10. The theory of acids & bases, weak or strong acids & bases

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

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

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

Part 10 introduces students to the more advanced theory of acids and bases. Acids are defined as proton donors and bases are defined as proton acceptors. The terms 'weak' and 'strong' are explained when referring to e.g. weak acids, strong acids, weak bases or strong bases. The theories of Arrhenius and Bronsted–Lowry are described and examples given and fully explained.  These revision notes on acid-base theory and the strength of acids and bases should prove useful for the new AQA chemistry, Edexcel chemistry & OCR chemistry GCSE (9–1, 9-5 & 5-1) science courses.

Note: aq or (aq) means water or aqueous solution.

GCSE/IGCSE Acids & 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

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

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

  • Water (aq) must be present for a substance to act as an acid or as a base (at least 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 sulfuric 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 'molecular formula' neutralisation equation might be ...

    • ACID + ALKALI ==> SALT + WATER

    • e.g.

    • top sub-indexhydrochloric 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)

    • 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.

      • So the salt is formed from the residual ions when all the hydrogen ions and hydroxide ions have reacted together to give 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) + sulfuric acid (acid) ==> copper(II) sulfate + water

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

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

      • The sulfate ion is a spectator ion.

    • 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 sulfate 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

    • 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)

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

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

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

      • Amphoteric means something that can act either as an acid or as a base.

      • To illustrate water functioning as both an acid and a base ...

      • 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(aq).

      • 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.

  • top sub-indexSeveral 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 are proton donors and bases are acceptors.

      • e.g. the reaction of ammonia with acids ...

      • in the reaction ...

      • HCl(g) + NH3(g) ==> NH4+Cl(s)

      • hydrogen chloride is the acid – a proton donor leaving the chloride ion Cl

      • AND

      • ammonia is the base – a proton acceptor to form the ammonium ion NH4+

    • 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.

  • Important NOTE on the pH Scale

    (i) pH is a measure of the hydrogen ion (H+) concentration

    The lower the pH, the higher the hydrogen ion concentration, the more acidic is the solution.

    I know this seems confusing, but that's the way the pH scale has been defined historically.

    (ii) Each pH unit change is equivalent to a 10x change in concentration of the hydrogen ion

    For example changing the pH of a solution from pH 3 to pH 2 makes the solution 10x more acidic.

    Changing a solution's pH from 1 to 3 makes it 100x less acidic (10 x 10), this is similar to comparing solutions of a strong acid with that of a weak acid - its all about the extent of ionisation and the resulting concentration of hydrogen ions - read on!

  • 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.

      • Ionisation in this context means on dissolving a substance in water it splits into ions or ions are formed.

      • Strong means a high degree of ionisation (~100%) and weak means a low degree of ionisation (often <<100%).

    • 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 completely different things!

    • Dilute and concentrated refer to the concentration of the acid or alkali in terms of how much of the original material is dissolved in water as measured by concentration e.g. molarity i.e. a little or a lot in a given volume of solution, low concentration or high concentration.  It is completely independent of what concentration of hydrogen ions (in an acid) or concentration of hydroxide ions (in an alkali) is formed when the substance is dissolved in water.

      • 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, nitric and sulfuric acids.

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

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

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

      • sulfuric acid: 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 e.g. pH 0–1.

      • All three of these acids are ~100% ionised (~100% dissociated into ions) in aqueous solution, which is why they are called strong acids.

      • For a given concentration, the lower the pH, the stronger the acid, irrespective of how strong or weak the acid is.

    • Examples of strong alkalis (soluble strong bases) are sodium hydroxide or potassium hydroxide etc. (usually Group 1 like NaOH & KOH etc. or Group 2 hydroxides like calcium hydroxide, Ca(OH)2).

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

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

      • calcium hydroxide Ca(OH)2(s) + aq ==> Ca2+(aq) + 2OH(aq)
      • The greater the concentration of hydroxide ions the higher the pH, so strong alkalis make the most alkaline solutions e.g. pH 13–14.
      • Again, all of these alkalis are ~100% ionised (dissociated into free ions) in aqueous solution, that's why they are described as strong bases (alkalis).
      • For a given concentration, the higher the pH, the stronger the alkali/base, irrespective of how strong or weak the soluble base (alkali) is.

  • top sub-indexA weak acid or alkali is only partially ionised in water (often just a few %)

    • Ionisation is usually far less than 100% for strong acids or alkalis, so far less hydrogen ions or hydroxide ions formed.

    • Examples of weak acids are ethanoic acid (in vinegar), citric acid (in citrus fruits) and carbonic acids ('soda water', fizzy drinks) and these weak acids ionise via reversible reactions and an equilibrium is reached when they dissolve in water.

    • 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.

      • Ethanoic acid dissociates (ionises) into the ethanoate ion and the hydrogen ion (proton).

    • 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) ...

      • ethanoic acid: CH3COOH(aq)(c) doc bCH3COO(aq) + H+(aq)

        • this gives a pH of around 3.

        • The ionisation of a weak acid is a reversible reaction and equilibrium is set up.

        • The equilibrium position is mainly on the left-hand side, the unionised (undissociated) ethanoic acid.

      • carbonic acid: H2CO3(aq)(c) doc bHCO3(aq) + H+(aq)

      • This gives water a pH of ~3-5, it would be lower in fizzy drinks that rainwater!

      • You can use a reversible sign like to indicate the weak acid equilibrium is much more on the left side than the right side.

      • carbon dioxide in water ('carbonic acid') forms a weakly acid solution.

        • The pH of unpolluted rainwater is about pH 5.5 due to carbon dioxide dissolved from the atmosphere.

        • The steady rise in atmospheric CO2 level, is causing concern to environmentalists because along with climate change, the oceans are becoming more acidic and this does affect marine ecosystems.

      • Many organic acids are weak acids e.g. citric acid from fruit.

    • An example of a weak soluble base (weak alkali) is ammonia solution, only about 2% changes to the ionic forms on the right of the equation as written below. Sodium carbonate is also a soluble weak base in aqueous solution. Again these weak soluble bases (alkalis) ionise via reversible reactions and an equilibrium is reached when they dissolve in water.

    • 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 is still over 7, typically pH 9 to pH 11) ...

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

        • Only about 2% of the ammonia ionises (dissociates) to form ammonium ions and hydroxide ions giving a pH of ~10.

        • The ionisation of a weak base/alkali is a reversible reaction and equilibrium is set up.

        • The equilibrium position is mainly on the left-hand side, the unionised ammonia.

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

      • both of which, when dissolved in water, produce hydroxide ions (OH) 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.

    • You could compare the pH of solutions of equal concentration (equal molarity) and measure the pH with an accurately calibrated pH meter or you could get a rough estimate from universal indicator solution or paper, but that's not very accurate.

    • The differences between strong and weak acids shows up in other sorts of experiments e.g.

    • The pH of solutions of equal concentration e.g. of molarity 1.0 mol/dm3

      • The pH of a strong acid might be pH 0-1 (hydrochloric, sulfuric or nitric acids).

      • The pH of a weak acid might be typically pH 3-6 (vinegar ~pH 3, carbonic acid pH 4-5).

    • 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!

      • You can repeat the experiment using calcium carbonate (limestone granules) instead of magnesium ribbon.

      • You can do simple rate of reaction experiments comparing has fast the gas is evolved from the reaction mixture.

        • The above links takes you to a page where the experimental procedures are described, little point in repeating them here.

      • Factors affecting the rates of Reaction - theory and methods of measuring the speed of a reaction (c) Doc Brown

        • The basic experimental procedure is shown in the diagram above.

      • The above graph shows the sort of results you might expect by adding the same masses of magnesium ribbon or calcium carbonate granules to the same volume of ethanoic acid, CH3COOH, or hydrochloric acid, HCl, of equal concentration e.g. both acids with a concentration of 1.0 mol/dm3.

      • The principal observation is that the rate of reaction with the strong hydrochloric acid is much greater than the rate of reaction of the weak ethanoic acid. You can tell this from the gradient of the rate of reaction, particularly at the start of the reaction.

        • The acid equilibrium in the case of hydrochloric acid, is 100% to the right in forming hydrogen ions.

          • HCl(aq) + aq ==> H+(aq) + Cl(aq)

        • In the case of ethanoic acid, the equilibrium is ~98% to the left-hand side, so far fewer hydrogen ions are produced to react with the magnesium or limestone.

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

        • These percentages mean that the hydrogen ion concentration in hydrochloric acid is about 50 times that in ethanoic acid (100% to 2%) AND it is the hydrogen ion that actually reacts directly with the metal or the carbonate.

          • So the rate observations can be interpreted as a function of the hydrogen ion concentration.

          • The greater the hydrogen ion concentration the faster the reaction because there is an increased probability of a hydrogen ion hitting the surface of the magnesium ribbon or a limestone granule i.e. a greater collision frequency, greater chance of a fruitful collision of the right kinetic energy to change reactants into products.

          • Therefore the collision frequency of the active ingredient (hydrogen ion) is much greater in the strong acid than the weak acid.

          • On this simple numerical basis, you might expect the hydrochloric acid to react 50 times faster than the ethanoic acid (100% to 2% ionisation), applying the simple concentration rule for reaction speed.

      • What you also find, in both cases, the same final volume of gas is the same in both cases, but the strong acid completes the reaction in a much shorter time.

        • The little arrows show when the reaction ceased in each case, that is when the graph of gas volume first becomes horizontal.

        • So, you might ask the question - if ethanoic acid only ionises about 2%, how can it still produce the same volume of hydrogen or carbon dioxide as hydrochloric acid by the end?

        • Well the answer is 'quite simple', the ionisation of the weak acid continues as hydrogen ions are used up.

        • As the hydrogen ions are used up, more ethanoic acid ionises to replace those used up in the reaction to try and maintain the position of the equilibrium.

        • This is a general rule about reversible reactions and a chemical equilibrium. The system will always try to maintain a balance in the ratio of reactants to products (reactants (c) doc b products), if you remove a component from the equilibrium, the system tries to replace it.

        • This contrasts with the strong acid solution where all possible hydrogen ions exist right from the start of the reaction.

      • The equations for these reactions producing the gases hydrogen or carbon dioxide are as follows ...

        • (fast) magnesium + hydrochloric acid  ==> magnesium chloride + hydrogen

          • Mg + 2HCl  ==>  MgCl2 + H2

        • (slow) magnesium + ethanoic acid  ==>  magnesium ethanoate

          • Mg + 2CH3COOH  ==> Mg(CH3COO)2 + H2

        • (fast) calcium carbonate + hydrochloric acid  ==> calcium chloride + water + carbon dioxide

          • CaCO3 + 2HCl ==> CaCl2 + H2O + CO2

        • (slow) calcium carbonate + ethanoic acid ==> calcium ethanoate + water + carbon dioxide

          • CaCO3 + 2CH3COOH ==> Ca(CH3COO)2 + H2O + CO2

        • -

    • 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 of hydrochloric acid and ethanoic acid, 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 length time at the same voltage (potential difference, p.d.) before measuring the gas volumes of hydrogen formed. (Electrolysis methods).

      • In the solution undergoing electrolysis, the current is carried by the charged particles e.g. hydrogen ions, hydroxide ions, chloride ions, ethanoate ions etc., therefore the greater the ionisation of the acid the greater the concentration of ions.

      • The greater concentration of ions in the strong acid solution reduces the electrical resistance and more current flows via the greater number of ions present to carry it, hence more hydrogen ions are reduced at the cathode to form hydrogen.

        • Setting the power supply to the same voltage and incorporating an ammeter into the electrolysis cell circuit, you can readily see the difference in the current readings between the strong acid solution (high A reading) and the weak acid solution (low A reading).

      • They will of course, both produce hydrogen in electrolysis cells, because all acids produce some hydrogen ions in water, but the rate of electrolysis is highly dependant on the hydrogen ion concentration, which in turn is dependant on the strength of the acid using solutions of the same concentration (as measured by molarity).

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

  • top sub-indexThe 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 0–2 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.

  • SUMMARY of the advanced BRONSTED LOWRY THEORY

    • 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

        • 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.

 



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