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What the exam asks for

• Convert among mass, moles, and particles with dimensional analysis; use isotope abundances to find average atomic mass.
• Infer empirical formulas from elemental mass percentages and distinguish a pure compound’s fixed ratios from mixture composition.
• Represent electron configurations for atoms and ions, and interpret photoelectron peak positions and heights.
• Use electron configuration and periodic position to explain trends, typical ionic charges, and related reactivity.

Common slips to watch for

• Reversing a conversion factor or omitting units can produce a plausible-looking but incorrect particle count.
• Using isotope mass numbers instead of isotope masses, or treating abundances as whole-number counts, skews an average.
• Reducing empirical-formula ratios too early or treating a mixture as if it had one fixed compound formula loses composition marks.
• Removing electrons from the wrong subshell in a cation, or explaining a trend using atomic size alone, weakens reasoning.

What the exam asks for

• Connect bond type and potential-energy curves to the particles’ interactions and the elements involved.
• Use particulate models of ionic solids, metals, and alloys to show arrangements and the nature of their interactions.
• Construct Lewis diagrams with appropriate valence electrons, bonds, lone pairs, and formal charges.
• Use resonance, VSEPR, bond order, polarity, and hybridization to connect electron arrangements to molecular structure.

Common slips to watch for

• Drawing isolated ion pairs rather than an extended ionic lattice misrepresents an ionic solid.
• Forgetting lone pairs or miscounting valence electrons leads to incorrect Lewis structures and geometries.
• Treating resonance structures as rapidly switching forms, rather than alternative representations of a delocalized structure, causes errors.
• Choosing a structure by minimizing formal charges without checking charge placement or electronegativity can favor the wrong diagram.

What the exam asks for

• Relate molecular structure and particle interactions to intermolecular forces, phase properties, solubility, and separation outcomes.
• Apply gas relationships and kinetic molecular theory to particle motion, energy distributions, and real-gas deviations.
• Calculate solution concentrations and gas quantities, and represent mixtures with particle diagrams that show relative amounts and interactions.
• Interpret electromagnetic radiation, photon energy, and absorbance in terms of molecular transitions and solution properties.

Common slips to watch for

• Ranking intermolecular forces without accounting for molecular shape or the specific particles present can give an incorrect comparison.
• Using gas-law quantities with inconsistent units, or assuming real gases behave ideally under all conditions, undermines calculations and explanations.
• Drawing solution particles without preserving relative concentration or distinguishing dissolved ions from intact molecules loses representation marks.
• Confusing wavelength and frequency trends, or changing concentration and path length in the same direction when applying Beer-Lambert relationships, leads to wrong conclusions.

What the exam asks for

• Distinguish evidence of chemical change from physical change and connect macroscopic observations to bond and particle interactions.
• Represent reactions with balanced molecular, complete ionic, and net ionic equations, and match equations to particulate models.
• Use reaction stoichiometry with masses, gas quantities, or solution concentrations; identify titration equivalence from reacting amounts.
• Classify acid-base, precipitation, and redox processes using proton transfer, insoluble products, and oxidation-number changes; balance redox with half-reactions.

Common slips to watch for

• Leaving spectator ions in a net ionic equation, or splitting a solid or weak electrolyte into ions, gives an incorrect representation.
• Using coefficients as mass ratios rather than mole ratios leads to incorrect stoichiometric results.
• Calling the indicator color change the equivalence point instead of using the reaction amounts confuses observation with stoichiometry.
• Changing atoms or charge inconsistently while balancing redox half-reactions fails conservation of both mass and charge.

What the exam asks for

• Infer rate laws and reaction orders from initial-rate data and concentration-versus-time patterns.
• Connect elementary-step stoichiometry to its rate law, and use a rate-limiting step or pre-equilibrium to derive an overall rate law.
• Interpret energy profiles by distinguishing activation energy from overall energy change, including for multistep mechanisms.
• Use collision theory to explain effects of concentration, temperature, orientation, and catalysts on reaction rate.

Common slips to watch for

• Assigning rate-law exponents from balanced overall-equation coefficients rather than experimental evidence or an elementary step.
• Confusing a reactant's disappearance rate with the reaction rate, or overlooking stoichiometric rate relationships.
• Reading activation energy as the energy change between reactants and products, or drawing the wrong number of peaks for a multistep path.
• Applying the pre-equilibrium approximation without relating the intermediate concentration to the initial-step equilibrium.

What the exam asks for

• Relate observed temperature changes to heat flow and whether a process is endothermic or exothermic.
• Calculate heat transfer using heat capacity, amount, temperature change, and molar enthalpy for phase changes or reactions.
• Determine reaction enthalpy from bond enthalpies or standard enthalpies of formation, showing balanced equations and units.
• Combine enthalpy steps with Hess’s law, reversing or scaling equations consistently and interpreting energy diagrams.

Common slips to watch for

• Using a temperature change with the wrong sign or assigning heat gained by the surroundings to the reacting system.
• Mixing up heat capacity and molar enthalpy, or omitting conversion between grams and moles.
• Reversing the bond-energy subtraction: energy for bonds broken minus energy for bonds formed.
• Reversing or scaling a Hess’s-law equation without making the corresponding change to its enthalpy.

What the exam asks for

• Write Q and K expressions from balanced equations, omitting pure solids and liquids and using concentration or pressure consistently.
• Use Q compared with K to determine reaction direction; distinguish that direction from the equilibrium amounts indicated by K’s size.
• Set up equilibrium-concentration calculations from initial amounts and stoichiometry, then check the result against K.
• Represent equilibrium and stress responses with particle models, and calculate solubility from Ksp, including a common ion.

Common slips to watch for

• Including pure solids or liquids in an equilibrium expression, or using exponents that do not match the balanced equation.
• Treating equal forward and reverse rates as equal concentrations, or assuming equilibrium means the reaction has stopped.
• Choosing the direction of change from K alone when Q is needed for the system’s current state.
• Forgetting salt-dissociation stoichiometry when relating molar solubility to Ksp, especially with a common ion.

What the exam asks for

• Calculate pH and pOH for strong and weak acids or bases, using equilibrium relationships for weak monoprotic species.
• Analyze acid-base mixtures, buffers, and titration curves by identifying major species and explaining changes during neutralization.
• Relate pH to pKa or pKb to identify predominant conjugate forms, and use conjugate-pair amounts to determine buffer pH and capacity.
• Use molecular and ionic structure to compare acid or base strength, and explain qualitative pH effects on salt solubility.

Common slips to watch for

• Using concentration directly as pH, or failing to account for strong-acid or strong-base stoichiometry before calculating pH.
• Applying a weak-acid equilibrium approximation when the mixture is a buffer or after strong acid-base neutralization has changed the major species.
• Reversing the ratio in the Henderson–Hasselbalch relationship or using the wrong conjugate pair in a buffer.
• Misidentifying the equivalence or half-equivalence point on a titration curve, or treating buffer capacity as independent of component amounts.

What the exam asks for

• Predict entropy-change signs and compare magnitudes across phase changes, dissolution, and changes in particle count.
• Calculate standard entropy changes from tabulated absolute entropies, using stoichiometric coefficients for every species.
• Connect ΔG°, K, and E°cell to favorability; distinguish thermodynamic favorability from reaction rate and explain how coupling can drive an unfavorable process.
• Interpret cell diagrams and half-reactions to identify electrode roles, electron flow, cell behavior, and the effects of nonstandard conditions.
• Use balanced electrochemical reactions, units, and electron stoichiometry to relate cell potential, free energy, and charge passed.

Common slips to watch for

• Omitting stoichiometric coefficients or reversing product and reactant terms when calculating ΔS°.
• Treating a favorable ΔG° as evidence that a reaction is fast, or assuming a catalyst changes the reaction’s thermodynamic favorability.
• Confusing anode and cathode roles or electron-flow direction, especially when switching between galvanic and electrolytic cells.
• Using concentrations in place of activities without accounting for reaction stoichiometry, or converting moles of electrons to charge incorrectly.