Edexcel iGCSE Chemistry

1:50 explain how the structures of diamond, graphite and C60 fullerene influence their physical properties, including electrical conductivity and hardness

Allotropes are different forms of the same element. Three different allotropes of carbon are shown here as examples: diamond, graphite and C60 fullerene.

 

Diamond is made up of only carbon atoms, in a giant 3D lattice, where each of those atoms has a strong covalent bonds to 4 other carbon. Every one of carbon’s 4 outer electrons is involved in one of these strong covalent bonds.

Diamond is extremely hard because it is a giant covalent structure with many strong covalent bonds.

Because it is hard, diamond is used in high speed cutting tools, eg diamond-tipped saws.

 

 

Graphite is also made of only carbon atoms, and is also a giant structure, but it is formed of layers where each carbon atom has a strong covalent bond to 3 other carbons. This means each carbon atom has one electron not involved in a covalent bond, and these electrons form a sea of delocalised electrons between the layers.

Even though it is a non-metal, graphite can conduct electricity because the delocalised electrons are free to move.

Each layer is a giant structure, with weak forces of attraction between the layers. These layers can easily slide over each other.

Graphite is soft and slippery because it has weak forces of attraction between layers. It is used as a lubricant and in pencils because it is soft and slippery.

 

 

 

C60 fullerene which is a simple molecular structure (also known as a buckyball) is also made of only carbon atoms, but it forms molecules of 60 carbon atoms. The molecule has weak intermolecular forces of attraction between them which take little energy to overcome. Hence C60 fullerene has a low melting point, and it is soft.

C60 fullerene cannot conduct electricity. Although in each molecule every carbon is only covalently bonded to 3 others and the other electrons are delocalised, these electrons cannot jump between different molecules.

 

 
2023-01-10T11:30:51+00:00Categories: (g) Covalent bonding, 1 Principles of chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , |

1:52 (Triple only) know how to represent a metallic lattice by a 2-D diagram

When metal atoms join together the outer electrons become ‘delocalised’ which means they are free to move throughout the whole structure.

Metals have a giant regular arrangement of layers of positive ions surrounded by a sea of delocalised electrons.

2019-06-21T14:39:25+00:00Categories: (h) Metallic bonding, 1 Principles of chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , |

1:54 (Triple only) explain typical physical properties of metals, including electrical conductivity and malleability

Metals are good conductors because they have delocalised electrons which are free to move.

 

Metals are malleable (can be hammered into shape) because they have layers of ions that can slide over each other.

2019-06-21T14:39:47+00:00Categories: (h) Metallic bonding, 1 Principles of chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , |

1:55 (Triple only) understand why covalent compounds do not conduct electricity

Electrical conductivity is the movement of charged particles.

In this case, charged particles means either delocalised electrons or ions.

These particles need to be free to move in a substance for that substance to be conductive.

Covalent compounds do not conduct electricity because there are no charged particles that are free to move.

2019-05-13T15:40:12+00:00Categories: (i) Electrolysis, 1 Principles of chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , |

1:56 (Triple only) understand why ionic compounds conduct electricity only when molten or in aqueous solution

Ionic compounds only conduct electricity only when molten or in solution.

When solid the ions are not free to move.

 

When molten or in solution the ions are free to move.
2019-05-13T15:39:56+00:00Categories: (i) Electrolysis, 1 Principles of chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , |

1:57 (Triple only) know that anion and cation are terms used to refer to negative and positive ions respectively

A negative ion is called an anion. Examples are the bromide ion (Br⁻) and the oxide ion (O²⁻).

A positive ion is called a cation. Examples are the sodium ion (Na⁺) and the aluminium ion (Al³⁺).

A trick to remember this is to write the ‘t’ as a ‘+’ in the word cation: ca+ion

2019-05-13T15:39:45+00:00Categories: (i) Electrolysis, 1 Principles of chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , |

1:58 (Triple only) describe experiments to investigate electrolysis, using inert electrodes, of molten compounds (including lead(II) bromide) and aqueous solutions (including sodium chloride, dilute sulfuric acid and copper(II) sulfate) and to predict the products

Electrolysis: The breaking down of a substance caused by passing an electric current through an ionic compound which is molten or in solution. New substances are formed.

 

ELECTROLYSIS OF MOLTEN IONIC COMPOUNDS

Example: The electrolysis of molten lead bromide (PbBr2)

  • Solid lead bromide is heated and becomes molten. Explanation: ions become free to move.

  • Electrodes attached to a power source are placed in the molten lead bromide. Explanation: these electrodes are made of either graphite or platinum because both conduct electricity and are fairly unreactive.
  • From the diagram, the left-hand electrode becomes positively charged, this is called the anode. The right-hand becomes negatively charged, this is called the cathode. Explanation: delocalised electrons flow from the anode to the cathode. 
  • At the anode a brown gas is given off. This is bromine gas (Br2(g)). Explanation: Negatively charged bromide ions are attracted to the anode (positive electrode). At the anode, bromide ions lose electrons (oxidation) and become bromine molecules.
  • At the cathode a shiny substance is formed. This is molten lead (Pb(l) ). Explanation: Positively charged lead ions are attracted to the cathode (negative electrode). At the cathode, lead ions gain electrons (reduction) and become lead atoms.

Overall Reaction

word equation:                    lead bromide      –>   lead + bromine

chemical equation:              PbBr2(l)              –>   Pb(l) +   Br2(g)

Remember

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

PANCAKE : Positive Anode, Negative Cathode

 

 

ELECTROLYSIS OF IONIC SOLUTIONS

Rules for working out elements formed from electrolysis of solutions

Follow these rules to decide which ions in solution will react at the electrodes:

At the cathode

Metal ions and hydrogen ions are positively charged. Whether you get the metal or hydrogen during electrolysis depends on the position of the metal in the reactivity series:

  • The metal will be produced if it is less reactive than hydrogen
  • Hydrogen gas (H2) will be produced if the metal is more reactive than hydrogen

 At the anode

At the anode, the product of electrolysis is always oxygen gas (O2) unless the solution contains a high concentration of Cl–, Br­- or I– ions, in which case a halogen is produced, e.g. chlorine gas (Cl2), bromine gas (Br2), and iodine gas (I2).

 

The electrolysis of sodium chloride solution (NaCl(aq))

  • Solid sodium chloride is dissolved in water. Explanation: The sodium ions and chloride ions become free to move.

  • The solution also contains hydrogen ions (H+) and hydroxide ions (OH–). Explanation: Water is a very weak electrolyte. It ionises very slightly to give hydrogen ions and hydroxide ions:

                              H2O(l) ⇋ H+(aq) + OH–(aq)

  • Chloride ions (Cl–) and hydroxide ions (OH–) are attracted to the anode.
  • Sodium ions (Na+) and hydrogen ions (H+) are attracted to the cathode.
  • At the anode a green gas is given off. This is chlorine gas (Cl2(g)). Explanation: chloride ions lose electrons (oxidation) and form molecules of chlorine. The chloride ions react at the anode instead of the hydroxide ions because the chloride ions are in higher concentration. The amount of chlorine gas produced might be lower than expected because chlorine is slightly soluble in water.

                    Electron half equation:   2Cl–(aq)    –>  Cl2 (g) + 2e–     

  • At the cathode a colourless gas is given off. This is hydrogen gas (H2(g)). Explanation: hydrogen ions gain electrons (reduction) and form molecules of hydrogen. The hydrogen ions react at the cathode because hydrogen is below sodium in the reactivity series.

                    Electron half equation:   2H+(aq) + 2e–  –>  H2 (g)

  • The solution at the end is sodium hydroxide (NaOH(aq)).

 

 

The electrolysis of copper sulfate solution (CuSO4(aq))

  • Copper sulfate solution is composed of copper ions (Cu2+), sulfate ions (SO42-), hydrogen ions (H+) and hydroxide ions (OH–). 
  • At the cathode a brown layer is formed. This is copper. Explanation: copper ions gain electrons (reduction) and form atoms of copper. The copper ions react at the cathode instead of hydrogen ions because copper is below hydrogen in the reactivity series.

                    Electron half-equation:   Cu2+(aq) + 2e–  –>  Cu (s)

  • At the anode, bubbles of gas are given off. This is oxygen gas (O2(g)). Explanation: hydroxide ions lose electrons (oxidation) and form molecules of oxygen and water. The hydroxide ions react at the anode instead of the sulfate ions because the hydroxide ions are less stable.

                    Electron half-equation:   4OH–(aq)     –>  O2 (g) + 2H2O(l) + 4e–      

 

 

The electrolysis of sulfuric acid (H2SO4(aq))

  • Sulfuric acid is composed of sulfate ions (SO42-), hydrogen ions (H+) and hydroxide ions (OH–).
  • At the cathode bubbles of gas are formed. This is hydrogen gas (H2(g)). Explanation: hydrogen ions gain electrons (reduction) and form molecules of hydrogen.

                    Electron half-equation:   2H+(aq) + 2e–  –>  H2(g)

  • At the anode, bubbles of gas are given off. This is oxygen gas (O2(g)). Explanation: hydroxide ions lose electrons (oxidation) and form molecules of oxygen and water. The hydroxide ions react at the anode instead of the sulfate ions because the hydroxide ions are less stable.

                    Electron half-equation:   4OH–(aq)     –>  O2 (g) + 2H2O(l) + 4e–      

  • Twice the volume of hydrogen gas is produce compared to oxygen gas. Explanation: from the two half equations, O2 needs 4e– but H2 only needs 2e– as can be seen from the equation

                              2H2O(l)  –>  2H2(g)  +  O2(g)

            There are twice the amount (in moles) of H2 compared to O2

2022-09-16T13:20:34+00:00Categories: (i) Electrolysis, 1 Principles of chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , |

1:59 (Triple only) write ionic half-equations representing the reactions at the electrodes during electrolysis and understand why these reactions are classified as oxidation or reduction

Oxidation: the loss of electrons or the gain of oxygen

Reduction: the gain of electrons or the loss of oxygen

 

Example: The electrolysis of lead (II) bromide, PbBr2

At the cathode (negative electrode):   Pb2+ (l) + 2e–  →  Pb (l)         (reduction)

At the anode (positive electrode):       2Br– (l)      →  Br2 (g) + 2e–       (oxidation)

 

Example: The electrolysis of aluminium oxide, Al2O3

At the cathode:   Al3+ + 3e–    →    Al         (reduction)

At the anode:      2O2-    →    O2 + 4e–       (oxidation)

 

Example: The electrolysis of sodium chloride solution (NaCl (aq))

At the cathode:   2H+ (aq) + 2e–  →  H2 (g)       (reduction)

At the anode:      2Cl– (aq)    →  Cl2 (g) + 2e–       (oxidation)

 

Example: The electrolysis of copper sulfate solution (CuSO4 (aq))

At the cathode:   Cu2+ (aq) + 2e–  →  Cu (s)      (reduction)

At the anode:      4OH– (aq)     →  O2 (g) + 2H2O (l) + 4e–       (oxidation)

2019-05-13T15:39:26+00:00Categories: (i) Electrolysis, 1 Principles of chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , |

1:60 (Triple only) practical: investigate the electrolysis of aqueous solutions

The diagram shows an electrolytic cell.

The electrolyte is an aqueous solution. For example it might be concentrated sodium chloride, NaCl (aq).

The test tubes over the electrodes must not completely cover them to make sure the ions are free to move throughout the solution.

In the case of NaCl (aq) bubbles of gas will be seen forming at the electrodes. These float up and collect in the test tubes when each gas can be tested to assess its identity.

 

 

2019-05-13T15:39:17+00:00Categories: (i) Electrolysis, 1 Principles of chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , |

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 the reaction of water with either potassium or sodium or lithium have the following similarities:

  1. fizzing (hydrogen is produced)
  2. metal floats and moves around on the water
  3. metal disappears

In each case a metal hydroxide solution is produced.

These similarities in the reactions provide evidence that the 3 metals are in the same group of the Periodic Table (i.e. have the same number of electrons in their outer shell).

2023-01-10T12:29:10+00:00Categories: (a) Group 1 (alkali metals) – lithium, sodium and potassium, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , , |

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

Sodium is the second alkali metal in the group. The reaction of sodium and water is more vigorous than lithium’s:

  1. fizzing (hydrogen gas is released)
  2. sodium floats and moves around on the water
  3. sodium melts into a silver-coloured ball
  4. sodium disappears

Potassium is the third alkali metal in the group. The reaction of potassium and water is more vigorous than sodium’s:

  1. fizzing (hydrogen gas is released)
  2. potassium floats and moves around on the water
  3. catches fire with a LILAC flame
  4. potassium disappears

When the group 1 metals react with air they oxidise, showing 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

2023-01-10T12:29:20+00:00Categories: (a) Group 1 (alkali metals) – lithium, sodium and potassium, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , , |

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)????
2023-01-10T12:29:32+00:00Categories: (a) Group 1 (alkali metals) – lithium, sodium and potassium, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , , |

2:04 (Triple only) explain the trend in reactivity in Group 1 in terms of electronic configurations

As you go down the group the outer electron lost from the group 1 metal is further from the nucleus therefore the electron is less attracted by the nucleus and therefore more easily lost.
2019-06-21T14:03:07+00:00Categories: (a) Group 1 (alkali metals) – lithium, sodium and potassium, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , |

2:05 know the colours, physical states (at room temperature) and trends in physical properties of chlorine, bromine and iodine

ElementColourState at room temp
Chlorine (Cl2)GreenGas
Bromine (Br2)Red-brownLiquid
Iodine (l2)GreySolid

Chlorine is a toxic gas, so should be handled in a fume cupboard.

2023-01-10T11:50:40+00:00Categories: (b) Group 7 (halogens) – chlorine, bromine and iodine, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , , |

2:06 use knowledge of trends in Group 7 to predict the properties of other halogens

If you look at the trends in the physical properties of the halogens, Cl2, Br2, I2 you can make predictions about the properties of the other halogens.

ElementColourState at room temp
Fluorine (F2)YellowGas
Astatine (At2)BlackSolid
2023-01-10T11:51:01+00:00Categories: (b) Group 7 (halogens) – chlorine, bromine and iodine, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , , |

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.

2023-01-10T11:51:13+00:00Categories: (b) Group 7 (halogens) – chlorine, bromine and iodine, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , , |

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.

2019-06-21T14:04:08+00:00Categories: (b) Group 7 (halogens) – chlorine, bromine and iodine, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , |

2:09 know the approximate percentages by volume of the four most abundant gases in dry air

Air is a mixture of different gases.

The abundance of gases in the air is as follows:

Gas% by volume
Nitrogen, N278.1
Oxygen, O221.0
Argon, Ar0.9
Carbon dioxide, CO20.04
2023-09-18T11:05:32+00:00Categories: (c) Gases in the atmosphere, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , |

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 the composition 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 and water are 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%.

 

2023-09-18T11:06:04+00:00Categories: (c) Gases in the atmosphere, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , , |

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)

2022-11-30T14:20:27+00:00Categories: (c) Gases in the atmosphere, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , |

2:12 describe the formation of carbon dioxide from the thermal decomposition of metal carbonates, including copper(II) carbonate

thermal decomposition is the process of breaking down by heating.

On heating metal carbonates thermal decompose into metal oxides and carbon dioxide.

Observation: green powder (CuCO3) changes to a black powder (CuO)

2022-11-28T13:09:36+00:00Categories: (c) Gases in the atmosphere, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , |

2:13 know that carbon dioxide is a greenhouse gas and that increasing amounts in the atmosphere may contribute to climate change

Carbon dioxide (CO2) is a greenhouse gas.

It absorbs infra-red radiation and therefore warms the atmosphere. This leads to global warming.

This may cause climate change.
2022-11-30T14:20:45+00:00Categories: (c) Gases in the atmosphere, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , |

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%.

 
2023-09-18T11:06:27+00:00Categories: (c) Gases in the atmosphere, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , , |

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

2023-01-10T12:30:26+00:00Categories: (d) Reactivity series, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , , |

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.
2023-01-10T12:30:47+00:00Categories: (d) Reactivity series, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , , |

2:17 know the order of reactivity of these metals: potassium, sodium, lithium, calcium, magnesium, aluminium, zinc, iron, copper, silver, gold

A more reactive metal will displace a less reactive metal.

In addition a more reactive metal will react more vigorously than a less reactive metal.

For example, potassium takes a shorter time to react than sodium:
2023-01-10T12:31:04+00:00Categories: (d) Reactivity series, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , , |

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

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.

 

2022-12-12T11:13:19+00:00Categories: (d) Reactivity series, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , |

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.

2023-01-10T12:31:29+00:00Categories: (d) Reactivity series, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , |

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 atom (Na). Another example is a chlorine atom (Cl) gaining an electron to become a chloride ion (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).

 

2023-04-19T10:58:30+00:00Categories: (d) Reactivity series, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , , |

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.

2023-01-10T12:32:04+00:00Categories: (d) Reactivity series, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , , |

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.

 

2019-05-13T15:38:43+00:00Categories: (e) Extraction and uses of metals, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , |

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.

2019-05-13T15:38:34+00:00Categories: (e) Extraction and uses of metals, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , |

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.

2019-05-13T15:38:25+00:00Categories: (e) Extraction and uses of metals, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , |

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
Aircrafts and cansLow 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.
2019-06-21T14:45:07+00:00Categories: (e) Extraction and uses of metals, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , |

2:26 (Triple only) know that an alloy is a mixture of a metal and one or more elements, usually other metals or carbon

An alloy is a mixture  of a metal with, usually, other metals or carbon.

For example, brass is a alloy of copper and zinc, and steel is an alloy of iron and carbon.

2019-06-21T14:45:47+00:00Categories: (e) Extraction and uses of metals, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , |

2:27 (Triple only) explain why alloys are harder than pure metals

Alloys are harder than the individual pure metals from which they are made.

In an alloy, the different elements have slightly different sized atoms. This breaks up the regular lattice arrangement and makes it more difficult for layers of ions to slide over each other.

 

2019-06-21T14:46:10+00:00Categories: (e) Extraction and uses of metals, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , |

2:28 describe the use of litmus, phenolphthalein and methyl orange to distinguish between acidic and alkaline solutions

Indicators are substances which change colour in response to a change in pH (acid or alkali).

IndicatorColour in acidic solution [H+]Colour in alkaline solution [OH-]
LitmusRedBlue
Methyl orangeRedYellow
PhenolphthaleinColourlessPink

Methyl orange is orange in a neutral solution
2023-01-13T22:06:03+00:00Categories: (f) Acids, alkalis and titrations, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , |

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

2023-01-10T11:38:25+00:00Categories: (f) Acids, alkalis and titrations, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , , , |

2:30 describe the use of Universal Indicator to measure the approximate pH value of an aqueous solution

An indicator is a substance that has more than one colour form depending on the pH.

Universal indicator is a mixture of different dyes which change colour in a gradual way over a range of pH.
2023-01-10T11:39:19+00:00Categories: (f) Acids, alkalis and titrations, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , , , |

2:31 know that acids in aqueous solution are a source of hydrogen ions and alkalis in a aqueous solution are a source of hydroxide ions

An acid is source of hydrogen ions (H+).

An alkali is source of hydroxide ions (OH–).
2023-01-10T11:39:28+00:00Categories: (f) Acids, alkalis and titrations, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , , , |

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

2023-09-18T11:07:11+00:00Categories: (f) Acids, alkalis and titrations, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , , , |

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.

2019-05-13T15:37:44+00:00Categories: (f) Acids, alkalis and titrations, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , |

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 ((NH4)2CO3)
hydroxides (OH-)insolublesodium hydroxide (NaOH), potassium hydroxide (KOH) and calcium hydroxide (Ca(OH)2) (calcium hydroxide is slightly soluble)
2023-01-10T11:40:19+00:00Categories: (g) Acids, bases and salt preparations, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , |

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.

2023-01-10T11:41:24+00:00Categories: (g) Acids, bases and salt preparations, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , , |

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.

2023-01-10T11:41:38+00:00Categories: (g) Acids, bases and salt preparations, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , , |

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)

2023-01-10T11:41:49+00:00Categories: (g) Acids, bases and salt preparations, 2 Inorganic chemistry, Edexcel iGCSE Chemistry|Tags: , , , , , , , , , |
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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

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