Revision:General Chemistry

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Chemistry Revision

Classifying Materials

Solids

• strong forces of attraction
• molecules in fixed positions
• don’t move from their positions-definite state and volume
• vibrate about their positions-the hotter the solid becomes the more they vibrate, this causes the solid to expand slightly when heated
• can’t be compressed
• generally very dense

Liquids

• some force of attraction between molecules
• free to move past each other but tend to stick together
• don’t keep a definite shape but do keep a definite volume
• molecules are constantly moving, the hotter the liquid the faster they move causing the liquid to expand slightly when heated
• can’t be compressed
• generally quite dense

Gases

• no forces of attraction
• free to move and only interact with each other when they collide
• don’t keep a definite shape or volume and will always expand to fill any container
• molecules constantly moving, the hotter the has becomes the faster they move causing the gas to expand or exert more pressure on the container
• can be compressed
• very low density

Melting

• when a solid is heated the heat energy goes to the molecules
• making them vibrate more and more
• eventually the strong forces of attraction are overcome and the molecules start to move around

Evaporation

• when a liquid is heated the heat energy goes to the molecules and makes them move faster
• some molecules move faster than others
• fast-moving molecules at the surface will overcome the forces of attraction from the other molecules and escape

Boiling

• when the liquid gets hot enough virtually all the molecules have enough speed and energy to overcome the forces and escape each other
• at this point big bubbles of gas form inside the liquid as the molecules break away from each other

Temperature is constant during changes of state (e.g. melting or boiling).

When a substance is melting or boiling all the heat energy supplied is used for breaking bonds rather than raising the temperature.

During freezing as the molecules fuse into a solid heat is given out as the bonds form so the temperature won’t go down until all the substance has turned to solid.

Atoms

The Mass Number - total number of protons and neutrons The Atomic Number - number or protons

Isotopes

Different atomic forms of the same element which have the same number of protons and a different number of neutrons

Ionic Bonding

Atoms gain or lose electrons to form ions which are then strongly attracted to one another.

Ionic Substances

• ionic bonds always produce giant ionic substances
• the ions form a closely packed regular lattice arrangement
• there are very strong ionic bonds between the all the oppositely charged atoms
• they have high melting points and high boiling points
• dissolve to form solutions that conduct electricity
• when dissolved the ions separate and are all free to move so they carry electric current
• conduct electricity when molten

Covalent Structures

Simple Molecular Substances

• the atoms from very strong covalent bonds creating small molecules of several atoms (strong intramolecular forces)
• however the forces of attraction between these molecules are very weak (weak intermolecular forces)
• the melting and boiling points are very low because the molecules are easily parted from each other
• most are gases and liquids at room temperature
• don’t usually dissolve in water

Giant Covalent Structures

• all the atoms are bonded to each other by strong covalent bonds
• very high melting and boiling points
• don’t conduct electricity
• insoluble in water

Diamond: each carbon atom forms four covalent bonds, has only covalent bonds which are very, very strong

Graphite: each carbon atom only forms three covalent bonds creating layers which are free to move over each other and leaving free electrons so it is the only non-metal which conducts electricity

Metallic Structures

• also consists of a giant structure
• have free electrons which produce all the properties of metals
• these conduct heat and electricity
• hold the atoms together in a regular structure
• also allow the atoms to slide over each other causing metals to be malleable

Common Tests

1. Chlorine bleaches damp litmus paper
2. Oxygen relights a glowing splint
3. Carbon Dioxide turns limewater milky
4. Water- has a boiling point of 100°C, turns whit anhydrous copper sulphate to blue hydrated copper sulphate, turns anhydrous cobalt chloride paper from blue to pink
5. Hydrogen- makes a lighted splint give off a squeaky pop

Earth Materials

Fractional Distillation of Crude Oil

• crude oil is a fossil fuel formed from the buried remains of plants and animals
• crude oil provides the fuel for most modern transport
• crude oil is a mixture of hydrocarbons of different sized molecules
• the bigger and longer the molecules the less runny the hydrocarbon is
• fractional distillation splits crude oil up into its separate fractions
• the shorter the molecules the lower the temperature at which that fraction condenses

General Practitioners Never Know Death Or Birth

Gas, Petrol, Naphtha, Kerosene, Diesel, Oil, Bitumen

Hydrocarbons

As the size of the molecule increases:

• the boiling point increases
• it becomes less flammable
• it gets more viscous (doesn’t flow so easily)
• it gets less volatile (doesn’t evaporate so easily)

Combustion of Hydrocarbons

• hydrocarbons are often used as fuels because they burn well
• the complete combustion of any hydrocarbon in oxygen will produce only carbon dioxide and water as waste products which are both quite clean and non-poisonous
• when there’s plenty of oxygen the gas burns with a clean blue flame
• if there isn’t enough oxygen incomplete combustion gives carbon monoxide and carbon as waste products
• a smoky yellow flame is an indicator of insufficient combustion

Cracking Hydrocarbons

• cracking is the splitting up of long chain hydrocarbons
• long chain hydrocarbons like tar aren’t very useful so they are cracked to become shorter molecules that are more useful
• cracking is a form of thermal decomposition which just means breaking molecules down into simpler molecules by heating them
• cracking also produces extra alkenes that are needed to make plastics

Industrial Conditions for Cracking

• vaporised hydrocarbons are passed over a powdered catalyst at about 400-700°C
• the catalyst used is aluminium oxide
• the long chain molecules split apart or crack on the surface of the bits of catalyst

Alkanes

• have single bonds
• called saturated hydrocarbons because they have no spare bonds left
• this is why they don’t decolourise bromine water because they do not form bonds with the bromide ions
• they won’t form polymers-no spare bonds
• burn cleanly producing carbon dioxide and water

• Methane - CH₄
• Ethane - C₂H₆
• Propane - C₃H₈
• Butane - C₄H₁₀

Alkenes

• have double bonds
• called unsaturated hydrocarbons because they have one spare bond left
• they decolourise bromine water
• form polymers by opening up their double bonds
• tend to burn with a smoky flame producing soot (carbon)

• Ethene - C₂H₄
• Propene - C₃H₆
• Butene - C₄H₈

Polymers and Plastics

• under pressure and with a catalyst polymerisation occurs
• many small ethenes form polyethene

Uses of Plastics

• Polythene
• very cheap and strong
• easily moulded

Polypropene

• forms strong fibres
• has high elasticity

Polystyrene

• cheap and easily moulded
• can be expanded into foam

PVC

• cheap and easily moulded
• can be expanded into foam

Most plastics aren’t biodegradable-they’re not broken down by micro-organisms so they don’t rot so it’s best to recycle plastics.

Metal Ores from the Ground

• a rock is a mixture of minerals
• a mineral is any solid element or compound found naturally in the Earth’s crust
• a metal ore is defined as a mineral or minerals which contain enough metal in them to make it worthwhile extracting metal
• metals are extracted from ores using carbon or electrolysis
• Iron ore is Haematite which is Iron (III) oxide (Fe2O3)
• Aluminium ore is Bauxite which is aluminium oxide (Al2O3)
• Copper ore is Malachite which is copper (II) carbonate (CUCO3)
• iron is extracted by chemical reduction using carbon or carbon monoxide
• aluminium is extracted by electrolysis
• gold is one of the few metals that is found as a pure metal rather than in an ore
• the more reactive a metal is the harder they are to extract from their ores

Reduction or Electrolysis

• metals higher than carbon in the reactivity series have to be extracted using electrolysis
• metals below carbon in the reactivity series have to be extracted by reduction using carbon

Extracting Iron: The Blast Furnace

• raw materials: iron ore, coke and limestone
• coke is almost pure carbon, this is for reducing the iron oxide to iron metal
• limestone takes away the impurities in the form of slag
• hot air is blasted into the furnace making the coke burn much faster than normal
• the coke burns and produces carbon dioxide (carbon + oxygen → carbon dioxide)
• the carbon dioxide then reacts with the unburnt coke to form carbon monoxide (carbon dioxide + carbon → carbon monoxide)
• the carbon monoxide then reduces the iron ore to iron (carbon monoxide + iron oxide → carbon dioxide + iron)
• the main impurity is sand which stays solid even at the high temperature, this is removed by limestone
• the limestone is decomposed by the heat into calcium oxide and carbon dioxide (limestone → calcium oxide + carbon dioxide)
• the calcium oxide then reacts with the sand to form slag which is molten (calcium oxide + silicon dioxide (sand) → molten slag
• the cooled slag is solid and is used for road building and fertiliser

Extracting Aluminium: Electrolysis

• bauxite is mined and purified to leave a white powder
• this is pure aluminium oxide which has a very high melting point of over 2000°C
• for electrolysis to work a molten state is required but heating to this temperature would be far too expensive
• instead the aluminium oxide is dissolved in molten cryolite (a less common ore of aluminium)
• this bring s the temperature required down to 900°C which is much cheaper and easier
• the electrodes are made of graphite
• the graphite anode needs replacing quite often because it keeps reacting to form carbon dioxide
• when the aluminium oxide is molten the positive aluminium ions are free to move and are attracted to the cathode
• they pick up electrons and turn into aluminium atoms, these sink to the bottom and flow out
• this is a REDOX reaction

Purifying Copper by Electrolysis

• copper is very unreactive and is extracted from its ore by reduction with carbon
• in a solution containing copper ions a cathode of pure copper and an anode of impure copper

Uses of Iron

• iron is cheap and strong
• however it is also heavy and prone to rusting
• iron is made into steel
• iron and steel are used for construction
• cars, lorries, trains and boats but not planes
• stainless steal doesn’t rust and is used for pans and for fixtures on boats

Uses of Aluminium

• has a low density, is strong and resistant to corrosion
• malleable
• doesn’t corrode due to the protective layer of oxide which always quickly covers it
• a good conductor of heat and electricity
• not as strong as steel and more expensive
• used for ladder, aeroplanes, car body panels, cans, window frames and big power cables

Uses of Copper

• good conductor, easily bent and doesn’t corrode
• ideal for water and gas shapes-can be bent by hand
• wires-easily bent around corners and very good conductor
• forms useful non-corroding allows such as brass and bronze
• quite expensive and not strong

Limestone

Building Material

• great for making into blocks for building with although acid rain can be a problem
• used for carving-statues and patterns
• crushed up into chippings and used for road surfacing

Neutralising Acid in lakes and soil

• ground into a powder can be used to neutralise acidity in lakes and soil caused by acid rain
• works better and faster if turned into slaked lime first

Slaked Lime

• heating the limestone (calcium carbonate) forms calcium oxide
• calcium oxide reacts violently with water to produce slaked lime (calcium hydroxide)

Cement

• clay contains aluminium and silicates and is dug out of the ground
• powdered clay and powdered limestone are heated in a rotating kiln to produce a complex mixture called cement
• when cement is mixed with water a slow chemical reaction takes place
• this causes the cement to gradually set hard
• cement is mixed with sand and chippings to make concrete
• concrete is a very quick and cheap way of constructing buildings

Glass

• heating up limestone and sand and soda until they melt forms glass when cooled

Making Ammonia: The Haber Process

• nitrogen and hydrogen are needed to make ammonia
• nitrogen is easily obtained from the air
• hydrogen is obtained from water (steam) and natural gas (methane)
• steam (H₂O) + methane (CH₄) → carbon dioxide (CO₂) + hydrogen (3H₂)
• hydrogen can also be obtained from crude oil
• the Haber Process is a reversible reaction
• N₂ + 3H₂ → 2NH₃
• the forward reaction is exothermic and the change in energy is negative
• the proportion of ammonia at equilibrium can only be increased by lowering the temperature.
• instead they raise the temperature and accept a reduced yield of ammonia
• the reason is that the higher temperature gives a much higher rate of reaction
• it is better to wait 20 seconds for a 10% yield than to wait 60 seconds for a 20% yield.
• the unused hydrogen and nitrogen are recycled so nothing is wasted
• the higher the pressure the more ammonia produced however this would be very expensive so they use 200-250 atmospheres
• 450°C is a temperature compromise
• the catalyst is iron

Using Ammonia to make Fertilisers

• ammonia gas reacts with oxygen over a hot platinum catalyst to form nitrogen monoxide
• 4NH₃ + 5O₂ → 4NO + 6H₂O
• the nitrogen monoxide then reacts with water and oxygen to form Nitric Acid
• 6NO + 3O₂ + 2H₂O → 4HNO₃ + 2NO
• nitric acid can be used to make ammonium nitrate fertiliser
• this is a straightforward reaction between an alkali and an acid giving a neutral salt
• ammonia + nitric acid → ammonium nitrate
• this is a good fertiliser because plants need nitrogen to make proteins

=Problems with excessive use of Nitrate Fertiliser

• can cause eutrophication
• if too many nitrates get into drinking water it can cause health problems as nitrates prevent the blood from carrying oxygen (children can turn blue and die)
• artificial nitrate fertilisers should be applied sparingly and not when it’s about to rain

Equations

Nine types of Chemical Change

• Thermal Decomposition- when a substance breaks down into simple substances when heated, often with the help of a catalyst
• Neutralisation- acid + alkali → salt + water
• Displacement- when a more reactive element reacts with a compound and pushes out a less reactive element
• Precipitation- when two solutions react and a solid form in the solution and sinks, the solid is said to precipitate out and is therefore called the precipitate
• Oxidation- the addition of oxygen but the general definition is the loss of electrons
• Reduction- the loss of oxygen and the gain of electrons
• Exothermic Reactions- these give out energy as heat
• Endothermic Reactions- these take in energy as heat, heat is needed to break chemical bonds
• Reversible Reactions- reactions that go both ways at the same time

Electrolysis and the Half Equations

• requires a liquid called the electrolyte which will conduct electricity
• these are usually free ions dissolved in water
• it’s the free ions which conduct the electricity

Percentage mass of an element = Ar x No of atoms (of that element) x 100 in a compound Mr (of whole compound)

Calculating Volumes

Volume of Gas (in cm3) = Mass of Gas 24 000 Mr of Gas

Electrolysis Calculations

• treat the two products each as one side of a usual equation
• then use the usual calculations

One mole of atoms or molecules of any substance will have a mass in grams equal to the Relative Formula Mass (Ar of Mr) for that substance

e.g. Carbon has an Ar of 12 so one mole of carbon weighs 12g Carbon Dioxide has an Mr of 44 so one mole of carbon dioxide will weigh 44g

Number of Moles = Mass in g Mr A one molar solution will contain one mole of an element or compound per litre

Number of Moles in a Given Volume = Volume in Litres x Molarity

Air and Rock

The Evolution of the Atmosphere

Phase 1

• the earth’s surface was originally molten for many millions of years
• eventually it cooled and a thin crust formed, but volcanoes kept erupting
• they mainly gave out carbon dioxide but also steam, ammonia and methane
• the early atmosphere was mainly carbon dioxide, carbon monoxide and water vapour
• there was virtually no oxygen
• the water vapour condensed to form the ocean

Phase 2

• green plants evolved over most of the earth
• a lot of the early carbon dioxide dissolved into the ocean
• the green plants steadily removed carbon dioxide and produced oxygen by photosynthesis
• much of the carbon dioxide from the air became locked up in fossil fuels and sedimentary rock
• methane and ammonia reacted with the oxygen releasing nitrogen gas
• ammonia was also converted into nitrates by nitrifying bacteria
• nitrogen gas was also released by living organisms like denitrifying bacteria

Phase 3

• the build up of oxygen in the atmosphere killed off early organisms that couldn’t tolerate it
• it also enabled the evolution of more complex organisms that made use of the oxygen
• the oxygen also created the ozone layer (O3) which blocked harmful rays from the sun and enabled even more complex organisms to survive
• there is virtually no carbon dioxide left now

Today’s Atmosphere

• 78% Nitrogen
• 21% Oxygen
• 0.04% Carbon Dioxide
• 1% Noble Gases (mainly Argon)

Finding the Percentage of Oxygen in the Air

• measure the initial volume of the air
• then push it back and forth over the heated copper
• when no more copper is turning black let it cool
• measure the amount of air left
• calculate the percentage

The Carbon Cycle

• one arrow going down-photosynthesis
• both plant and animal respiration return carbon dioxide to the air
• plants convert the carbon in the air into fats, carbohydrates and proteins
• these can be eaten, decay or be turned into useful products by people
• these plant and animal products ultimately decay or are burned and carbon dioxide is released
• animals also form shells which are eventually washed into the sea, these form limestone and chalk which when heated produce carbon dioxide

Acid Rain

• when fossil fuels are burned they mostly release carbon dioxide
• they also release sulphur dioxide and various nitrogen oxides-which are both harmful
• when these gases mix with clouds they form sulphuric acid and nitric acid
• this then falls as acid rain
• acid rain causes lakes to become acidic killing aquatic plants and animals
• it also kills trees and damages limestone and stone statues

The Greenhouse Effect

• this is causing the Earth to warm up very slowly
• it is mainly caused by an increase in the level of carbon dioxide in the atmosphere
• the carbon dioxide traps the heat that reaches the Earth from the sun
• this causes a change in climate
• it could possibly induce flooding and the melting of polar ice caps

The Ozone Layer

• ozone is a molecule made up of three oxygen atoms, O3
• there’s a layer of ozone high up in the atmosphere
• it absorbs harmful UV rays from the sun
• CFC gases react with ozone molecules and break them up
• the thinning of the ozone layer allows harmful UV rays to reach the surface of the earth
• CFCs are from aerosols

Rocks

• particles get washed to the sea and settle as sediment
• over millions of years these sediments get crushed into sedimentary rocks
• at first they get buried but they can either rise to the surface again to be discovered or can descend into the heat and pressure below
• at first they get buried but they can either rise to the surface again to be discovered or can descend into the heat and pressure below
• if they descend the heat and pressure alters their structure and they become metamorphic rocks
• these metamorphic rocks can either rise to the surface and be discovered or descend further and melt to become magma
• when magma reaches the surface it cools and sets and is called an igneous rock
• extrusive igneous rocks cool above the surface
• intrusive igneous rocks cool below the surface

Metamorphic Rocks

• Slate: formed from mudstone or clay, thin sheets that are ideal roofing material
• Marble: formed from limestone, decorative stone
• Schist: formed from very hot mudstone

Igneous Rocks

• Granite: an intrusive igneous rock with big crystals because it cools slowly, very hard and decorative
• Basalt: an extrusive igneous rock with small crystals because it cools quickly

Sedimentary Rocks

• only sedimentary rocks contain fossils
• fossils are a useful way of identifying rocks of the same age
• sandstone, limestone and mudstone are all sedimentary rocks
• Sandstone: commonly used for buildings
• Limestone: formed from seashells

Weathering

• the process of breaking rocks up

Physical Weathering

• rain water seeps into cracks in rocks
• if it gets very cold the water turns to ice and the expansion pushes the rocks apart
• eventually bits of rock breaks up as the water keeps thawing and refreezing

Chemical Weathering

• caused by acidic rain on limestone
• limestone will react with the acid and just dissolve away

Biological Weathering

• plants push their roots through cracks in rocks
• as the roots grow they gradually push the rocks apart

Erosion and Transport

• the wearing away of exposed rocks by any means
• transport is the process of carrying away of the rock fragments

The Water Cycle

1. Water evaporates off the sea
2. Water transpires from plants
3. It turns to clouds and falls as rain
4. Then it runs back to the sea

Periodic Trends

History of the Periodic Table

• Newlands’ Octaves
  • noticed that every eighth element had similar properties
  • left no gaps in his work so it was ignored.
• Dmitri Mendeleev
  • arranged them in order of atomic mass
  • left gaps in order to keep elements with the same properties in vertical groups
  • the gaps predicted the properties of unknown elements

The inner electron shells provide shielding so that the further away the outer electron shell is from the nucleus the more shielded it is from the attraction.

Group 0 – The Noble Gases

• all colourless, monatomic gases
• helium- has a very low density and won’t set on fire like Hydrogen does, used in airships and party balloons
• neon- when current is passed through it gives out a bright light
• argon- used in light bulbs-stops the hot filament from burning away
• all three are used in lasers too

As you go down the Group:

• the density increases (because the atomic mass increases)
• the boiling point increases

Group 1 – The Alkali Metals

• very reactive-have to be stored in oil because they would react with air or water, handled with forceps because they would burn the skin
• cut easily, they’re shiny when freshly cut but soon go dull as they react with the air
• have low density-they all float

As you go down the Group:

• become bigger atoms (extra electron shells)
• become more reactive because the outer electron is less easily lost
• have a higher density (greater atomic mass)
• become softer to cut
• have a lower melting point
• have a lower boiling point

Reactions of the Alkali Metals

• when lithium, sodium or potassium are put in water they react very vigorously
• they move around the surface fizzing furiously
• they produce hydrogen
• potassium gets hot enough to ignite it
• sodium and potassium melt in the heat of the reaction
• the form a hydroxide in solution
• they react with acids to form neutral salts
• all alkali metal compound are ionic so they form crystals like salt
• they also dissolve easily

Group VII – The Halogens

• all non-metals with colour vapours
• fluorine-very reactive, poisonous, yellow gas
• chlorine- fairly reactive, poisonous dense green gas
• bromine- dense, poisonous, red-brown volatile liquid
• iodine- dark grey crystalline solid or purple vapour
• all form diatomic molecules
• do both covalent and ionic bonding
• must always use a fume cupboard when using them

As you go down the Group:

• become bigger atoms
• become less reactive
• become darker in colour
• go from gas to solid
• have a higher melting point
• have a higher boiling point

Reactions of the Halogens

• the halogens react with metals to form salts
• more reactive halogens will displace less reactive ones
• halogens reaction with hydrogen
• hydrogen chloride gas dissolves in water to form hydrochloric acid
• HCl → H+ + Cl-
• the hydrogen ions make it acidic
• hydrogen bromide and hydrogen iodide also dissolve easily to form strong acids

Industrial Salt

• hot countries obtain salt by pouring sea water into a big, flat open tank and letting the sun evaporate the water to leave salt
• this does not work in cold countries because there is not enough sunshine
• in Britain, salt is extracted from underground mines especially in Cheshire
• this is called Rock salt and is a mixture of mainly sand and salt
• rock salt is used for de-icing roads-the salt melts the ice by lowering the freezing point of water to around -5°C
• when rock salt is filtered to give pure salt it is added to foods to enhance the flavour

Electrolysis of Brine

• salt dissolved in water is called brine
• when concentrated brine is electrolysed it gives three useful products:
• hydrogen gas is given off at the cathode
• chlorine gas is given off at the anode
• sodium hydroxide is left in solution

• Chlorine: used in bleach and for sterilising water
• Hydrogen: used in the Haber Process to make ammonia
• Sodium Hydroxide: used for soap and oven cleaner

Uses Of Halogens

• Fluorine
  • can be added to drinking water and toothpastes to help prevent tooth decay
• Chlorine
  • used in bleach
  • used for sterilising water
  • used to kill germs in swimming pools
• Iodine
  • used as an antiseptic
• Silver Halides
  • silver is very unreactive
  • it form halides but they’re very easily split up
  • visible light has enough energy to do so
  • photographic film is covered with colourless silver bromide
  • when light hits it the silver bromide splits up into silver and bromine
  • the silver metal appears black, the brighter the light the darker it goes
  • this produces a negative

Acids and Alkalis

• acids are substances which form H+ ions when added to water
• alkalis are substances which form OH- ions when added to water
• acid + alkali → salt + water
• indigestion is caused by too much hydrochloric acid in the stomach so indigestion tablets contains alkalis which neutralise the excess acid
• fields with acidic soils are improved by adding lime which is an alkali
• lakes affected by acid rain can also be neutralised by adding lime

Acids reacting with Metals

• acid + metal → salt + hydrogen
• the more reactive a metal the faster it will react
• any metal that is more reactive than hydrogen will react with acid

Acids with Oxides and Hydroxides

• some metal oxides and hydroxides dissolve in water to produce alkali solutions
• they react with acids to form a salt and water because they are alkali
• acid + metal oxide → salt + water
• acid + metal hydroxide → salt + water

Non-metal oxides

• the oxides of non-metals are usually acidic
• carbon dioxide
• sulphur dioxide
• nitrogen dioxide

Acids with Carbonates

• acid + carbonate → salt + water + carbon dioxide
• acid + hydrogencarbonate → salt + water + carbon dioxide

Acids with Ammonia

• dilute acid + ammonia → ammonium salt

Metals

• all conduct electricity
• all good conductors of heat
• are strong but also bendy and malleable
• shiny (when freshly cut or polished)
• all have high melting and boiling points
• can be mixed together to form useful alloys
• steel is an allow of iron and about 1% carbon-it is much less brittle than iron
• bronze is an alloy of copper and tin-it is harder than copper but still easily shaped

Reacting Metals with Water

• if a metal reacts with water it will always release hydrogen
• the more reactive metals react with cold water to form hydroxides
• the less reactive metals don’t react quickly with water but will react with steam to form oxides
• the upper third of the reactivity series will react with cold water
• the middle third of the reactivity series will react with steam
• the upper third of the reactivity series won’t react with water or steam

Transition Metals

• high melting points and high density
• good conductors of electricity
• very dense, strong and shiny
• transition metals and their compounds all make good catalysts
• iron is used in the Haber Process

8their compounds are very colourful

Non-Metals

• either dull, brittle solids or gases (excepting bromine which is liquid)
• poor conductors of heat
• don’t conduct electricity

Reaction Rates

• the speed of a reaction can be observed either by how quickly the reactants are used up or by how quickly the products are forming
• there are three different ways the speed of a reaction can be measured:
• Precipitation- when the product of a reaction is a precipitate which clouds the solution, observe a marker through the solution and measure how long it takes for it to disappear
• Change in mass- any reaction that produces a gas can be carried out on a mass balance and as the gas is released the mass disappearing is easily measured
• The volume of gas given off- this involves the use of a gas syringe to measure the volume of gas given off
• depends on the temperature, concentration, catalyst and surface area of particles

Collision Theory

• the rate of a reaction depends on how often and how hard the reacting particles collide with each other
• temperature increases the number of collisions because the particles have more energy and therefore move more quickly
• concentration (or pressure) increases the number of collisions-the more particles of reactant the more collisions happen, increased pressure means smaller space so more collisions
• surface area increases collisions-the larger the surface area to collide with the more successful collisions
• a catalyst increases the number of successful collisions by providing a surface for the reacting particles to stick to where they can bump into each other

Catalysts

• a substance which increases the speed of a reaction without being changed or used up in the reaction
• catalysts lower the activation energy of reactions making it easier for them to happen
• this means a lower temperature can be used
• catalysts work best when they have a big surface area
• catalysts help to reduce the cost in many industrial reactions simply because the plant does not need to operate for as long to produce the same amount of stuff
• a platinum catalyst is used in the production of nitric acid

Enzymes

• living things produce enzymes which act as catalysts to speed up chemical reactions without the need for high temperatures
• enzymes are used in biological washing powders
• dishwasher powders
• treatment of leather
• enzymes stop working by 0°C and this is why things are refrigerated

Use of Enzymes

Fermentation of Beer and Wine

• yeast cells convert sugar into carbon dioxide and alcohol
• they do this using the enzyme ZYMASE

Fermentation: Bread-making

• yeast cells use the enzyme zymase to break down sugar and produce gas throughout the mixture

Yoghurt and Cheese-making

• milk is mixed with specially grown cultures of bacteria
• various bacterial enzymes can be used in cheese making to produce different textures and tastes

Biological Detergents

• enzymes are the “biological” ingredients in washing powders
• they’re protease and lipase which digest the stains

Simple Reversible Reactions

• e.g. the thermal decomposition of ammonium chloride-heated to make ammonia gas and hydrochloric gas but when these cool they become ammonium chloride again
• e.g. the thermal decomposition of iodine crystals-they will turn into a purple vapour and then reform as grey crystals when they cool
• e.g. the thermal decomposition of hydrated copper sulphate-when heated the water evaporates and they become white anhydrous copper sulphate powder but if you add a couple of drops of water the blue crystals reappear

Equilibrium

• if a reversible reaction takes place in a closed system then a state of equilibrium will always be reached
• a dynamic equilibrium is one in which the reactions are still taking place in both directions but the effect is nothing because they are taking place at exactly the same rate in both directions
• the position of equilibrium depends very strongly on the temperature and pressure surrounding the reaction
• if we alter the temperature and pressure we can move the position of equilibrium to give more product and less reactants
• all reactions are exothermic in one direction and endothermic in another
• if we increase the temperature the endothermic reaction will increase to use up the extra heat
• if we decrease the temperature the exothermic reaction will increase to give out more heat
• many reactions have a greater volume on one side, either of products or reactants
• if we raise the pressure it will encourage the reaction which produces less volume
• if we lower the pressure it will encourage the reaction which produces more volume

Le Chatelier’s Principle

If you change the conditions the position of equilibrium will shift to oppose the change.

The Haber Process

• change in energy is negative
• the forward reaction is exothermic

Energy Transfer in Reactions

• an endothermic reaction is one which takes in energy from the surroundings usually in the form of heat and usually shown by a fall in temperature
• photosynthesis is endothermic
• thermal decomposition is endothermic
• an exothermic reaction is one which gives out heat to the surroundings usually in the form of heat and usually shown by a rise in temperature
• burning fuels is exothermic
• neutralisation reactions are exothermic
• energy must always be supplied to break bonds
• energy is always released when bond form
• bond breaking-endothermic
• bond forming-exothermic
• in an endothermic reaction the energy needed to break the bonds is greater than the energy released when new bonds are formed
• in an exothermic reaction the energy released in bond formation is greater than the energy used in breaking bonds

Energy Level Diagrams

Exothermic Reactions

• in exothermic reactions the energy change is negative
• the products are at a lower energy than the reactants

Endothermic Reactions

• in endothermic reactions the energy change is positive
• the products are higher than the reactants

Bond Energy Calculations

• break the bonds on the left hand side
• form the bonds on the right hand side
• take away the right hand side from the left hand side to find the change in energy
• a negative change is exothermic
• a positive change is endothermic

Comments

I've tried to structure these notes - they might not be sectioned in the best way so please feel fre to restructure them be changing the sizes of heading and the order etc.

Can anyone comment on their completeness?

Are diagrams needed anywhere?

Should the page be split up? or is it OK as one alrge article?