Section 2 g) Specfication

2.37 describe tests for the cations:
i Li+, Na+, K+, Ca2+ using flame tests
ii NH4 +, using sodium hydroxide solution and identifying the ammonia evolved
iii Cu2+, Fe2+ and Fe3+, using sodium hydroxide solution

2.38 describe tests for the anions:
i Cl- , Br- and I- , using dilute nitric acid and silver nitrate solution
ii SO4 2- , using dilute hydrochloric acid and barium chloride solution
iii CO3 2- , using dilute hydrochloric acid and identifying the carbon dioxide evolved

2.39 describe tests for the gases:
i hydrogen
ii oxygen
iii carbon dioxide
iv ammonia
v chlorine.

Section 2 f) Summary

Metals can be placed in an order of most to least reactive, called the reactivity series.

This can be remembered using an acronym:
Please
Stop
Calling
Me
A
Cranky
Zebra
I
Teach
Lions
How
Cars
Save
Gold
Printers

More reactive metals will displace less reactive ones in a compound (e.g. a metal salt or a metal oxide), which can be demonstrated by dissolving the compound in water to make it aqueous, then adding another metal to see if it reacts; if it does, then it is more reactive. If it doesn't, it's less reactive.
Reactivity series can also be determined by evaluating the intensity of reaction with water or acid.

Sacrificial protection
Iron is a useful metal, and it's on the less reactive side. But, when exposed to water and air, it will rust. It can be protected using grease, oil, paint or plastic, but another method is galvanising.
Galvanising is a form of sacrificial protection, wherein the iron is coated in a layer of zinc, which is a more reactive metal. This causes the oxygen and water to react with the zinc instead of the iron, protecting it. This method can be employed in different metals, but is called sacrificial protection and not galvanising in this case.

Section 2 f) Key Words

Galvanising: Coating iron in zinc as a form of sacrificial protection.

Oxidation: Gain of oxygen ions.

Reactivity series: The order in which metals will react with a non-metal, more reactive metals displacing less reactive metals.

Reduction: Loss of oxygen ions.

Rust: Compound formed when iron reacts with water and oxygen.

Sacrificial protection: Coating a less reactive metal in a more reactive metal, so the more reactive metal reacts with the surroundings and the less reactive metal is protected.

Section 2 f) Specification

2.29 understand that metals can be arranged in a reactivity series based on the reactions of the metals and their compounds: potassium, sodium, lithium, calcium, magnesium, aluminium, zinc, iron, copper, silver and gold

Some metals are more reactive than others, based on their chemical compositions. More reactive metals will displace less reactive metals in a compound, and they react more strongly. From reactions, we can determine the reactivity series (from most to least reactive):

Potassium K
Sodium Na
Lithium Li
Calcium Ca
Magnesium Mg
Aluminium Al
Zinc Zn
Iron Fe
Copper Cu
Silver Ag
Gold Au

2.30 describe how reactions with water and dilute acids can be used to deduce the following order of reactivity: potassium, sodium, lithium, calcium, magnesium, zinc, iron and copper

Potassium, sodium, and lithium are alkali metals, therefore their reactivity series is easily determined by reacting them with cold water. Calcium also reacts with water, but less so. Magnesium, zinc, and iron will only react with cold water very slowly, but will react with acids (e.g. hydrochloric acid) with varying degrees of intensities. 
Copper will not react with either an acid or water.
The more vigorous the reaction, the more reactive the metal.

2.31 deduce the position of a metal within the reactivity series using displacement reactions between metals and their oxides, and between metals and their salts in aqueous solutions

To carry out a reaction to determine reactivity series, dissolve a metal oxide or metal salt in water, then add another metal. If there is a reaction, the metal you introduced is more reactive. If there is no reaction, it is less reactive. This can be done multiple times with different metals to determine reactivity series.

2.32 understand oxidation and reduction as the addition and removal of oxygen respectively

OILRIG says Oxidation Is Loss of electrons and Reduction Is Gain of electrons, but it can also be in reference to loss or gain of an oxygen ion, as they are positively charged; they have more protons than electrons. Oxidation is gain of oxygen ions and reduction is loss of oxygen ions. This is easily memorable as 'oxidation' is rooted from the word oxygen.

2.33 understand the terms redox, oxidising agent, reducing agent

Redox means a reaction involving molecules that are reduced and molecules that are oxidised.
An oxidising agent is a less reactive metal in a compound that allows the other metal to be oxidised.
A reducing agent is a more reactive metal that reduces the other metal.

2.34 describe the conditions under which iron rusts

Under conditions where there is moisture and air; it requires both oxygen and water to rust.

2.35 describe how the rusting of iron may be prevented by grease, oil, paint, plastic and galvanising

Grease, oil, paint or plastic creates a layer between the iron and air and water. It repels water and doesn't allow air in.
Galvanising is a technique of coating iron with zinc, a more reactive metal. The air and water react with the zinc instead of the iron, and preventing the iron from coming into contact with the air.

2.36 understand the sacrificial protection of iron in terms of the reactivity series.

Sacrificial protection means a more reactive metal is made to coat a less reactive one, it then reacts with the air and water instead of the iron.

Section 2 e) Summary

When a metal reacts with an acid, it forms a salt and hydrogen. It is easily remembered by MASH
Metal
+
Acid
-->
Salt
+
Hydrogen

When hydrogen reacts with oxygen, it releases a lot of energy in the form of heat and light. This creates a flame, making it a combustion reaction.

2H2 + O2 --> 2H2O + energy

This reaction produces pure water as well as energy. We can test for water in a few different ways:

Testing for water

Physical test: Heat the water and see if it boils at 100℃, in which case it is pure.

Chemical test: Expose to anhydrous copper (II) sulphate, a white substance. It will turn blue in presence of water.

Section 2 e) Key Words

Chemical Test: A test involving the chemical molecules, and the way they react with other molecules. e.g. test for water using anhydrous copper (ii) sulphate.

Combustion: Reaction with oxygen that produces a flame (a lot of heat and light energy is released)

Physical Test: A test involving the physical properties of a molecule, e.g. boiling point of water.

Section 2 e) Specification

2.25 describe the reactions of dilute hydrochloric and dilute sulfuric acids with magnesium, aluminium, zinc and iron

The rule "MASH" tells us that when a metal reacts with an acid, a salt will be produced, as well as hydrogen.

Metal + Acid --> Salt + Hydrogen

e.g.

Mg(s) + 2HCl(aq) --> MgCl2 + H2

2.26 describe the combustion of hydrogen

The combustion of hydrogen is its reaction with oxygen. It is exothermic, so a lot of energy is produced, as well as pure water.

2H2 + O2 --> 2H2O

2.27 describe the use of anhydrous copper(II) sulfate in the chemical test for water

If anhydrous copper sulphate changes from white to blue, it has become hydrous copper sulphate and it contains water.

2.28 describe a physical test to show whether water is pure.

Heat the water until it boils, with a thermometer. If it boiled at 100℃ exactly, it is pure water.

Section 2 d) Summary

The Earth's atmosphere is made up of many different gases, including

• Nitrogen (78%)
• Oxygen (21%)
• Argon (0.9%)
• Carbon dioxide, Water vapour, etc. (>0.1%)

The last category includes greenhouse gases: Gases that trap the sun's energy inside the atmosphere, causing the greenhouse effect. This is usually a good thing; the greenhouse effect keeps us alive. But excess greenhouse gases are being released into the atmosphere, increasing the volume constantly (due to burning fossil fuels, etc.). This causes the enhanced greenhouse effect, where too much of the sun's energy is trapped and causes a heating effect called climate change. This is a dangerous process that is threatening the world we live in, and must be stopped. 

We can work out the percentage volume of oxygen in air with a simple experiment. 

1. Place a tube containing copper in the middle of two gas syringes (containing a known volume of air), attached so the ends are sealed. 
2. Gently heat the copper with a bunsen burner, while slowly pressing the syringes, alternating, and keeping an eye on the volume. 
3. Once the volume stops changing when you press the syringe through, turn off the bunsen burner and wait for it to cool. 
4. Compare the volume you started with, and the volume you ended up with, and calculate the percentage. This is how much oxygen was in there. 

Oxides can be formed by burning elements, for example:

• Burning magnesium forms magnesium oxide, a basic compound
• Burning sulphur forms sulphur dioxide, an acidic compound that can be dissolved to form sulphuric acid
• Burning carbon to form carbon dioxide, an acidic compound which has many uses but also contributes greatly to the enhanced greenhouse effect. 

Carbon dioxide is water-soluble, making it useful for carbonating drinks. The carbon dioxide is dissolved under high pressure, but as this is a reversible reaction when the pressure is released bubbles form.
It is also denser than air, making it useful for smothering fires. CO2 is used in many fire extinguishers for this reason.

Carbon dioxide can be formed by reacting hydrochloric acid and calcium carbonate:
Hydrochloric acid + Calcium carbonate --> Carbon dioxide + Calcium chloride + Water

HCl(aq) + CaCO3(s) --> CO2(g) +CaCl2(s) + H2O(l)

Calcium carbonate could be used as marble or limestone, and dropped into a sealed flask of dilute hydrochloric acid in small pieces. A delivery tube could be placed in the end of the bung to allow for the gas to be collected in the downwards displacement method.

Carbon dioxide can also be formed through the thermal decomposition of a metal carbonate, for example:
Copper (II) Carbonate --> Carbon Dioxide + Copper Oxide
CuCO3 --> CO2 + CuO

Another decomposition reaction is hydrogen peroxide heated with manganese (IV) oxide.
Hydrogen peroxide --> Water + Oxygen

Section 2 d) Specification

2.16 recall the gases present in air and their approximate percentage by volume

Oxygen - 21%
Nitrogen - 78%
Argon - 0.9%
Other (carbon dioxide, water vapour, etc.) - >0.1%

2.17 explain how experiments involving the reactions of elements such as copper, iron and phosphorus with air can be used to investigate the percentage by volume of oxygen in air

If you have a known volume of air, then remove the oxygen by reacting it with an excess of another element (copper, iron, phosphorus) to create a solid, the volume of gas will change. This change can then be used to calculate the percentage of oxygen in the air. This can be done multiple times, and an average found, to increase the accuracy of the results.

2.18 describe the laboratory preparation of oxygen from hydrogen peroxide, using manganese(IV) oxide as a catalyst

Hydrogen peroxide can be decomposed simply by heating with a catalyst - Manganese oxide
Hydrogen peroxide --> Oxygen + Water
The oxygen bubbles created can be collected using the downwards displacement method.

2.19 describe the reactions of magnesium, carbon and sulfur with oxygen in air, and the acid-base character of the oxides produced

Magnesium burns in air, reacting with oxygen, to form magnesium oxide, a basic substance, as it is a metal oxide.

Magnesium + Oxygen --> Magnesium Oxide

Carbon and sulphur are both non-metals, that react with air, giving out heat and light, to form acidic non-metal oxides.

Carbon + Oxygen --> Carbon dioxide

Sulphur + Oxygen --> Sulphur dioxide

2.20 describe the laboratory preparation of carbon dioxide from calcium carbonate and dilute hydrochloric acid

Hydrochloric acid + Calcium carbonate --> Carbon dioxide + Calcium chloride + Water

HCl(aq) + CaCO3(s) --> CO2(g) +CaCl2(s) + H2O(l)

Calcium carbonate could be used as marble or limestone, and dropped into a sealed flask of dilute hydrochloric acid in small pieces. A delivery tube could be placed in the end of the bung to allow for the gas to be collected in the downwards displacement method. 

2.21 describe the formation of carbon dioxide from the thermal decomposition of metal carbonates such as copper(II) carbonate

Copper (II) Carbonate --> Carbon Dioxide + Copper Oxide
CuCO3 --> CO2 + CuO

2.22 describe the properties of carbon dioxide, limited to its solubility and density

Carbon dioxide is a relatively dense gas, it is denser than air. It is water-soluble at high pressure, so when bubbled through water carbonic acid can be formed, and it turns lime water cloudy when dissolved.

2.23 explain the use of carbon dioxide in carbonating drinks and in fire extinguishers, in terms of its solubility and density

When carbon dioxide is dissolved into a liquid at a high pressure, it carbonates it. This reaction is reversible, so this means that when it is returned to atmospheric pressure, bubbles of carbon dioxide form again, and this is how we get fizzy drinks.

Carbon dioxide is used in fire extinguishers because it is denser than air, and is therefore good at smothering fires and preventing oxygen from reaching them. The carbon dioxide sinks over the fire in a blanket and stops oxygen from reaching it, which is necessary for combustion.

2.24 understand that carbon dioxide is a greenhouse gas and may contribute to climate change.

Carbon dioxide is a greenhouse gas, meaning that it traps the sun's energy in the atmosphere, creating the enhanced greenhouse effect with excess. This excess heat that is not able to escape the Earth is a major cause of climate change. 

Section 2 c) Summary

The group 7 elements, also known as halogens, are F, fluorine, Cl, chlorine, Br, bromine, I, iodine, and At, astatine. They all have antimicrobial properties, but in larger doses are toxic.

F - Fluorine - a pale yellow gas
Cl - Chlorine - a green gas
Br - Bromine - an orange liquid
I - Iodine - a purple solid
At - Astatine - a black solid

The halogens are in the same group, so have similar properties that show trends. As you go down the group, the elements get darker in colour, less reactive, and have a higher melting and boiling point.
The reason for this is because as the molecules increase in size, the distance between the valence electrons and the nucleus increases, weakening the forces of attraction and making it more difficult for the atom to attract another electron. Additionally, the increase in molecule size means the attraction between the molecules is more difficult to break, causing the increase in melting and boiling point.

This topic focuses primarily on fluorine, chlorine and bromine. Their reactivity series can be determined by combining a metal halide and aqueous halide, and seeing if a reaction takes place. By adding methylbenzene, we can see which molecules are present (purple is iodine, yellow is bromine)

This shows that chlorine is the most reactive of the three, and iodine the least. These are displacement reactions, where the less reactive halogen is replaced by the more reactive halogen. The more reactive one is reduced, it gains electrons, and the less reactive one is oxidised, it loses electrons.

Reactions 

 The hydrogen halides formed in reaction with hydrogen can be bubbled through water, which is a polar substance. This causes the molecules to dissociate, and the H+ and halide- ions split, the H+ ions being acidic. This creates an acid, e.g. hydrochloric acid with chlorine, hydrobromic acid with bromine. This only works because water is a polar substance; it has charged ends. Non-polar substances, such as methylbenzene will not cause the compound to dissociate as the charge is evenly distributed.

Section 2 c) Key Words

Displacement reaction: A reaction in which a less reactive molecule is replaced with a more reactive molecule.

Dissociation: the splitting of a molecule into smaller molecules, atoms, or ions, especially by a reversible process

Group 7: The seventh group of the periodic table. Elements in this group have 7 valence electrons and are known as halogens. They share similar properties due to their similar electronic configurations.

Halogen: A group 7 element. Examples include chlorine and bromine.

Non-polar: A substance in which the electrons are shared equally between the nuclei, resulting in an even distribution of charge.

Oxidation: Loss of electrons

Polar: A substance in which the molecules are arranged so one end has a positive charge and on has a negative charge.

Reactivity series: The order of reactivity

Redox Reaction: A reaction in which electron(s) are transferred from one molecule to another. The molecule losing an electron is oxidised, and the one gaining is reduced.

Reduction: Gain of electrons

Section 2 c) Specification

2.9 recall the colours and physical states of the elements at room temperature

Fluorine, F : Pale yellow gas
Chlorine, Cl : Pale green gas
Bromine, Br : Orange liquid
Iodine, I : Purple solid
Astatine, At : Black solid

2.10 make predictions about the properties of other halogens in this group

Based on the information we know about fluorine, chlorine and bromine we can assume that as we travel down the group

• Reactivity decreases
• Melting and boiling point increase
• Elements get darker in colour

They all have antimicrobial properties in small doses, but are toxic in large doses.

2.11 understand the difference between hydrogen chloride gas and hydrochloric acid

Hydrogen chloride gas is the product of the reaction between hydrogen and chlorine. This can be dissolved in water to make it aqueous, causing the ions to be dissociated - detatched. This creates Cl - ions, as well as acidic H + ions, creating hydrochloric acid.

2.12 explain, in terms of dissociation, why hydrogen chloride is acidic in water but not in methylbenzene

Water is a polar substance, and methylbenzene is non-polar. Polar substances cause ionic bonding to be separated (dissociated), as the anions are attracted to the positive poles of the molecules, and the cations are attracted to the negative poles. This causes them to split. Non-polar substances aren't able to dissociate compounds in this way.

2.13 describe the relative reactivities of the elements in Group 7

They become less reactive as you go down the group, so larger elements are less reactive because the force of attraction between the nucleus and the valence electrons is weaker, so it is less able to attract another electron.

2.14 describe experiments to demonstrate that a more reactive halogen will displace a less reactive halogen from a solution of one of its salts

A more reactive halogen will displace a less reactive halogen bonded as a salt. For example, a sodium halide solution could be created for each of the halogens (potassium fluoride, potassium chloride, potassium bromide, etc.), then reacted with a halide solution, to discover the reactivity series:

Methylbenzene turns purple in presence of iodine, and yellow in presence of bromine. It can be added as a layer of indicator, showing us which substances are present (when bonded, does not affect indicator).
This series of experiments shows us that chlorine is the most reactive of these halogens, and iodine is the least, as chlorine water had the most displacements and iodine the least.

2.15 understand these displacement reactions as redox reactions.

Redox reactions are reactions that involve the loss and gain of electrons:
Oxidation
Is
Loss of electrons
Reduction
Is
Gain of electrons

When one more reactive halogen displaces a less reactive halogen, the more reactive halogen is reduced, and the less reactive one is oxidised.

As you can see in the above equations, more reactive chlorine reacts with sodium bromide. Bromine is less reactive than chlorine, so it displaces it. In this process, the bromine loses an electron and bonds with itself to form diatomic molecules. The chlorine molecules bond with sodium, and gain an electron, reducing them and giving them a negative charge.

Section 2 b) Summary

The elements of group 1 are:

• Li, Lithium
• Na, Sodium
• K, Potassium
• Rb, Rubidium
• Cs, Caesium
• Fr, Francium

They each have 1 valence electron, making them quite reactive. As you go down the group, from lithium to francium, the reactivity increases. This is because the number of shells of electrons increases, making the distance from nucleus to valence electrons further and decreasing the strength of the forces of electrostatic attraction. Because of these weaker forces, the atoms are able to lose their valence electrons more easily, making them more reactive. 

Group 1 elements react quickly and vigorously with cold water, indicating just how reactive they are. They must be stored under oil so they are unable to react with water or air while in storage. 

Reactions with cold water:

Lithium floats on the surface of the water due to its low density, gently fizzing and giving off hydrogen. It gradually reacts, forming an ionic compound that dissolves in water to form a lithium hydroxide solution. 

Sodium, like lithium, floats on the surface of the water. The heat from the reaction is enough to melt it, so as it moves around it forms a sphere, gradually decreasing in size. A white trail of sodium hydroxide follows this ball as it moves (due to hydrogen production propelling it), then dissolves to form a sodium hydroxide solution. It lets of yellow sparks as it reacts. 

Potassium reacts more violently than both elements before it. The reaction is faster, and the hydrogen produced is caught fire to, burning with a purple flame due to potassium contamination. Its product, potassium hydroxide, also dissolves. 

Rubidium, Caesium and Francium all react too violently to be demonstrated in a school, so this video shows how they (and the other ones too, with inaccuracy about francium which is unable to be obtained to react with water) react. 

Interesting (but unnecessary) information on Francium here (Francium is less reaction than caesium?)

Section 2 b) Key Words

Electrostatic attraction: Attraction between two charged particles, e.g. nuclei and electrons

Group 1: The first group of the periodic table. Each element has 1 valence electron and is extremely reactive. Examples include lithium, sodium, potassium, caesium, francium

Lithium: Li, the least reactive group 1 element. Has two shells of electrons, 3 protons, and a relative atomic mass of 7.

Potassium: K, the third element of group 1. Has four shells of electrons, 19 protons and a relative atomic mass of 39.

Sodium: Na, the second element of group 1. Has three shells of electrons, 11 protons and a relative atomic mass of 23.

Valence electrons: Outer shell electrons.

Section 2 b) Specification

2.6 describe the reactions of these elements with water and understand that the reactions provide a basis for their recognition as a family of elements

Lithium, Li: Fizzes vigorously, moves around the surface of the water.
Sodium, Na: Fizzes violently, moves quickly on the water, forms a ball shape & yellow sparks.
Potassium, K: Produces a purple flame, violently moves around the surface and pops with the force of hydrogen production.

These elements each react vigorously with cold water (due to their similar electronic configurations), an indication that they are the most reactive group of elements. They produce hydrogen gas as well as a metal hydroxide.

2.7 describe the relative reactivities of the elements in Group 1

The reactions start small at the top of the table, but gradually increase in violence as you travel down the table as the atoms are bigger and less able to hold onto their valence electron. Francium reacts extremely violently and explosively when reacting with cold water, whereas lithium just fizzes.

2.8 explain the relative reactivities of the elements in Group 1 in terms of distance between the outer electrons and the nucleus.

The elements get more reactive as you go down the group because the increase in distance between the valence electron and nucleus is increased due to increased number of shells, resulting in a weaker force of electrostatic attraction and a more reactive element.

Section 2 a) Summary

This is the most widely recognised version of the Periodic Table of Elements. At IGCSE level, the transition metals are not really relevant, so we look at a version like this:

The main difference is just that the middle groups are not counted, making it easier to recognise which group each element falls into. The Periodic table is organised based on the properties of the elements. A line can be drawn from between boron and aluminium to between polonium and astatine to separate metals and non-metals. The left side, and the majority of the table, is classified as metals. The right is non-metals.
Metals are characterised by:

• Basic (Alkali) hydroxides
• Electrical and heat conductivity
• Giant metallic structure/metallic bonds
• Ability to form cations

Non-metals are characterised by:

• Acidic hydroxides
• Insulating properties (except carbon)
• Covalent bonding
• Ability to form anions

Group 1 - Alkali metals
The first group is the alkali metals. They each have 1 valence electron. Reactivity increases as the periods increase, because the outer shells increase in distance for the nucleus and the forces of attraction weaken, so the element is able to lose this electron more easily.

Alkali metals react with cold water, an indication that they are extremely reactive.
Lithium, Li: Vigorous reaction, gives off bubbles of hydrogen gas, floats on the surface until the product created dissolves.
Sodium, Na: Faster, more vigorous reaction than lithium, melts into a sphere and lets off some yellow sparks.
Potassium, K: Violent reaction, creates a purple flame and pops with the force of hydrogen production.
The larger elements; Rb, Cs and Fr react so violently that it is unsafe to have in a school - they combust and create large explosions.
They also decrease in melting and boiling points as the atoms increase in size.

Group 7 - Halogens
Halogens are the seventh group of elements. They are non-metals that exist naturally as diatomic molecules. They each have 7 valence electrons; one short of a full outer shell. They all have antimicrobial properties in small quantities, but in large doses they are extremely toxic.

Halogens react with iron wool:

And with hydrogen gas:

As they become larger (descend down the periods?), Halogens become less reactive, darker in colour, more dense, and have a higher melting and boiling point.


Group 0 - Noble gases
Noble gases are the last group of the periodic table. Each of the elements has a full outer shell of electrons, so they are inert. Noble gases are the only elements that can form stable single-atom molecules at room temperature and pressure.
Noble gases include: Helium, neon, argon, krypton, xenon and radon.

Section 2 a) Key Words

Acid: A substance with a low pH (less than 7). The lower the pH, the more acidic (due to more H+ ions). Can neutralise bases and form salts.

Alkali metals: Metals in group 1 of the periodic table; they have 1 valence electron. They are highly reactive and form very strongly basic hydroxides.

Base: A substance with a high pH (more than 7). The higher the pH, the more basic (due to more OH- ions). Also known as alkalies, they can neutralise acids and form salts.

Diatomic molecules: Molecules that are comprised of two atoms.

Group: The columns of the periodic table. Elements of the same group have the same number of valence electrons.

Halogens: Elements in the 7th group of the periodic table. They have seven valence electrons, and are all non-metals, whose hydroxides are strongly acidic.

Inert: A substance that does not react. Noble gases are unreactive, therefore they are inert.

Metal: Elements that conduct electricity and heat, and whose hydroxides are basic. They bond together in a giant metallic crystal, with a sea of free electrons.

Noble gas: An element in the last group of the periodic table, group 0 or group 8. It is inert, and has a full outer shell of electrons

Non-metal: Elements that generally do not conduct electricity (carbon is an exception). Form covalent bonds involving a shared pair of electrons, and their hydroxides are acidic.

Period: The rows of the periodic table. Indicates how many shells of electrons there are.

Periodic table: A table of organisation for the chemical elements, developed by Mendeleev  (though not quite as we know it today, and there are many other versions)

Valence electrons: Outer-shell electrons

Section 2 a) Specification

2.1 understand the terms group and period 

The period is the row that the element is in, it indicates how many shells of electrons it has.

The group is the column in which the element can be found. It indicates how many valence electrons the element has.

2.2 recall the positions of metals and non-metals in the Periodic Table 

The non-metals lie to the right of the red line, and metals to the left:

2.3 explain the classification of elements as metals or non-metals on the basis of their electrical conductivity and the acid-base character of their oxides 

Metals are characterised by their conductive abilities, and that their oxides dissolve in water to form bases.
Non-metals are insulators (except for carbon), and their oxides dissolve in water to form acids.

2.4 understand why elements in the same group of the Periodic Table have similar chemical properties 

Elements in the same group have the same number of valence electrons: they have to gain or lose the same number of electrons. This means their chemical properties are similar, but as the periods change the characteristic gradually changes. (e.g. reactivity increases, the colour gradually darkens, etc.)

2.5 understand that the noble gases (Group 0) are a family of inert gases and explain their lack of reactivity in terms of their electronic configurations.

Noble gases have full outer shells, so they are unreactive. Reactive elements are unstable, and react to achieve a full outer shell, but because noble gases have full outer shells, they don't have electrons to gain or lose.

Section 2 Specification

a) The Periodic Table 

2.1 understand the terms group and period

2.2 recall the positions of metals and non-metals in the Periodic Table

2.3 explain the classification of elements as metals or non-metals on the basis of their electrical conductivity and the acid-base character of their oxides

2.4 understand why elements in the same group of the Periodic Table have similar chemical properties

2.5 understand that the noble gases (Group 0) are a family of inert gases and explain their lack of reactivity in terms of their electronic configurations.


b) Group 1 elements — lithium, sodium and potassium 

2.6 describe the reactions of these elements with water and understand that the reactions provide a basis for their recognition as a family of elements

2.7 describe the relative reactivities of the elements in Group 1

2.8 explain the relative reactivities of the elements in Group 1 in terms of distance between the outer electrons and the nucleus.


c) Group 7 elements — chlorine, bromine and iodine 

2.9 recall the colours and physical states of the elements at room temperature

2.10 make predictions about the properties of other halogens in this group

2.11 understand the difference between hydrogen chloride gas and hydrochloric acid

2.12 explain, in terms of dissociation, why hydrogen chloride is acidic in water but not in methylbenzene

2.13 describe the relative reactivities of the elements in Group 7

2.14 describe experiments to demonstrate that a more reactive halogen will displace a less reactive halogen from a solution of one of its salts

2.15 understand these displacement reactions as redox reactions.


d) Oxygen and oxides 

2.16 recall the gases present in air and their approximate percentage by volume

2.17 explain how experiments involving the reactions of elements such as copper, iron and phosphorus with air can be used to investigate the percentage by volume of oxygen in air

2.18 describe the laboratory preparation of oxygen from hydrogen peroxide, using manganese(IV) oxide as a catalyst

2.19 describe the reactions of magnesium, carbon and sulfur with oxygen in air, and the acid-base character of the oxides produced

2.20 describe the laboratory preparation of carbon dioxide from calcium carbonate and dilute hydrochloric acid

2.21 describe the formation of carbon dioxide from the thermal decomposition of metal carbonates such as copper(II) carbonate

2.22 describe the properties of carbon dioxide, limited to its solubility and density

2.23 explain the use of carbon dioxide in carbonating drinks and in fire extinguishers, in terms of its solubility and density

2.24 understand that carbon dioxide is a greenhouse gas and may contribute to climate change.


e) Hydrogen and water

2.25 describe the reactions of dilute hydrochloric and dilute sulfuric acids with magnesium, aluminium, zinc and iron

2.26 describe the combustion of hydrogen

2.27 describe the use of anhydrous copper(II) sulfate in the chemical test for water

2.28 describe a physical test to show whether water is pure.


f) Reactivity series 

2.29 understand that metals can be arranged in a reactivity series based on the reactions of the metals and their compounds: potassium, sodium, lithium, calcium, magnesium, aluminium, zinc, iron, copper, silver and gold

2.30 describe how reactions with water and dilute acids can be used to deduce the following order of reactivity: potassium, sodium, lithium, calcium, magnesium, zinc, iron and copper

2.31 deduce the position of a metal within the reactivity series using displacement reactions between metals and their oxides, and between metals and their salts in aqueous solutions

2.32 understand oxidation and reduction as the addition and removal of oxygen respectively

2.33 understand the terms redox, oxidising agent, reducing agent

2.34 describe the conditions under which iron rusts

2.35 describe how the rusting of iron may be prevented by grease, oil, paint, plastic and galvanising

2.36 understand the sacrificial protection of iron in terms of the reactivity series.


g) Tests for ions and gases 

2.37 describe tests for the cations:
i Li+, Na+, K+, Ca2+ using flame tests
ii NH4 +, using sodium hydroxide solution and identifying the ammonia evolved
iii Cu2+, Fe2+ and Fe3+, using sodium hydroxide solution

2.38 describe tests for the anions:
i Cl- , Br- and I- , using dilute nitric acid and silver nitrate solution
ii SO4 2- , using dilute hydrochloric acid and barium chloride solution
iii CO3 2- , using dilute hydrochloric acid and identifying the carbon dioxide evolved

2.39 describe tests for the gases:
i hydrogen
ii oxygen
iii carbon dioxide
iv ammonia
v chlorine.