Topic: Salts preparation

1:31 understand how the formulae of simple compounds can be obtained experimentally, including metal oxides, water and salts containing water of crystallisation

Finding the formula of a metal oxide experimentally

The formulae of metal oxides can be found experimentally by reacting a metal with oxygen and recording the mass changes.

Example: When magnesium is burned in air, it reacts with oxygen (O2) to form magnesium oxide (MgO).

Method:
• Weigh a crucible and lid
• Place the magnesium ribbon in the crucible, replace the lid, and reweigh
• Calculate the mass of magnesium
   (mass of crucible + lid + Magnesium – mass of crucible + lid)
• Heat the crucible with lid on until the magnesium burns
   (lid prevents magnesium oxide escaping therefore ensuring accurate results)
• Lift the lid from time to time (this allows air to enter)
• Stop heating when there is no sign of further reaction
   (this ensures all Mg has reacted)
• Allow to cool and reweigh
• Repeat the heating , cooling and reweigh until two consecutive masses are the same
   (this ensures all Mg has reacted and therefore the results will be accurate)
• Calculate the mass of magnesium oxide formed (mass of crucible + lid + Magnesium oxide – mass of crucible + lid)

 

Finding the formula of a salt containing water of crystallisation

When some substances crystallise from solution, water becomes chemically bound up with the salt.  This is called water of crystallisation and the salt is said to be hydrated. For example, hydrated copper sulfate has the formula  CuSO4.5H2O  which formula indicates that for every CuSO4 in a crystal there are five water (H2O) molecules.

When you heat a salt that contains water of crystallisation, the water is driven off leaving the anhydrous (without water) salt behind. If the hydrated copper sulfate (CuSO4.5H2O) are strongly heated in a crucible then they will break down and the water lost, leaving behind anhydrous copper sulfate (CuSO4). The method followed is similar to that for metal oxides, as shown above. The difference of mass before and after heating is the mass of the water lost. These mass numbers can be used to obtain the formula of the salt.

2:34 know the general rules for predicting the solubility of ionic compounds in water: common sodium, potassium and ammonium compounds are soluble, all nitrates are soluble, common chlorides are soluble, except those of silver and lead(II), common sulfates are soluble, except for those of barium, calcium and lead(II), common carbonates are insoluble, except for those of sodium, potassium and ammonium, common hydroxides are insoluble except for those of sodium, potassium and calcium (calcium hydroxide is slightly soluble)

SaltSolubilityExceptions
sodium (Na+), potassium (K+) and ammonium (NH4+)solublenone
nitrates (NO3-)solublenone
chlorides (Cl-)solublesilver chloride (AgCl) and lead (II) chloride (PbCl2)
sulfates (SO42-)solublebarium sulfate (BaSO4), calcium sulfate (CaSO4) and lead (II) sulfate (PbSO4)
carbonates (CO32-)insolublesodium carbonate (Na2CO3), potassium carbonate (K2CO3) and ammonium carbonate ((NH3)2CO3)
hydroxides (OH-)insolublesodium hydroxide (NaOH), potassium hydroxide (KOH) and calcium hydroxide (Ca(OH)2) (calcium hydroxide is slightly soluble)

2:39 describe an experiment to prepare a pure, dry sample of a soluble salt, starting from an insoluble reactant

Excess Solid Method:

Preparing pure dry crystals of copper sulfate (CuSO4) from copper oxide (CuO) and sulfuric acid (H2SO4)

StepExplanation
Heat acid (H2SO4) in a beakerSpeeds up the rate of reaction
Add base (CuO) until in excess (no more copper oxide dissolves) and stir with glass rodNeutralises all the acid
Filter the mixture using filter paper and funnelRemoves any excess copper oxide
Heat the filtered solution (CuSO4)Hot saturated solution forms
Allow the solution to cool so that hydrated crystals formCopper(II) sulfate is less soluble in cold water
Remove the crystals by filtration and wash with distilled waterRemoves any impurities
Dry by leaving in a warm placeEvaporates the water

 

2:40 (Triple only) describe an experiment to prepare a pure, dry sample of a soluble salt, starting from an acid and alkali

Titration Method:

Preparing pure dry crystals of sodium chloride (NaCl) from hydrochloric acid (HCl) and sodium hydroxide (NaOH)

Before the salt preparation is carried out using the below method, the volume of acid that exactly reacts with 25cm3 of the alkali is found by titration using methyl orange indicator.

StepExplanation
Pipette 25cm3 of alkali (NaOH) into a conical flaskAccurately measures the alkali (NaOH)
Do not add indicatorPrevents contamination of the pure crystals with indicator
Using the titration values, titrate the known volume acid (HCl) into conical flask containing alkaliExactly neutralises all of the alkali (NaOH)
Transfer to an evaporating basin & heat the solutionForms a hot saturated solution (NaCl(aq))
Allow the solution to cool so that hydrated crystals formSodium chloride is less soluble in cold water
Remove the crystals by filtration and wash with distilled waterRemoves any impurities
Dry by leaving in a warm placeEvaporates the water

(Note – This process could be reversed with the acid in the pipette and the alkali in the burette)

How to select the right method for preparing a salt:

2:41 (Triple only) describe an experiment to prepare a pure, dry sample of an insoluble salt, starting from two soluble reactants

Precipitation Method:

Preparing pure dry crystals of silver chloride (AgCl) from silver nitrate solution (AgNO3) and potassium chloride solution (KCl)

StepExplanation
Mix the two salt solutions together in a beakerForms a precipitate of an insoluble salt (AgCl)
Stir with glass rodMake sure all reactants have reacted
Filter using filter paper and funnelCollect the precipitate (AgCl)
Wash with distilled waterRemoves any the other soluble salts (KNO3)
Dry by leaving in a warm placeEvaporates the water

2:42 practical: prepare a sample of pure, dry hydrated copper(II) sulfate crystals starting from copper(II) oxide

Excess Solid Method:

Preparing pure dry crystals of copper sulfate (CuSO4) from copper oxide (CuO) and sulfuric acid (H2SO4)

StepExplanation
Heat acid (H2SO4) in a beakerSpeeds up the rate of reaction
Add base (CuO) until in excess (no more copper oxide dissolves) and stir with glass rodNeutralises all the acid
Filter the mixture using filter paper and funnelRemoves any excess copper oxide
Heat the filtered solution (CuSO4)Hot saturated solution forms
Allow the solution to cool so that hydrated crystals formCopper(II) sulfate is less soluble in cold water
Remove the crystals by filtration and wash with distilled waterRemoves any impurities
Dry by leaving in a warm placeEvaporates the water

 

2:43 (Triple only) practical: prepare a sample of pure, dry lead(II) sulfate

Objective: prepare a pure, dry sample of lead (II) sulfate (PbSO₄).

Preparing a pure, dry sample of lead (II) sulfate (PbSO₄) from lead (II) nitrate solution (Pb(NO₃)₂) and sodium sulfate solution (Na₂SO₄).

      Pb(NO₃)₂ (aq)      +      Na₂SO₄ (aq)      →        PbSO₄ (s)      +      2NaNO₃ (aq)

  1. Mix similar volumes lead nitrate solution and sodium sulfate solution in a beaker. The precise volumes do not matter since any excess will be removed later.
  2. A white precipitate of lead (II) sulfate will form.
  3. The reaction mixture is filtered.
  4. The residue left on the filter paper is washed with distilled water several times to remove impurities.
  5. The residue is then moved to a warm oven to dry.

 

2:44 describe tests for these gases: hydrogen, oxygen, carbon dioxide, ammonia, chlorine

Tests for gases

GasTestResult if gas present
hydrogen (H2)Use a lit splintGas pops
oxygen (O2)Use a glowing splintGlowing splint relights
carbon dioxide (CO2)Bubble the gas through limewaterLimewater turns cloudy
ammonia (NH3)Use red litmus paperTurns damp red litmus paper blue
chlorine (Cl2)Use moist litmus paperTurns moist litmus paper white (bleaches)

2:45 describe how to carry out a flame test

A flame test is used to show the presence of certain metal ions (cations) in a compound.

  • A platinum or nichrome wire is dipped into concentrated hydrochloric acid to remove any impurities.
  • The wire is dipped into the salt being tested so some salt sticks to the end.
  • The wire and salt are held in a non-luminous (roaring) bunsen burner flame.
  • The colour is observed.

 

2:46 know the colours formed in flame tests for these cations: Li⁺ is red, Na⁺ is yellow, K⁺ is lilac, Ca²⁺ is orange-red, Cu²⁺ is blue-green

When put into a roaring bunsen burner flame on a nichrome wire, compounds containing certain cations will give specific colours as follows.

IonColour in flame test
lithium (Li⁺)red
sodium (Na⁺)yellow
potassium (K⁺)lilac
calcium (Ca²⁺)orange-red
copper (II) (Cu²⁺)blue-green

2:47 describe tests for these cations: NH₄⁺ using sodium hydroxide solution and identifying the gas evolved, Cu²⁺, Fe²⁺ and Fe³⁺ using sodium hydroxide solution

Describe tests for the cation NH4+, using sodium hydroxide solution and identifying the ammonia evolved

 

Describe tests for the cations Cu2+, Fe2+ and Fe3+, using sodium hydroxide solution

First, add sodium hydroxide (NaOH), then observe the colour:

2:48 describe tests for these anions: Cl⁻, Br⁻ and I⁻ using acidified silver nitrate solution, SO₄²⁻ using acidified barium chloride solution, CO₃²⁻ using hydrochloric acid and identifying the gas evolved

Describe tests for anions: Halide ions (Cl, Br and I)

 

Describe tests for anions: Sulfate ions (SO42)

 

Describe tests for anions: Carbonate ions (CO32-)

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Section 1: Principles of chemistry

      a) States of matter

      b) Atoms

      c) Atomic structure

     d) Relative formula masses and molar volumes of gases

     e) Chemical formulae and chemical equations

     f) Ionic compounds

     g) Covalent substances

     h) Metallic crystals

     i) Electrolysis

 Section 2: Chemistry of the elements

     a) The Periodic Table

     b) Group 1 elements: lithium, sodium and potassium

     c) Group 7 elements: chlorine, bromine and iodine

     d) Oxygen and oxides

     e) Hydrogen and water

     f) Reactivity series

     g) Tests for ions and gases

Section 3: Organic chemistry

     a) Introduction

     b) Alkanes

     c) Alkenes

     d) Ethanol

Section 4: Physical chemistry

     a) Acids, alkalis and salts

     b) Energetics

     c) Rates of reaction

     d) Equilibria

Section 5: Chemistry in industry

     a) Extraction and uses of metals

     b) Crude oil

     c) Synthetic polymers

     d) The industrial manufacture of chemicals