JK-TK/Newdata huggingface dataset:
The liquid does not return to the solid state unless a catalyst is used. The specific catalysts are not usually available in school laboratories. GENERAL EXAMINATION REQUIREMENTS
Apparatus List
This list given below has been drawn up in order to give guidance to schools concerning the apparatus that is expected to be generally available per student for examination purposes. The list is not intended to be exhaustive: in particular, items (such as Bunsen burners, tripods) that are commonly regarded as standard equipment in a chemical laboratory are not included in this list. One burette, 50 cm3
A measuring cylinder, 50 cm³ or 25 cm³
A filter funnel
One pipette, 25cm3
Two conical flasks within the range 150cm3 0r 250cm3
A beaker, squat form with lip: 250 cm³
A thermometer, -10°C to + 110°C at 1 °C
A polystyrene or other plastic beaker of approximate capacity 150 cm³
Clocks (or wall-clock) to measure to an accuracy of about 1s. (Where clocks are specified, candidates may use their own wristwatch if they prefer)
Wash bottle
Test tubes (some of which should be Pyrex or hard glass). Approximately
125 mm x 16 mm
Boiling tubes, approximately 150 mm x 25 mm
Stirring rod
ACCURACY
Unless a question instructs candidates differently they should assume that readings from equipment and apparatus ought to be made with the following precision:
Burette readings should be to the nearest 0.05 cm3
Weighings should be made to 0.01 g or 0.001 g depending on the precision of the balance
0 to 1000C thermometers should be read to the nearest 0.50C and 0 to 500C thermometers to the nearest 0.20C
Timers will normally be read to the nearest second. It is important that when candidates record reading they include the appropriate number of decimal places. For example a burette reading of exactly 24.7 cm3 should be recorded in a results table as 24.70cm3. A temperature of reading of exactly 350C should be recorded as 35.50C.
It is important that when candidates record reading they include the appropriate number of decimal places. For example a burette reading of exactly 24.7 cm3 should be recorded in a results table as 24.70cm3. A temperature of reading of exactly 350C should be recorded as 35.50C. When titres have to be averaged, it is important that the mean is expressed to either the nearest 0.05cm3 or to the second decimal point. For example, if a candidate records four different titres as listed below, the mean can be worked out. 26.50cm3 26.25cm3 26.60cm3 26.65cm3
The candidate is expected to ignore the second titre and average the remaining three
26.50 + 26.60 + 26.65 = 26.583 this should be recorded as 26.60cm3
3
(to the nearest 0.05cm3)
In general, a final should always be given to the same number of significant figures as is suggested in the exercise. A significant proportion of marks for quantitative exercise will be awarded for accuracy. The marks will be awarded by comparing the candidate’s results with the teacher’s reports. CALCULATIONS
Usually calculations will be structured. Candidates will be led through a series of steps leading to a final value. Since most of the marks for these steps will be for a correct method rather than the numerical answer, it is important that candidates include their working even if this seems to be trivial. No marks can be awarded for an incorrect answer without working but a correct method followed by an incorrect answer will receive credit. GRAPHS
Some exercises in practical chemistry will require candidates to treat their readings graphically. The question will however instruct the candidate which axes to use for each quantity being plotted. Some useful points for candidates to keep in mind when constructing graphs are listed below. Candidates will normally be instructed to put the dependent variable, the quantity being measured e.g. temperature on the y- axis. The predetermined quantity e.g. volume will be on the x-axis. The scales should be chosen so that the results are spread out as far apart as the size of the grid allows but not at the expense of using a sensible scale. For example using 1cm to represent 3 units might spread the readings better than using 1 cm to represent 4 units but the scale may be difficult to read. It is always advisable ton use even scales.
The scales should be chosen so that the results are spread out as far apart as the size of the grid allows but not at the expense of using a sensible scale. For example using 1cm to represent 3 units might spread the readings better than using 1 cm to represent 4 units but the scale may be difficult to read. It is always advisable ton use even scales. The origin (0,0) need not necessarily be included on either scale if it is not relevant. For example if temperature readings between 210C and 280C are plotted, there is no need to begin the axes at zero. The axes must be clearly labelled with the quantity being plotted e.g. mass and its units e.g. kilograms
The points plotted may be joined with a straight line or a smooth curve. Since readings are all subject to experimental error, the line drawn may not necessarily pass through every point. Points should never be joined by a series of short straight lines. Doing a Titration
Begin by preparing your burette, as described on the burette page. Your burette should be conditioned and filled with titrant solution. You should check for air bubbles and leaks, before proceeding with the titration.Take an initial volume reading and record it in your notebook. Before beginning a titration, you should always calculate the expected endpoint volume. Prepare the solution to be analysed by placing it in a clean Erlenmeyer flask or beaker. If your sample is a solid, make sure it is completely dissolved. Put a magnetic stirrer in the flask and add indicator. Use the burette to deliver a stream of titrant to within a couple of mL of your expected endpoint. You will see the indicator change colour when the titrant hits the solution in the flask, but the colour change disappears upon stirring.Approach the endpoint more slowly and watch the colour of your flask carefully. Use a wash bottle to rinse the sides of the flask and the tip of the burette; to be sure all titrant is mixed in the flask. As you approach the endpoint, you may need to add a partial drop of titrant. You can do this with a rapid spin of a Teflon stopcock or by partially opening the stopcock and rinsing the partial drop into the flask with a wash bottle. Ask your TA to demonstrate these techniques for you, in the lab.Make sure you know what the endpoint should look like. For phenolphthalein, the endpoint is the first permanent pale pink.
You can do this with a rapid spin of a Teflon stopcock or by partially opening the stopcock and rinsing the partial drop into the flask with a wash bottle. Ask your TA to demonstrate these techniques for you, in the lab.Make sure you know what the endpoint should look like. For phenolphthalein, the endpoint is the first permanent pale pink. The pale pink fades in 10 to 20 minutes.If you think you might have reached the endpoint, you can record the volume reading and add another partial drop. Sometimes it is easier to tell when you have gone past the endpoint. When you have reached the endpoint, read the final volume in the burette and record it in your notebook. Subtract the initial volume to determine the amount of titrant delivered. Use this, the concentration of the titrant, and the stoichiometry of the titration reaction to calculate the number of moles of reactant in your analyte solution. Indicators
An indicator is a substance used in titrations, which has one colour in the presence of an excess of one reagent and a different colour in the presence of an excess of the other. Examples of indicators include;
Methyl orange
Phenolphthalein
Methyl red
Bromothymol blue
The pH scale
A much more useful measure of the strength of an acid solution was worked out by the Danish biochemist S. Sorensen. He worked in the laboratories of the Carlsberg breweries and was interested in checking the acidity of beer. The scale he introduced was the pH scale. The scale runs from 1 to 14, and the following general rules apply. Rules for the pH scale:
Acid have a pH less than 7,
The more acidic a solution, the lower the pH. Neutral substances, such as pure water, have a pH of 7,
Alkalis have a pH greater than 7,
The more alkaline a solution, the higher the pH. The pH of a solution can be measured in several ways. Universal indicator papers that are sensitive over the full range of values can be used. Some colour changes of common indicators are shown below. Indicator colour changes
Candidates may be asked to carry out exercises involving:
Simple quantitative experiments involving the measurement of volumes:
Speeds of reaction. Measurement of temperature based on a thermometer with 10C graduations. Problems of an investigatory nature, possibly including suitable organic compounds. Filtration.
Measurement of temperature based on a thermometer with 10C graduations. Problems of an investigatory nature, possibly including suitable organic compounds. Filtration. Identification of ions and gasses as specified in the curriculum. HEALTH AND SAFETY
Candidates must follow the health and Safety policy normally operates in their laboratories when carrying out the practical Examination. Eye protection must always be worn. Laboratory overalls are recommended. All substances should be regarded as being potentially toxic and hazardous. Hazard labels (e.g. flammable) should be read and appropriate precautions (e.g. keep liquid away from flame) taken. All substances spilled on the skin should be rinsed off immediately. Chemicals must never be tasted Gases and vapours should never be smelt unless the question instructs the candidates to do so and then only with great care. 2.0 QUANTITATIVE ANALYSIS
2.1 Molar solutions and volumetric analysis. Amount of substance
In quantitative analysis, it is important to measure the amounts or moles of reacting substances accurately. To do that, we must use values of relative atomic masses expressed on a periodic table. If we are given the mass of a compound, we can determine the number of moles. The first step is to calculate the mass of one mole of the compound by summing up the relative atomic masses of the constituent atoms. E.g. 1 mole of calcium carbonate (CaCO3) will have a mass of 40 + 12 + 48 = 100g
10g of CaCO3 will contain 10/100 = 0.1moles
concentration of a solution
It is very useful to be know exactly how much of a dissolved substance is present in a solution of particular concentration or volume of a solution. So we need a standard way of comparing the concentrations of solutions. The concentration of a solution is determined from the number of moles of solute dissolved in one litre (1dm3) of solution. If one mole of a solute is dissolved in water and the volume of the solution made up to 1litre (1dm3), this solution in known as a molar solution or 1M solution. If two moles are dissolved in 1litre (1dm3), the resulting solution is 2M and so on.
The concentration of a solution is determined from the number of moles of solute dissolved in one litre (1dm3) of solution. If one mole of a solute is dissolved in water and the volume of the solution made up to 1litre (1dm3), this solution in known as a molar solution or 1M solution. If two moles are dissolved in 1litre (1dm3), the resulting solution is 2M and so on. You need to be able to calculate
(i) The number of moles or mass of substance in an aqueous solution of given volume and concentration
(ii) The concentration of an aqueous solution given the amount of substance and volume of water. (iii) Use the equation: molarity of Z = moles of Z / volume in dm3
Remember: moles Z = mass Z / formula mass of Z
Concentration can also be expressed directly in grams per litre (1dm3). 9.8g of H2SO4 can be expressed either as 0.1M or 9.8g/dm3
Molarity (M) = g/dm3 / formula mass
Example 1: 5.95g of potassium bromide were dissolved in 400cm3 of water. Calculate its molarity. [Ar values: K = 39, Br = 80]
Moles = mass / formula mass, (KBr = 39 + 80 = 119)
mol KBr = 5.95/119 = 0.05 mol
400 cm3 = 400/1000 = 0.4 dm3
molarity = moles of solute / volume of solution
molarity of KBr solution = 0.05/0.4 = 0.125M
2.3 Volumetric calculations (acid-alkali titrations)
Chemical Equations
These balancing numbers have an additional meaning where these amount calculations are concerned. The balancing numbers give the numbers of moles present for each chemical involved in the reaction. For example, sodium chloride may be prepared by the reaction of sodium hydroxide and hydrochloric acid, according to the following equation -
NaOH (aq) + HCl (aq) NaCl (aq) + H2O (l)
The equation specifies what amounts of sodium hydroxide and hydrochloric acid will react together and what amounts of products are produced.
[Ar values: K = 39, Br = 80]
Moles = mass / formula mass, (KBr = 39 + 80 = 119)
mol KBr = 5.95/119 = 0.05 mol
400 cm3 = 400/1000 = 0.4 dm3
molarity = moles of solute / volume of solution
molarity of KBr solution = 0.05/0.4 = 0.125M
2.3 Volumetric calculations (acid-alkali titrations)
Chemical Equations
These balancing numbers have an additional meaning where these amount calculations are concerned. The balancing numbers give the numbers of moles present for each chemical involved in the reaction. For example, sodium chloride may be prepared by the reaction of sodium hydroxide and hydrochloric acid, according to the following equation -
NaOH (aq) + HCl (aq) NaCl (aq) + H2O (l)
The equation specifies what amounts of sodium hydroxide and hydrochloric acid will react together and what amounts of products are produced. 1NaOH (aq) + 1HCl (aq) 1NaCl (aq) + 1H2O (l)
This equation states that 1 mol of sodium hydroxide and 1 mol of hydrochloric acid will react together to give of 1 mol sodium chloride and of 1 mol water. The ratio of NaOH: HCl: NaCl: H2O is 1: 1: 1: 1. What this means is that if the amount of any one of the components in the above reaction is known then the others may be worked out from the ratio above. Or, if the masses of chemicals reacting together are known then amounts can be calculated and the balancing numbers deduced from the amounts reacting together. 2.4 Titration Experiment
Titrations can be used to find the concentration of an acid or alkali from the relative volumes used and the concentration of one of the two reactants. Titration involves the neutralisation of an acid with an alkali or a soluble carbonate. Since both reactants and products are colourless, an indicator is used to find the neutralisation point or end point i.e. the point at which the acid has been neutralised. This process is done in three stages. The acid is poured into the burette.
the point at which the acid has been neutralised. This process is done in three stages. The acid is poured into the burette. A known volume of the alkali is transferred using a pipette into four separate conical flasks. Two or three drops of suitable indicator are added to each of the conical flasks. The acid is run into the flask until the indicator just changes colour. The difference between the initial and final burette readings gives the volume of the acid used commonly known as the titre. The first titration usually gives and approximate end point and is treated as the trial. The experiment is repeated for each of the other conical flask, to try to obtain the end point accurately. This is usually done by running out the acid to a point one unit away from the trial, the going drop-by-drop until an accurate end point is obtained. The volume of the acid required for each change is read off and recorded in a table similar to the one below. PROCEDURE DURING TITRATION
Titrations require continuous shaking of the conical flask and its contents. Check the quantity of the pipette as indicated on the bulb of the pipette and remember to record it in the appropriate place. As you release the liquid from a pipette into a conical flask, one should not blow out the last drop remaining in the jet. Before you take any reading from the burette after filling it, first allow the solution to run out to fill the tap and jet of the burette, and then you begin taking your readings. Always write the reading immediately you take them. When completing the table of results, you will be expected to complete all columns as accurately as the as the limits of the apparatus can allow e.g. burette used is usually read to the nearest read to the nearest 0.05 cm 3, pipette is accurate to 0.05 cm 3 (1 drop). Use minimum amount of indicator possible (2 – 3 drops) and recognize the end point has been reached e.g. when the colour just changes. Show the values that can be averaged to obtain an acceptable value for use in calculations (only those values within 0.2 cm 3 should be averaged). Attempt to work out the questions from the first principles and not use the formula method, which has its own limitations.
when the colour just changes. Show the values that can be averaged to obtain an acceptable value for use in calculations (only those values within 0.2 cm 3 should be averaged). Attempt to work out the questions from the first principles and not use the formula method, which has its own limitations. GENERAL NOTES FOR QUALITATIVE AND QUANTITIVE ANALYSIS
Indicators
An indicator is a substance used in titrations, which has one colour in the presence of an excess of one reagent and a different colour in the presence of an excess of the other. Examples of indicators include;
Methyl orange
Phenolphthalein
Methyl red
Bromothymol blue
Indicator colour changes
Candidates may be asked to carry out exercises involving:
Simple quantitative experiments involving the measurement of volumes:
Speeds of reaction. Measurement of temperature based on a thermometer with 10C graduations. Problems of an investigatory nature, possibly including suitable organic compounds. Filtration. Identification of ions and gasses as specified in the curriculum. QUALITATIVE ANALYSIS
Qualitative analysis is mainly about identification of substances. No emphasis is laid on amount or quantity. In qualitative analysis exercises, candidates should use approximately 1cm depth of a solution (1-2cm3) for each test and add reagents slowly, ensuring good mixing, until no further change is seen. Candidates should indicate at what stage a change occurs, writing any deductions alongside the observations on which they are based. Answers should include details of colour changes and precipitates formed and the names and chemical tests for any gases evolved. Marks for deductions or conclusions can only be gained if the appropriate observations are recorded. QUALITATIVE ANALYSIS NOTES
TABLE 1:TESTING FOR GASES
TABLE 2: TEST FOR ANIONS
TABLE 3: TESTS FOR CATIONS
TABLE 4: ORGANIC TESTS
TABLE 5: MISCELLANEOUS TESTS
TABLE 6. FLAME TESTS
A tungsten wire loop is first dipped into some concentrated hydrochloric acid to dissolve any oxides and hence clean the wire. It is then dipped into salt powder and introduced into a colourless Bunsen burner flame. The colour is then observed and inferences made.
FLAME TESTS
A tungsten wire loop is first dipped into some concentrated hydrochloric acid to dissolve any oxides and hence clean the wire. It is then dipped into salt powder and introduced into a colourless Bunsen burner flame. The colour is then observed and inferences made. TABLE 7: HEAT ANALYSIS
TABLE 8: SOLUBLE AND INSOLUBLE SALTS
TABLE 9: GENERAL PRELIMINARY TESTS
TABLE 10 : IGNITION ANALYSIS
Candidates may be asked to heat an unknown alone in an ignition tube. The colour changes or the identity of any gases evolved may provide evidence as to the identity of the unknown. TABLE 10: RESIDUE WHEN HOT AND COLD
Residue remaining after ignition may have a different colour when hot and cold
TABLE 11: HEATING A SOLID
Gases or vapours may be evolved during heating of the solid
TABLE 12: SOLUBILITY
The patterns of solubility for various types of salts
KEY
Ppt = precipitate
Sol = solution
Insol = insoluble
Xs = excess
Dil = dilute
1.HYDROGEN CHLORIDE
Physical Properties
Chemical Properties
Footnotes:
Dilute hydrochloric acid is one of the three common dilute acids used in the laboratory. Concentrated hydrochloric acid is used in the manufacture of many chemicals. 2.CARBON DIOXIDE
Physical Properties
Chemical Properties
Footnotes:
Makes up 0.04% of the gases in the air. Rainwater is slightly acidic because carbon dioxide dissolves in it. Needed by green plants during photosynthesis. Product of respiration in living things. Product of combustion
Is one of the "greenhouse" gases present in the atmosphere. Used in fire extinguishers since it is heavier than air and forms a "blanket" around the fire. This prevents combustion since carbon dioxide does not support it and the presence of carbon dioxide stops oxygen reaching the combustible material. 3.CHLORINE GAS
Physical Properties
Chemical Properties
Footnotes:
One of the "family" of halogen gases (iodine and bromine are in the same family)
Used in water purification. Used to make bleaching powder, disinfectants and antiseptics
Also used to make some explosives, poison gases and pesticides.
This prevents combustion since carbon dioxide does not support it and the presence of carbon dioxide stops oxygen reaching the combustible material.

3.CHLORINE GAS
Physical Properties
Chemical Properties
Footnotes:
One of the "family" of halogen gases (iodine and bromine are in the same family)
Used in water purification.

Used to make bleaching powder, disinfectants and antiseptics
Also used to make some explosives, poison gases and pesticides.

4.HYDROGEN CHLORIDE GAS
Physical Properties
Chemical Properties
Footnotes:
Dilute hydrochloric acid is one of the three common dilute acids used in the laboratory.

Concentrated hydrochloric acid is used in the manufacture of many chemicals.

5.HYDROGEN GAS
Physical Properties
Chemical Properties
Footnotes:
The lightest gas known.

Once used in airships but replaced by helium which is not explosive.

Used to make ammonia which is needed in the manufacture of fertilizers and explosives.

6.NITROGEN
Physical Properties
Chemical Properties
Footnotes:
Makes up around 79% of the gases in the air.

Used to make ammonia gas, which in turn is used to make explosives and fertilizers.

Used in light bulbs and thermometers because it is not reactive.

Also used as the atmosphere in rooms where explosives are stored.

7.OXYGEN
Physical Properties
Chemical Properties
Footnotes:
Makes up 20% of the gases in the air.

Needed by the majority of living organisms for respiration
Is produced by green plants as a by-product of photosynthesis.

8.SULPHUR DIOXIDE GAS
Physical Properties
Chemical Properties
Footnotes:
It is used as a bleaching agent.

It is one of the gases responsible for air pollution.
18.0.0 ACIDS, BASES AND SALTS (25 LESSONS) A.ACIDS AND BASES At a school laboratory: (i)An acid may be defined as a substance that turn litmus red. (ii)A base may be defined as a substance that turn litmus blue. Litmus is a lichen found mainly in West Africa. It changes its colour depending on whether the solution it is in, is basic/alkaline or acidic.It is thus able to identify/show whether 1. An acid is a substance that dissolves in water to form H+/H3O+ as the only positive ion/cation. This is called the Arrhenius definition of an acid. From this definition, an acid dissociate/ionize in water releasing H+ thus: HCl(aq) -> H+ (aq) + Cl- (aq) HNO3(aq) -> H+ (aq) + NO3- (aq) CH3COOH(aq) -> H+ (aq) + CH3COO-(aq) H2SO4(aq) -> 2H+ (aq) + SO42-(aq) H2CO3(aq) -> 2H+ (aq) + CO32-(aq) H3PO4(aq) -> 3H+ (aq) + PO43-(aq) 2.A base is a substance which dissolves in water to form OH- as the only negatively charged ion/anion. This is called Arrhenius definition of a base. From this definition, a base dissociate/ionize in water releasing OH- thus: KOH(aq) -> K+(aq) + OH-(aq) NaOH(aq) -> Na+(aq) + OH-(aq) NH4OH(aq) -> NH4+(aq) + OH-(aq) Ca(OH)2(aq) -> Ca2+(aq) + 2OH-(aq) Mg(OH)2(aq) -> Mg2+(aq) + 2OH-(aq) 3. An acid is a proton donor. A base is a proton acceptor. This is called Bronsted-Lowry definition of acids and bases. From this definition, an acid donates H+ . H+ has no electrons and neutrons .It contains only a proton. Examples I.