 All salts are ionic compounds.  A salt is formed when a metallic ion or an ammonium ion (NH 4 + ) replaces one or more hydrogen ions of an acid. HClNaCl HNO 3 NH 4 NO 3 replaced by Find out what is an acid salt!

 Where have you heard of salts?  Salts are essential to animal life in small quantities, but in large excess will be very harmful.  Are all salts salty?  Group I ions similar in size to sodium tend to give salty taste. Which ions do you think give a salty taste?

 How do you form CaSO 4 ?  Acid + Base  Salt + water CaSO 4 Comes from base Comes from acid One base that can be used is Calcium hydroxide. One acid that can be used is Sulfuric acid. Ca(OH) 2 + H 2 SO 4  CaSO 4 + 2H 2 O Can you use calcium carbonate or calcium?

 How do you form NaNO 3 ?  Acid + Base  Salt + Water NaNO 3 Comes from base Comes from acid One base that can be used is Sodium hydroxide. One acid that can be used is Nitric acid. NaOH + HNO 3  NaNO 3 + H 2 O

Negative ionAcid used SO 4 2- (sulfate ion) NO 3 - (nitrate ion) Cl - (chloride ion) Sulfuric acid (H 2 SO 4 ) Nitric acid (HNO 3 ) Hydrochloric acid (HCl)

 What acids do you use to get the following salts?  Sodium nitrate  Potassium phosphate  Ammonium ethanoate  Copper(II) iodide  Sodium citrate  Aluminium sulfate

 Many salts combine with water molecules to form crystals.  These water molecules are called water of crystallization.  Salts that contain water of crystallization are called hydrated salts.  Salts that do not contain water of crystallization are called anhydrous salts.

Name of saltFormula of anhydrous salt Formula of hydrated salt copper(II) sulfateCuSO 4 CuSO 4.5H 2 O magnesium sulfateMgSO 4 MgSO 4.7H 2 O sodium carbonateNa 2 CO 3 Na 2 CO 3.10H 2 O zinc sulfateZnSO 4 ZnSO 4.7H 2 O What does the ‘dot’ mean?

 Heating a hydrated salt removes water of crystallization. CuSO 4.5H 2 O Heat CuSO 4 + 5H 2 O hydrated hydrated copper(II) sulfate anhydrous copper(II) sulfate + water Heat

 Cobalt(II) chloride CoCl 2.6H 2 O Heat CoCl 2 + 6H 2 O hydrated hydrated cobalt(II) chloride anhydrous cobalt(II) chloride + water Heat Find out the use of cobalt(II) chloride paper in the lab.

Chlorides/ bromides/ iodides Sulfates Carbonates Nitrates All are soluble except: lead(II) chloride/brom ide/iodide silver chloride/brom ide/iodide All are soluble except: barium sulfate calcium sulfate lead(II) sulfate All are NOT soluble except: S odium carbonate P otassium carbonate A mmonium carbonate All are soluble. S odium salts P otassium salts A mmonium salts All are soluble.

 There are 3 ways of preparing salts. But to choose which way depends on:  1. Whether the salt is soluble in water?  2. Whether the starting materials are soluble in water?

Methods of preparing salts Is the salt soluble? No Method 3: Precipitation Yes Reaction with acids Acid + metal Acid + base Acid + carbonate Are the starting materials soluble? YesNo Method 1: Reaction of acids with insoluble substances Acid + metal Acid + base Acid + carbonate Method 2: Titration Acid + alkali Acid + carbonate solution 1)Filter the mixture 2)Collect filtrate Salt solution Salt crystals (dry with filter paper) 1)Concentrate 2)Crystallize 3)Filter

 Recap:  Acid + Metal oxide/hydroxide  Salt + Water  Acid + Carbonate  Salt + Water + Carbon dioxide  Acid + Metal  Salt + Hydrogen gas  In Method 1, all the substances in red are insoluble.

 The acid is reacted with an excess of the substances (metal, carbonate or base).  Why? Acid Insoluble base

 Acid + Metal  Salt + Hydrogen gas ZnSO 4 Comes from metalComes from acid Zinc is insoluble in water and reacts with sulfuric acid. We can use zinc here.

sulfuric acid zinc sulfate solution + unreacted zinc Keep adding zinc until no more effervescence is observed. zinc sulfate solution 1)Filter the mixture 2)Collect filtrate zinc sulfate crystals 1)Concentrate 2)Crystallize 3)Filter

 Only moderately reactive metals like zinc, aluminum, magnesium and iron can be used.  Not suitable for 1. Very reactive metals such as sodium, potassium and calcium. Why? 2. Unreactive metals like copper and silver. Why?

 Acid + Metal oxide/ hydroxide  Salt + Water CuSO 4 Comes from metal oxideComes from acid Copper(II) oxide is insoluble in water and reacts with sulfuric acid. We can use Copper(II) oxide here. Why can’t we use copper metal?

sulfuric acid Copper(II) sulfate solution + unreacted Copper(II) oxide Keep adding copper(II) oxide until no more effervescence is observed. Copper(II) sulfate solution 1)Filter the mixture 2)Collect filtrate Copper(II) sulfate crystals 1)Concentrate 2)Crystallize 3)Filter

 Acid + Carbonate  Salt + Carbon dioxide + Water MgCl 2 Comes from carbonateComes from acid Magnesium carbonate is insoluble in water and reacts with hydrochloric acid. NOTE: ALL carbonates are insoluble except potassium, sodium and ammonium carbonate!

hydrochloric acid Magnesium chloride solution + unreacted magnesium carbonate Keep adding magnesium carbonate until no more effervescence is observed. Magnesium chloride solution 1)Filter the mixture 2)Collect filtrate Magnesium chloride crystals 1)Concentrate 2)Crystallize 3)Filter

Filter the crystals and dry them by squeezing them between sheets of filter paper. Pour the filtrate into an evaporating dish and heat to evaporate most of the water. This produces a hot saturated solution of the salt. Let the solution cool. Remove the excess metal/metal carbonate/ base by filtering and collect the filtrate. The filtrate contains the solution of the salt you want.) Add the metal/metal carbonate/ base slowly with stirring to hot acid (what acid depends on what salts you want) until no more dissolves. (This means all the acid is used up )

 Notice that all sodium, potassium and ammonium salts are SOLUBLE in water.  So you cannot use Method 1 for any of such salts! Why?

 To tell when all the acid has been completely used up, we have to use titration, by using an indicator.  What is an indicator? IndicatorColour in acidic solution pH range at which indicator changes colour Colour in alkaline solution methyl orangered3 – 5yellow screened methyl orange violet3 – 5green litmusred5 – 8blue bromothymol blueyellow6 – 8blue phenolphthaleincolourless8 – 10pink

Fill up a burette with dilute nitric acid and note down the initial burette reading (V 1 cm 3 ). V 1 cm 3 Pipette 25.0 cm 3 of dilute sodium hydroxide into a conical flask. Add one or two drops of methyl orange to the NaOH solution. The solution turns yellow. Add dilute HNO 3 from the burette slowly until the solution turns orange permanently. This is the end-point. The acid is all used up. V 2 cm 3 Record the final burette reading (V 2 cm 3 ). Hence, the volume of acid required for complete neutralization = (V 2 – V 1 ) cm 3.

1. Pipette 25.0cm 3 of NaOH into a beaker. Then add (V 2 – V 1 ) cm 3 of dilute nitric acid from the burette. This time do not add indicator. Why? 2. Heat to evaporate the water till it is saturated. 3. Cool the saturated solution so that the salt crystallizes. 4. Filter to collect the crystals. 5. Dry the crystals between a few sheets of filter paper.

 Simulation Simulation

Record the final burette reading (V2 cm3 ). Hence, the volume of acid required for complete neutralization = (V 2 – V 1 ) cm3. Add dilute HNO3 from the burette slowly until the solution turns orange permanently. This is the end-point. Add one or two drops of methyl orange to the alkali solution. The solution turns yellow. (Note, if you have a strong acid and weak base, you use methyl orange, if you have a strong base and weak acid, use phenolphthalein. If both are strong, you can use either indicator. Pipette 25.0 cm 3 of dilute alkali (depending on what salt you want) into a conical flask. Fill up a burette with dilute acid (depending on what salt you want) and note down the initial burette reading (V1 cm3 ).

Dry the crystals between a few sheets of filter paper. Filter to collect the crystals. Cool the saturated solution so that the salt crystallizes. Heat to evaporate the water till it is saturated. Pipette 25.0cm 3 of NaOH into a beaker. Then add (V2 – V1) cm 3 of dilute nitric acid from the burette.

 Easiest to prepare  Just need to use precipitation  Mix a solution containing the positive ions of the salt with another solution containing the negative ions of the salt to be prepared.  What salts are insoluble?

 Using lead(II) nitrate and dilute sulfuric acid 1. First, pour 50 cm 3 of lead(II) nitrate solution into a small beaker. Add sulfuric acid (in excess) and stir until no more precipitate forms. 2. Filter and collect precipitate. 3. Wash the precipitate with a small amount of distilled water to remove impurities. 4. Allow the precipitate to dry on filter paper.

Allow the precipitate to dry on filter paper. Wash the precipitate with a small amount of distilled water to remove impurities. Filter and collect precipitate. Add another reagent (again depending on what salt you want) and stir until no more precipitate forms. First, pour 50 cm 3 of one reagent (depending on what salt you want) into a small beaker.

 How do you get the following salts: 1. Magnesium sulfate 2. Lead(II) chloride 3. Potassium nitrate 4. Sodium sulfate 5. Zinc nitrate

 Are these salts soluble? 1. iron(III) nitrate 2. potassium carbonate 3. sodium ethanoate 4. silver chloride 5. lead(II) nitrate 6. copper(II) carbonate 7. ammonium iodide 8. titanium(IV) chloride 9. barium sulfate YesYesYesNoYesNoYesYesNo

 Which method will you use to get the following salts: 1. Magnesium sulfate 2. Lead(II) chloride 3. Potassium nitrate 4. Sodium sulfate 5. Copper(II) chloride 6. Lead(II) carbonate 7. Silver chloride 8. Zinc chloride

 When an acid Z is added to a solution of lead(II) nitrate, an insoluble precipitate is formed.  When Z is added to a solution of silver nitrate, an insoluble precipitate is formed too.  What acid could Z be? A) hydrochloric acidB) sulfuric acid C) nitric acid

 A metal oxide A dissolves in sulfuric acid, hydrochloric acid and nitric acid and does NOT give any precipitate with any of the acids. Which of the following could be A? A) Barium oxideB) Calcium oxide C) Silver oxideD) Sodium oxide