UNIVERSITI TUNKU ABDUL RAHMAN
FACULTY OF SCIENCE
DEPARTMENT OF CHEMICAL SCIENCE
PERAK CAMPUS

BACHELOR OF SCIENCE (HONS) CHEMISTRY
YEAR 1
UDEC 1134
CHEMISTRY LABORATORY I
Name: Ong En-Ming
Student ID: 1404405
Title of Experiment: Determination of The Enthalpy (Heat) of Reaction of A Monobasic Acid with Sodium Hydroxide
No. of Experiment: 12
Date of Experiment: 2/2/2016
Date of Submission: 24/2/2016
Group members: 1)Chong Chi Wei
2)Mong Lai Wan
3)Ang Yen Yuan
Name of Lecture: Dr. Sim Yoke Leng
Title: Determination of the Enthalpy (Heat) of Reaction of A Monobasic
Acid with Sodium Hydroxide

Objective: 1. To understand the enthalpy chemistry. 2. To determine the calorimeter constant. 3.To determine the enthalpy of reaction of acid-base reactions.

Introduction: Heat is associated with nearly all chemical reactions. In such instances, the reaction either liberates heat (exothermic) or absorbs heat (endothermic). When a reaction is carried out under constant pressure (as in an open beaker) the heat associated with the reaction is known as enthalpy. The symbol for enthalpy is ΔH. It is most often too difficult to direct measure the enthalpy change for a reaction. What can be done is to measure the heat changes that occur in the surrounding by monitoring temperature changes. Conducting a reaction between two substances in aqueous solution, allows the enthalpy of the reaction to be indirectly calculated with the following equation. q=m x c x Δ t The term q represents the heat energy that is gained or lost. C is the specific heat of water, m is the mass of the water and Δ t is the temperature change of the reaction mixture. The specific heat and mass of water will either gain or loss heat energy in a reaction that occur in aqueous solution. Normally, the change in enthalpy that occur as a result of a chemical reaction is numerically equal to the heat of reaction under constant (atmospheric) pressure conditions (Δ H=q). The heat of reaction is conveniently measured adiabatically in a Dewar calorimeter by the rise or fall in temperature of the products produced by the reactions in solution. The “calorimeter constant” must first be determined. This is the quantity of heat required to increase the temperature of the calorimeter and its constant by 1°C. The constant is measured by supplying the calorimeter and contents with a definite known quantity of heat. This can be done electrically or by adding a known amount of concentration sulphuric acid. Hess’s Law This principle (Hess’s Law) states that the enthalpy change for a reaction depends on the products and reactants and is independent of the pathway pr the number of steps between reactants and products. In the other words, if a reaction is carried out in a series of steps, H for the reaction will be equal to the sum of the enthalpy changes for the individual steps. Therefore, the enthalpy change for a given reaction is calculated by adding the individual chemical equations and taking the sum of the enthalpy changes associated with each of these individual chemical equations.

Apparatus: Dewar flask, stopwatch, thermometer (5 to + 50 °C, graduated in 1/10 °C), graduated pipette fitted with a suction bulb, 10 cm3 graduated cylinder, Materials: concentrated sulphuric acid (specific gravity 1.84, 98.5% H2 SO4),conc. nitric acid, 1M sodium hydroxide, 0.1 M hydrochloric acid, methyl orange indicator. Procedure: A) Calorimeter constant 1. 100 cm3 of water is pipetted into the Dewar flask. The water is stirred slowly and regularly with a 1/10 °C thermometer. 2. The temperature is noted at intervals of 1 minute over a period of five minutes or so. 3. At the end of this period introduce into the calorimeter about 2 cm3 of conc. sulphuric acid using a graduated cylinder. 4. The temperature is continued recording at 1 minute intervals whilst continuing to stir until the rate fall of temperature has become constant about 10 minutes. 5. The contents of the calorimeter are allowed to cool. 6. About 25 cm3 of the solution is titrated against 1 M sodium hydroxide to determine the molarity of the solution. B) Enthalpy of reaction 1. In the calorimeter of known constant 50 cm3 of 1 M sodium hydroxide and 50 cm3 of water are mixed. 2. The temperature observation is made as already described. 3. At a known time, 5 cm3 of 10 M nitric acid is added from a graduated pipette. (Prepare 10 M acid by diluting 65 cm3 conc. nitric acid to 100 cm3 with distilled water). 4. The temperature is continued observations. 5. The solution in the calorimeter is added a few drops of methyl orange indicator to ensure that the solution is added. 6. If the solution is alkaline, titrate it in the calorimeter with 0.1 M hydrochloric acid to determine the exact amount of alkaline neutralized during the experiment. ii) The experiment described in (i) above is repeated but substitute 100 cm3 of distilled water for the (50 cm3 1 M sodium hydroxide + 50cm3 water) mixture. Results and Calculation: Part A: For calorimeter constant Time (min) | 0 | 1 | 2 | 3 | 4 | 5 | Temp. (°C) | 25 | 25 | 25 | 25 | 25 | 25 |

After added conc. sulphuric acid Time (min) | 0 | 0.5 | 1.0 | 1.5 | 2.0 | 2.5 | 3.0 | 3.5 | 4.0 | 4.5 | 5.0 | Temp. (°C) | 28 | 29 | 29 | 29 | 29 | 29 | 29 | 29 | 29 | 29 | 29 |

Initial reading of burette = 25.0 cm3
Final reading of burette = 39.6 cm3
Volume of NaOH used = 39.6 cm3 - 25.0 cm3 = 14.6 cm3

Different in temperature (Δ t) = T2 –T1 =(29℃ - 25℃) =4℃
Part B: For Enthalpy of reaction

Time (min) | 0 | 1 | 2 | 3 | 4 | 5 | Temp. (°C) | 24 | 24 | 24 | 24 | 24 | 24 |

After added nitric acid Time (min) | 0 | 0.5 | 1.0 | 1.5 | 2.0 | 2.5 | 3.0 | 3.5 | 4.0 | 4.5 | 5.0 | Temp. (°C) | 27 | 29 | 29 | 29 | 29 | 29 | 29 | 29 | 29 | 29 | 29 |

Different in temperature (Δ t) = T2 –T1 =(29℃ – 24℃) =5 ℃
(ii)
Time (min) | 0 | 1 | 2 | 3 | 4 | 5 | Temp. (°C) | 23 | 23 | 23 | 23 | 23 | 23 |

Time (min) | 0 | 0.5 | 1.0 | 1.5 | 2.0 | 2.5 | 3.0 | 3.5 | 4.0 | 4.5 | 5.0 | Temp. (°C) | 24 | 24 | 24 | 24 | 24 | 24 | 24 | 24 | 24 | 24 | 24 |

Different in temperature (Δ t) = T2 –T1 =(24℃– 23℃) =1℃
Calculations:
Part A: Determination of Calorimeter constant. i) Based on the table one, a straight line graph has been drawn starting from the origin. ii) The theoretical molarity of the sulphuric acid was first calculated as shown below and then trace through the graph drawn to determine the heat liberated, Q. a) 98.5% concentrated sulphuric acid means 98.5 mL of sulphuric acid in 100ml. b) 2 cm3 of sulphuric acid has a density of 1.84 g/cm3 and molar mass of 98.079g/ mol.
Density= massvolume mass = density × volume mass=1.84gmL×98.5 mL=181.24 g number of moles=massmolar mass number of moles=181.24 g98.079 g/mol=1.8479 mol molar concentration=1.8479 mol100×10-3 L=18.479 mol/ L

After diluting by adding 100 cm3, the molar concentration is m1v1=m2v2 18.479×2=m×102.00 m=0.3623 mol/L
∴Therefore, the heat liberated found by tracing from the graph is 2.575kJ

iii) a) The experimental molarity of the sulphuric calculated from the results obtained through titration against
2NaOH + H2SO4 Na2SO4 + 2H2O nH2SO4n(NaOH)=12 nH2SO4 = 0.5 n(NaOH) nH2SO4=0.5 25.00 ×10-3 L(1.0 mol/L) nH2SO4=0.0125 mole molarity of H2SO4=0.0125mole50.00 ×10-3mL molarity of H2SO4=0.250 mol/ L

∴Therefore, the heat liberated found by tracing from the graph is 1.775kJ.

iv) Percentage difference in molarity of H2SO4=0.3623-0.2500.3623 ×100%
Percentage difference in molarity of H2SO4 = 31%

v) Calorimeter constant
C=Qexperimentm∆T

C= 1.775kJ100.00mL+2.00mL( 4oC)
C=0.004350 kJ g-1 oC-1

Part B (ii): Determination of the heat of dilution, Qd Q = mcΔT3 =Δ Hd mcΔT3 =Δ Hd Δ Hd = 100.00+5.00g 0.004350 kJ g-1 oC-1(1oC) Δ Hd = 0.45675 kJ
∴Since it is an exothermic reaction, Δ Hd is - 0.2676 kJ

Part B (i): Determination of the heat of neutralisation, Qn Q = mcΔT2 =Δ Hd +Δ Hn mcΔT2 =Δ Hd +Δ Hn
(100 + 5)g 0.004350 kJ g-1 oC-1(5℃) = 0.45675 kJ +Δ Hn
Δ Hn = 2.28375 kJ - 0.45675 kJ
Δ Hn = 1.827 kJ ∴Since it is an exothermic reaction, Δ Hn is - 1.827 kJ.

Discussion： The fundamental change common to all neutralization reactions is the reaction between H+ (aq) ions from the acid and OH-(aq) ions from the alkali to form water molecules. H+ (aq) + OH-(aq) H2O (l) Strong acids and strong alkalis dissociate almost completely in aqueous solution. This is confirmed by the fact that the standard molar enthalpy change of neutralization of any strong acid and any strong base is almost constant, that is, -57.3 kj mol-1. NaOH(aq) + HCl(aq) NaCl(aq) + H2O (l) Δ H = -57.3 kJmol-1 ½ Ca (OH)2(aq) + HNO3(aq) ½ Ca(NO3)2(aq) ++ H2O (l) ΔH = -57.3 kJmol-1 When sodium hydroxide ( a strong alkali ) and hydrochloric acid ( a strong acid ) are mixed, the following reaction take place. Na+(aq) + OH-(aq) + H+(aq) + Cl-(aq) Na+(aq) + Cl-(aq) + H2O (l) The products of the reaction are sodium chloride (salt) and water. However, Na+(aq) ions and Cl-(aq) ions do not take part in the above reaction. So in fact, the standard enthalpy change of neutralization between a strong acid and a strong alkali is the standard enthalpy change of formation of 18 g of water from 1.0 g of H+ (aq) ions and 17 g of OH-(aq) ions.
H+(aq) + OH-(aq) H2O (l) ΔH = -57.3 kJmol-1
In this experiment, the reaction between two acids such as sulphuric acid and nitric acid is an exothermic reaction. This reaction is exothermic due to the O-H bond in H3O+ is stronger than the O-H bond found in H2SO4 and HNO3. The stronger the bond formation, the higher the heat energy will be released in order to become more stable. In addition, nitric acid is a monobasic acid because it can only ionised 1 time when added into water with the equation. HNO3 + H2O → H3O+ + NO3- Oppositely, sulphuric acid is a diprotic acid which can ionized 2 times due to the 2 hydrogen molecules. H2SO4 + H2O → H3O+ + HSO4- HSO4- + H2O → H3O+ + SO4- In this experiment, the reaction between nitric acid and sodium hydroxide is an acid base reaction and also an exothermic reaction. The acid base reaction will released heat due to the formation of water molecules. The equation between nitric acid and sodium hydroxide is shown at below: HNO3 + NaOH→ H2O + NaNO3 H+ + OH- → H2O(ionic equation) Exothermic reaction is a chemical reaction that releases energy by light or heat and it is oppositely of an endothermic reaction. In this experiment, a closed condition is required to carry out because as to prevent heat lost from the system. But due to inconvenience, the cover of the Dewar flask (calorimeter) has been opened for the solution to be stirred. From this, there was an open system which resulted heat lost in the process. As the experiment being carried out, more heat lost occurred in this process. This causes an error which classified as systematic error. Besides, concentrated sulphuric acid is a strong acid with the molecular structure of H2SO4, after it was added by amount of water , the solution will become a dilute sulphuric acid ,while for the sodium hydroxide which is a strong alkali, after it was added by amount of water, the solution will become a dilute sodium hydroxide. For the methyl orange indicator it used as indicator in the titration. Methyl orange is used in titrations because of its clear and distinct colour change. Methyl orange was used to determine the acidic and the alkaline of the solution. When the solution is an acidic solution, this indicator solution will turn into red colour, for contrast, if the solution is an alkaline solution, this indicator will change from red colour to yellow colour. The time taken for the color change of this indicator was depends on the acidic or alkaline of the solution. The strongest the acid, the time taken for the indicator color shown is shorter. For the first part of experiment, to determine the calorimeter constant, the conc. sulphuric acid which process to dilute sulphuric acid, the temperature of dilute sulphuric acid was noted at intervals of 1 minute over a period of five minute. After this time taken, the solution was titrated using 1 M of sodium hydroxide. A few drops of methyl orange indicator were used to determine the acidic or alkaline of solution. The methyl orange indicator was change the color from red color to yellow color, this show that the solution become alkaline after adding of 14.6 cm3 of 1 M sodium hydroxide. This result show that the dilute sulphuric acid was became a weak acid which does not dissociate completely, therefore when the acid react, some of the heat energy liberated during neutralization is absorbed in the dissociation of the acid. CH3COOH CH3COO-(aq+H+(aq) ΔH= positive For the second part of experiment, the 50cm3 1M sodium hydroxide +50cm3 water also noted it temperature at known time, the solution was added by 10 M nitric acid. The temperature observation was continued. A few drops of methyl orange indicator were added and the color of indicator turn into red color which means the solution after mixed was become acidic solution. On the other hand, the 50 cm3 of water was changed to 100 cm3 of water and the procedure was repeated as described above. The color of methyl orange indicator also turn into red color, which means that the solution remains was acidic solution. There are some precaution steps that needed to be followed in this experiment. Firstly, we must handle the thermometer carefully because the thermometer bulb is very fragile. Besides, the sulphuric acid and sodium hydroxide is very toxic, must avoid by touching by bare hand and avoid contact with eyes. In addition, the reading of the thermometer must read perpendicularly with the thermometer to prevent of misleading records. Moreover, all the materials needed must take it from fume hood. Lastly, stay away to any high air current area to prevent heat lost of the calorimeter. Conclusion: In conclusion, the objectives in this experiment have been achieved. The calorimeter constant used in this experiment 0.004350 kJ g-1 oC-1. The enthalpy of the acid base reaction has also been determined with the value of -1.827 kJ. The process of heat of neutralization is an exothermic reaction which will denoted with a negative sign and this show that heat is released to the surrounding. Dewar flask is not so suitable using in this experiment because it cannot provide a better close system and it can be replaced by bom calorimeter (it can provide a better close system).

References:

Anonymous. (2004). Thermodynamics- Enthalpy. Retrieved 2 20, 2016, fromhttp://www.chem.tamu.edu/class/majors/tutorialnotefiles/enthalpy.m UDEC 1224 Chemistry laboratory I. (2016). Experiment 12: Determination of The Enthalpy (Heat) of Reaction of A Monobasic Acid with Sodium Hydroxide: Lab Manual for Bachelor of Science (Hons) Chemistry Year 1.Kampar Utar. (pp. 23-24) [ONLINE] Available at: http://wble-pk.utar.edu.my/file.php/12032/UDEC_1224_-_Chemistry_Laboratory_I_Manual.pdf [accessed on 20 February 2016].

Martin S. Silberberg. (2009). Chemistry, 5th edition, New York: McGraw-Hill Higher Education Chemwiki.ucdavis.edu, (2013). Exothermic vs. Endothermic and K - Chemwiki. [online] Available at: http://chemwiki.ucdavis.edu/Physical_Chemistry/Equilibria/Le_Chatelier's_Principle/ Effect_Of_Temperature_On_Equilibrium_Composition/Exothermic_Versus_Endothe rmic_And_K [Accessed 20Feb. 2016].