Enthalpy Of Solution Lab Report Assessed on: * Data Collection and Processing (DCP) * Conclusion and Evaluation (CE)

Usman Omid Chemistry IB 1A

Enthalpy Of Solution

_Introduction_

AIM

The aim of this investigation was to determine the enthalpy of the solution created when dissolving Sodium Hydroxide (NaOH) in water (H2O).

THEORY

When a solid ionic compound is dissolved in water a change of enthalpy is involved. _The enthalpy of a solution_, ΔHsol is the enthalpy change when one mole of the substance is dissolved in water to form aqueous solution. �

CHEMICAL

Sodium Hydroxide (NaOH)

PROCEDURE

The recommended amount of compound ito use is around 0.1 moles.

Add water into your caleriometer

Add the solid salt and stir the mixture thoroughly and continuously.

Calculate ΔHsol. Ignore the heat capacity of the caleriometer.

_Results_

RAW DATA

Tab.1. The different data required for the calculations.

Trial

Mass of NaOH / g (± 0.001 g)

Volume of H2O / ml (± 1 ml)

Initial Temperature / °C (± 1 °C)

Final Temperature / °C (± 1 °C)

1

3.954

50

19

39

2

4.027

70

19

33

3

3.991

80

19

31

4

3.988

100

19

28

One qualitative data that might have affected the results is that the Sodium Hydroxide (NaOH) seemed to start a slow melting process as soon as it left its container. This was a qualified guess based on that it started to sparkle and looked like it was melting.

Another one that probably did affect the results was the fact that the containers in which we mixed the water (H2O) and Sodium Hydroxide (NaOH) in had holes. And this increases the energy lost to the surroundings in the form of heat, which affects the temperature. �

PROCESSED DATA

Calculations

Theory states:

ΔH= -Q/n

Q = m * c * ΔT �

Where in this case m = mass of water, c = specific heat capacity of water and T in Kelvin.

In order to know whether mass should be in kg or g we look at the c-value.

c= 4.18 JK-1g-1.

And Q is measured in Joules.

J = g * JK-1g-1 * K.

m must be in grams since J=J, and not J = kJ like it would have been with kg.

We can assume that the density of water is 1000g / dm3. �

NaOH has got a molar mass of (MNa + MO + MH) 23 + 16 + 1 = 40 g / mole.

The Procedure recommended the use of 0.1 moles of NaOH.

n = m/M

m = M * n.

m = 0.1 mole * 40 g / mole = 4 g of NaOH.

The temperature is to be in Kelvin and not Celsius Degrees.

K = °C + 273. �

The calculations of the different uncertainties when multiplication or division is needed will be calculated by using the percentage uncertainty of the quantity and then multiplying the percentage with the value after the conversion so that the percentage uncertainty is the same.

When addition or subtraction is needed, the uncertainties will just be added.

In converting from Celsius degrees to Kelvin we do not need to change the uncertainty since the scale is the same.

The uncertainty of the number of moles NaOH that was used is so small it is negligible

Example: In trial one the initial temperature is 19 °C ± 1 °C.

The uncertainty is 1.

19 °C = 292K

The uncertainty = ±1K

Temprature = 292K ± 1 K

To find the uncertainty in the enthalpy change, the uncertainties of the other values must be considered. Since in the calculation of the enthalpy change only multiply or divide, the different percentage uncertainties are required. To give an example, the calculations on the uncertainty of the enthalpy change of trial one is given. The uncertainty of the number of moles is so small it is almost negligible, which is the reason it will not be presented in the table 2. But just for the record, it will be present in the following calculation.

The percentage uncertainties for m, c and Δ T are required since Δ Q = m* c * Δ T

Percentage uncertainty of number of moles of NaOH = (0.001/4) * 100 = (2.5 * 10-4) *100 =

(2.5 * 10-2)%.

Percentage uncertainty of mass of H2O = (1/50) * 100 = 0.02 * 100 = 2%.

Percentage uncertainty of Δ Temperature =( 2/20) * 100 = 10%.

Then the percentage uncertainties are added.

10% + 2% + 0.025% = 12.025%.

Then the uncertainty for ΔQ = 0.12025 * 4180 = ± 502.645 J.

And the uncertainty for ΔH = 502.645J / -0.1 = ± 5026.45 J/ moles.

Tab.2. The Enthalpy of the solution and the data required to calculate it.

Trial

Number of NaOH / Moles

Mass of H2O / (g ± 1 g)

Δ Temperature / K ± 2K

Heat Energy / J

Enthalpy / J /mole

1

0.1

050

20

4180

-41800

2

0.1

070

14

4100

-41000

3

0.1

080

12

4010

-40100

4

0.1

100

09

3760

-37600

In order to compare the results of the different trials only the Enthalpy is needed to be observed since all the other data is used to get that value. As the table indicates, the calculated enthalpy values are very similar except in trial 4.

Graph.1. How the mass of H2O affects the Δ Temperature when dissolving NaOH in H2O and using a constant mass of NaOH.

As the mass increases, the Δ Temperature decreases.

The mean value of the Enthalpy of the solution values is ((-41800 + -41000 + -40100 + -37600)/4 = ) -40100 J / mole. The mixture between NaOH and H2O is an exothermic reaction since it gives away energy in the form of heat and increases the temperature.

The means the random uncertainty is

(12.025% + 15.739% + 17.942% + 23.247%) / 4 = 17. 238%.

So the mean enthalpy of solution value for when NaOH reacts with H2O is:

-40 kJ ± 7 kJ.

_Conclusion and Evaluation_

So when NaOH reacts with H2O, the enthalpy of the solution is -40 kJ ± 7 kJ.

Since the enthalpy value is negative, it means that energy is lost, probably due to heat to the surrounding, which in this case was the water. Since it gives away energy, it is an exothermic reaction.

The value of -40 kJ ± 7 kJ can be supported by the literary values.

The literary values of the enthalpy in a NaOH and H2O is -44.4 kJ/mol. �

The experimental value was -40 kJ ± 7 kJ.

It seems as the experimental value is very good since the literary values is within the range of the experimental values uncertainty.

ERRORS

Total percentage error =

Total percentage error= ((-44.4 - -40)/ -44.4) * 100% = 9.9%.

Systematic error = Total Percentage error - Random Error.

Systematic error = 9.9 - 17.2 = -7.3%.

The systematic errors are smaller than the random errors in this experiment. This is probably because the apparatus were very simple and there were many conversions of units in the calculations and also because three different measured quantities were used to calculate the ΔQ. And when you add and multiply and divide the values the uncertainties stack on top of each other and grow. And there were many calculations and measurements in this experiment which is why the random errors were so big. The systematic errors are in most cases not very big since after some lab experience one does not make simple mistakes.

EVALUATING PROCEDURE AND APPARATUS

The procedure had a lot of weaknesses and limitations. To be able to see the limitations and weaknesses easier a list would be necessary.

We used a Styrofoam cup with a hole in the lid. This has a lot of heat loss to the surroundings since heat is easily radiated to the surrounding air. So the temperature was probably greatly affected because of this. (very significant since A LOT heat is lost to the surroundings.

We had to stir the salt in the water. This generates heat. This affects the results. (not very significant since it should not generate very much heat.)

The NaOH was in crystal form and not a fine salt. This leads to that it has less surface area and so the reaction goes slower. When the reaction goes slower it takes more time to carry out the experiment and more heat can radiate to the surrounding air and the solution has to be stirred for a longer time. (quite significant since the reaction time is important to minimize the heat loss to the surroundings)

The specific heat capacity of the water might alter when the salt is added to it. Then the c-value in the formula will change (not very significant since it will only alter it very little).

When stirring the solution and trying to keep the lid on most people failed. The lid moved and at some points fell off. This lead to an even greater heat loss to the surroundings.

Trying to add the salt into the water while holding the thermometer and stirring at the same time was hard. This lead to a lot of pauses and in some cases salt was dropped on the table as well affecting the results.

Although there are some weaknesses and one is a very significant one, the experimental value compared to the literature value was not very different. Also, the weaknesses might cancel out a little since stirring the solution generates heat and the poor apparatus makes the solution loose heat to the surrounding air. So the procedure could give a result with good quality, like in this case. And the weigh showed up to 3 decimal numbers, which lead to that the percentage uncertainty was almost negligible. So the apparatus were precise and accurate up to some level. The thermometer on the other hand showed no decimal numbers (it was not a digital electronic one). This was the main reason to the very high random errors. This also lead to that the experiment results were not very precise since the random errors were 17.2%! This is not very accurate. But since the results were accurate there can not have been very many systematic errors or other errors since the experimental value agreed a lot with the literary value.

IMPROVEMENTS

Since it had weaknesses which were mainly due to the poor apparatus, there is room for easy improvements in most cases. A list will be easier to read.

If we had used a caleriometer the heat loss would not have been near the heat loss that was present in this experiment.

Taping the lid on top of the cup before stirring and then maybe even taping the thermometer (which we stirred the solution with) through the hole and cover the hole with tape would have helped a lot. One would still be able to stir the solution.

If we had started with an initial temperature higher than 19 Celsius degrees the salt crystals would have dissolved by itself and generating heat by stirring would not have been a problem. But in that case a box that would have a temperature the same as the solutions initial temperature would have been needed since otherwise a lot of heat from the hot water would have been lost to the surrounding.

If the individuals that carried out the experiment had worked two and two the problem with doing many things at the same time would not have been a problem.

The systematic error was not very high at all but could have been reduced if we had done many more trials. If we had been given more time then there would have been time for more trials and the mean value ought to have less systematic errors then.

If we had been in a proper and better lab room then we would have had a greater access to precise and accurate apparatus which would have lead to less random errors. If we had used another method, such as using hydrochloric acid (HCl) instead of water or maybe water with higher initial temperature then the reaction time problem would have been reduced.

� Cited from Chemistry handout "ENTHALPY OF SOLUTION" received from the teacher 19/01 2009

� The Lanly Company. Updated 06/19 2007. The Physics Of Heat Processing. Lanly. Retrieved February 04, 2009, from http://www.lanly.com/heating.htm

� Chemistry 3rd Edition by John Green and Sadru Damji. IBO 2007. ISBN 978-1-876659-08-0. Page 138.

� 1998-2008 Roger Walker. Information retrieved 04/01 -2009. Page last modified : 3 rd. September 2008 .From �HYPERLINK "http://www.simetric.co.uk/si_water.htm"�http://www.simetric.co.uk/si_water.htm�

� �HYPERLINK "http://ostermiller.org/calc/temperature.html"�http://ostermiller.org/calc/temperature.html� Copyright Stephen Ostemiller 2001-2006. Page visited 04/02 2009.

� �HYPERLINK "http://www.mindspring.com/~drwolfe/WWWolfe_dat_enthalpies.htm"�http://www.mindspring.com/~drwolfe/WWWolfe_dat_enthalpies.htm�

Page visited 05/02-2009. WWWolfe Enthalpies.

�PAGE � �PAGE �- 1 -� Katedralskolan � TIME @ "d/M yyyy" �8/3 2009�

Cited: from Chemistry handout "ENTHALPY OF SOLUTION" received from the teacher 19/01 2009 � The Lanly Company. Updated 06/19 2007. The Physics Of Heat Processing. Lanly. Retrieved February 04, 2009, from http://www.lanly.com/heating.htm � Chemistry 3rd Edition by John Green and Sadru Damji. IBO 2007. ISBN 978-1-876659-08-0. Page 138. � 1998-2008 Roger Walker. Information retrieved 04/01 -2009. Page last modified : 3 rd. September 2008 .From �HYPERLINK "http://www.simetric.co.uk/si_water.htm"�http://www.simetric.co.uk/si_water.htm� � �HYPERLINK "http://ostermiller.org/calc/temperature.html"�http://ostermiller.org/calc/temperature.html� Copyright Stephen Ostemiller 2001-2006. Page visited 04/02 2009. � �HYPERLINK "http://www.mindspring.com/~drwolfe/WWWolfe_dat_enthalpies.htm"�http://www.mindspring.com/~drwolfe/WWWolfe_dat_enthalpies.htm� Page visited 05/02-2009. WWWolfe Enthalpies. �PAGE � �PAGE �- 1 -� Katedralskolan � TIME @ "d/M yyyy" �8/3 2009�