Thursday, June 2, 2011

Thermodynamic Study on reaction between an Acid and a Base

Objectives

1. To study the enthalpy chemistry

2. To study the exothermic reactions

3. To determine the calorimeter constant

4. To determine the enthalpy of reaction of acid-base reactions

Introduction

Enthalpy is defined as the total amount of total energy of a thermodynamic system. The enthalpy of a system includes the internal energy and the work done of the system. The work done is equals to the product of the volume of the system multiplied by the pressure exerted on it by its surrounding, as shown in the following:

H = U + w

H = U + pV

where H = Enthalpy

          U = Internal Energy

          W = work done

           p = Pressure

           V = Volume

Enthalpy of a thermodynamic system is usually measured in S.I. unit of Joule, J or kiloJoule, kJ. However, we often measure the change in enthalpy, ΔH instead of measuring the value of enthalpy of the system, H because the total enthalpy of the system cannot be measured directly.

Change in enthalpy, ΔH is defined by the following equation:

ΔH = HFinal – HInitial

ΔH = Enthalpy change

HFinal = Final enthalpy of the thermodynamic system. In a chemical reaction, HFinal is the enthalpy of the products.

HInitial = Initial enthalpy of the thermodynamic system. HInitial is the enthalpy of the reactants in a thermodynamic reaction.

In the endothermic reaction, the magnitude of a thermodynamic reaction is shown in positive. By contrast, the exothermic reaction shows in negative value. Exothermic reactions involve heat energy transferred from the system into its surroundings, causing the temperature of the surroundings to rise. On the contrary, an endothermic reaction involves the energy acquired from the surrounding into the system.

The change in enthalpy of a thermodynamic system either endothermic or exothermic reaction is numerically equals to the magnitude of heat of the reaction which is ΔH = q. The heat of reaction is easily to be measured adiabatically by using a Dewar flask. The increase or decrease in temperature of the products produced by the reaction in solution is being measured and compared with the initial temperature of the system. The difference in the initial value of enthalpy and final value of the enthalpy is the change of the enthalpy. The Dewar flask is being used because its design is to preserve the heat from loss or minimize heat loss to the surrounding from the system. In addition, an isolated system is also needed in this experiment to obtain a more accurate data. Each Dewar flask used in the experiment has different calorimeter constant since each Dewar flask has difference in materials used. In order to measure the total amount of heat in a chemical reaction, the calorimeter constant ( ) must be firstly determined. The calorimeter constant is defined as the quantity of heat required to increase the temperature of the calorimeter and its content by 1 °C.

Ccal = ∆H / ∆T

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 concentrated sulphuric acid.

Results & Calculation

Part 1 Calorimeter Constant

From Graph 1,

Δ T1 = 31.0°C – 22.8°C

= 8.2 °C

From Table 4,

Volume of sodium hydroxide solution used = 34.2 cm3 – 15.0 cm3

= 19.2 cm3

Moles of sodium hydroxide used = 1M x (19.2/1000) dm3

= 0.0192 moles

H2SO4 + 2 NaOH àNa2SO4 + 2H2O

Moles of sulphuric acid used = 1/2 x 0.0192 moles

= 0.0096 moles

Molarity of sulphuric acid solution = 0.0096 moles / (25/1000) dm3

= 0.384M

By using the values given, 0.384M in Graph 2, we can estimate the amount of heat liberated which is 2.775 kJ. Since the heat liberated by the dilution of sulphuric acid is then absorbed by the Dewar flask, so the value of the ΔH is in positive value.

From Graph 2,

Ccal = ΔH1 / ΔT1

= 2.775kJ / 8.2 °C

= 0.338 kJ / °C

Therefore, the calorimeter constant has a value of 0.338 kJ / °C.

Part 2 Enthalpy of Reaction

Part I

HNO3 + NaOH à NaNO3 + H2O

From Graph 3,

ΔT2 = 29.5°C – 22.0°C

= 7.5 °C

ΔH2 = Ccal x ΔT2

= 0.338 kJ / °C x 7.5 °C

= 2.535 kJ

The amount of heat absorbed by the Dewar flask after adding nitric acid to the mixture (50 cm3 of water + 50 cm3 of 1M sodium hydroxide) is 2.535 kJ. This means the total enthalpy of reaction, ΔH has a value of -2.535 kJ because the heat is being released from the reaction. However, the enthalpy changes during the dilution of nitric acid also need to be considered to obtain more accurate result.

Part II

From Graph 4,

ΔT3 = 23.0°C – 21.5°C

= 1.5 °C

ΔH3 = Ccal x ΔT3

= 0.338 kJ / °C x 1.5 °C

= 0.507 kJ

Amount of heat absorbed by the Dewar flask after adding nitric acid to the 100 cm3 of water is 0.507 kJ. This means the dilution of nitric acid has liberated heat to the surrounding.

ΔHdilution = -ΔH3

= -0.507 kJ

ΔHReaction = ΔH2 - ΔHDilution

= [-2.535 – (-0.507)] kJ

= -2.028 kJ

Therefore, the enthalpy of reaction of sodium hydroxide with nitric acid is -2.028 kJ.

Discussion

In this experiment, the calorimeter constant for the Dewar flask used has a value of 0.338 kJ / °C. In another word, this means that for every 0.338 kJ of energy absorbed by the Dewar flask, the contents of the Dewar flask will increase by 1 °C. In order to obtain an accurate result, the Dewar flask must be an isolated system. This is to ensure that all the heat released to or absorbed from the surrounding remain inside the Dewar flask but not from outside of the flask. However, ideal Dewar flask is does not exist in the world that makes sure no heat will be loss from it. This is because heat can travel through vacuum by radiation which is uncontrollable. So, the Dewar flask can only able to minimize the radiation of heat at the same time prevent conduction and convection of heat to the surrounding.

Besides, the top of the Dewar flask is open in this experiment to ease the measurement of temperature of the contents in the Dewar flask. This could cause the error in data because the system is an open system instead of an isolated system. The open system allows some of the heat loss to or absorbed from the surroundings which caused the change in temperature, ∆T to be inaccurate. During the titration, methyl orange was introduced into the solution before titration is carried out. The purpose of using methyl orange is to act as an indicator of the equivalence point of the titration. This is because it has a sharper end point and the change of colour is easily to be observed. The methyl orange is red in colour when the pH value is below 3.1 and yellow when above 4.4. The pH values which lie between 3.1 and 4.4 will show the mixture of the two colours.

From the experiment, both the reaction between sodium hydroxide with sulphuric acid and sodium hydroxide with nitric acid are exothermic reactions. This is because the temperature differences where the temperature of the mixture increased after reactions take place. The heat is released from the system to its surrounding which caused the surrounding temperature to rise. In part 2, we do not take the value of ΔH as the enthalpy of reaction between sodium hydroxide and nitric acid because dilution of nitric acid also can cause change in enthalpy. Instead, we take into consideration of enthalpy change of dilution of nitric acid to obtain a more accurate result.

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