Topic 4 - Acids, Bases and Salts

• Indicators and the pH scale

  • Indicators change colour when added to alkalis and acids
  • Litmus is the most well-known indicator
  • It turns red in acid and blue in alkalis
  • Universal indicators are most commonly used in the laboratory
  • When added to a solution, it changes colour that shows the pH of the solution

  Acidic - pH<7

  Neutral - pH 7

  Alkalis - pH>7
• The problem with indicators + solutions

  Problem: Colour changes are difficult to identify

  Solutions:

  1. Use a white tile - can see colour change quicker
  2. pH probes - with a pH probe a neutral substance has a pH of exactly 7.0 so it is precise and easy to notice
  3. Measuring temperature - all neutralisation reactions are exothermic -> we can find the point of neutralisation by finding the volume of acid or base added at the highest temperature

  Note:

  • The temperature rise between a substance and a strong acid is greater than with a weak acid
• Acids and Alkalis are commonly used in industry and at home
• Acids produce hydrogen ions, H+, when they dissolve in water e.g hydrochloric acid
• A base is chemically opposite to an acid and a base that dissolves in water is an alkali
• Alkalis produce hydroxide ions, OH-, when they dissolve in water e.g sodium hydroxide
• Acids and Alkalis are used in vinegar or lemon juice

  However, some are dangerous and have the corrosive hazard warning symbol on their containers
• Naming salts


  First name: Metal

  • Metal
  • Metal Oxide/Hydroxide
  • Metal Carbonate
  • Alkali

  Surname: Salts

  • Hydrochloric acid (HCl) - chloride
  • Sulfuric acid (H2SO4) - sulfate
  • Nitric acid (HNO3) - nitrate
• Reactions and observations between Metals/Alkalis and Acids

  Reactions:

  Metal + Acid -> Salt + Hydrogen

  Metal Oxide/Hydroxide + Acid -> Salt + Water

  Metal Carbonate + Acid -> Salt + Water + Carbon Dioxide

  Alkali + Acid -> Salt + Water

  Observations: (in order)

  1. Bubbling, exothermic reaction, lit splint test -> squeaky pop
  2. Exothermic reaction
  3. Limewater turns milky/cloudy, exothermic reaction
  4. Exothermic reaction, fizzing, bubbling

  • Bases + Acid -> no fizzing, temperature rise
  • The more reactive the metal, the faster the reaction, so more bubbling occurs and a greater temperature rise
• Neutralisation

  • Occurs when acid and alkalis 'cancel eachother' out and salt and water is formed
  • The H+ ions from the acid react with the OH- ions from the alkali to form water

  Ionic equation:

  H^+ (aq) + OH^- (aq) -> H2O (l)
• Dissociation 


  • Strong acids fully dissociate in water whereas weak acids only partially dissociate

  Example: HCl -> H^+ + Cl^-

  • Stronger acids have less pH and more H+ ions
  • Strong acids react more quickly than weak acids
• Preparation of a salt from a metal or insoluble base/carbonate

  1. Neutralise the acid and make the salt

  2. Filtration

  3. Evaporation to form salt
• Preparation of a salt from a metal or insoluble base/carbonate

  Step 1 - Neutralise the acid and make the salt

  1. Excess metal/base/carbonate is added to the acid to make sure all the acid has reacted and been used up

  2. Heating and stirring helps the process

  3. For metals and metal carbonates, the fizzing stops when all acid has been used up
• Preparation of a salt from a metal or insoluble base/carbonate

  Step 2 - Filtration

  1. The mixture is filtered using a filter funnel and filter paper

  2. Excess solid remains in the filter paper

  3. Salt solution passes through into the evaporating basin
• Preparation of a salt from a metal or insoluble base/carbonate

  Step 3 - Evaporation to form salt

  1. Salt crystals are collected from the solution by evaporation

  2. Solution is heated to evaporate water

  3. Size of crystals depends on rate of evaporation
• Preparation of a salt from an alkali or soluble carbonate

  Step 1 - Neutralise acid by titration

  Example Method:

  1. Measure exactly 25cm^3 of alkali into a clean conical flask
  2. Add a few drops of indicator to the flask
  3. Place the flask onto a white tile
  4. Fill the burette with acid
  5. Slowly add the acid from the burette to the alkali until the indicator changes colour
  6. Record the volume of acid added to the flask
• Preparation of a salt from an alkali or soluble carbonate

  Step 2 - Make the salt

  • Repeat steps 1-6 of titration without indicator and adding the same volume of acid from the burette and the same volume of alkali to the flask
• Preparation of a salt from an alkali or soluble carbonate

  Step 3 - Evaporation to form salt

  • Heat water until all of the water has been evaporated
• Purpose of titration


  • To find the point of neutralisation
• Purpose of a rough titration


  • To get an approximate end point to speed up the titrations
• Moles, Concentration, Volume equation

  Moles = Concentration (mol dm^-3) x Volume (dm^3)

  Must divide cm^3 value by 1000 to get dm^3 for volume
• Moles, Mr, Mass equation

  Mass = Mr x Moles

  n = Moles

  M = Mass

  Mr = Molar Mass (formula mass)