IB Chemistry Internal Assessment 2.docx

8/10/2019 IB Chemistry Internal Assessment 2.docx

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Ahmet Ulusoy College

How Can Pure Substances Be Differentiated?

IB Chemistry HL Internal Assessment

Candidate Name: evval Beli

Session Number: 006615-006


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evval Beli006615-006

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bonded to a very electronegative oxygen atom therefore these molecules are attracted to each

other by a particularly strong type of intermolecular force called a hydrogen bond. Hydrogen

bonds are very strong forms of intermolecular attraction and cause the molecules containing

these bonds to have significantly higher boiling points than the hydrides that do not have these

bonds. (Brown & Ford, 2009)

When a substance, table salt in this case, is dissolved in water, some of the hydrogen bonds in

water and ionic bonds in table salt (NaCl) would be broken. Positive ends of water molecules

(H+) will be attracted to the negative ions in table salt (Cl-) and the negative ends (O2-) will

attract positive ions (Na+) in NaCl. These attractions are caused by ion-dipole bonds which

are even stronger than the hydrogen bonds. Therefore, with this alteration in the type of

intermolecular bonds, a mixture, salty water, will form and the melting and boiling points of

water will change.

Diagram 1:Ionic compound of NaCl dissolving in distilled water. a) Undissolved NaCl

crystal surrounded by water molecules. b) Water molecules associate with Na+ and Cl- ions.

c) Na+ions (showed in green) are attracted to hydrogen atoms in water molecules, Cl-ions

(showed in purple) are attracted to oxygen atoms in water molecules. (Halifax Regional

School Board, 2011)


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Design

Variables

Independent Variable

Concentration of distilled water and salt solutions will be altered.

Dependent Variable

Boiling point of solutions with different salt concentrations will be measured.

Controlled Variables

Same hot plate is used to heat each solution and all substances are heated at the

same temperature.

Amount of water in the erlenmeyer flask is kept constant at 100 cm3throughout

the experiment. Escape of water vapour was prevented via using a rubber stopper.

Room temperature was constant in the laboratory environment. Windows were

closed; heaters and air conditioners were not turned on. Hence, the boiling points

of the solutions can more accurately be recorded.

Measuring the Variables

Apparatus

Erlenmeyer flask, 250 cm3(10 cm3) (1)

Thermometer (0.5C) (1)

Stirring rod (1)

Right angled glass tube (1)

Rubber stopper, with two holes (1)

Watch glass (2)

Hot plate (1)


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Beaker, 100 cm3(5 cm3) (1)

Digital balance (0.01 g)

Materials and Safety Equipment

Distilled water (100 cm3)

Table salt (30 g)

Safety googles

Lab coat

Setting up Experiment

1. Clean all glassware and rinse them with distilled water.

2. Pour 100 cm3of distilled water into the 250

cm3erlenmeyer flask. Then add some small

porcelain chips into the flask to regulate

boiling.

3. Lubricate the holes of the rubber stopper

with glycerol. Connect the thermometer and

right-angled glass tube with rubber stopper,

then close the flask with rubber stopper.

4.

Put the erlenmeyer flask on the hot plate and

fix thermometer to the support rod.

Caution:The thermometer should not touch

the bottom of the erlenmeyer flask to get

more realistic temperature reading for the

solution. Picture 1: Completely set-up experiment.


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Method

1. Turn on the hot plate and heat the water until temperature rises up to 90C. Record

this temporature as initial temperature.

2. Start the timer and read the temperature every minute for ten minutes. Record the

readings.

3. Turn off the hot plate and let the water cool down to

70C.

4. Take out the rubber stopper, add 10 g of table salt into

the water and mix it with stirring rod.

5. Close the flask with rubber stopper which is connected

to thermometer and right-angled glass tube.

6. Turn on the hot plate and reheat the solution up to 90C. Record this temperature as

initial temperature.

7.

Start the timer and read the temperature every minute for ten minutes. Record the

readings.

8. Turn off the hot plate and let the water cool down to 70C.

9. Take out the rubber stopper, add 20 g of table salt into the water and mix it with

stirring rod. Repeat steps 5 to 7 and record the readings.

Caution:Do not allow water vapour to escape the erlenmeyer flask throughout the

experiment to maintain the volume of water.


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Results

Raw Data

Other than the recorded data, additional observations have been made during the experimental

process. These observations are as follows:

Thermometer in the 100 cm3of distilled water showed 100C six minutes and

twenty one seconds after the timer has been started. Temperature stayed

constant at 100C thereafter.

Salty water with 10 g salt started forming bubbles about six minutes after the

timer has been started, indicating boiling, and the temperature stabilized at

103.7C after two minutes.

Temperature increase of salty water with 30 g salt slowed down after six

minutes but has not stabilized during the experiment.

Distilled water has been observed to heat up more quickly with addition of salt.

Time (min)(1 s) 0 1 2 3 4 5 6 7 8 9 10

Tem

perature

(0.5

C)

Distilled

water

90 91.4 93 94.6 97.2 98.7 99.3 100 100 100 100

Salty

Water 10 g

salt

90 92.1 94.3 95.8 98 100.6 102.5 103.3 103.6 103.7 103.7

30 gsalt

90 93.6 96.4 97.9 100.2 103.2 106.1 107.4 108 108.3 108.5

Table 1:Temperature readings of distilled water and solutions with different salt

concentrations.


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Processed Data

Temperature of 100 cm3distilled water increased gradually until it reached 100C around

seven minutes, suggesting that water started to boil at this temperature. To be able to visualize

the overall trend in the temperature change of distilled water during experiment, data points

have been plotted on graph:

Graph 1: Temperature readings of 100 cm3 distilled water taken for ten minutes.

Even though the trend of Graph 1 is a gradual increase, stabilization is observed in data points

following the point (7,100). Temperature (y-axis) does not show any variation after that point,

suggesting that the line of best-fit should show a similar pattern. Logarithmic curve would

represent the stabilization over time accurately; however it would be insufficient to show the

difference in temperature between minutes 0 and 1. Thus, inverse exponential curve is used to


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show the overall trend of temperature readings of 100 cm3distilled water. RMSE (Root Mean

Square Error) value of the curve in Graph 1 is fairly low, 0.7478, indicating a good fit.

As water is heated, the particles move faster and so the temperature increases. Some

molecules will have enough energy to break away from the surface of the liquid so some

water evaporates. At the boiling point of water, there is sufficient energy to break all the

intermolecular bonds. The added energy is used for this process, not to increase the kinetic

energy, and so the temperature remains constant. (Brown & Ford, 2009)

Graph 2:Heating curve of water. More energy is required during evaporation than melting

as all intermolecular bonds are broken during this process. (The University of Texas, 2013)

As seen in Graph 2, temperature of water stops increasing while changing states. It is very

well known that water boils at 100C under standard conditions. 100 cm3distilled water in the

experiment started boiling in between sixth and seventh minutes, and its temperature

remained at 100C beyond that point. Hence, it can be concluded that the boiling point of

distilled water has been successfully observed.


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To be able to compare differences between temperature change of distilled water and salty

water more efficiently, data points for 10 and 30 g salty water solutions in Table 1 have been

plotted on separate graphs, using best-fit inverse exponential curve again.

Graph 3: Temperature change of 10 g salt and 100 cm3distilled water solution.

In Graph 3, the curve seems steeper than in Graph 1, thus it can be concluded that the boiling

point of salty water with 10 g table salt is higher than the boiling point of distilled water.

RMSE value of the curve in Graph 3 is 0.8439, which is a bit higher than the value in Graph 1

but not high enough to consider the curve as unfit. To be able to give reason to the difference

between the boiling points of saline solutions and distilled water, the concentrations of salty

water solutions have been calculated:

Molarity (M) =

n(NaCl) (10 g) =

= 0.17 mol


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Molarity=

= 1.7 moldm-3

% concentation=

100 = %10

As seen in Graph 3, temperature increase stabilized about minute 8 and was constant at

103.7C for nearly 3 minutes at the end of the experiment. Thus, it can be concluded that %10

saline water solutions boiling point is 3.7C higher than distilled waters.Comparing

literature values (Lide, 2005) with the results obtained in this experiment:

29.2 g of salt in 1 kg water is known to increase the boiling point by 0.5C.

1 kg=1 L, hence 2.92 g of salt in 100 cm3water raises the boiling point by 0.5C.

For salty water with 10 g salt:

=

x = 1.71C

The boiling point of the salty water solution with 10 g salt in the experiment was

approximately 2 C higher than the literature value. This situation is most probably caused by

random and systematic errors and will be investigated in evaluation.


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Graph 4: Temperature change of 30 g salt and 100 cm3distilled water solution.

The best-fit inverse exponential curve in Graph 4 seems to be even steeper that the curve in

Graph 3. RMSE value has also increased to 0.9016, but again, this is not a significant change.

Another point that differed Graph 4 from graphs 1 and 3 was that there was no discrete

stabilization in this graph. This was an unexpected result and could be linked to time

limitations of the experiment. Even though there was not a stable temperature recorded to

consider as the boiling point, temperature reading in minute 10 could be used to approximate

the difference between graphs.

n(NaCl) (30 g) =

= 0.51 mol

Molarity=

= 5.1 moldm-3

% concentation=

100 = %30


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Therefore, it could be said that %30 saline solutionsboiling point is approximately 8.5C

higher than distilled waters.Comparing literature values with the results obtained in this

experiment:

For salty water with 30 g salt:

=

x = 5.14C

The boiling point of the salty water solution with 30 g salt in the experiment was

approximately 3.36C higher than the literature value. Seeing a proportionate increase

between the values obtained for the solutions with 10 g and 30 g salt, the presence of

systematic error(s) can be assumed.


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Thedifferenceintheoveralltrendofgraphsof

solutionswithd

ifferentconcentrationscanbe

observedwhenassembledonthesamepage.

Curv

es

inthegraphsget

steeperasmoresaltisaddedto

distilledwater.T

hisshowsthattheboilingpoint

increasesproportionatelywiththeconcentrationof

tablesaltinwater.


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Discussion

Concentration of mixtures that have been observed in this experiment showed differences in

the boiling points. The 30% salty water solution increased the boiling point of water more

than the 10% salty water. Pure substances have literature boiling points dependent on their

intermolecular forces which can be used to differentiate them from mixtures. Referring back

to background information, it could be said that more hydrogen bonds have been disrupted

and replaced by ion-dipole bonds. Hence, this experiment suggests that an increase in the

ratio would raise the boiling point proportionately.

Conclusion

Considering all data given and assessed above, an answer to the research question of this

experiment could be given. All pure substances have certain boiling points and so

experimenters can use these points to identify an unknown liquid or to test the purity of items.

Boiling point is directly related to the intermolecular attractions within a substance.

Interference of other substances disrupts the bonds between molecules of pure substances,

causing variations in the boiling points. Concentration of these additional substances has also

been observed to alter boiling points. Even though the boiling points of mixtures vary, the rate

of temperature increase stabilizes for all substances after a while.


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Evaluation

Several problems have been encountered due to limitations of equipment and method of the

experiment. Thermometer was only able to measure up to 110C, which was a temperature

nearly reached with 30 g salty water solution. One of the holes of the rubber stopper used to

close the erlenmeyer flask was too wide, so a piece of plastic had to be stuck next to the

thermometer in the hole. Some water vapour escaped when rubber stopper has been taken out

to add table salt in the erlenmeyer flask, and even though a slight amount has escaped, this

situation has caused the volume of distilled water used to decrease, affecting the result of the

experiment.

Unexpectedly, no stabilization has been observed in the graph of 30% saline water (Graph 4).

This situation has been caused by the limitation in the duration of the experiment. If the time

measured after the solutions reached 90C was longer, stabilization over time could more

easily be observed in graphs 1, 3 and 4. Also, even though the temperature and pressure of the

room was kept constant, the experiment was performed in an open environment; hence

random errors will be attained no matter the number times the experiment was performed.

Even though the best-fit curve found for the graphs, the inverse exponential curve, has

consistently reflected the trend and the RMSE values were fairly low; there were some data

points which significantly deviated from the curve. These variations are most probably caused

by random errors and thus, do not affect the curvesoverall fit to the graph.

Deviations from the literature values can be caused by many factors. Table salt used in this

experiment was for commercial use; hence its purity may differ from other brands and NaCl

for experimental use.


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Improvements

Factor Possible Effect on the

Experiment

Improvement

Rubber stopper with

different sized holes

Caused water vapour to escape

from the flask, leading to a slight

decrease in the boiling point.

Using a rubber stopper

with properly sized holes

to be able to fit the

thermometer.

Thermometers limit was

nearly reached

Thermometer could have been

damaged, not only constraining

experimenter to redo the

experiment, but also causing

leakage of poisonous mercury to

the environment.

Using a thermometer with

wider range in order to

observe higher boiling

points.

Expected stabilization

could not have been

observed in Graph 4 due

to limits in the duration of

the experiment

Boiling point of distilled water

with 30 g salt could not have been

observed properly, leading to

difficulties in answering the

research question.

Extending the experiment

to 15-20 minutes in order

to observe the overall trend

on the graphs more

accurately.

Possible inaccurate

observations of the

experimenter

Cause the experimenter to enter

inaccurate data in the data table,

consequently increasing the

percentage error for the processed

data.

Making more than one

experimenter measure the

temperature and/or using

an ebulliometer.

Heat loss to the

environment

Temperature readings in each

minute will show slight

differences when the experiment

is repeated because some amount

of heat from the hot plate will

escape to the air.

Repeating experiment in a

closed, insulated system.


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Works Cited

Brown, C., & Ford, M. (2009). Higher Level Chemistry.Malaysia: Pearson Education Limited.

Halifax Regional School Board. (2011, March). Halifax Regional School Board Teacher Webspace.Retrieved October 01, 2013, from

http://hrsbstaff.ednet.ns.ca/benoitn/adv%20chem%2011/homework/2011/chem11_hw_march2011

.htm.

Lide, D. R. (2005). CRC Handbook of Chemistry, Internet Version.Boca Raton, FL, USA: CRC Press.

The University of Texas. (2013). Heating Curves (revisited). Retrieved September 23, 2013, from

http://ch302.cm.utexas.edu/physEQ/physical/selector.php?name=heat-curves-revisit.

Warning signs were retrieved from:

http://www.seton.net.au/signs-labels/labels/safety-labels/mandatory-pictos/international-labels-lab-

coat-picto-s9334.html

http://www.jactone.com/health-safety-signs/personal-protection/eye-protection/wear-goggles-

symbol.html

http://www.gsbhealthandsafetysigns.co.uk/caution-hot-water-sign.html on October 4, 2013
http://www.seton.net.au/signs-labels/labels/safety-labels/mandatory-pictos/international-labels-lab-coat-picto-s9334.htmlhttp://www.seton.net.au/signs-labels/labels/safety-labels/mandatory-pictos/international-labels-lab-coat-picto-s9334.htmlhttp://www.jactone.com/health-safety-signs/personal-protection/eye-protection/wear-goggles-symbol.htmlhttp://www.jactone.com/health-safety-signs/personal-protection/eye-protection/wear-goggles-symbol.htmlhttp://www.gsbhealthandsafetysigns.co.uk/caution-hot-water-sign.htmlhttp://www.gsbhealthandsafetysigns.co.uk/caution-hot-water-sign.htmlhttp://www.jactone.com/health-safety-signs/personal-protection/eye-protection/wear-goggles-symbol.htmlhttp://www.jactone.com/health-safety-signs/personal-protection/eye-protection/wear-goggles-symbol.htmlhttp://www.seton.net.au/signs-labels/labels/safety-labels/mandatory-pictos/international-labels-lab-coat-picto-s9334.htmlhttp://www.seton.net.au/signs-labels/labels/safety-labels/mandatory-pictos/international-labels-lab-coat-picto-s9334.html

IB Chemistry Internal Assessment 2.docx

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