Chapter 12 Salts (MC2013) PDF

CHAPTER 12 Salts © 2013 Marshall Cavendish International (Singapore) Private Limited Chapter 12 Salts 12.1 Salts 12.2 Preparing Salts 12.3 Qualitative Analysis 2 12.1 Salts Learning Outcomes At the end of this section, you shou...

CHAPTER 12 Salts © 2013 Marshall Cavendish International (Singapore) Private Limited Chapter 12 Salts 12.1 Salts 12.2 Preparing Salts 12.3 Qualitative Analysis 2 12.1 Salts Learning Outcomes At the end of this section, you should be able to: define a salt; describe the general rules for the solubility of common salts in water. 3 12.1 Salts What is a Salt? A salt is a compound formed when the hydrogen ion in an acid is replaced by a metallic ion or an ammonium ion. H+ E.g. Na+, K+, Mg2+, Zn2+, Al3+ NH4+ 4 12.1 Salts What is a Salt? A salt is made up of two parts. CATION ANION metallic ion non-metallic ion ammonium ion comes from the comes from the acid base Chloride from HCl Nitrate from HNO3 Sulfate from H2SO4 5 12.1 Salts Reactions in Which Salts Can Be Made Example 1: Zinc sulfate Zn(OH)2(s) + H2SO4(aq) → ZnSO4(aq) + 2H2O(l) ZnSO4 cation anion 2+ 2– Zn SO4 Comes from the base, Comes from the acid, Zn(OH)2 H2SO4 6 12.1 Salts Reactions in Which Salts Can Be Made Example 2: Sodium chloride NaOH + HCl → NaCl + H2O ________ H+ of HCl is replaced by ________. Na+ 7 12.1 Salts Reactions in Which Salts Can Be Made Possible reactants Salt formed Metal/ Carbonate Acid (for the cation) (for the anion) zinc (Zn) hydrochloric acid zinc chloride (HCl) (ZnCl2) copper(II) carbonate nitric acid (HNO3) copper(II) nitrate (CuCO3) (Cu(NO3)2) magnesium oxide sulfuric acid (H2SO4) magnesium sulfate (MgO) (MgSO4) aqueous ammonia sulfuric acid (H2SO4) ammonium sulfate (NH3(aq)) ((NH4)2SO4) potassium hydroxide phosphoric acid potassium (KOH) (H3PO4) phosphate (K3PO4) 8 12.1 Salts Water of Crystallisation Water is present in the crystals of certain compounds. Gives a compound its crystalline properties Easily removed by heating heat hydrated salt → anhydrous salt + water 9  12.1 Salts Water of Crystallisation Example: Copper(II) sulfate Anhydrous salts do Hydrated salts water removed not contain water of contain water of by heating crystallisation. crystallisation. white powder, blue crystals, copper(II) sulfate hydrated copper(II) sulfate add water CuSO4 CuSO4.5H2O The amount of water crystallised is indicated 10 after the dot ‘.’ in its chemical formula. 12.1 Salts Solubility of Salts All salts containing Na+, K+, and NH4+ ions are soluble in water. E.g. NaCl, K2CO3 and (NH4)2SO4. All nitrates are soluble. E.g. Pb(NO3)2, KNO3, Zn(NO3)2 and AgNO3. 11 12.1 Salts Solubility of Salts All chlorides are soluble except PbCl2 and AgCl. All sulfates are soluble except BaSO4, CaSO4 and PbSO4. All carbonates are insoluble except Na2CO3, K2CO3 and (NH4)2CO3. 12 12.1 Salts Solubility of Salts (Summary) Soluble salts Insoluble salts All nitrates All sodium salts All potassium salts except All carbonates All ammonium salts Silver chloride, AgCl All chlorides except Lead(II) chloride, PbCl2 Barium sulfate, BaSO4 All sulfates except Lead(II) sulfate, PbSO4 Calcium sulfate, CaSO4 13 Chapter 12 Salts 12.1 Salts 12.2 Preparing Salts 12.3 Qualitative Analysis 14 12.2 Preparing Salts Learning Outcomes At the end of this section, you should be able to: suggest a suitable method and the starting materials for preparing a salt; describe methods of separation and purification used in preparing salts. 15 12.2 Preparing Salts Before preparing a salt, consider: 1. Is the salt soluble in water? 2. Are the starting materials (reactants) soluble in water? 16 12.2 Preparing Salts 1. Salt is soluble 2. Starting material is insoluble Method 1 (Soluble salt): Reaction of acid + insoluble metal Reaction of acid + insoluble base Reaction of acid + insoluble carbonate 17 12.2 Preparing Salts Method 1 E.g. Mg, Al, Zn, Fe acid + metal E.g. MgO, ZnO, CuO, acid + insoluble base Fe(OH)3, Cu(OH)2 acid + carbonate E.g. HCl, HNO3, H2SO4 E.g. MgCO3, ZnCO3, CuCO3 18 12.2 Preparing Salts Method 1 acid + metal → salt + hydrogen gas acid + insoluble base → salt + water acid + carbonate → salt + water + carbon dioxide gas 19 12.2 Preparing Salts Method 1 (Soluble Salt): Reacting an Acid with a Metal, Insoluble Base or Insoluble Carbonate Acid Soluble salt a. H2SO4(aq) + Zn(s) → ZnSO (aq) + H (g) 4 2 b. H2SO4(aq) + CuO(s) → CuSO4(aq) + H2O(l) c. HCl(aq) + MgCO3(s) → MgCl2(aq) + H2O(l) + CO2(g) Insoluble metal, base or carbonate used in excess: ensures that all the acid is used up. insoluble: can be filtered from the salt solution at the 20 end of the reaction. 12.2 Preparing Salts Method 1: Reacting an Acid with a Metal Example: Preparation of zinc sulfate H2SO4(aq) + Zn(s) → ZnSO4(aq) + H2(g) acid metal salt hydrogen zinc powder zinc sulfate solution excess zinc dilute sulfuric zinc sulfate acid solution 1. Add excess Zn powder into 2. Filter to remove excess dilute H2SO4. Stir until Zn powder. Collect the effervescence stops. Why? filtrate, ZnSO4 solution. 21 12.2 Preparing Salts Method 1: Reacting an Acid with a Metal Example: Preparation of zinc sulfate H2SO4(aq) + Zn(s) → ZnSO4(aq) + H2(g) acid metal salt hydrogen zinc sulfate solution glass rod crystals 3. Heat filtrate to obtain 4. Test for saturation. The concentrated ZnSO4 solution is saturated when solution. crystals form on glass rod. 22 12.2 Preparing Salts Method 1: Reacting an Acid with a Metal Example: Preparation of zinc sulfate H2SO4(aq) + Zn(s) → ZnSO4(aq) + H2(g) acid metal salt hydrogen filter paper saturated zinc sulfate solution zinc sulfate zinc sulfate crystals crystals (pure) URL 5. Leave the 6. Filter to obtain crystals. solution to cool 7. Wash with cold, distilled water. Pat and crystallise. dry between pieces of filter paper. 23 12.2 Preparing Salts Method 1: Reacting an Acid with a Metal This method is not suitable for some metals. Why? Some metals are too reactive. Potassium, sodium, calcium react too violently with acids! Some metals are unreactive. ✓ This method is suitable for moderately Copper, silver reactive metals do not react with acids! such as Zn, Mg, 24 Al and Fe. 12.2 Preparing Salts Method 1: Reacting an Acid with an Insoluble Base Example: Preparation of copper(II) sulfate H2SO4(aq) + CuO(s) → CuSO4(aq) + H2O(l) acid insoluble base salt water 1. React excess CuO with dilute H2SO4. Some heating is required. Stir until no more CuO dissolves. 2. Filter the mixture. 3. Collect the filtrate. 4. Heat the filtrate. 1. Cool and crystallise. 2. Filter, wash and dry. 25 12.2 Preparing Salts Method 1: Reacting an Acid with an Insoluble Carbonate Example: Preparation of magnesium chloride HCl (aq) + MgCO3(s) → MgCl2(aq) + H2O(l) + CO2(g) acid carbonate salt water carbon dioxide 1. React excess MgCO3 with dilute HCl until effervescence stops. 2. Filter the mixture. 3. Collect the filtrate. 4. Heat the filtrate. 5. Cool and crystallise. 26 6. Filter, wash and dry. 12.2 Preparing Salts 1. Salt is soluble 2. Starting material are both soluble Method 2 (Soluble Salt): Titration Reaction of acid + soluble base Suitable for preparing sodium, potassium and ammonium salts 27 12.2 Preparing Salts Method 2 (Soluble Salt): Titration acid + soluble base E.g. HCl, HNO3, H2SO4 E.g. NaOH, KOH, NH3(aq) Part 1: Titrate to determine volumes of reactants required using a suitable indicator. Part 2: Use the above volumes of reactants to prepare sodium nitrate. The indicator is not used here. 28 12.2 Preparing Salts Method 2: Reacting an Acid with a Soluble Base HNO3(aq) + NaOH(aq) → NaNO3(aq) + H2O(l) acid soluble base salt water Give examples of two salts that can prepared using this method. Write balanced chemical equations for them. HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l) H2SO4(aq) + 2KOH(aq) → K2SO4(aq) + 2H2O(l) 29  12.2 Preparing Salts Method 2: Reacting an Acid with a Soluble Base Example: Preparation of sodium nitrate Part 1: Titration to determine volumes of reactants HNO3(aq) + NaOH(aq) → NaNO3(aq) + H2O(l) 1. Fill up a burette with dilute HNO3. Note the initial burette reading (V1 cm3). retort burette stand 2. Pipette 25.0 cm3 of aqueous NaOH solution into a conical flask. 3. Add 1−2 drops of methyl orange (indicator) to the conical flask NaOH solution. 30 The solution turns yellow.  12.2 Preparing Salts Method 2: Reacting an Acid with a Soluble Base Example: Preparation of sodium nitrate Part 1: Titration to determine volumes of reactants 4. Add dilute HNO3 slowly from the burette until the solution just turns orange. This is the retort burette end-point. stand 5. Stop adding HNO3. Record the final burette reading (V2 cm3). Volume of acid required for conical flask complete neutralisation = (V2 – V1) cm3. 31 12.2 Preparing Salts Method 2: Reacting an Acid with a Soluble Base Example: Preparation of sodium nitrate Part 2: Use known volumes (from Part 1) of reactants to prepare NaNO3 1. Pipette 25.0 cm3 of NaOH solution into a beaker. Do not add indicator. Why? 2. Add (V2 – V1) cm3 of dilute HNO3 from the burette. 3. Heat to saturate the solution. 4. Cool and crystallise to obtain NaNO3. 5. Filter and dry. 32 12.2 Preparing Salts Comparison Between Method 1 and Method 2 (Titration) Method 1 Method 2 Uses one soluble Uses two soluble substance (acid) substances (acid and base) Residue of insoluble NO residue seen when substance can be seen reaction is complete when reaction is complete We use an indicator to tell us when the reaction is complete. 33 12.2 Preparing Salts Method 3 (Insoluble Salt): Precipitation The following salts can be prepared by this method: Insoluble salts Barium sulfate Lead(II) chloride All carbonates Calcium sulfate Silver chloride except carbonates Lead(II) sulfate of sodium, potassium and ammonium 34 12.2 Preparing Salts Method 3 (Insoluble Salt): Precipitation This method can be represented by the equation: AB(aq) + CD(aq) → AD(s) + CB(aq) Ionic equation for this reaction: A+(aq) + D−(aq) → AD(s) E.g. AgNO3(aq) + NaCl (aq) → AgCl (s) + NaNO3 (aq) Pb(NO3)2(aq) + Na2SO4(aq) → PbSO4(s) + 2NaNO3(aq) 35 12.2 Preparing Salts Method 3 (Insoluble Salt): Precipitation Example: Preparation of barium sulfate Ba(NO3)2(aq) + Na2SO4(aq) → BaSO4(s) + 2NaNO3(aq) barium nitrate sodium sulfate white precipitate of barium sulfate in sodium nitrate solution 36 12.2 Preparing Salts Method 3 (Insoluble Salt): Precipitation Example: Preparation of barium sulfate sodium nitrate barium solution + sulfate sodium excess sodium barium (impure) sulfate sulfate nitrate solution solution solution 1. Add Na2SO4 solution to 2. Filter to obtain the Ba(NO3)2 in a beaker. precipitate. A white precipitate of BaSO4 forms. 37 12.2 Preparing Salts Method 3 (Insoluble Salt): Precipitation Example: Preparation of barium sulfate barium sulfate filter paper (pure) barium sulfate (impure) 3. Wash the precipitate 4. Leave the with some cold, precipitate of distilled water to BaSO4 to dry. remove impurities. 38 Chapter 12 Salts Summary Map (Preparing Salts) 39 Chapter 12 Salts 12.1 Salts 12.2 Preparing Salts 12.3 Qualitative Analysis 40 12.3 Qualitative Analysis Learning Outcomes At the end of this section, you should be able to: describe the tests to identify aqueous cations; describe the tests to identify anions; describe the tests to identify gases. 41 12.3 Qualitative Analysis What is Qualitative Analysis (QA)? It is a process used by a chemist to identify the cations and anions in an unknown solution. 42 12.3 Qualitative Analysis Identifying Cations Add NaOH/NH3(aq): − Most cations give precipitates with alkalis, NaOH/NH3(aq), except Na+, K+ and NH4+. A cation can be identified by making the following observations: 1. The colour of the precipitate produced 2. Whether the precipitate is soluble or insoluble in excess NaOH/NH3(aq) 3. Whether ammonia gas is liberated on addition of NaOH solution 43 12.3 Qualitative Analysis What is the precipitate formed? It is the hydroxide of the metal ion. Example: A solution containing Cu2+ forms copper(II) hydroxide which is a light blue precipitate. URL Cu2+(aq) + 2OH−(aq) → Cu(OH)2(s) from solution from NaOH light blue of unknown or NH3(aq) precipitate substance 44 12.3 Qualitative Analysis Identifying Cations Cation Test with NaOH(aq) Test with NH3(aq) Al3+ A white precipitate is A white precipitate is formed which dissolves formed. in excess NaOH(aq) to The precipitate is give a colourless insoluble in excess solution. NH3(aq). Cu2+ A light blue precipitate is A light blue precipitate formed which is is formed which insoluble in excess dissolves in excess NaOH(aq). NH3(aq) to form a deep blue solution. Some precipitates dissolve because they form soluble compounds with excess NaOH or NH3(aq). 45 12.3 Qualitative Analysis Aqueous ammonia, NH3(aq) Identifying Cations Cation On adding a On adding excess few drops Example: Cu2+ Light blue Precipitate Observation for NH3(aq) precipitate dissolves in excess test on Cu2+ to form a deep blue solution. Record your observation as: A light blue precipitate is formed. It is soluble in excess NH3(aq) to give a deep blue solution. 46 12.3 Qualitative Analysis Identifying Cations Summary of tests with NaOH(aq) Sodium hydroxide solution, NaOH(aq) Cation On adding a few drops On adding excess Zinc ion, White precipitate Precipitate dissolves in excess to Zn2+ form a colourless solution. Aluminium White precipitate Precipitate dissolves in excess to ion, Al3+ form a colourless solution. Lead(II) White precipitate Precipitate dissolves in excess to ion, Pb2+ form a colourless solution. Calcium White precipitate Precipitate is insoluble in excess. ion, Ca2+ 47 12.3 Qualitative Analysis Identifying Cations Summary of tests with NaOH(aq) Sodium hydroxide solution, NaOH(aq) Cation On adding a few drops On adding excess Copper(II) Light blue precipitate Precipitate is insoluble in ion, Cu2+ excess. Iron(II) ion, Green precipitate Precipitate is insoluble in Fe2+ excess. Iron(III) ion, Reddish-brown Precipitate is insoluble in Fe3+ precipitate excess. Ammonium No precipitate. No change is observed. ion, NH4+ On heating, ammonia gas is given off. Ammonia gas turns 48 moist red litmus paper blue. 12.3 Qualitative Analysis Identifying Cations Summary of tests with NH3(aq) Aqueous ammonia, NH3(aq) Cation On adding a few drops On adding excess Zinc ion, White precipitate Precipitate dissolves in excess Zn2+ to form a colourless solution. Aluminium White precipitate Precipitate is insoluble in ion, Al3+ excess. Lead(II) ion, White precipitate Precipitate is insoluble in Pb2+ excess. Calcium No precipitate No precipitate ion, Ca2+ 49 12.3 Qualitative Analysis Identifying Cations Summary of tests with NH3(aq) Sodium hydroxide solution, NaOH(aq) Cation On adding a few drops On adding excess Copper(II) Light blue precipitate Precipitate dissolves in excess ion, Cu2+ to form a deep blue solution. Iron(II) ion, Green precipitate Precipitate is insoluble in Fe2+ excess. Iron(III) ion, Reddish-brown Precipitate is insoluble in Fe3+ precipitate excess. 50  12.3 Qualitative Analysis Summary Map (Identifying Cations) 51  12.3 Qualitative Analysis Summary Map (Identifying Cations) 52 12.3 Qualitative Analysis Identifying Anions Anion Test Observations for positive test and inference Carbonate Add dilute hydrochloric Effervescence is observed. ion, CO32– acid. Gas given off forms a white Pass the gas given off precipitate with limewater. URL 1 into limewater. Carbon dioxide gas is given off. Nitrate ion, Add sodium hydroxide Effervescence is observed. NO3– solution, then add a piece of aluminium foil. Warm the mixture. The moist red litmus paper Test the gas given off turns blue. Ammonia gas is URL 2 with a piece of moist red given off. litmus paper. 53 12.3 Qualitative Analysis Identifying Anions Anion Test Observations for positive test and inference Sulfate Add dilute nitric acid, A white precipitate of ion, SO42– then add barium barium sulfate is formed. nitrate solution. URL 1 Chloride Add dilute nitric acid, A white precipitate of ion, Cl– then add silver silver chloride is formed. nitrate solution. URL 2 Iodide Add dilute nitric acid, A yellow precipitate of ion, I– then add silver silver iodide is formed. nitrate solution. URL 3 54  12.3 Qualitative Analysis Summary Map (Identifying Anions) 55 12.3 Qualitative Analysis Identifying Gases A gas is given off during a chemical reaction when effervescence is observed in a liquid; the colour or odour of a gas is detected; a solid substance is heated (sometimes). 56 12.3 Qualitative Analysis Identifying Gases Gas Colour and Test Observations odour Hydrogen, Colourless Place a lighted splint The lighted H2 and at the mouth of the splint is odourless test tube. extinguished with a ‘pop’ URL sound. pop lighted splint 57 12.3 Qualitative Analysis Identifying Gases Gas Colour and Test Observations odour Oxygen, Colourless Insert a glowing The glowing O2 and splint into the test splint is odourless tube. rekindled (i.e. catches glowing splint fire again). 58 12.3 Qualitative Analysis Identifying Gases Gas Colour and Test Observations odour Carbon Colourless Bubble gas A white dioxide, and through limewater. precipitate is CO2 odourless formed. The precipitate dissolves upon limewater further bubbling. carbon dioxide 59 12.3 Qualitative Analysis Identifying Gases Gas Colour and Test Observations odour Chlorine, Greenish-y Place a piece of The moist blue Cl2 ellow gas moist blue litmus litmus paper with a paper at the turns red, and is pungent mouth of the test then bleached. URL smell tube. moist blue litmus paper 60 12.3 Qualitative Analysis Identifying Gases Gas Colour and Test Observations odour Sulfur Colourless Place a piece of The purple dioxide, gas with a filter paper soaked acidified SO2 pungent with acidified potassium smell potassium manganate(VII) manganate(VII) at turns the mouth of the test colourless. tube. filter paper soaked with acidified KMnO4 61 12.3 Qualitative Analysis Identifying Gases Gas Colour and Test Observations odour Ammonia, Colourless Place a piece of The moist red NH3 gas with a moist red litmus litmus paper pungent paper at the mouth turns blue. smell of the test tube. moist red URL litmus paper 62  12.3 Qualitative Analysis Summary Map (Identifying Gases) 63 12.3 Qualitative Analysis Tests for the Presence of Water Water is given off when hydrated salts are heated. Example: CuSO4.7H2O → CuSO4 + 7H2O Colourless liquid Condenses at mouth of the test tube 64 12.3 Qualitative Analysis Tests for the Presence of Water 1. Test with anhydrous cobalt(II) chloride Water will change the colour of dry cobalt(II) chloride paper from blue to pink. 2. Test with anhydrous copper(II) sulfate. Water will change the colour of anhydrous copper(II) sulfate from white to blue. Note: These two tests only show the presence of water. They cannot be used to test for the purity of water. 65 Chapter 12 Salts Concept Map 66 Chapter 12 Salts The URLs are valid as at 15 October 2012. Acknowledgements (slide 1) © Marshall Cavendish International (Singapore) (slide 10) © Marshall Cavendish International (Singapore) (slide 25) copper sulfate © Stephanb | Wikimedia Commons | CC BY-SA 3.0 (http://creativecommons.org/licenses/by-sa/3.0/deed.en) (slide 26) magnesium chloride © Walkerma | Wikimedia Commons | Public Domain (slide 36) © Marshall Cavendish International (Singapore) (slide 38) lead(II) sulfate © Walkerma | Wikimedia Commons | Public Domain 67