative atom ( $F$ , $O$ , or $N$ ) to get very close. * Requirements: Found in molecules containing an $H-F$ , $H-O$ , or $H-N$ bond. * Relative Strength: Generally stronger than London dispersion or simple dipole-dipole forces. * Solubility: Explains why polar molecules with $F$ , $O$ , or $N$ usually dissolve in water; they can hydrogen bond with water molecules. * Examples for practice: * $NH_3$ exhibits hydrogen bonding, while $NO_2$ only exhibits dipole-dipole. * $NH_3(g)$ dissolves in water because it forms hydrogen bonds with water; $PH_3(g)$ does not because it lacks the highly electronegative atom required for hydrogen bonding. * Ranking by boiling point: CO_2 < NO_2 < H_2O.
Ionic Forces * Visualization (Magnesium Oxide): Ions stick together because of electrostatic forces. * Strength: Ionic forces are the strongest of the discussed interparticle forces, leading to the highest melting and boiling points. * AP Chemistry Context: While often categorized separately from IMFs, it is crucial to recognize they are stronger than LDFs, dipole-dipole, and H-bonds.
Summary Hierarchy of Forces 1. Ionic Forces (Strongest). Tiebreaker: Charge differential ( $Q_1Q_2$ ), then ionic size ( $r^2$ ). 2. Hydrogen Bonds (Second strongest: $H-F$ , $H-O$ , $H-N$ ). 3. Dipole-dipole forces (Polar molecules). 4. London Dispersion Forces (Weakest; most dominant in nonpolar molecules). Tiebreaker: Higher number of electrons leads to higher polarizability and stronger forces.

Topic 3.2 – Properties of Solids

Vapor Pressure * Definition: The pressure exerted by a gas back down on the liquid from which it evaporated. * Temperature Correlation: As temperature increases, the vapor pressure of a liquid increases. * IMF Correlation: Liquids with weaker intermolecular forces will have higher vapor pressures (as they evaporate more easily). * Boiling Point Correlation: Liquids with weaker intermolecular forces will have a lower boiling point.
Ionic Solids * Structure: Composed of ionic compounds in a crystal lattice. * Physical State: Brittle and can be easily shattered. * Conductivity: Only conduct electricity when dissolved in water ( $aq$ ) or melted ( $l$ ).
Molecular Solids * Structure: Composed of nonmetal covalent compounds (individual molecules). * Strength: The forces holding these molecules together are relatively weak. * Properties: Normally have somewhat low boiling and melting points. * Examples: Ice ( $H_2O$ ), naphthalene (moth balls), many polymers.
Covalent Network Solids * Structure: A continuous network of covalent bonds creates an extremely rigid structure. * Properties: Very difficult to melt; extremely high melting points and boiling points. Bonds are often oriented in multiple directions. * Examples: Diamond and graphite (pure $C$ ), silicon dioxide ( $SiO_2$ , such as quartz/sand), and silicon carbide ( $SiC$ ). * Comparison: $CO_2$ has a low boiling point because it is a molecular solid held by weak IMFs; $SiO_2$ has a high boiling point because it is a covalent network solid.

Topic 3.3 – Solids, Liquids, and Gases

Solids * True solids are crystalline in structure (orderly, repeating 3D patterns). * Amorphous materials lack this long-range order (e.g., glass, plastic).
Gases * Structure: Gas molecules are far apart from each other. * Compressibility: Because of the distance between particles, gases can be squeezed closer together (compressed). * Container: A gas is the only main state of matter that can expand to fill a container.
Four Properties of Gases 1. Pressure ( $P$ ): Determined by how often the gas molecules collide with the container walls. 2. Volume ( $V$ ): The amount of space the gas takes up. 3. Temperature ( $T$ ): A measure of the average kinetic energy ( $mass \times velocity$ ) of the molecules. 4. Number of Gas Molecules ( $n$ ): Usually measured in moles.

Topic 3.4 – Ideal Gas Law

Fundamental Gas Laws * Boyle’s Law: The volume of a gas is inversely proportional to its pressure. $P_1V_1 = P_2V_2$ . * Charles’ Law: The volume of a gas is directly proportional to its Kelvin temperature. $\frac{V_1}{T_1} = \frac{V_2}{T_2}$ . * Gay-Lussac’s Law: The pressure of a gas is directly proportional to its Kelvin temperature. $\frac{P_1}{T_1} = \frac{P_2}{T_2}$ . * Combined Gas Law: $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$ .
Ideal Gas Law * Formula: $PV = nRT$ . * Variables: $P$ (Pressure), $V$ (Volume), $n$ (moles), $R$ (Ideal Gas Constant), $T$ (Temperature in Kelvins).
Dalton’s Law of Partial Pressures * Concept: The total pressure of a mixture of gases is the sum of the pressures each gas would exert if it were alone. * Calculation with Mole Fraction: $P_{Gas} = X_{Gas} \times P_{total}$ , where $X$ is the mole fraction. * Summation Formula: $P_{Total} = P_1 + P_2 + …$ * Collection over Water: When a gas is collected over water, the pressure of the dry gas must be calculated by subtracting the vapor pressure of water from the total atmospheric pressure: $P_{gas} = P_{total} - P_{H_2O}$ .

Topic 3.5 – Kinetic-Molecular Theory

Maxwell-Boltzmann Distribution Curve * Observations: At higher temperatures, more molecules have higher speeds and, thus, greater kinetic energy. * Area: The area underneath each curve is the same because the total number of molecules remains constant. * Curve Shift: At higher temperatures (e.g., 750 K), the curve becomes flatter and shifts to the right.
Heat & Temperature * Temperature: A numerical measure of the average kinetic energy in the particles of a material. Measured in Kelvins. If kinetic energy doubles, the Kelvin temperature doubles. * Heat: The form of energy transferred between two systems at different temperatures; also called thermal energy. Measured in Joules ( $J$ ).
Graham’s Law of Effusion * Effusion: The escape of gas molecules through a tiny hole in a material. * Concept: More massive (heavier) molecules move more slowly than less massive (lighter) molecules. * Formula: $\frac{Rate_A}{Rate_B} = \sqrt{\frac{M_B}{M_A}}$ . * Example: Helium effuses faster than Nitrogen because it has a smaller molecular mass.

Topic 3.6 – Deviation from Ideal Gas Law

Ideal Gas Assumptions * Molecules have no volume. * Molecules have no attraction or repulsion for each other.
Real Gas Behavior * Real gases can deviate from ideal behavior under high pressure and low temperature. * Small nonpolar molecules ( $He$ , $H_2$ , $Ne$ ) are closest to ideal behavior because they take up very little volume and have weak attractions. * Van der Waals Equation: $(P + \frac{n^2a}{V^2})(V - nb) = nRT$ . * Factor $a$ : Corrects for intermolecular forces of attraction. * Factor $b$ : Corrects for the volume of the gas molecules themselves.

Topic 3.7 – Solutions and Mixtures

Types of Mixtures * Heterogeneous: Individual components are visible to the naked eye (e.g., oil and water). * Homogeneous: Components are mixed together evenly; individual components cannot be seen. Also called solutions.
Molarity ( $M$ ) * Definition: The primary measure of solution concentration in chemistry. * Formula: $Molarity = \frac{moles\,of\,solute}{liters\,of\,solution}$ . * Calculations involve converting mass to moles and using the total volume of the solution.
Dilution * Equation: $M_1V_1 = M_2V_2$ . * Safety Note: Never add water directly to concentrated acid; always add acid to water.

Topic 3.8 – Representations of Solutions

Components of a Solution * Solvent: The medium into which a solute is dissolved. Water is the "universal solvent." * Solute: A substance that dissolves into a medium.
Solution States * Dilute: Relatively small amount of solute. * Concentrated: Relatively large amount of solute. * Saturated: Contains the maximum amount of solute for that temperature. * Supersaturated: Temporarily contains more than the maximum amount of solute for that temperature.
Mole Fraction ( $X$ ) * Formula: $X_{substance} = \frac{moles_{substance}}{moles_{total}}$ .
Electrolytes * Dissociation: Ionic compounds dissolve when polar water molecules surround and pull ions away from the crystal lattice. * Ion-dipole Force: Dissolution occurs only if the ion-dipole force is greater than the electrostatic force holding the compound together. * Strong Electrolytes: Conduct electricity strongly. Includes soluble ionic compounds (e.g., $NaCl$ , $Zn(NO_3)<em>2$ ), strong acids (7 total), and strong bases (8 total). * Weak Electrolytes: Ionize slightly; conduct poorly. Includes weak acids and weak bases (e.g., $NH_3$ ). * Nonelectrolytes: Hardally ionize; do not conduct. Includes molecular compounds like sugar ( $C</em>{12}H_{22}O_{11}$ ) and water.

Topic 3.9 – Separation of Solutions and Mixtures

Distillation: Separates components based on differences in boiling points.
Chromatography: Separates components based on differences in polarity and their affinity for a stationary vs. mobile phase.

Topic 3.10 – Solubility

General Principle: "Like dissolves like" * Polar solutes dissolve in polar solvents ( $H_2O$ and $C_2H_5OH$ ). * Nonpolar solutes dissolve in nonpolar solvents ( $CH_4$ and $C_6H_6$ in $CCl_4$ ).
Solubility and Temperature * Solids: Solubility usually increases as temperature increases. * Gases: Solubility decreases as the solution gets warmer (e.g., thermal pollution in lakes).
Solubility and Pressure (Gases) * Henry’s Law: The solubility of a gas in a liquid is directly proportional to the pressure of the gas above the solution.

Topic 3.11 – Spectroscopy and the Electromagnetic Spectrum

Planck’s Postulate: Electromagnetic energy (light) consists of discrete packets called quanta (photons).
Radiation Effects on Matter * Ultraviolet (UV) or Visible Light: Causes electrons to transition between energy levels. * Infrared (IR) Radiation: Causes molecules to vibrate. * Microwave Radiation: Causes molecules to rotate.

Topic 3.12 – Properties of Photons

Wave Speed: All electromagnetic waves move at the speed of light ( $c = 3.00 \times 10^8\,m/s$ ).
Wavelength and Frequency: $c = \lambda \nu$ .
Energy of a Photon: $E = h \nu$ . ( $h = 6.626 \times 10^{-34}\,J \cdot s$ ).
Example Calculation: To find energy from wavelength, first find frequency ( $\nu = \frac{c}{\lambda}$ ), then find energy ( $E = h\nu$ ).

Topic 3.13 – Beer-Lambert Law

Methodology: Spectroscopy uses electromagnetic radiation to determine a solution's concentration based on the intensity of light absorbed.
Beer-Lambert Law Formula: $A = \epsilon b c$ . * $A$ : Absorbance (unitless). * $\epsilon$ : Molar absorptivity (constant for a substance at a specific wavelength). * $b$ : Path length (usually constant at $1\,cm$ ). * $c$ : Concentration (Molarity).
Direct Proportionality: Absorbance is directly proportional to concentration ( $A \propto c$ ).
Standard Curve: A graph of Absorbance vs. Concentration is used to find the unknown concentration of a sample.
Sources of Error: * Fingerprints on cuvette or contamination: Can increase absorbance. * Distilled water in the cuvette (dilution): Decreases absorbance. * Incorrect path length: Longer cuvettes increase absorbance.
Colored Ions: 1. Copper ions ( $Cu^{2+}$ ): Blue. 2. Nickel ions ( $Ni^{2+}$ ): Green. 3. Iron ions ( $Fe^{3+}$ ): Yellow/Brown (Standard FeSCN complex is dark red). 4. Chromium ions: Orange/Yellow. 5. Cobalt ions ( $Co^{2+}$ ): Pink/Red.