Subject: Science Years : 9 -11 Topic: C1 and C2 – States ...

Subject: Science Years: 9-11 Topic: C1 and C2 – States of matter and separating mixtures Paper: 3

Lesson Sequence 1. States of Matter 2. Mixtures 3. Filtration and

Crystallisation 4. Paper

Chromatography 5. Distillation 6. CP – Investigating Inks 7. Drinking Water

Core Texts Edexcel Combined Science Text book pages 146-159

State Changes: Evaporation/ boiling: Liquid to gas Melting: Solid to liquid Sublimation: Solid to gas Condensing: Gas to liquid Freezing: Liquid to solid Deposition: Gas to solid

Identifying Pure Substances/ Mixtures We can identify a pure substance or a mixture by its melting point. If it has a specific melting point it is pure. If it has a melting range it is a mixture.

Separating Mixtures Crystallisation Solutes (dissolved substances) can be separated from a solution by evaporation. Example: Separating salt from sea water. 1. Evaporate the solvent until it reaches a point

where there is as much salt in the water that can dissolve (saturated).

2. Leave the water to evaporate, which will then form crystals.

Filtration This is separating an insoluble substance from a solution. Use a filter funnel lined with filter paper. The filter paper has tiny holes which won’t allow larger insoluble substances to fit through. This leaves the solid as a residue on the filter paper. Paper chromatography This is a process that works because some compounds dissolve better in a solvent than others. When the solvent moves along a strip of paper, it carries the different substances at different speeds, separating them. The paper is the stationary phase, the solvent is the mobile phase. Distillation A mixture is evaporated, and then the solvent is condensed in a condenser.

Rf = distance moved by sample ÷ distance moved by solvent

Drinking Water: Purifying Sea Water: Using simple distillation. Sea water is heated so that water vapour leaves. The water vapour is then condensed. Requires a lot of energy. Water for Chemical Analysis: Should not contain any dissolved salts otherwise incorrect results can be obtained. Tap water can contain small amounts of these salts. Water for Drinking: Most of this comes from aquifers in the UK. This needs to be dealt with through sieving (remove large object), sedimentation (small particles settle out), filtration through sand and gravel. Chlorine is added to kill microorganisms.

Key Words States of matter

The three forms that a substance can have (solid, liquid and gas).

Desalination Producing fresh drinking water through separating the water from salts.

Still Apparatus used to carry out distillation.

Melting point

The point in which solid turns into a liquid (or vice versa).

Filtration Using a filter to separate insoluble substances from a liquid.

Mixture

A substance containing two or more different substances that aren’t joined together.

Boiling point

The temperature when liquid turns into a gas (or vice versa).

Insoluble A substance that can’t be dissolved in a certain liquid.

Aquifers Underground layer of rock containing groundwater.

Pure A single substance that does not have anything else mixed with it.

Crystallisation Separating the solute from a solution through evaporation.

Sedimentation The process in which rock grains and insoluble substances sink to the bottom of a liquid.

Subject: Science Years: 9-11 Topic: C3 – Atomic Structure Paper: 3 and 4

Lesson Sequence 1. Structure of an atom 2. Atomic number and

mass number 3. Isotopes


Key Words Atoms Smallest neutral part of an

element that can take part in chemical reactions.

Relative mass

The mass of something relative to the mass of something else.

Relative atomic mass

The mean mass of an atom relative to the mass of an atom of C-12.

Element A simple substance made up of only one type of atom.

Atomic number

The number of protons in the nucleus of an atom.

Nucleus The central part of an atom or ion.

Subatomic particles

Particles that are smaller than an atom.

Mass number

Total number of protons and neutrons found in an atom

Electron shells

Areas around a nucleus that can be occupied by electrons.

Relative charge

The charge of something in relation to the charge of something else.

Isotopes Atoms on an element with the same number of protons but different number of neutrons

Nucleons Subatomic particles found in the nucleus.

Atomic Structure John Dalton’s theory All matter is made up of tiny particles called

atoms. Atoms are tiny, hard spheres that can’t be

broken down into smaller parts Atoms cannot be created or destroyed. Atoms of an element are all identical. How Dalton’s ideas have changed since then Electrons were discovered in a series of

investigations. This was the first sign of subatomic particles.

Further subatomic particles have since been discovered (protons and neutrons).

At the centre of all atoms is a nucleus, which is surrounded by fast moving electrons in electron shells.

Atoms always have equal numbers of protons and electrons as they have no overall charge.

The overall diameter of an atom is MUCH bigger than the nucleus of an atom (up to 100 000 times bigger.

Subatomic particle

Relative charge

Relative mass

Location

Proton +1 1 Nucleus Neutron 0 1 Nucleus Electron -1 1/1835 Electron

shell

Atomic Number and Mass Number Experiments carried out by Ernest Rutherford

suggested that most of an atom is empty space, with a small positive nucleus that contains most of the mass of the atom.

Atomic number (Symbol - Z) Atoms in the periodic table are organised by

atomic number. This is unique for each atom. The atomic number of an element is the same

as the number of protons in the nucleus. As atoms don’t have an overall charge, the

number of protons is the same as the number of electrons. Therefore, the atomic number is the same as the number of electrons in the shells (not true for ions)

Mass number (Symbol - A) The mass number of an atom is the same

as the total number of protons and neutrons an element has.

This is due to the mass of an electron being negligible.

Isotopes The discovery of neutrons explains

why some atoms of the same element have different masses.

We refer to specific isotopes by adding its mass number to the end of its name (e.g. Lithium-6 and Lithium-7)

The number of neutrons in an isotope can be calculated using the mass number and atomic number.

Number of

neutrons

=

Mass

number

-

Atomic number

Relative atomic mass (RAM) is the number found in the periodic table. This takes into account the relativity of the isotopes found for an element.

RAM = Total mass of the atoms Total number of atoms

Subject: Science Years: 9-11 Topic: C4 – The Periodic Table

Lesson Sequence 1. Elements and the

periodic table 2. Atomic number and

the periodic table 3. Electronic

configurations and the periodic table


Key Words Periodic table: The chart in which the elements are arranged in order of increasing atomic number.

Atomic number: The number of protons in the nucleus of an atom

Electronic configuration: The arrangement of electrons in shells around the nucleus of an atom

Chemical properties: How a substance reacts with other substances

Period: Horizontal rows in the periodic table. Electron shells: areas around a nucleus that can be occupied by electrons.

Physical properties: A description of how a material behaves and responds to force and energy.

Group: Vertical column in the periodic table. Relative atomic mass: The mean mass of an atom relative to the mass of an atom of C-12

Elements and the early periodic table Mendeleev’s Table Dmitri Mendeleev was a Russian

chemist who constructed an early version of the periodic table.

He arranged the elements in order of increasing relative atomic masses (unlike the modern day periodic table).

He left gaps in his tables for elements that hadn’t been discovered yet.

He also arranged the elements by their chemical properties and physical properties (which sometimes meant that elements were swapped around)

Vertical columns contained elements with increasing relative atomic masses, and horizontal rows contained elements with similar chemical properties.

Mendeleev used the gaps in his periodic table to predict the properties of yet to be discovered elements.

Atomic number and the periodic table The modern day periodic table is

arranged in order of increasing atomic number.

Explaining the shape of the periodic table Periods: these are the rows of the

periodic table. The elements in a period are in order of increasing atomic number.

Groups: The groups are the vertical columns. Elements with similar properties (both chemical and physical) are placed in the same group.

Non-metals are on the right of the table. Metals are the rest of the table.

Electronic configurations Electrons fill the electron shells, filling

the first shell first, then the second, and so on…

The electronic configuration of an atom is a diagram that represents the number of electrons on each shell, with the shells drawn as circles and the electrons as dots or crosses.

Each shell can contain different numbers of electrons. For the first 3 shells (the ones you need to know):

o The first shell can hold up to 2 electrons

o The second and third shells can hold up to 8 electrons.

Example electron configuration.

This can also be written as 2.1 The number of occupied shells is equal

to the period number of the element in the periodic table.

The number of electrons in the outer shell is equal to the group number of the element.

Subject: Science Years: 9-11 Topic: C5, C6 and C7 – Chemical bonding

Lesson Sequence 1. Ionic bonds 2. Ionic lattices 3. Properties of ionic

compounds 4. Covalent bonds 5. Molecular compounds 6. Allotropes of carbon 7. Properties of metals 8. Bonding models


Key Words Ions: Atoms with a charge Crystals: Solids that are made of regular

repeating structures Anode: Positive electrode

Malleable: A substance that can be hammered into shape.

Polyatomic ions: A group of atoms that have an overall charge

Molecular formula: The formula showing the actual number of atoms

Electrostatic forces: Force of attraction between oppositely charged particles.

Aqueous: Mixture formed when a substance is dissolved in water

Valency: The number of covalent bonds that an atom can form.

Lattice: An arrangement of many particles that are bonded together in a fixed, regular grid-like pattern

Cathode: Negative electrode Allotrope: Different structures of the same element.

Ionic bonds Ionic bonds take place between metals and

non-metals Atoms are more stable if they have a full

outer shell of electrons They can gain a full outer shell by either

gaining or losing electrons, forming ions. Cations: Positively charged ions. Formed if

an atom loses electrons. They have less electrons than protons.

Anions: Negatively charged ions. Formed if an atom gains electrons. They have more electrons than protons.

Metals form cations, non-metals form anions.

Electrostatic forces hold the oppositely charged ions together, forming ionic bonds.

When ionic bonds form, electrons are transferred from one atom (the metal) to another atom (the non-metal). E.g. when sodium chloride forms, one electron moves from the chlorine atom to the sodium atom.

Charges of groups 1, 2, 6 and 7 Group 1: lose 1 electron. Forms a 1+ cation. Group 2: loses 2 electrons. Forms a 2+

cation. Group 6: gains 2 electrons. Forms a 2-

anion. Group 7: gains 1 electron. Forms a 1- anion.

Ionic lattices Ionic compounds are held

together by strong electrostatic forces between oppositely charged ions.

These strong ionic bonds allow billions ions to be packed together in a repeating arrangement. This is a lattice.

Working out ionic formulae Ionic compounds are electrically

neutral, so must have the same number of positive charges as negative charges. You can work it out using the ion formulae below

Ion Formula Group 1 metal Group 2 metal Group 3 metal

Group 7 ion Group 6 ion Group 5 ion Ammonium

Nitrate Hydroxide Carbonate

Sulfate Sulfite

M+

M2+ M3+ X- X2- X3-

NH4+

NO3-

OH- CO2

2- SO4

2- SO3

2-

Ionic compound names If an ionic compound ends in –ide, it

contains an anion that only has one atom in it (e.g. magnesium oxide contains an O2- anion)

If an ionic compound ends in –ate, it’s anion contains at least two atoms, with one of them oxygen (e.g. sodium carbonate contains the anion CO3

2-) (hydroxide is the only exception)

Properties of ionic compounds High melting and boiling points: this is

due to strong electrostatic forces, requiring a large amount of energy to separate the ions.

Electrical conductivity: Ionic compounds can’t conduct electricity when solids, although they can when aqueous or molten. Electrical conductivity requires the movement of charge, which is only possible when the compound is not solid.

Covalent bonds Occur between non-metals. Formed through the sharing of

outer electrons. If atoms share 1 pair of electron

they form a single bond. If atoms share 2 pairs of

electrons, they form a double bond.

Covalent bonds continued There are strong electrostatic forces

between atoms in the molecule. There are weak intermolecular forces

between atoms in a molecule.

Dot and cross diagrams These are diagrams that are drawn to show

how bonds are formed. Covalent bonds You need to remember the below compounds: Compound Dot and cross diagram

H2

HCl

H2O

CH4

O2

CO2

Dot and cross diagrams Ionic compounds Below is an example of a dot and cross diagram for an ionic compound

Working out molecular formula You can use this information to work out the

ratio of atoms in a molecule. Group Outer

electrons Bonds formed

Valency

4 4 4 4 5 5 3 3 6 6 2 2 7 7 1 1

Allotropes of carbon Fullerenes: each carbon atom is bonded to three other carbon

atoms. Can be tubular (nanotubes) or spherical. Weak intermolecular forces between molecules, so have low melting points. They are soft and slippery.

Graphene: not a simple molecule. A sheet of carbon atoms with no fixed formula. One atom thick, so is very thin, but is also very strong. Electrical conductor as electrons are free to move.

Diamond: A giant structure. Each carbon atom is bonded to another 4 carbon atoms. Diamond is very hard due to a rigid network of carbon atoms. Diamond is an electrical insulator.

Graphite: A giant structure. Arrangements of sheets of carbons, where each carbon is strongly bonded to 3 other carbons. There are delocalised electrons that are free to move. Due to the delocalised electrons, graphite can conduct electricity. Graphite is used as a lubricant as the layers are able to slide over each other.

Molecular properties Low melting and boiling points due to weak

intermolecular forces between molecules. Electrical insulators due to lack of moving

charge. Polymers: Simple molecules (monomers) can

join to form a chain of molecules, called a polymer. These can be different lengths, and can lead to different properties.

Properties of metals All metals have similar properties:

o Solids with high melting points: due to strong electrostatic attractions between positive metal ions and negative delocalised electrons.

o Shiny (when polished) o Malleable: the sea of electrons hold the ions together so

the metal changes shape when hit, not breaking. o High density o Good conductors of electricity: the delocalised electrons

move when a potential difference is applied through the metal.

Metals are in the structure of a giant lattice. Metal atoms have one, two or three electrons in their outer shell, which are lost from each atom and become free to move randomly throughout the metal. The electrostatic attraction between the metal ion and electron is known as metallic bonding.

Problems with bonding models Dot and cross diagrams don’t show the structure formed. Suggest

electrons from different atoms are different to each other. Metallic model doesn’t show ions vibrating all the time. 3D ball and stick models show the atoms too far away from each

other, and bonds aren’t a “stick”

Subject: Science Years: 9-11 Topic: C8 – Acids and alkalis Lesson Sequence 1. Acids, alkalis and indicators 2. Looking at acids 3. Bases and salts 4. Preparing Copper sulfate 5. Alkalis and balancing equations 6. Investigating neutralisation 7. Alkalis and neutralisation 8. Reactions of acids 9. Solubility

Key Assessments Core Practical: Preparing copper sulfate Core Practical: Investigating neutralisation


Key Words Acid: A solution that has a pH of less than 7 Concentrated: Containing a large amount of

solute dissolved Titration: A technique that is used to find the exact volumes of solutions reacting with each other

Alkali: A solution that has a pH greater than 7 Dilute: A low concentration of a solute Spectator ions: Ions that don’t take part in a reaction

Neutral: A liquid that is neither acidic nor alkaline

Dissociate: Breaking up a compound into simple components

Precipitation: A reaction in which a precipitate is formed

Indicator: A substance that changes colour depending on the pH of a substance

Base: A substance that reacts with acid to form only water and a salt

Burette: Piece of apparatus used to measure the volume of a solution

Indicator Litmus Methyl orange

Phenolphthalein

Colour in alkali

Blue Yellow Pink

Colour in acid

Red Red Colourless

The effect of concentration of ions on pH:

Higher concentration of H+ ions: the more acidic the solution, the lower the pH Higher concentration of OH- ions: the more alkaline the solution, the higher the pH

Concentration (g dm-3) = Amount dissolved (g) Volume of solution (dm3)

Strong vs Weak Acids:

Strong acid: dissociate completely into ions when they dissolve in water high concentration of H+ ions Weak acid: do not dissociate completely into ions in a solution

State symbols: (s) – solid (l) – liquid (g) – gas (aq) – dissolved in water

Steps for preparing a soluble salt:

1) Add excess metal oxide to acid 2) Warm the mixture in a water

bath to speed up the reaction 3) Filter to remove the unreacted

solid from the solution 4) Heat to evaporate water slowly

and concentrate the salt solution 5) Leave to evaporate the water

slowly

General equations for reactions involving acids:

Metal + acid Salt + hydrogen Metal oxide + acid Salt + water Metal hydroxide + acid Salt + water Metal carbonate + acid salt + water + carbon dioxide

Compounds Soluble Insoluble Nitrates All compounds None Chlorides Most Silver and lead Sulfates Most Lead, barium and calcium Carbonates Sodium, potassium

and ammonium Most

Hydroxides Sodium, potassium and ammonium

Most

All sodium, potassium and ammonium salts are soluble

Steps for preparing an insoluble salt:

1) Mix the two solutions 2) Filter the mixture 3) Rinse the beaker with distilled

water and pour through the funnel

4) Pour distilled water through the precipitate in the funnel

5) Carefully remove the filter paper and dry it in a warm oven

Subject: Science Years: 9-11 Topic: C9 – Calculations involving masses

Lesson Sequence 1. Masses and empirical

formulae 2. Conservation of mass 3. Moles

Key Assessments Core Practical: N/A


Key Words Empirical formula: The formula showing the simplest whole number ratio of each atom

Concentration: The amount of a solute dissolved in a certain volume of solvent.

Limiting reactant: The reactant that determines the amount of product in a chemical reaction.

Molecular formula: The formula showing the actual number of atoms of each element.

Closed system: When substances can’t enter or leave an observed environment.

Mole: A mole of something is 6 x 1023 of it. 1 mole is equal to the Ar in g.

Relative atomic mass: The mass of something compared to the mass of something else.

Non-enclosed system: Where substances can enter or leave an observed environment.

Stoichiometry: The molar ratio of reactants and products in a chemical reaction.

Conservation of mass: The total mass of products is the same as the total mass of the reactants

Avogadro constant: The number of particles in one mole of a substance (6.02 x 1023)

Calculating the empirical formula (from molecular formula) Simplify it in to the smallest whole number ratio. e.g. C4H10 would become C2H5, because 4:10 can be simplified to 2:5

Calculating relative formula mass (Mr) This is the sum of the relative atomic masses (Ar) of all the atoms or ions in its formula. e.g. Calculate the Mr of carbon dioxide (CO2) = Ar(C) + (2 x Ar(O)) = 12 + (2 x 16) = 44 Therefore, the Mr of CO2 is 44

Calculating empirical formula (from data) 1. Write the question in ratio format. 2. Divide each mass by the Ar of each element 3. Divide both sides by the smallest answer from

step 2 4. Write the empirical formula using the

numbers from step 3. Using the example of ethane, which contains 80g of carbon and 20g of hydrogen.

C : H 80g :20g

80÷12 = 6.67 : 20÷1 = 20 6.67÷6.67=1 : 20÷6.67 = 3

CH3

Finding empirical formula experimentally The example you need to know about: Magnesium oxide 1. Measure the mass of a crucible with the lid on. 2. Measure the mass of a piece of magnesium ribbon

and place into the crucible. 3. Heat the crucible over a Bunsen burner, lifting the

lid slightly occasionally to increase the oxygen supplied to the metal.

4. Measure the mass of the crucible and lid, with the magnesium oxide in it.

5. Work out the mass of the magnesium oxide by subtracting the mass from step 1 from the mass from step 4. This can be used to work out the mass of oxygen that has reacted.

Concentration (gdm-3) = mass of solute (g) ÷ volume of solution (dm3)

Calculating reacting masses 1. Write a balanced

equation 2. Underline the subject

(The chemical you have most information about) and the target (the chemical you are working out)

3. Find the Mr of these 4. Write the number of

each present 5. Multiply the Mr by the

number present. 6. Divide the target by the

subject to get a ratio. 7. Multiply the mass of the

subject by the ratio.

Number of moles = Mass of substance (g) ÷ Ar/ Mr

Calculations involving moles Working out the limiting reactant: calculate the number of moles of each reactant. Whichever there is the least of is the limiting reactant. Working out the mass of something produced: Multiply the moles by the Mr/ Ar of the product. Stoichiometry: Calculate the number of moles of each reactant. Place these into the simplest ratio.

Calculating the molecular formula from empirical formula: Divide the empirical formula by the mass given in the question. Multiply the empirical formula by this number.

Subject: Science Years: 9-11 Topic: C10 – Electrolysis, C11 – Metals, C12 – Reversible Reactions

Lesson Sequence 1. Electrolysis 2. CP – Electrolysis of copper

sulfate solution 3. Products from electrolysis 4. Reactivity 5. Ores 6. Oxidation and reduction 7. Life cycle assessment 8. Dynamic Equilibrium

Key Assessments Core Practical: Electrolysis of copper sulfate solution


Key Words Electrolysis: Using electricity to split up a compound.

Native state: When an element is not combined with other elements in a compound.

Closed system: Substances can’t enter or leave an observed environment

Electrolyte: A liquid that is electrolysed. Ore: A rock that contains enough metal compound to be worthwhile extracting.

Open system: A system into or from which substances can enter or leave.

Reactivity Series: A list of metals in order of reactivity, with the most reactive at the top.

Redox: A reaction where both oxidation and reduction occur

Corrosion: The reaction between metals and oxygen.

Displacement Reaction: When a more reactive element takes the place of a less reactive element in a compound.

Dynamic Equilibrium: When the forwards and backwards reactions in a reversible chemical reaction are occurring at the same rate.

Rusting: The corrosion of iron to form iron oxide. (ONLY APPLIES TO IRON)

Electrolysis Cathode: Negative electrode. Cations

are attracted here. Anode: Positive electrode. Anions are

attracted here. Direct current is needed for electrolysis. OILRIG: Oxidation Is Loss (of electrons),

Reduction Is Gain (of electrons) Transfer of electrons changes charged

ions into atoms or molecules. This causes a chemical change.

ReduCtion takes place at the Cathode. o ½ equation: M+

(aq) + e- M(s) o E.g. Zn2+ + 2e- Zn

OxidAtion takes place at the Anode. o ½ equation: 2X-

(aq) X2(g) + 2e- o E.g. 2Cl-(aq) Cl2(g) + 2e-

Core Practical: Electrolysis of copper sulfate solution Using copper electrodes: Cathode: decreases in size. Impurities

are collected underneath. Anode: Pure copper forms here. Using graphite (inert) electrodes: Cathode: bubbles are seen. CO2 is

formed as oxygen reacts with carbon. Anode: Copper is formed. The anode

turns a bronze colour.

Products from electrolysis: General Electrolysis of molten salt: Decomposes into its elements. Metal is

produced at the cathode, non-metal at the anode.

Electrolysis of salt solutions: The ions of salt are present, along with

the ions of water (H+ and OH-) Which substances are given off is

decided by how readily the ions are discharged.

Cations order of discharge (form most likely to least likely): Cu2+, H+, Na+

Anions order of discharge (from most likely to least likely): Cl-, OH-, SO4

2- Products from electrolysis: Examples you need to know Copper sulfate solution: hydrogen at the

cathode, carbon dioxide at the anode. Sodium chloride solution: Hydrogen at

the cathode, chlorine at the anode. Sodium sulfate solution: hydrogen at

the cathode, oxygen at the anode. Water acidified with sulfuric acid:

hydrogen at the cathode, oxygen at the anode.

Molten lead bromide: Lead at the cathode, bromine at the anode.

Reactivity of metals Metal Reactivity

Potassium (K) Sodium (Na) Calcium (Ca)

Magnesium (Mg) Aluminium (Al)

Zinc (Zn) Iron (Fe)

Copper (Cu) Silver (Ag) Gold (Au)

Most reactive

Least reactive The more reactive a metal is, the

more likely it is to form cations. Reactions with water K, Na and Ca all react with cold

water to form hydrogen and a metal hydroxide.

Mg, Al, Zn and Fe react very slowly with cold water, but react with steam to form hydrogen and a metal oxide.

Cu, Ag and Au don’t react with water at all.

Reactions with dilute acid K and Na react violently. Ca, Mg, Al, Zn and Fe react to form

hydrogen and a salt. Cu, Ag and Au don’t react.

Reactions of metals: symbol equations Reactions with cold water 2K(s) + 2H2O(l) 2KOH(aq) + H2(g) 2Na(s) + 2H2O(l) 2NaOH(aq) + H2(g) Ca(s) + 2H2O(l) Ca(OH)2(aq) + H2(g)

Reactions with steam 2Mg(s) + H2O(g) 2MgO(s) + H2(g) 2Al(s) + 3H2O(g) Al2O3(s) + 3H2(g)

2Zn(s) + H2O(g) 2ZnO(s) + H2(g) 2Fe + H O 2FeO + H Displacement reactions We can use the reactivity series to

predict whether reactions will take place.

A metal will react with compounds of metals below it in the reactivity series.

These reactions are examples of displacement reactions.

Displacement reactions are examples of redox reactions.

The more reactive metal loses electrons during a displacement reaction – oxidation.

The less reactive metal (the metal in the compound) gains electrons - reduction.

Half Equations You need write half equations for

displacement reactions. Spectator ions: the non-metal ions.

They are spectator ions as they don’t change.

When a metal goes from an atom to a compound, it loses electrons:

o E.g. Zn(s) Zn2+(aq) + 2e-

When a metal goes from a compound to an atom, it gains electrons:

o E.g. Cu2+(aq) + 2e- Cu(s)

Extracting metals Some metals are very unreactive, which means they are

found naturally as their native state. Most metals are found in rocks called ores. These contain

metals in a compound (this tends to be a metal oxide). Extracting metals with carbon:

o This is for any metal less reactive than carbon. o The ore is heated with carbon. o Carbon displaces the metal from its oxide,

forming the metal and carbon dioxide. Extracting metals through electrolysis of molten ores:

o Used for metals that are more reactive than carbon.

o Splits the metal oxide into the metal and oxygen. o Much more expensive than heating with carbon.

Bioleaching: o Bacteria is grown on a low grade ore (doesn’t

contain a lot of metal ions) o Bacteria produces a solution containing metal

ions (called a leachate) o Metal is extracted from the leachate through

displacement using scrap iron, then purified by electrolysis.

o Used for copper, nickel, zinc and cobalt. o Advantages: no harmful gases produced, less

damage caused to the landscape, conserves supplies of higher grade ores, doesn’t require high temperatures

o Disadvantages: slow, toxic substances can be produced (which damages the environment)

Phytoextraction: o Growing plants that absorb metal compounds.

These are burnt to form ash, from which the metal is extracted.

o Advantages the same as bioleaching, except that it does need high temperatures, and that it can extract metals from contaminated soils.

o Disadvantages: slow, more expensive than mining, dependent on plant growth.

Oxidation and reduction Heating with carbon: carbon is oxidised, metal is reduced Electrolysis: Metal is reduced, oxygen is oxidised

o ½ equation at cathode: M+ + e- M o ½ equation at anode: 2O2- O2 + 4e-

Recycling metals Advantages: Natural reserves of metal ores last longer Need to mine ores is reduced Less pollution may be produced. Many metals need less energy to recycle than to extract. Less waste metal ends up in landfill sites. Disadvantages: Costs/ energy of collecting, transporting and sorting

Dynamic Equilibrium Some reactions are reversible. In a closed system, these

reactions will reach a dynamic equilibrium.

Affecting the position of the equilibrium Changing temperature:

o ↑: shifts to the endothermic direction o ↓: Shi s to the exothermic direc on.

Changing gas pressure: o ↑: shi s to the direc on forming fewer gas

molecules o ↓: Shi s to the direc on forming more gas

molecules Changing concentration:

o ↑: shi s in the direc on to use up the substance that’s been added

o ↓: Shi s in the direc on to form more of the substance that’s been removed

Subject: Science Years: 9-11 Topic: C13 – Groups in the Periodic Table

Lesson Sequence 1. Group 1 2. Group 7 3. Halogen Reactivity 4. Group 0

Key Assessments N/A


Key Words Reduction: A reaction in which oxygen is lost from a chemical or electrons are gained by an atom

Displacement Reactions: When a more reactive element displaces a less reactive element from one of its compounds

Oxidation: A reaction in which oxygen is added to a chemical or where electrons are lost from an atom

Diatomic: Two atoms chemically bonded. Group: Vertical column in the periodic table. Noble Gases: An unreactive gas in group 0 of the periodic table

Disinfectants: Something that destroys or neutralises disease-carrying microorganisms

Inert: Does not react Alkali Metals: A group of very reactive metals found in group 1 of the periodic table

Bleaches: A substance that takes the colour out of something

Redox: A reaction where both oxidation and reduction take place

Halogens: An element in group 7 of the periodic table

Group 1 – The Alkali Metals Physical properties: malleable, conduct

electricity, relatively low melting points, soft and easy to cut.

Reactions with oxygen: o Metal + Oxygen Metal oxide o 4M (s) + O2 (g) M2O (s)

Reactions with water: o Metal + Water

Metal Hydroxide + Hydrogen 2M(s) + 2H2O(l) 2LiOH(aq) + H2 (g)

Gets more reactive as you go down the group. Lithium + water Bubbles fiercely

on the surface Sodium + water Melts into ball

and fizzes about the surface

Potassium + water

Bursts into flame and flies about the surface

Explaining the reactivity: When they react, their outer electron

is transferred to a non-metal. As we go down the group, the atoms

get larger as an extra electron shell is added. The force of attraction between the nucleus and the electron becomes weaker as they get further apart.

The effect of concentration of ions on pH:

Higher concentration of H+ ions: the more acidic the solution, the lower the pH Higher concentration of OH- ions: the more alkaline the solution, the higher the pH

Group 7 – The Halogens All exist as diatomic molecules. Are toxic and corrosive. Density increases down the group. Halogen Appearance Chlorine Green gas Bromine Brown liquid Iodine Purple/ black solid Reactions with metals:

Halogen + metal metal halide X2(g) + M(s) MX2(s)

Used as disinfectants and bleaches. Reactions with hydrogen: Hydrogen + Halogen Hydrogen halide

o H2(g) + X2(g) 2HX(g) When hydrogen halides are dissolved

in water they form acidic solutions. They get more reactive as you go up

the group. Order of reactivity can be shown

through displacement reactions. Group 7 atoms gain one electron when

they react. The closer the electron is to the nucleus, the easier it is to gain the electron.

As you go down the group, the atoms get larger. Therefore the electrons become further away from the nucleus.

Group 0 – The Noble Gases Properties: are colourless, have very

low melting and boiling points, are poor conductors of heat and electricity.

They are inert. They exist as single atoms, because

they do not form bonds easily with other atoms.

Full outer shell of electrons. This is why they are unreactive.

Krypton: used in photography, produces a brilliant white light when electricity is passed through it.

Argon: Denser than air. Added to the space above the wine in wine barrels to stop oxygen reacting with the wine.

Helium: Low density and non-flammable. Used in weather balloons and airships.

Neon: Produces a distinctive red-orange light when electricity is passed through it. Used in making long lasting illuminated signs.

Subject: Science Years: 9-11 Topic: C14 and C15 – Chemical reactions

Lesson Sequence 1. Rates of reaction 2. Factors affecting

reaction rates 3. CP – Investigating

reaction rates 4. Catalysts and

activation energy 5. Exothermic and

endothermic reactions 6. Energy changes in

reactions

Key Assessments Core Practical: Investigating Reaction rates


Key Words Rate: How quickly something happens. Catalysts: A substance that speeds up the

reaction without itself being used up. Works by reducing the activation energy.

Displacement: When a more reactive element displaces a less reactive element from one of its compounds

Variables: Factors that can change. Endothermic: When energy is transferred from the surroundings to the products.

Bond energy: Energy needed to break one mole of a specific covalent bond.

Activation Energy: The minimum about of energy needed to start a reaction.

Reaction Profiles: A diagram that shows how energy changes during a chemical reaction.

Enzymes: A biological catalyst.

Exothermic: When energy is transferred to the surroundings from the reactants.

Neutralisation: When an acid reacts with a base to produce salt and water only.

Active Site: The space in an enzyme where the substrate fits during a reaction.

Measuring Reaction Rates This can be done by measuring how the amount of reactants or products changes with time. Measuring the rate of reactions that give off a gas 1. You can measure the volume of gas being

given off during the reaction. This can be done using either a gas syringe or a measuring cylinder.

2. You can measure the rate at which there is a change in mass. Carry out the experiment in a flask placed on a mass balance.

Measuring the rate of reaction when there is a precipitate forming 1. Place the mixture on top of a cross that has

been drawn on a piece of paper, and time how long it takes the cross to disappear.

What needs to happen for a chemical reaction to take place? For a reaction to happen reactant particles need to collide with enough energy (the activation energy) to react.

Factors Affecting Reaction Rates Concentration Change: Increasing the concentration of solutions increases the rate of reaction Explanation: There are more reacting particles in the same volume so collisions are more frequent Surface Area Change: Increasing the surface area to volume ratio increases the reaction rate Explanation: There is more surface for collisions to occur on, so they occur more frequently. Pressure of gases Change: Increasing the pressure of gases increases the rate of reaction. Explanation: The reactant particles are squeezed closer so collisions occur more frequently. Temperature Change: Increasing the temperature increases the rate of reaction. Explanation: The reactant particles speed up and have more energy. They collide more often and more particles have enough energy to react when they collide.

Reaction profiles Catalysts

Exothermic

Endothermic

Exothermic and Endothermic Reactions that are always exothermic: neutralisation, combustion and displacement Determining whether a reaction is exothermic or endothermic: Record the temperature before and after a reaction has taken place. A temperature increase means the reaction is exothermic; a temperature decrease means the reaction is endothermic. Breaking and making bonds: Energy is transferred to the reactants to break the bonds, this is endothermic. Energy is transferred to the surroundings, this is exothermic.

Calculating bond energy 1. Calculate the total energy in (bonds broken) 2. Calculate the total energy out (bonds made) 3. Energy change = energy in – energy out Negative value = exothermic Positive value = endothermic

Subject: Science Years: 9-11 Topic: C16 – Fuels

Lesson Sequence 1. Hydrocarbons in crude oil

and natural gas 2. Fractional distillation of

crude oil 3. The alkane homologous

series 4. Complete and incomplete

combustion 5. Combustible fuels and

pollution 6. Breaking down

hydrocarbons


Key Words Natural gas: Fossil fuel formed from the remains of microscopic dead sea organisms.

Viscosity: How thick or runny a liquid is. Impurities: Unwanted substances found mixed in useful substances.

Crude oil: Mixture of hydrocarbons formed from dead microscopic organisms.

Ignite: To start burning Weathering: The breaking down of rocks in chemical, physical and biological methods.

Hydrocarbon: A compound containing hydrogen and carbon only.

Feedstock: Raw material. Pollutants: A substance that harms living organisms when released in the environment.

Homologous series: A family of compounds that have the same general formula and properties

Combustion: A reaction between a fuel and oxygen which produces a large amount of energy.

Cracking: A reaction where large alkane molecules are split into two or more smaller alkanes and alkenes.

Crude oil and natural gas Crude oil and natural gas are both natural

resources formed from the ancient remains of microscopic animals and plants that once lived in the sea.

These are finite resources, because they are not made any more.

Crude oil is a mixture of hydrocarbons. They are present as both rings and chains. The number of carbons can vary.

Crude oil is very important in the manufacturing of:

1. Fuels for vehicles, aircraft, ships, heating and power stations

2. Feedstock for the petrochemical industry

The fuels we get from crude oil and natural gas are non-renewable.

Fractional distillation To make crude oil useful, it needs to be separated into simpler, more useful

mixtures. This separation is done using fractional distillation. This uses the different

fractions (groups of hydrocarbons) boiling points. Steps for fractional distillation of crude oil:

1. Heat the crude oil strongly to evaporate it. Pipe the hot vapours into a fractionating column.

2. The column is hottest at the bottom and coldest at the top 3. The vapours rise through the column and cool down 4. The vapours condense when they reach a part of the column that is cool

enough 5. The liquid falls into a tray and is piped away 6. The vapours with the lowest boiling points do not condense at all and

leave at the top as mixtures of gas. 7. Bitumen has the highest boiling point and leaves at the bottom as a hot

liquid.

Properties and uses of the fractions from crude oil Fraction Number

of atoms

Boiling point

Ease of ignition

Viscosity Uses

Gases Smallest

Largest

Lowest

Highest

Easy

Hard

Lowest

Highest

Domestic cooking/ heating Petrol Fuel for cars Kerosene Fuel for aircraft Diesel oil Fuel for some cars and trains Fuel oil Fuel for ships and power stations Bitumen Surfacing roads and roofs

Alkanes formulae and structure Name Molecular

formula Structural formula

Methane CH4

Ethane C2H6

Propane C3H8

Alkanes The formulae of

neighbouring compounds differs by CH2

General formula: CnH2n+2 Gradual variation in physical

properties Similar chemical properties Saturated

Complete combustion During complete combustion of a

hydrocarbon: o Only carbon dioxide and water are

produced o Energy is given out

General word equation for combustion of a hydrocarbon: Fuel + oxygen carbon dioxide + water

Below is a diagram that shows apparatus that can be used to investigate the products of combustion. The pump draws the products towards it. The ice cools and condenses the water vapour passing through the tube. White anhydrous copper sulfate turns blue, showing the presence of water. The limewater turns milky, showing the presence of carbon dioxide.

Incomplete combustion Incomplete combustion occurs when there is a

limited supply of air, so there isn’t enough oxygen for the fuel to react with.

When a hydrocarbon undergoes incomplete combustion;

o water is produced o energy is given out (less than complete

combustion) o carbon monoxide (CO) and carbon are

produced The carbon atoms are released as smoke or soot. Problems with incomplete combustion Carbon monoxide is a toxic gas. This is because it

binds with haemoglobin in the red blood cells, preventing oxygen combining.

Symptoms of carbon monoxide poisoning include sleepiness, unconsciousness and in extreme cases death.

Soot can block pipe carrying waste gases away from an appliance. It blackens buildings, and can also cause breathing problems if it collects in the lungs.

Acid rain Rain is naturally acidic, due to carbon dioxide from the air

dissolving in it. Acid rain has a pH lower than 5.2. This is caused mainly by sulfur

dioxide. Hydrocarbons may contain sulfur compounds, which occur

naturally as impurities. Most are removed at oil refineries. When the hydrocarbons are burned, the sulfur reacts with

oxygen, producing sulfur dioxide gas. Sulfur dioxide dissolves in the water in the clouds to form a

mixture of acids. When sulfur dioxide dissolves it forms sulphurous oxide, which can then oxidise with oxygen in the air to form sulfuric acid.

Problems Crops: Acid rain can cause the pH of soil to decrease. This can

cause crops to not grow as well. Lakes and rivers: Excess acidity can prevent fish eggs from

hatching, and it can kill fish and insects. Weathering: Acid rain can increase the rate of weathering of

buildings made from limestone or marble, as sulfuric acid reacts with calcium carbonate.

Corrosion: Acid rain increases the rate of corrosion of metals, such as the iron in steal.

Oxides of nitrogen Car engines can be hot enough for nitrogen

and oxygen in the air inside the engine to react together. This produces oxides of nitrogen (NOx), which are pollutants.

NOx can cause acid rain. Nitrogen dioxide can cause respiratory diseases, such as bronchitis.

Breaking down hydrocarbons The volume of each fraction from crude oil

doesn’t usually match with the demand of the fractions.

There tends to be more long hydrocarbons, which are less useful, compared to short hydrocarbons, which are more useful.

Cracking is used to get around this problem. Cracking involves breaking down larger

hydrocarbons into smaller hydrocarbons, including an alkene (a compound with a C=C double bond) as well as an alkane.

Crude oil fractions are heated to evaporate them. They are then passed over a catalyst containing aluminium oxide and heated to 650oC.

Alkanes and alkenes Alkanes and alkenes

are different homologous series.

They are both hydrocarbons.

Alkanes are saturated (all the carbon atoms are joined by single bonds, C-C)

Alkenes are unsaturated (they contain a carbon-carbon double bond, C=C).

Fuels for cars Petrol is in high demand for use as a

fuel. This is because: o It’s a liquid at room

temperature, so easy to store. o It is easily ignited. o Its combustion releases a lot of

energy. Hydrogen can also be used as a fuel. It is

often a by-product of cracking. o Its combustion doesn’t produce

carbon dioxide, providing environmental benefits.

o Easily ignited too, as well as producing a lot of energy.

o Gas at room temperature, so difficult to store.

Subject: Science Years: 9-11 Topic: C17 – The atmosphere

Lesson Sequence 1. The early atmosphere 2. The changing atmosphere 3. The atmosphere today 4. Climate change


Key Words Composition: The way in which a mixture is made up.

Photosynthesis: A series of enzyme-catalysed reactions carried out in the green parts of plants.

Greenhouse gases: A gas that helps trap heat in the atmosphere. Carbon dioxide, methane and water vapour are the main examples.

Atmosphere: The gases in the environment. Infrared: Electromagnetic radiation that we feel as heat.

Greenhouse effect: The warming effect caused by greenhouse gases

Volcanic activity: The activity of a volcano. Emits: Gives out Correlation: Relationship between two variables.

Hypotheses: A proposed explanation for something.

Absorb: Soak up or take in Causal link: When something causes something else to happen

The early atmosphere Composition The early atmosphere was composed

of: o Mainly carbon dioxide o Small amounts of water vapour

and other gases. o Little or no oxygen

There was a lot of volcanic activity, which produced most of these gases.

Evidence Looking at what gases are released from

volcanoes in the present day. Looking at the atmosphere of similar

planets (Venus and Mars), which are mainly made of carbon dioxide.

Evidence for the lack of oxygen: o Oxygen is not produced by

volcanoes o Iron pyrite (which is broken

down by oxygen) is found in ancient rocks.

The changing atmosphere Oceans Around 4 billion years ago the

oceans began to form. This happened when the Earth

started to cool down. The water vapour in the atmosphere condensed and fell to Earth, forming the oceans.

Changes caused by this: o Carbon dioxide dissolved

in the oceans, reducing the levels of carbon dioxide in the atmosphere.

o Sea creatures dissolved the carbon dioxide to form shells, allowing the ocean to dissolve more carbon dioxide

Photosynthesis Photosynthetic organisms

(bacteria and primitive plants) started to grow and photosynthesis.

Changes caused by this: o Oxygen levels in the

atmosphere increase. o Carbon dioxide levels in

the atmosphere decrease.

Testing for oxygen Light a splint, and blow it out. This

should leave a glowing splint. When placed in a container with pure

oxygen, the splint will relight.

The greenhouse effect Energy from the Sun is transferred to the

Earth through waves. Energy gets absorbed by the Earth’s

surface. The warm Earth emits waves. Gases in the

air can absorb energy from these waves. The gases re-emit the energy, with some of

it returning to the Earth’s surface and warming it.

Climate change There is evidence that supports the idea

that human activity is increasing the greenhouse effect.

This is thought to be causing climate change.

Carbon dioxide levels have been increasing since 1850. This ties in with when fossil fuels were started to be burned.

There is a strong correlation between CO2 levels and surface temperature. This doesn’t mean that there is a causal link.

Earlier measurements were not as accurate as measurements today, as modern thermometers have a greater resolution.

Reasons for extra CO2 in atmosphere: Burning fossils fuels, release of methane from extracting oil and gas, and from cattle and paddy fields

Climate change Effects of climate change - Rising average temperatures, causing ice poles to melt. This will raise sea levels, causing flooding. - Animals may become extinct if habitats are destroyed. - More extreme weather conditions - Increasing acidity of the oceans, harming organisms and bleaching coral. Limiting the impact - Increase the use of renewable energy resources. - Capture CO2 from the air and bury it underground. - Adapting to new conditions. - Reflect sunlight back into space.

Subject: Science Years : 9 -11 Topic: C1 and C2 – States ...

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