CHEM 1112 Final Exam Lab Concepts and Calculations Study Guide

Overview of Key Laboratory Topics for CHEM 1112 Final Exam

This study guide summarizes the main concepts and calculations from the eight laboratory experiments covered in CHEM 1112. Each section provides essential definitions, principles, and example calculations relevant to general chemistry topics, with a focus on laboratory applications.

Experiment 1: Rate of Chemical Reactions

Understanding Reaction Rates

• Reaction Rate: The change in concentration of a reactant or product per unit time.

• Rate Law: An equation that relates the reaction rate to the concentrations of reactants, typically in the form $\text{Rate} = k [A]^m [B]^n$.

• Order of Reaction: The exponents m and n in the rate law, determined experimentally.

• Factors Affecting Rate: Concentration, temperature, presence of a catalyst, and surface area.

Example: For a reaction where doubling [A] doubles the rate, the reaction is first order in A.

Additional info: The rate constant k increases with temperature, as described by the Arrhenius equation: $k = A e^{-E_a/(RT)}$.

Experiment 2: Properties of Systems in Equilibrium – Le Chatelier’s Principle

Le Chatelier’s Principle and Equilibrium Shifts

• Chemical Equilibrium: The state where the rates of the forward and reverse reactions are equal.

• Le Chatelier’s Principle: If a system at equilibrium is disturbed, it will shift to counteract the disturbance.

• Disturbances: Changes in concentration, temperature, or pressure.

Example: Adding more reactant shifts equilibrium toward products.

Additional info: For gaseous equilibria, increasing pressure favors the side with fewer moles of gas.

Experiment 3: Determination of Equilibrium Constant

Calculating the Equilibrium Constant (K)

• Equilibrium Constant (K): A ratio of product concentrations to reactant concentrations at equilibrium, each raised to the power of their coefficients.

• General Expression: $K = \frac{[C]^c [D]^d}{[A]^a [B]^b}$ for the reaction $aA + bB \rightleftharpoons cC + dD$.

• Experimental Determination: Measure concentrations at equilibrium and substitute into the K expression.

Example: If $[A]=0.2$ M, $[B]=0.3$ M, $[C]=0.4$ M, $[D]=0.5$ M, then $K = \frac{(0.4)(0.5)}{(0.2)(0.3)} = 3.33$.

Experiment 4: pH Measurements: Buffers and Their Properties

Buffer Solutions and pH Calculations

• Buffer: A solution that resists changes in pH upon addition of small amounts of acid or base, typically made from a weak acid and its conjugate base.

• Henderson-Hasselbalch Equation: $\mathrm{pH} = \mathrm{p}K_a + \log \left( \frac{[A^-]}{[HA]} \right)$

• Buffer Capacity: The amount of acid or base a buffer can neutralize before pH changes significantly.

Example: For a buffer with $[A^-]=0.1$ M, $[HA]=0.1$ M, and $pK_a=4.75$, $\mathrm{pH}=4.75$.

Experiment 9: Acid-Base Titration

Titration Curves and Calculations

• Titration: Gradual addition of a solution of known concentration (titrant) to a solution of unknown concentration until the reaction is complete.

• Equivalence Point: The point at which stoichiometrically equivalent amounts of acid and base have reacted.

• Indicator: A substance that changes color at (or near) the equivalence point.

• Calculation: $M_1V_1 = M_2V_2$ for monoprotic acid-base titrations.

Example: 25.0 mL of 0.100 M HCl requires 25.0 mL of 0.100 M NaOH to reach equivalence.

Handout: Determination of Solubility Product of Ca(OH)2

Solubility Product Constant (Ksp)

• Ksp: The equilibrium constant for the dissolution of a sparingly soluble salt.

• For Ca(OH)2: $\mathrm{Ca(OH)_2 (s) \rightleftharpoons Ca^{2+} (aq) + 2OH^- (aq)}$

• Ksp Expression: $K_{sp} = [Ca^{2+}][OH^-]^2$

• Experimental Determination: Measure [Ca2+] and [OH-] in a saturated solution and substitute into the Ksp expression.

Example: If $[Ca^{2+}] = 0.010$ M and $[OH^-] = 0.020$ M, then $K_{sp} = (0.010) \times (0.020)^2 = 4.0 \times 10^{-6}$.

Experiment 10: Enthalpy (Calorimetry)

Measuring Enthalpy Changes

• Calorimetry: The measurement of heat flow in a chemical reaction.

• Enthalpy Change (ΔH): The heat absorbed or released at constant pressure.

• Calculation: $q = m c \Delta T$, where q = heat, m = mass, c = specific heat, $\Delta T$ = temperature change.

• For reactions: $\Delta H = -q_{\text{solution}}$ (assuming no heat loss to surroundings).

Example: Mixing 50.0 g water ($c = 4.18$ J/g·°C) with a temperature increase of 2.0°C: $q = 50.0 \times 4.18 \times 2.0 = 418$ J.

Experiment 12: Voltaic Cells (Electrochemistry)

Electrochemical Cells and Cell Potentials

• Voltaic (Galvanic) Cell: A device that converts chemical energy into electrical energy via a spontaneous redox reaction.

• Cell Notation: Anode | Anode solution || Cathode solution | Cathode

• Cell Potential (Ecell): $E_{cell} = E_{cathode} - E_{anode}$

• Nernst Equation: $E = E^\circ - \frac{0.0592}{n} \log Q$ (at 25°C), where n = number of electrons transferred, Q = reaction quotient.

Example: For a Zn/Cu cell: $E_{cell} = E^\circ_{Cu^{2+}/Cu} - E^\circ_{Zn^{2+}/Zn}$

Summary Table: Key Laboratory Concepts

Experiment	Main Concept	Key Equation
1. Rate of Reaction	Reaction rate, rate law	$\text{Rate} = k [A]^m [B]^n$
2. Equilibrium/Le Chatelier	Equilibrium shifts	Le Chatelier’s Principle
3. Equilibrium Constant	Calculating K	$K = \frac{[C]^c [D]^d}{[A]^a [B]^b}$
4. Buffers	Buffer pH	$\mathrm{pH} = \mathrm{p}K_a + \log \left( \frac{[A^-]}{[HA]} \right)$
9. Titration	Equivalence point	$M_1V_1 = M_2V_2$
Handout: Ksp	Solubility product	$K_{sp} = [Ca^{2+}][OH^-]^2$
10. Enthalpy	Calorimetry	$q = m c \Delta T$
12. Voltaic Cells	Cell potential	$E_{cell} = E_{cathode} - E_{anode}$