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pH scale, neutralisation, ionic theory of acids & alkalis

Doc Brown's Chemistry  GCSE/IGCSE Science–Chemistry Revision Notes

pH scale of acidity and alkalinity, acids, bases–alkalis, salts and neutralisation

2. pH Scale, indicators, acids, alkalis (bases), neutralisation its ionic theory introduction

This page introduces and explains the pH scale measuring the relative acidity and alkalinity of aqueous solutions, that is solutions of substances dissolved in water. The use of indicators is described and several well known indicators are tabulated showing their different colours in solutions of different pH. The ionic theory of acids, bases and neutralisation is simply described and why explains why solutions are either acid, neutral or alkaline.

GCSE/IGCSE Acid & Alkalis revision notes sub–index: Index of all pH, Acids, Alkalis, Salts Notes 1. Examples of everyday acids, alkalis, salts, pH of solution, hazard warning signs : 2. pH scale, indicators, ionic theory of acids–alkali neutralisation : 4. Reactions of acids with metals/oxides/hydroxides/carbonates, neutralisation reactions : 5. Reactions of bases–alkalis like ammonia & sodium hydroxide : 6. Four methods of making salts : 7. Changes in pH in a neutralisation, choice and use of indicators : 8. Important formulae of compounds, salt solubility and water of crystallisation : 10. More on Acid–Base Theory and Weak and Strong Acids

See also Advanced Level Chemistry Students Acid–Base Revision Notes – use index

2. The pH scale, indicators acids, alkalis (bases), neutralisation & ionic theory


 

2a. Introduction to the pH scale What is the pH scale?

(c) doc b

The colours observed in solutions when universal indicator is added

  • The pH scale is a measure of the relative acidity or alkalinity of a solution (see diagram).

    • So, knowing the pH of a solution, you know how acid or alkaline it is by reference to the pH scale (diagram above) or whether the solution is neutral.

    • The smaller the pH number, the more acid it is, the greater the pH number, the more alkaline it is, and if the pH is close to 7, you have a more or less neutral solution that has neither acidic or alkaline chemical properties.

    • Lots of examples of solution pH values are tabulated with everyday examples of acid/alkaline chemistry are described in section 1.

  • pH can be approximately measured using indicator solution by putting a few drops of universal indicator into a solution and comparing the colour formed with a standard chart (picture above).

  • You can also used paper impregnated with an indicator solution (pH paper), the paper is dipped in the solution and again the colour matched with a pH chart.

    • This is quite handy for testing soil mixed and shaken with water.

    • You can get special soil testing kits which use indicator solution and the colour of the indicator in the water is matched with a chart after the soil has settled out.

  • pH can be very accurately measured with a special instrument called a pH meter using a glass electrode probe which is calibrated with standard buffer solutions of accurately known pH (see photographs and note at the end of the page).

  • WHAT IS AN INDICATOR?

  • An indicator is a dye substance or mixture of coloured substances that when added to the solution gives a different colour depending on the pH of the solution.

    • Universal indicator solution or paper, is prepared from mixing several indicators to give a variety of colours to match a wide range of pH values from very acid to very alkaline.

    • The mixture of dyes responds to changes in pH, so depending on what the pH is, i.e. how acid, how alkaline or neutral the solution is, the indicator tells which it is.

      • Not only that, an indicator like universal indicator' can tell you how strongly acid or strongly alkaline the solution is by giving you the pH to about the nearest indicator.

      • A calibrated instrument called a pH meter can give the pH to two decimal places.

        • pH meters are calibrated using buffer solutions which have an accurately known pH.

    • It is a very handy indicator for showing whether the solution is very weakly/strongly acidic (pH <7) or alkaline (pH > 7) or neutral (pH = 7) and gives the pH to the nearest pH unit.

    • (c) doc b

    • The diagram above gives the sort of range of colours you get from using universal indicator, which is a complex mixture of different dye molecules that respond to changes in pH.

  • Theoretically there is no limit to the pH scale, but most solutions are between pH 0 and pH 14.

    • For example, looking at the 'extremes', 1M hydrochloric acid (HCl) has a pH of 0 and 10M HCl has a pH of –1 and these would be described as strongly acidic solutions.

    • 1M sodium hydroxide (NaOH) has a pH of 14, but 10M potassium hydroxide (KOH) has a pH of 15 and these would be described as strongly alkaline solutions.

    • The closer the pH is to 7, the less strong is the acid or alkali.

    • However the solubility limits of substances in water ensures that its almost impossible to get below –1 or above 15 and most laboratory measurements will be in the range pH 1 to pH 14

    • .
  • Note 1: M is the old shorthand for solubility in mol/litre or mol dm–3.

  • Note 2: The pH scale is known as a logarithmic scale of base 10.

    • At GCSE/IGCSE level, to put it more simply, a change of one pH unit means a 10x change in the acidity or alkalinity of the solution

      • e.g. from pH 5 to pH 2 means an increase in acidity of 1000x

      • or to change from pH 13 to pH 11 means to become 100x less alkaline.).

Other common indicators used in the laboratory (c) doc b often used in titrations – e.g. salt preparation (a)

Indicator

colour in acid pH<7 colour in neutral pH=7 colour in alkali pH >7

litmus

red 'purple' blue
phenolphthalein* colourless colourless >9 pink
methyl orange* <3.5 red, orange about pH 5, > 6 yellow yellow yellow
methyl red* <5 red, orange, >6 yellow yellow yellow
bromothymol blue* <6 yellow green >8 blue

 


 

2b. Introduction to Acid–Base (including Alkalis) Theory including Neutralisation

  • Water is a neutral liquid with a pH of 7 (green with universal indicator).

  • When a substance dissolves in water it forms an aqueous (aq) solution that may be acidic, neutral or alkaline.

  • Acidic solutions have a pH of less than 7, and the lower the number, the stronger the acid it, or the more acidic the solution.

    • The colour can range from orange–yellow (pH 3–6) for partially ionised weak acids like ethanoic acid (vinegar) and carbonated water.

    • Strong acids like hydrochloric, sulphuric and nitric are fully ionised and give a pH 1 or less and a red colour with universal indicator or litmus paper.

  • Neutral solutions have a pH of 7. These are quite often solutions of salts, which are themselves formed from neutralising acids and bases.

  • The 'opposite' of an acid is called a base. Some bases are soluble in water to give alkaline solutions – these are known as alkalis.

  • Alkaline solutions have a pH of over 7 and the higher the pH the stronger is the alkali, the more alkaline is the solution.

    • Weak alkalis (soluble bases) like ammonia give a pH of 10–11 but strong alkalis (soluble bases) like sodium hydroxide give a pH of 13–14.

    • Alkalis give blue–purple–violet colour with universal indicator or litmus paper.

  • NEUTRALISATION usually involves mixing an acid (pH <7) with a base or alkali (pH > 7) which react to form a neutral SALT solution of pH ~7

    • in general the word equation for a neutralisation reaction is

      • ACID + BASE/ALKALI ===> SALT + WATER

      • An alkali is a soluble base, an insoluble base is NOT an alkali.

      • All bases, soluble or insoluble reaction with acids in a neutralisation reaction to form a salt like compound.

 


 

2c. More advanced Acid–Base Theory

  • THE IONIC THEORY of ACIDS and ALKALIS – a brief introduction and a few technical terms

    • The proton (H+) donation–acceptance theory of acids and bases (Bronsted–Lowry) is covered in Section 10.

      • Part 10 "More on acid–base theory",

      • but here, I'm explaining the theory in the simplest way with the minimum of detail.

      • Ions are charged particles that carry an overall net positive electric charge e.g. 2+, +, – or 2– etc.

      • When a substance dissolves in water the total number of positive charges on the positive ions must equal the total number of negative charges on the negative ions.

    • Acids are substances that form hydrogen ions, H+(aq), when dissolved in water e.g.

      • hydrochloric acid HCl gives H+(aq) and Cl(aq) ions in water.

        • (aqueous solution of hydrogen ions and chloride ions, pH reduced to <7)

      • sulfuric/sulphuric acid H2SO4 gives 2H+(aq) and SO42–(aq) ions in water.

        • (aqueous solution of hydrogen ions and sulfate ions, pH reduced to <7)

      • nitric acid HNO3 gives H+(aq) and NO3(aq) ions in water.

        • (aqueous solution of hydrogen ions and nitrate ions)

    • Alkalis are substances that form hydroxide ions (OH(aq)) in water e.g.

      • sodium hydroxide NaOH gives Na+(aq) and OH(aq) ions in water.

        • (aqueous solution of sodium ions and hydroxide ions, pH increased to >7)

      • calcium hydroxide Ca(OH)2 gives Ca2+(aq) and 2OH(aq)  ions in water.

        • (aqueous solution of calcium ions and hydroxide ions, pH increased to >7))

      • Note that an alkali is a base soluble in water.

        • An insoluble base like copper(II) oxide, CuO, will NOT affect the pH of water (pH 7 neutral), i.e. it will not cause the formation of either hydrogen ions or hydroxide ions on mixing with water BUT it will still neutralise acids in forming a soluble salt.

    • The majority of liquid water consists of covalent H2O molecules, but there are trace quantities of H+ and OH ions from the self–ionisation of water,

      • H2O(l) H+(aq) + OH(aq)

        • Only about 1 in 200 million water molecules does this!, the reaction is reversible (hence the (c) doc b sign), so the longer half–arrow to the left tells you that most water remains as water molecules!

        • Also note that hydrogen ion is sometimes described as a proton.

      • BUT, logically, this means that the hydrogen ion concentration must equal the hydroxide ion concentration, so they are of equal concentration and so water is neutral at pH 7.

    • In acid solutions there are more H+ ions than OH ions, so an excess of hydrogen ions makes the solution acidic with a pH of less than 7.

    • In alkali solutions there are more OH ions than H+ ions, so an excess of hydroxide ions makes the solution an alkaline with a pH of over 7.

    • When alkalis and acids react, the 'general word' or 'molecular formula' equation might be for NEUTRALISATION ...

      • ACID + ALKALI ==> SALT + WATER

      • e.g.

      • hydrochloric acid + sodium hydroxide ==> sodium chloride + water

      • HCl(aq) + NaOH(aq) ==> NaCl(aq) + H2O(l)

      • BUT the ionic equation for ANY neutralisation involving the reaction between an acid and alkali is

      • hydrogen ion + hydroxide ion ==> water

      • H+(aq)  + OH(aq)  ==> H2O(l)

      • Because all acids form hydrogen ions in water and all alkalis (soluble bases) form hydroxide ions in water.

      • So the 'acidic' hydrogen ions cancel out the 'alkaline' hydroxide ions by combining to form neutral water, AND give a neutral solution of a salt.

      • and, in this case, the remaining ions e.g. sodium Na+(aq) and chloride Cl(aq) become the salt crystals of sodium chloride NaCl(s) on evaporating the water.

        • So the salt is formed from the residual ions when all the hydrogen ions and hydroxide ions have reacted.

        • In this simple case the sodium ions and chloride ions don't take part in the reaction and are known as spectator ions.

        • BUT, on evaporation of the solution, the sodium ions and chloride ions will come together and crystallise out of solution as the 'salt' sodium chloride.

    • BASES e.g. oxides, hydroxides and carbonates, are substances that react and neutralise acids to form salts and water.

      • Bases which are soluble in water are called alkalis e.g. NaOH sodium hydroxide, KOH potassium hydroxide or Ca(OH)2 calcium hydroxide.

        • The reaction described above is a simple and good example of an acid neutralising an alkali.

      • Bases which are water insoluble include CuO copper(II) oxide, MgO magnesium oxide and these will also react and dissolve in acids to form salt solutions e.g.

        • ACID + BASE ==> SALT + WATER

        • copper oxide + sulfuric acid ==> copper sulfate + water

        • H2SO4(aq) + CuO(s) ==> CuSO4(aq) + H2O(l)

    • After a neutralisation, the salt solutions consist of a mixture of positive and negative ions (and their names are in the salt name!) e.g.

      • sodium chloride (NaCl) is a mixture of  Na+ and Cl ions in the ratio 1:1 (from hydrochloric acid and sodium hydroxide)

      • calcium chloride (CaCl2) is a mix of Ca2+ and Cl ions of ratio 1:2 (from hydrochloric acid and calcium oxide/hydroxide)

      • magnesium nitrate (Mg(NO3)2) is a mix of Mg2+ and NO3 ions in the ratio 1:2 (from nitric acid and magnesium oxide/hydroxide)

      • aluminium sulphate (Al2(SO4)3) consists of Al3+ and SO42– ions in the ratio 2:3 (from sulfuric acid and aluminium oxide/hydroxide)

        • and when the water is evaporated the oppositely charged ions combine to form the crystalline salt (names above).

    • More in the notes (c) doc b GCSE/IGCSE chemistry higher level notes for the advanced proton/hydrogen ion theory of acids and bases (Bronsted–Lowry theory

APPENDIX - ACCURATELY MEASURING THE pH OF A SOLUTION

You can measure the pH of a solution very accurately using a pH meter and a glass membrane pH probe.

The pH meter is calibrated against a standard buffer solution of accurately known pH

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