Enthalpy of Neutralization (Calorimeter)

A calorimeter measures heat changes in chemical reactions. Simple calorimeter: Polystyrene cup with lid, thermometer, and stirrer. Polystyrene provides good insulation to minimize heat loss. Reaction procedure: Measure known volumes of acid and base separately, record initial temperatures, mix in calorimeter, record maximum temperature reached. Key assumptions: No heat loss to surroundings, specific heat capacity of solution equals water (4.18 J/g°C), density of solution equals water (1 g/mL). Calculations: Total mass = volume acid + volume base (assuming density 1 g/mL), Heat evolved: q = m × c × ΔT, Moles water formed = moles H⁺ or OH⁻ (limiting reagent), ΔH = -q / moles water.

For strong acid-strong base reactions, the enthalpy is approximately -57.3 kJ/mol. Examples: HCl + NaOH → NaCl + H₂O, ΔH = -57.3 kJ/mol; HNO₃ + KOH → KNO₃ + H₂O, ΔH = -57.3 kJ/mol; H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O, ΔH = -57.3 kJ/mol per mole H₂O. The consistency arises because strong acids and bases are completely ionized in solution, so the net reaction is always: H⁺(aq) + OH⁻(aq) → H₂O(l). The heat comes from the formation of O-H bonds in water molecules and the neutralization of charges. The experimental value may vary slightly due to specific heat capacities of solutions and heat losses.

Weak acids/bases show different enthalpy values: Weak acid-strong base: Less exothermic than -57.3 kJ/mol because energy is required to ionize the weak acid. Example: CH₃COOH + NaOH → CH₃COONa + H₂O, ΔH ≈ -55 kJ/mol. Strong acid-weak base: Also less exothermic due to ionization energy of weak base. Example: HCl + NH₄OH → NH₄Cl + H₂O, ΔH ≈ -52 kJ/mol. Weak acid-weak base: Least exothermic, around -50 kJ/mol. The difference from -57.3 kJ/mol represents the ionization enthalpy of the weak electrolyte. This principle allows determination of ionization constants from calorimetric data.

Step 1: Measure 50 mL of 1M HCl into calorimeter, record temperature T₁. Step 2: Measure 50 mL of 1M NaOH, record temperature T₂. Step 3: Average initial temperature = (T₁ + T₂)/2. Step 4: Quickly add base to acid, stir, record maximum temperature T₃. Step 5: ΔT = T₃ - average initial temperature. Step 6: Total volume = 100 mL, mass = 100 g (assuming density 1 g/mL). Step 7: q = 100 × 4.18 × ΔT J. Step 8: Moles water = 0.05 (from 50 mL of 1M solutions). Step 9: ΔH = -q / 0.05 J/mol = -q / 50 kJ/mol. Example: If ΔT = 6.8°C, q = 100 × 4.18 × 6.8 = 2842 J, ΔH = -2842 / 50 = -56.8 kJ/mol.

Heat losses: To minimize, use insulated calorimeter, perform quickly, cover calorimeter. Heat capacity: Assumption that solution has same heat capacity as water introduces small error. Temperature measurement: Use accurate thermometer, read at eye level, stir continuously for uniform temperature. Concentration accuracy: Use standardized solutions. Reaction completeness: Ensure stoichiometric proportions. Evaporation: Keep covered to prevent cooling by evaporation. Heat of dilution: Acid and base solutions may have significant heats of dilution. Corrections: Account for heat capacity of calorimeter (calibrate with known reaction). Typical experimental error: ±1-2 kJ/mol.

Enthalpy of neutralization measurements are used for: Thermodynamic studies - understanding energy changes in reactions, Acid-base characterization - distinguishing strong and weak electrolytes, Calorimeter calibration - using known reactions to determine heat capacity, Industrial processes - designing neutralization reactors with heat management, Environmental applications - treating acidic or basic waste streams, Educational purposes - teaching thermochemistry principles. The consistency of -57.3 kJ/mol for strong acid-strong base reactions demonstrates the concept of net ionic equations and provides a fundamental reference point in thermochemistry.