Enthalpy change in chemical reactions

ENTHALPY CHANGE IN CHEMICAL REACTIONS

Purpose

This experimentwas conducted to study thermal changes in chemical reactions. Throughthis, the theory behind the heat transfers and wells definition ofterms is mastered. In order to achieve understanding of the enthalpychanges during chemical reactions, the explanations of the absorptionor evolution of heat energy by chemical systems are studied. Theexperiment is also conducted to study the factors behind thedeviation of experimental enthalpy values from the theoreticalvalues. In order to achieve all this, the enthalpy change for thereaction between sodium hydroxide and hydrochloric acid, and heat ofsolution of potassium nitrate are studied.

Reactions areusually accompanied with energy changes. Change of the thermalproperties of reactants or a system can be determined through thehelp of an apparatus known as the calorimeter. Calorimetry is thetechnique that is employed in experiments to measure the amount ofheat a system evolves, absorbs or transfers to its surrounding(Hemmings, H. C. 2005). Zielenkiewicz, W., &amp Margas, E. (2004)explain that the laws and relations in the theory of heat transferare employed in the analysis of the thermal processes that take placein a calorimeter its environment.

Calorimetry canbe used in finding the unknown amount of heat and/or more informationabout the heat flow in a given chemical or physical process, as inthis experiment. Through this, the heat capacity of various objectsin a system that evolves or absorbed heat can be determined. Heatcapacity, also called thermal capacity, is defined by Hemmings, H. C.(2005) as the amount of heat required to raise the temperature of abody by one kelvin, 1K. When two objects A and B of different heatcapacities are brought together in a closed system, both objects willattain the same temperature (Ramsden, E. N. 2000). For example,object A has a lower temperature and gains resulting in an increasein its temperature. This can be denoted by the equation below

Heat capacity=(heat supplied)/(temperature rise)

Heat capacity andspecific heat capacity various with different objects. The heat thata unit mass of the object absorbs in order to raise its temperatureby one kelvin is referred to as specific heat capacity. It isrepresented by the equation (Ramsden, E. N. 2000)

Specific heatcapacity= (heat capacity)/(mass of the object)

Mass X specificheat capacity (c) = (heat supplied)/(rise in temperature)

Heat = Mass (m) Xspecific heat capacity (c) X temperature change (ΔT)

ΔH=m x c x ΔT

The amount ofheat object B requires in order to raise the temperature of object Aby one kelvin is known as a calorie. Calorie and heat of reactiondiffer in the sign as one is associated with loss of heat energywhile the other is associated with heat gain. They are represented bythe equation

qrxn = -qcal

Chemicalreactions are accompanied with reactants and products. The differencein enthalpies of the products and reactants in a chemical reaction ata constant pressure when one mole of the reactants completely reactsis referred to as enthalpy change. Enthalpy change is represented byΔH

ΔH= (qrxn(enthalpy of products-enthalpy of reactants))/(moles of reactants)

Enthalpy changecan be either exothermic or endothermic. Endothermic reaction is thatwhich absorbs heat energy from the surrounding and is positive whileexothermic reaction gives off heat energy and is thus negative(Palanna, O. G. 2009).

Experimental

The first part ofthis experiment was involved with the heat associated with thereaction between hydrochloric acid and sodium hydroxide. 75ml of 2MHCl was poured into the calorimeter, and its temperature measured andrecorded before adding an equal amount of 2M NaOH. The mixture wasthoroughly stirred with a thermometer. The highest temperature wasobserved and recorded, and the rate at which the temperaturedecreases as well as the low temperature was recorded. The heatabsorbed was calculated (laboratory manual).

Second part: 150ml water (500C) was measured by the calorimeter (insulating Styrofoamcup) and 25.0g of KNO3 added into the water while vigorouslystirring. The temperature of the system was recorded after every 30seconds until a uniform solution was achieved. The heat of solutionwas calculated using the data collected (laboratory manual).

Data and results

The followingtable shows the results for the experiment for the reaction betweenHCl and NaOH.

A

Acid

Base

Total Vol.(ml)

Ti (°C)

Tf (°C)

Hrxn (J)

H/mol (J/mol)

HCL

NaOH

150 ml

21.2 c

34.5 c

-8347.1

-55,647.3

The Hrxn (J)was arrived at by

Finding the totalmass of the reactants in grams and the calculating the change intemperature (34.5 -21.2=13.30C). The specific heat capacity of wateris used in the calculation

Hrxn (J) = 150X 4.18 X 13.3

=-8339.1 J

Number of molesof 2M HCl in 75ml is 0.15 moles ((75 x 2)/1000)

=-8339.1/0.15

H/mol =-55, 594 J/mol

CalculatingUncertainty

T= 34.5-21.2=13.3 0C

δ(T)=√((0.2)2+(0.2)2)

= 0.28

Hence, T is13.3 ± 0.28

δ m=√((2g)2+(2g)2)

=2.8g

Uncertainty inthe heat of reaction is

δ qrxn= -8339.1J x √((2.8/150)2 + (0.28/13.3)2)

=-8339.1 J x√ (0.000348 + 0.000443)

=-8339.1 J x0.028

= -234 J

The error incalculating H/mol is approx.. 3%. The calorimeter heat value istherefore 8339.1 ± 234 J.

volumeMass (g) Ti (°C) Tf (°C) Hsoln (J) H/mol (J/mol)

KNO3 150 ml24.9 22.9 13 7245.5 1767.2

The heat absorbedby water plus the heat lost by the salt is called the heat ofsolution.

Total mass of thesolution is 150 + 24.9= 174.9g.

Hsoln (J)=174.9 x 4.18 x 9.9

= -7237.71 J

H/mol (J/mol)= (-7237.71 J)/(moles of KNO3)

Moles ofKNO3= 24.9/(39+(16×3)+14)

=0.2465 moles

H/mol (J/mol)=(-7237.71)/0.2465

= -29,354.9 J/mol.

Discussion

The two reactionsin this experiment were accompanied by a rise in temperature and thusindicating that they were exothermic reactions. Exothermic reactionslose thermal energy to its surrounding and, therefore, the enthalpyassociated with the chemical reactions was negative. The transfer ofheat or enthalpy change was achieved by the need for the system toattain an equilibrium within the system and with its environment(Hemmings, H. C. 2005).

The experimentwas conducted using insulating Styrofoam cup (calorimeter), and itwas assumed that there was no heat energy lost to the surrounding.The heat absorbed by the insulating Styrofoam cup was not accountedin the calculations of heat change as it was assumed that anegligible amount of heat was absorbed by this object. The heat thatthe chemical reactions dissipate or absorb is significantlyinfluenced by the ions of the reacting chemicals. Any deteriorationin the composition of the reactants would alter the expected results.The purity of the chemical reactants was not taken into considerationin the calculations. The heat capacities of NaOH and HCl were notaccounted for when calculating the enthalpy change for the reactionbetween HCL and NaOH but instead that of water was used. Since thevolume of KNO3was less than that of water, the enthalpy change forthe formation of KNO3solution was calculated using the specific heatcapacity of water. Therefore, the values calculated in thisexperiment deviated from theoretical ones.

A good example isthe calculation of the heat of solution of KNO3. Theoretically,enthalpy of solution should be calculated using heat of fusion ofKNO3, K+ and NO3-(Palanna, O. G. 2009):

Hsoln = Hf(K+(aq))+ Hf (NO3-(aq)) – Hf(KNO3(s))

=-252.4kJ +-207.4kJ – (-494.6kJ)

=34.8kJ

The possiblesources of error in the experiment include faulty of the apparatus,purity of the chemicals or incorrect measuring equipment. Theobserver had in the important role of ensuring that accurate readingsof the mass, temperature and volume were taken. Careless reading mayhave resulted in a deviation from the expected results as it providedthe data used in the calculations. The outcome of the experimentcould have been improved through the use of calorimeter that is ingood condition, well-functioning weighing machines, and care whentaking the reading and also throughout the experiment. However, theconcept of calorimetry has been well understood along with the theorybehind the reactions and heat transfers that bring about the enthalpychange.

References

Cheng, S. Z. D. (2002).&nbspHandbook of thermal analysis andcalorimetry: Volume 3. Amsterdam: Elsevier.

Hemmings, H. C. (2005).&nbspFoundations of anesthesia: Basicsciences for clinical practice. St. Louis, Mo. London: ElsevierMosby.

Mortimer, R. G. (2008).&nbspPhysical chemistry. Amsterdam:Academic Press/Elsevier.

Palanna, O. G. (2009).&nbspEngineering chemistry. New Delhi:Tata MrGraw-Hill.

Ramsden, E. N. (2000).&nbspA-level chemistry. Cheltenham:Stanley Thornes.

Zielenkiewicz, W., &amp Margas, E. (2004).&nbspTheory ofCalorimetry. Dordrecht: Kluwer Academic Publishers.