CHEM126 Experiment Review

Experiment 1: Excel & Boyle’s Law

Goal: Learn to use Excel for processing experimental data and investigate Boyle's Law by analyzing how the pressure and volume of a gas relate.

Concept Explanation: Boyle’s Law describes the relationship between pressure (P) and volume (V) of a fixed amount of gas at a constant temperature. According to this law, pressure and volume are inversely proportional. This means:

• If pressure increases, the gas particles are compressed into a smaller volume.

• If pressure decreases, gas particles have more space to move, so volume increases.

Why It Works: As gas particles move, they collide with the container walls. The more frequent these collisions, the greater the pressure. If volume is reduced while the number of particles and temperature are kept constant, collisions become more frequent → pressure increases. This constant relationship is expressed by:

• P * V = k (k is a constant for fixed n and T)

Graphing V vs 1/P: You graph volume on the y-axis and 1/P on the x-axis. The inverse relationship means this graph should be linear if Boyle’s Law holds. A straight line confirms that V is proportional to 1/P, validating the theory with experimental data.

Mathematical Form:

• P * V = k

• V = k / P

Skills Developed:

• Data handling and graphing in Excel

• Experimental verification of theoretical relationships

• Understanding pressure-volume behavior of gases

Experiment 2: General Lab Techniques

Goal: Learn proper usage of essential lab equipment for accurate quantitative measurements.

Concept Explanation: Reliable data in chemistry comes from precise and accurate measurements. Accuracy refers to how close a measurement is to the true value, while precision reflects consistency across repeated measurements.

Why It Works: Each piece of lab glassware is designed for a specific accuracy. For example:

• Volumetric flasks deliver fixed, extremely accurate volumes.

• Graduated cylinders allow for quick estimates but less precise.

• Analytical balances ensure tiny mass differences are detected.

Core Techniques Include:

• Rinsing pipettes with solution to prevent dilution from residual water

• Reading the bottom of the meniscus to ensure volume accuracy

• Titration endpoints indicated by a clear color change with an indicator (e.g., phenolphthalein)

Skills Developed:

• Consistency in measurement

• Recognizing and minimizing error sources

• Foundation for quantitative lab analysis

Experiment 3: Rate Law of the Crystal Violet Reaction

Goal: Determine the rate law and reaction orders for the fading of crystal violet (CV⁺) in basic solution.

Concept Explanation: Reaction rate depends on how frequently reacting molecules collide with sufficient energy. The rate law shows how concentration affects rate:

• Rate = k * [CV⁺]^n * [OH⁻]^m

Since [OH⁻] is in excess, its concentration doesn’t change much → pseudo rate law simplifies analysis:

• Rate = k' * [CV⁺]^n

Why Graphing Absorbance Works: CV⁺ is purple, so its concentration can be tracked by its absorbance. The Beer-Lambert Law:

• A = ε * l * c → shows that absorbance is directly proportional to concentration.

Thus, tracking A vs time is like tracking [CV⁺] vs time. You can determine reaction order (n) by testing different graphs:

• If A vs t is linear → zero order (rate independent of [CV⁺])

• If ln(A) vs t is linear → first order (rate proportional to [CV⁺])

• If 1/A vs t is linear → second order

Determining m: Using two different OH⁻ concentrations gives different pseudo rate constants:

• m = log(k₂'/k₁') / log([OH⁻]₂ / [OH⁻]₁)

Skills Developed:

• How to extract rate law from data

• Use of absorbance to track reaction progress

• Application of logarithmic and exponential relationships to chemical systems

Experiment 4: Iodine Clock Reaction

Goal: Measure the rate of the oxidation of iodide by hydrogen peroxide and determine activation energy using reaction timing and temperature variation.

Concept Explanation: In a clock reaction, a visible change (like color) occurs after a fixed amount of product has formed. This provides a sharp, measurable indication of reaction progress. Here, iodide (I⁻) is oxidized by hydrogen peroxide (H₂O₂) to form iodine (I₂), which then reacts with starch to form a dark blue complex.

To delay the color change, thiosulfate (S₂O₃²⁻) is added to consume I₂ as it forms. The time taken for the blue color to appear corresponds to the time needed for a fixed amount of I₂ to accumulate after all thiosulfate is used up.

Why It Works: Since you know the stoichiometry and volume, you can calculate how much reactant has been converted in that time, allowing rate calculations.

Rate Law:

• Rate = k * [H₂O₂]^a * [I⁻]^b * [H⁺]^c

• With [I⁻] and [H⁺] in excess, the simplified form is:

  • Rate = k' * [H₂O₂]

Using Temperature to Determine Activation Energy: By changing temperature and measuring the new rate constants (k'), you use the Arrhenius equation:

• ln(k') = -Ea / (R * T) + ln(A)

• Plotting ln(k') vs 1/T yields a straight line whose slope is -Ea/R

Skills Developed:

• Use of reaction timing to analyze rates

• Understanding of temperature’s effect on rate

• Use of integrated rate laws and Arrhenius equation

Experiment 5: Beer-Lambert Law & Equilibrium Constant

Goal: Use spectrophotometry to determine the concentration of FeSCN²⁺ and calculate the equilibrium constant (K) for the formation reaction.

Concept Explanation: When Fe³⁺ and SCN⁻ are mixed, they form the red-colored FeSCN²⁺ complex. The concentration of this product can be tracked via its absorbance using the Beer-Lambert Law:

• A = ε * l * c Where A is absorbance, ε is the molar absorptivity, l is path length (usually 1 cm), and c is concentration.

Why It Works: FeSCN²⁺ absorbs strongly at 470 nm, so its concentration can be directly inferred from its absorbance.

Calibration Curve: Prepare standard solutions with known FeSCN²⁺ concentrations and measure absorbance. Plot A vs c → linear relationship. The slope gives ε.

Equilibrium Constant Calculation: In mixtures of Fe³⁺ and SCN⁻, some remains unreacted. Using ICE tables and Beer-Lambert-based [FeSCN²⁺], calculate:

• K = [FeSCN²⁺] / ([Fe³⁺][SCN⁻])

Skills Developed:

• Construction of calibration curves

• Application of equilibrium concepts

• Spectrophotometric data interpretation

Experiment 6: Solubility Product of KHT

Goal: Determine the solubility product constant (Ksp) of potassium hydrogen tartrate (KHT) by titrating the hydrogen tartrate ion in a saturated solution.

Concept Explanation: Sparingly soluble salts like KHT do not dissolve completely in water. When they reach equilibrium between their solid and dissolved forms, the solution is said to be saturated. The extent to which the salt dissolves is governed by the solubility product constant (Ksp).

Why It Works: For KHT:

• KHT(s) ⇌ K⁺(aq) + HT⁻(aq)

Because this dissociation is 1:1, the concentration of K⁺ is equal to that of HT⁻ at equilibrium. If we can determine one of these concentrations, we know the other, and thus:

• Ksp = [K⁺][HT⁻] = s²

Titration Connection: HT⁻ is a weak acid anion. We can determine its concentration by titrating it with standardized NaOH:

• HT⁻ + OH⁻ → T²⁻ + H₂O

The amount of NaOH added corresponds to the number of moles of HT⁻ in the sample. Using the volume of sample titrated, we determine [HT⁻] and therefore [K⁺].

Mathematical Relationships:

• moles HT⁻ = volume NaOH (L) * molarity NaOH

• [HT⁻] = moles / volume sample (L)

• Ksp = [HT⁻]²

Skills Developed:

• Connecting titration to solubility

• Calculating equilibrium constants for slightly soluble salts

• Using stoichiometry in equilibrium contexts

Experiment 7: Titration and pH Curves

Goal: Use a pH meter to study titration curves of strong acid/strong base and weak acid/strong base systems. Determine pKa and Ka of the weak acid.

Concept Explanation: Titration curves graph the pH of a solution as titrant is added. The shape of the curve reveals the acid-base strength and stoichiometry. The point at which moles of acid equal moles of base is the equivalence point.

Why It Works: In a weak acid/strong base titration, the pH rises gradually at first due to buffering. Near equivalence, pH rises sharply. Halfway to the equivalence point, [HA] = [A⁻], and thus:

• pH = pKa

This allows for experimental determination of pKa, which leads to:

• Ka = 10^(-pKa)

Graphical Method:

• Use pH vs. volume added

• Find half-equivalence point by identifying where volume added is half that of the equivalence point

• Read pH → pKa

Skills Developed:

• Interpreting titration curves

• Determining acid dissociation constants

• Understanding buffering regions

Experiment 8: Buffers

Goal: Demonstrate how buffered solutions resist pH changes by comparing their response to added acid/base with that of pure water.

Concept Explanation: A buffer solution contains a weak acid and its conjugate base. This equilibrium system resists changes in pH by reacting with added H⁺ or OH⁻ ions.

Why It Works: Buffer equation:

• HA ⇌ H⁺ + A⁻

When H⁺ is added, A⁻ neutralizes it. When OH⁻ is added, HA donates H⁺ to neutralize the base. The balance between [HA] and [A⁻] keeps pH relatively constant.

Mathematical Support (Henderson-Hasselbalch):

• pH = pKa + log([A⁻]/[HA])

Adding strong acid increases [HA] and decreases [A⁻] slightly → small pH change. Water, lacking this buffer capacity, undergoes drastic pH changes.

Skills Developed:

• Visualizing buffer action

• Quantitative pH prediction using HH equation

• Experimental comparison of buffered vs unbuffered systems

Experiment 9: Determining ΔH⁰ and ΔS⁰ of FeSCN²⁺ Formation

Goal: Use temperature-dependent equilibrium data to calculate the enthalpy (ΔH⁰) and entropy (ΔS⁰) changes of the Fe³⁺ + SCN⁻ ⇌ FeSCN²⁺ equilibrium.

Concept Explanation: The position of equilibrium shifts with temperature based on thermodynamic properties. By measuring the equilibrium constant K at various temperatures and applying the van’t Hoff equation:

• ln(K) = -ΔH⁰ / (R * T) + ΔS⁰ / R

Why It Works: The equation is derived from combining Gibbs free energy (ΔG = ΔH - TΔS) with the equilibrium condition ΔG = -RT ln(K).

• A plot of ln(K) vs 1/T gives a straight line.

• Slope = -ΔH⁰/R

• Intercept = ΔS⁰/R

Experimental Strategy:

• Use colorimeter to measure [FeSCN²⁺] at various temperatures

• Calculate [Fe³⁺] and [SCN⁻] from initial and reacted values

• Determine K and build van’t Hoff plot

Skills Developed:

• Applying thermodynamic theory to real systems

• Using spectrophotometry for equilibrium analysis

• Constructing and interpreting linearized equations

Experiment 10: Galvanic Cells

Goal: Construct and analyze galvanic (voltaic) cells to measure E°cell and understand redox reactions.

Concept Explanation: Galvanic cells convert chemical energy into electrical energy through redox reactions. They consist of two half-cells linked by a salt bridge, allowing ion flow, and an external wire, allowing electron flow.

Why It Works: The difference in potential energy between the two electrodes drives electron movement. The potential difference is calculated from standard reduction potentials:

• E°cell = E°cathode - E°anode

Experimental Approach:

• Use different metals/electrolytes (e.g., Zn/Cu)

• Measure voltage with voltmeter

• Compare to theoretical E°cell using standard reduction tables

Optional Extension – Nernst Equation: If concentrations deviate from standard conditions:

• E = E° - (0.0591 / n) * log(Q)

Skills Developed:

• Understanding electron transfer

• Measuring cell potential

• Connecting chemical reactivity to electrical output

Experiment 11: Electrolysis of Water

Goal: Demonstrate the use of electricity to drive a non-spontaneous redox reaction by electrolyzing water.

Concept Explanation: Unlike galvanic cells, electrolytic cells use electrical energy to drive chemical change. Water does not spontaneously split into H₂ and O₂ but can be forced to do so with an external power supply.

Half-Reactions:

• Cathode (reduction): 2 H⁺ + 2 e⁻ → H₂(g)

• Anode (oxidation): 2 H₂O → O₂(g) + 4 H⁺ + 4 e⁻

Why It Works: Applying voltage provides the energy needed to overcome the positive ΔG of the reaction. Acidified water ensures conductivity via free H⁺ and SO₄²⁻ ions.

Observations:

• Gases collected at each electrode

• H₂ volume is double that of O₂ (2:1 molar ratio)

Skills Developed:

• Understanding forced redox reactions

• Observing stoichiometry in gas evolution

• Applying Faraday’s laws qualitatively