WIS Chemistry Department

NAME ………………………………………………………………………………………….

TEACHER ……………………………………………………………………………………..

If found please return to West Island School Chemistry Department

WELL DONE – YOU’VE CHOSEN CHEMISTRY AS AN IB SUBJECT!

Now that you have shown extremely good taste in selecting Chemistry for the next two years, there is an introductory phase to the course that you will have to go through. A sort of Boot Camp, if you will.

Please remember – material that was covered in the Chemistry part of your IGCSE course is ASSUMED to be understood already. Important concepts that should have remained in your cranium through the long summer months include:

Basic periodic table arrangement Calculation of formula masses Calculation of moles based on masses of solids Writing and balancing chemical equations Calculation of the concentrations of solutions Expressing answers in standard form (from Mathematics) Significant figures (from Mathematics)

All of these will be reviewed during the course, but only very briefly. It’s up to you to ensure that you feel confident with all of the above i.e. through the novel concept of SELF STUDY (get used to it – not just in Chemistry but all your IB subjects).

A basic outline for the first 10 lessons of your Chemistry course (SL or HL) is as follows:

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Lessons 1 & 2:

Practical X - An investigation into acid/base reactions and indicators.

This will serve as an introduction to one of the most important techniques – titrations - used in Years 12 and 13 which will help you to prepare for the practical assessments you will face. This will also give your teacher an opportunity to assess your current level of confidence and ability when doing experiments (don’t worry, it doesn’t count for anything – yet). -

Lessons 3 & 4

Theory – Measurements and Uncertainties

This is actually one of the topics included in the IB course (topic 11). It is covered at this early stage as understanding of error analysis, accuracy, precision etc. is essential to perform any experiment at IB level to an acceptable standard.

Lessons 5 & 6

Practical Y – Identifying the formula of a compound

This activity allows you to find the formula of a compound from your own experimental data – a technique that you may recall from year 10. It also prompts you to think about the errors in your measurements and the concepts of precision and accuracy. This experiment will really test your practical skills as a high standard of accuracy is required during the heating process.

Lessons 7 & 8

Theory – Criteria used for internal assessment of experimental work

Internal assessment of your experimental work will make up 24% (that’s almost one quarter for those a bit slow on the uptake) of your final IB level. It is vital that you understand how your work will be assessed in order for you to gain as high a mark as possible. You will need your copy of the Chemistry Guide for these lessons.

Lessons 9 & 10

Practical Z – An investigation into solubility

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A Year 12 look at this relatively straightforward GCSE concept. You will need to expand your use of terminology to an IB standard and develop further practical skills. This will provide an opportunity to peer mark your work as per the IB criteria.

The Appendices

In addition, there are several appendices of relevant information that will be of use as you proceed with the experiments during your course. Much of the content of the appendices is common sense, but this appears to be lacking in the average Year 12 student so a refresher is appropriate. You may work through these appendices with your teacher in class, and refer to them during later experimental work.

Why study Chemistry? Well, take a look at the following Q & A from a well-respected researcher with a doctorate in Chemistry:

Question: Is there really a chemistry of love?

Answer: I don't think there are any magic love potions that you can use to make someone fall in love, but chemistry does play an important role in how a relationship progresses. First, there's attraction. Nonverbal communication plays a big part in initial attraction and some of this communication may involve pheromones, a form of chemical communication. Did you know that lust is characterized by high levels of testosterone? The sweaty palms and pounding heart of infatuation are caused by higher than normal levels of norepinepherine. Meanwhile, the 'high' of being in love is due to a rush of phenylethylamine and dopamine. All is not lost once the honeymoon is over. Lasting love confers chemical benefits in the form of stabilized production of serotonin and oxytocin. Can infidelity be blamed on chemistry? Perhaps, in part. Researchers have found that suppression of vasopressin can cause males (voles, anyway) to abandon their love nest and seek new mates. Hey, you gotta have chemistry! (Anne Marie Helmenstine, PhD)

That’s right – love easily explained as a big bag of chemicals. What did Physics ever give us? Gravity – woo hoo.

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Practical X - An investigation into acid / base reactions and indicators.

In this investigation we find out why certain indicators work well for titrations in some instances but not in others. You will titrate an acid with a soluble base (alkali) and use an indicator to determine the equivalence point – when the acid has JUST been neutralised. The two indicators most commonly used are methyl orange and phenolphthalein. First we find out why universal indicator is not a good indicator for titration work.

Indicator colour in acid colour in alkaliMethyl orange red orange Phenolphthalein colourless pink

Instructions:

Part A

1. Rinse a clean burette with 10 cm3 of distilled water then with 10 cm3of the solution of sodium hydroxide. Make sure the section of the burette below the tap is rinsed then emptied into a labeled waste beaker.

2. Rinse a 25 cm3 pipette in the same way to prepare for use with the hydrochloric acid solution.

3. Fill the burette with the sodium hydroxide solution and zero the liquid level, making sure that the tap area is full and that there are no bubbles which would compromise the accuracy of your work.

4. Pipette 25 cm3of the hydrochloric acid into a clean conical flask.

(We do not rinse the flask with the acid. Why not?)

_________________________________________________________________________________________________________________________________________________________________________________

5. Add 2-3 drops of universal indicator to the flask and swirl. Make sure that the colour is visible but not intense. Add a little more indicator if the colour cannot be seen.

6. Add the alkali from the burette into the acid in the flask a little at a time and swirl after each addition. You should add about 1 cm3 at a time at the beginning and systematically reduce the volume added as the equivalence point is approached. We should never add a large volume of one reactant to another without mixing thoroughly as in some cases this can result in inaccuracy.

7. Continue until you think that the equivalence point has been reached. Record the volume of sodium hydroxide used in a result table. How

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confident can you be of an accurate result using this indicator? Explain why.

8. Repeat the titration but using methyl orange as your indicator. Record the amount of sodium hydroxide needed to neutralize 25 cm3 of the hydrochloric acid and comment on the likely accuracy of your result.

9. Finally repeat once again but using phenolphthalein as your indicator. Determine the difference in volume between your 3 results and comment on which indicator(s) are acceptable for this acid/base pair.

Part B

1. Rinse your pipette with distilled water and then with the ethanoic acid solution and pipette 25cm3 of this into a clean conical flask.

2. Add methyl orange as indicator and then titrate to an end point.

Comment on your result and the confidence you have in your value.

3. Repeat the operation using phenolphthalein as indicator and compare your results. Can both indicators be confidently used for this acid base pair?

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Results table

Conclusions drawn: _____________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________

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Theory – Measurements and Uncertainties

If a reported measurement is to be useful, it must include some indication of its uncertainty. The complete expression of a measured quantity must include the number value and the unit and also show how reliable the number is. This section discusses how measurements are reported and used so that the reliability of the numbers is understood by everyone.

Accuracy and Precision

The quality of a measurement depends on the measuring instrument and the skill of the person making the measurement. Uncertainty in measurements can result from limitations in accuracy or limitations in precision. The limitations may be those of the instrument or the experimenter, or both.

Accuracy refers to the closeness of a measurement to the true or accepted value of the quantity measured. An electronic balance that is often calibrated with a standard mass is likely to be more accurate than an old, mechanical balance that has been dropped many times.

Precision refers to the agreement among the numerical values of a set of measurements of the same quantity made in the same way. A chemist who frequently carries out a complex experiment is likely to have more precise results than someone just learning the experiment.

Suppose it is necessary to determine the density of a sample of chloroform (CHCI3) at 20oC. One chemist obtained values of 1.495g/ml, 1.476g/ml, and 1.485g/ml for the density of the sample. These measurements vary widely; the precision is poor. Another chemist repeated the measurements three times using the same equipment as the first chemist and found densities of 1.487g/ml, 1.490g/ml, and 1.488g/ml. This group of measurements has good precision. Comparing these values with the accepted value of 1.489g/ml for chloroform at 20oC shows that this group of measurements also has good accuracy.

If a measurement can be compared to the correct value, its accuracy can be judged using percent error. Percent error is calculated as follows:

Value accepted – Value experimental

Percent Error = x100% Value accepted

Errors associated with apparatus:

Whenever it is available manufacturer's information (labels) should be reported according to the specifications for the particular class of instrument e.g. on pipettes, burettes and measuring cylinders if obvious.

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If no error is stated on the apparatus, the following guidelines should be followed:

a) For most of the scale divisions (analogue equipment), the reading uncertainty is equal to ± half scale division. Thermometers, measuring cylinders and burettes are examples of such apparatus

For example, if a thermometer goes up in increments of 0.1oC then the error will be ± 0.05oC (NB: readings and errors must be given to the same number of decimal places, so a reading of 24.1oC should be recorded as 24.10 ± 0.05oC ).

b) For a digital readout, the absolute uncertainty is equal to the least count (not one half the least count) or more specifically, the absolute uncertainty is equal to the unit of the last place of the readout. Balances and pH meters are examples of apparatus where this is appropriate.

For example, if a digital balance displays 63.24 grams for the mass of an object, then the error = 0.01 grams. So the measured mass of the body would be: (63.24 ± 0.01) grams.

Find as many different pieces of apparatus as possible in the laboratory and

state the errors associated with them, classified as Manufacturer’s, Digital or

Analogue:

________________________ ________________________________________________ ________________________________________________ ________________________________________________ ________________________________________________ ________________________________________________ ________________________________________________ ________________________

The following pages cover material that is stated in the actual IB Chemistry syllabus for topic 11 – this could be discussed in class or set as self study.

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Topic 11 Measurement and Data Processing

Pages 75 – 77 of the Chemistry for the IB Diploma IB Study guide (Neuss) should also be referred to.

11.1 Uncertainty and Error

Precision and accuracy:

Precision: A measure of the amount of human error in taking a reading or a measurement.

Accuracy: The equipment used will have a limit to the accuracy with which measurements or readings can be taken. Most lab equipment has a standard level of accuracy.

Therefore measurements can be precise but may not be accurate.

E.g. the volume of solution in a pipette may be recorded precisely using the meniscus of the liquid, however it will still be inaccurate due to the limitations of the accuracy of the pipette itself.

Random Uncertainties:

By combining together the precision of the measurement with the accuracy of the equipment the total random uncertainty can be calculated for each reading taken.

This is given as a + and – as it could be above or below the actual reading.

E.g. Standard accuracy of 25cm3 pipette = ± 0.05 cm3

Estimation of precision with which readings are taken = ± 0.1 cm3

As opposed to estimating precision, the usual practice is to take a sufficient number of repeats and obtain a RANGE of values for any specific reading. Other random errors may arise from difficulty in determining colour changes etc.

Total random uncertainty = ± 0.15 cm3

(Random uncertainties are estimated and presented during DATA COLLECTION)11.2 Uncertainties in Calculated Results

Processing Random Uncertainties

Random uncertainties can be presented as % uncertainty.

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E.g. In a simple titration:

Uncertainty due to 25cm3 pipette = ± 0.15 cm3

Uncertainty due to burette = ± 0.05 cm3 for a titre reading of 14.25 cm3

% Uncertainty due to pipette = (0.15 / 25) x 100 = ± 0.6 %% Uncertainty due to burette = (0.05 / 14.25) x 100 = ± 0.4%

Total random uncertainties = 0.6 + 0.4 = ± 1.0%

Note: Round up % uncertainties to 1 or 2 sig figs depending on circumstances.

Use of Significant Figures

When an answer is calculated, awareness must be displayed that it has an uncertainty.

This can be done by choosing an appropriate number of significant figures.

E.g. Calculated concentration of acid by titration = 1.3245 mol dm-3 ± 1.3%

Since 1.3% of 1.3245 = ± 0.017 mol dm-3

Therefore the concentration could be: 1.3245 + 0.017 = 1.3415 mol dm-3

or: 1.3245 – 0.017 = 1.3245 mol dm-3

Clearly there is uncertainty in third significant figure of the answer.

The concentration should therefore be expressed as: 1.32 ± 0.02 mol dm-3

Or as a range: 1.342 – 1.325 mol dm-3.

(Random uncertainties are processed during DATA PROCESSING)

Systematic Errors

Systematic errors are qualitative explanations describing why the answer given by experiment differs from the actual literature/theoretical value.

If the experimental answer is below the literature value, the systematic errors attempt to explain why this may be. These explanations may be in the form of:

Human errors (e.g. misinterpretation of a colour change at end point) Equipment errors (e.g. the solid didn’t filter successfully, instrument was

calibrated incorrectly)

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Chemical errors (e.g. there were small amounts of other substances in the solution that the acid may have reacted with, incomplete combustion)

There does not need to be any quantitative estimation of such errors.

(Systematic errors are explained and presented in the EVALUATION).

11.3 Graphical Techniques

Construct a variety of types of graph to represent experimental data.

Choose appropriate scales for graphs. This can affect the accuracy of any calculations (e.g. gradients) using the graph.

Use lines of best fit that can either be straight or smooth curves. Try to get the same number of points on each side of the line.

Interpret the shape of a graph and use to show proportional or inversely proportional relationships.

Although uncertainty bars are not specifically required in IB chemistry, HL students are encouraged to show some appreciation of the uncertainty in graphs.

HL students are encouraged to use simple uncertainty bars and if a gradient is calculated, appreciate the uncertainty in the gradient (max/min).

For further clarification, please read through the following extract from the IB

teacher’s guide (even they need help, sometimes).

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Errors and uncertainties in chemistry internal assessment

The consideration and appreciation of the significance of the concepts of errors and uncertainties helps to develop skills of inquiry and thinking that are not only relevant to the group 4 experimental sciences. The evaluation of the reliability of the data upon which conclusions can be drawn is at the heart of a wider scientific method that IB students consider in other areas such as group 3 individuals and societies, and theory of knowledge. They then may apply this in their subsequent educational, professional and personal lives.

Errors and uncertainties are addressed in “Topic 11: Measurement and data processing” of the subject guide and this topic can be very effectively treated through the practical scheme of work.

The treatment of errors and uncertainties is directly relevant in the internal assessment criteria of:

data collection and processing, aspects 1 and 3 (recording raw data and presenting processed data)

conclusion and evaluation, aspects 1, 2 and 3 (concluding, evaluating procedure(s), and improving the investigation).

Expectations at standard level and higher level

The expectations with respect to errors and uncertainties in internal assessment are the same for both standard and higher level students and are supported by topic 11 of the subject guide.

Within internal assessment students should be able to do the following:

make a quantitative record of uncertainty range (±) (data collection and processing: aspect 1)

state the results of calculations to the appropriate number of significant figures. The number of significant figures in any answer should reflect the number of significant figures in the given data (data collection and processing: aspect 3).

propagate uncertainties through a calculation so as to determine the uncertainties in calculated results and to state them as absolute and/or percentage uncertainties (this applies to both higher and standard level students). Only a simple treatment is required. For functions such as addition and subtraction absolute uncertainties can be added. For multiplication, division and powers, percentage uncertainties can be added. If one uncertainty is much larger than others, the approximate uncertainty in the calculated result can be taken as due to that quantity alone (data collection and processing: aspect 3).

determine from graphs physical quantities (with units) by measuring and interpreting a slope (gradient) or intercept. When constructing graphs from experimental data, students should make an appropriate choice of axes and scale, and the plotting of points should be clear and accurate. (Millimetre square graph paper or software is appropriate. Quantitative measurements should not be made from sketch graphs.) The uncertainty requirement can be satisfied by drawing best-fit curves or straight lines through data points on the graph (data collection and processing: aspect 3). (Note: Chemistry students at SL and HL are not expected to construct uncertainty bars on their graphs and may achieve “complete” for aspect 3 of data collection and processing without them. However, students, probably

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those who also study IB physics, often construct error bars and there is no requirement to discourage them from doing so.)

justify their conclusion by discussing whether systematic errors or further random errors were encountered. The direction of any systematic errors should be appreciated. The percentage error should be compared with the total estimated random error as derived from the propagation of uncertainties (conclusion and evaluation: aspect 1).

comment about the precision and accuracy of the measurements when evaluating their procedure (conclusion and evaluation: aspect 2)

suggest how the effects of random uncertainties may be reduced and systematic errors be eliminated. Students should be aware that random, but not systematic, errors are reduced by repeating readings (conclusion and evaluation: aspect 3).

Explaining terms and concepts(a) Random and systematic error

Systematic errors arise from a problem in the experimental set-up that results in the measured values always deviating from the “true” value in the same direction, that is, always higher or always lower. Examples of causes of systematic error are miscalibration of a measuring device or poor insulation in calorimetry experiments.

Random errors arise from the imprecision of measurements and can lead to readings being above or below the “true” value. Random errors can be reduced with the use of more precise measuring equipment or its effect minimized through repeat measurements so that the random errors cancel out.

(b) Accuracy and precision

Accuracy is how close a measured value is to the correct value, whereas precision indicates how many significant figures there are in a measurement. For example, a mercury thermometer could measure the normal boiling temperature of water as 99.5° C (±0.5° C) whereas a data probe recorded it as 98.15° C (±0.05° C). In this case the mercury thermometer is more accurate whereas the data probe is more precise. Students should appreciate the difference between the two concepts (topic 11.1.2).

(c) Uncertainties in raw data

When numerical data is collected, values cannot be determined exactly, regardless of the nature of the scale or the instrument. If the mass of an object is determined with a digital balance reading to 0.1 g, the actual value lies in a range above and below the reading. This range is the uncertainty of the measurement. If the same object is measured on a balance reading to 0.001 g, the uncertainty is reduced, but it can never be completely eliminated. When recording raw data, estimated uncertainties should be indicated for all measurements.

There are different conventions for recording uncertainties in raw data.

The simplest is the least count, which simply reflects the smallest division of the scale, for example ±0.01 g on a top pan balance.

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The instrument limit of error: this is usually no greater than the least count and is often a fraction of the least count value. For example, a burette is often read to half of the least count division. This would mean that a burette value of 34.1 cm3 becomes 34.10 cm3

(±0.05 cm3). Note that the volume value is now cited to one extra decimal place so as to be consistent with the uncertainty.

The estimated uncertainty takes into account the concepts of least count and instrument limit of error but also, where relevant, higher levels of uncertainty as indicated by an instrument manufacturer, or qualitative considerations such as parallax problems in reading a burette scale, reaction time in starting and stopping a timer, random fluctuation in a voltmeter read-out, or difficulties in knowing just when a colour change has been completed in a rate experiment or titration. Students should do their best to quantify these observations into the estimated uncertainty.

In chemistry internal assessment it is not specified which protocol is preferred and a moderator will support a teacher when it is clear that recording uncertainties has been required and the uncertainties recorded are of a sensible and consistent magnitude.

(d) Propagating errors

Random errors (uncertainties) in raw data feed through a calculation to give an error in the final calculated result. There is a range of protocols for propagating errors. A simple protocol is as follows:

1. When adding or subtracting quantities, then the absolute uncertainties are added.

For example, if the initial and final burette readings in a titration each have an uncertainty of ±0.05 cm3 then the propagated uncertainty for the total volume is (±0.05 cm3) + (±0.05 cm3) = (±0.10 cm3).

2. When multiplying or dividing quantities, then the percent (or fractional) uncertainties are added.

For example,

molarity of NaOH(aq) = 1.00 M (±0.05 M) percent uncertainty = [0.05/1.00]×100 = 5%

volume of NaOH(aq) = 10.00 cm3 (±0.10 cm3)percent uncertainty = [0.10/10.00]×100 = 1%

Therefore, calculated moles of NaOH in solution = 1.00×[10.00/1000] = 0.0100 moles (±6%)

The student may convert the calculated total percent uncertainty back into an absolute error or leave it as a percentage.

Note: A common protocol is that the final total percent uncertainty should be cited to no more than one significant figure if it is greater than or equal to 2% and to no more than two significant figures if it is less than 2%.

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There are other protocols for combining uncertainties such as “root sum of square” calculations. These are not required in IB chemistry but are acceptable if presented by a student.

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(e) Repeated measurements

Repeated measurements can lead to an average value for a calculated quantity. The final answer could be given to the propagated error of the component values in the average.

For example,

Δ Hmean = 102 kJ mol–1 (±10%)

Δ Hmean = [+100 kJ mol–1 (±10%) + 110 kJ mol–1 (±10%) + 108 kJ mol–1 (±10%)] / 3

This is more appropriate than adding the percent errors to generate 30%, since that would be completely contrary to the purpose of repeating measurements. A more rigorous method for treating repeated measurements is to calculate standard deviations and standard errors (the standard deviation divided by the square root of the number of trials). These statistical techniques are more appropriate to large-scale studies with many calculated results to average. This is not common in IB chemistry and is therefore not a requirement in chemistry internal assessment.

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Practical Y - Identifying the formula of a compound

This activity allows you to find the formula of a compound from your own experimental data. It also helps you to think about the uncertainties in your measurements and it gives you practice at converting mass to amount of substance.

Requirements

copper strip (approximately 15 cm x 1 cm) emery paper dry filter paper access to a balance iodine crystals (0.3 g) boiling tube Bunsen burner

Introduction

You know for a fact that the formula of water is H2O. But did you know that, like every fact in chemistry, it has been experimentally determined; it has come from observation and measurement?

One way to determine the formula of water is to break it down into its elements _ hydrogen and oxygen - in an electrolysis experiment:

water hydrogen + oxygen

The hydrogen and oxygen formed can be collected and the volume of each gas measured. The mass of each gas produced can be calculated from the volumes and then converted to the amount in moles. The simplest ratio of moles of atoms can be found and thus the empirical formula of water.

Experiment Mass of Mass ofHydrogen (g) Oxygen (g)

1 0.020 0.1682 0.017 0.1393 0.021 0.157

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The Table on the previous page gives the results of three experiments. Use these data to find the empirical formula of water. The relative atomic masses of hydrogen and oxygen are 1 and 16 respectively (Ar: H, 1; 0,16). The rest of this activity will help you to find experimentally the formula of another compound - a yellow compound formed when copper reacts directly with iodine vapour. In the water example above the compound was broken down into its elements. In this example, you will make a compound from its elements.

The exact masses which react together (reacting masses) can be measured and then used to find the empirical formula of the compound.

What you do

1. Clean a strip of copper with emery paper and wipe it with a piece of dry filter paper. Weigh the copper strip.

2. You need to use about 0.3g of iodine crystals so first weigh a dry boiling tube. (You may find it helpful to stand the boiling tube in a beaker so that it does not roll.) Add a small quantity of iodine to the boiling tube and reweigh it. (CARE Iodine is harmful. Avoid skin contact and do not inhale the vapour.) Keep adding iodine crystals until you have about 0.3g in the tube. It is not necessary to know the exact mass of iodine in the tube, so you do not need to record the mass of the tube or the iodine. 3. Bend one end of the copper strip so that it fits over the lip of the boiling tube and the other end of the copper is about 2cm above the iodine crystals (Figure 1).

HEAT

Figure 1

4. Gently heat the copper nearest to the iodine. This should be done in a fume cupboard. 5. Continue heating until no more iodine vapour is seen but do not heat so strongly that iodine vapour escapes from the tube. Keep moving the tube in and out of the flame to avoid overheating it.

6. Allow the tube to cool. Carefully remove the copper strip and reweigh it.

7. Remove the. yellow coating of copper iodide from the surface of the copper by scraping gently with a spatula onto a piece of scrap paper and

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Copper strip

Boiling tube

reweigh the copper strip.

Record your results

Initial mass of copper strip = gMass of copper strip + copper iodide = gFinal mass of copper strip = g

Using your results

8. What mass of copper has reacted with the iodine?

9. What mass of iodine did it react with?

10. How many moles of copper atoms reacted? (Ar: Cu, 63.5)

11. How many moles of iodine atoms reacted? (Ar: I, 127)

12. How many moles of iodine atoms combine with one mole of copper

atoms?

13. What formula does this indicate for the copper iodide?

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Evaluating your results and procedures

As a matter of routine, on completing an experiment you should review and evaluate the procedures you used so that you can identify any that could have been sources of error. Where your experiment involves measurements, you should consider whether the sources of error in your procedure would tend to make your values larger or smaller. For example, if it was difficult to remove all of the copper iodide then the final mass of copper would be too high.

As well as evaluating your procedures it is also important to bear in mind that when any physical measurement is made there is always some kind of uncertainty associated with the value obtained. This is sometimes referred to as experimental error, although it is not a mistake.

A source of uncertainty is the precision of the instrument being used. Every measuring instrument is designed to measure to a certain level of precision. For example, the copper iodide activity involves several weighings using a balance. The commonest balance you will use reads to 2 decimal places but there is an uncertainty associated with the second place. Remember – balances give DIGITAL readings.

For example, a reading of 11.46g means that the mass is between 11.45g and 11.47g. This can be written as 11.46 ± 0.01g .

± 0.01g is the precision error of the balance.

To compare the importance of the precision errors for different measurements they are usually expressed as percentage errors.

error x 100Percentage error = reading

In this case, the percentage error = 0.01g x 100= 0.08%11.46 g

QUESTIONS

a. What was the percentage error for each of the three weighings that you did?

b. How is the percentage error related to the value of the mass recorded?

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c. If you wanted to reduce the percentage error for the initial mass of copper, would you use a smaller or larger piece of copper?

d. Identify stages in your procedure which could have led to errors. How would these errors have influenced the answer you obtained for I2?

e. Look at all the sources of error associated with weighing and with carrying out your experiment. Which one(s) do you think had the greatest effect on your final result? If you were to repeat the experiment, write down any changes you would make to how you would carry it out or with which parts of the procedure you would take extra care.

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Theory - Criteria used for internal assessment of experimental work As stated previously, internal assessment of your experimental work will

constitute 24% of your overall IB level. The criteria for marking that is

followed may appear confusing at first, but you will have plenty of

opportunities to gain marks in each of the respective disciplines – do not be

discouraged if your initial few marks are a little low! As you become more

familiar with the criteria, so your marks will improve (Year 13 is when you

expect to gain most of the marks that will be used in your final level).

In the next two lessons, you will take a look at the marking criteria as well as

some examples of past students’ work – along with the marks allocated.

For these lessons you will need the following:

IB Diploma Chemistry Guide

Practical Scheme of Work Record (PSOW)

A manila folder

If you do not have these three items, ask your teacher nicely for them

(manners cost nothing, you know).

The following extract gives some clearly set out examples of students’ work,

along with the allocated marks as explained by the IB Diploma assessors.

Note carefully the layout of results, the decimal places, the significant

figures, the errors associated, etc. It is attention to detail and GLP (Good

Laboratory Practice) that will lead to level 6s being obtained.

REMEMBER – always refer to the marking criteria in your Chemistry guide when writing up practical work.

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Interpreting the relevant assessment criteriaData collection and processing: Aspect 1 (recording raw data)

Sample extracts of typical student work from an experiment on volumetric analysis in acid-base titration are shown in tables 1–3.

Table 1: DCP aspect 1 = “partial”

Table 1

a. Titration of standard NaOH against vinegar

Final volume / cm3 42.5 41.5

Initial volume / cm3 2.5 1

Volume of base required / cm3 40.0 40.5

Colour of solution at end point light pink dark pink

Some appropriate raw data is recorded but there are no uncertainties and the number of decimal places is inconsistent. Either of these factors reduces the level of achievement for aspect 1 of data collection and processing to partial/1.

Table 2: DCP aspect 1 = “partial”

Table 2

b. Titration of standard HCl against NaOH

Run 1 Run 2 Run 3

Initial volume / cm3(±0.1 cm3) 0.0 2.7 1.0

Final volume / cm3(±0.1 cm3) 42.2 42.7 41.5

Volume of base required / cm3(±0.2 cm3) 42.2 40.0 40.5

Some appropriate raw data is recorded with units and uncertainties. However, relevant qualitative observations were not recorded and the level of achievement for aspect 1 of data collection and processing is partial/1.

Table 3: DCP aspect 1 = “complete”

Table 3

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c. Titration of 5.00 cm3 vinegar against the standardized NaOH

  Trial 1 Trial 2 Trial 3

Initial volume / cm3 (±0.05 cm3) 1.00 2.55 0.00

Final volume / cm3 (±0.05 cm3) 42.50 43.25 40.50

Total volume of base required / cm3 (±0.1 cm3) 41.5 40.7 40.5

Colours of solutions: acid, base and phenolphthalein indicator were all colourless. At the end point, the rough trial was dark pink. The other two trials were only slightly pink at the end point.

The student records appropriate qualitative and quantitative raw data, including units and uncertainties. The level of achievement for aspect 1 of data collection and processing is complete/2.

The following examples of data collection and processing (see tables 4–6) are from a gas law experiment.

Table 4: DCP aspect 1 = “complete”

Table 4

Temperature T / °C ± 0.2° C Height of column h / mm ± 0.5 mm

10.5 58.0

20.3 60.5

30.0 61.0

39.9 64.0

50.1 64.5

60.2 67.5

70.7 68.0

80.8 71.0

90.0 71.5

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In table 4, the student designed the data table and correctly recorded the raw data, including units and uncertainties. The level of achievement for aspect 1 of data collection and processing is complete/2.

Table 5: DCP aspect 1 = “partial”

Table 5

Temperature (T) Height of column (h)

10.5 58.0

20.3 60.5

30.0 61.0

39.9 64.0

50.1 64.5

60.2 67.5

70.7 68.0

80.8 71.0

90.0 71.5

In table 5, units and uncertainties are not included. The level of achievement for aspect 1 of data collection and processing is partial/1.

Table 6: DCP aspect 1 = “partial”

Table 6

Temperature Height of column

10.5 58

20.3 60.5

30 61

39.9 64

26

50 64.5

60.2 67.5

70.7 68

80.8 71

90 71.5

In table 6, units and uncertainties are not included, and the data is recorded in an inconsistent manner. Significant digits are not appreciated. The level of achievement for aspect 1 of data collection and processing is partial/1.

Note: In investigations where a very large amount of data is recorded (probably by a data logger), it may be more appropriate to present the data as a graph. The uncertainties should be recorded on the axis labels and any qualitative observations recorded as annotations on or below the graph.

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Data collection and processing: Aspect 3 (presenting processed data)

Figures 1 and 2 show graphs of the gas law data from table 4.

Figures 1 and 2: DCP aspect 3 = “complete”

Figure 1

Figure 2

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Figure 1 is a graph of the gas law data showing the significant uncertainty. The computer drew the uncertainty bars based on the student entering the correct information, which in this case was 0.5 mm for each value. Figure 2 does not show the uncertainty bars. In chemistry, students are not expected to construct uncertainty bars. In both graphs the title is given (although it should be more explicit), and the student has labelled the axes and included units. The level of achievement for aspect 3 of data collection and processing for both graphs is complete/2.

Figure 3: DCP aspect 3 = “partial”

Figure 3

In figure 3, the student does not include a title for the graph, and the units are missing. The level of achievement for aspect 3 of data collection and processing is partial/1.

Figure 4: DCP aspect 3 = “not at all”

Figure 4

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In the examples shown in figure 4, the first student has failed to draw a best-fit line graph and the second has drawn no line at all. The units and the titles of the graphs are missing. In the second graph, poor use is made of the x-axis scale/ The level of achievement for aspect 3 of data collection and processing for both graphs is not at all/0.

Conclusion and evaluation: Aspects 1, 2 and 3 (concluding, evaluating procedure and improving the investigation)

When attempting to measure an already known and accepted value of a physical quantity, such as the charge of an electron, the melting point of a substance, or the value of the ideal gas constant, students can make two types of comments:

1. The error in the measurement can be expressed by comparing the experimental value with the textbook or literature value. Perhaps a student measured the value of the ideal gas constant R = 8.11 kPa dm3 mol–1 K–

1, and the accepted value is 8.314 kPa dm3 mol–1 K–1. The error (a measure of accuracy, not precision) is 2.45% of the accepted value. This sounds good, but if, in fact, the experimental uncertainty is only 2%, random errors alone cannot explain the difference, and some systematic error(s) must be present.

2. The experimental results fail to meet the accepted value (a more relevant comment). The experimental range does not include the accepted value. The experimental value has an uncertainty of only 2%. A critical student would appreciate that they must have missed something here. There must be more uncertainty and/or errors than acknowledged.

In addition to the above two types of comment, students may also comment on errors in the assumptions of the theory being tested, and errors in the method and equipment being used. Two typical examples of student work are given in figures 5 and 6.

Figure 5: CE aspect 1 = “partial”, aspect 2 = “not at all”

Figure 5

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Conclusion and evaluation: Intermolecular bonds are being broken and formed which consumes energy. There is a definite correlation between the melting point and the freezing point of a substance. If good data is collected, the melting point should be the same as the freezing point. A substance should melt, go from solid to liquid, at the same temperature that it freezes, goes from liquid to solid. Our experiment proved this is true because, while freezing, the freezing point was found to be 55 °C, and when melting, the melting point was also found to be 55 °C (see graph).

The student states a conclusion that has some validity. No comparison is made with the literature value. There is no evaluation of the procedure and results. The level of achievement for aspect 1 of conclusion and evaluation is partial/1. The level of achievement for aspect 2 is not at all/0.

Figure 6: CE aspects 1, 2 and 3 = “complete”

Figure 6

Melting point = freezing point = 55.0 ± 0.5 °C

Conclusion and evaluation: Literature value of melting point of para-dichlorobenzene = 53.1° C (Handbook of Chemistry and Physics).

The fact that % difference > % uncertainty means random errors alone cannot explain the difference and some systematic error(s) must be present.

Melting point (or freezing point) is the temperature at which the solid and the liquid are in equilibrium with each other: (s) ⇌ (l). This is the temperature at which there is no change in

31

kinetic energy (no change in temperature), but a change in potential energy. The value suggests a small degree of systematic error in comparison with the literature value as random errors alone are unable to explain the percentage difference.

Evaluation of procedure and modifications:

1.Duplicate readings were not taken. Other groups of students had % uncertainty > % difference, that is, in their case random errors could explain the % difference, so repeating the investigation is important.

2.How accurate was the thermometer? It should have been calibrated. In order to eliminate any systematic errors due to the use of a particular thermometer, calibration against the boiling point of water (at 1 atmosphere) or better still against a solid of known melting point (close to the melting point of the sample) should be done.

3.The sample in the test tube was not as large as in other groups. Thus the temperature rises/falls were much faster than for other groups. A greater quantity of solid, plus use of a more accurate thermometer (not 0.5°C divisions, but the longer one used by some groups) would have provided more accurate results.

The level of achievement for aspects 1, 2 and 3 of conclusion and evaluation is complete/2.

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Practical Z – An investigation into solubility

You are going to investigate the way that the solubility of a salt changes when we change the temperature, and to graph your results. Marks will be awarded for the way in which you carry out the experiment. This means CONTROLLED heating, CAREFUL measurement of masses and volumes and GOOD OBSERVATION as well as the ability to calculate the solubility and draw a graph.

1. Weigh out exactly 5g of potassium CHLORATE on to a square of paper, and make sure that all of the crystals go in to a dry boiling tube.

2. Clamp the tube and contents onto a stand, and arrange it so that it can be lowered easily in and out of a 250 ml beaker half full of tap water. Marks will be awarded for how well you can operate these pieces of equipment.

3. Add 10 ml. (10g) of PURE water using the measuring cylinder or syringe provided. The salt will not dissolve in 10g of water at this temperature no matter how much it is stirred, but will dissolve if heated to a higher temperature.

4. Raise the temperature of the tube gradually by placing it into the beaker of water and heating the beaker of water. Watch carefully until the solid completely dissolves, then IMMEDIATELY turn off your Bunsen and raise the tube from the hot water by raising the clamp. [MAKE SURE YOU PRACTICE DOING THIS A COUPLE OF TIMES BEFORE STARTING THE EXPERIMENT!] There is no need to try to record the temperature at which the solid just dissolved; it is much more accurate to record the temperature at which precipitation occurs because the temperature has cooled.

5. As the tube cools watch out for the appearance of small crystals. This occurs when the solution is just saturated. It is a good idea to CAREFULLY stir the solution as it cools with the thermometer, and be ready to record the EXACT temperature at which this 'snowing' occurs. If you are not paying full attention you will miss this!

6. When you have read the temperature record it in the appropriate place in

the results table.

7. Now add a further 5ml (5g) of pure water using the syringe. Stir with the thermometer CAREFULLY and check that the salt does not dissolve. Now place the tube and warm the solution again until the salt just dissolves. This should be at a lower temperature than before, because you now have more water present to dissolve the salt.

8. Repeat operations 5 and 6 to find out the new temperature at which the solution becomes saturated.

9. Add a further 5ml of pure water and repeat operations 4, 5, and 6 to obtain

33

a new temperature for the saturation of 5g of your salt in 20ml (20g) of water.

10. Repeat until you have a total of 30ml of water in your solution, or until the solute remains dissolved even at room temperature. You should have 6 temperatures written down. Now try to do the calculations set out below.

RESULTS TABLE – Recorded as per IB Marking Criteria, with all associated errors detailed

Calculations

WORK OUT the solubility of your solute in g per l00g of water at each of the temperatures recorded. Propagate errors as appropriate.

Graph your results by hand (line of best fit) or using a software programme like Excel (if using EXCEL then graph must be an X-Y scatter with a trend line).

Plot solubility of KClO3 per 100g. water against temperature.

e.g. at 78 °C 5g of KClO3 dissolved in 10g of water

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therefore the solubility of KClO3 at 78°C is 5g x l00g/l0 g

= 50g per l00g water.

Question: What factors affect the solubility’s of the ions shown in your

results?

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Appendix 1: Questions on Techniques & Safety

Section A: Identify the following pieces of equipment.

Section B: Answer the following questions in complete sentences.

1. As soon as you enter the lab, what safety equipment should you put on immediately?

2. Before doing an experiment, what should you read and discuss with your teacher?

3. Before you light a burner, what safety precautions should always be followed?

4. What type of flame is preferred for laboratory work and why?

5. Why are broken glassware, chemicals, matches, etc. never thrown into a wastepaper basket?

36

6. Why should you never touch chemicals with your hands?

7. What precaution can you take against chemical contamination in reagent bottles?

8. Why are chemicals and hot objects never placed directly on a balance pan?

9. List three instruments used in the laboratory for measuring small quantities of liquids. What precautions should be taken when filling a burette with liquid?

10. How can a liquid be transferred from a beaker to a funnel without spattering and without running down the outside wall of the beaker?

11. What procedure is used to dilute a concentrated acid?

12. How would you dispose of a spillage of concentrated sulphuric acid?

Section C: Hazard Symbols

Identify each of these hazard symbols:

Appendix 2: Techniques and Safety

Decanting and Transferring Liquids

1. The safest way of transferring a liquid from one test tube to another is shown in Figure S-1. The liquid is

37

transferred at arm's length with the elbows slightly bent. This position enables you to see what you are doing and still maintain steady control.

Figure S-1.

2. Sometimes liquids contain particles of insoluble solids that sink to the bottom of a test tube or beaker. Use one of the methods shown below to separate a supernatant (the clear fluid) from insoluble solids.

(a) Figure S-2 shows the proper method of decanting a supernatant liquid in a test tube.

Figure S – 2

(b) Figure S-3 shows the proper method of decanting a supernatant liquid in a beaker by using a stirring rod.

The rod should touch the wall of the receiving vessel. Hold the stirring rod against the lip of the beaker containing the supernatant liquid. As you pour, the liquid will run down the rod and drop off into the beaker resting below. In this way the liquid will not run down the side of the beaker from which you are pouring.

38

Heating Substances and Evaporating Solutions

1. Use care in selecting glassware for high-temperature heating. The glassware should be Pyrex or a similar heat-resistant type.

2. When heating substances in glassware by means of a gas flame, use a ceramic-centered wire gauze to protect glassware from direct contact with the flame. These wire gauzes can withstand extremely high temperatures and will help prevent glassware from breaking.

Figure S-4 shows the proper setup for evaporating a solution over a water bath.

Figure S-4

3. In some experiments you are required to heat a substance to high temperatures in a porcelain crucible. Figure S-5 shows the proper apparatus setup used to accomplish this task.

4. Figure S-6 shows the proper setup for evaporating a solution in a porcelain evaporating dish with a watch glass cover that prevents spattering.

5. Heated glassware, porcelain, and iron rings look cool several seconds after they are removed from a heat source, but can still burn your skin for several minutes. Use heat-safety items such as safety tongs, heat-resistant mittens and pads, aprons, rubber gloves, and safety goggles whenever you handle this apparatus.

Figure S-5

6. You can test the temperature of questionable beakers, ring stands, wire gauzes, or other pieces of apparatus that have been heated by holding the back of your hand close to their surfaces before grasping them.

39

Any heat generated from the hot surfaces will be felt - DO NOT TOUCH. Allow plenty of time for the apparatus to cool before handling.

How to Pour Liquid from a Reagent Bottle

1. Read the label at least three times before using the contents of a reagent bottle.

2. Never lay the stopper of a reagent bottle on the lab table. Remove the stopper by grasping the stopper between two fingers.

3. Pick up the reagent bottle making sure the label is toward the palm of your hand. Note that the stopper is still between the fingers.

4. When pouring a caustic or corrosive liquid into a beaker, use a stirring rod to avoid drips and spills. Hold the stirring rod against the lip of the reagent bottle. Estimate the amount of liquid you need and pour this along the rod into the beaker.

5. Extra precaution should be taken when handling a bottle of acid. Remember these two important rules: (a) Never add water to any concentrated acid, particularly sulphuric acid, because of splashing and heat generation, (b) To dilute any acid, add the acid to water in small quantities, stirring slowly and constantly.

Remember the "triple A's"—Always Add Acid to water.

6. Replace the stopper on the reagent bottle after you are finished pouring. 7. Examine the outside of the reagent bottle for any liquid

that has dripped down the bottle or spilled on the counter top. Your teacher will show you the proper procedures for cleaning up a chemical spill.

8. Never pour reagents back into stock bottles. At the end of the experiment, any excess chemicals should be properly discarded under the direction of your teacher.

How to Heat Material in a Test Tube

40

Note:You will more likely use a

tripod, rather than an iron ringon a clamp stand, but otherwise the set-up is

the same.

1. Check to see that the test tube is PYREX or a similar heat-resistant type.

2. Always use a test tube holder or clamp when heating the test tube.

3. Never point a heated test tube at anyone, because the liquid may splash out of the test tube.

4. Never look down into the test tube while heating it.

5. Do not heat any one spot on the test tube. Heat the test tube from the upper portions of the tube downward and continuously move the test tube. Otherwise pressure from a vapour meeting a layer of liquid above it m may cause the bottom of the tube to blow out.

How to Use a Mortar and Pestle

1. A mortar and pestle should be used for grinding only one substance at a time.

2. Never use a mortar and pestle for simultaneously mixing different substances.

3. Place the substance to be broken up into the mortar.

3. Pound the substance with the pestle and grind to pulverize.

5. Remove the powdered substance with a spatula.CAUTION Do not blow into the mortar to remove any remaining powder, since dust may get into eyes and nasal passages.

Testing an Odour Safely

41

Figure S-11

1. Test for the odour of gases by wafting your hand over the test tube and cautiously sniffing the fumes.

2. Do not inhale any fumes directly.

3. Use a fume cupboard whenever poisonous or irritating fumes are evolved.

DO NOT waft and sniff poisonous or irritating fumes.

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Liquid Measurements – Measuring Cylinders, Pipettes & Burettes.

1. For approximate measurements of liquids, a measuring cylinder such as the one shown to the right is generally used. These cylinders are usually graduated in millilitres (ml). Such a measuring cylinder may read from 0 to 10 ml, 0 to 25 ml, or more, from bottom to top.

2. A pipette or a burette is used for more accurate measurements. Pipettes are made in many sizes and are used to deliver measured volumes of liquids. A pipette is fitted with a suction bulb (pipette filler) which is used to withdraw air from the pipette while drawing up the liquid to be measured. Always use the suction bulb.

3. Burettes, fitted with a tap, are used for delivering any desired quantity of liquid up to the capacity of the burette. Many burettes are graduated in tenths of millilitres. When using a burette, follow these steps:

a) Clamp the burette in position on a clamp stand. See Figure 1-10.

b) While filling the burette, place it on a stool, so that it is below eye-level.

c) Place a beaker, 250-ml, at the bottom of the burette. The beaker serves to catch any liquid that will be drawn off.

d) Pour into a 100-ml beaker a quantity of the liquid you want to measure from the liquid's reagent bottle. Remember to carefully check the label of the reagent bottle before removing any liquid.

e) Fill the burette with the liquid and then draw off enough liquid to fill the tip below the tap and bring the level of the liquid down to the zero. The height at which the liquid stands is then read accurately. (You may need to flick the tap with your fingers to remove air bubbles.)

4. Observe that the surface of a liquid in the burette is slightly curved. Such a curved surface is called a meniscus. You read to the bottom of the meniscus, as shown in Figure 1-11. This is the line AB. If you read the markings at the top of the meniscus, CD, you will get an incorrect reading.

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Measuringcylinder

Pipette

Burette

5. After you have taken your first burette reading, as directed, open the tap and draw off as many millilitres of the liquid as you wish. The exact amount drawn off is equal to the difference between your first and final burette readings.

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Notes:All the glassware used here may be

graduated in cubic centimetres (cm3) instead of millilitres (ml). 1ml = 1cm3

Glassware should be read to an accuracy of half of one division e.g. if a measuring

cylinder is graduated in 1cm3 steps, give the reading to the nearest 0.5 cm3. i.e. read 14.0

cm3 or 17.5 cm3

Appendix 3: Table of Ion valencies.

+1 (I) +2 (II) +3 (III) +4 (IV) -3 -2 -1

H+

Li+

Na+

K+

Rb+

Cs+

NH4+

(ammonium)

Cu+

(copper(I))

Mg2+

Ca2+

Sr2+

Ba2+

Cu2+

(copper(II))

Fe2+ (iron(II))

Mn2+

(manganese(II))

Ni2+

(nickel(II))

Zn2+

Sn2+ (tin(II))

Pb2+ (lead(II))

Al3+

Fe3+ (iron(III))

Cr3+

(chromium(III))

V3+

(vanadium(III))

Sn4+ (tin(IV)

Pb4+

(lead(IV))

Mn4+

(manganese(IV))

N3- (nitride)

P3-

(phosphide)

PO43-

(phosphate(V))

O2-

S2- (sulphide)

SO32-

(sulphate(IV) – sulphite)

SO42-

(sulphate(VI) – sulphate)

CO32-

(carbonate)

Cr2O72-

(dichromate(VI))

F-

Cl-

Br-

I-

NO2-

(nitrate(III) - nitrite)

NO3-

(nitrate(V) – nitrate)

MnO4-

(manganate(VII))

ClO-

(chlorate(I) - hypochlorite)

ClO3-

(chlorate(V))

IO3- (iodate(V))

Notes:Some d-block metals, and group 4 elements have more than one possible oxidation state. When writing the name the oxidation state is indicated using Roman numerals e.g. tin(II) chloride, iron(III) sulphate

Names in italics are old-fashioned, but are still commonly used. The old convention was to use –ite to indicate the lower oxidation state and –ate to indicate the higher oxidation state. So nitrate(III) was nitrite, and nitrate(V) was simply nitrate. If nitrate, sulphate etc appear without a Roman numeral, assume the higher oxidation state.

Appendix 4: Organic Functional Groups

Many important organic chemistry molecules contain oxygen or nitrogen. It's essential to memorise the names and structures of these functional groups

45

Other functional groups that you will become familiar with include:

Alkanes Alkenes Alkynes Haloalkanes (Halogenoalkanes) Aromatic Compounds that include a Benzene ring

46