Enthalpy of Neutralization Virtual lab

Name: Lachlan Campbell Partner: Nathan Hannam Student No: 20279617 Student No: Lab Section: 40 Bench # (on computer screen): 36 Experiment 11: (ENT) Enthalpy of Neutralization Purpose The purpose of this experiment is to analyze both quantitatively and qualitatively the effects of combining a strong acid and a strong base solution, weak acid and strong base, and a strong acid and weak base. Following the rules of calorimetry, the enthalpy change can be calculated for the above reactions. Introduction Identifying whether a reaction is spontaneous is important because it determines what reactions are likely to occur naturally, and whether the products or the reactants are favoured. The spontaneity of a reaction can be determined by Gibbs free energy with the variable  G. The equation below shows how it can be used. ∆G=∆H−T∆S The change in the state of disorder (or entropy change) is defined with the variable ∆S. T is the temperature that the reaction takes place at. Finally, ∆H displays the enthalpy change in the reaction. If the free energy is negative, then the reaction is spontaneous and positive means the reaction requires energy to start. This experiment will go over both exothermic and endothermic reactions, so being able to use another formula (shown below) may be helpful for the experiment. ∆G ° = Σ∆G f − Σ ∆G i This formula uses the exact free energy of the reactants and subtracts them from the products, overall calculating the final change in free energy. Procedure Part A 1) Set up the calorimeter. Fold paper towel around a smaller beaker and place inside a larger beaker. Stir bar in smaller beaker on top of the magnetic stirrer. Place a cardboard lid on top of the calorimeter. 2) Obtain 50 mL solutions of NaOH and HCl and place them in separate beakers. 1

3) Ready logger pro to collect temperature over a time of 1200s 4) Mix both solutions, while having one of the temperature probes in an empty beaker to act as a control. 5) Once you have obtained data that looks to be a good straight line, then rinse the beakers and prepare for Part B. Part B: 1) Follow the same steps for Part A except with different solutions. First, 50 mL of NaOH and HOAc and then 50 mL of NH 3 and HCl. 2) Make sure to graph the data using a temperature time graph from the data collected on logger pro. Data and Observations Questions 1. Write the balanced overall and net ionic equations for the: a. Neutralization - Strong Acid/Strong Base (overall and net ionic) Overall Equation: HCl (aq) + NaOH (aq) NaCl (aq) + H 2 O (l) Net Ionic Equation: H ( aq ) + ¿¿ + OH ( aq ) − ¿¿ H 2 O (l) b. Neutralization - Strong Base/Weak Acid (overall and net ionic) Overall Equation: NaOH (aq) + CH 3 COOH (aq) H 2 O (l) + CH 3 COONa (aq) Net Ionic Equation: CH 3 COOH (aq) + OH − ¿¿ H 2 O (l) + CH 3 COO − ¿¿ (aq) c. Neutralization - Weak Base/Strong Acid (overall and net ionic) 0verall Equation: NH 3(aq) + HCl (aq) NH 4 Cl (aq) Net Ionic Equation: NH 3(aq) + H + ¿¿ (aq) NH 4 + ¿¿ (aq) 2. Use the temperature data you collected to compute the enthalpies for the two reactions: Put your answers in kJ/mol. Reaction 1: ∆ r H ° = − C cal ∆T n C = Q / m  T = 337.15 J 2

∆ r H = ∆ H ( products ) − ∆ H ( reactants ) = (-486.01 + - 285.8) – (-485.8 + -229.994) = -56.016 ∆ S = Σ Products − Σ Reactants = (86.6+69.91) –(178.7 -10.75) =-11.44 ∆G = ∆ H − T ∆ S = (-56.016) – (298.15) (-0.01144) = -52.61 kJ/mol Reaction 3: ∆ r H = ∆ H ( products ) − ∆ H ( reactants ) = (-133.077) – (-46) = - 87.077 kJ/mol ∆ S = Σ Products − Σ Reactants = (105.5) – (130.6) = -25.1 ∆G = ∆ H − T ∆ S = (-87.077) – (298.15) (-0.0251) = -79.5934 kJ/mol 4. Write the balanced equation for the equilibrium of your weak acid or weak base (whichever one you used) in water (a K a equation, NOT a neutralization equation). Calculate its Standard Gibbs energy using the thermodynamic data supplied (Error: Reference source not found) with Equation 49. CH 3 COOH (aq) + H 2 O (l) CH 3 CO O − ¿ ¿ (aq) + H 3 O (aq) K a = ¿¿ = = 1.8 × 10 -5 ∆G ° = Σ∆G ( products ) − Σ ∆G ( reactants ) = (-369.31) – (-396.5 + -237.1) = 264.29 kJ/mol 5. Comment on the spontaneity of the following chemical reactions: a. Compare the two neutralization reactions (strong acid/strong base with the strong/weak combination) you did. 4

The reaction between the strong acids and bases was more spontaneous than those of the weak/strong combinations. Since the weak acids and bases caused only a limited dissociation in water, less heat was released in those reactions, whereas with a strong base and a strong acid, much more of the product would be dissociated, equally out the pH of the two reactants, and producing a lot of heat. The bigger the change in temperature the larger negative value of Gibbs free energy. b. Compare the strong weak combination with the simple dissolution of your weak acid or base in water (calculated in question 4.). The reaction of the weak base in water resulted in what is called an endothermic reaction since there is a positive delta G value. As we know, the strong/weak combination is an exothermic reaction. This is because when a weak base only reacts with water, the water’s priority is to help dissociate the weak base and ends up absorbing a lot of energy. c. Which is more spontaneous, the acid/base reaction or the weak (acid or base) with water? Why is your selected weak (acid or base) called weak? The acid and base reactions produce a negative delta G value which means that the products are favourable, and the reaction is spontaneous. Like stated above, the weak base reaction causes a positive delta g, so the reaction is actually non-spontaneous. A weak base is called ‘weak’ because it lacks the ability to fully dissociate in water, which means that not all the hydroxide ions are donated, and many still contain their ionic charge. 6. List 2 sources of systematic error. One potential source for error in this lab could be the set-up of the calorimeter. It is imperative that little heat is lost within the system, and if there are any leaks or improper materials like carboard, it could leak and cause drastically different results. Another potential source of systematic error was the amount of either acid or base mixed. If there was a discrepancy and too much acid was added versus a smaller amount of base, then eventually the products would be more acidic rather than neutralized, overall affecting the results. References [Each reference listed here should have a number that corresponds to the superscripted number in the body of the report where this reference is pertinent.] 1. Department of Chemistry. Queens University. (2021-2022). Chemistry 112 First Year Lab Manual. 2. 19. 4: Entropy changes in chemical reactions . (2014, November 20). Chemistry LibreTexts. https://chem.libretexts.org/Bookshelves/General_Chemistry/Map %3A_Chemistry_-_The_Central_Science_(Brown_et_al.)/ 5

0 20 40 60 80 100 120 140 160 180 200 220 240 260 280 300 320 340 360 380 400 420 440 460 19.5 20.5 21.5 22.5 23.5 24.5 25.5 26.5 NH3 + HCl Series1 Linear (Series1) Series2 Moving average (Series2) Time (s) Temperature (C) 7