Welcome to the Practical Guidelines for Cambridge International AS Level Chemistry (9701). In this guide, you will find helpful instructions and tips for the practical component of the AS Level Chemistry exam. These guidelines are designed to assist you in performing experiments and understanding the requirements of the practical paper. By following these instructions, you can better prepare for success in your chemistry practical and exams.

Titration:

A burette is a long, thin tube used in chemistry labs to measure and dispense precise amounts of liquids. It has markings along its length, like a ruler, so scientists can see exactly how much liquid they’re using. Burettes are commonly used in experiments where accuracy is important, like titrations, where you need to mix specific amounts of chemicals together. They usually have a stopcock at the bottom to control the flow of liquid. 

A burette can be started from any point. Typically, we use a 50.0 cm³ burette. Initially, we fill the burette to 0.0 cm³. We consider 0.0 cm³ as the initial reading. The final reading can vary, but in most cases, it falls within the range of 23.0 cm³ to 27.0 cm³. The difference between the final and initial burette readings is called the titre. For example, if the rough titre was 23.0 cm³, you can start the next titration from 23.0 cm³ because the final burette reading of the next titration falls within the 50 cm³ range of the burette. However, if the rough titre is greater than 25.0 cm³, you should start the next titration from 0.0 cm³.

The correct burette reading is 27.50 cm³ 26.50 cm³. Always start to read from the upper side of the burette. You must consider the lower meniscus for the colourless solution. It means to read the mark along the bottom of the liquid surface as shown in the upper diagram.

The two best titrations must be consecutive and within 0.1 cm³ of each other. If the first titration was 24.40 cm³, the second was 24.10 cm³, and the third was 24.50 cm³, then the first and third titers are the best titers, and the reading is nearer. 

If the first two titers are within 0.1 cm3 then there is no need for third titration. If you think that any reading is not correct because it is out of pattern or too far away from the line of graph circle that point on the graph, again take that reading on the same table but on a new row, and then plot a new point.

All readings on the table must be consistent. All your burette readings must be given to the nearest 0.05 cm3. For example, if your volume comes out to be 24.1 cm³, then don’t write it as 24.1 cm³ rather write it as 24.10 cm³.

Always beware of endpoints. Whenever you know the endpoint is near, add the solution in the burette drop by drop and close tap when the endpoint is reached without overly titrating.

All readings should be in a single organized table. 1 MARK would be deducted otherwise. An example of a good table for recording readings while performing titrations is given below (you may memorize this table!) :

Redox Titration:

The titration of acidified potassium permanganate (KMnO₄) with iron (II) salt is a common laboratory procedure used to determine the concentration of iron (II) ions in a solution. Here’s a step-by-step procedure:

Materials:

1. Acidified potassium permanganate solution (KMnO₄)

2. Iron (II) salt solution

3. Burette

4. Pipette

5. Conical flask

6. Distilled water

7. Indicator (optional) – typically a starch solution or a suitable indicator like ferroin.

Procedure:

1. Preparation of Solutions:

   – Prepare a standard solution of acidified potassium permanganate by dissolving a known mass of KMnO₄ in distilled water and adding sulfuric acid to acidify the solution.

   – Prepare a solution of iron (II) salt of known concentration. This could be ferrous ammonium sulfate, ferrous sulfate, or any other suitable iron (II) salt.

2. Standardization of KMnO₄:

   – Pipette a known volume (e.g., 25 mL) of the iron (II) salt solution into a clean conical flask.

   – Add a few drops of indicator, if using.

   – Titrate the iron (II) solution with acidified KMnO₄ solution from the burette until a persistent pink coloration is obtained. This indicates the endpoint of the titration.

   – Record the volume of KMnO₄ solution used.

   – Repeat the titration until concordant results are obtained.

3. Titration of Iron (II) Salt:

   – Pipette a known volume (e.g., 25 mL) of the iron (II) salt solution into a clean conical flask.

   – Add a few drops of indicator, if using.

   – Titrate the iron (II) solution with acidified KMnO₄ solution from the burette until a persistent color change occurs. The endpoint is typically signaled by a color change from pink to colorless due to the reduction of permanganate ions to manganese (II) ions.

   – Record the volume of KMnO₄ solution used.

4. Calculation:

   – Use the volume of acidified KMnO₄ solution and its concentration to calculate the number of moles of KMnO₄ used in the titration.

   – Since the balanced chemical equation for the reaction between Fe(II) and KMnO₄ is known, you can use stoichiometry to determine the number of moles of Fe(II) ions present in the solution.

   – From the number of moles of Fe(II) ions and the volume of the solution, calculate the concentration of Fe(II) ions in the original solution.

Always remember to handle chemicals with care, wear appropriate safety gear, and perform the titration in a well-ventilated area.

Thiosulfate titration

The thiosulfate titration method is commonly used to determine the concentration of an oxidizing agent, in this case, acidified potassium iodate KIO₃. The reaction involved is:

IO₃^– + 5I^– + 6H⁺ =  3I₂ + 3H₂O

Here’s a step-by-step procedure for conducting the thiosulfate titration for acidified potassium iodate and potassium iodide:

Equipment and Reagents Required:

1. Acidified potassium iodate solution (containing excess potassium iodide)

2. Sodium thiosulfate solution (standardized)

3. Starch solution (indicator)

4. Dilute sulfuric acid (H₂SO₄)

5. Burette

6. Pipette

7. Conical flask

8. Bunsen burner

9. White tile or paper

Procedure:

1. Standardization of Sodium Thiosulfate Solution:

   – Pipette a known volume (usually 10 mL) of the acidified potassium iodate solution into a conical flask.

   – Add a few drops of starch solution to the conical flask.

   – Titrate the solution with sodium thiosulfate solution from the burette until the blue color just disappears, indicating the endpoint.

   – Record the volume of sodium thiosulfate solution used.

   – Repeat the titration until concordant results are obtained.

   – Calculate the molarity of the sodium thiosulfate solution using the volume and the known concentration of the potassium iodate solution.

2. Titration of Acidified Potassium Iodate Solution:

   – Pipette a known volume (usually 10 mL) of the acidified potassium iodate solution into a conical flask.

   – Add a few drops of starch solution to the conical flask.

   – Add a few milliliters of dilute sulfuric acid (H₂SO₄) to acidify the solution and release iodine.

   – Titrate the solution with the standardized sodium thiosulfate solution from the burette until the blue color just disappears, indicating the endpoint.

   – Record the volume of sodium thiosulfate solution used.

   – Repeat the titration until concordant results are obtained.

Calculations:

– Calculate the amount of IO₃^– reacted using the stoichiometry of the reaction.

– Use the volume and molarity of the sodium thiosulfate solution to calculate the amount of IO₃^– present in the original solution.

– From the amount of IO₃^–, calculate its concentration in the original acidified potassium iodate solution.

Precautions:

– Ensure proper mixing of solutions by swirling the flask gently during titration.

– Add sulfuric acid carefully to avoid splashes.

– Ensure that the starch solution is added after the addition of acid to prevent premature reaction.

– Perform the titration under consistent lighting conditions to observe the color change accurately.

This procedure provides a quantitative method for determining the concentration of acidified potassium iodate solution using sodium thiosulfate solution as the titrant.

Measuring Mass:

Using the Electronic Balance:

Before measuring any mass, press the “Zero” button to make sure that the initial reading is zero.

Make sure that the electronic balance is clean with no residues on it. Blow lightly if you feel there are any.

The precision of the electronic balance is up to 2 decimal places e.g. 5.27 g.

Units: Electronic balance is measured in grams, so always write units (“g”, “/g”, or “grams”) beside every reading.

All recordings should be made in a single table. Otherwise, your 1 Mark will be at stake.

As already stated, the headings and their units have separate marks. Write mass/g or mass (g) but m/g or mass g are not allowed. The full name of the quantity must be used.

Measuring Temperature:

A laboratory thermometer is a special thermometer used in science labs to measure temperature. Here are some key points:

1. Purpose: It helps scientists measure the temperature of substances accurately during experiments.

2. Design: It looks like a regular thermometer but is designed to be more precise and often has a narrower range.

3. Material: It’s usually made of glass or metal, and the liquid inside (usually mercury or alcohol) expands and contracts with temperature changes.

4. Scale: It typically has a scale marked in degrees Celsius or Fahrenheit, depending on the region’s preference.

5. Usage: Scientists use it by immersing the thermometer in the substance they want to measure the temperature of.

6. Handling: They need to be handled carefully to avoid breakage, and they’re often calibrated regularly to ensure accuracy.

There are only a few tips relating to this measurement. First, always stir the mixture before recording the temperature. Second,  when measuring temperature, make sure that the thermometer is not taken out of the solution.

Calculations:

This is an easier yet trickier part of the paper. Mostly, it will involve calculations regarding moles, enthalpy changes etc.

In this part of the paper, keep the following things in mind:

Formulae of moles: You should understand the relationships between moles and concentration, moles and volume, moles and relative molecular mass, moles and gas volume, etc. If you need help, you can go to the ‘Chemical Calculations’ lesson of the Cambridge International OL Course.

3 significant figures: Calculation should be done to 3 significant figures. Remember that there’s a difference between 3 significant figures and 4 significant figures. 25.6 is 3 significant figures and 25.60 is 4 significant figures & 2 decimal places. 25.659 is 25.7 when read to 3 significant figures. (1 Mark)

Working: Show your working in calculations, and the key steps in your reasoning.

Calculations involving enthalpy: In calculations involving enthalpy, never forget to put a negative sign (-) before exothermic, and a positive sign (+) before endothermic reactions.

Errors & Improvements (Enthalpy experiments):

Graphs:

Layout:

The best-fit line has to be drawn, which means that the line must be an average of all the values. The scattering of points must be uniform: a number of points above and below the line must be the same. The axes must be labeled with the quantity and its units. Values on axes must be written. While plotting, odd scales such as using 10 blocks to represent 3 units is not acceptable. All of the points must be plotted. However, if you suspect one of them is wrong (anomalous point), plot it, and then you MUST circle it, otherwise, you will lose a mark if you do not circle it.

Your graph should cover more than half of the provided graph paper along both axes. For example, if the graph paper provided has 12 big squares along the y-axis and 12 along the x-axis, then your graph should cover at least 6 boxes along the x-axis and 6 boxes along the y-axis.

Gradient:

When calculating the gradient for a line, always remember to draw a large triangle on the graph paper (1 Mark). The triangle should be drawn next to the points where the gradient has to be calculated as shown on the right.

Extrapolating graphs:

There are some practicals where you are required to take maximum temperature change. The change in temperature increases and then falls. By drawing two lines and extrapolating them you need to find the maximum temperature change and where it occurs. In others, you need to extrapolate the line where the value is constant after certain readings and you need to extrapolate the two lines and find where they intersect.

Mark scheme:

Layout: axes labeled with quantities + units.

Layout: line is best fit + intersects + plotted close to 1 mm of the value from table

Quality: close to 1 degree of the max temp of the supervisor.

Marks of these can vary depending on the question and its demand. From the samples given below, you can see examples of drawing best-­fit as well as extrapolating.

Salt Analysis:

Always write observations by looking at observations mentioned in the salt analysis data given at the back of the question paper. Find a close match to your observations from the sheet, and copy the same wordings of the observations given in the datasheet. There will be rare cases when you’ll be unable to find a match. In that case, write down whatever you see.

Moreover, always write proper bench reagents when you are asked to state certain reagents. Never write like H⁺¹ or Cr₂O₇^-2, rather write names of proper bench reagents like HCl(aq) or HNO₃(aq) or K₂Cr₂O_7, etc.

When you stir the boiling tube, don’t let the stirrer touch or strike the bottom of the tube. Whenever heating let the tube give time to heat. First heat gently then strongly. Keep the tube oblique. This way your tube will never break. It’s what my experience has taught me. When adding NaOH or NH3 for ion testing, add a few drops first (DO NOT ADD A DROPPER FULL OF REAGENT AT ONCE). Your eye must be close to the top level of the reagent in a test tube. When a precipitate is observed, add the reagent in excess to almost ⅔ of the test tube. Don’t use more than 1 cm³ of the reagent because it will have more precipitate which will take more NaOH or NH₃ to dissolve; this can confuse you into thinking that ppt . is insoluble.

Distinguishing between Pb⁺² / Al⁺³ :

NaOH / NH₃ reagents when used for identification of these ions give similar results (observations).

PbCl₂, PbI₂, PbSO₄, PbCr₂O₇ or PbCrO₄, are insoluble. Use the following reagents; HCl, KI, K₂Cr₂O_7, or any other reagent that contains the ions mentioned above. The insolubility of the afore-mentioned lead compounds can help us distinguish between the Pb⁺² and Al⁺³ ions.

Also, take care of other possible precipitates that might confuse your results. For example, BaCl₂ (aq) contains Cl^-1 and can be used to test for Pb⁺² ions but the presence of Ba⁺² ions makes precipitates of its own, so always use reagents that have Na⁺¹, K⁺¹, or H⁺¹ ions which always make soluble compounds, reducing the possibility of any other precipitates, except for the precipitates formed by Pb⁺².

The good news! Presently lead ion is not included in the syllabus.

Distinguishing between Ba⁺² / NH₄⁺¹ :

Both give no precipitate with NaOH or NH₃ except that NH₃ is produced on warming NH₄⁺¹ with NaOH. Ba⁺² can be identified by H₂SO₄, it will give white precipitate.

Test for Manganese Mn⁺²:

Mn⁺² can identified with NH₃ / NaOH. With these reagents, it has a white/pale brown ppt. which turns brown when in contact with air and is insoluble in excess of the reagent. There would be brown residues floating on the top surface and on the sides of the test tube, and white ppt./light brown ppt. at the bottom.

Tests for Cu⁺²:

With NaOH: pale blue ppt. insoluble in excess.

With NH₃: Blue ppt. which dissolves and forms a dark blue solution in excess. It will be hard to dissolve the precipitate if too much Cu⁺² is present in the test tube, so use very small quantity of Cu⁺² (less than 1 cm³) or use a lot of NH³ (fill the entire test tube) and shake vigorously to dissolve this precipitate.

Problems with Al⁺³ ions test with NaOH:

The precipitate formed by Al⁺³ is very soluble and disappears very quickly. Students can easily make the mistake of not noticing any precipitate, and writing down that no change occurred. Use a very tiny quantity of NaOH at first, just a few drops (put one drop, shake it lightly, then put another, and so on), and a small white ppt. will form floating on the surface of the solution, which would dissolve very quickly if a very small amount of NaOH is added.

The identification of the other cations and anions is easy. For them just refer to the salt analysis notes given at the end of the paper.

Identification of gases:

Whenever a gas evolves, there is some sort of effervescence (bubbles form) that occurs in the solution. Whenever you notice such a thing, just put your thumb on the top of the test tube. If the pressure builds up, there’ll be definitely some sort of gas evolving.

Now the thing is, how to identify them?

CO₃^-2 :

If an acid is added or is present in the test tube and you see vigorous effervescence, then definitely it’s CO2 that’s evolving. If you have time, just to counter-check, test it with lime water, it will definitely turn milky.  The effervescence produced is similar to gas bubbles in coke. Generally produced when metal carbonates react with acids

NO₂^-1 :

Whenever an acid is added, put your thumb top of the tube, and allow pressure to build up. The tube will turn pale brown and when you release your thumb and allow gas to escape, then a pale brown gas will release. If it occurs then definitely NO₂^-1 is present in the solution. This pale brown gas is also very visible if seen in front of a white background. The gas is especially very visible when the reactants are thrown in the white sink and you will notice brown vapors in the sink easily.

SO₃^-2 :

 Another gas that is produced on the addition of dilute acids is SO₂ which indicates the presence of SO₃^-2 ions. SO₂ gas is colorless and acidic and is produced when dilute acid is added to sulfite SO₃^-2 ions. If a damp blue litmus paper is placed at the mouth of the tube, then it will turn red. Damp litmus paper must not touch the test tube itself as it might contain an acid. Note, that damp blue litmus paper will turn red when NO₂ gas is produced but NO₂ is pale brown and can be distinguished from SO₂. Another test for SO₂ gas is that it smells of rotten eggs or burnt matches. It can also be distinguished by dipping a paper in K₂Cr₂O₇ and then placing it at the mouth of the test tube. This paper will turn from orange to green.

NO₃^-1, NO₂^-1 :

To test for these ions, NaOH is added followed by the addition of Al foil and heated. When bubbles start to form (vigorous bubbling), put a damp red litmus paper near the mouth of the test tube. NH₃ is liberated if these ions are present, and it turns damp red litmus paper blue.

Always use damp red litmus paper, by making the litmus paper wet. Nothing happens if the litmus is not damp! Make sure that the litmus paper never touches the test tube, because the test tube might contain an alkali which will turn the litmus paper blue. A lot of students make the mistake of allowing the litmus paper to touch the top of the test tube, and in many cases, an alkali is present in the test tube which makes the litmus paper blue. So, keep the litmus paper a fair distance (1 cm) away from the test tube.

Students should be able to distinguish between a red litmus paper from a blue litmus paper. Red litmus paper is pale pink, and blue litmus paper is pale blue. Some students also make the mistake of using the cover paper of the litmus paper stack which is also pink (Avoid silly mistakes)

Identification for hydrogen ( H₂ ) gas:

Metal + Acid —> Salt + H₂

Use the above equation to detect the hydrogen gas.

If you are adding metal, and a gas is produced, then you don’t necessarily need to test for Hydrogen gas, if you see effervescence, then it is obviously hydrogen.

Just to confirm if you have time, test it. Hydrogen gas produces a pop sound when burnt with a lighted splint. The only way it produces a pop sound when enough pressure is built up in the test tube. Put your thumb on top of the test tube and allow pressure to build up and only then light it.

Rate of reaction experiments:

In the Cambridge International A Level Chemistry curriculum, the Rate of Reaction experiment is a fundamental practical activity designed to help students understand the factors influencing the speed of a chemical reaction. Here’s a general outline of how such an experiment might be conducted:

The experiment you’re referring to is typically conducted to investigate the effect of varying the concentration of one reactant on the rate of reaction. In this case, the reactants are hydrochloric acid (HCl) and sodium thiosulfate (Na₂S₂O₃).

Here’s how the experiment is usually set up:

1. Apparatus Setup:

   – A conical flask is placed on a piece of paper with a black cross drawn on it.

   – Hydrochloric acid (HCl) of fixed concentration is measured and poured into the flask.

   – Sodium thiosulfate solution of different concentrations is measured using a graduated cylinder.

   – A stopwatch or a clock with a second hand is used to measure time.

2. Reaction:

   – The sodium thiosulfate solution is poured into the conical flask containing the HCl.

   – Upon mixing, a reaction occurs between the HCl and the sodium thiosulfate, producing a yellow precipitate of sulfur (due to the formation of sulfur dioxide gas).

   – As the reaction progresses, the yellow sulfur precipitate gradually obscures the black cross on the paper.

3. Observations:

   – The time taken for the cross to be obscured from view is recorded.

   – This time is an indicator of the rate of reaction: the faster the cross is obscured, the faster the reaction occurs.

4. Variables:

   – The independent variable is the concentration of sodium thiosulfate solution. Different concentrations are tested.

   – The dependent variable is the time taken for the cross to disappear. This time will vary depending on the concentration of sodium thiosulfate.

   – Other variables such as temperature, pressure, and volume should be kept constant to ensure accurate results.

5. Data Collection and Analysis:

   – Data is collected for each trial, recording the time taken for the cross to disappear with each concentration of sodium thiosulfate.

   – A graph is then plotted with concentration on the x-axis and time taken for the cross to disappear on the y-axis.

   – The graph typically shows that as the concentration of sodium thiosulfate increases, the time taken for the cross to disappear decreases. This indicates that higher concentrations lead to faster reaction rates.

6. Conclusion:

   – Based on the results, conclusions are drawn about the effect of concentration on the rate of reaction. Typically, it’s observed that increasing the concentration of sodium thiosulfate leads to a faster rate of reaction, as more reactant particles are available to collide and react in a given time.

Heat Experiments:

The experiment related  to  heat is often known as the “heat experiment” or “thermal decomposition experiment.” It’s commonly performed in chemistry labs to determine the relative atomic mass of a metal in a metal salt.

Here’s a general outline of the procedure:

1. Preparation of the Metal Salt: Start by obtaining a known quantity of the metal salt you want to analyze. This salt is usually a metal carbonate or metal sulfate.

2. Weighing: Weigh the empty crucible and lid accurately using a balance. Then, add a known mass of the metal salt to the crucible and reweigh it to determine the mass of the metal salt.

3. Heating: Place the crucible containing the metal salt on a tripod and heat it gently using a Bunsen burner. The purpose of heating is to decompose the metal salt into its metal oxide and other products.

4. Observation and Collection of Data: As the metal salt decomposes, observe any changes in the appearance of the substance. Often, you’ll see color changes or the release of gases. Record these observations.

5. Cooling and Weighing: Once the heating is complete and no further changes are observed, allow the crucible to cool down. After cooling, reweigh the crucible with its contents.

6. Calculations: Calculate the mass of the metal oxide formed by subtracting the mass of the crucible and any remaining residue from the mass of the crucible and metal salt before heating.

7. Determining the Relative Atomic Mass (RAM): From the mass of the metal oxide formed, you can determine the relative atomic mass of the metal in the metal salt using stoichiometry and the balanced chemical equation for the decomposition reaction.

   For example, if you start with a known mass of metal carbonate (MCO3) and the balanced equation for its decomposition is:

 MCO₃  ® MO + CO₂

where M  represents the metal, then the relative atomic mass of the metal can be calculated using the stoichiometry of the reaction.

This experiment relies on the principle of conservation of mass and stoichiometry. By measuring the mass of the metal oxide formed and knowing the stoichiometry of the reaction, you can determine the relative atomic mass of the metal present in the metal salt.

Organic Chemistry:

Testing for Carbonyl compounds: ketones and aldehydes.

Tollens Reagent:

Tollens reagent is made by mixing AgNO₃ and NH₃. It gives a black precipitate with Aldehyde which has a silvery mirror floating on top. The observation should be that the silver mirror is obtained with Tollens Reagent. Most of the time this silver mirror will not be visible, so black precipitate is enough to test for the presence of aldehyde.

Fehling Solution:

Fehling solution also tests for the presence of Aldehyde. Aldehyde is added to Fehling Solution and heated lightly. A red/brown precipitate is obtained

2,4 DNPH:

 This is an orange-colored solution, which has a strong acid in it. Be careful when using this. 2,4 DNPH has forms a yellow or orange precipitate with carbonyl compounds (both ketones and aldehydes). Remember that anything, when it is added to 2,4 DNPH, will turn yellow because it has a yellow color. So you should be looking for yellow precipitates and ignore the color of the solution.

Carboxylic Acid Test:

If you add Na₂CO₃ and vigorous effervescence is observed, then the compound present is a carboxylic acid.

Potassium Di Chromate:

It will turn from orange to green with alcohols and aldehydes but the mixture has to be gently heated otherwise the color change wouldn’t be visible. If it is strongly heated then aldehyde and alcohols will evaporate. You should also not add Potassium di chromate in excess, as a lot of it will not get reduced and you will get a mixture of green and orange which would be very hard to distinguish.

Potassium Manganate (VII):

If it is added to a solution, and the solution is then warmed in a water bath; if the purple color of the solution disappears then in the solution either an aldehyde or alcohol is present.

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