CHAPTER 12 Salts 2013 Marshall Cavendish International Singapore

CHAPTER 12 Salts © 2013 Marshall Cavendish International (Singapore) Private Limited © 2014 Marshall Cavendish Education Pte Ltd (Formerly known as 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. E. g. Na+, K+, Mg 2+, Zn 2+, Al 3+ H+ NH 4+ 4

12. 1 Salts What is a Salt? A salt is made up of two parts. CATION ANION • metallic ion • ammonium ion • comes from the base • non-metallic ion • comes from the acid Chloride from HCl Nitrate from HNO 3 Sulfate from H 2 SO 4 5

12. 1 Salts Reactions in Which Salts Can Be Made Example 1: Zinc sulfate Zn(OH)2(s) + H 2 SO 4(aq) → Zn. SO 4(aq) + 2 H 2 O(l) Zn. SO 4 cation anion Zn 2+ SO 42– Comes from the base, Zn(OH)2 Comes from the acid, H 2 SO 4 6

12. 1 Salts Reactions in Which Salts Can Be Made Example 2: Sodium chloride Na. OH + HCl → Na. Cl + H 2 O H+ of HCl is replaced by ____. Na+ ____ 7

12. 1 Salts Reactions in Which Salts Can Be Made Possible reactants Salt formed Metal/ Carbonate (for the cation) Acid (for the anion) zinc (Zn) hydrochloric acid (HCl) zinc chloride (Zn. Cl 2) copper(II) carbonate (Cu. CO 3) nitric acid (HNO 3) copper(II) nitrate (Cu(NO 3)2) magnesium oxide (Mg. O) sulfuric acid (H 2 SO 4) magnesium sulfate (Mg. SO 4) aqueous ammonia (NH 3(aq)) sulfuric acid (H 2 SO 4) ammonium sulfate ((NH 4)2 SO 4) potassium hydroxide (KOH) phosphoric acid (H 3 PO 4) potassium phosphate (K 3 PO 4) 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 hydrated salt heat → anhydrous salt + water 9

12. 1 Salts Water of Crystallisation Example: Copper(II) sulfate Hydrated salts contain water of crystallisation. blue crystals, hydrated copper(II) sulfate Cu. SO 4. 5 H 2 O water removed by heating add water Anhydrous salts do not contain water of crystallisation. white powder, copper(II) sulfate Cu. SO 4 The amount of water crystallised is indicated after the dot ‘. ’ in its chemical formula. 10

12. 1 Salts Solubility of Salts • All salts containing Na+, K+, and NH 4+ ions are soluble in water. E. g. Na. Cl, K 2 CO 3 and (NH 4)2 SO 4. • All nitrates are soluble. E. g. Pb(NO 3)2, KNO 3, Zn(NO 3)2 and Ag. NO 3. 11

12. 1 Salts Solubility of Salts • All chlorides are soluble except Pb. Cl 2 and Ag. Cl. • All sulfates are soluble except Ba. SO 4, Ca. SO 4 and Pb. SO 4. • All carbonates are insoluble except Na 2 CO 3, K 2 CO 3 and (NH 4)2 CO 3. 12

12. 1 Salts Solubility of Salts (Summary) Soluble salts Insoluble salts All nitrates All sodium salts All potassium salts All ammonium salts All chlorides All sulfates except All carbonates except Silver chloride, Ag. Cl Lead(II) chloride, Pb. Cl 2 except Barium sulfate, Ba. SO 4 Lead(II) sulfate, Pb. SO 4 Calcium sulfate, Ca. SO 4 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 • acid + insoluble base • acid + carbonate E. g. HCl, HNO 3, H 2 SO 4 E. g. Mg. O, Zn. O, Cu. O, Fe(OH)3, Cu(OH)2 E. g. Mg. CO 3, Zn. CO 3, Cu. CO 3 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 Soluble salt Acid a. H 2 SO 4(aq) + Zn(s) → Zn. SO 4(aq) + H 2(g) b. H 2 SO 4(aq) + Cu. O(s) → Cu. SO 4(aq) + H 2 O(l) c. HCl(aq) + Mg. CO 3(s) → Mg. Cl 2(aq) + H 2 O(l) + CO 2(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 end of the reaction. 20

12. 2 Preparing Salts Method 1: Reacting an Acid with a Metal Example: Preparation of zinc sulfate H 2 SO 4(aq) + Zn(s) acid metal zinc powder → Zn. SO 4(aq) + H 2(g) salt hydrogen zinc sulfate solution excess zinc dilute sulfuric acid 1. Add excess Zn powder into dilute H 2 SO 4. Stir until effervescence stops. Why? zinc sulfate solution 2. Filter to remove excess Zn powder. Collect the filtrate, Zn. SO 4 solution. 21

12. 2 Preparing Salts Method 1: Reacting an Acid with a Metal Example: Preparation of zinc sulfate H 2 SO 4(aq) + Zn(s) acid metal zinc sulfate solution → Zn. SO 4(aq) + H 2(g) salt hydrogen glass rod crystals 3. Heat filtrate to obtain concentrated Zn. SO 4 solution. 4. Test for saturation. The solution is saturated when crystals form on glass rod. 22

12. 2 Preparing Salts Method 1: Reacting an Acid with a Metal Example: Preparation of zinc sulfate H 2 SO 4(aq) + Zn(s) acid saturated zinc sulfate solution zinc sulfate crystals 5. Leave the solution to cool and crystallise. metal → Zn. SO 4(aq) + H 2(g) salt hydrogen filter paper zinc sulfate crystals (pure) 6. Filter to obtain crystals. 7. Wash with cold, distilled water. Pat dry between pieces of filter paper. URL 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. Copper, silver do not react with acids! ✓ This method is suitable for moderately reactive metals such as Zn, Mg, Al and Fe. 24

12. 2 Preparing Salts Method 1: Reacting an Acid with an Insoluble Base Example: Preparation of copper(II) sulfate H 2 SO 4(aq) + Cu. O(s) acid → Cu. SO 4(aq) + H 2 O(l) insoluble base 1. React excess Cu. O with dilute H 2 SO 4. Some heating is required. Stir until no more Cu. O dissolves. 2. Filter the mixture. 3. Collect the filtrate. 4. Heat the filtrate. 1. Cool and crystallise. 2. Filter, wash and dry. salt water 25

12. 2 Preparing Salts Method 1: Reacting an Acid with an Insoluble Carbonate Example: Preparation of magnesium chloride HCl (aq) + Mg. CO 3(s) → Mg. Cl 2(aq) + H 2 O(l) + CO 2(g) acid carbonate salt 1. React excess Mg. CO 3 with dilute HCl until effervescence stops. 2. Filter the mixture. 3. Collect the filtrate. 4. Heat the filtrate. 1. 5. Cool and crystallise. 2. 6. Filter, wash and dry. water carbon dioxide 26

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 + E. g. HCl, HNO 3, H 2 SO 4 soluble base E. g. Na. OH, KOH, NH 3(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 HNO 3(aq) + Na. OH(aq) → Na. NO 3(aq) + H 2 O(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) + Na. OH(aq) → Na. Cl(aq) + H 2 O(l) H 2 SO 4(aq) + 2 KOH(aq) → K 2 SO 4(aq) + 2 H 2 O(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 HNO 3(aq) + Na. OH(aq) → Na. NO 3(aq) + H 2 O(l) retort stand burette conical flask 1. Fill up a burette with dilute HNO 3. Note the initial burette reading (V 1 cm 3). 2. Pipette 25. 0 cm 3 of aqueous Na. OH solution into a conical flask. 3. Add 1− 2 drops of methyl orange (indicator) to the Na. OH solution. The solution turns yellow. 30

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 retort stand burette 4. Add dilute HNO 3 slowly from the burette until the solution just turns orange. This is the end-point. 5. Stop adding HNO 3. Record the final burette reading (V 2 cm 3). conical flask Volume of acid required for complete neutralisation = (V 2 – V 1) cm 3. 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 Na. NO 3 1. Pipette 25. 0 cm 3 of Na. OH solution into a beaker. Do not add indicator. Why? 2. Add (V 2 – V 1) cm 3 of dilute HNO 3 from the burette. 3. Heat to saturate the solution. 4. Cool and crystallise to obtain Na. NO 3. 5. Filter and dry. 32

12. 2 Preparing Salts Comparison Between Method 1 and Method 2 (Titration) Method 1 Method 2 Uses one soluble substance (acid) Residue of insoluble substance can be seen when reaction is complete Uses two soluble substances (acid and base) NO residue seen 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 Calcium sulfate Silver chloride Lead(II) sulfate All carbonates except carbonates 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. Ag. NO 3(aq) + Na. Cl (aq) → Ag. Cl (s) + Na. NO 3 (aq) Pb(NO 3)2(aq) + Na 2 SO 4(aq) → Pb. SO 4(s) + 2 Na. NO 3(aq) 35

12. 2 Preparing Salts Method 3 (Insoluble Salt): Precipitation Example: Preparation of barium sulfate Ba(NO 3)2(aq) + Na 2 SO 4(aq) → Ba. SO 4(s) + 2 Na. NO 3(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 barium nitrate solution sodium sulfate solution 1. Add Na 2 SO 4 solution to Ba(NO 3)2 in a beaker. A white precipitate of Ba. SO 4 forms. barium sulfate (impure) sodium nitrate solution + excess sodium sulfate solution 2. Filter to obtain the precipitate. 37

12. 2 Preparing Salts Method 3 (Insoluble Salt): Precipitation Example: Preparation of barium sulfate filter paper barium sulfate (pure) barium sulfate (impure) 3. Wash the precipitate with some cold, distilled water to remove impurities. 4. Leave the precipitate of Ba. SO 4 to dry. 38

12. 2 Preparing 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 Na. OH/NH 3(aq): − Most cations give precipitates with alkalis, Na. OH/NH 3(aq), except Na+, K+ and NH 4+. • 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 Na. OH/NH 3(aq) 3. Whether ammonia gas is liberated on addition of Na. OH solution 43

12. 3 Qualitative Analysis What is the precipitate formed? • It is the hydroxide of the metal ion. • Example: A solution containing Cu 2+ forms copper(II) hydroxide which is a light blue precipitate. URL Cu 2+(aq) + 2 OH−(aq) → Cu(OH)2(s) from solution of unknown substance from Na. OH or NH 3(aq) light blue precipitate 44

12. 3 Qualitative Analysis Identifying Cations Cation Al 3+ Cu 2+ Test with Na. OH(aq) Test with NH 3(aq) A white precipitate is formed which dissolves in excess Na. OH(aq) to give a colourless solution. A light blue precipitate is formed which is insoluble in excess Na. OH(aq). A white precipitate is formed. The precipitate is insoluble in excess NH 3(aq). A light blue precipitate is formed which dissolves in excess NH 3(aq) to form a deep blue solution. Some precipitates dissolve because they form soluble compounds with excess Na. OH or NH 3(aq). 45

12. 3 Qualitative Analysis Identifying Cations Example: Observation for NH 3(aq) test on Cu 2+ Aqueous ammonia, NH 3(aq) Cation Cu 2+ On adding a few drops Light blue precipitate On adding excess Precipitate dissolves in excess to form a deep blue solution. Record your observation as: A light blue precipitate is formed. It is soluble in excess NH 3(aq) to give a deep blue solution. 46

12. 3 Qualitative Analysis Identifying Cations Summary of tests with Na. OH(aq) Sodium hydroxide solution, Na. OH(aq) Cation On adding a few drops On adding excess Zinc ion, Zn 2+ White precipitate Precipitate dissolves in excess to form a colourless solution. Aluminium ion, Al 3+ White precipitate Precipitate dissolves in excess to form a colourless solution. Lead(II) ion, Pb 2+ White precipitate Precipitate dissolves in excess to form a colourless solution. Calcium ion, Ca 2+ White precipitate Precipitate is insoluble in excess. 47

12. 3 Qualitative Analysis Identifying Cations Summary of tests with Na. OH(aq) Sodium hydroxide solution, Na. OH(aq) Cation On adding a few drops On adding excess Copper(II) ion, Cu 2+ Light blue precipitate Precipitate is insoluble in excess. Iron(II) ion, Fe 2+ Green precipitate Precipitate is insoluble in excess. Iron(III) ion, Fe 3+ Reddish-brown precipitate Precipitate is insoluble in excess. Ammonium ion, NH 4+ No precipitate. No change is observed. On heating, ammonia gas is given off. Ammonia gas turns moist red litmus paper blue. 48

12. 3 Qualitative Analysis Identifying Cations Summary of tests with NH 3(aq) Aqueous ammonia, NH 3(aq) Cation On adding a few drops On adding excess Zinc ion, Zn 2+ White precipitate Precipitate dissolves in excess to form a colourless solution. Aluminium ion, Al 3+ White precipitate Precipitate is insoluble in excess. Lead(II) ion, White precipitate Pb 2+ Precipitate is insoluble in excess. Calcium ion, Ca 2+ No precipitate 49

12. 3 Qualitative Analysis Identifying Cations Summary of tests with NH 3(aq) Sodium hydroxide solution, Na. OH(aq) Cation On adding a few drops On adding excess Copper(II) ion, Cu 2+ Light blue precipitate Precipitate dissolves in excess to form a deep blue solution. Iron(II) ion, Fe 2+ Green precipitate Precipitate is insoluble in excess. Iron(III) ion, Reddish-brown Fe 3+ precipitate Precipitate is insoluble in 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 Carbonate ion, CO 32– URL 1 Nitrate ion, NO 3– URL 2 Test Observations for positive test and inference Add dilute hydrochloric acid. Pass the gas given off into limewater. Effervescence is observed. Gas given off forms a white precipitate with limewater. Carbon dioxide gas is given off. Add sodium hydroxide solution, then add a piece of aluminium foil. Warm the mixture. Test the gas given off with a piece of moist red litmus paper. Effervescence is observed. The moist red litmus paper turns blue. Ammonia gas is given off. 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, SO 42– then add barium sulfate is formed. nitrate solution. URL 1 Chloride ion, Cl– Add dilute nitric acid, A white precipitate of then add silver chloride is formed. nitrate solution. Iodide ion, Add dilute nitric acid, A yellow precipitate of I– then add silver iodide is formed. nitrate solution. URL 2 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 odour Test Observations Hydrogen, Colourless Place a lighted splint The lighted H 2 and at the mouth of the splint is odourless test tube. extinguished with a ‘pop’ sound. URL pop lighted splint 57

12. 3 Qualitative Analysis Identifying Gases Gas Oxygen, O 2 Colour and odour Test Colourless Insert a glowing and splint into the test odourless tube. glowing splint Observations The glowing splint is rekindled (i. e. catches fire again). 58

12. 3 Qualitative Analysis Identifying Gases Gas Carbon dioxide, CO 2 Colour and odour Test Observations Colourless Bubble gas A white and through limewater. precipitate is odourless formed. The precipitate dissolves upon limewater further bubbling. carbon dioxide 59

12. 3 Qualitative Analysis Identifying Gases Gas Chlorine, Cl 2 URL Colour and odour Greenishyellow gas with a pungent smell Test Place a piece of moist blue litmus paper at the mouth of the test tube. Observations The moist blue litmus paper turns red, and is then bleached. moist blue litmus paper 60

12. 3 Qualitative Analysis Identifying Gases Gas Colour and odour Sulfur Colourless dioxide, gas with a SO 2 pungent smell Test Observations Place a piece of filter paper soaked with acidified potassium manganate(VII) at the mouth of the test tube. The purple acidified potassium manganate(VII) turns colourless. filter paper soaked with acidified KMn. O 4 61

12. 3 Qualitative Analysis Identifying Gases Gas Colour and odour Test Observations Ammonia, Colourless Place a piece of The moist red NH 3 gas with a moist red litmus paper pungent paper at the mouth turns blue. smell of the test tube. URL moist red litmus paper 62

12. 3 Qualitative Analysis Summary Map (Identifying Gases) 63

12. 3 Qualitative Analysis Tests for the Presence of Water is given off when hydrated salts are heated. Example: Cu. SO 4. 7 H 2 O → Cu. SO 4 + 7 H 2 O • 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 November 2014. 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