chnage of standard gibbs free energy for the decomposition of sodium hydrogen carbonate

EXPERIMENT 1

:

CHANGE OF STANDARD GIBBS FREE ENERGY FOR THE DECOMPOSITION OF SODIUM HYDROGEN CARBONATE

OBJECTIVE To determine the change of standard Gibbs free energy for the decomposition of sodium hydrogen carbonate from the change of standard enthalpy and the change of standard entropy. INTRODUCTION Gibbs free energy combines enthalpy and entropy into a single value. The change of free energy is equal to the sum of its enthalpy plus the product of the temperature and entropy of the system. ΔG can also predict the direction of the chemical reaction. If ΔG is positive then the reaction is nonspontaneous. If it is negative, then it is spontaneous. In 1875, Josiah Gibbs introduced a thermodynamic quantity combining enthalpy and entropy into a single value called Gibbs free energy. This quantity can be defined as:

G=H−TS G=U+PV−TS where

     

U = internal energy (SI unit: joule) P = pressure (SI unit: pascal) V = volume (SI unit: m3) T = temperature (SI unit: kelvin) S = entropy (SI unit: joule/kelvin) H = enthalpy (SI unit: joule) Gibbs Free Energy (G) - The energy associated with a chemical reaction that can be used to do work. The free energy of a system is the sum of its enthalpy (H) plus the product of the temperature (Kelvin) and the entropy (S) of the system:

G=H−TS Free energy of reaction (ΔG) In chemical reactions involving the changes in thermodynamic quantities we often use another variation of this equation: [1]

The sign of delta G can allow us to predict the direction of a chemical reaction:

Terminology:

 

If ΔH

< 0 and ΔS

> 0 then the reaction will be SPONTANEOUS (ΔG

If ΔH > 0 and entropy ΔS temperature.

Standard-state free energy of reaction (



< 0 ) at any temperature.

< 0 then the reaction will be NON SPONTANEOUS (ΔG

> 0 ) at any

G )

The free energy of reaction at standard state conditions:

By determining the quantity of experimentally according to Hess’s Law and extraction of from the standard data, of the reaction can be estimated easily. In our experiment, a double vacuum stainless steel calorimeter is used to measure the heat released or absorbed by decomposition of NaHCO3 . the thermo chemical equation for the decomposition of NaHCO3 (s) is

( )

( )

()

( )

However in the laboratory, the quantity of of NaHCO3 (s) cannot be determined directly. Instead, two separate experiments in determining the quantities for for the reaction of NaHCO3 (s) and estimation of

(aq) are carried out. Hess’s law is applied in

( ) , respectively with of NaHCO3 (s) .

The reaction are as follow :

( ) ( )

( ) ( )

( ) ( )

( ) ( )

() ()

A simple double-wall vacuum steel calorimeter is used for determining the quantity of . It may serve as the constant-pressure calorimeter (at atmospheric pressure under the experimental conditions). The enthalpy (H) is defined as H = U + PV

A simple double-wall vacuum steel calorimeter is used for determining the quantity of

. It may

serve as the constant-pressure calorimeter (at atmospheric pressure under the experimental conditions). The enthalpy (H) is defined as H = U + PV A change in enthalpy is equal to the heat supplied at constant pressure to a system in the case where the system does no additional work. dH = dq For measurable change,

A known amount of NaHCO3 (s) or Na2CO3 (s) is added to an excess of H2SO4 (aq) and the change in temperature ( ) is measured. The heat released or absorbed by each reaction ( ) is calculated by using the formula :

Where m denotes mass of the solution (in unit g) in the calorimeter by assuming no change in the volume of the solution and the density of the solution is 1 g mL -1 and (specific heat capacity of the solution at constant pressure) = 4.18 J g -1 -1. Express quantity in unit kJ. The quantity for the double-wall vacuum stainless steel calorimeter is small and any heat absorbed is negligible. Quantity is negative for exothermic reaction (increase in temperature during the experiment) and is positive for endothermic reaction (decrease in temperature during experiment). CHEMICALS Sodium hydrogencarbonate, NaHCO3 Sodium carbonate, Na2CO3 1 M sulphuric acid, H2SO4 APPARATUS 50 mL clear glass volumetric pipette 250 mL double-wall vacuum stainless steel calorimeter with double-wall cover General purpose and mercury filled thermometer 0 – 100 , gradually every 1 100 mL beaker Stopwatch

PROCEDURE 1. 4.0 to 4.5 g of NaHCO3 was measured exactly. 2. 50 mL of 1 M H2SO4 was transferred into the double-wall vacuum stainless steel calorimeter using a volumetric pipette. The cover of the double wall and the thermometer were replaced. 3. The temperature of H2SO4 was recorded for every one minute for 4 minutes. At the 5 th, the NaHCO3 was quickly transferred into the H2SO4 . 4. The cover was replaced and the content in the calorimeter was being stirred carefully using thermometer. 5. The temperature was recorded for the next 4 minutes every 10 seconds. 6. Step 1 to 5 was repeated with 3.0 to 3.5 g of Na2CO3 RESULT Mass of NaHCO3 Mass of Na2CO3 Room temperature

: : :

4.4309 g 3.2432 g 28.5

NaHCO3 with H2SO4 T( ) Time (s) 60 30 120 30 180 30 240 30 310 25 320 25 330 25 340 25 350 25 360 25 370 25 380 25 390 24 400 24 410 24 420 24 430 24 440 24 450 24 460 24 470 24 480 24 490 24 500 24 510 24 520 24 530 24 540 24

Na2CO3 with H2SO4 T( ) Time (s) 60 29 120 29 180 29 240 29 310 30 320 30 330 30 340 30 350 30 360 30 370 30 380 30 390 30 400 30 410 30 420 30 430 30 440 30 450 31 460 31 470 31 480 31 490 31 500 31 510 31 520 31 530 31 540 31

CALCULATION Heat change of NaHCO3 with H2SO4

(

)

Na2CO3 with H2SO4

(

( ) ( )

( ) ( )

( ) ( )

⁄

(

(

(

( )

(

)

( )

)

(

()

( ) ( )

⁄

)

() ()

)

)

( )

( )

( )

)

()

(

( )

)

()

( )

∑

( )

∑

( )

()

(

)

( ) (

)

Heat required for decomposition of one moles NaHCO3 or 42.85kJ/mole NaHCO3 and the process is endothermic.

∑

( ) (

∑

( )

() )

( ) (

)

∑

∑

(

)

(

(

)

)

DISCUSSION Recalling the condition for spontaneous change ΔG = ΔH – TΔS < 0 It is apparent that the temperature dependence of ΔG depends almost entirely on the entropy change associated with the process. We say "almost" because the values of ΔH and ΔS are themselves slightly temperature dependent; both gradually increase with temperature. In the equation above, the sign of the entropy change determines whether the reaction becomes more or less spontaneous as the temperature is raised. For any given reaction, the sign of ΔH can also be positive or negative. positive value while the

Presented a

gives out a negative value. Apparently in our experiment, the ΔH is

less than 0 while ΔS is greater than 0. Under these conditions, both the ΔH and TΔS terms will be negative, so ΔG will be negative regardless of the temperature. An exothermic reaction whose entropy increases will be spontaneous at all temperatures. In order to make use of Gibbs energies to predict chemical changes, we need to know the free energies of the individual components of the reaction. For this purpose we had combine the standard enthalpy of formation and the standard entropy of a substance to get its standard free energy of formation

We then determine the standard Gibbs energy of the reaction according to

∑

∑

As for the sources of error in this experiment, we believe that the major reason was heat lost as we added sodium hydrogen carbonate into the H2SO4 in the double-wall vacuum steel calorimeter. The heat lost here relatively alter the value of

q. Besides, we also struggle to literally stir the content

of the calorimeter while having to take the reading of the thermometer each 10 seconds which might cause small faults in the temperature reading. CONCLUSION

The change of standard Gibbs free energy for the decomposition of sodium hydrogen carbonate from the change of standard enthalpy and the change of standard entropy is obtained. The value is meanwhile the theoretical value is found to be The experiment would be lot more accurate and successful if the major error occurring is avoided.