Acids

3.1 Recall that acids in solution are sources of hydrogen ions and alkalis in solution are sources of hydroxide ions
  • Acids produce H+ ions in aqueous solutions
  • Alkalis produce OH- ions in aqueous solutions
3.2 Recall that a neutral solution has a pH of 7 and that acidic solutions have lower pH values and alkaline solutions higher pH values
  • The pH scale (0 to 14) measures the acidy or alkalinity of a solution, and can be measured using universal indicator of a pH probe
    • pH 7 is neutral
    • < pH 7 is acidic
    • > pH 7 is alkaline

  • Acids are from 0-6PH
  • Alkalis are from 8-14PH
  • Water is 7PH
3.3 Recall the effect of acids and alkalis on indicators, including litmus, methyl orange and phenolphthalein
  • Phenolphthalein
    • Alkaline = pink
    • Acidic = colourless
  • Methyl orange
    • Alkaline = yellow
    • Acidic = red
  • Litmus
    • Litmus solution
      • Alkaline = blue
      • Acidic = red
    • Litmus paper
      • Blue litmus paper goes red in acidic & stays blue in alkaline
      • Red litmus paper goes blue in alkaline & stays red in acidic
3.4 (HT only) Recall that the higher the concentration of hydrogen ions in an acidic solution, the lower the pH; and the higher the concentration of hydroxide ions in an alkaline solution, the higher the pH
  • Higher concentration of H+ ions means that there is a larger amount of H+ ions per unit volume, therefore the acid gives off more H+ ions and is therefore more acidic, thus having a lower pH
  • Exact same for OH- ions and alkalis
3.5 (HT only) Recall that as hydrogen ion concentration in a solution increases by a factor of 10, the pH of the solution decreases by 1
  • As the pH decreases by one unit, the H+ conc. of the solution increases by a factor of 10.
3.6 Core Practical: Investigate the change in pH on adding powdered calcium hydroxide or calcium oxide to a fixed volume of dilute hydrochloric acid
3.7 (HT only) Explain the terms dilute and concentrated, with respect to amount of substances in solution
  • Strong and weak is NOT the same as concentrated and dilute – the latter refers to the amount of substance whereas, the former refers to the above – the H+ ion conc. in aq. solutions
  • Concentrated = more amount of substances in solution
  • Dilute = less amount of substances in solution
3.8 (HT only) Explain the terms weak and strong acids, with respect to the degree of dissociation into ions
  • Strong acid = completely ionised in aqueous solution
    • Hydrochloric, nitric and sulfuric acids
  • Weak acid = partially ionised in aqueous solution
    • Ethanoic, citric and carbonic acids
  • Stronger an acid, lower the pH (for a given conc. of aq. solutions)
3.9 Recall that a base is any substance that reacts with an acid to form a salt and water only

ACID + BASE -> SALT + WATER

3.10 Recall that alkalis are soluble bases
  • Examples of alkalis are soluble metal hydroxides

ACID + ALKALI -> SALT + WATER

3.11 Explain the general reactions of aqueous solutions of acids with: metals, metal oxides, metal hydroxides, and metal carbonates (all) to produce salts

ACID + METAL -> SALT + HYDROGEN

ACID + CARBONATE -> SALT + WATER + CO2

  • Metal oxides are normally bases (because insoluble)
  • Metal hydroxides are bases/alkalis if insoluble/soluble
3.12 Describe the chemical test for: hydrogen and carbon dioxide (using limewater) 
  • Test for hydrogen
    • Use a burning splint held at the open end of a test tube of the gas
      • Creates a ‘squeaky pop’ sound
  • Test for carbon dioxide
    • Calcium hydroxide (aq) (also known as lime water), bubble the gas through the limewater and it will turn milky (cloudy)
3.13 Describe a neutralisation reaction as a reaction between an acid and a base
3.14 Explain an acid-alkali neutralisation as a reaction in which hydrogen ions (H+) from the acid react with hydroxide ions (OH-) from the alkali to form water
  • H+(aq) + OH-(aq) -> H2O(l) is the ionic equation of any neutralization reaction
3.15 Explain why, if soluble salts are prepared from an acid and an insoluble reactant: excess of the reactant is added, the excess reactant is removed, and the solution remaining is only salt and water
  • Excess of the reactant is added
    • Add the base to the warm acid until no more will dissolve and you have some base left over, this is to ensure your volume of acid reacts completely, you then remove the excess reactant by filtration to leave only salt and water
      • You can use this leftover salt solution to obtain the salt by evaporating water
3.16 Explain why, if soluble salts are prepared from an acid and a soluble reactant: titration must be used, the acid and the soluble reactant are then mixed in the correct proportions, and the solution remaining, after reaction, is only salt and water
  • Titration must be used since both reactants are liquids, thus you must measure the exact amount of volumes that react
  • You use an indicator to ensure that the acid and soluble reactant are then mixed in the correct proportions, by stopping the burette at the first sight of a colour change
  • The exact amount of acid has thus been added to the soluble reactant, meaning that the leftover solution is only salt and water, no acid or alkali, because they have been completely neutralised
3.17 Core practical: Investigate the preparation of pure, dry hydrated copper sulfate crystals starting from copper oxide including the use of a water bath
3.18 Describe how to carry out an acid-alkali titration, using burette, pipette and a suitable indicator, to prepare a pure, dry salt
How to carry out a titration:
  1. Wash burette using dilute hydrochloric acid and then water
  2. Fill burette to 100cm3 with acid with the meniscus’ base on the 100cm3 line
  3. Use 25cm3 pipette to add 25cm3 of alkali into a conical flask, drawing alkali into the pipette using a pipette filler
  4. Add a few drops of a suitable indicator to the conical flask (eg: phenolthalein which is pink when alkaline and colourless when acidic)
  5. Add acid from burette to alkali until end-point is reached (as shown by indicator)
  6. The titre (volume of acid needed to exactly neutralise the acid) is the difference between the first (100cm3) and second readings on the burette)
  7. Repeat the experiment to gain more precise results
  8. To prepare a pure, dry salt – you warm the salt solution to evaporate the water
  9. Crystals form
3.19 Recall the general rules which describe the solubility of common types of substances in water: all common sodium, potassium and ammonium salts are soluble, all nitrates are soluble, common chlorides are soluble except those of silver and lead, common sulfates are soluble except those of lead, barium and calcium, and common carbonates and hydroxides are insoluble except those of sodium, potassium and ammonium
3.20 Predict, using solubility rules, whether or not a precipitate will be formed when named solutions are mixed together, naming the precipitate if any

*use information given in 3.19 to predict, knowing the acid reactions

3.21 Describe the method used to prepare a pure, dry sample of an insoluble salt
  • An insoluble salt is formed as a precipitate, which means it is a solid in the water/solution.
    • To separate it, you filter the solution