Michael Farabaugh 2026 AP Chemistry Practice Exam Study Guide

General Information and Exam Conditions

• Total Content: Michael Farabaugh's 2026 AP Chemistry Practice Exam (Multiple Choice Section).

• Total Questions: 60 Multiple Choice Questions.

• Time Limit: 1 hour and 30 minutes.

• Standard Conditions: Unless otherwise specified, calculations assume a temperature of $298\,K$, a pressure of $1.0\,atm$, and aqueous solutions.

• Formula References: All calculations and justifications utilize standard chemical principles, including the Ideal Gas Law ($PV = nRT$), thermodynamics ($\Delta G = \Delta H - T\Delta S$), and stoichiometry.

Gases and Kinetic Molecular Theory

• Partial Pressure in Mixtures:     * In a mixture of $Ar(g)$ and $O_2(g)$, the partial pressure of a gas is proportional to its mole fraction ($X_i$).     * Formula: $P_i = X_i \times P_{total}$.     * Example: A particle diagram shows 4 $Ar$ atoms and 2 $O_2$ molecules (total 6 particles). With a total pressure of $3.0\,atm$, the partial pressure of $O_2$ is $P_{O2} = \frac{2}{6} \times 3.0\,atm = 1.0\,atm$.

• Addition of Non-Reactive Gases:     * Adding an inert gas (like $Ne$) to a rigid container at constant temperature increases the total pressure ($P_{total}$) but does not change the partial pressures of the existing gases ($P_i$) because the moles, temperature, and volume for those specific gases remain constant.

• Deviations from Ideal Behavior:     * Gases deviate most from ideal behavior at high pressures and low temperatures.     * Molecular factors: High intermolecular forces (IMF) and large molecular volumes increase deviations. $CF_4$ is more likely to deviate than $H_2$, $O_2$, or $CH_4$ due to having a larger, more polarizable electron cloud resulting in stronger London dispersion forces.

• Maxwell-Boltzmann Distribution:     * Particle speed distributions depend on temperature and molar mass.     * At the same temperature, heavier gases have a lower average speed and a narrower distribution (Curve 1) compared to lighter gases (Curve 2).     * Example: $Ar(g)$ ($MW \approx 40\,g/mol$) at $300\,K$ would be represented by a curve shifted left compared to $Ne(g)$ ($MW \approx 20\,g/mol$) at $300\,K$.

Kinetics and Reaction Rates

• Factors Affecting Rate:     * Concentration: Increasing the concentration of reactants (e.g., $HCl(aq)$) generally increases the rate and decreases the time to completion.     * Surface Area: Fine powders react faster than single chunks due to increased contact area.     * Temperature: Higher temperatures increase kinetic energy and collision frequency.

• Rate Laws:     * For a second-order reaction (e.g., $2NO_2(g) \rightarrow 2NO(g) + O_2(g)$), the rate is proportional to the square of the concentration: $\text{Rate} = k[NO_2]^2$.     * If concentration triples ($0.020\,M \rightarrow 0.060\,M$), the rate increases by a factor of $3^2 = 9$.

• Reaction Mechanisms:     * The rate-determining step (RDS) is the slowest step, represented by the highest activation energy peak on an energy profile.     * Thermodynamics of the overall reaction depends on the energy difference between reactants and products ($\Delta H_{obs} = E_{products} - E_{reactants}$).

• Half-life ($t_{1/2}$):     * A constant half-life is a signature of first-order kinetics.     * Rate constant ($k$) for first-order reactions has units of $s^{-1}$.

Atomic Structure and Periodic Trends

• Ionization Energy:     * First ionization energy generally increases across a period (left to right) and decreases down a group.     * Example: $Sr$ has a larger first ionization energy than $Rb$, $Cs$, or $Ba$.

• Atomic and Ionic Radii:     * Radii increase down a group and decrease across a period due to effective nuclear charge ($Z_{eff}$).     * Anions (e.g., $S^{2-}$) are larger than their neutral atoms; cations are smaller.     * Order: Cl < S < S^{2-}.

• Photoelectron Spectroscopy (PES):     * Peak position (binding energy) corresponds to the attraction of electrons to the nucleus.     * $1s$ electrons are closest to the nucleus and have the highest binding energy (peaks appearing further left on the x-axis).

• Mass Spectrometry:     * Determines relative abundance of isotopes.     * Example: If $^{85}Rb$ has a $72\%$ abundance and $^{87}Rb$ has a $28\%$ abundance, the peak at 85 will be roughly 2.5 times higher than the peak at 87.

Chemical Bonding and Intermolecular Forces

• VSEPR and Geometry:     * T-shaped Geometry: Results from 5 electron domains ($AX_3E_2$), such as in $ClF_3$.

• Lattice Enthalpy:     * Strength of ionic bonds in a solid crystal. Calculated via Coulomb's Law: $F \propto \frac{q_1 q_2}{r^2}$.     * Higher charges ($Mg^{2+}$ vs $Na^+$) and smaller radii ($F^-$ vs $Br^-$) result in stronger attractions. $MgF_2$ has stronger attractions than $NaF$, $NaBr$, or $MgBr_2$.

• Resonance and Bond Order:     * Molecules with resonance (like $NO_2^-$) have intermediate bond orders. If there is one single bond and one double bond in two resonance structures, the bond order is $1.5$.

• Intermolecular Forces (IMFs) and Boiling Points:     * London Dispersion Forces: Present in all molecules; strength increases with polarizability (larger electron clouds). $CH_3I$ (BP $316\,K$) has a higher boiling point than $CH_3Cl$ (BP $195\,K$) because Iodine has a larger, more polarizable electron cloud.     * Dipole-Dipole: Present in polar molecules (e.g., $CH_2F_2$).     * Hydrogen Bonding: Occurs when H is bonded to N, O, or F. This significantly raises boiling points (e.g., alcohols vs ethers).

• Chromatography:     * Separation based on differential affinity for stationary and mobile phases.     * In a polar stationary phase with a nonpolar solvent, a nonpolar molecule (like benzene) will travel faster/further than a polar molecule (like phenol).

Thermochemistry and Thermodynamics

• Enthalpy (\Delta H):     * Bond breaking is endothermic (\Delta H > 0); bond formation is exothermic (\Delta H < 0).

• Standard Enthalpy of Formation (\Delta H_f^o):     * $\Delta H_{rxn}^o = \Sigma \Delta H_f^o (products) - \Sigma \Delta H_f^o (reactants)$.     * For $CCl_4(g) + 4HF(g) \rightarrow CF_4(g) + 4HCl(g)$, with $\Delta H^o = -106\,kJ/mol_{rxn}$, $\Delta H_f^o$ for $CCl_4(g)$ is calculated to be $-103\,kJ/mol$.

• Entropy (\Delta S):     * Related to disorder. Processes increasing moles of gas have \Delta S > 0.     * If moles of gas remain constant ($CH_4(g) + 2O_2(g) \rightarrow CO_2(g) + 2H_2O(g)$), $\Delta S$ is close to zero.

• Gibbs Free Energy and Favorability:     * $\Delta G^o = \Delta H^o - T\Delta S^o$.     * Thermodynamically favorable reactions have \Delta G^o < 0.     * If \Delta H^o < 0 and \Delta S^o < 0, the reaction is favorable at low temperatures but unfavorable at high temperatures (e.g., $NH_3(g) + HCl(g) \rightarrow NH_4Cl(s)$).

• Specific Heat Calculations:     * $q = mc\Delta T$.     * Example: A $71.0\,g$ sample of $Al$ absorbing $6.24\,kJ$: $\Delta T = \frac{6240\,J}{71.0\,g \times 0.896\,J/g\cdot^oC} \approx 98.1^oC$. Final temperature: $22.0 + 98.1 = 120.1^oC$.

Chemical Equilibrium

• Equilibrium Constants ($K_p$ and $K_c$):     * Expression excludes solids and liquids. For $CH_3NH_2(aq) + H_2O(l) \rightleftharpoons CH_3NH_3^+(aq) + OH^-(aq)$, $K = \frac{[CH_3NH_3^+][OH^-]}{[CH_3NH_2]}$.

• The Reaction Quotient ($Q$):     * If Q < K, the reaction shifts right (toward products).     * If Q > K, the reaction shifts left (toward reactants).

• Le Chatelier’s Principle:     * Increasing temperature for an exothermic reaction (\Delta H < 0) shifts equilibrium to the left, increasing reactants and decreasing the value of $K_p$.

• Thermodynamics and Equilibrium:     * $\Delta G^o = -RT\ln(K)$.     * Large $K$ values (K > 1) correspond to positive $E^o_{cell}$ and negative $\Delta G^o$.

Acids, Bases, and Solubility

• Weak vs. Strong Acids:     * For a strong acid, $pH = -\log[Acid]$. If observed pH > -\log[Acid], the acid is weak.     * Sample calculation: For a $0.20\,M$ acid with $pH = 2.7$, $[H^+] = 10^{-2.7} \approx 2 \times 10^{-3} M$. $K_a = \frac{(2 \times 10^{-3})^2}{0.20} = 2 \times 10^{-5}$.

• Titrations:     * The equivalence point volume depends on the total moles of acid ($n = M \times V$).     * At the half-equivalence point, $pH = pK_a$.     * Doubling the titrant concentration (e.g., $0.0750\,M \rightarrow 0.150\,M\,NaOH$) halves the volume required to reach the equivalence point.

• Buffers:     * Buffer capacity depends on the concentration of the buffer components. A $1.0\,M$ buffer has a higher capacity and smaller $pH$ change than a $0.10\,M$ buffer when a base is added.

• Salt pH:     * Salts of weak acids (like $NaF$) are basic because the conjugate base ($F^-$) hydrolyzes water to produce $OH^-$: $F^- + H_2O \rightleftharpoons HF + OH^-$.

• Solubility Product ($K_{sp}$):     * Solubility (molarity of a saturated solution) is an intensive property and does not change with volume. However, the total number of moles of dissolved ions increases with volume in a saturated solution.

Stoichiometry and Solutions

• Limiting Reactants:     * Determined by the mole ratio in the balanced equation.     * Example: $4Cr + 3O_2 \rightarrow 2Cr_2O_3$. Given $52\,g\,Cr$ ($1.0\,mol$) and $32\,g\,O_2$ ($1.0\,mol$). $1.0\,mol\,Cr$ requires $0.75\,mol\,O_2$. Since we have excess $O_2$, $Cr$ is the limiting reactant. Maximum $Cr_2O_3$ produced is $0.5\,mol$, or $76\,g$.

• Particulate Models of Solutions:     * Ion-dipole interactions: Positive ions ($K^+$) are surrounded by the oxygen ends (partial negative) of water; negative ions ($Cl^-$) are surrounded by the hydrogen ends (partial positive).

• Percent Composition:     * $\text{Mass \%} = \frac{\text{Mass of element}}{\text{Total mass of compound}} \times 100$.     * $NH_3$ has a higher $\text{N \%}$ ($\approx 82\%$) than $NO_2$, $N_2O$, or $HNO_3$.