Chem Lab Practical

Reaction Rate & Thiosulfate Coupling

Rate = -(½) · (Δ[S₂O₃²⁻] / Reaction time(Δt))

General Rate Law

Rate = k[Fe³⁺]ᵐ[I⁻]ⁿ

Finding Orders (m and n) via Log Plots

ln(Rate) = m · ln([Fe³⁺]) + ln(constant)

Slopes typically round to whole numbers (such as 1 or 2)

Graph 1 (ln(Rate) vs. ln([Fe³⁺])): The slope equals the reaction order m with respect to Fe³⁺

Graph 2 (ln(Rate) vs. ln([I⁻])): The slope equals the reaction order n with respect to I⁻.

Pseudo-First Order Kinetics

Because [H₂O] and [OH⁻] are in huge excess compared to [dye], their concentrations remain effectively constant, making the reaction behave as first-order with respect to the dye

Rate = k_obs[dye], where k_obs = k₁[H₂O] + k₂[OH⁻]

Van 't Hoff Equation

ln(Ksp) = (-ΔH° / R)(1 / T) + (ΔS° / R), where T is in Kelvin and R = 8.314 J/(K·mol)

On a graph of ln(Ksp) vs. 1/T, Slope = -ΔH° / R., ΔH° = -Slope × R.

Freezing Point Depression Formula

ΔTf = i · Kf · m

ΔTf = Tf(solvent) - Tf(solution) (always expressed as a positive temperature drop).

Step-by-Step Molar Mass Calculation

Find ΔTf from cooling curves.

Calculate molality: m = ΔTf / (i · Kf).

Calculate moles of solute: Moles = m × kg of solvent.

Calculate molar mass: Molar Mass = Mass of solute (g) / Moles of solute.

Semiconductor (TiO₂)

Coated on the negative electrode, accepts excited electrons from the dye and conducts them into the external circuit.

Natural Dye (Anthocyanin from raspberry juice)

Absorbs sunlight photons to excite electrons and injects them into TiO₂.

Counter Electrode (Carbon/Soot)

Positive electrode that completes the circuit.

Electrolyte redox mediator in ethylene glycol)

Iodide (I⁻) donates electrons to regenerate the oxidized dye, creating a continuous regenerative cycle. I⁻(reduced)/I2⁻(oxidized)

Power (Watts)

Voltage (V) × Current (A)

Efficiency (%)

(Electrical Power Output / Incoming Solar Power) × 100%

Empirical Formula Calculation

Convert mass percent (or experimental mass) of Cu²⁺, SO₄²⁻, NH₃, and H₂O (found by subtracting the other mass percents from 100%) into moles, divide by the smallest mole value, and round to whole numbers.

Molarity of Acetic Acid

M_acid · V_acid = M_base · V_base

Percent (w/v) Acetic Acid

[Mass of CH₃COOH (g) / Volume of vinegar (mL)] × 100% (using molar mass 60.05 g/mol)

Ka from Halfway Point)

If the equivalence point is at 18.50 mL, the halfway point is 9.25 mL. If pH at the halfway point is 4.10, then:, pKa = 4.10, Ka = 10⁻⁴·¹⁰ = 7.9 × 10⁻⁵

Determining Molarity from Mass Data

1. Mass dissolved = Initial solid mass - Remaining solid mass after dissolution.

2. Moles dissolved = Mass dissolved / 261.34 g/mol.

3. Molarity (M) = Moles dissolved / 0.0500 L (for a 50.0 mL sample).

Spectrophotometry & Beer-Lambert Law

Measured at maximum absorbance wavelength λmax = 470 nm (using 0.5 M HNO₃ as the reference blank).

Graphical Determination of Keq

Linear regression yields a straight line where Slope = -Keq (so Keq = -Slope).

Experiment 1: Chemical Kinetics

1. Mix KI, Na₂S₂O₃, starch indicator, and water, then add Fe(NO₃)₃ and start timing

2. Stop the timer when the solution turns dark blue-black to calculate the reaction rate

3. Repeat while varying [Fe³⁺] and [I⁻] separately, then use ln(Rate) vs. ln[reactant] plots to find the reaction orders

Experiment 2: The Rate of Decolorization

1. Mix phosphate buffer solutions to set [OH⁻] and add methyl green dye

2. Measure absorbance over time in a spectrophotometer as the green color fades

3. Plot −ln(A) vs. time to get k_obs, then plot k_obs vs. [OH⁻] across different buffers to get k₂

Experiment 3: Borax Solubility

1. Saturate water with excess solid borax at several different temperatures

2. Pipet out a sample of the clear liquid (no solid) and titrate with HCl using bromocresol green to find molar solubility and Ksp

3. Plot ln(Ksp) vs. 1/T to calculate ΔH° from the slope

Experiment 4: Freezing Point Depression

1. Measure the cooling curve of pure water in an ice-salt bath to find its normal freezing point

2. Dissolve a weighed solute in water and measure its new, lower freezing point

3. Use ΔTf to calculate molality, moles, and the solute's molar mass

Experiment 5: Preparation of a Solar Cell

1. Coat a conductive glass slide with TiO₂ paste, bake it, and soak it in raspberry juice (anthocyanin dye)

2. Coat a second conductive slide with carbon soot from a flame

3. Clamp both slides together with iodine electrolyte in between and measure voltage and current under light

Experiment 6: Synthesis & Analysis of a Copper Coordination Compound

1. Synthesis: React copper(II) sulfate with ammonia and ethanol, then vacuum-filter and dry the crystals

2. Analysis: Find mass percentages of sulfate (precipitate with BaCl₂), ammonia (titrate with HCl), and copper (absorbance via Beer's Law) to determine the empirical formula

Experiment 7: Analysis of Acetic Acid in Vinegar

1. Dilute a measured sample of vinegar with water and add phenolphthalein indicator

2. Titrate with standardized NaOH until a faint, permanent pink endpoint appears

3. Use M₁V₁ = M₂V₂ to find vinegar's molarity and convert to percent (w/v)

Experiment 8: pH Titration of Vinegar

1. Place a calibrated pH probe in a measured sample of vinegar and record the initial pH

2. Add NaOH in small increments while recording pH to map the full titration curve

3. Use the steepest jump (equivalence point) to find molarity and the halfway-point pH to find pKa and Ka

Experiment 9: Solubility & Ksp of Barium Nitrate

1. Stir excess solid Ba(NO₃)₂ in distilled water until saturated, then decant off the liquid

2. Evaporate the remaining wet solid to dryness and weigh it to find how much dissolved, then calculate Ksp

3. Repeat using HNO₃ instead of water to observe lower solubility from the common ion effect

Experiment 10: Complex Ion Equilibrium of Fe(III) and Thiocyanate Ions

1. Prepare a dilute KSCN solution in HNO₃ and zero the spectrometer using HNO₃ as a blank

2. Add small, equal increments of Fe(NO₃)₃ to the beaker, measuring absorbance after each addition and returning the sample

3. Plot the linearized Beer's Law data and take the negative slope to find Keq