# GCSE & A Level Revision Notes

## Edexcel IGCSE - Seperate Science

### Section 1 (1.1-1.57)

#### 1.1

Solids: Regular arrangement and little movement (only vibrate around a fixed point) and have low energy, not enough to overcome intermolecular forces of attraction.

Liquids: Weaker forces of intermolecular attraction and are free to move around but still close together, constantly move in random motion.

Gases: Very weak forces of intermolecular attraction – move freely but the particles in gases travel in straight lines. Move faster when heated and have more pressure the hotter they get due to kinetic energy.

#### 1.2

Solid –melting→ liquid –boiling→ gas
←Freezing-   ←condensing-

Movement from solid straight to gas and vice versa is known as sublimation. To move from a solid to a liquid or a liquid to a gas more energy has to be added to overcome the forces of intermolecular attraction.

#### 1.3

Solid → Liquid: Particles have more energy so there are weaker forces of intermolecular attraction and therefore it moves more and more freely – the arrangement is less regular. Liquids → Gases: Particles have even more energy so move around in fast random motion. The arrangement is completely irregular and there are only very weak forces of attraction.

### Dilution of coloured solutions

#### 1.4

• Potassium Manganite (VII) and water – purple colour spreads out which results in a pale, even coloured solution
• Tube – one end contains cotton wool soaked in aqueous ammonia and the other end is soaked in hydrochloric. A white ring of ammonium chloride forms nearer the HCl because ammonia particles are smaller and lighter and therefore the ammonia moves more.
• Fill half a jar with bromine gas and the other half with air, separated by a glass plate. When the plate is removed the bromine will diffuse into the air, resulting in an even colour throughout.

#### 1.5

An atom is made out of protons and neutrons in a nucleus surrounded by electrons orbiting in shells. A molecule is a group of atoms that are chemical bonded.

#### 1.6

An element contains only one type of atom and is pure. Two or more elements that are chemically combined make a compound, and a mixture is a group of two or more elements that have been physically mixed but not chemically bonded.

### Separation Techniques

#### 1.7

Simple Distillation – Used for separating out liquid from a solution (a solute and a solvent.) The solution is heated so that the part with the lowest boiling point evaporates, usually pure water in most solutions. This runs down the condensing tube, which is surrounded by a cold water jacket. This condenses and runs into a conical flask.

Fractional Distillation – Used to separate out a mixture of liquids such as crude oil. There is a fractionating column filled with glass rods, which highlight the differences in boiling points of the different liquids. The compounds with the highest boiling points turn back into liquids first, where they are tapped off.

Filtration – Used to separate insoluble solids from liquids. Filter paper is put into a funnel and then the mixture is poured through, leaving the solid behind.

Crystallization – Used to separate out a solute from a solution. The solution is heated over a Bunsen burner until crystals are just starting to form on the very edges, and then it is placed in a cooling dish and left in a warm place to allow the rest of the liquid to evaporate naturally.

Chromatography – Used to separate out dyes into different parts. The solvent is left just under the dyes (so they don’t simply dissolve into it) and then it travels up the paper, carrying the dyes with it. The most soluble dyes travel the furthest.

#### 1.8

Chromatography can be used to determine what a dye is composed of. For example, if dye A is your unknown dye, and you test it against 4 colours, whichever colour the dye travels the same distance as, Dye A contains.

### Atomic Structure

#### 1.9

Atoms consist of a central nucleus containing protons and electrons, which are surrounded by electrons that orbit in shells.

#### 1.10

Protons have a charge of positive one (+1) and a relative mass of one, neutrons are neutral (0) and have a relative mass of one and electrons have a charge of negative one (-1) and have a relative mass of 1/2000.

#### 1.11

Atomic Number: The number of protons in the nucleus of an atom, this determines which element it is.

Mass Number: The number of protons and neutrons in an atom. This makes up the nucleus, and is an average calculation based on relative abundance.

Isotopes: Atoms with a different number of neutrons but the same number of electrons.

Relative Atomic Mass: A weighted average of all the isotopes of an element. How heavy the nucleus of an atom is compared to 1/12 of Carbon 12.

#### 1.12

To find relative atomic mass based on relative abundance:

1. Multiply each mass by its relative abundance (usually a percentage)
3. Divide them by the total of their relative abundances

#### 1.13

The modern periodic table is arranged in order of atomic number, and new elements are being made, which have to be added to the end of the periodic table.

#### 1.14

An atom has the same number of protons as neutrons. By looking at the atomic number of an element, you can also discover the number of electrons that it has. The first shell contains 2 electrons; the next shell has 8 and then 8 again after that.

#### 1.15

The group to which an element belongs corresponds to the number of electrons that it has in its outer shell. The period to which it belongs corresponds to the number of shells of electrons there are orbiting the nucleus.

#### 1.16

To calculate the relative formula mass of formulae from atomic masses, simply add together the relative atomic masses. E.g. CO2 is made up of C (12) and 2O (16x2=32) → 12+32 = 44

#### 1.17

A mole is representative of the amount of a substance that is available, in terms of particles to react. It is calculated by doing mass divided by RAM.

#### 1.18

A mole is also the Avogadro number (6.02 x1023) particles that are contained in a substance.

#### 1.19

The number of moles in a substance can be calculated using these molar triangles. The one on the left is for solutions, the middle one is for gasses and the right hand one is for dry matter.

#### 1.20

The molar volume of a gas is the amount of space that it takes up per mole. This is 24 dm3 or 24000 cm3 at room temperature and pressure.

### Chemical Formulae

#### 1.21

Both sides of an equation must be balanced, i.e. they must each contain the same amount of each element, due to the law of conservation of mass.

#### 1.22

State symbols are used to represent the states of matter that the products and reactants are in during the equation.

• Solid (s)
• Liquid (l)
• Aqueous (aq)
• Gas (g)

#### 1.23

Find the weight of Mg and MgO (as an example) and then use the molar triangle to calculate masses of each elements with water of crystallization can be found by weighing an anhydrous salt such as CUSO4 then a hydrated version (CuSO4 - 5H2O) and measure the difference. You can then use this to find the formulae.

#### 1.24

An empirical formula is the lowest ratio that a formula can be put into, for example, the empirical formula for C2H4Br2 would be CH2Br. It can be calculated using experimental data.

1. Find moles of each element (mass/RAM)
2. Ratio moles (divide by the smallest number)
3. Record formula To multiply this to the molecular formula, find the RFM of the empirical formula, and divide it into the RFM. If you get 2, then the whole formula needs to be multiplied by 2.

#### 1.25

You can find the reacting masses of elements due to the law of conservation of mass. You do this by finding the moles of the original substance, ratio-ing the moles and then multiplying up the moles of the product to find the mass of product.

1. Calculate RAM of reactant (mass given)
2. Find moles of reactant
3. Ratio moles (reactant: product)
4. Calculate mass of product

#### 1.26

You can calculate how efficient a reacting process was by its percentage yield:

1. Calculate expected yield
2. Divide actual yield by calculated yield
3. Multiply by 100

#### 1.27

Volume = moles x molar volume E.g. what volume does 0.65 moles of gas occupy? → 0.65 x 24000 = 15600 cm3

### Ionic Compounds

#### 1.28

Ions are formed when an atom becomes either positively or negatively charged. This is due to the loss or gain of electrons, which are negatively charged.

#### 1.29

REDOX – OILRIG Oxidation is the loss (of electrons) and reduction is the gain (of electrons.) When one of these processes happens so does the other, because the electrons need somewhere to come from and go to, the do not just appear or disappear.

#### 1.30

There are many ions which are common, and which have different valances. Here are some of the common ones that appear frequently:

• K+
• Na+
• Li+
• H+
• Mg 2+
• Ca 2+
• Al 3+
• Cl-
• Br-
• I-
• F-
• OH-
• NO3-
• SO4 2-
• CO3 2-

#### 1.31

You can deduce the charge of different ions from the electronic configuration of the neutral atom that it comes from and vice versa. E.g. calcium’s electronic configuration is [2,8,8,2], which means it wants to lose 2 electrons from its outer shell – the charge of its ion is 2+.

#### 1.32

Group 1 has one outer shell electron, which means it wants to give one electron away so that it has a full outer shell, whereas group 7 has 7 outer shell electrons, which means an atom in this group needs to gain one electron to have a full outer shell. Group 2 loses 2 electrons so that it has a full outer shell, whereas group 6 wants to gain 2 electrons.

Example: Sodium [2, 8, 1] bonding with Chlorine [2,8,7] – Sodium gives one electron to chlorine so they become Na+ and Cl-, with configurations of [2,8]+ and [2,8,8]-

If the valance of the elements is different, sometimes two of one ion is needed to bond with one of another.

#### 1.33

Ionic bonding is the strong electrostatic attraction of oppositely charged ions, such as in NaCl, which is made up of negative chlorine ions and positive sodium ions.

#### 1.34

Ionic compounds generally have high melting and boiling points. This is because of the strong electrostatic forces of attraction that create one giant lattice structure. All the forces of attraction in this structure must be overcome at once to allow the substance to change state, which takes a lot of energy.

#### 1.35

The higher the ionic charge, the stronger the forces of attraction between the different ions. E.g. if there are two ions with a charge of one, one positive and the other negative, then although the resulting charge is neutral, there is a resultant attraction due to difference in charge of 2. If there are two ions with a charge of 2, one positive and one negative then the resultant difference will be 4.

#### 1.36

An ionic crystal is one molecule. It is a giant, three-dimensional lattice that is held together by forces of attraction between oppositely charged ions.

#### 1.37

The ionic lattice can be drawn like this. The corners of the cube in this case are the Cl- ions and the ions in the middle are the Na+ ions.

### Covalent Substances

#### 1.38

The formation of a covalent bond is the sharing of one or more pairs of electrons with one electron from each pair coming from each atom.

#### 1.39

Covalent bonding is a strong attraction between the bonding pair of electrons and the nuclei of both the atoms involved in the bonding.

#### 1.40

Oxygen shares two pairs of electrons in its covalent bond, which is called a double bond, and nitrogen shares three, making it a triple bond.

#### 1.41

There are two types of covalent structures, simple discrete molecules, such as water, which have much lower melting points and giant covalent structures.

#### 1.42

Simple molecular structures have relatively low melting and boiling points and therefore are almost always found as gases and liquids. This is due to the weak forces of electrostatic attraction between the molecules, which require little force to overcome. There is no bond between each molecule, only an intermolecular force, which is weak.

#### 1.43

Giant covalent structures are almost always solids as they have very high melting and boiling points. This is because they have a crystalline structure, which means each crystal is one giant molecule, and all the bonds must be broken at once for it to become a liquid. The electrostatic forces of attraction within the molecule are very strong.

#### 1.44

Structure of Diamond Structure of Graphite

#### 1.45

Diamond is very hard, due to the four strong covalent bonds on each carbon atom. This means that it is very good for cutting, as it will not break – it is the hardest natural substance on Earth. Graphite is used as an industrial lubricant because of its layered structure. This allows one layer to slide over another quire easily, which limits the amount of friction.

### Metallic Crystals

#### 1.46

Metals can be described as a giant structure of positive electrons, which is surrounded by a sea of delocalized electrons.

#### 1.47

Metals are malleable and good conductors of heat and electricity, due to the sea of delocalized electrons. Only the metal ions are fixed in place, which means that they are malleable and also good for heat transfer because the ions can move. This also means that they are good conductors because the ions are free to carry the charge.

### Electrolysis

#### 1.48

Electric current is the flow of charged particles – either ions or electrons.

#### 1.49

Covalent compounds are unable to conduct electricity because they have no free ions or electrons that can carry the charge.

#### 1.50

Ionic compounds do not conduct electricity when solid because the ions are set in a fixed structure and therefore there are no particles free to carry the charge. However, when they are molten or in solution, the ions are free to move and therefore they can conduct electricity.

#### 1.51

An electrolyte is a substance that contains ions in solution. You can test for whether or not a liquid is an electrolyte by putting two carbon electrodes into a beaker of the liquid, ensuring that they do not touch. These electrodes should be connected to a light bulb and a power source. The bulb will light up if the liquid is an electrolyte because the ions in the solution will carry the charge in the gap between the two electrodes.

#### 1.52

Electrolysis involves the formation of new substances when electricity is put through an ionic compound as it overcomes the electrostatic forces of attraction in the ionic substance.

#### 1.53

To investigate the electrolysis of molten salts, set up a solution of the salt, such as copper chloride, which is used in this example, or lead bromide. Inert electrons must be used so that they do not react with any of the substances and only play a role in the transfer of electrons.

Copper chloride solution will separate into copper (Cu) and chlorine (Cl2) whereas lead bromide will separate into lead (Pb) and Bromine (Br2)

#### 1.54

You can also do experiments to discover the results of aqueous solutions such as sodium chloride, copper sulphate and sulphuric acid. This is done by, as in the example before, using inert electrodes to transfer the electrons through the solution. The positively charged ion compound will form an atom at the cathode, whereas the negatively charged ions will form atoms at the anode. If the metal in solution is more reactive than hydrogen, then the hydrogen (H+ ions from water) will form the product, as the metal will bond with oxygen.

In these cases, these are the products:

• Sodium Chloride: H2 at the cathode and Cl2 at the anode.
• Copper Sulphate: Cu at the cathode and O2 at the anode.
• Sulphuric acid: H2 at the cathode and O2 at the anode.

The products can be tested for using common lab tests.

#### 1.55

Ionic equations involved only the ions involved in reduction and oxidation. The spectator ions are not recorded.

E.g. Al3+ + 3e- → Al

#### 1.56

One faraday is representative of one mole of electrons. (96,500 coulombs)

#### 1.57

A current of 0.2 Amps is passed through CuSO4 for 2 hours, how much copper will be produced?

1. Half equation Cu2+ + 2e- → Cu

2. Work out coulombs of electrons flowing Q= IT Time = 2x60x60 Q= 0.2 x 7200= 1440C

3. Convert C into moles of electrons Moles= C/Faraday Moles= 1440/96500 Moles= 0.015

4. Ratio Moles Cu2+ + 2e- → Cu For every 2 moles of electrons, there will be one Cu Ratio = 2:1

5. Work out moles of product using ratio Moles of electrons x ratio 0.015x0.5= 0.0075 Moles of Cu

6. Convert moles into mass Moles x Mr 0.0075 x 63.5= 0.48g of copper