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Acids, Bases & Salts

Acids, Bases & Salts. Characteristic Properties of Acids and Bases The pH scale, pH indicators and measurement of pH Classification of Oxides Salts: Uses and Preparation. ACIDS.

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Acids, Bases & Salts

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  1. Acids, Bases & Salts • Characteristic Properties of Acids and Bases • The pH scale, pH indicators and measurement of pH • Classification of Oxides • Salts: Uses and Preparation

  2. ACIDS • An acid is a substance which when dissolved in water produces hydrogen ions (H+) as the only positive ion in solution. Types of acids Mineral acids or Inorganic acids (from mineral elements or inorganic material) Organic acids (from plant or animal material) • acetic acid (vinegar or ethanoic acid) • formic acid (from ants) • ascorbic acid (vitamin C) • citric acid • tartaric acid • amino acids (subunits of proteins) • lactic acid (fermented milk) • hydrochloric acid • sulfuric acid • nitric acid • phosphoric acids • carbonic acid

  3. Pure acids • Pure acids do not have water added to them. ie. HCl(l), HCl(g) H2SO4(l), HNO3(l) • Pure acids exist as covalent molecules. • They only ionise when they dissolve in water. • As they do not possess hydrogen ions, they do not exhibit acidic properties. • No water => no hydrogen ions • Water must be present for an acid to exhibit its acidic properties.

  4. H+ SO42- SO42- H+ H+ SO42- CH3COOH H+ CH3COO- CH3COOH Strong and Weak acids • A strong acid is an acid which undergoes complete ionization in water (ie. all the molecules ionize to produce H+ ions. There are no molecules left in solution.) • H2SO4(aq) 2H+(aq) + SO42-(aq) • eg. H2SO4,HCl, HNO3 • A weak acid is an acid which undergoes only partial ionization in water. (ie, most of the acid molecules remain uncharged in water, only a few molecules ionize to release H+ ions) • CH3COOH(aq) CH3COO-(aq) + H+(aq) • H2SO3(aq) 2H+(aq) + SO32-(aq) • H2CO3(aq) 2H+(aq) + CO32-(aq)

  5. Basicity of an Acid • The basicity of an acid refers to the number of hydrogen ions that can be produced per molecule of acid when dissolved in water. • Example: Sulfuric acid is a dibasic acid as it produces two hydrogen ions per molecule of acid when dissolved in water. • Basicity and strength of some common acids

  6. Properties of acids • Dilute acids have a sour taste. • Acids are soluble in water and give solutions with a pH below 7. They give charatcteristic colours with indicators.They turn litmus red. • The concentrated forms of strong acids like hydrochloric acid and sulphuric acid are corrosive. • Aqueous solutions of acids are able to conduct electricity due to the presence of mobile ions

  7. Reactions of Acids • Reaction with metals • Dilute acids reacts with metals above hydrogen in the electrochemical series to liberate hydrogen. • Nitric acid (HNO3) doesn't usually form hydrogen with a metal, instead you get brown fumes of nitrogen dioxide! but you still get the metal nitrate salt. • exceptions: Al + H2SO4 X (Al is unreactive due to oxide layer) • Ca(s) + H2SO4(aq) CaSO4(s) + H2(g) • Ba(s) + H2SO4(aq) BaSO4(s) + H2(g) • Pb(s) + H2SO4(aq) PbSO4(s) + H2(g) • Salt formed is insoluble and coats metal preventing further reaction.

  8. Reactions of Acids • Reaction with carbonates • carbonate + acid salt + carbon dioxide + water • exception: • Aluminium carbonate does not exist. • CaCO3(s) + H2SO4(aq) CaSO4(s) + H2O(l) + CO2(g) Poor yield - similarly with PbCO3 and BaCO3. • Reaction with bases(metal oxides or hydroxides) • Dilute acids neutralise bases to form a salt and water only. • eg. 2NaOH(s) + H2SO4(aq) Na2SO4(aq) + 2H2O(l)

  9. Comparison of the properties of strong and weak acids • Property Strong Acid Weak Acid • effect on litmus red pink • action on magnesiumrapid effervescence slow effervescence of hydrogen gas of hydrogen gas • action on Na2CO3rapid effervescence slow effervescence of carbon dioxide of carbon dioxide • conductivity conducts electricity poor conductor

  10. Uses of acids • Preserving food • eg. ethanoic acid, benzoic acid, citric acid • Manufacture of products like paint, detergent and fertilizers. • eg. sulfuric acid, hydrochloric acid and nitric acid. • Batteries for vehicles • eg. sulfuric acid • removing rust from iron or steel • eg. sulfuric acid • making rubber from latex • eg methanoic acid • Pure fruit juice • eg. citric acid and tartaric acid

  11. Bases and Alkalis • A base is a substance which can react with acids to form salt and water only. eg. CuO, MgO, CaO, KOH • Bases are generally metal oxides or hydroxides. • NOTE: All basic oxides are ‘soluble’ in acids (react). But only a few are soluble in water. Basic oxides which are soluble in water form hydroxides called alkalis. • An alkali is a basic oxide which is soluble in water. • eg. sodium hydroxide, potassium hydroxide, calcium hydroxide, barium hydroxide & aqueous ammonia. • Neutralization is the process whereby an acid reacts completely with an appropriate amount of alkali to form a salt and water only. • H+ (aq) + OH- (aq) H2O (l)

  12. OH- Na+ Na+ OH- NH3 NH4+ NH3 OH- Strong and weak alkalis • A strong alkali is one that undergoes complete ionization in water. • NaOH(aq) → Na+(aq) + OH-(aq) • KOH(aq) → K+(aq) + OH-(aq) • A weak alkali is one that undergoes only partial ionization when dissolved in water. • NH3(g) + H2O NH3(aq) + H2O • NH3(aq) + H2O NH4+(aq) + OH-(aq) Note: Calcium hydroxide is considered as a weak alkali as it is only slightly soluble in water.

  13. Properties of alkalis • Alkalis have a bitter taste. • Alkalis are soapy to the touch. • Alkalis have a pH above 7 • They turn litmus blue. • Concentrated forms of potassium hydroxide and sodium hydroxide are known to be corrosive. • Aqueous solutions of alkalis are able to conduct electricity due to the presence of mobile ions.

  14. Reactions of bases / alkalis • Action on acids • All bases react with acids to form a salt and water only. • Base + Acid Salt + water • MgO + 2HNO3 Mg(NO3)2 + H2O • KOH + HCl KCl + H2O • All bases react with acidic gases to form a salt ( + water for alkalis) • CaO + CO2 CaCO3 • CuO + SO2 CuSO3 • 2NaOH + CO2 Na2CO3 + H2O

  15. Reactions of bases / alkalis • Alkalis react with ammonium salts in the presence of heat to produce ammonia (NH3) gas. • alkali + ammonium salt ammonia + salt + water • NaOH + NH4NO3 NH3 + NaNO3 + H2O • Ca(OH)2 + 2NH4Cl 2NH3 + 2H2O + CaCl2 • Alkalis precipitate many insoluble hydroxides from their salts • Alkali + salt solution precipitation reaction

  16. Point toNote: • Alkalis dissolve in water to produce hydroxide ions. • It is the hydroxide ions that give alkalis their properties

  17. Uses of alkalis • Neutralise acids • eg. tooth paste, antacids • Dissolve dirt and grease • eg. soap, detergents

  18. pH scale • A scale to measure the acidity or alkalinity of a solution. • The scale ranges from 0 - 14. • pH less than 7 => an acidic solution • pH = 7 => a neutral solution • pH more than 7 => an alkaline solution • Strong acids : eg. HCl, H2SO4 pH = 1 or 2 • Weak acids : eg. H2CO3 ,H2SO3 pH = 5 or 6 • Weak alkalis:eg. NH3(aq) pH = 9 or 10 • Strong alkalis : eg. NaOH, KOH pH = 13 or 14

  19. pH Indicators • Indicators are compounds which change colour in accordance to the pH of the medium. • Indicators are commonly used for determining the pH of colourless liquids. • Although they are not very accurate or sensitive , they give quicker results. • Each indicator has an acidic and alkaline colour and a pH at which it changes colour . Indicator Colour in medium which is pH at which strongly acidic strongly alkaline water colour changes methyl orange pink yellow yellow 4 (orange) litmus red blue red 8 (purple) phenolphthalein colourless pink colourless 10 (colourless) screened red green green 4 (grey) methyl orange

  20. Colour Changes of some Indicators pH 3.7 methyl orange pH 3.7 screened methyl orange pH 9.3 phenolphthalein pH 7 litmus pH 1 2 3 4 5 6 7 8 9 10 11 12 13 14 Universal indicator

  21. pH range 1 2 3 4 5 6 7 8 9 10 11 12 13 14 colour red orange yellow green blue purple/violet Universal Indicator • It is made up of a mixture of indicators working at different pH ranges. • The universal indicator is also known as pH indicator. • The pH of a solution can be measured by dipping a piece of universal indicator paper (pH paper) in the solution and comparing the colour obtained with a standard colour chart. Alternatively a pH meter can be used

  22. OXIDES Acidic Basic Amphoteric Neutral • Oxides of non metals • turn moist litmus red • reacts with alkalis to form salt & water • e.g. SO2, CO2, NO2 • Oxides of metals (Group 1 & II) • reacts with acids to form salt & water • e.g. Na2O, MgO, CuO • Oxides of metals • reacts with acids and alkalis to form salt & water • e.g. Al2O3, ZnO, PbO • Oxides of non metals • no effect on litmus • e.g. H2O, NO, CO

  23. Salts • A salt is a substance formed when any of the replacable hydrogen ions in an acid have been partly or completely replaced by an equivalent number of metal or ammonium ions. • Normal salts are formed when all the replaceble hydrogen ions in the acid has been completely replaced by metallic or ammonium ions • HCl + NaOH NaCl + H2O • H2SO4 + ZnO ZnSO4 + H2O • Acid salts are formed when the replaceble hydrogen ions in the acid have been partially replaced by metallic or ammonium ions. • H2SO4 + KOH KHSO4 + H2O • H2CO3 + NaOH NaHCO3 + H2O

  24. Normal salts and acid salts • Acid Acid Salts Normal Salts • HCl NaCl, CaCl2, NH4Cl • H2SO4 KHSO4, Mg(HSO4)2 K2SO4, CuSO4, NH4(SO4)2 • H3PO4 NaH2PO4, Na2HPO4 Na3PO4,, (NH4)2PO4 • HNO3 NaNO3 Mg(NO3)2, • Al(NO3)3, NH4NO3 • H2CO3 NaHCO3, Mg(HCO3)2 K2CO3, CaCO3, NH4(CO3)2

  25. Salt Preparation Soluble salts Insoluble salts Na+, K+ NH4+ salts Soluble salts that are NOT Na+, K+ NH4+ salts Precipitation Titration acid + soluble base /oxide/ carbonate Excess solid (metal, base, carbonate) + acid

  26. Points to note • The procedure used in preparing salts are such that the salts obtained are of a high degree of purity. • eg. A + B salt X + D • The method used in making salt X must be such that salt X obtained is free off excess reagent A and B and by-product D. • The method used depends on whether the salts are soluble or insoluble in water.

  27. Preparation of soluble salts • Action of an acid on a metal. • acid + metal salt + hydrogen gas • Action of an acid on an insoluble base. • acid + base salt + water • Action of an acid on an insoluble carbonate. • acid + carbonate salt + water + carbon dioxide • Action of an acid on a soluble base or carbonate (ie.Neutralization of an acid by a base - titration) • acid + base salt + water • acid + carbonate salt + water + carbon dioxide

  28. Preparation of soluble salts that do not contain Na+, K+ or NH4+ Adding Excess Solid to a fixed volume of Acid • Reacting an acid with a metal or with an insoluble base or carbonate. • This method is suitable as one of the reagents used is in the solid state and when added in excess can be easily removed by filtration

  29. The required volume of acid is measured out into the beaker with a measuring cylinder. The metal, oxide, hydroxide or carbonate is weighed out and added in small portions to the acid in the beaker with stirring. • The mixture maybe heated to speed up the reaction. When no more of the solid dissolves it means the acid is neutralised. • The hot solution is filtered to remove the excess solid • The filtrate is retained and heated in an evaporating dish to saturate it. The solution is left to cool and crystallise. • The crystals are filtered out and rinsed quickly with cold distilled water & dried between pieces of dry filter paper.

  30. Example: Preparation of Copper (II) Sulfate • Reactants: Copper (II) oxide & Sulfuric acid • Reaction: • CuO(s) + H2SO4(aq) CuSO4 (aq) + H2O(l) • Procedure: • Add copper (II) oxide powder until excess to a fixed volume of warm dilute sulfuric acid i.e. till the copper (II) oxide forms a residue at the base of the vessel. • Filter the mixture to remove the excess copper (II) oxide. • Heat the filtrate, which is copper(II) sulfate solution until it is saturated and allow it to cool. Crystals of copper (II) sulfate will form. • Filter off the crystals and dry them between two pieces of filter paper.

  31. Preparation of Na+, K+ , NH4+ salts • A known volume of alkali is pipetted into a conical flask and screened methyl orange indicator added. • The acid is titrated with the alkali in the burette until the indicator turns from green to grey. • The volume of acid needed for neutralisation is then noted, this is called the endpoint. • Steps (1-3) are repeated with both known volumes mixed together BUT without the contaminating indicator. i.e. no indicator is added • The solution is transferred to an evaporating dish and heated to saturate it. • The solution is left to cool to complete the crystallisation. • The residual liquid can be filtered away and the crystals can be carefully collected and rinsed & dried between 2 pieces of filter paper.

  32. In this method, both the reactants used are in the aqueous state thus excess reagent cannot be easily removed from the product which is in the same state. Thus exact quantities of reactants must be used to ensure that product is not contaminated by excess reagent. • The two main steps involved are: • Finding the exact quantity of reactant required for reaction by titration • Using these quantities to prepare the salt. • This method of neutralising an acid with a soluble base (e.g. sodium hydroxide) is called a titration.

  33. Example: Preparation of Sodium Chloride • Reactants: sodium hydroxide hydrochloric acid • Reaction: NaOH(aq) + HCl(aq) NaCl(aq) + H2O(l) • Procedure: • Pipette a fixed volume of sodium hydroxide into a conical flask. Add a few drops of methyl orange indicator. Slowly run the hydrochloric acid from a burette till one drop of acid added turns the indicator from yellow to orange. • Note the volume of acid added. • Repeat the whole procedure without adding the indicator. • Evaporate the solution obtained till it is saturated and allow it to cool. Crystals of sodium chloride will form. • Filter off the crystals and rinse quickly with cold distilled water & dry them between pieces of dry filter paper.

  34. Preparation of insoluble salts • Precipitation reactions • These reactions involve the addition of two soluble compounds; one containing the metallic ion and the other ion of the desired salt. • Exchange of ions between the two compounds occurs. • Formation of desired salt as a precipitate and the by-product as an aqueous solution. • The desired salt is obtained by filtration. • The salt is washed thoroughly to remove any excess reagent or by product, then dried by pressing between pieces of filter paper • eg. CuSO4 + 2NaOH Cu(OH)2 + Na2SO4 AgNO3 + NaCl AgCl + NaNO3 • Note:the reagents used must be in the form of aqueous solutions so that the insoluble salt formed being the only solid present can be completely separated from the other chemicals present. • To prepare insoluble salt XY, use: X nitrate + sodium Y insoluble salt XY + sodium nitrate

  35. Example: Preparation of Lead (II) sulfate • Reactants: sodium sulfate , lead(II) nitrate • Reaction: • Na2SO4 (aq) + Pb(NO3) 2 (aq) NaNO3(aq) + PbSO4(s) • Procedure: • Mix aqueous solutions of sodium sulfate and lead(II) nitrate together. • Filter the mixture to remove the precipitate of lead(II) sulfate formed. • Wash the residue thoroughly with distilled water. • Dry the salt between two pieces of filter paper.

  36. N S • Direct combination of two elements • This method is confined to binary salts only ie. salts made of two elements only. • eg. 2Fe(s) + 3Cl2(g) 2FeCl3 An excess of chlorine is added. • eg. Fe(s) + S(s) FeS(s) • An excess of iron is added and the excess reagent csn be removed using a magnet.

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