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.

10(a) More on defining acids and bases - the results of ionisation

• 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!).

• A substance dissolving in water and splitting up into ions undergoes an example of ionisation.

• 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: H₂SO_4(l) + aq ==> 2H⁺_(aq) + SO₄^2–_(aq)

    • Note: + aq is simply 'shorthand' for adding to water!

• 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)

  • Reminder: An alkali is a soluble base.

• When alkalis and acids react, the 'general word' and '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)

  • 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.

10(b) A more advanced acid–base theory

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

  • The hydrogen ion (H⁺) is equivalent to a proton - you can use either term.

  • e.g. here the hydroxide ion is the base and accepts a proton from an acid to form water (as in neutralisation).

    • H⁺_(aq) + OH^–_(aq) ==> H₂O_(l)

    • This simple ionic equation represents what happens when any soluble alkali (metal hydroxides like NaOH) is mixed with any acid (like HCl etc.).

    • Water is formed, and the ions left in solution constitute the ions that make up a salt that can be crystallised out of solution by evaporating the water.

  • 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, because non formed!)

    • HCl_(g) + NH_3(g) ==> NH₄⁺Cl^–_(s)

    • You can write an almost identical equation when aqueous ammonia is mixed with dilute hydrochloric acid.

    • HCl_(aq) + NH_3(aq) ==> NH₄Cl_(aq)

    • You can evaporate and crystallise out ammonium chloride crystals.

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

  • CuO_(s) + H₂SO_4(aq) ==> CuSO_4(aq) + H₂O_(l)

  • ionically it is: Cu²⁺O^2–_(s) + 2H⁺_(aq) ==> Cu²⁺_(aq) + H₂O_(l) 

    • The copper oxide is the base because it accepts hydrogen ions from the acid.

    • The sulfate ion is a spectator ion.

    • Copper(II) oxide is an insoluble base, therefore can't be an alkali.

  • You can analyse what happens when a carbonate dissolves in an acid e.g.

  • magnesium carbonate + sulfuric acid ==> magnesium sulfate + water + carbon dioxide

    • MgCO₃ + H₂SO₄ ==> MgSO₄ + H₂O + CO₂

    • Sulfuric acid donates hydrogen ions to the base ion - the carbonate ion,

    • so in terms of ions, the reaction between ANY carbonate and ANY acid can be written as:

    • 2H⁺  +  CO₃^2-  ===> H₂O  +  CO₂

  • 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)

    • (i) HNO₃ ==> NaNO₃  

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

    • e.g. sodium nitrate from neutralising nitric acid with sodium hydroxide.

    • (ii) HCl ==> NaCl  

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

    • e.g. sodium chloride from neutralising hydrochloric acid with sodium hydroxide.

    • (iii) H₂SO₄ ==> NaHSO₄ ==> Na₂SO₄  

    • two sulfate salts possible from two replaceable 'hydrogens'

    • (they are actually called sodium hydrogensulfate from half-neutralising sulfuric acid and sodium sulphate from complete neutralisation with e.g. sodium hydroxide)

    • (iv) H₃PO₄ ==> KH₂PO₄ ==> K₂HPO₄ ==> K₃PO₄ 

    • three phosphate salts possible from three replaceable 'hydrogens', so the sequence gives the basic formula of each successive salt as each hydrogen of the acid is replaced.

• 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 water molecules does this!, the reaction is reversible ( sign), so the longer half–arrow to the left tells you that most water remains as water molecules!]

    • BUT there are equal concentrations of hydrogen ions and hydroxide ions, 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)

      •  more simply: HCl_(g) + _aq) ==> H⁺_(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

      • 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.

• 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 are proton donors and bases are acceptors.

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

    • in the reaction ...

    • HCl_(g) + NH_3(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 at the time there was no pre–existing (and proven) theory of ion formation.

10(c) Some extra important notes on the pH scale that you need to know

(a) 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.

 

(b) 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!

 

(c) The pH scale and acid concentration (advanced level students only)

Its quite easy to relate the pH of an acid solution from its concentration of hydrogen ions

(without going too much into the maths!).

 

The concentration of hydrogen ions (mol/dm³) = 10^-pH  (ONLY advanced level students need to know all this maths)

So, if the pH is 3.0, the hydrogen ion concentration = 10^-3 or 0.001 mol/dm^-3

If the pH is 1.0, the hydrogen ion concentration = 10^-1 or 0.1 mol/dm^-3

You can then work the argument the other way: (pH = -log₁₀(H+ concentration))

If the concentration of hydrochloric acid is 0.5 mol/dm³, strong acid,

so hydrogen ion concentration is also 0.5 mol/dm³

the pH is -log₁₀(0.5) = 0.30

Please note that GCSE students do not have to do these calculations - but perhaps make an estimate?

Advanced A level pre-university chemistry notes - equilibrium section 5 (NOT for GCSE!)

5.1 Lewis and Bronsted–Lowry acid–base theories

5.2 self–ionisation of water and pH scale

5.3 strong acids - examples - Ph calculations

5.4 weak acids - examples & pH-K_a-pK_a calculations

5.5 strong bases -examples -pH calculations

5.6 weak bases - examples & pH-K_b-pK_b calculations

10(d) Defining and explaining weak and strong acids or alkalis (and avoid a confusion!)

• 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 positive and negative ions.

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

    • Examples are explained below, but after avoiding misunderstanding and misusing terms like dilute and concentrated.

  • BUT FIRST ....

  • Do not confuse the terms weak and strong, which is 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.

    • The original material might be: HCl, HNO₃, H₂SO₄, NaOH, Ca(OH)₂, and dilute and concentrated refer to how many particles of the formula are in a given volume e.g.

    •  

    • The above diagrams illustrate a less concentrated (more dilute) solution on the left and a more concentrated solution (less dilute) on the right. The fact that the diagrams show a mixture of chemicals A and B is irrelevant.

  • 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 - but that's when the terms 'strong' and 'weak' are necessary.

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

• So, what do we mean by the terms 'weak' and 'strong' when talking about acids (and soluble bases-alkalis)

  • A lot to take in reading section 10(d)

• A strong acid or strong 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)

      • the molecules of HCl are completely dissociated into hydrogen ions and chloride ions. (aq just means water)

    • nitric acid: HNO_3(l) + aq ==> H⁺_(aq) + NO₃^–_(aq)

      • the molecules of HNO₃ are completely dissociated into hydrogen ions and nitrate ions.

    • sulfuric acid: H₂SO_4(l) + aq ==> 2H⁺_(aq) + SO₄^2–_(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 for a given concentration.

    • 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.

    • Its all about the extent to which the original acid molecules are dissociated into ions.

  • 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.

• A 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)CH₃COO^–_(aq) + H⁺_(aq)

      • undissociated acid ~98% <===> dissociated acid ions ~2%

      • 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: H₂CO_3(aq)HCO₃^–_(aq) + H⁺_(aq)

      • undissociated acid <===> dissociated acid ions

      • 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

        • (H₂O + CO₂ ==> H₂CO₃).

      • 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) ...

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

      • Only about 2% of the ammonia ionises 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: CO₃^2– + H₂O_(l) 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!

• Take care in combing the two sets of terms e.g. you can have:

  • a dilute solution of a weak acid e.g. 0.1 mol/dm³ solution of ethanoic acid

    • CH₃COOH_(aq)CH₃COO^–_(aq) + H⁺_(aq)

  • a concentrated solution of a weak acid e.g. a 10 mol/dm³ solution of ethanoic acid

  • a dilute solution of a strong acid e.g. a 0.1 mol/dm³ solution of hydrochloric acid

    • HCl_(g) + aq ==> H⁺_(aq) + Cl^–_(aq)

  • or a concentrated solution of a strong acid e.g. a 10 mol/dm³ solution of hydrochloric acid

10(e) Experimental evidence to support the idea of weak acids versus strong acids

• How can you distinguish a weak acid from a strong acid?

• 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 ...

    • (i) Compare the pH of solutions of equal concentration (equal molarity) and measure the pH with an accurately calibrated pH meter.

    • (ii) 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).

    • See section 1. Examples of everyday acids, alkalis, salts, pH of solution

  • The rate of reaction with metals.  (1 molar means 1.0 mol/dm³, sometimes written as 1M for shorthand))

    • If you put magnesium ribbon into 1 molar solutions of hydrochloric acid (strong, high % ionisation so high H⁺_(aq) concentration) and 1 molar solution of 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.

• 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)CH₃COO^–_(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 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 + hydrogen

    • 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 greater rate 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).

    • From the strong acid solution, you should get a greater volume of hydrogen in the same time.

  • 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.

10(f) pH titration curves for weak acids and weak soluble bases (alkalis) - relating to feasibility of a quantitative titration

In section 7. pH titration curves were introduced by looking what happens when you add a strong acid to a strong alkali.

For titrations involving weak acids or weak soluble bases (weak alkalis) things are a bit more complicated.

You need to add sections together from the graph lines to describe what happens for a particular combination AND order of addition.

 

Graph A adding an alkali to an acid - pH changes

Lines (1) + (4) is adding a strong alkali to a strong acid - this titration gives a distinct end point in a titration

e.g. adding sodium hydroxide solution to hydrochloric acid

Lines (1) + (3) is adding a weak alkali to a strong acid - this titration gives a poorer end point in a titration

e.g. adding ammonia solution to hydrochloric acid

Lines (2) + (4) is adding a strong alkali to a weak acid - this titration gives a poorer end point in a titration

e.g. adding sodium hydroxide solution to ethanoic acid solution

Lines (2) + (3) is adding a weak alkali to a weak base - this titration gives a very poor end point in a titration - not practical

e.g. adding aqueous ammonia to ethanoic acid solution

Graph B adding an acid to an alkali - pH changes

Lines (1) + (4) is adding a strong acid to a strong alkali - this titration gives a distinct end point in a titration

e.g. adding hydrochloric acid to sodium hydroxide solution

Lines (1) + (3) is adding a weak acid to a strong alkali - this titration gives a poorer end point in a titration

e.g. adding ethanoic acid solution to sodium hydroxide solution

Lines (2) + (4) is adding a strong acid to a weak alkali - this titration gives a poorer end point in a titration

e.g. adding hydrochloric acid to aqueous ammonia solution

Lines (2) + (3) is adding a weak acid to a weak alkali - this titration gives a very poor end point in a titration - not practical

e.g. adding ethanoic acid solution to ammonia solution.

10(g) A summary of some points about pH, acids and alkalis

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

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

• In general a 'rough' guide to the 'strength' of an acid or alkali:

  • 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 (hydrogen ion 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.

10(h) Some key revision points about this page

Key definitions and core ideas

• Acid — a substance that produces hydrogen ions H+ in aqueous solution; acids have pH less than 7.
• Base — a substance that can neutralise an acid.
• Alkali — a soluble base that produces hydroxide ions OH– in aqueous solution; alkalis have pH greater than 7.
• Neutralisation — acid + base → salt + water; observed in titrations and salt preparation.
• Indicator and pH — indicators change colour to show whether a solution is acidic, neutral or alkaline; pH is a numerical measure of acidity/basicity from 0–14 in GCSE-level treatment.

Strong and weak acids and alkalis — conceptual explanation

• Strong acid — a substance that ionises almost completely in water so nearly all molecules donate H+; e.g., hydrochloric acid is treated as a strong acid at GCSE level. Solutions of strong acids produce a high concentration of H+ for a given concentration of acid.
• Weak acid — a substance that only partially ionises in water; only a fraction of molecules donate H+, so H+ concentration is much lower at the same molar concentration.
• Strong alkali — a soluble base that dissociates almost completely to give OH– in solution (e.g., concentrated sodium hydroxide).
• Weak alkali — a soluble base that partially ionises to give fewer OH– ions per mole in solution (e.g., ammonia in water produces some OH– by reacting with water).

Important practical consequence: strength is about degree of ionisation, not concentration.

A dilute strong acid can have fewer H⁺ ions than a concentrated weak acid.

Worked ideas students must be able to do

• Describe and give ionic equations for neutralisation (e.g., H+ + OH– → H₂O).
• Use pH scale qualitatively and read indicator colours.
• Distinguish between strong/weak by ionisation behaviour and predict relative conductivity and reactivity trends.
• Perform simple titration calculations at higher tiers (moles, concentrations) and deduce concentration of an acid or alkali from titration results.
• Explain buffer action qualitatively and predict the pH effect of adding small amounts of acid or base for buffer solutions on higher-tier syllabuses.

Student exam tips

• Learn the precise definition: strength = degree of ionisation; concentration = amount of substance per volume. Test yourself with examples (e.g., 0.1 mol dm–3 HCl versus 1.0 mol dm–3 ethanoic acid).
• Use ionic equations for neutralisation; they are shorter and show the key H+ and OH– behaviour.
• For titrations, practise mole calculations until switching between mol, concentration (mol dm–3) and volume is automatic.
• Memorise typical strong acids (HCl, HNO₃, H₂SO₄) and common alkalis (NaOH, KOH, but NH₃ is a weak alkali - weak base). Knowing examples saves time in exams.
• When asked about conductivity, link it to number of ions: strong acids/alkalis conduct better than weak ones at the same concentration.
• For choosing an indicator, focus on the pH at the equivalence point and match an indicator whose colour change range straddles that pH. Titration curve practice helps here.

Common misconceptions and how to avoid them

• Misconception: “Strong means concentrated.” Correction: strength is ionisation; concentration is how much dissolved substance there is. Always state which you mean.
• Misconception: “All bases are insoluble” or “all alkalis are bases.” Correction: bases can be insoluble (e.g., copper(II) oxide), alkalis are only those bases soluble in water and producing OH–.
• Misconception: “Neutralisation always makes a neutral solution.” Correction: the salt produced can give acidic, basic or neutral solutions depending on the parent acid/base strengths; predict by comparing conjugate acid/base strengths (higher-tier).
• Misconception: “Indicators change colour because acid consumes them.” Correction: indicators are weak acids/bases that change colour depending on their protonation state; they report pH, they are not consumed in the acid–base chemistry except in extreme stoichiometric terms.
• Misconception: “pH 7 is always safe or harmless.” Correction: pH is simply a measure of [H+]; many neutral solutions can still be reactive or unsafe chemically for reasons other than pH.

Quick exam checklist

• Write ionic equations for H+ and OH– where appropriate.
• Always state whether you discuss strength or concentration.
• For titrations show clear working: balanced reaction, moles = concentration × volume, show substitute numbers and final significant figures.
• If asked to predict salt solution pH, state which parent acid and base were strong/weak and give the reasoning.
• For higher-tier questions on Ka/pKa or buffers, show the equilibrium expression and explain qualitatively how buffer components remove added H+ or OH–.

Short summary for rapid revision

• Acids produce H+ in water; alkalis produce OH– in water.
• Strength = degree of ionisation; strong acids/alkalis ionise almost completely; weak ones only partially ionise.
• Concentration is distinct from strength; always check what the question asks.
• Neutralisation yields a salt and water; use ionic equations to simplify answers.
• Higher-tier topics include Ka/pKa, buffers and titration curves; practise calculations and indicator choice.

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

• Further work–study links:

  • Acid–Alkali neutralisation titration calculations are on the ChemicalCalculations page, especially sections 11 and 12.

  • The Acids, Bases, pH page section 2. lists common indicators.
  • and the Acids, Bases, pH page section 7. explains the changes in pH in the titration.
  • Advanced level theory of indicators and titrations and Advanced acid–alkali titration questions

  • Advanced A level chemistry notes - equilibrium section 5

    5.1 Lewis and Bronsted–Lowry acid–base theories

    5.2 self–ionisation of water and pH scale

    5.3 strong acids - examples - pH calculations

    5.4 weak acids - examples & pH-K_a-pK_a calculations

    5.5 strong bases -examples -pH calculations

    5.6 weak bases - examples & pH-K_b-pK_b calculations