2 Inorganic chemistry

2:01 understand how the similarities in the reactions of lithium, sodium and potassium with water provide evidence for their recognition as a family of elements

Group 1 metals, such as potassium, sodium and lithium, react with water to produce a metal hydroxide and hydrogen. For example:

          lithium   +   water   →   lithium hydroxide   +   hydrogen

          2Li (s)   +   2H₂O (l)   →   2LiOH (aq)   +   H₂ (g)

The observations for this reaction of lithium and water are:

  1. Fizzing
  2. Lithium floats and moves around on the water
  3. Lithium disappears

The reactions of sodium or potassium and water are somewhat similar, giving those observations and producing a metal hydroxide and water.

This similarity in reactions provides evidence that the 3 metals are in the same family of elements (group of the Periodic Table).

2:02 understand how the differences between the reactions of lithium, sodium and potassium with air and water provide evidence for the trend in reactivity in Group 1

Lithium is the first element in group 1 of the Periodic Table. The observations for the reaction of lithium and water are:

  1. fizzing (hydrogen gas is released)
  2. Lithium floats and moves around on the water
  3. Lithium disappears

The reactions of sodium or potassium and water are similar but more vigorous:

  • When sodium reacts with water, the sodium also melts.
  • When potassium reacts with water, the potassium also catches fire with a lilac flame.

When the group 1 metals react with air they show a similar trend in reactivity as we go down the group of the Periodic Table.

Therefore, as we go down group 1 (increasing atomic number), the elements become more reactive: Li<Na<K<Rb<Cs<Fr

2:03 use knowledge of trends in Group 1 to predict the properties of other alkali metals

From the data in the table, it is possible to deduce the properties of francium from the trends in the other group 1 metals.

For example, we can predict that francium will have a melting point around 20⁰C and a density of just over 2g/cm³.

We can also predict that francium will react violently with water, producing francium hydroxide and hydrogen.

Alkali metalMelting point (⁰C)Density (g/cm³)Reaction with waterProducts
lithium (Li)1810.53fizzinglithium hydroxide + hydrogen
sodium (Na)980.97rapid fizzingsodium hydroxide + hydrogen
potassium (K)630.86vigorous fizzing and lilac flamepotassium hydroxide + hydrogen
rubidium (Rb)391.53?rubidium hydroxide + hydrogen
caesium(Cs)291.88?caesium hydroxide + hydrogen
francium (Fr)????

2:07 understand how displacement reactions involving halogens and halides provide evidence for the trend in reactivity in Group 7

Group 7 elements are called the Halogens. As you go up group 7 (decreasing atomic number), the elements become more reactive. For example, fluorine is the most reactive and astatine is the least reactive.

 

A more reactive halogen will displace a less reactive halogen, e.g. chlorine will displace bromine:

By reacting a halogen solution with a potassium halide solution and making observations, the order of their reactivity can be deduced:

 Potassium chloride, KCl(aq)Potassium bromide, KBr(aq)Potassium iodide, KI(aq)
Chlorine, Cl2(aq)No changeColourless to orangeColourless to brown
Bromine, Br2(aq)No changeNo changeColourless to brown
Iodine, I2(aq)No changeNo changeNo change

From the above results, chlorine displaces both bromine and iodine, and bromine displaces iodine. Therefore the order of reactivity is: chlorine is more reactive than bromine, which in turn is more reactive than iodine.

2:08 (Triple only) explain the trend in reactivity in Group 7 in terms of electronic configurations

The higher up we go in group 7 (halogens) of the periodic table, the more reactive the element. The explanation concerns how readily these elements form ions, by attracting a passing electron to fill the outer shell.

In fluorine the outer electron shell is very close to the positively charged nucleus, so the attraction between this nucleus and the negatively charged electrons is very strong. This means fluorine is very reactive indeed.

However, for iodine the outer electron shell is much further from the nucleus so the attraction is weaker. This means iodine is less reactive.

2:10 understand how to determine the percentage by volume of oxygen in air using experiments involving the reactions of metals (e.g. iron) and non-metals (e.g. phosphorus) with air

The following 3 experiments can be used to determine that oxygen (O2) makes up approximately 20% by volume of air.

Copper

The copper is in excess and uses up the oxygen to form copper oxide (CuO).

All the oxygen in the air is therefore used up, and so the volume of the air decreases by about 20% (the percentage of oxygen in air).

 

Iron

The iron reacts with the oxygen in the air (rusting).

As long as the iron, oxygen and water are all in excess, the total volume of air enclosed by the apparatus decreases by about a fifth (20%) over several days.

 

Phosphorus

The phosphorus is lit with a hot wire.

It reacts with the oxygen in the air and causes the water level in the bell jar to rise by about 20%.

 

2:11 describe the combustion of elements in oxygen, including magnesium, hydrogen and sulfur

Magnesium reacts with oxygen producing a bright white flame leaving behind a white ash of magnesium oxide.

          magnesium   +   oxygen   →   magnesium oxide

          2Mg (s)   +   O₂ (g)   →   2MgO

MgO is a base, which can react with an acid to give a salt and water.

 

Hydrogen reacts with oxygen in an explosive reaction. This is the basis of the ‘squeak pop’ test for hydrogen in test tube. With larger quantities of hydrogen this explosion can be dangerous.

          hydrogen   +   oxygen   →   water

          2H₂ (g)   +   O₂ (g)   →   2H₂O (l)

 

Sulfur reacts with oxygen producing a blue flame.

          sulfur   +   oxygen   →   sulfur dioxide

          S (s)   +   O₂ (g)   →   SO₂ (g)

When sulfur dioxide (SO₂) dissolves in water it forms an acidic solution of sulfurous acid:

          SO₂ (g)   +   H₂O (l)   →   H₂SO₃ (aq)

2:14 Practical: determine the approximate percentage by volume of oxygen in air using a metal or a non-metal

The following 3 experiments can be used to determine that oxygen (O2) makes up approximately 20% by volume of air.

Copper

The copper is in excess and uses up the oxygen to form copper oxide (CuO).

All the oxygen in the air is therefore used up, and so the volume of the air decreases by about 20% (the percentage of oxygen in air).

 

Iron

The iron reacts with the oxygen in the air (rusting).

As long as the iron, oxygen and water are all in excess, the total volume of air enclosed by the apparatus decreases by about a fifth (20%) over several days.

 

Phosphorus

The phosphorus is lit with a hot wire.

It reacts with the oxygen in the air and causes the water level in the bell jar to rise by about 20%.

 

2:15 understand how metals can be arranged in a reactivity series based on their reactions with: water and dilute hydrochloric or sulfuric acid

Some metals are more reactive than others.

The order of reactivity can be determined by adding acid to different metals and observing the rate of reaction.

For example, when hydrochloric acid is added to iron (Fe) then bubbles of hydrogen are produced slowly. However, if the same acid is added to zinc (Zn) then bubbles will be produced more quickly. This tells us that zinc is more reactive than iron.

Instead of using acid, water can be used to test the relative reactivity of metals. However, many metals are too low in the reactivity series to react with water

2:16 understand how metals can be arranged in a reactivity series based on their displacement reactions between: metals and metal oxides, metals and aqueous solutions of metal salts

A metal will displace another metal from its oxide that is lower in the reactivity series. For example, a reaction with magnesium and copper (II) oxide will result in the magnesium displacing the copper from its oxide:

A metal will also displace another metal from its salt that is lower in the reactivity series. For example, the reaction between zinc and copper (II) sulfate solution will result in zinc displacing the copper from its salt:

The blue colour of the copper (II) sulfate solution fades as colourless zinc sulfate solution is formed.

2:19 understand how the rusting of iron may be prevented by: barrier methods, galvanising and sacrificial protection

Barrier Methods: Rusting may be prevented by stopping the water and oxygen getting to the iron with a barrier of grease, oil, paint or plastic.

Galvanising: (coating in zinc) also prevents water and oxygen getting to the iron, but with galvanising even if the barrier is broken the more reactive zinc corrodes before the less reactive iron. During the process, the zinc loses electrons to form zinc ions.

Sacrificial Protection: Zinc blocks are attached to iron boat hulls and underground pipelines to act as sacrificial anodes. Zinc is more reactive than iron, so oxygen in the air reacts with the zinc to form a layer of zinc oxide instead of the iron.

2:20 in terms of gain or loss of oxygen and loss or gain of electrons, understand the terms: oxidation, reduction, redox, oxidising agent, reducing agent, in terms of gain or loss of oxygen and loss or gain of electrons

Oxidation

  • Oxidation is the loss of electrons. For example a sodium atom (Na) loses an electron to become a sodium ion (Na⁺). Another example is a chloride ion (Cl⁻) losing an electron to become a chlorine atom (Cl).
  • Another definition of oxidation is the gain of oxygen. For example if carbon combines with oxygen to form carbon dioxide, the carbon is being oxidised.

 

Reduction

  • Reduction is the gain of electrons. For example a sodium ion (Na⁺) gains an electron to become a sodium ion (Na). Another example is a chlorine atom (Cl) gaining an electron to become a chloride atom (Cl⁻).
  • Another definition of reduction is the loss of oxygen. For example when aluminium oxide is broken down to produce aluminium and oxygen, the aluminium is being reduced.

 

Redox: A reaction involving oxidation and reduction.

A good way to remember the definitions of oxidation and reduction in terms of electrons is:

  • OILRIG : Oxidation Is the Loss of electrons and Reduction Is the Gain of electrons

 

Oxidising agent: A substance that gives oxygen or removes electrons (it is itself reduced).

 

Reducing agent: A substance that takes oxygen or gives electrons (it is itself oxidised).

 

2:21 practical: investigate reactions between dilute hydrochloric and sulfuric acids and metals (e.g. magnesium, zinc and iron)

Metals which are above hydrogen in the reactivity series will react with dilute hydrochloric or sulfuric acid to produce a salt and hydrogen.

metal   +   acid   →   salt   +   hydrogen

For example:

         magnesium   +   hydrochloric acid   →   magnesium chloride   +   hydrogen

         Mg (s)         +         2HCl (aq)         →         MgCl₂ (aq)         +         H₂ (g)

This is a displacement reaction.

Image result for magnesium + hydrochloric acid

There is a rapid fizzing and a colourless gas is produced. This gas pops with a lighted splint, showing the gas is hydrogen.

The reaction mixture becomes warm as heat is produced (exothermic).

The magnesium disappears to leave a colourless solution of magnesium chloride.

If more reactive metals are used instead of magnesium the reaction will be faster so the fizzing will be more vigorous and more heat will be produced.

2:22 (Triple only) know that most metals are extracted from ores found in the Earth’s crust and that unreactive metals are often found as the uncombined element

Most metals are found in the Earth’s crust combined with other elements. Such compounds are found in rocks called ore, rocks from which it is worthwhile to extract a metal.

A few very unreactive metals, such as gold, are found native which means they are found in the Earth’s crust as the uncombined element.

 

2:23 (Triple only) explain how the method of extraction of a metal is related to its position in the reactivity series, illustrated by carbon extraction for iron and electrolysis for aluminium

Extraction of a metal from its ore typically involves removing oxygen from metal oxides.

 

If the ore contains a metal which is below carbon in the reactivity series then the metal is extracted by reaction with carbon in a displacement reaction.

 

If the ore contains a metal which is above carbon in the reactivity series then electrolysis (or reaction with a more reactive metal) is used to extract the metal.

2:24 (Triple only) be able to comment on a metal extraction process, given appropriate information

Extraction of a metal from its ore typically involves removing oxygen from metal oxides.

 

If the ore contains a metal which is below carbon in the reactivity series then the metal is extracted by reaction with carbon in a displacement reaction.

 

If the ore contains a metal which is above carbon in the reactivity series then electrolysis (or reaction with a more reactive metal) is used to extract the metal.

2:25 (Triple only) explain the uses of aluminium, copper, iron and steel in terms of their properties the types of steel will be limited to low-carbon (mild), high-carbon and stainless

Aluminium
UseProperty
AircraftLow density / resists corrosion
Power cablesConducts electricity / ductile
Pots and pansLow density / strong (when alloyed) / good conductor of electricity and heat

Aluminium resists corrosion because it has a very thin, but very strong, layer of aluminium oxide on the surface.

Copper 
UseProperty
Electrical wiresvery good conductor of electricity and ductile
Pots and pansvery good conductor of heat / very unreactive / malleable
Water pipesunreactive / malleable
Surfaces in hospitalsantimicrobial properties / malleable
Iron
UseProperty
BuildingsStrong
SaucepansConducts heat / high melting point / malleable
Steel  
Type of steelIron mixed withSome uses
Mild steelup to 0.25% carbonnails, car bodies, ship building, girders
High-carbon steel0.6%-1.2% carboncutting tools, masonry nails
Stainless steelChromium (and nickel)cutlery, cooking utensils, kitchen sinks

Mild steel is a strong material that can easily be hammered into various shapes (malleable). It rusts easily.

High-carbon steel is harder than mild steel but more brittle (not as malleable).

Stainless steel forms a strong, protective oxide layer so is very resistant to corrosion.

2:29 understand how to use the pH scale, from 0–14, can be used to classify solutions as strongly acidic (0–3), weakly acidic (4–6), neutral (7), weakly alkaline (8–10) and strongly alkaline (11–14)

The pH scale ranges from 0 to 14, and tells you how acidic or how alkaline a solution is.

 strongly acidicweakly acidicneutralweakly alkalinestrongly alkaline
pH0-34-678-1011-14

2:32 know that bases can neutralise acids

Metal oxides, metal hydroxides and ammonia (NH₃) are called bases.

Bases neutralise acids by combining with the hydrogen ions in them.

The key reaction is:

   acid             +             base             →            salt             +             water

An example of this is:

   sulfuric acid   +   copper oxide   →   copper sulfate   +   water

   H₂SO₄          +          CuO          →          CuSO₄          +          H₂O

2:33 (Triple only) describe how to carry out an acid-alkali titration

Titration is used to find out precisely how much acid neutralises a certain volume of alkali (or vice versa).

The diagram shows the titration method for a neutralisation reaction between hydrochloric acid and sodium hydroxide, using phenolphthalein as an indicator. The indicator changes colour when neutralisation occurs.

The conical flask is swirled to mix the solutions each time alkali is added. When reading the burette it is important to be aware that the numbers on the scale increase from top to bottom. Readings are usually recorded to the nearest 0.05cm³ so all readings should be written down with 2 decimal places. The second decimal place is given as a ‘0’ if the level of the solution is on a line, or ‘5’ if it is between the lines. The volume of alkali added is calculated by subtracting the final reading from the initial reading. Various indicators can be used such as phenolphthalein or methyl orange. However universal indicator should not be used since it has a wide range of colours rather than one specific colour change so it would be unclear when the precise endpoint of titration was achieved.

This process is repeated a number of times. The first time it is done roughly to get a good approximation of how much alkali needs to be added. On subsequent attempts, the alkali is added very slowly when approaching the correct volume.

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)

2:35 understand acids and bases in terms of proton transfer

An acid is a proton (H⁺) donor.

A base is a proton (H⁺) acceptor.

 

A proton is the same as a hydrogen ion. A good way to think about that is to realise that a hydrogen atom is just one proton and zero neutrons surrounded by only one electron. If that atom becomes an ion by the removal of the electron, then only one proton is left.

 

When sulfuric acid reacts with copper (II) oxide (CuO):

Cu²⁺O²⁻ (s)         +         H₂SO₄ (aq)         →         Cu²⁺ (aq)         +         SO₄²⁻ (aq)         +         H₂O (l)

H₂SO₄ is an acid. It donates protons (H⁺) to CuO, the base.

2:36 understand that an acid is a proton donor and a base is a proton acceptor

An acid is a proton donor.

A base is a proton acceptor.

 

A proton is the same as a hydrogen ion. A good way to think about that is to realise that a hydrogen atom is just one proton and zero neutrons surrounded by only one electron. If that atom becomes an ion by the removal of the electron, then only one proton is left.

2:37 describe the reactions of hydrochloric acid, sulfuric acid and nitric acid with metals, bases and metal carbonates (excluding the reactions between nitric acid and metals) to form salts

Acid reactions summary

         alkali      +      acid      →      water      +      salt

         base      +      acid      →      water      +      salt

         carbonate      +      acid      →      water      +      salt      +      carbon dioxide

         metal   +   acid   →   salt   +   hydrogen

To assist remembering this list, many pupils find it useful to remember this horrid looking but very effective mnemonic:

         AAWS

         BAWS

         CAWS CoD

         MASH

Acids are a source of hydrogen ions (H⁺) when in solution. When the hydrogen in an acid is replaced by a metal, the compound is called a salt. The name of the salt depends on the acid used. For example if sulfuric acid is used then a sulfate salt will be formed.

Parent acidFormulaSaltFormula ion
sulfuric acidH2SO4sulfateSO42-
hydrochloric acidHClchlorideCl-
nitric acidHNO3nitrateNO3-

 

Acid + Alkali   and   Acid + Base

A base is a substance that can neutralise an acid, forming a salt and water only.

Alkalis are soluble bases. When they react with acids, a salt and water is formed. The salt formed is often as a colourless solution. Alkalis are a source of hydroxide ions (OH⁻) when in solution.

         alkali      +      acid      →      water      +      salt

         base      +      acid      →      water      +      salt

Examples of acid + alkali reactions:

  •          sodium hydroxide   +   hydrochloric acid   →   sodium chloride   +   water
  •          NaOH (aq)         +         HCl (aq)         →         NaCl (aq)         +         H₂O (l)
  •          potassium hydroxide   +   sulfuric acid   →   potassium sulfate   +   water
  •          2KOH (aq)         +         H₂SO₄ (aq)         →         K₂SO₄ (aq)         +         2H₂O (l)

Example of an acid + base reaction:

         CuO (s)         +         H₂SO₄ (aq)         →         CuSO₄ (aq)         +         H₂O (l)

 

Acid + Carbonate

         carbonate      +      acid      →      water      +      salt      +      carbon dioxide

A carbonate is a compound made up of metal ions and carbonate ions. Examples of metal carbonates are sodium carbonate, copper carbonate and magnesium carbonate.

When carbonates react with acids, bubbling is observed which is the carbon dioxide being produced. If the acid is in excess the carbonate will disappear.

Examples of acid + carbonate reactions:

  •          calcium carbonate   +   hydrochloric acid   →   calcium chloride   +   water   +   carbon dioxide
  •          CaCO₃ (s)         +         2HCl (aq)         →         CaCl₂ (aq)         +         H₂O (l)         +         CO₂ (g)
  •          potassium carbonate   +   hydrochloric acid   →   potassium chloride   +   water   +   carbon dioxide
  •          K₂CO₃ (aq)         +         2HCl (aq)         →         2KCl (aq)         +         H₂O (l)         +         CO₂ (g)

 

Acid + Metal

         metal   +   acid   →   salt   +   hydrogen

Metals will react with an acid if the metal is above hydrogen in the reactivity series.

When metals react with acids, bubbling is observed which is the hydrogen being produced. If the acid is in excess the metal will disappear.

Examples of acid + metal reactions:

  •          magnesium   +   sulfuric acid   →   magnesium sulfate   +   hydrogen
  •          Mg (s)         +         H₂SO₄ (aq)         →         MgSO₄ (aq)         +         H₂ (g)
  •          aluminium   +   hydrochloric acid   →   aluminium chloride   +   hydrogen
  •          2Al (s)         +         6HCl (aq)         →         2AlCl₃ (aq)         +         3H₂ (g)
  •          copper   +   hydrochloric acid   →   no reaction (since copper is below hydrogen in the reactivity series)

2:38 know that metal oxides, metal hydroxides and ammonia can act as bases, and that alkalis are bases that are soluble in water

A base is a substance that neutralises an acid by combining with the hydrogen ions in them to produce water.

A base usually means a metal oxide, a metal hydroxide or ammonia.

Alkalis are bases which are soluble in water.

 

Some metal oxides are soluble in water and react with it to form solutions of metal hydroxides. For example:

Na₂O (s)         +         H₂O (l)         →         2NaOH (aq)

Apart from this and other group 1 oxides (such as potassium oxide) most other metal oxides are not soluble in water.

One exception is calcium oxide which does dissolve slightly in water to form calcium hydroxide (known as limewater):

CaO (s)         +         H₂O (l)         →         Ca(OH)₂ (aq)

 

Ammonia is another base. Ammonia reacts with water to form ammonium ions and hydroxide ions:

NH₃ (aq)         +         H₂O (l)         ⇋         NH₄⁺ (aq)         +         OH⁻ (aq)

 

All the solutions produced here contain hydroxide ions (OH⁻) so they are all alkalis.

 

 

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-)

Select a set of flashcards to study:

     Terminology

     Skills and equipment

     Remove Flashcards

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