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CHEMISTRY REVISION
CC1A STATES OF MATTER
PARTICLE MODEL – explains state changes in a substance, in terms of arrangement, movement
and energy stored in its particles

State changes –




Physical changes – can be reversed
Chemical properties of substances don’t change
Only rearrangement, movement and stored energy changed





Particles are attracted to one another by weak forces of attractions/ covalent bonds
Bonds breaking – energy needed from surroundings (endothermic)
Bonds forming – energy released into surroundings (exothermic)

HEATING CURVE

ATTRACTIVE FORCES – the weak forces of attraction between two molecules
MELTING POINT - temperature at which a substance changes from the solid state to the liquid
state when heated or from the liquid to solid state when cooled
BOIILING POINT – temperature at which the liquid boils, bonds being broken


State of a substance can be predicted using its melting/boiling point and temperature

CC2A MIXTURES –
PURE – a single substance with a fixed composition, that doesn’t have anything else mixed
within it (pure gold) – can be atoms, compounds or elements
MIXTURE – two or more substances jumbled together but not chemically bonded
...
0

Uses of chromatography –




To distinguish between pure and impure substances
To identify substances by comparing patter on chromatogram with other known
substances
Identify using Rf values

CC2D DISTILLATION
DISTILLATION – the separation of liquid components from one another using evaporation,
condensation, and collection
Simple distillation 




Simple distillation used to separate one liquid component from another
Not entirely successful as some gas can get lost/not condensed
To reduce risk of over boiling, anti-bumping particles are used
Liquid evaporates, other components left behind, vapour condensed and collected

Fractional distillation 





Used to separate a substance of two or more components/gases
Original mixture separates into various fractions/parts
Fraction with the lowest boiling point collected first
as column is heated, temperature gradient established
Hottest at bottom, coolest at top as fraction with the lowest boiling point will rise and
condense at the top
- Used to separate crude oil
- Produce alcohol products (whiskey/vodka)
- To separate gases (air) – cooled and turned into liquid at -200 degrees

simple distillation
fractional distillation

CC2d CORE PRACTICAL – INVESTIGATING INKS
1
...
METHOD (chromatography using ink from distillation) 






Draw pencil line on chromatogram (2cm from bottom)
Add small spot of ink to pencil line
Add water to container (depth of 1cm)
Place paper into container and support paper with paper clip
All water to travel through paper
Measure distance and collect Rf value for each coloured substance

APPARATUS –






Beaker
Solvent (water)
Chromatogram
Ink
Paper clip

PRECAUTIONS –





Use anti-bump granules to prevent the ink over-boiling
Use Bunsen burner appropriately – clear area around it
Draw start line in pencil so it doesn’t move from solvent
Place solvent beneath start line so process is fair

CC2e DRINKING WATER
DESALINATION – producing clean and pure water by the process of distillation


Sea water is heated to evaporate the water, leaving behind the salt
...


CHEMICAL ANALYSIS – chemical reactions or machinery to identify and measure the
substances within it
WATER FOR DRINKING –




Raw material found in rivers, lakes and aquifers (rocks with water underneath)
Often stored in reservoirs
Can contain leaves, twigs, insoluble particles including grit, soluble substances including
salts and pesticides, as well as harmful bacteria and micro-organisms

PROCESS TO REMOVE IMPURITES FROM FRESH WATER –
1
...

3
...


Screening using a sieve
Sedimentation (letting the different components rest and large insoluble particles gone)
Filtration (small insoluble particles removed)
Chlorination (chlorine used to kill micro-organism and soluble particles)

CC3A STRUCTURE OF AN ATOM








Atoms cant be created nor destroyed
Consists of subatomic particles
The dense nucleus of an atom consists of neutrons and protons
Electrons orbit the nucleus
The space in-between consists of nothing
Majority of atom therefore consists of nothing
Atoms within an element are identical, however differ between each element

SUBATOMIC
PARTICLE
Proton
Neutron
Electron

RELATIVE MASS
1
1
Negligible (nothing)

RELATIVE CHARGE
+1
0
-1

CHARGE
Positive
Neutral
Negative

CC3B – ATOMIC AND MASS NUMBER
Mass number/ atomic mass = number of protons +
number of neutrons
Number of protons = number of electrons
Atomic number = number of protons/electrons
Number of neutrons = mass number − atomic number
(BIGGER NUMBER WILL ALWAYS BE THE ATOMIC MASS)




Initially placed in order of masses in periodic table
Rearranged so those of similar properties were together
Modern periodic table – ordered by atomic number

CC3C - ISOTOPES
ISOTOPES – different forms of the same atom, number of protons remains the same however
atomic mass changes due to change in neutrons
RELATIVE ATOMIC MASS/ RAM – mass of an element in correspondence to another


Nucleus can be split using nuclear fusion (firing neutrons at an atom) and produces
large sums of energy
ABUNDANCE = RAM x 100/ mass
RAM = total mass of atoms/ number of atoms

RAM = (Abundance of Isotope 1 X mass of Isotope 1) +(Abundance of Isotope 2 X mass of
Isotope 2) / 100
CC4A – ELEMENTS AND PERIOD TABLE OF ELEMENTS




Mendeleev arranged elements in order of increasing RAM
Left gaps for elements that were yet to be discovered
Swapped elements to better suit their chemical properties – placed iodine after
tellurium so it lined up with fluorine, chlorine and bromine
- Vertical columns – groups (INCREASE IN ATOMIC MASS AS YOU GO DOWN)
- Horizontal columns – periods (INCREASE IN ATOMIC NUMBER AS YOU GO
ACROSS – SIMILAR PROPERTIES

Group 1 / 2 = alkali metals (LEFT)
Group 2-3 = transitional metals
Group 4-7 = non-metals
Group 7 = halogens
Group 0 = noble gases (RIGHT)
CC4B - ATOMIC NUMBER AND THE PERIODIC TABLE




End of 19th century, noble gases discovered
Hadn’t predicted existence due to being inert
Pair reversals (iodine and tellurium) not explained until 1913, Moseley

Atomic number –







Moseley showed an elements position was based on its physical properties
Fired high-energy electrons at elements, making them give off x-rays
Discovered for every step increase of atomic number, change in energy of x-rays
Realised atomic number was equal to number of positive charges in nucleus
Proton then discovered shortly after
Therefore proved atomic number to be equal to number of protons in a nucleus

CC4C - ELECTRON CONFIGURATION



Electrons occupy electron shells, arranged around the nucleus
Arrangement is known as electric configuration

For first 20 elements



1st shell = 2 electrons
2cd shell = 8 electrons
3rd shell = 8 electrons




Vertical column (group) indicates
how many electrons are in outer shell
Horizontal column (period)
indicates how many shells there are
Can be calculated using atomic number of an element – fill shells

ATOM - what makes up an element, consists of protons, neutrons, and electrons on the outer
shells
ELEMENT –substance made out of the same atoms (o2)
COMPOUND – substance made out of different atoms that are chemically bonded (co2, H2O)
MIXTURE- substance consisting of different atoms that aren’t chemically bonded (Salt water)

CC5A – IONIC BONDING
ION – atom with an unequal amount of protons and electrons, establishing a charge









Atoms are usually neutral, however this is can sometimes be altered due to chemical
reactions
Loss and gain of protons and electrons are attempts to complete outer shell
Electrostatic forces of attraction between negative and positive atoms
Hold oppositely charged ions together and form ionic bond
LESS ELECTRONS THAN PROTONS = POSTIVE

Cation

MORE ELECTRONS THAN PROTONS = NEGATIVE

Anion

Atoms with nearly full or nearly empty outer-shells will form into ions the quickest
Most ionic bonds occur between a metal and a non-metal
Ion formed is dependent on elements position in period table and number of outer
electrons

CC5B - IONIC LATICIES
IONIC BONDING – when an atom gains from or loses electrons to another atom during a
chemical reaction

(If more atoms removed than need to be added, you need to use another atom from the product)
Group on
periodic
table
Change
formed

1

2

3

4

5

6

7

0

1+

2+

3+

4-/4+

3-

2-

1-

N/A

LATTICE STRUCUTRE – arrangement of many particles that are bonded together in a fixed,
regular, grid-like pattern and held together by electrostatic forces of attraction


Ionic compounds often form crystals, making them extremely hard and strong

CC5C - PROPERTIES OF IONIC COMPOUNDS



Consist of many ions, held together by electrostatic forces of attraction/bonds
Held together in a lattice structure

Melting and boiling points –




A lot of energy needed to overcome strong electrostatic forces
Require high temperatures to break bonds
Those with more than one charge have an even high boiling/melting point

Electrical conductivity –







When molten/dissolved in water (aqueous solution), conduct electricity
Do not conduct electricity in a solid state
Must contain charged particles and they must be free to move
Ions need carry the current to establish conductivity
Anions attracted to anode (negative ions to positive charge)
Cations attracted to cathode (positive ion to negative charge)

CC6A - COVALENT BONDING
IONIC BONDS – metal to non-metal (one element has too loose electron in order to give to
other), strong electrostatic bonds between oppositely charged ions
COVALENT BONDS – non-metal to non-metal in attempt to complete outer shells by sharing
electrons
MOLECULAR SUBSTANCES – groups of atoms held together by strong covalent bonds

(Cl2)

CH4, methane

Single covalent bond – only one electron is shared
Double covalent bond – two electrons are shared


Used to stabilise atoms by completing their outer shell

MOLECULAR SUBSTANCES – group of atoms held together by
covalent bonds
VALENCY – number of covalent bonds formed by atoms / number of
electrons needed to complete the outer shell

(You can also swap the valencies between elements)

CC7A - MOLECULAR COMPOUNDS
COMPOUNDS – more than one different type of atom joined by bonds
MOLECULES – distinct group of atoms joined together by covalent bonds (can be different
types), all have a covalent, simple molecular structure
MELTING AND BOILING POINTS –
INTERMOLECULAR FORCES – weak forces of attraction between molecules




These forces must be overcome for a substance to change state
The smaller and simpler the molecule, the lower the melting and boiling point
less energy being needed to overcome bonds

CONDUCTION OF ELECTRICITY –
ELECTRIC CURRENT – flow of charged particles




simple molecules have no charge due to them not having an electric current
due to covalent bonds, the strong forces between the positive nuclei and negative
electrons hold electrons in place
this prevents them from moving and carrying a current

POLYMERS –
MONOMER – small simple molecule that can join in chain to form polymer
POLYMER – monomers joined together in a chain by intermolecular forces


the longer the polymer, the more intermolecular forces, meaning the higher the
melting and boiling point

(between molecules, electrostatic forces, between monomers, intermolecular forces (weaker) )
CC7B – ALLOTROPES OF CARBON
ALLOTROPE – different structural form of the same element (influences properties and uses)
Material
Diamond

Structure
Giant covalent structure

Bonding
Linked by covalent bonds – 4
bonds per atom

Graphite

Giant covalent structure

Graphene

Giant covalent structure,
no fixed formula, 2D
hexagonal sheet
Simple molecular
structure - fullerenes

Linked by covalent bonds – 3
bonds per atom, weak
intermolecular forces
Strong covalent bonds – 3
bonds per atom

Buckminster
fullerene – c60

Linked by covalent bonds – 3
bonds per atom, weak
intermolecular forces

Properties
High melting and boiling
points, doesn’t conduct
electricity, extremely hard
High melting and boiling
points, soft and slippery

Uses
Jewellery, used for cutting
things

Lightest material (one
atom thick), strong,
flexible, transparent
Low, melting/sublimation
point, soft and slippery,
very strong, unstable at
high temperatures

Good electrical conductor

Electrodes for electrolysis,
lubricant

Lubricant

Buckminster fullerene (Bucky ball) –

Diamond -

Graphene –

Graphite -

CC7C – PROPERTIES OF METALS
Properties 





Malleable – used for infrastructure/building
Good conductors of electricity – used as wires in circuits (copper)
Shiny – making jewellery
Very high melting and boiling points
High density

Metallic bonding 



Metals tend to have 1-3 electrons on their outer shell, which become delocalised
meaning they are free to move around
This establishes positive ions (due to loss of electrons) in a sea of negative, delocalised
electrons
Electrostatic attraction between cations and anions

Metallic structure –




Atoms packed together in lattice structure
Stacked layers of positive ions
Sea of delocalised negative electrons that move around feely

1
...
Malleable – layers of ions slide over one another preventing them from breaking, sea of
electrons and electrostatic force of attraction hold metal together, keeping its shape,
making it bendable
3
...
Dot-to-dot :
 shows how electrons are shared in covalent bonds
 Doesn’t show structure formed
 Suggests all the electrons are different when they are all the same
2
...
3D ball and stick:
 Show which toms are joined together
 Show shape of structure
 Show atoms are too far apart
 Not actually sticks holding bonds together
CC8A - ACIDS AND ALKALIS




All aqueous solutions (dissolved in water) are either acidic, alkaline or neutral
Universal dictator turn acids red, alkaline purple, and neutrals green
HAZARDOUS TO ENVIRONMENT – avoid putting down sink

Do not consume or
inject – wash hands
once touched

Avoid direct contact with source
(wear gloves and protect skin)

IONS – charged particles due to a loss or gain of electrons
POLYATOMIC IONS – groups of charged particles due to a loss or gain of electrons



Acids produce an excess of H+ when dissolved in water (hydrogen)
Alkalis produce an excess of OH- when dissolved in water (hydroxide)

Keep away from
all heat sources

CC8B – LOOKING AT ACIDS
-

The higher the concentration of hydrogen ions (H+) the more acidic it is
The higher the concentration of hydroxide ions (OH-), the more alkaline it is
Neutral solutions will have equal concentrations of hydrogen and hydroxide
The stronger the acid, the more the acid will dissociate into ions in aqueous sol
...

2
...

4
...


Add excess base to acid to ensure all acid is used up
Warm to fasten reaction
Filter to remove precipitant (unwanted solid), leaving salt and water
Evaporate water using Bunsen burner (crystallisation)
Leave in warm, dry place for crystals to form
CC8C PREPARING COPPER SULFATE CORE PRACTICAL

METHOD –







Measure 20cm3 of dilute sulfuric acid using measuring cylinder, pour into conical flask
Warm acid in water bath to 50 degrees
Add a little copper oxide powder to acid and stir
If all reacts and disappears, add more – stop when in excess
Filter and place filtrate in evaporating basin
Heat and pour solution into watch glass, leave water to evaporate and form salts

APPARATUS –







Copper oxide (base), sulphuric acid
Water bath
Heater/Bunsen burner
Thermometer
Filtering equipment
Conical flask and measuring cylinder

CC8D – ALKALIS AND BALANCING EQUATIONS




Base is a substance that reacts with an acid to form salt and water
Many bases are insoluble
A soluble base is known is an alkali (pH above 7) – group 1 and 2 hydroxides, sodium
and calcium
CC8D INVESTIGATING NEUTRALISATION CORE PRACTICAL

METHOD –







Use a measuring cylinder to add 50cm3 to dilute hydrochloric acid to a beaker
Estimate and record pH of beaker
Place universal indicated paper onto white tile and drop solution, record pH
Then measure 0
...
4 g of calcium hydroxide has been added

APPARATUS –







Beaker
Dilute hydrochloric acid
Universal indicator paper
White tile
Calcium hydroxide
Glass rod

CC8E - ALKALIS AND NEUTRALISATION
1
...
Sodium hydroxide + other alkali (hydroxide ions) = NaOH = Na + OH
NEUTRALISATION – hydrogen ions form acid and react with hydroxide ions that form alkali,
correct proportions balance each other out, forming water





All aqueous solutions of acids contain hydrogen ions, and these are displaced during a
reaction
The negative ions do not change
When salt is formed, hydrogen ions in the acid are replaced by metal/ammonium ions
The hydrogen ions react with the oxide ions to form water molecules
(MgO + H2 SO4 = MgSO4 + H2O – Magnesium displaced hydrogen ions)
(WATER - H + OH = H2O)

BASE – any substance that reacts with an acid to form a salt + water (insoluble in water)
ALKALINE – ionic salt that is soluble in water (often group 1 & 2)

TITRATION –







Use a burette to measure correct amount of acid and store in there
Use a pipette to measure alkali and store in conical flask
Add a few drops of indicator to alkali (methyl orange or phenolphthalein)
Adjust tap to slowly and accurately drop acid into alkali
As soon as alkali has changed colour, reaction has occurred
Calculate how much acid was used by its initial and final value
(Evaporate neutralised solution and leave to form dry, pure salt)

BURETTE – allows drop by drop of acid into alkaline to neutralise
-

Alkali with acid to form soluble salt and water
Without burette, salt could be contaminated due to excess of acid into alkaline

METALS REACTING WITH ACIDS AND CARBONATES Test for Hydrogen –




Collect gas in test tube
Hold lit splint at end of tube
Hydrogen will burn with an explosive/squeaky pop if present

Test for carbon dioxide –



Bubble gas through limewater using delivery tube
If present, limewater turns from clear to cloudy (salt forming – produces precipitate of
calcium carbonate)

CC8F – REACTIONS OF ACIDS WITH METALS AND CARBONATES
Acids and metals –




The lower the metal is within the reactive series, the less it will react with dilute acids
(copper and silver don’t react)
The higher the metal is within the reactive series, the more it will react with dilute acids
(potassium and sodium)
Those in the middle react steadily with dilute acids

EFFERVESENCE - presence of hydrogen gas bubbles produced
(metal + acid = salt + hydrogen)
SPECTATOR IONS – other ions formed in an acid that don’t change (sodium and nitrate)
HALF EQUATION – hydrogen ions gain electrons to form hydrogen molecules
(2H + 2E = H2)
OXIDATION – loss of electrons in a reaction (magnesium)

REDUCTION – gain of electrons in a reaction
General equation for reaction of metal carbonate and acid –
“Calcium carbonate + hydrochloric acid = calcium carbonate + water + carbon dioxide”
IONIC EQUATION – metal atoms react with hydrogen atoms to form metal ions and hydrogen
molecules - write out all aqueous substances as ions (split up so operate individually), cancel
out ions that are present on both sides due to being spectator ions
// CaCO3 (s) + 2H (aq) + 2Cl (aq) = Ca2+ (aq) + 2Cl (aq) + H 2O (l) + CO2 (g)
(to fully form equation, cancel out spectator ions, in this case 2Cl)



H+ = Lost electron – oxidation
Cl2- = lost two electrons – oxidation
- Calcium = neutral to positive
- Hydrogen = positive to neutral

CC8G - SOLUBILITY
PRECIPITATION REACTION – whereby soluble substances in a solution cause insoluble
precipitates to form
PRECIPITANT – substance to be deposited in a solid form from a solution
SOLUBLE – dissolves in water
INSOLUBLE – doesn’t dissolve in water
Soluble in water –






All common sodium, potassium and ammonium salts
All nitrates
Most chlorides
Most sulphates
Sodium, potassium and ammonium carbonates and hydroxides

Insoluble in water –




Silver and lead chlorides
Lead, barium and calcium sulphates
Most carbonates and hydroxides

Precipitation method can be used for form an insoluble salt 

e
...

2
...

4
...
MASS OF EACH COMPOUND/RELATIVE ATOMIC MASS
2
...
MULTIPLY NUMBERS IF NOT WHOLE NUMBERS
MASS
RELATIVE ATOMIC MASS
MASS/RAM
DIVISION BY SMALLEST
NUMBER

CA
10g
40
0
...
25/0
...
8g
35
...
5
0
...
25 = 2

CA = 1

CL = 2
= CaCl2

MOLECULAR FORMULA - how many atoms of each element are actually in a molecule
RELATIVE FORMULA MASS/ EMPIRIAL FORMULA MASS
MULTIPLY EMPIRACLE FORMULA BY THE RESULT
1
...

3
...

5
...
02 x 1023)

e
...
of moles in 10g of calcium carbonate (CaCo3) –
ATOMIC MASS - CaCo3 = 40 + 12 + (16x3) = 100
Mass/molar mass = no
...
1 mol

CC10A - ELECTROLYSIS


when forming an oxide, ionic bond is formed
“Magnesium + oxygen = magnesium oxide”
(Magnesium loses two electrons to the oxygen atoms, becoming Mg2+)
(Oxygen has been reduced, gaining two electrons from magnesium, becoming O2-)

REDOX REACTIONS – whereby bother reduction and oxidation occur, involving transfer of
electrons
“2Mg (s) + O2 (g) = 2MgO (s)”
Oxidation: magnesium, Mg -> Mg2+ + 4e- (used to have)
Reduction: oxygen, O2 + 4e- -> 2O2IONIC COMPOUNDS –





made up of positive metal ions and negative non-metal ions
two ions attract one another due to strong electrostatic force of attraction
requires great energy to split
sodium chloride – positive sodium ions (cations)
- Negative chlorine ions (anions

MOLTEN – solid melted down – complete separation
-

when in a solid form, particles can`t move so substance cant transfer electricity
when in a molten or liquid form, particles can move so transfer a current

AQUEOUS – dissolved in water, H+ and OH_ ions incorporated
Cathode:
-

The metal will be produced if it is less reactive than hydrogen
Hydrogen will be produced if the metal is more reactive than hydrogen
“2H+ + 2E -> H2”

-

If non-metal isn’t a halogen/group 7, oxygen formed
“4OH- -> O2 + 2H20 + 4E-“
If halogen present, halogen formed at anode

Anode:

ELECTROLYSIS – an ionic compound conducting electricity, current causing substance to split
up and form new substance (breaks down into primary elements)
-

purifying metals
extracting reactive metals
produces chlorine, hydrogen and sodium hydroxide

ELECTROLYTE – solution consisting of positive cations and negative anions, separate from one
another
ELECTRODES – two conductive poles where the reactions occur

1
...
Anions -> positive electrode (anode)
- loses electrons, oxidation
CC10A PURIFYING COPPER CORE PRACTICAL
(Copper sulphate solution)


Two electrodes – pure copper -> negative electrode (cathode)
- Impure copper -> positive electrode (anode)

METHOD –









Use two pieces of clean copper foil – label one anode and one cathode
Measure and record masses of the two
Set up electrolysis circuit
Turn on power and adjust variable resistor to give current of 0
...
3, 0
...
5










Electrolyte consisting of copper sulphate solution
Copper ions become ionised (positive), due to gaining electron from sulphur, move
towards cathode
Mass of anode decreases, oxidation occurs
Mass of cathode increases, reduction occurs
As current increases, change in mass increases
Some mass lost due to impurities collecting as mud at bottom of beaker
Goggles needed to protect eyes from harmful solution
Propanone used to completely sterilise the anode and cathode to enable accurate
measurement

Current (amps)
0
...
3
0
...
5

Change in mass at anode (g)
-0
...
13
-0
...
21

Change in mass at cathode (g)
+0
...
11
+0
...
17

CC10B - PRODUCTS OF ELECTROLYSIS
LEAD BROMIDE

Cation -> cathode (metal)

anion -> anode (non-metal)

Reduction

oxidation

“PbBr2 (l) -> Pb (l) + Br2 (g) “
IONIC SALTS –




When electrolysing, electrodes are made of something inert (unreactive – group 0/8)
When salt is electrolysed, ions become discharged
Molten salt always decomposes into elements
METAL –> CATHODE (reduction, 2:1)
NON-METAL -> ANODE (oxidation, 1:2)

POTASSIUM CHLORIDE

Cation -> cathode (metal)
K+ (l) + e- -> k (l)

anion -> anode (non-metal)
2cl- (l) -> cl2 (g) + 2e

IONIC COMPOUNDS DISSOLVED IN WATER –





Water can be ionised too
H+ and OH- ions present in water
Products formed at electrodes depend on whether water ions
discharge more easily than salt ions
Ionic compound with metal more reactive than hydrogen will
be replaced with H+ (produced instead of metal) due to being
discharged more difficultly

Sodium chloride -> Na+ more reactive than H+ therefore H+ replaces
Na+ and goes to negative electrode, Na+ stays in solution
CC11A – REACTIVITY
REACTIVITY SERIES – list of metals in order of reactivity, most reactive at the top
SPECTATOR IONS – ions that remain the same during a reaction
HALF EQUATIONS – way of representing the change of electrons (from ionic equation)
DISPLACEMENT REACTION – whereby a more reactive metal takes place of a less reactive metal
compound, being a redox reaction (both oxidation and reduction occur)
Metals + cold water -> hydrogen + metal hydroxide
Metals + steam -> hydrogen + sold metal oxide
Metals + dilute acid -> hydrogen + salt solution (first name from metal, second from acid)

Equation – Zn + CuSo4 -> Cu + ZnSo4 (zinc displaced copper)
Ionic – Zn + Cu2+ + So42+ -> Cu + Zn2+ + So42Zn + Cu2+ -> Cu + Zn2+
Half-equation – Zn -> Zn2+ + 2e (OXIDATION)
Cu2+ + 2e -> Cu (REDUCTION)

CC11B/C – ORES AND OXIDATION/REDUCTION
ORE – rock that contains enough of a compound to extract a metal for profit
EXTRACTION – the method of which is used to obtain a metal from ores/compounds
NATIVE STATE – whereby something isn’t combined with any other elements/impurities



Unreactive metals (gold/platinum) found naturally in their native state
More reactive metals (oxides, chlorides and carbonates )have reacted with other
elements to form compounds in rocks
Methods of extraction –
1
...
Electrolysis –
 For those higher than carbon in reactive series
 Electricity passed through molten, ionic compounds
 Decomposes into separate elements by breaking of bonds
 A lot of energy and money required to perform
 Only used for very reactive metals
“Aluminium oxide -> aluminium + oxide/ 2Al2 -> 4Al + 3O2”
Obtaining Aluminium from electrolysis 



Al3+ ions attracted to cathode (negative) – gain 2 electrons, undergoing reduction
O2-, Oxide ions attracted to anode (positive), looses 2 electrons, undergoing oxidation
At high temperature of 1000 degrees, oxygen reacts with graphite/carbon at the anode
to form carbon dioxide

3
...
Phytoextraction  Growing plants that absorb metal compounds
 Then burnt to ash, from which the metal is extracted

CORROSION – reaction of metal and oxygen, making metal weaker over time – metal oxidised
due to gain of oxygen (the more reactive a metal is, the faster it corrodes)
RUSTING - corrosion of iron, also requiring water as well as oxygen
TARNISH – a protective oxide layer formed on some metals, such as aluminium, preventing
further corrosion
CC11D – LIFE CYCLE ASSESSMENT AND RECYCLING
RECYCLING – the conversion/manufacturing of waste into reusable material




Every UK household estimated to produce 1 tonne of waste each year
45% is recycled – EU target to reach 70% by 2030
2 million tonnes of waste electrical equipment – worth £1 billion and £36 M of Al

Recycling metals –









Many metals can be recycled by melting them down into something new
Natural reserves of metal ores will last longer
Need to mine ores reduced, less damage to landscapes and noise/dust pollution
Less pollution may be produced (less sulfur dioxide from metal sulphide extraction)
Many metals need less energy to recycle them than to extract new metal from ore
Less waste metal ends up in landfill sites
Costly and high energy needed for collecting, transporting and sorting
Can be more costly than just extracting new metals

Life cycle assessment –




LCA can be carried out to work out an environmental impact of a product
Helps decide whether recycling and manufacturing of a product is worthwhile
Can also be used to compare effect of using different materials for the same product

CC12A – DYNAMIC EQUILIBRIUM
REVERSIBLE REACTIONS – a chemical reaction which can work in both reactions, whereby the
products can react to reform the products “⇌”
DYNAMIC EQUILIBRIUM – reached when a reaction is still occurring however substances within
it remain balanced – when the forwards and backwards reactions in a reversible reaction are
occurring at the same time
CLOSED SYSTEM – no loss of reactants and products, nothing added into observed environment
due to being enclosed
OPEN SYSTEM – system of which substances can leave and enter (open test tube)
EQUILIBRIUM POSITION – whereby percentage of products and reactants are at equilibrium



Dynamic equilibrium doesn’t always occur at the “half way point” within a reaction
Can be shifted by the alteration of temperature, gas pressure and concentration

Haber process –

Forward reaction: exothermic
Backward reaction: endothermic
- All in gas form
- Reactants have higher pressure of 4 molecules
- Product has a lower pressure of 2 molecules
-

-








Process of manufacturing ammonia for industrial use
Needed for fertilisers, cleaning fluids and floor wax
Involve reversible reaction of nitrogen and hydrogen
Nitrogen obtained from the air
Hydrogen obtained from reaction of steam and methane
Conditions in process are chosen to favour forward reaction and make a larger yield
production, at the cheapest cost

Conditions –




450 degrees – low as possible to favour forward reaction (exothermic) and increase
product formed, can’t be too low as heat needed for efficiency of reaction
Pressure of 200 atmospheres – high as possible to favour forwards reaction towards
product which has lower pressure
Use of iron catalyst – speeds up reaction further

Process –









Steam reacted with methane to produce hydrogen
Hydrogen and nitrogen mixed (obtained from air)
Gases compressed to 200 atmospheres
Heated to 450 degrees
Passed over an iron catalyst
Ammonia gas produced, cooled to liquid
Liquid pumped off to be a solid
Unreacted nitrogen and hydrogen recycled

CC13A - GROUP 1
Group 1 = alkali metals (all have similar properties)
-

Very reactive with water and oxygen (stored in oil)
Relatively low melting points
Soft and easily cut
Form compounds with non-metals
Malleable
Conduct electricity
Reactivity increase as you go down the series

1
...
Chlorine – sodium chloride used in salts and in pools to clean water
3
...
Iodine – sodium iodide used in salts and to prevent iodine deficiency
(All can be used as detergents/bleach due to toxic/cleansing properties)

Increasing reactivity further down –








Atoms loose electrons on their outer shell in order to become stable
Force of attraction between negative electrons and positive nucleus
Those higher up have a stronger force of attraction due to the electrons being nearer to
nucleus
This makes the bond harder to break, and their reactivity weaker
The further down you go, the higher the reactivity becomes
The distance between electrons and the nucleus increase, making the force weaker
The bond is broken more easily, causing reactions to be fiercer
(distance increases, force decreases so ions form less easily and reactivity decreases)

Reactions with water(alkali metal + water = metal hydroxide + hydrogen)
1
...
Sodium + water – melts into ball and fizzes around surface
2Na (s) + 2H2 O (l) = 2NaOH (aq) + H2 (g)
3
...
Lithium + oxygen –
4Li (s) + O2 (g) = 2Li 2O2 (s)
2
...
Potassium + oxygen –
4K(s) + O2 (g) = 2K 2O2 (s)

CC13B - GROUP 7
Group 7 = halogens
-

All have seven electrons on outer shell, meaning they gain one to become stable (1-)
Bad conductors
Diatomic structure (two atoms held by single, covalent bond)
Often used as disinfectants/bleach/cleaning products
Reactivity decreases as you go down series due to the increased distance making
the electrostatic force of attraction stronger
harder to break and weaker reactivity (due to trying to gain electron)

(Fluorine, chlorine, bromine, iodine, astatine)
HALOGEN
Fluorine
chlorine
Bromine
Iodine


RELATIVE
SIZE
-1
0
1
2

MELTING
POINT
-220
-101
-7
144

BOILING
POINT
-118
-34
59
184

STATE AT
ROOM
TEMP
Gas
Gas
Liquid
Solid

React with metals forming ionic compounds/salts (contain halide ions – x-)

Reactions with hydrogen –





Halogens react with non-metals by sharing electrons and forming covalent compounds
These gases are extremely soluble in water and dissolve to produce acids (aqueous)
Halogen + hydrogen = hydrogen halides
Can convert hydrogen halide to its acid by dissolving it in water
“H2 + Cl2 = 2HCl” – hydrochloric acid (aqueous when dissolved in water)

Test for chlorine –



Turns blue litmus paper red, then white
Turns bleach white

CC13C - HALOGEN REACTIVITY
Halogens + metal = halide salts
DISPLACEMENT REACTION – whereby a more reactive element replaces a less reactive element
in a compound (more reactive halogen replaces less reactive to form halide compound)
“Chlorine + sodium bromide = bromide + sodium chloride”
REDOX REACTION – reaction in which both oxidation and reduction occurs – (OIL,RIG)
OXIDATION – the loss of electrons + gain of oxygen
REDUCTION – the gain of electrons
“2Na (s) + Br2 (g) = 2NaBr (s)”
-

Sodium oxidised, lost electron to bromine, bromine reduced, gained electron from sodium

CC13D - GROUP 0 – NOBEL GASES











8 electrons in outer shell – no delocalised electrons
Monatomic (singular atoms)
Inert (non-reactive) – no electrons to lose or gain
Non-metals
Low boiling and melting points
Bad conductors of electricity (no free electrons to move around)
Reactivity remains the same throughout series
As you move down series, density increases due to more outer shells
Used to be group zero however helium didn’t fit into requirement
Weren’t in first periodic table due to being so unreactive, couldn’t be
detected

1
...
ARGON – used in wine barrels to prevent wine oxidising (more dense than air)
3
...
NEON – long lasting illuminated signs (produces red/orange light when current passed
through)

CC14A - RATES OF REACTION
CHEMICAL REACTION - when one or more reactants form one or more products
RATE OF REACTION - speed of which reactants are turned into products (frequency of collisions
and amount of energy needed)
Chemical reactions –






Colour change
New product formed (precipitate – two soluble substances producing insoluble solid)
Effervescence (gas formed)
Irreversible
Temperature change
(No mass lost or gained in reaction, only rearranged)

-

Rusting/eroding = slow reaction
Explosions/ potassium & water = quick reaction





Steeper the slope – faster the reaction (gradient)
No slope/flat line – all energy used up and reaction complete
Greatest speed of reactant in beginning due to highest concentration of reactants
available
Concentration of reactants decrease whilst products increase
Reactions don’t proceed at steady rate
Gas syringes used to measure rate of reaction as traps product and measures production
in given time





CC14B - FACTORS THAT AFFECT RATE OF REACTION


For reaction to occur, atoms of reactants must collide with one another with enough
energy

ACTIVATION ENERGY – minimum amount of energy required for a reaction to occur

Factors that affect rate of reaction –
1
...
Temperature – the greater the temperature, faster rate of reaction (particles provided
with more energy to collide into others)
3
...
Pressure – the greater the pressure, faster the reaction (particles more tightly compact
and have higher chance of bumping into one another)
CC14B INVESTIGATING REACTION RATES CORE PRACTIAL
(Investigating how surface area affects rate of reaction)
“CaCo3 + 2HCl -> CaCl2 + H2O + CO2”







Set up apparatus
Measure 40 Cl of HCl into conical flask
Weigh 2g of each sized marble chips (powder, small, medium, large)
Add marble chips into HCl and immediately place stopper over, connecting to syringe
Measure gas every 30 seconds
Note volume at the start and end

-

Smaller the chips, greater the surface area, quicker the reaction





Dependant variable – amount of gas produced
Independent variable – surface area of marble chips
Controlled variables – concentration of HCl and mass of marble chips
CB14B CORE PRACTICAL 2
(Investigating how concentration affects rate of reaction)









Set up apparatus
Measure 50 Ml of thiosulfate of solution into conical flask
Measure 5 Ml of 0
...
5 and 1
...
Enough energy required to break initial bonds (activation energy), taken in
2
...
BOILING POINT – larger the hydrocarbon/ longer the chain, higher the boiling point –
more, stronger intermolecular bonds, harder to break
2
...
EASE OF IGNITION/FLAMMABILITY - larger the hydrocarbon/ longer the chain,
harder to set alight/smaller ease of ignition – more, stronger intermolecular bonds takes
longer to break apart and alight

CC16C - THE

ALKANE
HOMOLOGOUS

SERIES
ALKANE – hydrocarbon compounds whereby each carbon atom has made a single and complete
bond with four other atoms, forming a homologous series


each bond used for separate atom and no new ones to be made, therefore saturated

HOMOLOGOUS SERIES – all molecules follow a similar molecular formula whereby:





The molecular formulae of neighbouring compounds differ by CH2
They have the same general formula
They show a gradual variation in physical properties (e
...
2, results from carbon dioxide, sulphur
dioxide and other non-mental oxides forming an acid solution with water
-

results in soil becoming too acidic, crops being unable to grow well, reduced soil
fertility
damages aquatic environments, prevents eggs hatching and kills animals
increases rate of weathering (limestone, marble, iron, breaks down structures
increases rate of corrosion of metals (iron)

Sulfur dioxide –





hydrocarbon fuels, petrol/diesel oil, may contain sulphur compounds
these occur naturally as impurities and aren’t deliberately added
most are removed at oil refineries in attempt to reduce environmental impact
when the fuels are burnt, they react with oxygen to form sulphur dioxide gas

Oxides of nitrogen 





Car engines are “internal combustion” engines – fuel mixed with air and ignited inside
Causes temperatures high enough for nitrogen and air inside to react together
Reaction produces various oxides of nitrogen
...
5 billion years ago, and to have changed
significantly since then
Asteroids collided with one another and created large sums of heat energy trapped in
the earth`s core (molten lava)
Its surface was initially molten, and cooled to form rocks
Atmosphere consisted of carbon dioxide, water vapour , ammonia and methane
There was no oxygen or water
Volcanoes release large amounts of carbon dioxide and water vapour, and smaller
amounts of other gases – nitrogen and hydrogen sulphide
Some think the Earths early atmosphere consisted of nitrogen instead of carbon dioxide
in relation to Titan (moon of Saturn)

The oceans –




Around 4 billion years ago, the earth cooled down significantly
This caused water vapour in the atmosphere to condense to liquid water and formed
ocean bodies
Simple plants evolved

Introduction of oxygen –








Iron Pyrite can be used to prove the initial absence of oxygen
Broken down by oxygen, and therefore is only formed when no oxygen is present
Can be found in very ancient rocks
Around 2
...
8 BYA, simple life forms (tiny microbes) emerged in oceans
Some of the earliest photosynthetic organisms known as cyanobacteria - live in deep
shallow waters
Grow in huge colonies and produce sticky mucus
Traps layer of sand grains and other sediments
Organisms need to move around layer to get sunlight – which then build up overtime to
form rocky shapes, known as stromatolites
These are over 3 billion years old
Other bacteria deep within ocean – no access to sunlight so rely on other chemicals for
energy as a matter of chemosynthesis
Near hydrogen vents to permit this
Provide evidence of photosynthetic organisms living at this time
Thought that these microorganisms cause a rise in levels in oxygens and atmosphere
Cyanobacteria then evolved into other life forms
Land plants evolved 500 million years ago, furtherly decreasing carbon dioxide levels
and increasing oxygen levels

Nitrogen –





Always been present within the atmosphere
Relatively unreactive there accumulated over the years
Being released through volcanoes
Used up within haber process to produce ammonia artificially

Test for:
Hydrogen – lit splint into test tube, squeaky pop if present
Carbon dioxide – limewater turns from clear to cloudy if present
Oxygen – reignites a glowing splint if present
CC17C – THE ATMOSPHERE TODAY
CAUSAL LINK – whereby one factor is responsible for causing the other
GLOBAL WARMING – the gradual increase in overall temperature of the earth’s atmosphere,
attributed to the greenhouse effect caused by increased levels of co2 and other pollutants
CLIMATE CHANGE – a change in climate patterns, mainly due to increased levels of co2
Greenhouse effect –












In order to maintain constant temperature, earth must absorb same amount of energy
that is radiated out – greenhouse effect to do so
Energy from sun is transferred to earth by waves (infrared radiation)
Some is absorbed by the earth’s surface, warming it up
The earth emits energy in the form of infrared waves, some back into space
Some gases in air also absorb energy transferred by these waves
Greenhouse gases re-emit back to earth’s surface
warms it up as the greenhouse effect
Include carbon dioxide, methane and water vapour
Without them, surface would be -18 degrees rather than 14
Desirable for the most part – earth inhabitable otherwise
In enhancement, becomes a problem

Correlation and climate change –









Evidence proving human activity to increase greenhouse effect, causing global warming
Thought to cause climate change as well
Increase of burning of fossil fuels since 1850
During this, co2 levels have increased
Also risen temperature of earth’s surface
Strong correlation between carbon dioxide and temperature levels
However not a causal link
Co2 absorbs infrared – reduction of waves laving surface and therefore increases in temp

Collecting evidence –




Amount of co2 measured by monitoring stations around the world
Concentration of gas trapped in ice cores used (doesn’t show cause, only causal link)
Continuous temperature measurements from around the world date back to 1880




Modern thermometers less prone to error and have greater resolution
Sensors and satellites can also be used

Name of gas in air
Nitrogen – N2

% in air
78

Oxygen – o2

21

Argon – Ar

0
...
04

Purposes/uses
used in fertiliser and food
packaging, unreactive
Used in respiration and
burning, reactive
Used in light bulbs, unreactive
– noble gas
Used in photosynthesis and
fizzy drinks, unreactive

CC17D – CLIMATE CHANGE
Methane –







Methane molecules in atmosphere absorb infrared ration and emit some back to earth
as photons
Methane is the most powerful greenhouse gas
Main component of natural gas
Released into atmosphere by rice farming, cattle/ livestock farming, oil and gas
exploration and rotting materials in landfill sites
Bacteria in cattle`s stomach released containing methane
Soil bacteria in landfill sits and rice “paddy” fields produce methane

Effects of climate change –










Ice at south pole and glaciers to melt
Raise of sea levels – flooding in some areas
Loss of sea ice at the poles
Extinction of animals as they move to find cooler habitats
Some areas will become drier and other wetter
Extreme weather events (heavy rainfall, powerful storms and heat waves)
Changes will affect wildlife and growth of crops
Acidification of oceans can cause death of sea life and interrupt food chains
UK now gets less rain in summer and more in winter than it used to

Limiting the impact –




Renewable energy resources (might not be enough to mitigate evident effects)
Global engineering solutions – reflecting sunlight back into space or capturing carbon
dioxide, burry underground
Building flood defences, dams and irrigation systems – damage habitats, expensive and
may not work



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