Doc Brown's Chemistry Revision Study Notes

The pH scale of acidity and alkalinity, acids, alkalis, salts and neutralisation

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

Advanced Level Chemistry Acid-Base Revision Notes - use index

GCSE Sub-index: Index of all pH, Acids, Alkalis, Salts Notes 1. Examples of acid-alkali chemistry : 2. pH scale, indicators, ionic theory of acids-alkali neutralisation : 3. pH examples of acid, neutral or alkaline solutions : 4. Acid reactions with metals/oxides/hydroxides/carbonates and neutralisation reactions : 5. Reactions of bases-alkalis like sodium hydroxide : 6. Four methods of making salts : 7. Changes in pH in a neutralisation : 8. Important formulae, salt solubility and water of crystallisation : 9. Further examples of word/symbol equations for salt preparations : 10. More on Acid-Base Theory and Weak and Strong Acids

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

• 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 which is calibrated with buffer solutions of accurately known pH.

• An indicator is a substance or mixture of 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 the 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.

• 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.

  • 1M sodium hydroxide (NaOH) has a pH of 14, but 10M potassium hydroxide (KOH) has a pH of 15.

  • 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.).

• 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.

  • The colour can range from orange-yellow (pH 3-6) for partially ionised weak acids like ethanoic acid (vinegar)  to 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.

  • 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 eqution for a neutralisation reaction is

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

• More details below

• 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. "More on acid-base theory"

  • 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)

    • sulfuric/sulphuric acid H₂SO₄ gives 2H⁺_(aq) and SO₄^2-_(aq) ions in water (aqueous solution)

    • nitric acid HNO₃ gives H⁺_(aq) and NO₃^-_(aq) ions in water (aqueous solution)

  • 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)

    • calcium hydroxide Ca(OH)₂ gives Ca²⁺_(aq) and 2OH^-_(aq)  ions in water (aqueous solution)

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

  • 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, BUT they are of equal concentration and so water is neutral at pH 7.

  • In acid solutions there are more H⁺ ions than OH^- ions.

  • In alkaline solution there are more OH^- ions than H⁺ ions.

  • 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 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.

    • 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.

  • 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.

    • Bases which are water insoluble include CuO copper(II) oxide, MgO magnesium oxide.

  • 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

    • calcium chloride (CaCl₂) is a mix of Ca²⁺ and Cl^- ions of ration 1:2

    • magnesium nitrate (Mg(NO₃)₂) is a mix of Mg²⁺ and NO₃^- ions in the ratio 1:2

    • aluminium sulphate (Al₂(SO₄)₃) consists of Al³⁺ and SO₄^2- ions in the ratio 2:3

  • See other GCSE/IGCSE chemistry higher level notes for the advanced proton/hydrogen ion theory of acids and bases (Bronsted-Lowry theory)

    • The main concept of the advanced Bronsted-Lowry theory is ...

    • a Bronsted-Lowry acid is defined as a proton donor (H⁺),

    • and a Bronsted-Lowry base is defined as a proton acceptor e.g. two examples

      • (i) hydrogen chloride gas + ammonia gas ==> ammonium chloride solid

      • HCl(g) + NH₃(g) ==> NH₄Cl(s)

      • acidic hydrogen chloride gives a proton to the ammonia molecule base to give the ammonium ion (NH₄⁺).

      • (ii) copper oxide dissolves in acid solutions

      • copper(II) oxide + sulfuric acid ==> copper(II) sulfate + water.

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

      • copper oxide is the base because it reacts with protons from the acid to form water.

    • Incidentally water is a neutral oxide because its pH is 7.

    • However water is an amphoteric oxide i.e. it reacts as both a proton acceptor and a proton donator.

      • e.g. water acting as a base - proton acceptor with a stronger acid like the hydrogen chloride gas

        •  HCl_(g) + H₂O_(l) ==> H₃O⁺_(aq) + Cl^-_(aq)

        • This is how hydrochloric acid is formed which you write simply as HCl.

      • e.g. water acting as an acid - proton donor with a weak BUT stronger base like the alkaline gas ammonia

        • NH_3(aq) + H₂O_(l) NH₄⁺_(aq) + OH^-_(aq)

        • This is why ammonium solution is alkaline - sometimes wrongly called 'ammonium hydroxide' instead of aqueous ammonia.

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