Topic 8 Acids, Bases, and Salt preparations

2.28 Litmus paper in acidic and alkali condition

• Acid → pink/red

• Alkali → blue

Phenophthalein indicator in acidic and alkali condition

• Acid→ colourless (markscheme ignores colourless)

• Alkali→ pink

Methyl Orange

Acid → Red

neutral → orange

Alkali → yellow

2.39 pH scale

0-3 strongly acidic

4-6 weakly acidic

7 neutral

8-10 weakly acidic

11-14 strongly alkali

2.30 Universal indicator to measure the pH of an aqueous solution

Universal indicator paper:

• Used for dark solutions where adding liquid would obscure the result.

• Add drops of the solution to universal indicator paper, observe + compare the colour to the pH chart to find the approximate pH value

Universal indicator solution:

• Used for clear solutions

• Add few drops of universal indicator into solution and oberserve+compare the colour to the pH chart to find the approximate pH value

2.31 Acids and Alkalis in aqueous solution

Acids are a source of(release) hydrogen H⁺ ions in an aqueous solution.

• Example of Acids: HCL, H₂SO_4, HNO₃
  → They all contain hydrogen ions.

Alkalis are a source of(release) hydroxide OH^- ions in an aqueous solution.

• Alkalis are bases that dissolve in water such as NaOH, KOH, Ca(OH)₂
  →They all contain hydroxide ions.

2.32 Alkalis and Acids mixed together

• Alkalis can neutralise acids

• (A neutralisation reaction occurs when a base/alkali is added to an acid to form a salt + water.)

• the ionic equation for any alkali-acid neutralisation reaction is: H + (aq) + OH- (aq) -> H2O(l)

2.35 + 2.36 Understand Acids and Bases in term of proton transfer

• Acids are proton donors/they donate protons ⁺ (H⁺ ions are considered protons because they have no electrons in their outershell and only a proton in its nucleus)

• Bases are proton acceptors/they accept protons⁺

2.38 What can act as bases and what are alkalis?

• Metal oxide, metal hydroxide and ammonia (NH₃) can act as bases

• Alkalis are bases that are soluble in water

• (Metal carbonates can also act as bases since they neutralise acids)

Reaction between acid and metal OR base(metal oxide or hydroxide) OR metal carbonate

• Acid + Metal → salt + hydrogen (H₂)

• Acid + Base → salt + water

• (Acid + Ammonia → Ammonium salt + water)

• Acid + Metal Carbonate → salt + water + carbon dioxide

Find the formula of salt formed

• Name of the salt = metal chloride/metal nitrate/metal sulfate

• take the metal ion in the base/metal/carbonate e.g. NaOH → Na+

• take the ion part of the acid which isn’t hydrogen
  -for HCl= Cl
  -for HNO3= NO3
  -for H2SO4= SO4

• balance the charges (crossing method)

2.34 Solubility rules

• Sodium(Na), Potassium(K) and Ammonium(NH4) compounds (salt, carbonate, hydroxide) are soluble

• ALL nitrates are soluble

• MOST chlorides, bromides, iodides are soluble, except lead and silver

• MOST sulfates are soluble, except lead, barium and calcium

• MOST carbonates are insoluble, except sodium, potassium, ammonium

• MOST hydroxides are insoluble, except sodium, potassium, ammonium (and calcium, which slightly soluble)

Precipitate

• A precipitate reaction occurs when a solid forms from 2 solutions reacting together.

• The solid formed is called precipitate.

• You need to determine what the precipitate is from a given reaction based on solubility rules.

• e.g. ammonium carbonate + calcium nitrate → ammonium nitrate + calcium carbonate

• Calcium carbonate is insoluble(like most carbonates), so it is the precipitate. Ammonium nitrate is soluble.

2.39 Preparing a salt from an insoluble reactant(metal carbonate)

1. Transfer acid to a beaker and heat the acid with a Bunsen burner(→ when the acid is heated, the particles have more kinetic anergy and collide more frequently, increasing the rate of reaction)

2. Add excess insoluble base to the acid (to make sure all the acid reacts)

3. When the solution stops bubbling (or metal carbonate appears at the bottom), it means that no more CO2 is produced so all of the acid has reacted.

4. Filter to remove unreacted, excess base (excess metal carbonate stays in the filter paper)

5. Transfer the filtered solution into an evaporating basin. Heat the solution until 1/3 of the volume is reduced (water evaporates)/crystals appear on the side/When glass rod is put into solution, crystals appear on it→ this means the solution is saturated, so overheating may dehydrate the crystal. (crystals contain water of crystallisation, overheating causes the water to evaporate so crystals become powder)

6. Transfer the evaporating basin to a windowsill for a few days to allow the solution to cool and crystallisation to occur

7. Filter the crystals from the remaining liquid and rinse with distilled water

8. Dry between filter paper

2.42 Prepare a sample of pure, dry hydrated copper(II) sulfate crystals from copper(II)oxide

1. Transfer sulfuric acid to a beaker and heat it with a Bunsen Burner. This increases the rate of reaction as molecules have more energy and collide more.

2. Add excess copper oxide to the acid to ensure all the acid has reacts. If copper oxide stops dissolving or appears at the bottom, it means that all the acid has reacted.

3. Filter the mixture to remove excess, unreacted copper oxide, which will stay in the filter paper.

4. Transfer the filtrate into an evaporating basin and heat it until 1/3 of the volume is reduced or when the solution is saturated (put glass rod into solution, if crystals appear on it, the solution is saturated).

5. Transfer the evaporating basin to a windowsill for a few days to allow the solution to cool and crystallisation to occur.

6. Filter the crystals from the remaining liquid, rinse it in distilled water.

7. Dry it between filter paper.