2. A general introduction to connecting pH scale, measuring pH, what is an acid?, what is a base?, what is a salt? What is neutralisation? - salt formation, a simple ionic theory of acids & alkalis

• When water is the solvent dissolving a material, acidic, neutral or alkaline, the resulting mixture is generally referred to as an aqueous solution of whatever was dissolved.

• 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).

2c. 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.

  • 

  • 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 (H⁺ ion) or alkalinity (OH^- ion) of the concentration 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.).

See often used in titrations – e.g. salt preparation method (a)

* Used in titrations - see section 10, litmus is not a good indicator for use in titrations.

Despite the wide use and convenience of indicators the most accurate way of measuring pH is electronically using a special pH meter - illustrated below.

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 very accurately known pH

• 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.

  • If the solution of the substance has a pH of 7 it is described as neutral.

• 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.

  • An acid is a substance that reacts with metals and bases to form a salt.

  • In more advanced theory an acid is defined as a proton donor ...

    •  ... because an acid produces aqueous hydrogen ions in water, H⁺(aq) giving a pH of <7.

    • See Acid reactions with metals/oxides/hydroxides/carbonates, neutralisation reactions for examples.

  • 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, sulfuric 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.

  • A base is a chemical that reacts with acids to form a salt e.g. a metal oxide, hydroxide or carbonate.

    • Bases can be insoluble in water (like copper oxide) or soluble bases like sodium hydroxide - and these soluble bases are called alkalis.

  • In more advanced theory a base is defined as a proton acceptor:

    • A base accepts a hydrogen ion H⁺ from an acid in aqueous solution giving a pH of >7.

    • See Reactions of bases-alkalis like sodium hydroxide for examples.

• 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.

  • This will be extended later to include the neutralisation of carbonates

  • ACID + CARBONATE ==> SALT + WATER + CARBON DIOXIDE

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

  • For even more detail than here see More on Acid–Base Theory and Weak and Strong Acids

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

    • Part 10 "Even 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 or negative 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/sulfuric acid H₂SO₄ gives 2H⁺_(aq) and SO₄^2–_(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)

    • The greater the concentration of hydrogen ions, the more acidic the solution, the lower the pH.

  • Alkalis are soluble base 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 greater the concentration of hydroxide ions, the more alkaline the solution, the higher the pH.

  • 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 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 ...

    • In general: ACID  + BASE  ==>  SALT  +  WATER

    • For a soluble base you can write the word equation as:

    • 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

      • 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

      • H₂SO_4(aq) + CuO_(s) ==> CuSO_4(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 combination of  Na⁺ and Cl^– ions in the ratio 1:1 (from hydrochloric acid and sodium hydroxide)

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

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

    • aluminium sulfate (Al₂(SO₄)₃) consists of a combination of Al³⁺ and SO₄^2– 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).

      • Note how the electric charges of the positive and negative ions must balance each other out (total + = total - ion charges).

2f. An important NOTE on the pH Scale and what happens in simple neutralisation experiments

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

The lower the pH, the higher the hydrogen ion concentration, the more acid 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 4 to pH 3 makes the solution 10x more acidic.

Changing a solution's pH from 4 to 6 makes it 100x less acidic (10 x 10).

• Simple neutralisation experiments

  • (1) Starting with 25 cm³ of a dilute alkali solution (e.g. sodium hydroxide), add some universal indicator solution (blue-violet in excess alkali).

    • Add slowly in small portions, until in excess (well over >25 cm³), a dilute acid solution of similar concentration (e.g. hydrochloric acid).

    • The colour will change from blue to green (neutral) to red (excess acid), with other colours in between.

    • You can follow the rise in pH with a colour chart, or you use universal indicator paper or a pH meter instead of the universal indicator solution.

    • You should find that the pH falls from ~pH 13 to pH 7 (solution neutralised) and then falling further to ~pH 1 with excess acid.

    • If you plot a graph it should roughly look like the red line right diagram by adding acid to the alkali as the pH falls.

    • You can do this experiment very accurately using a pipette, burette and pH meter.

    • You can also reverse the experiment by adding the alkali to the acid and produce the blue graph line and see the pH rise.

  • (2) Starting with 25-50 cm³ of a dilute hydrochloric acid solution, add powdered calcium oxide (lime) or calcium hydroxide (slaked lime) in small portions at a time.

    • Again, follow the pH as it rises from pH 0-1 to pH 7 (neutralised) and then rising further to ~pH 13 with excess of the calcium oxide/hydroxide alkali.

    • If you plot a graph it should roughly look like the blue line (above right diagram).

• See section 7. for more advanced details on Changes in pH in a neutralisation reaction, choice and use of indicators

2g. Here's a comprehensive set of summary revision points tailored for students preparing for GCSE and IGCSE Chemistry across major UK exam boards (WJEC, CCEA, CIE, AQA, Edexcel, OCR Gateway, OCR 21st Century). These notes cover the pH scale, acid-alkali reactions, and include exam tips and common misconceptions.

The pH Scale: Summary Notes

What is pH?

• The pH scale measures how acidic or alkaline a solution is.
• It ranges from 0 to 14:
  • pH 0–6: Acidic
  • pH 7: Neutral (e.g. pure water)
  • pH 8–14: Alkaline (basic)

Ion Concentration

• Acids release hydrogen ions (H⁺) in water.
• Alkalis release hydroxide ions (OH⁻) in water.
• The higher the concentration of H⁺, the lower the pH.
• The higher the concentration of OH⁻, the higher the pH.

Indicators

Mixing Acids and Alkalis: Neutralisation

What Happens?

• A neutralisation reaction occurs:
  • Acid (H⁺) + Alkali (OH⁻) → Water (H₂O)
  • A salt is also formed depending on the acid and alkali used.

General Equation

Acid + Base → Salt + Water

Example:

HCl + NaOH → NaCl + H2O

Common Salts Formed

Exam Tips

• Always name the salt correctly based on the acid and alkali used.
• Use ionic equations for neutralisation:
  H⁺ + OH⁻ → H2O
• Remember indicators and their colour changes.
• Know strong versus weak acids (dealt with in Part 10)
  • Strong acids fully ionise (e.g. HCl).
  • Weak acids partially ionise (e.g. ethanoic acid).

Common Misconceptions about pH, acids, bases and alkalis

• Not all acid reactions are neutralisations – e.g. acid + metal does not produce water.
• pH 0 is not “more acidic” than pH 1 by one unit – it's 10× more acidic (logarithmic scale).
• Alkalis are not just “bases” – they are soluble bases.
• Neutralisation doesn’t always mean pH 7 – depends on the relative strength and concentration of acid and alkali.