Comprehensive CHEM II Laboratory Review

Thin Layer Chromatography (TLC)

PURPOSE
• Separate tiny quantities of compounds; forensic use for dye ID.
• Calculate Rf (distance spot ÷ distance solvent front). • Relate Rf to polarity & IM-forces.
HARDWARE / PHASES
• Stationary phase: \text{SiO}2 or Al a on glass/plastic (≈0.25 mm). • Mobile phase: chosen solvent drawn upward by capillarity. • Greater attraction for solvent → larger Rf; greater attraction for plate → smaller R_f.
PROCEDURE
• Draw start line 1 cm from bottom, 5 spots (crime sample + 4 suspects).
• Spot ≤ 2 mm; darken by re-spotting.
• Place TLC into chamber with solvent below start line; saturate vapor with filter paper.
• Develop until solvent ≈1 cm from top; mark solvent front immediately.
• Visualize (visible, UV, or stain); measure all distances; compute R_f.
PRE-LAB CONCEPTS
• Tight, small spot needed; size influences resolution.
• Starting line & solvent level critical.
• Do not let solvent run off plate before marking front.

OBJECTIVES
• Locate triple, boiling & freezing points; sketch P–T diagram.
• Reinforce concepts of vapor pressure, phase coexistence, Tb, Tf.
TRIPLE-POINT APPARATUS
• Test tube w/ t-BuOH + boiling chips, thermometer through #4 stopper, vacuum line via water aspirator.
• Under ~42.4\,\text{mmHg} observe simultaneous boiling & freezing → record T_{tp}.
BOILING POINT
• Heat 600 mL water bath; immerse tube; read constant vapor-temperature (≈ normal T_b at lab P).
FREEZING POINT
• Cold-water (≈18 °\mathrm C) bath; stir continually; record plateau temperature while solid forms.
DATA ANALYSIS
• Plot \ln P vs 1/T or simple sketch marking solid, liquid, gas fields; label all three key T’s.
KEY EQUATIONS
• Clausius–Clapeyron \displaystyle \ln P = -\frac{\Delta H_{vap}}{R}\,\frac1T + C

THEORY
• Colligative property: \Delta Tb = Kb m (independent of solute identity).
• m = molality = \dfrac{n_{sol}}{\text{kg solvent}}.
PROCEDURE
1. Find pure ethanol T_b in 90 °C water bath.
2. Prepare 1 m, 2 m, 3 m ethylene-glycol / ethanol solutions (density 0.789\,\text{g mL}^{-1}).
3. Record elevated Tb; plot \Delta Tb vs m → slope =K_b^{\,(EtOH)}.
4. Dissolve unknown sample with mass equal to 1 m EG mass; measure \Delta Tb; compute molality then moles; M = m{unk}/n_{unk}.
SAFETY: ethanol & ether are flammable; EG toxic.

RATE LAW
\text{Rate}=k[\text{dye}]^{a}[\text{OCl}^-]^{b}
• Pseudo-order: large excess NaOCl ⇒ k' = k[\text{OCl}^-]^b.
EXPERIMENT
• Microlab spectro-kinetics at 525 nm; follow fade of FD&C Red #3 or Blue #1.
• Plot \ln[\text{dye}] vs t (1st-order) and 1/[\text{dye}] vs t (2nd-order).
• Best linear fit → order in dye; slope =-k'.
• Repeat with new [OCl⁻]; compare k' values ⇒ determine b & true k.

REACTION
\ce{Fe^{3+} + SCN^-
EQUILIBRIUM CONSTANT
• Measure absorbance at 468 nm; \varepsilon=7260\,\text{M}^{-1}\text{cm}^{-1}.
• [\text{FeSCN}^{2+}] = A/(\varepsilon b), use RICE to back-calculate reactant eq. conc.; compute K_c (triplicate trials, 𝑇 ≈ rt).
LE CHÂTELIER TESTS
• Add various salts/temperature changes; note colour shift: darker → right (products), paler → left (reactants).

STOICHIOMETRY
\ce{MnO4^- + 5Fe^{2+} + 8H^+ → Mn^{2+} + 5Fe^{3+} + 4H2O}
• Endpoint: first persistent faint pink \text{MnO}_4^- excess.
ANALYSIS STEPS
1. Dissolve crushed tablet in \ce{H2SO4}.
2. Titrate 10 mL aliquots w/ standard 0.0020\,\text M KMnO_4 (triplicate).
3. Moles \ce{Fe^{2+}} = \dfrac{1}{5} n_{\ce{MnO4^-}}.
4. Convert to mg Fe; %Fe = (mg Fe / tablet mass)×100.

EQUILIBRIUM
\ce{[Co(H2O)6]^{2+} + 4Cl^-
DATA COLLECTION
• Record absorbance at \lambda{max} (≈ 650 nm for blue species) at 10° C steps (10–40 °C). • K = \dfrac{[\text{CoCl}4^{2-}]}{[\text{Co(H}2\text O)6^{2+}][Cl^-]^4} (Cl⁻ large excess ⇒ constant).
• Plot \ln K vs 1/T : slope =-\Delta H^\circ/R; intercept = \Delta S^\circ/R.
• \Delta G^\circ = -RT\ln K at each T.

WEAK-ACID Ka Ka = \frac{[H_3O^+][A^-]}{[HA]} ;
% ionization increases on dilution.
STRONG ACID–STRONG BASE CURVE
• S-shaped; equivalence pH=7.
• Half-equivalence of weak acid: pH = pK_a.
BUFFER
\text{pH}=pKa+\log\frac{[\text{base}]}{[\text{acid}]} (Henderson-Hasselbalch). • Effective range pKa ±1.

Diprotic example: \ce{H2A}
• Two equivalence points; volumes ratio 1 : 1.
• K{a1} from pH at 1st half-eq; K{a2} from 2nd.

EQUILIBRIUM
\ce{Ca(OH)2(s)
TITRIMETRIC DETERMINATION
1. Filter saturated Ca(OH)₂ (discard first 5 mL).
2. Titrate 10.00 mL aliquots with 0.0500\,\text M HCl using phenolphthalein.
3. n{\text{OH}^-}=n{HCl}; [OH^-]=n/V; [Ca^{2+}] = 0.5[OH^-].
4. Calculate K_{sp} and molar solubility s=[Ca^{2+}]; average over 3 trials.
COMMON-ION TEST
• Add drops of \ce{NaOH}, \ce{CaCl2}, \ce{HCl}, etc. to aliquots; monitor pH shifts:
– Add OH⁻ or Ca²⁺ → ↓solubility (shift left).
– Add H⁺ or precipitating \ce{Ag^+} (removes OH⁻) → ↑solubility (shift right).

STANDARD CELL POTENTIAL
E^\circ{cell}=E^\circ{\text{cath}}-E^\circ_{\text{anode}}
• Positive ⇒ spontaneous (voltaic).
LINE NOTATION
\text{Zn}|\text{Zn}^{2+}||\text{Cu}^{2+}|\text{Cu}
ELECTROLYTIC CALCULATIONS (Faraday’s Law)
• Charge Q=It, 1\,\text F = 96485\,\text C\,\text{mol}^{-1}\,e^-.
• Mass plated m = \dfrac{ItM}{nF}.
ELECTROPLATING EXAMPLE
• Ni²⁺ + 2e⁻ → Ni(s) on Cu; at 5 V, 600 s, measured 1.02\times10^{-3}\,\text{mol }e^- ⇒ Q ≈ 98 C, I ≈ 0.16 A, theoretical Ni mass \approx 0.030 g (close to observed 0.029 g).

Discard initial filtrate when filtering saturated solutions (filter paper adsorption).
Deviations: DEVi = |xi - \bar x|; AD = \frac{\sum DEV}{n}.
Maintains CO₂-free basic solutions; quick titration reduces error.
Safety: acids corrosive, bases caustic, KMnO₄ strong oxidizer, Cr⁶⁺ toxic, heavy-metal waste collected.