RGS 2025 Y4 QA Notes

1 QA Notes (I) – Introduction and Laboratory Techniques Qualitative analysis (QA) is the process of identifying the cations and anions present in unknown substances using chemical reactions and test reagents . Unlike quantitative analysis, where the amount of a particular chemical is measured, a qualitative analysis scheme simply confirms t he presence or absence of certain chemicals or ions. The chemical reactions used to identify the cations and anions mostly involve observable changes including: • production of a precipitate, • evolution of a gas, or • visible colour change 1.1 Preliminary examination Before carrying out any chemical reacti on, a preliminary examination to note the appearance, colour and solubility of the unknown substance often gives clues to the identity of the unknown chemical or the element present in the substance. (a) Colour • In general, compounds of transition metals have characteristic colours. • Group I, II and III metals generally form white salts that dissolve to form colourless solutions. • However, as can be seen from Table 1 , there are several possibilities for a give n appearance. The appearance is not conclusive and you should carry out further chemical tests to arrive at your final conclusion about the identity of the sample. Table 1: Characteristic colours of some chemicals Colour Inferences Colourless Dilute acids , alkalis and solutions of salts of Group I, II and III metals White Solid salts of Na+, K+, NH4+, Ca2+, Zn2+, Pb2+, Al3+ Black CuO, CuS, CoO, FeO, FeS, PbS, MnO2, I2 crystals Grey Metals e.g. Zn, Fe, Al in powder form Dark green Chromium salts Light green Iron(II) salts Blue or bluish green Hydrated copper(II) salts Yellow or brown Solutions of iron(III) salts (Fe3+), PbI2, AgI Pale pink Manganese(II) salts (Mn2+) Purple KMnO4 RAFFLES GIRLS’ SCHOOL (SECONDARY) RAFFLES PROGRAMME YEAR 4 CHEMISTRY Name : _________________________ ( ) Class : _______ Date : _____________ Established in 1879

2 • Crystalline salts sometimes show different colours when in hydrate d and anhydrous states. For example, CuSO4.5H2O is blue and anhydrous CuSO4 is white in colour. (b) Solubility of ionic compounds Below is a summary of the solubility of common substances in water. The following are soluble in water: 1. Salts of NH4+, K+, Na+ 2. Nitrates 3. Hydrogencarbonates 4. Sulfates except PbSO4, BaSO4, CaSO4 (sparingly soluble) 5. Halides including Chlorides Bromides Iodides except PbCl2, AgCl except PbBr2, AgBr except Pbl2, Agl The following are insoluble in water: 1. Carbonates except those of Na+, K+, NH4+ CaO, Ca(OH)2, CaS, CaSO3 are slightly soluble. 2. Hydroxides 3. Oxides 4. Sulfides (S2-) 5. Sulfites (SO32-) 1.2 Types of Reactions involved in QA After a preliminary examination of the app earance and solubility of the unknown substance, chemical reactions such as precipitation reactions, acid -base reactions, thermal decomposition and redox reactions are carried out to identify the unknown chemical or ion(s) present in the substance. In a confirmatory test, each ion has a unique response to an added chemical such as a color change of the solution or the formation of a precipitate. For example, if an acid is added to the solution and a gas that forms a white precipitate in limewater is evolv ed, we can conclude that the solution contains carbonate ions (CO 32- ions). This test described is a confirmatory test for the presence of carbonate ions. Hence, observations specific to certain types of chemical reactions allow us to infer and make conclusions about the identity of specific ions.

3 (a) Precipitation Reaction • Precipitation reaction is one that results in the formation of an insoluble product when two solutions are mixed together. • The insoluble solid formed separates from the solution and is called a precipitate. For example, when we mix lead(II) nitrate solution and potassium iodide solutio n, a yellow precipitate of lead(II) iodide forms. • In the boxes below, draw the changes in the arrangement of the ions when precipitation takes place between lead(II) nitrate and potassium iodide solutions. lead(II) nitrate solution Potassium iodide solution What happens when the two solutions are mixed together? Chemical equation: Pb(NO3)2 (aq) + 2KI (aq) → PbI2 (s) + 2KNO3 (aq) Ionic equation: Pb2+ (aq) + 2I- (aq) → PbI2 (s) (b) Acid-base neutralization reaction (Revision) • An acid is a substance that dissociates in water to produce hydro gen ions . Common laboratory acids include dilute hydrochloric acid, dilute nitr ic acid and dilute sul furic acid. • A base is a substance that reacts with acids to produce a salt and water. Bases can be metal oxides or hydroxides. Examples of bases are iron(II) oxide, iron(II) hydroxide, zinc oxide, zinc hydroxide, copper(II) oxide and copper(II) hydroxides. • An alka li is a substance that dissociates in water to produce hydroxide ions. Common alkalis in the laboratory are sodium hydroxide solution, potassium hydroxide solution and aqueous ammonia.

4 (c) Thermal Decomposition • Thermal decomposition is a r eaction in which a single compound breaks up into two or more simpler compounds or elements when heated. The decompo sition reaction is irreversible. Example 1 Copper(II) carbonate decomposes on heating to form copper (II) oxide and carbon dioxide: CuCO3 (s) → CuO (s) + CO2 (g) Example 2 Zinc nitrate decomposes on heating to give zinc oxide, nitrogen dioxide and oxygen: Zn(NO3)2 (s) → 2ZnO (s) + 4NO2 (g) + O2 (g) (d) Redox reaction – to be covered in next topic 1.3 Laboratory Techniques (a) Quantity of Substances Used There is a common misconception that using large quantities of chemic als for a reaction would produce better results. In fact, using excessive amounts of chemicals could lead to a waste of time as it takes a longer time for the reagents to react. It may even lead to wrong results e.g. if too much of salt is present, its solubility in excess reagent will not be observed. Students are recommended to use the following guidelines, unless otherwise specified: i. If tests are carried out with the unknown in its solid form, the amount used should be just enough to fill the hemisphere at the bottom of a test tube. ii. If tests are carried out with a solution of the unknown, a stock solution can first be prepared. The stock solution may be prepared by filling the bottom hemisphere of a test tube with the solid and dissolving it in half a test tube of distilled water . For each test, fill the test tube to a depth of approximately 1 cm with the stock solution. Why is tap water not used to prepare stock solution? Tap water may contain ions such as Ca 2+, Mg 2+, NO3-, SO42- and Cl - and using it to prepare stock solution could lead to ‘false-positive’ results.

5 (b) Handling of Reagents • Use a spatula to transfer solid chemicals. Avoid spilling on your hands, clothes or table tops as some chemicals cause stains or burns. • Before using a solution from a reagent bottle, read the label to avoid using the wrong reagent. Hold the bottle on the side of the label to avoid spilling the solution ove r the label. Then pour the solution carefully into the test tube. Always remember to replace the stopper after the solution has been used. • It is essential that all reagents should be free from contamination. In transferring a liquid reagent from a dropper into the test tube containing the solution of the unknown, hold the dropper just above the mouth of the test tube and allow the reagent to drop into the test tube. Do not leave the dropper on the bench or allow the dropper tip to touch anything except for the reagent itself.