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.

 

Section 1 i) Summary

Electrolysis is defined as chemical decomposition produced by passing an electric current through a liquid or solution containing ions. This process by which we can obtain the individual elements from an ionic compound has truly revolutionised the world.

Ionic compounds do not conduct electricity while solid. This is because the ions are not free to move. When aqueous, the ions are dissociated and therefore are able to move freely. When molten, the bonds between the ions are broken and they are able to move freely. Electricity is able to pass through because there are free-moving charged particles that can carry current.
Ionic compounds have very high melting points, and because of this the solid is often dissolved in water to create an aqueous electrolyte. When molten ionic compounds are electrolysed, the rules are simple: Metal goes to the cathode, Non-metal to the anode. But with aqueous solutions, it is less simple.

At the anode: If the non-metal is a halogen, it will form at the anode. If it is not, hydrogen gas will form: 2H+ + 2e- --> H2

At the cathode: If the metal is a low reactivity metal, it will coat the cathode. If it is a high reactivity metal, oxygen gas will form: 4OH- --> O2 + 4H2O + 4e-


The amount of a substance removed can be calculated by how much charge has passed through.
1 mole of electrons is equal to 1 faraday, which is 96500 coulombs.
First, you have to know the ionic half equation, the current passing through, and the time it takes.
Then, you can calculate the charge with this equation:
Charge (in coulombs) = Current (in amps) x Time (in seconds)
Q = I T
Divide the charge you have just calculated by 96500 to get the number of faradays, which is how many moles of electrons transferred.
Then look at the equation to see how many moles of electrons there are for each mole of the product, and divide accordingly.
If you are trying to find the mass, multiply by Mr, and you will have mass in grams
If you are trying to find the gas volume, multiply by 24000 to get volume in cm3.
See examples in the specification here

Industry
One of the most common industrial uses of electrolysis is the electrolysis of brine, or sodium chloride. This is called chlor-alkali industry.

The products of this reaction are used for many different things:

Chlorine: killing bacteria, making bleach, making hydrochloric acid, etc.

Hydrogen: used in making ammonia, making margarine, etc.

Sodium hydroxide: Making soap, paper, ceramics, etc.

The ionic half-equations are:

2Cl- --> Cl2 + 2e-

2H+ + 2e- --> H2

Because two gases are formed, a diaphragm is put in place between the two electrodes to prevent them from mixing and reacting.

Section 1 i) Key Words

Anode: Positive electrode. Attracts negative anions.

Anions: Negatively charged ions

Cathode: Negative electrode. Attracts positive cations.

Cations: Positively charged ions

Electrode: A conductor placed in an electrolyte to electrolyse it

Electrolysis: The process by which a molten or aqueous ionic compound is separated by passing a current through.

Electrolyte: A molten or aqueous ionic compound that conducts electricity.

Half equation: An ionic half equation includes only half a chemical reaction, it just shows what is happening at the individual electrodes with the transfer of electrons.

Section 1 i) Specification

1.48 understand that an electric current is a flow of electrons or ions

Electric current is the movement of charged particles in a sort of stream. Usually it is in reference to electrons, but free ions can also conduct electricity/carry charge, as they are charged particles.

1.49 understand why covalent compounds do not conduct electricity

In covalent compounds, there are no free charged particles (no delocalised electrons, no ions) available to carry charge.

1.50 understand why ionic compounds conduct electricity only when molten or in
solution

Ionic compounds, when solid, do not conduct electricity. This is because the ions are not free to move; they are held in place by strong ionic bonding. When molten or aqueous, the ions become free and can then carry electrical charge.

1.51 describe experiments to distinguish between electrolytes and nonelectrolytes

Make a simple circuit involving a buzzer or a lightbulb, then, in series with the chosen component, place a beaker containing the liquid in question with electrodes attached into the circuit. If the component works, it is an electrolyte, because it allowed current to flow through. If the component does not work, it is a nonelectrolyte as it has broken the connection of the circuit.

1.52 understand that electrolysis involves the formation of new substances when
ionic compounds conduct electricity

Electrolysis is a process designed to separate ionic compounds and re-form them back into their individual elements.

1.53 describe experiments to investigate electrolysis, using inert electrodes, of
molten salts such as lead(II) bromide and predict the products

1. Place carbon electrodes into a beaker containing lead bromide, so they are submerged. 
2. Hook the electrodes up to a cell and allow current to flow for some time.
3. At the cathode (negative electrode), positive metal ions will be reduced, in this case lead will form. 
4. At the anode (positive electrode), negative non-metal ions will be attracted and oxidised, in this case bromine will form. 

1.54 describe experiments to investigate electrolysis, using inert
electrodes, of aqueous solutions such as sodium chloride, copper(II)
sulfate and dilute sulfuric acid and predict the products

Sodium chloride

1. Place carbon electrodes into a beaker containing lead bromide, so they are submerged. 
2. Hook the electrodes up to a cell and allow current to flow for some time.
3. At the cathode (negative electrode), positive hydrogen ions will be reduced, and bubbles of hydrogen gas will form
4. At the anode (positive electrode), negative halide ions will be attracted and oxidised, in this case chlorine will form. 

Copper (II) sulphate

1. Place carbon electrodes into a beaker containing lead bromide, so they are submerged. 
2. Hook the electrodes up to a cell and allow current to flow for some time.
3. At the cathode (negative electrode), positive copper ions will be reduced, and copper will coat the electrode.
4. At the anode (positive electrode), negative hydroxide ions will be attracted and oxidised, and bubbles of oxygen gas will form. 

Dilute sulphuric acid

1. Place carbon electrodes into a beaker containing lead bromide, so they are submerged. 
2. Hook the electrodes up to a cell and allow current to flow for some time.
3. At the cathode (negative electrode), positive hydrogen ions will be reduced, and bubbles of hydrogen gas will form
4. At the anode (positive electrode), negative hydroxide ions will be attracted and oxidised, and bubbles of oxygen gas will form.

1.55 write ionic half-equations representing the reactions at the electrodes during
electrolysis

Lead bromide
Cathode: Pb2+ + 2e- --> Pb
Anode: 2Br- --> Br2 + 2e-

Sodium chloride
Cathode: 2H+ + 2e- --> H2
Anode: 2Cl- --> Cl2 + 2e-

Copper (II) sulphate
Cathode: Cu2+ + 2e- --> Cu
Anode: 4OH- --> O2 + H2O + 4e-

Dilute Sulphuric acid
Cathode: 2H+ + 2e- --> H2
Anode: 4OH- --> O2 +H2O + 4e-

1.56 recall that one faraday represents one mole of electrons

One faraday = 96500 coulombs = 1 mole of electrons

1.57 calculate the amounts of the products of the electrolysis of molten
salts and aqueous solutions.

If a current of 1.5A is passed through lead bromide for 1 hour, how much lead will form?
Half equation: Pb2+ + 2e- --> Pb
Coulombs = Current (in amperes) x Time (in seconds) = 5400 C

Conversion into moles of electrons:
1 mole = 1 faraday = 96500 coulombs
5400 / 96500 = 54/965 mol

There are 2 moles of electrons for every mole of lead, so divide by 2 = 27/965 mol
Then multiply by Mr to find mass 27/965 x 207 = 5.79 g 


If a current of 2A is passed through sodium chloride for 2 hours, what volume of chlorine will form?
Half equation: 2Cl- --> Cl2 + 2e-
Q = I T = 14400 C

1 mole = 1 faraday - 96500 C
14400 / 96500 = 144/965 mol

2 electrons per atom, so divide by 2 = 72/965 mol
Then multiply by 24 000 to find volume in cm3 72/965 x 24000 = 1790.67 cm3


If a current of 4.2A is passed through copper (II) sulphate for 45 minutes, what mass of copper will be formed?
Half equation: Cu2+ + 2e- --> Cu
Q = I T = 11340C

11340 / 96500 = 567/4825 mol

2 electrons per atom, so divide by 2 = 567/9650 mol
Then multiply by Mr 567/9650 x 63.5 = 3.73 g


If a current of 0.2A is passed through dilute sulphuric acid for 5 hours and 45 minutes, what volume of hydrogen is produced?
Half equation: 2H+ + 2e- --> H2
Q = I T = (5 x 60 + 45) x 60 x 0.2 = 4140C

4140 / 96500 = 207/4825 mol

2 moles of electrons per mole of hydrogen, so divide by 2 = 207/9650 mol
Then multiply by 24 000 to get the volume in cubic centimetres 207/9650 x 24 000 = 514.82 cm3

Section 1 h) Summary

When metals bond, they form a giant metallic structure, also known as a metallic crystal. They are made up of positively charged nuclei surrounded by a delocalised electrons, sometimes referred to as a 'sea of free electrons'. These free electrons allow metals to conduct electricity, as they are able to move and carry charge freely. 

This structure is a giant structure, meaning it is constantly repeating regularly. It is in sort of layers which mean that when force is applied, the ions will slide, and the metal will bend, not break. 

The strong forces of attraction between the electrons and charged metal ions are difficult to break; they require a lot of energy. This means that metals have high melting and boiling points. 

Metals are also:

• Shiny/lustrous: The giant structure means it has flat, smooth surfaces that reflect light.
• Hard: The giant structure is densely packed and difficult to dent
• High density: The ions are packed closely together, only tiny electrons keeping them apart
• Strong: The giant structure is difficult to break 

Section 1 h) Key Words

Delocalised electrons: Electrons that are not tied to a specific nucleus, they are free to move around and allow electricity to be conducted.

Ductile: Able to be stretched out/pulled out into a wire. Metals are ductile.

Giant structure: A constantly repeating lattice structure of molecules, it doesn't end until there aren't enough molecules to add more.

Malleable: Easily bent when force is applied, the opposite of brittle.

Section 1 h) Specification

1.46 understand that a metal can be described as a giant structure of positive
ions surrounded by a sea of delocalised electrons

In metallic bonding, the positive ions are regularly arranged in a giant structure. These are surrounded by a sea of delocalised electrons, without which the cations would repel each other; the attraction between the positive ions and electrons holds the structure together.

1.47 explain the electrical conductivity and malleability of a metal in terms of its
structure and bonding.

Metal is a good conductor of both electricity and heat. This is because of its sea of delocalised electrons that are free to move and carry charge.

Metal is malleable and ductile because of its giant structure. It is in layers that can slip past one another when force is applied, as opposed to snapping. 

Section 1 g) Summary

Covalent bonding is the bonding of two non-metal atoms. It involves a shared pair of electrons. The attraction between the nucleus of each of the atoms and the electrons of the other atom creates a very strong covalent bond.

Simple covalent structures usually have low melting and boiling points, and are often fluid at room temperature, because the forces between the molecules are weak and don't require much energy to break. They have weak intermolecular forces.
Examples include: water, oxygen, carbon dioxide, hydrogen chloride, etc.

Giant covalent structures have very high melting and boiling points, because they have so many strong covalent bonds holding the atoms together. It requires a lot of energy to break that many covalent bonds, so they have high melting points.
Examples include: diamond (carbon), graphite (carbon), silicon dioxide, etc.

Dot and cross diagrams can be used to represent covalent bonding. Here are the most important ones for this topic:

Methane

Water

Hydrogen Chloride

Chlorine

Hydrogen

Ammonia

Oxygen

Nitrogen

Carbon Dioxide

Ethane

Ethene

Section 1 g) Key Words

Allotrope: The different molecular forms in which an atom can exist.

Covalent bond: A bond involving the attraction of a pair of electrons and a pair of nuclei from two non-metal atoms

Diamond: An allotrope of carbon with a tetrahedral structure. Used in jewellery and in power tools, as it is the hardest material naturally occurring on Earth and can cut through virtually any solid.

Intermolecular forces: The weak forces of attraction between covalent molecules

Graphite: An allotrope of carbon in a layered structure that easily slides. Used in pencils and as a lubricant, as it transfers easily and slides easily.

Tetrahedral: An atom directly bonded with four other atoms: it has four covalent bonds. Structure of diamond.

Section 1 g) Specification

1.38 describe the formation of a covalent bond by the sharing of a pair of
electrons between two atoms

A covalent bond is formed between two non-metals and involves the atoms sharing one or more pairs of electrons

1.39 understand covalent bonding as a strong attraction between the bonding
pair of electrons and the nuclei of the atoms involved in the bond

In a covalent bond, the nucleus of each atom is attracted to the other atom's outer shell electron(s).

1.40 explain, using dot and cross diagrams, the formation of covalent compounds
by electron sharing for the following substances:
i hydrogen
ii chlorine
iii hydrogen chloride
iv water
v methane
vi ammonia
vii oxygen
viii nitrogen
ix carbon dioxide
x ethane
xi ethene

1.41 understand that substances with simple molecular structures are gases or
liquids, or solids with low melting points

Simple covalent compounds have strong bonds, but very weak intermolecular forces. These don't require much energy to break, leading to low melting and boiling points. As a result, most simple covalent compounds are gas or liquid at room temperature.

1.42 explain why substances with simple molecular structures have low melting
and boiling points in terms of the relatively weak forces between the
molecules

There aren't very strong forces of attraction between covalent molecules, they have weak intermolecular forces that are easily broken; they don't require much energy to break. This means that they have low melting and boiling points.

1.43 explain the high melting and boiling points of substances with giant covalent
structures in terms of the breaking of many strong covalent bonds

In giant covalent structures, each atom is bonded to another with by strong covalent bonds. These are not easily broken, they require a lot of energy to break. Additionally, in a giant structure there are so many bonds to break that it requires even more. This results in very high melting and boiling points.

1.44 draw diagrams representing the positions of the atoms in diamond
and graphite

1.45 explain how the uses of diamond and graphite depend on their
structures, limited to graphite as a lubricant and diamond in cutting.

In diamond, the carbon atoms are arranged in a tetrahedronal structure, each atom with four strong covalent bonds. This makes it very hard, with a very high melting and boiling point. It has lots of very strong bonds. Its giant structure means that it forms in crystals that reflect light and let it pass through. It is so hard that its use in industrial saws etc. is ideal, it can cut through almost anything.

Graphite is arranged in layers held together by weak intermolecular forces. This means that when force is applied, the layers slide over each other easily, making graphite a good lubricant.

Both diamond and graphite are allotropes of carbon, but buckminsterfullerene is another one, it's ball shaped, made of 60 carbon atoms, and used for lubrication, drug administration, and in catalysts. 

Section 1 f) Summary

Ionic compounds are formed when a metal and non-metal meet, and there is a transfer of electrons. This causes both atoms to become oppositely charged ions, which experience strong forces of electrostatic attraction towards each other.

Ionic compounds are not easily broken apart, their strong forces of attraction mean they have high melting and boiling points: they require a lot of energy to break the bonds.
The compounds, as a solid, form a giant ionic lattice structure. It is a lattice, meaning it is regular which results in the solids forming into translucent, geometric crystals (think of salt).

This can be represented at a molecular level using the following diagram:

The positive and negative ions alternate, because like attractions repel and opposites attract. Therefore, when force is applied to an ionic solid, it breaks easily: it is brittle. This is because when the ions are forced to move, they don't slide over each other easily, they get close to a similar charge, causing a force of repulsion and making it break apart easily.

It is important to know the charges of ions seen commonly in this topic. We can deduce the charge based on the number of outer shell (valence) electrons (if it has two, it will lose them, if it has six, it will gain two, etc.)
The following periodic table is a useful guide on the trends:

Section 1 f) Key Words

Anion: Negatively charged ion

Cation: Positively charged ion

Ion: Charged particle, atom with extra electron/s or missing electron/s

Oxidation: Loss of electrons

Reduction: Gain of electrons

Section 1 f) Specification

1.28 describe the formation of ions by the gain or loss of electrons

An ion is formed when an atom loses an electron (It is reduced) or gains an electron (it is oxidised). This change in electron number means the atom has gained a charge. If it has been oxidised, it will become a negative ion: an anion. If it has been reduced, it will become a positive ion: a cation. In ionic bonding, one or more electrons is transferred from one atom to another, turning both into oppositely charged ions which are strongly attracted to one another.

1.29 understand oxidation as the loss of electrons and reduction as the gain of
electrons

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

1.30 recall the charges of common ions in this specification

Lithium: Li+
Chlorine: Cl-
Fluorine: F-
Hydrogen: H+
Potassium: K+
Magnesium: Mg2+
Oxygen: O2-

General rule: non-metals are negative, metals are positive

1.31 deduce the charge of an ion from the electronic configuration of the atom
from which the ion is formed

You can look at the outer shell of an atom to determine its ionic charge. If the atom has 2 outer shell electrons, it will lose them and have a charge of 2+. If it has 5 outer shell electrons, it will gain 3, leaving it with a charge of 3-.

1.32 explain, using dot and cross diagrams, the formation of ionic compounds by
electron transfer, limited to combinations of elements from Groups 1, 2, 3
and 5, 6, 7

Ionic bonding is between a metal and a non-metal, the outer shell electrons transferring from the metal to the non metal to give both a full outer shell.

1.33 understand ionic bonding as a strong electrostatic attraction between
oppositely charged ions

The ions become oppositely charged when the electrons are transferred. They experience strong forces of attraction, because opposite charges attract. This forms a strong ionic bond.

1.34 understand that ionic compounds have high melting and boiling points
because of strong electrostatic forces between oppositely charged ions

After transferring electrons, atoms become charged ions which experience strong attraction to opposite charges. This means the electrons are bound by strong electrostatic forces of attraction in this structure:

The strong forces of attraction require a lot of energy to be broken, which is why they have a high melting and boiling point.

1.35 understand the relationship between ionic charge and the melting
point and boiling point of an ionic compound

The bigger the charge, the higher the melting and boiling point as higher charge means stronger forces of attraction.

1.36 describe an ionic crystal as a giant three-dimensional lattice
structure held together by the attraction between oppositely
charged ions

Ionic crystals are held together by the strong forces of attraction between the positively and negatively charged ions. This forms a giant structure, meaning it repeats over and over, and it is a three-dimensional lattice, meaning it is a regular structure that has gaps etc. that allow light to pass through; this is why salt crystals are translucent.

1.37 draw a diagram to represent the positions of the ions in a crystal of
sodium chloride.

Green: Chlorine
White/Grey: Sodium

Section 1 e) Key Words

Balanced equation: A chemical equation that is written so the numbers of molecules do not change from one side to the other of the equation.

Concentration: How many molecules of a substance there are in a fixed volume, how dilute it is.

Crucible: A piece of apparatus - a ceramic "pot" that can be used for heating substances to high temperatures. Includes a lid.

Percentage Yield: The percentage of a substance you collect as a percentage of how much was created in a reaction.

State Symbol: Indicates the state of a chemical at room temperature in a reaction (unless otherwise specified), e.g. (g) for gas, (aq) for aqueous. Found in a chemical equation written in the bottom right corner of a molecule.

Yield: The amount of a substance you retrieve from an experiment/chemical reaction

Section 1 e) Summary

This topic is about chemical formulae and equations, how to write and balance them correctly, as well as how to calculate reacting masses.

In a chemical reaction, the number and types of particles should be the same on both sides. But as different molecules require different numbers of each element, the ratio of each molecule created is relative. We must balance the equation to ensure the number of particles, for example:

H2 + O2 --> H2O
The above equation is unbalanced. There are two particles of hydrogen, and two particles of oxygen on the reactants side, but on the products side there are two particles of hydrogen and only one particle of oxygen.
We can remedy this by first changing the particle with the deficit:
H2 + O2 --> 2H2O
Now there are two oxygen particles on either side of the equation. But there are also double the number of hydrogen particles in the products than in the reactants, but this is easily fixed:
2H2 +O2 --> 2H2O
Now there are 4 hydrogen and 2 oxygen on both sides of the equation: it is balanced.

We can use balanced chemical equations to figure out the reacting masses, as it gives us a ratio of the moles, which when multiplied by Mr gives the mass, as shown in this formula triangle:

For example, if we have this balanced equation:

Fe2O3 + 3CO --> 2Fe + 3CO2

And we are told that there is 75 grams of iron oxide, and that we have to calculate what mass of iron can be derived from this with excess carbon monoxide, it can be calculated as follows:

Mr of Iron Oxide: 160

Mass of Iron Oxide: 75 g

75 / 160 = 15/32 mol

Ratio of Fe2O3 : Fe is 1:2

15/32 x 2 = 15/16 mol

Mr of Fe: 56 

56 x 15/16 = 52.5g

52.5 grams of iron can be derived from this reaction 

However, the actual iron we get from this reaction in reality is lower. This is because some may be lost in separation etc., this is called the yield. 

Percentage yield is how much is actually collected of the amount that is expected, and can be calculated with the following equation:

Yield = (mass of product collected x 100) / mass of product calculated

This tells us what percentage of the product is lost in this method, and can be used to compare different methods, as well as give a more accurate prediction of what the actual yield will be. 

For this topic, it is important to know the specifics of two different reactions, which simply demonstrate that you know how to calculate reacting masses and percentage yield:

Metal Oxide
Weigh a crucible with its lid, then place magnesium inside and weigh again, recording the measurements.
Heat it strongly over a roaring bunsen flame, and gently open the lid with tongs momentarily, to allow oxygen to get in without letting the magnesium oxide escape.
Weigh it every so often so you can see how the mass is changing. When the mass stops changing, the reaction is complete.
Record the final mass of the magnesium oxide and crucible, and use the numbers you recorded earlier to determine the mass of magnesium and the mass of oxygen.
Using each element's relative atomic mass, the empirical formula can be determined.

Salt (crystallisation)
Weigh an evaporating basin, then add hydrated copper sulfate and weigh again.
Heat the solution over a bunsen burner, gently stirring to ensure even heating.
Stir until the solution loses colour, indicating all water has been lost.
Weigh again to find the mass of water lost, and also the mass of anhydrous copper sulfate.
Divide mass by Mr of the copper sulfate, and do the same with water.
Simplify this to find the ratio of water to copper sulfate, and round to the nearest whole number,
e.g. CuSO4 ● 5H2O

Section 1 e) Specification

1.21 write word equations and balanced chemical equations to represent the
reactions studied in this specification

Copper + Oxygen --> Copper Oxide
2Cu + O2 --> 2CuO

Carbon + Oxygen --> Carbon Dioxide
C + O2 --> CO2

Hydrogen + Oxygen --> Water (Dihydrogen Monoxide)
2H2 + O2 --> 2H2O

1.22 use the state symbols (s), (l), (g) and (aq) in chemical equations to
represent solids, liquids, gases and aqueous solutions respectively

(s) - the substance is solid at room temp (unless otherwise specified in the question)
(l) - the substance is liquid at room temp (unless otherwise specified in the question)
(g) - the substance is gas at room temp (unless otherwise specified in the question)
(aq) - the substance is dissolved in water

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

Through experiments to remove an element from a compound, the chemical formula can be derived. The reactants are first weighed, and then the mixture is split (You could work out a metal oxide or salt)

Metal Oxide
Weigh a crucible with its lid, then place magnesium inside and weigh again, recording the measurements.
Heat it strongly over a roaring bunsen flame, and gently open the lid with tongs momentarily, to allow oxygen to get in without letting the magnesium oxide escape.
Weigh it every so often so you can see how the mass is changing. When the mass stops changing, the reaction is complete.
Record the final mass of the magnesium oxide and crucible, and use the numbers you recorded earlier to determine the mass of magnesium and the mass of oxygen.
Using each element's relative atomic mass, the empirical formula can be determined.

Salt (crystallisation)
Weigh an evaporating basin, then add hydrated copper sulfate and weigh again.
Heat the solution over a bunsen burner, gently stirring to ensure even heating.
Stir until the solution loses colour, indicating all water has been lost.
Weigh again to find the mass of water lost, and also the mass of anhydrous copper sulfate.
Divide mass by Mr of the copper sulfate, and do the same with water.
Simplify this to find the ratio of water to copper sulfate, and round to the nearest whole number,
e.g. CuSO4 ● 5H2O

1.24 calculate empirical and molecular formulae from experimental data

By weighing compounds, then extracting an element from them through a chemical reaction and recording the change in mass, then making these masses into percentages and dividing by Ar, we can find the number of atoms of it in a molecule. Empirical and molecular formulae of compounds and elements can be found this way.

1.25 calculate reacting masses using experimental data and chemical equations

The mass of the reactants is always equal to the mass of the products. Using this principle, applied to equations, we can calculate the different masses.

1.26 calculate percentage yield

(yield retrieved  x 100 )/ full potential yield = percentage yield

The expected or potential yield can be calculated in a reacting mass calculation, however we are not always able to recover the full yield. The amount we retrieve can be calculated as a percentage of the full potential yield using the above equation.

1.27 carry out mole calculations using volumes and molar concentrations.

Using the above formula triangle, the moles, volume and concentration (molar volume) can be calculated in a chemical equation.