7. Acids, Bases and Salts Acids
According to the Bronsted-Lowry theory, an a cid is a proton donor. ( A A proton is a hydrogen cation, H + ). Classifying Acids
Strength
Strong Acids: A strong acid dissociates to a large ex tent in a solution. Almost all the acid molecules of a strong strong acid dissociate to form H+ ions. Examples: Hydrochloric acid, sulphuric acid and nitric acid
Weak Acid: A weak acid dissociates only to a small extent, and, therefore, can provide p rovide only a low concentration of hydrogen ions. Examples: Carbonic acid, acetic acid and oxalic acid are weak acids.
Basicity Basicity of an acid: The basicity of an acid is the number of hydrogen ions produced when one molecule of acid ionizes in water.
Mono-basic acids: acids:
Acids which on ionization produces on hydronium ion in water are termed mono-basic acids. Example: HCl
Di-basic acids: acids:
Acids which on ionization produces two hydronium ions are called di-basic acids. Example: H SO , H CO .etc
7. Acids, Bases and Salts Bases
According to the Bronsted-Lowry theory, an acid is a proton acceptor. Classifying Bases
Strength
Strong base: A strong base dissociates almost completely in a solution. Example: NaOH, KOH
Weak base: A weak base is not able to dissociate completely, giving a low concentration of hydroxyl ions in a solution. Example: NH4OH, NH3
Acidity Acidity of bases: It is the number of hydroxyl groups present in one molecule of a base.
Mono acidic bases: Bases which produces only one hydroxide (OH-) ion in aqueous solutions are called mono acidic bases. Example: NaOH, KOH
Di acidic bases: Bases which produces two h ydroxide ions in aqueous solutions are called di acidic bases. Example: Ca(OH)2, Mg(OH)2
Tri acid bases: Bases which produces three h ydroxide ions in aqueous solutions are called tri acidic bases. Example: Al(OH)3, Fe(OH)3
ACID/BASE PROPERTIES OF OXIDES An oxide i a compound f
med b twee
d anoth
l ment. Oxides show acid/base
7. Acids, Bases and Salts (2) Basic oxides:- These are oxides of metals which react with acids to form a salt and water only. Examples of basic oxides are magnesium oxide (MgO), calcium oxide (CaO), iron(II) oxide(FeO) MgO (s) + H2SO4(aq)
→ Mg SO4(aq) + H2O(l)
CaO (s) + 2HCl(aq)
→ Ca Cl2(aq) + H2O(l)
FeO (s) + HNO3(aq)
→ Fe(NO3)2 (aq) + H2O(l)
(3) Neutral Oxides:- These are oxides of non-metals which react with neither acids or bases . Examples of acidic oxides are Carbon monoxide (CO), nitrogen monoxide (NO), N2O.
Amphoteric Oxides:- These are oxides of some metals which react with both acids and bases.
Examples of amphoteric oxides are aluminium oxide (Al 2O3), lead(II) oxide (PbO), zinc oxide(ZnO) PbO (s) + HNO3(aq)
→ Pb( NO3)2(aq) + H2O(l)
PbO (s) + NaOH(aq)
→ Na2PbO2(aq) + H2O(l) Sodium plumbate
7. Acids, Bases and Salts Recognizing Acids and Alkalis
We can identify acids and alkalis by the use of indicators. An indicator is a substance (usually a dye) which has one colour in acidic solutions and another colour in alkaline solutions. Indicator
Colour in acidic solutions
Litmus Methyl Orange Phenolphthalein Screened methyl orange Bromothymol Blue
Red Pink/red Colourless Red Yellow
Colour in alkaline solutions. Blue Yellow Pink Green Blue
Note: In most instances Litmus comes in two forms, Red litmus paper and Blue litmus paper Acids turn blue litmus paper red Alkalis turn red litmus paper blue
There are two other chemical tests which can be used to identify a substance as an acid:
7. Acids, Bases and Salts The pH of a substance is a measure of how acidic or how alkaline a solution is. pH is measured using the pH scale.
The pH scale ranges from 0 to 14. A range of 0-6 is acidic, a range of 7 is neutral and a range of 8-14 is basic.
The pH scale is used along with Universal Indicator to determine the pH of a solution. Universal indicator is a mixtures of dyes which gives a particular colour for a specific pH range.
7. Acids, Bases and Salts 2. Reaction with Metals Acids react with metals to form a salt and h ydrogen gas. E.g. 2HCl (aq) + Zn(s) → ZnCl2 (s) + H2 (g) 3. Reaction with Carbonates Acids react with carbonates to form a salt, carbon dioxide and water. E.g. 2H3PO4 (aq) + 3Na2CO3 (aq) → 2Na3PO4 (aq) + 3CO2 (g) + 3H2O (l) HCl (aq) + NaHCO3 (aq) → NaCl (aq) + CO2 (aq) + H2O (l)
Reactions of Bases
1. Reaction with Acids Bases with acids to form a salt and water o nly. E.g. H2SO4 (aq) + KOH (aq) → K 2SO4 (aq) + H2O (l) 2. Reaction with Ammonium Salts Bases react with ammonium salts to form a salt, ammonia and water. NaOH (aq) + NH4Cl (aq) → NaCl (aq) + NH3 (g) + H2O (l)
Salts
7. Acids, Bases and Salts Acid Salts These are formed when some of the hydrogen ions in the acid have been replaced by metal or ammonium ions. The acids used to form these salts must have more than one replaceable hydrogen ions. These acids are dibasic and tribasic. Examples of these types of salts are potassium hydrogensulphate, KHSO4; sodium hydrogencarbonate, NaHCO3 and potassium dihydrogen phosphate, KH2PO4.
Acid salts react similarly to acids because of the presence o f replaceable hydrogen ions in them. As a result of this they can be distinguished from normal salts.
Chemical tests to identify acid salts: (i)
React the acid salt with a reactive metal. Hydrogen gas and a salt are formed.
(ii)
React with a carbonate. Carbon dioxide, a salt and water are formed.
Hydrated salts
7. Acids, Bases and Salts The solubility of salts
All nitrates are soluble.
All sodium, potassium and ammonium salts are soluble.
All chlorides are soluble except silver chloride. Lead chloride is o nly soluble in hot water.
All sulphates are soluble except for barium sulphate and lead sulphate. Calcium sulphate and silver sulphate are slightly soluble.
All carbonates and phosphates are insoluble except sodium, potassium and ammonium carbonates and phosphates.
All ethanoates are soluble except silver ethanoate which is sparingly soluble.
Note: Solubility of other substances.
All metal oxides are insoluble except sodium oxide and potassium oxide. Calcium oxide is slightly soluble.
All hydroxide are insoluble except sodium h ydroxide, potassium hydroxide and ammonium hydroxide. Calcium hydroxide is slightly soluble.
7. Acids, Bases and Salts Method 1. In a beaker add barium chloride to sodium sulphate 2. Filter the mixture using a filter funnel and filter paper. 3. Collect the residue (which is the insoluble salt) and wash it with distilled water to remove any impurities. 4. Allow the salt to dry.
Preparation of Soluble Salts
Preparation of Binary Anhydrous Salts These salts are prepared by Direct Combination. For example anhydrous iron (II) chloride is prepared by burning iron in chlorine gas. The reaction is given by: Fe(s) + Cl2 (g) → FeCl2 (s)
Soluble salts can also be prepared by the reaction of an acid with a reactive metal, an insoluble
7. Acids, Bases and Salts Preparation
Metals + acid
Salts prepared Soluble salts of the reactive metals, Mg, Al, Zn, Fe.
Starting materials Appropriate metal to provide cations and appropriate acid to provide anions.
Method
Insoluble carbonate + acid
Soluble salts Appropriate except sodium, carbonate to potassium and provide cations ammonium and appropriate salts. acid to provide anions.
Insoluble base + acid
Soluble salts Insoluble base to except sodium, provide cations potassium and and acid to ammonium provide anions. salts.
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Examples
Starting materials and equations for examples Zn(s) + 2HCl(aq) → ZnCl2(aq)
Add metal to fixed volume of acid until effervescence stops and metal present in excess. Filter to remove excess metal, collect filtrate. Evaporate some water. Leave to crystallize. Rinse and dry crystals.
Zinc chloride
Add insoluble carbonate to fixed volume of acid until effervescence stops and carbonate present in excess. Heat if necessary. Filter to remove excess carbonate, collect filtrate. Evaporate some water. Leave to crystallize. Rinse and dry crystals.
Calcium nitrate
CaCO3(s) + HNO3(aq) → Ca(NO3)2(aq) + CO2(g) + H2O(l)
Add insoluble base to fixed volume of acid until base present in excess. Stir and heat if necessary. Filter to remove excess base, collect filtrate. Evaporate some water. Leave to crystallize. Rinse and dry crystals.
Copper sulphate
CuO(s) + H2SO4(aq) → CuSO4(aq)+ H2O(l)