Inorganic analysis

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					Unit 3: Laboratory chemistry I

Inorganic analysis
Solubility in water
• All compounds of group I metals and ammonium salts are soluble • All nitrates, ethanoates, hydrogencarbonates and hydrogensulphates (i.e. salts with large singly-charged anions) are soluble • Gp I (hydr)oxides and Gp II (hydr)oxides (increasingly well down the group) dissolve to form alkaline solutions • Group II sulphates become less soluble down the group

Cation testing
Flame test Dip a nichrome wire in concentrated hydrochloric acid and then place in the edge of a medium Bunsen flame. Repeat until the wire gives no colour to the flame. Dip the wire into HCl and then into a sample of the unknown solid before placing in the flame again and note any colour imparted to the flame. Metal ion Lithium (Li ) 2+ Calcium (Ca )
+

Flame colour Crimson red Brick red

Metal ion Sodium (Na ) 2+ Strontium (Sr )
+

Flame colour Yellow-orange Crimson red

Metal ion Potassium (K ) 2+ Barium (Ba )
+

Flame colour Lilac Apple green

Action of sodium hydroxide solution Dissolve a little of the unknown solid to make a solution. To this solution, add a few drops of sodium hydroxide solution and note any observation. Then add an excess of NaOH(aq) and note any change. Most metal hydroxides are insoluble, so the appearance of a precipitate (ppt) serves to identify the cation. e.g. MgSO4(aq) + 2NaOH(aq) → Mg(OH)2(s) + Na2SO4(aq) 2+ in ionic form Mg (aq) + 2 OH (aq) → Mg(OH)2(s) Amphoteric metal hydroxides re-dissolve in excess NaOH(aq). Add a few drops of dilute sodium Add an excess of dilute sodium Cation hydroxide solution hydroxide solution 2+ Magnesium (Mg ) White ppt No further change 2+ Calcium (Ca ) Possible cloudy white ppt No change 3+ Iron(III) (Fe ) Brown ppt No further change 2+ Iron(II) (Fe ) Dull green ppt No further change 2+ Copper (Cu ) Bright blue ppt No further change 2+ Zinc (Zn ) 3+ White jelly-like ppt Ppt re-dissolves (amphoteric) Aluminium (Al ) 2+ Lead (Pb ) + Sodium (Na ) No change No change + Potassium (K ) No change No change + Ammonium (NH4 ) Ammonia gas given off on warming (see Tests for Gases)

Anion testing
Most tests for anions depend on the formation of an insoluble precipitate containing the anion. Anion Test Result 2Carbonate (CO3 ) and To the solid or a solution of the Effervescence producing a colourless hydrogencarbonate unknown add a few drops of dilute gas (CO2, see Tests for Gases) which turns lime water cloudy (HCO3 ) hydrochloric acid White ppt of silver chloride – soluble on Chloride (Cl ) addition of dilute NH3(aq) (bung/shake) To a solution of the unknown add 2-3 drops of dil. nitric acid (to react with any Cream ppt of silver bromide – soluble Bromide (Br ) carbonate present) followed by 2-3 on shaking with conc NH3(aq) drops of silver nitrate solution Yellow ppt of silver iodide– insoluble on Iodide (I ) shaking with conc. NH3(aq) To a solution of the unknown add 2-3 White ppt of barium sulphate drops of dil. HCl (to react with any N.B. barium carbonate is also insoluble 2Sulphate(VI) (SO4 ) carbonate) followed by 2-3 drops of but would effervesce when the acid was barium chloride solution added To the solid or a solution of the Colourless acidic gas given off (SO2) 2Sulphate(IV) (SO3 ) 3 unknown add 1 cm of dil. HCl and which turns acidified potassium (or sulphite) warm dichromate green (see Tests for Gases) To a solution of the unknown add a little Colourless gas with a characteristic aq. NaOH, some powdered aluminium smell and which turns moist red litmus Nitrate(V) (NO3 ) (or Devarda’s alloy) and warm gently paper blue given off (NH3) Effervescence to give a brown acidic Nitrate(III) (NO2 ) To the solid or a concentrated solution gas (NO2, see Tests for Gases) and a 3 (or nitrite) of the unknown add 1 cm of dil. HCl pale blue solution
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Tests for gases
Gas Oxygen (O2) Hydrogen (H2) Carbon dioxide (CO2) Appearance Colourless Odourless Colourless Odourless Courless Odourless Colourless Characteristic smell (smelling salts, window cleaner) Colourless Fumes in moist air Acrid Pale green/yellow ‘Antiseptic’ smell Red-brown fumes condensing to liquid ‘Antiseptic’ smell Purple fumes condensing to black solid Red/brown Acrid Test Insert a glowing splint into the mouth of the test tube Hold a lighted splint over the mouth of the test tube Bubble the gas through a small amount of lime water (calcium hydroxide solution) i) place a piece of damp red litmus paper over the mouth of the test tube ii) place a bottle of conc. HCl near the mouth of the test tube Place a bottle of conc. NH3 near the mouth of the test tube Place a piece of damp blue litmus paper over the mouth of the test tube Place a piece of damp starchiodide paper over the mouth of the test tube Place a piece of damp blue starch paper over the mouth of the test tube Place a piece of damp blue litmus paper over the mouth of the test tube i) Place a piece of damp blue litmus paper over the mouth of the test tube ii) Place a piece filter paper soaked in acidified potassium dichromate solution over the mouth of the test tube Place a piece filter paper soaked in lead ethanoate solution over the mouth of the test tube Result Splint re-lights Splint ‘pops’ as hydrogen ignites in air Lime water turns ‘cloudy’ due to formation of sparingly soluble calcium hydroxide i) red litmus turns blue

Ammonia (NH3)

ii) white clouds of ammonium chloride formed White clouds of ammonium halide formed Blue litmus paper turns red (briefly) and then white (bleached) Red-brown fumes turn white paper blue-black (Br2 oxidises I to I2) Purple fumes turn white paper blue-black Red/brown gas turns blue litmus paper red i) Colourless gas turns blue litmus paper red ii) Orange dichromate paper turns green due to reduction of 23+ Cr2O7 to Cr White paper turns black due to formation of insoluble lead(II) sulphide

Hydrogen halides (HCl, HBr, HI) Chlorine (Cl2) Bromine (Br2) Iodine (I2) Nitrogen dioxide (NO2)

Sulphur dioxide (SO2)

Colourless Acrid

Hydrogen sulphide (H2S)

Colourless ‘Bad eggs’ smell

Action of heat on solids
Group I and Group II carbonates Group II carbonates decompose on strong enough heating to form the metal oxide and carbon dioxide. Thermal stability increases down a group (as decomposition is favoured when there is a highly polarising cation to polarise 2the CO3 ion), thus MgCO3 decomposes at low temperature whilst BaCO3 is highly resistant to decomposition. MgCO3(s) → MgO(s) + CO2(g) + Group I carbonates (with the exception of the most polarising, Li ) do not decompose on heating. Li2CO3(s) → Li2O(s) + CO2(g) Group I and Group II nitrates There is a similar trend in thermal stability to the carbonates (for the ‘polarising power’ argument) Group II nitrates decompose with increasing difficulty down the group to form the metal oxide, nitrogen dioxide and oxygen (see Tests for Gases). 2 Mg(NO3)2(s) → 2 MgO(s) + 4 NO2(g) + O2(g) In Group I, lithium nitrate decomposes in a similar manner (i.e. forming NO2 and O2) whilst the others decompose with increasing difficulty down the group to form the metal nitrite and oxygen gas (see Tests for Gases). 2 KNO3(s) → 2 KNO2(s) + O2(g) Hydrated nitrates may initially dissolve in their own water of crystallisation which subsequently evaporates.

Action of concentrated sulphuric acid on solid Group I halides
• Chlorides give of hydrogen chloride (see Tests for Gases) as conc. H2SO4 is not reduced by Cl • Bromides give of fuming hydrogen bromide, red-brown bromine gas and SO2 (see Tests for Gases) o H2SO4 (S = +6) is reduced to SO2 (+4) whilst Br (-1) is oxidised to Br2 (0) • Iodides give of hydrogen iodide, black I2 solid (purple vapour on heating) with SO2 and H2S (see Tests for Gases) o H2SO4 (S = +6) is reduced to SO2 (+4), S8 (0) and H2S (-2) whilst I (-1) is oxidised to I2 (0)
30/01/2007
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Description: Inorganic analysis