Qualitative+Analysis+-+Chemical+Detective+Work

Carry out qualitative analysis: An Internal Assessment · //Apply solubility rules and write ionic equations to describe the formation of precipitates and complex ions //
 * //A table of ions will not be provided. //
 * //Appropriate solubility rules and a flow chart to assist in determining the unknown ions ////will be provided. //
 * //The application of solubility rules could involve the prediction of whether a precipitate forms or the identification of any precipitate formed. //
 * //In identifying the ions, it is expected that experimental observations will be reported. //
 * //Ions to be identified will be limited to Ag+, Al3+, Ba2+, Cu2+, Fe2+, Fe3+, Mg2+, Pb2+, Na+, NH4+, Zn2+, Cl //// - ////, CO32 //// - ////, I //// - ////, NO3 //// - ////, OH //// - ////, SO42 //// - ////, Na+ and NO3 //// - //// are identified by a process of elimination. NH4+ will be identified using its reaction with NaOH. //
 * //Complex ions will be limited to those formed when OH-(aq) or NH3(aq) react with cations listed above, such as Ag(NH3)2+, Al(OH)4 //// - ////<span style="font-family: Arial,sans-serif; font-size: 11pt;">, Zn(OH)42 //// - ////<span style="font-family: Arial,sans-serif; font-size: 11pt;">, Zn(NH3)42+, Cu(NH3)42+. //

<span style="font-family: Arial,sans-serif; font-size: 11pt;">The water found in our lakes, rivers and oceans contains many different dissolved substances. Weathering and erosion washes many materials into the waterways and out to sea, making the seawater, in particular, salty. The most abundant ions in seawater are Cl // - //<span style="font-family: Arial,sans-serif; font-size: 11pt;">, Na+, K+, Mg2+, Ca2+, and SO42 // - //<span style="font-family: Arial,sans-serif; font-size: 11pt;">. The concentrations of ions in water near land masses varies more widely than in deep ocean water.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Seawater – what’s in it? **

<span style="font-family: Arial,sans-serif; font-size: 11pt;">The dissolved solids can be isolated by evaporating the water. As salt water is evaporated the least soluble salt, usually calcium carbonate, will crystallise first. If the volume of solution is further reduced, the substance of next lowest solubility will separate, and so on. This is called fractional crystallisation.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Ionic solids contain a 3-dimensional network of positive and negative ions held together by the strong electrostatic attraction between these oppositely charged ions (ionic bonds). The formula for an ionic compound is an empirical formula i.e. it gives the ratio of cations : anions in the network so that overall the crystal is neutral. <span style="font-family: Arial,sans-serif; font-size: 11pt;"> e.g. NaCl has a ratio 1:1 of Na+ : Cl- <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Al2O3 has a ratio of 2:3 of Al3+ : O2 - <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Mg(OH)2 has a ratio 1 : 2 of Mg2+: OH // - // <span style="font-family: Arial,sans-serif; font-size: 11pt;">e.g.calcium nitrate has a formula Ca(NO3)2
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Formulae of Ionic Solids **
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Note **<span style="font-family: Arial,sans-serif; font-size: 11pt;">: Where a formula requires more than 1 of a polyatomic ion then brackets must be used.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Because water has polar bonds and an unsymmetrical shape (see later) the overall molecule has a charge separation. i.e. the H end of the molecule has a partial positive charge and the O end has a partial negative charge. When water comes in contact with the ions in an ionic crystal, each ion is attracted to the oppositely charged end of the water molecules as in the diagram below. <span style="font-family: Arial,sans-serif; font-size: 11pt;">In many cases these new attractions are strong enough to overcome the strong ionic bonds holding the ionic crystal together and it will dissolve. In other cases the new attractions are not strong enough and the ionic solid does not dissolve. i.e. it is insoluble**.**
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Water as a polar solvent **



=<span style="font-family: Arial,sans-serif; font-size: 11pt;">Solubility = <span style="font-family: Arial,sans-serif; font-size: 11pt;">Soluble ionic compounds dissolve readily in water forming a solution containing dissolved ions that are kept apart from each other by the attractions to the water molecules between them. <span style="font-family: Arial,sans-serif; font-size: 11pt;">The solution of a soluble ionic solid will be clear. It may be either colourless or coloured, depending on the particular ions present. Sodium chloride is a soluble solid. <span style="font-family: Arial,sans-serif; font-size: 11pt;"> H2O <span style="font-family: Arial,sans-serif; font-size: 11pt;"> NaCl(//s//) ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Na+(//aq//) + Cl - <span style="font-family: Arial,sans-serif; font-size: 11pt;">(//aq//) =<span style="font-family: Arial,sans-serif; font-size: 11pt;">Precipitation = <span style="font-family: Arial,sans-serif; font-size: 11pt;">Ionic compounds which do not readily dissolve are said to be insoluble e.g. CaCO3. This also means that if a solution containing Ca2+ ions is mixed with a solution containing CO32- ions, a precipitate of CaCO3will form. Only relatively insoluble compounds form precipitates. <span style="font-family: Arial,sans-serif; font-size: 11pt;">We say a precipitate has formed if, on mixing 2 clear solutions, the mixture goes cloudy or solid settles in the bottom of the container. Ca(OH)2 and Ba(OH)2 are slightly soluble. A solution containing partially dissolved Ca(OH)2 is called limewater. <span style="font-family: Arial,sans-serif; font-size: 11pt;">The solubility rules are summarised in the table below: <span style="font-family: Arial,sans-serif;"> || ||= **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Insoluble ** <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">(Form Precipitates) || <span style="font-family: Arial,sans-serif; font-size: 11pt;">e.g. Na2CO3, KOH ||=  || <span style="font-family: Arial,sans-serif; font-size: 11pt;">e.g. (NH4)2SO4 ||=  || <span style="font-family: Arial,sans-serif; font-size: 11pt;">e.g. Cu(NO3)2 ||=  || <span style="font-family: Arial,sans-serif; font-size: 11pt;">e.g. FeSO4 ||= **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Insoluble sulfates **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> are CaSO4 BaSO4, PbSO4 || <span style="font-family: Arial,sans-serif; font-size: 11pt;">e.g. FeCl3, CuBr2 ||= **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Insoluble halides **<span style="font-family: Arial,sans-serif; font-size: 11pt;">are those of Ag+, Pb2+ e.g. AgCl, AgI ||
 * ~ <span style="font-family: Arial,sans-serif;">
 * = **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Soluble ** <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">(Never Precipitate) <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">
 * = <span style="font-family: Arial,sans-serif; font-size: 11pt;">All **Group 1** compounds
 * = <span style="font-family: Arial,sans-serif; font-size: 11pt;">All **ammonium** compounds
 * = <span style="font-family: Arial,sans-serif; font-size: 11pt;">All **nitrate** compounds
 * = <span style="font-family: Arial,sans-serif; font-size: 11pt;">Most **sulfates**
 * = <span style="font-family: Arial,sans-serif; font-size: 11pt;">Most **halides**
 * =  ||= <span style="font-family: Arial,sans-serif; font-size: 11pt;">All **carbonates** e.g. BaCO3

<span style="font-family: Arial,sans-serif; font-size: 11pt;"> ||
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">EXCEPTthose of Group 1 and NH4+ ** ||
 * =  ||= <span style="font-family: Arial,sans-serif; font-size: 11pt;">All **oxides, hydroxides** e.g. CuO, Al(OH)3 **EXCEPTthose of Group 1 and NH4+**
 * =  ||= <span style="font-family: Arial,sans-serif; font-size: 11pt;">All **sulfides** e.g. FeS
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">EXCEPTthose of Group 1 and NH4+ ** ||

<span style="font-family: Arial,sans-serif; font-size: 11pt;">The formation of a precipitate, and its colour, can be used to identify ions present in solution. This identification of species in a solution is called **qualitative analysis**. <span style="font-family: Arial,sans-serif; font-size: 11pt;">Some important points in this practical procedure are given below.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">1) **Reaction of carbonates with acid.** <span style="font-family: Arial,sans-serif; font-size: 11pt;">Metal carbonates react with acid. This means they dissolve in acid giving off carbon dioxide and forming a solution of the metal salt in water. Fizzing on addition of acid is commonly used to identify a carbonate.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">2) **Hydroxide precipitates.** <span style="font-family: Arial,sans-serif; font-size: 11pt;">Many metal ions in solution can be identified by the addition of sodium hydroxide and the formation of the corresponding metal hydroxide precipitate. <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Fe2+(//aq//) + 2OH - <span style="font-family: Arial,sans-serif; font-size: 11pt;">(//aq//) ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Fe(OH)2(//s//) green ppt <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Cu2+(//aq//) + 2OH - <span style="font-family: Arial,sans-serif; font-size: 11pt;">(//aq//) ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Cu(OH)2(//s//) light blue ppt <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Fe3+(//aq//) + 3OH - <span style="font-family: Arial,sans-serif; font-size: 11pt;">(//aq//) ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Fe(OH)3(//s)// red-brown ppt <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Ag+(//aq//) + 2OH - <span style="font-family: Arial,sans-serif; font-size: 11pt;">(//aq//) ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Ag2O(//s//) + H2O(//l//) brown ppt
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Note: **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> The spectator ions e.g. Na+ have been omitted from the following ionic equations.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Hydroxide precipitates of Mg2+, Al3+, Zn2+, Ba2+ and Pb2+ are all white. <span style="font-family: Arial,sans-serif; font-size: 11pt;">(i) Na+ and K+ don’t form precipitates with OH - <span style="font-family: Arial,sans-serif; font-size: 11pt;"> and therefore cannot be identified in this way. <span style="font-family: Arial,sans-serif; font-size: 11pt;">(ii) Ammonium ion, NH4+, can be identified when OH - <span style="font-family: Arial,sans-serif; font-size: 11pt;"> is added because, although no precipitate forms, the NH4+ is converted to NH3 gas which (although soluble) can be detected by its unpleasant smell and the fact it is a basic gas that turns moist pink litmus blue.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Note: **

<span style="font-family: Arial,sans-serif; font-size: 11pt;">3) **Amphoteric hydroxides.** <span style="font-family: Arial,sans-serif; font-size: 11pt;">Almost all insoluble hydroxides are classed as **basic** and will therefore dissolve in acid to form a <span style="font-family: Arial,sans-serif; font-size: 11pt;">solution of the metal ion in water. <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Fe(OH)3(//s//) + 3H+(//aq//) ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Fe3+(//aq//) + 3H2O(//l//) <span style="font-family: Arial,sans-serif; font-size: 11pt;">Those hydroxides which are amphoteric, Al(OH)3 and Zn(OH)2 ,will dissolve in both acid (as above) and base. If excess sodium hydroxide is added the precipitate disappears forming a colourless <span style="font-family: Arial,sans-serif; font-size: 11pt;">solution of a complex ion, [Al(OH)4] - <span style="font-family: Arial,sans-serif; font-size: 11pt;"> or [Zn(OH)4]2 - <span style="font-family: Arial,sans-serif; font-size: 11pt;">. <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Al(OH)3 + OH - ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> [Al(OH)4] - <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Similarly Zn(OH)2 + 2OH - ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> [Zn(OH)4]2 - <span style="font-family: Arial,sans-serif; font-size: 11pt;"> white ppt colourless solution <span style="font-family: Arial,sans-serif; font-size: 11pt;">In contrast the white insoluble solid Mg(OH)2 **does not** form a complex ion with hydroxide and so does not redissolve in excess hydroxide.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">4) **Formation of complex ions with ammonia.** <span style="font-family: Arial,sans-serif; font-size: 11pt;">A solution of ammonia is also basic i.e. it contains OH- ions. This means addition of ammonia solution to a solution containing metal ions that form insoluble hydroxides, will also produce a precipitate as above. There are 2 important cases where addition of excess ammonia causes the precipitate to redissolve forming a solution of a complex ion. These are as follows: <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Cu(OH)2(//s//) + 4NH3(//aq//) ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> [Cu(NH3)4]2+(//aq//) + 2OH - <span style="font-family: Arial,sans-serif; font-size: 11pt;">(//aq//) <span style="font-family: Arial,sans-serif; font-size: 11pt;"> light blue ppt dark blue solution <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Zn(OH)2(//s//) + 4NH3(//aq//) ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> [Zn(NH3)4]2+(//aq//) + 2OH - <span style="font-family: Arial,sans-serif; font-size: 11pt;">(//aq//) <span style="font-family: Arial,sans-serif; font-size: 11pt;"> white ppt colourless solution
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Note: **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> Because Al3+ does not form a complex ion with ammonia this can be used to distinguish between it and Zn2+.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">5) **Identification of chloride ions.** <span style="font-family: Arial,sans-serif; font-size: 11pt;">When silver ions are added to a solution containing chloride ions (or vice-versa) a white precipitate of AgCl forms. This redissolves when excess dilute ammonia solution is added, due to the formation of the colourless, complex ion [Ag(NH3]2+. <span style="font-family: Arial,sans-serif; font-size: 11pt;"> AgCl(//s//) + 2NH3(//aq//) ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> [Ag(NH3)2]+(//aq//) + Cl - <span style="font-family: Arial,sans-serif; font-size: 11pt;">(//aq//) <span style="font-family: Arial,sans-serif; font-size: 11pt;">AgBr is cream coloured and dissolves in concentrated ammonia. AgI is pale yellow and does not redissolve in ammonia.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">6) **Identification of sulfate ions** <span style="font-family: Arial,sans-serif; font-size: 11pt;">The presence of sulfate ions in an aqueous solution can be shown by the formation of a white <span style="font-family: Arial,sans-serif; font-size: 11pt;">precipitate on addition of an aqueous solution of barium nitrate or barium chloride. <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Ba2+(//aq//) + SO42 - <span style="font-family: Arial,sans-serif; font-size: 11pt;">(//aq//) ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> BaSO4(//s//) <span style="font-family: Arial,sans-serif; font-size: 11pt;">This precipitate does not dissolve in acid - as it would if the precipitate was BaCO3 or Ba(OH)2. <span style="font-family: Arial,sans-serif; font-size: 11pt;">BaCO3(//s//) + HCl(//aq//) ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Ba2+(//aq//) + 2Cl - <span style="font-family: Arial,sans-serif; font-size: 11pt;">(//aq//) + CO2(//g//) + H2O

<span style="font-family: Arial,sans-serif; font-size: 11pt;">7) **Identification of Fe3+ ions** <span style="font-family: Arial,sans-serif; font-size: 11pt;">The presence of Fe3+ ions can be detected by the formation of a rust coloured precipitate on addition of base (OH - <span style="font-family: Arial,sans-serif; font-size: 11pt;">). It can be confirmed by addition of a solution of potassium thiocyanate to the original metal ion solution. The Fe3+ reacts with SCN - <span style="font-family: Arial,sans-serif; font-size: 11pt;"> ion to form a blood red solution containing the complex ion FeSCN2+. This is not a precipitate!

<span style="font-family: Arial,sans-serif; font-size: 11pt;"> Fe3+(//aq//) + SCN - <span style="font-family: Arial,sans-serif; font-size: 11pt;">(//aq//) ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> FeSCN2+(//aq//)