General Chemistry (CHM-113/115)

Chemistry explains the world.

CHM-113/115 provides critical knowledge for science careers.

Quantitative study of reactants and products.

Accuracy matters. Check math setup carefully.

State where forward/reverse reaction rates are equal.

Dynamic State

Reactions continue, but net concentrations constant.

Equilibrium Constant (K)

Ratio of product concentrations to reactant concentrations (raised to stoichiometric powers) at equilibrium. $K_c$ (molarity), $K_p$ (partial pressures).

• $K > 1$: Products favored.
• $K < 1$: Reactants favored.
• $K \approx 1$: Significant amounts of both.
• Only T changes K. Pure solids/liquids excluded.

Reaction Quotient (Q)

Same form as K, but uses non-equilibrium concentrations. Compare Q to K to predict shift direction.

• $Q < K$: Shift right (products).
• $Q > K$: Shift left (reactants).
• $Q = K$: At equilibrium.

ICE Tables

Method (Initial, Change, Equilibrium) to calculate equilibrium concentrations given K and initial conditions.

Predicts equilibrium shifts due to stress.

1. Concentration Changes

Add reactant -> shift right. Add product -> shift left. Remove reactant -> shift left. Remove product -> shift right.

2. Pressure/Volume Changes (Gases)

Increase P (decrease V) -> shift to side with fewer gas moles. Decrease P (increase V) -> shift to side with more gas moles. No shift if gas moles equal.

3. Temperature Changes

Treat heat as reactant (endothermic, $\Delta H > 0$) or product (exothermic, $\Delta H < 0$). Add heat -> shift away from heat. Remove heat -> shift toward heat. Only T changes K value.

4. Catalysts

Speed up both forward/reverse rates equally. Reach equilibrium faster, but no shift in position, no change in K.

Understanding shifts is key for reaction control. Problems complex? Consider science assignment help.

Study of energy transformations in reactions.

1. First Law

Energy conservation. Internal energy change $\Delta E = q + w$ (heat + work).

2. Enthalpy ($\Delta H$)

Heat change at constant pressure. Negative $\Delta H$ = exothermic (releases heat). Positive $\Delta H$ = endothermic (absorbs heat). Hess’s Law calculates $\Delta H$ for overall reaction from steps.

3. Second Law & Entropy ($\Delta S$)

Entropy measures disorder/randomness. Universe entropy increases for spontaneous processes. Positive $\Delta S$ = increased disorder.

4. Gibbs Free Energy ($\Delta G$)

Determines spontaneity at constant T, P. $\Delta G = \Delta H – T\Delta S$.

• $\Delta G < 0$: Spontaneous (product-favored).
• $\Delta G > 0$: Non-spontaneous (reactant-favored).
• $\Delta G = 0$: Equilibrium.
• Relates to equilibrium constant: $\Delta G^\circ = -RT \ln K$.

State Functions

Properties depending only on current state, not path (E, H, S, G). Heat (q) and work (w) are path-dependent.

Standard States

Defined conditions (usually 1 atm, 298 K, 1 M conc) for comparing thermodynamic data ($\Delta H^\circ, \Delta S^\circ, \Delta G^\circ$).

Calculating $\Delta H^\circ_{rxn}$

Using standard enthalpies of formation ($\Delta H^\circ_f$): $\Delta H^\circ_{rxn} = \sum n \Delta H^\circ_f (\text{products}) – \sum m \Delta H^\circ_f (\text{reactants})$. Similar calculations for $\Delta S^\circ$ and $\Delta G^\circ$.

Careful with signs, units (kJ vs J). Need review? Check Thermochemistry Resources.

Key equilibrium topic.

• Definitions: Arrhenius, Brønsted-Lowry (proton donor/acceptor), Lewis (electron pair acceptor/donor).
• Strong vs Weak: Extent of dissociation. Strong acids/bases dissociate completely. Weak establish equilibrium ($K_a, K_b$).
• pH Scale: Logarithmic measure of $[H^+]$. $pH = -\log[H^+]$. $pOH = -\log[OH^-]$. $pH + pOH = 14$ at 25°C.
• Buffers: Resist pH change. Weak acid/conjugate base or weak base/conjugate acid pair. Henderson-Hasselbalch equation.
• Titrations: Determine unknown concentration using known solution. Equivalence point.

Calculations involve equilibrium principles.

Practical application of concepts.

• Safety: Goggles, procedures, waste disposal. Critical.
• Techniques: Titration, spectroscopy, calorimetry, chromatography basics.
• Measurement: Accurate use of balances, glassware (pipettes, burettes). Significant figures.
• Observation: Recording qualitative/quantitative data precisely.
• Data Analysis: Calculations, graphing, error analysis (statistical analysis).
• Lab Reports: Clear communication of procedure, results, conclusions.

Labs reinforce theory, build essential skills.

Avoid common errors:

• Unit Conversion Errors (esp. kJ vs J)
• Incorrect Significant Figures
• Stoichiometry: Not balancing equation, wrong mole ratios
• Equilibrium: Forgetting coefficients as powers in K, wrong shift prediction
• Thermodynamics: Sign errors ($\Delta H, \Delta G$), confusing state/path functions
• Acids/Bases: Confusing strong/weak, pH/pOH calculation errors
• Lab: Safety violations, measurement errors, poor data recording

Attention to detail, practice (Cognitive Load Reduction).