* KS4 Science GCSE Chemistry Notes 10. More on acid-base theory and weak & strong acids at Doc Brown's

DOC BROWN'S HOMEPAGE and WELCOME ALPHABETICAL SITE INDEX for chemistry KS3 SCIENCE QUIZZES and WORKSHEETS (~US grades 6-8) KS4 Science GCSE/IGCSE CHEMISTRY NOTES (~US grades 8-10) KS4 Science GCSE/IGCSE CHEMISTRY QUIZZES and WORKSHEETS (~US grades 8-10) KS4 Science GCSE/IGCSE CHEMISTRY SYLLABUS HELP LINKS (~US grades 8-10) ADVANCED LEVEL CHEMISTRY NOTES (~US grades 11-12) ADVANCED LEVEL CHEMISTRY QUIZZES and WORKSHEETS (~US grades 11-12) ADVANCED LEVEL CHEMISTRY SYLLABUS HELP LINKS (~US grades 11-12)  
TEACH YOURSELF CHEMISTRY

 Doc Brown's Chemistry

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

10. More on acid-base theory and weak & strong acids

Revision Notes KS4 Science IGCSE/O level/GCSE Chemistry Information Study Notes for revising for AQA GCSE Science, Edexcel 360Science/IGCSE Chemistry & OCR 21stC Science, OCR Gateway Science  (revise courses equal to US grades 9-10)

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

10. More on Acid-Base Theory and Weak and Strong Acids

  • Some compounds react will water to produce acidic or alkaline solutions.

  • Water must be present for a substance to act as an acid or as a base (usually at gcse level!).

  • Acids in aqueous solution produce hydrogen H+ ions. The H+ ion is a proton. In water this proton is hydrated (associated with water and more correctly expressed as H3O+(aq)) but H+(aq) is adequate here. The greater the concentration of hydrogen ions the more acid the solution and the lower the pH.

    • e.g. hydrochloric acid: HCl(g) + aq ==> H+(aq) + Cl-(aq)

    • or sulphuric acid: H2SO4(l) + aq ==> 2H+(aq) + SO42-(aq)

  • Alkalis in aqueous solution produce OH-(aq) hydroxide ions.  The greater the concentration of hydroxide ions the more alkaline the solution and the higher the pH.

    • e.g. sodium hydroxide: NaOH(s) + aq ==> Na+(aq) + OH-(aq)

    • or calcium hydroxide: Ca(OH)2(s) + aq ==> Ca2+(aq) + 2OH-(aq)

  • When alkalis and acids react, the 'general word' and e.g. 'molecular formula' neutralisation equation might be ...

    • 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

    • H+(aq)  + OH-(aq)  ==> H2O(l)

    • and the remaining ions e.g. Na+(aq) and Cl-(aq) become the salt crystals NaCl(s) on evaporating the water.

  • Acids can be defined as proton donors. A base can be defined as a proton acceptor (Bronsted-Lowry theory).

    • e.g. here the hydroxide ion is the base and accepts a proton from an acid.

      • H+(aq) + OH-(aq) ==> H2O(l)

    • or here the hydrogen chloride is the acid and the ammonia is the base when ammonium chloride is formed when the two gases are mixed. The acid hydrogen chloride donates a proton to the base ammonia. (note: no water present!)

      • HCl(g) + NH3(g) ==> NH4+Cl-(s)

    • or copper(II) oxide (base) + sulphuric acid (acid) ==> copper(II) sulphate + water

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

    • ionically it is: Cu2+O2-(s) + 2H+(aq) ==> Cu2+(aq) + H2O(l) 

    • Acids are characterised by having at least one replaceable hydrogen atom in forming a salt, the H is replaced by a metal ion (Na+, Mg2+ etc.) or the ammonium ion (NH4+):

      • e.g. for acid => sodium salt or salts (from Na2O, NaOH, NaHCO3 or Na2CO3)

        • HNO3 ==> NaNO3  

          • only one nitrate salt possible from one replaceable 'hydrogen'

        • HCl ==> NaCl  

          • only one chloride salt possible from one replaceable 'hydrogen'

        • H2SO4 ==> NaHSO4 ==> Na2SO4  

          • two sulphate salts possible from two replaceable 'hydrogens'

        • H3PO4 ==> KH2PO4 ==> K2HPO4 ==> K3PO4 

          • three phosphate salts possible from three replaceable 'hydrogens'

  • 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

        • NH3(aq) + H2O(l) (c) doc b NH4+(aq) + OH-(aq)

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

  • Several scientists have made contributions to ionic and acid-base theory e.g.

    • Arrhenius (1887), was one of the first scientists to suggest that substances could split into free positive and negative ions when dissolved in water, the so called 'electrolytic dissociation' giving rise to electrically conducting solutions. His theory was considered a bit revolutionary, and he was given a low rating for his PhD at Paris at first! - however the 'professors' recanted when other scientists decided it was a good idea and in 1903 he was awarded the Nobel Prize for his ionic theory work! 

    • Lowry and Bronsted (1923) took further the work of Arrhenius and applied ionic theory to the concept of acids and bases - that is, that acids and bases are proton donors and acceptors (see above).  It should be noted that the work of Arrhenius took much longer to be accepted than the work of Lowry and Bronsted because there was no pre-existing (and proven) theory of ion formation.

  • Acids and alkalis are further classified by the extent of their ionisation in water.

    • They are described as strong or weak depending on their degree of ionisation in water.

    • Do not confuse the terms weak and strong about how far the 'molecules' become ionised in water with the terms dilute and concentrated, they mean different things!

    • Dilute and concentrated refer to the concentration of the acid or alkali in terms of how much (i.e. a little or a lot) of the original material is dissolved in water as measured by concentration e.g. molarity.

      • You need to read on and then return here to clarify the points.

  • A strong acid or alkali is one that is that is nearly or completely 100% ionised in water (not an equilibrium situation)

    • examples of strong acids are hydrochloric, sulphuric and nitric acids.

      • e.g. the maximum (or nearly) hydrogen ion concentration results in the lowest pH ...

      • nitric acid is: HNO3(l) + aq ==> H+(aq) + NO3-(aq)

      • and sulphuric acid is: H2SO4(l) + aq ==> 2H+(aq) + SO42-(aq)

      • The greater the concentration of hydrogen ions the lower the pH, so strong acids make the most acidic solutions.

    • examples of strong alkalis (soluble strong bases) are sodium hydroxide or potassium hydroxide etc. (usually Group 1 or 2 hydroxides).

      • e.g. the maximum (or nearly) hydroxide ion concentration results in the highest pH ...

      • potassium hydroxide is: KOH(s) + aq ==> K+(aq) + OH-(aq)

      • or strontium hydroxide is: Sr(OH)2(s) + aq ==> Sr2+(aq) + 2OH-(aq)
      • The greater the concentration of hydroxide ions the higher the pH, so strong alkalis make the most alkaline solutions.

  • A weak acid or alkali is only partially ionised in water.

    • examples of weak acids are ethanoic, citric and carbonic acids.

    • e.g. for ethanoic about 2% ionises (forward reaction to the right), the equilibrium lies mainly to the un-ionised form on the left and for the weaker carbonic acid even less is ionised. So only a relatively low concentration of free hydrogen ions form giving a less acidic higher pH solution than strong acids (but pH still less than 7) ...

      • CH3COOH(aq)(c) doc bCH3COO-(aq) + H+(aq)

      • H2CO3(aq)(c) doc bHCO3-(aq) + H+(aq)

    • An example of a weak alkali/base (weak soluble base) is ammonia solution, about 2% changes to the ionic forms on the right. So only a relatively low concentration of free hydroxide ions form giving a less alkaline solution, so the pH is less than a strong base/alkali (but pH still over 7) ...

      • NH3(aq) + H2O(l)(c) doc bNH4+(aq) + OH-(aq)

      • or sodium carbonate: CO32- + H2O(l) (c) doc b HCO3-(aq) + OH-(aq)

      • both of which, when dissolved in water, produce hydroxide ions giving an alkaline solution, despite the fact that OH doesn't appear in their formulae!

  • You can distinguish between strong and weak acids of the same concentration by using the pH scale and observations from a variety of experiments support the low or high of ionisation theory.

    • e.g. by the rate of reaction with metals.

      • If you put magnesium ribbon into 1 molar solutions of hydrochloric acid (strong, high % ionisation so high H+(aq) concentration) and ethanoic acid (weak, low percentage ionization so much lower H+(aq) concentration), you can see the difference in the fast and slow 'fizzing' rates!

    • Since stronger/weak acid solutions (or alkalis) contain more/less hydrogen ions, they are better/poorer conductors of electricity.

      • e.g. If you carry out electrolysis experiments with the same two solutions, you get a much greater volume of hydrogen collected at the cathode from the hydrochloric acid compared to the ethanoic acid.

      • You must use solutions of the same concentration and electrolysed them for the same time before measuring the gas volumes (Electrolysis methods 1a and 1b).

    • Remember that its the H+ ion that is the active chemical species in acid solutions NOT a 'HCl' or a 'H2SO4' or a 'CH3COOH' molecule.

    • More on pH scale and indicators in section 2.

  • The pH is dependent on the relative concentrations of the H+(aq) and the OH-(aq) concentrations.

    • a high H+(aq) concentration means a low pH

      • and low OH-(aq) concentration, usually strong acid

    • lower H+(aq) concentration means higher pH and higher OH-(aq) concentration, less acid

    • a high OH-(aq) concentration means a high pH

      • and low H+(aq) concentration, usually strong base/alkali

    • lower OH-(aq) concentration means lower pH and higher H+(aq) concentration, less alkaline

  • In general:

    • pH 1-2 strong acids

    • pH 3-6 weak acids

    • pH 7 neutral

    • pH 8-11 weak base/alkali

    • pH 12-14 strong base/alkali

  • Neutralisation ionically is: H+(aq) + OH-(aq) ==> H2O(l) (exothermic)

    • The pH of a solution, or determining the neutralisation point, can be measured with

      • an indicator colour comparison card or indicator added to a titration

      • a pH meter which is calibrated with 'buffer solutions' of exactly know pH.

    • When mixing an acid and alkali the neutralisation end-point can also be determined by

      • the point of maximum temperature rise.

  • Further work-study links:

 

Advanced Chemistry Page Index and LinksWebsite content copyright © Dr W P Brown 2000-2010 All rights reserved on revision notes, puzzles, quizzes, worksheets, x-words etc. * Copying of website material is not permitted * I do not personally endorse the adverts - but they do pay for the site!
TEACH YOURSELF CHEMISTRY