C4,5,6

• when heated, some elements emit distinctive colours e.g.
• Lithium: red
• Sodium: Na
• Potassium: K 
• Each element gives a characteristic line spectrum
• When heated, electrons release energy as light, the wavelengths can be recordedas a line spectrum, different elements emit different wavelengths, due to their electron arrangement
• Line spectrums can therefore be used to identify different elements- SPECTROSCOPY
• Line spectrums have identified new elements e.g. Caesium, rubidium, helium. 

• Dobereiner- triads- 1800s: was only able to measure atomic mass, so the known elements were arranged this way. He started to put them into groups of three based on their chemical properties. Each middle element had an atomic mass that was the average of the other two. 
• Newlands- law of octaves. He noticed that when in order of atomic mass, every eighth element had similar properties . Sets of eights= Newlands octaves, Didn't work because he left no gaps, work critcized by the Chemical Society in 1865 because the groups contained elements that didn't have similar properties, mixed metals and non-metals, didn't leave any gaps 
• Dmitri Mendeleev- 1869- used 50ish known elements and arranged them into the Table of Elements with gaps. Put them in order of atomic mass, but left gaps. Gaps were clever because they predicted properties of undiscovered elements. When the missing were found, they fitted the pattern. 

• In order of increasing proton number, means there are repeating patterns in the properties of elements. 
• Useful for dividing metals and non-metals (metals=left) 
• Similar properties in form groups (vertical columns)
• GROUP NO. TELLS YOU HOW MANY ELECTRONS ON THE OUTER SHELL
• If you know the properties of one element in a group you can predict others' properties
• Can predict reactivity-elements in group 1 react more vigorously as you go down the group. Group 7= reactivity decreases as you go down the group. 
• Rows= periods, each new period represents another full shell of electrons
• R.A.M= total number of protons and neutrons in nucleus
• Proton no.= how many protons and neutrons there are

ELECTRON SHELL RULES

• always occupy shells (AKA energy levels) 
• lowest are always filled first
• only a certain number are allowed per shall (1st=2, 2nd=8, 3rd= 8) 
• much happier when they have full electron shells
• most atoms= outer shell is not full, makes them want to react
• Electron arrangement determines its chemical properties

Working out configurations:

• need to know for first 20 elements
• E.g. nitrogen= 2,5 because the proton number if 7, so 7 electrons and only two allowed in the  first shell. 

• Ions= made when atoms lose or gain electrons- charged particles
• can be made from single atoms or groups of atoms, atoms lose or gain electrons to try and get a full outer shell. Shell with one electron will want to get rid of it
• In group one they all have one electron on their outer shell, so keen to get rid of it and form positive ions and in group seven, outer shells are nearly full so they want to gain an extra electron to form negative ions
• Ionic bonding- transferring electrons- oppositely charged ions are strongly attracted to one another, they tend to lead at ions of opposite charge and form an ionic bond- 
• E.g. sodium and chlorine- sodium atoms gives its other electron and becomes Na+, chlorine picks up the spare elctron and becomes Cl- ion.
• Compounds formed between group one and group 7 elements are held together by ionic bonds- they are ionic compounds-these form a regular lattice
• Solid ionic ocmpounds are made up of a giant lattice of ions, each lattice forms a crystal. When ionic compounds are molten or are dissolved in water they can conduct electricity because the ions are able to move. 
• The fact molten compounds of metals and non-metals can conduct electricity is evidence that they are made up of ions. 

• The charges in an ionic compound add up to to zero
• Some metals can form ions with different charges, the number in brackets after the name tells you the size of the positive charge on the ion- e.g. iron(ll) ion has a charge of 2+, so it's Fe2+.
• Same size charge= easy- e.g. find the formula for lithium fluoride- lithium ion is Li+ and fluoride is F-, so the formula must be LiF. 
• Harder ones: different size charges- need to add some numbers- e.g. find the formula for calcium chloride- a calcium ion is CA2+ and a chloride ion is Cl-  to balance it you need two chloride ions to every one calcium ion, so the formula must be CaCl2 
• Another example: find the formula for iron(lll) sulfate. An iron(lll) ion is Fe3+ and a sulfate ion is SO4^-2. To balance the total charge you need two iron (lll) ions to every three sulfate ions. So the formula must be Fe2(SO4)3. 
• You can also work out the charges on ions if you know the formula of a salt and the charge on one of the atoms- example: Find the charge in the oxide ion in K2O if the charge on the potassium ion is 1+. There are two potassium ions (2x1+)=2+. The charges must balance so the charge on the oxygen ion must be 2-. 

• LITHIUM, SODIUM, POTASSIUM, RUBIDIUM, CAESIUM, FRANCIUM
• All have one outer shell electron, v reactive and gives them all similar properties
• when they react they all form similar compounds
• shiny when freshly cut but oxidise and tarnish in moist air
• As you go down the group they become more reactive because the outer electron is more easily lost because it is further from nucleus, they have a higher density because the atoms have more mass, they have a lower melting point and a lower boiling point. 
• Reaction with cold water prodcues hydrogen gas..
• when lithium, sodium or potassium are put in water they react very vigorously, moving around the surface, fizzinf furiously. They produce hydrogen, potassium gets hot enough to ignite it, if it hasn't already been ignited, a lighted splint  wll indicate hydrogen by producing a "squeaky pop". The reaction makes an alkaline solution, this is why group 1 is known as the alkali metals. A hydroxide of the metal forms e.g. NaOH, KOH and LiOH. 
• Reaction with chlorine produces salts, alkalis react vigorously with chlorine, the reaction produces colourless crystalline salts e.g. sodium chloride. 

• fluorine, chlorine, bromine, iodine astaline
• all have seven outer electrons, v. reactive, gives them similar properties and means they form similar compounds to each other. Halogens form diatomic molecules which are pairs of atoms. As you go down the group they become less reactive, because outer electrons are further fromnucleus so additional electrons are attracted less, have a higher melting point and a higher boiling point. 
• Halogens are all non-metals with coloured vapours: fluorine= v. reactive, yellow, poisonous gas @ room temp. and pressure, chlorine= fairly reactive, poisonous dense green gas @ room temp and pressure, bromine= dense, poisonous, orange volatile liquid @ room temp and pressure, it forms an orange gas. Iodine= dark grey crystalline solid @ room temp, or a purple vapour (halogens go from gases to solids down the group) 
• Halogens become less reactive down the group..
• reactions w. alkali metals: react to form salts called metal halides (NaCl, KBr, LiI) Reactions are less vigorous as you go down the group
• w. iron: form coloured solids called iron halides, less vig. further down the group. 
• Displacement reaction: where one element displaces a less reactive element from a compound- the more reactive halogen will react with the alkali metal. 

Hazard symbols: Oxidising, highly flammable, toxic, explosive, corrosive. 

Alkali metals: group 1 are really reactive and can combust spontaneously, if they come into contact w. water vapour in the air there can be a violent reaction so they are stored in oil to prevent this. Never touched w. bare hands bc sweat is enough to cause a reaction that will produce heat and a corrosive hydroxide. Every piece of apparatus must be completely dry. The alkaline solutions they form are corrosive and may cause blistering. 

Halogens- group 7 elements are also harmful, chlorine and iodine are very toxic. Fluorine is the most reactive halogen, too dangeroud to use inside the lab. Liquid bromine is corrosive and so contact w. skin must be avoided. Halogens have poisonous vapours that irritate the respiratory system and eyes, must be used inside a fume cupboard so you don't breathe in the fumes. 

• 78% Nitrogen, 21% Oxygen, 1% Argon and 0.05% Carbon dioxide.
• Most of the gases are molecular substances.
• Molecular substances have low melting and boiling points because the forces inbetween the molecules are very weak, only a little bit of energy is needed to overcome the force. This means they are usually gases of liquids at room temperature. 
• Pure molecular substances do not conduct electricity because their molecules are not charged, there are no free electrons or ions.
• Most non-metal elements and most compounds formed from non-metal elements are molecular. 

• Covalent bonds means that both atoms feel they have a full outer shell
• each covalent bond provides one extra electron for each atom
• each atom involved has to make enough covalent bonds to fill up its outer shell
• the atoms bond due to the electrostatic attraction between the positive nuclei and the negative electrons
• Hydrogen (H2) needs just one extra electron to fill its outer shell, so two hydrogen atoms share their outer electron so that they have a full outer shell and a covalent bond is formed
• Carbon dioxide (CO2) Carbon needs four more electrons to fill it up, oxygen needs two, so two double covalent bonds form, a double covalent bond has two shared pairs of electrons. 

• Hydrosphere consists of all water in the oceans, seas, lakes, rivers, puddles etc. 
• Contains any compounds that are dissolved in water
• many of these compounds are ionic compounds called salts (e.g. NaCl, MgCl2, KBr) 
• Solid ionic compounds form crystals, they are made of charged particles called ions, oppositely charged ions are v. strongly attracted to eacg other and a giant lattice builds up. There are v strong ionic bonds between all ions. A single crystal of salt is one giant lattice, which is why salt crystals tend to be in a cuboid shape. 
• Ionic compounds have high melting and boiling points, the forces of attraction between ions is v strong, it takes a lot of energy to overcome these forces and melt the compound, and even more energy to boil it. So ionic ocmpounds are solids at room temp. 
• Ionic compounds conduct electricity when dissolved or molten- it dissolves, the ions seperate and are all free to move in the solution, meaning they are able to carry an electric current. Similarly, when an ionic compound melts the ions are again  free to move, so they'll carry an electric current. When an ionic compound is a solid, the ions are not free to move and so an electrical current can't pass through the substance. 

• Flame tests- compounds of some metals give a characteristic colour when heated, so you can use them to detectand identify different ions. 
• PRECIPITATION REACTION: WHERE TWO SOLUTIONS REACT TO FORM AN INSOLUBLE SOLID COMPOUND CALLED A PRECIPITATE.
• Add sodium hydroxide and look for a coloured precipitate: many metal hydroxides are insoluble and precipitate out of solution when you add an alkali, some of these hydroxides have a characteristic colour. So in this test you add a few drops of sodium hydroxide solution to a solution of the mystery compound, and if you get a coloured insoluble hydroxide you can identify the metal ion that was in the compound.
• Ionic equations only show the useful parts of reactions, they just show what you're interested in. 

• Hydrochloric acid can help detect carbonates- with dilute hydrochloric acid, carbonates (CO3^-2) will fizz because they give off carbon dioxide. You can test for CO2 using limewater, carbon dioxide turns limewater cloudy- bubble the gas through a test tube of limeater and if the water goes cloudy you've identified a carbonate ion. 
• Test for sulfates with HCl and barium chloride- to identify a sulfate ion (SO4^2-) add dilute HCl, followed by barum chloride solution (BaCl2). A white precipitate of barium sulfate means the original compound was a sulfate. The hydrochloric acid is added to get rid of any traces of carbonate ions before you do the test- these would produce a precipitate, so they'd confuse the results. 
• Test for halides (Cl-, Br-, I-) with nitric acid and silver nitrate. To identify a halide ion, add dilute nitric acid (HNO3) followed by silver nitrate solution(AgNO3) A chloride gives a white precipitate of silver chloride, a bromide gives a cream precipitate of silver bromide and an iodine gives a yelow precipitate of silver iodide. Acid is added to get rid of carbonate ions before the test, you use nitric acid though, not HCl. 

• Lithosphere is the Earth's rigid outer layer- the crust and part of the mantle. It is made of a mixture if minerals, often containing silicon, oxygen and aluminium. Different kinds of rick contain different elements and minerals. 
• Carbon forms two types of giant covalent structures- diamond and graphite- both found in Earth's crust. Diamond: carbon atoms in diamond each from four covalent bonds in a v. rigid structure. This structure makes diamond the hardest natural substance, the strong covalent bonds give it a very high melting point and it does not conduct electricity because it has no free electrons, even when it is molten, it's insoluble in water. 
• Graphite- made from carbon but it has a different giant covalent structure, each carbon atom only forms three covalent bonds, creating sheets of carbon atoms which are free to slide over each other. This makes it slippy and a good lubricant, the layers are held together so loosely they can be rubbed off onto paper- pencils. V. high metling point, covalent bonds needs loads of energy to break. Conducts electricity because only three of carbon's four outer electrons form covalent bonds, so one is free to move- it is used for electrodes. 
• Slicon dioxide is also a giant covalent structure. Most of the silicon and oxygen found in the crust exists as silicon dioxide (aka silica- what sand is made of) Each grain is one giant structure of Si and O. it has similar properties to diamond because it has a similar structure. 

Ores

• contain enough minerals to make extraction worthwhile
• minerals=solid elements and compounds, rocks are made of minerals
• metal ores- rocks that contain varying amounts of minerals from which metals can be extracted. Ore is often an oxide of the metal e.g. a type of iron ore is haematite, iron (lll) oxide and a type of copper ore is choleopyrite- copper iron sulfide
• Some metals need large amounts of ore extracted just to obtain small percentages-copper!
• More reactive metals are harder to get- most metals need to be extracted using a chemical reaction and more reactive metals (e.g. sodium) are harder to extract.
• A common way of extracting a metal from its ore is by chemical eduction using C or CO
• When an ore is reduced, oxygen is removed e.g. Fe2O3 + 3CO-> 2Fe + 3CO2
• When a metal oxide loses its oxygen, it is reduced, carbon gains oxygen and is oxidised
• How reactive the metal is compared to carbon determines whether it can be reduced with C or CO- less reactive metals then carbon can be extracted by heating with carbon because it can take oxygen away from metals which are less reactive than itself. Metals more reactive than carbon can't be extracted by reduction, they have to be extracted by electrolysis. 

• Needs a liquid to conduct electricity (the electrolyte). These are usually free ions dissolved in water or molten ionic compounds, the free ions are what conduct electricity and allows it all to work
• For the electrical circuit to be complete, there needs to be a flow of electrons. In electrolysis, electrons are taken away from ions at the positive electrode and given electrons at the negative electrode. The ions then become normal molecules. 
• Electrolysis removes aluminium from its ore (bauxite) which contains aluminium oxide. Molten aluminium oxide contains free ions- so it'll conduct electricity. Positive Al3+ ions will pick up three electrons at the negative electrode and negative O2- ions will lose two electrons at the positive electrode, after they'll combine to make O2 molecules. 
• Complete equation for decomposition of alumium oxide= 2Al2O3->4Al+3O2
• At neg electrode: AL3+3e- -> Al (reduction)
• At pos electrode: 2O2- -> O2 +4e- (oxidation) 
• Non-metals form negative ions, metals form positive ions

RAM (Ar) is the top number on an element in periodic table, the bigger one. 

RFM (Mr) add all the RAMs together

Brackets.. e.g... Mg(OH)2 The two means there is two of everything inside the brackets.

Mr and Ar are used to calculate how much metal can be extracted

• e.g. how much coppper can be extracted from 800g of CuO
• STEP ONE: calc proportion of Cu in CuO: Ar of Cu x  no. of atoms of Cu/Mr of  CuO
• STEP TWO: times answer by the mass of copper oxide

• metal properties are due to the free electrons
• metals consist of a giant structure, metallic bonds involve the 'free-electrons' which produce all properties of metals
• the free electrons come from the outer shell of every metal atom in the structure
• the positively charged metal ions are held together in a crystal by the free electrons that can move
• they are good conductors of electricity and heat, bc free electrons carry them both through the material, good for saucepan base or electrical wires. 
• most metals are strong and malleable- metallic bonds mean metals have a high tensile strength. The layers of atoms in a metal can slide over eachother, making metals malleable. 
• generally they have high metlting and boiling points bc metallic bonds are very strong and need lots of energy to break them. 
• Environmental impact: ores are finite sources, people have to balance social, economics and environmental effects of mining. Deep mine shafts can collapse but social and eco benefits. Recycling metals is VERY important bc mining and extracting uses lots of energy (fossil fuels) which are RUNNING OUT and contribute to other problems. Recycling only uses a tiny amount of energy- saves money, conserves needs and cuts down amount in landfill. 

• The chemical industry makes useful products- food additives, cleaning and decorating products, drugs, fertilisers etc. 
• chemists need to figure out how to make them in a way that produces highest yield, do this by controlling the rate of reaction, they also think about environment, choosing processes w. low impact, v. huge industry, produced on a huge scale e.g. sulfuric acid, and small scale, e.g. pharmaceuticals- but they are equally as important. 
• Sectors- C.I. makes up a large part of our economy, lots of different uses and areas within C.I. e.g. pharmaceuticals. 

• pure acidic compounds can be solids, liquids or gases (e.g. citric acid (s), sulfuric acid (l), hydrogen chloride (g)) 
• common alkalis include sodium hydroxide, potassium hydroxide and calcium hydroxide.
• Indicators and pH meters are used to determine pH
• Indicators contain a die that changes colour depending on pH
• Litmus paper turns red if acidic and blue if alkaline
• Universal indicator- combination of dies, good for estimating pH of a solution
• pH meters used to measure the pH of a substance- usually consists of a probe which is dipped into a substance and pH gives a reading of the pH- more accurate than indicators
• Neutralisation reactions happen between acids and alkalis to make salts
• Acid- a substance with a pH less than 7, produce aqueous hydrogen ions, H+, in water
• Alkali- substance with a pH greater than 7, produces aqueous hydroxide ions, OH-, in water
• Acid + alkali -> salt + water
• H+ and OH- ions react to make water
• H+ +OH- -> H2O

• Acid + metal -> salt + hydrogen
• more reactive metal= faster the reaction will go- v. reactive, react explosively
• Copper does not react with dilute acids because it is less reactive than hydrogen
• speeed is indicated by rate at which bubbles of hydrogen are given off
• H is confirmed by burning splint test giving "squeaky pop"
• name of salt produced depends on metal and acid used
• HCl acid will always produce chloride salts, 
• Sulfuric acid will always produce sulfate salts

• metal oxides and metal hydroxides react with acids
• acid + metal oxide -> salt + water
• acid + metal hydroxide -> salt +water
• combination of metal and acid decide the salt e.g.
• Hydrochloric acid + copper oxide -> copper chloride and water 
• sulfuric acid + calcium hydroxide -> calcium sulfate + water 
• metal carbonates give salt, water and carbon dioxide
• acid + metal carbonate -> salt, water and carbon dioxide 
• reaction is similar to any other neutralisation reaction but carbonates give off carbon dioxide too e.g. hydrochloric acid + sodium carbonate -> sodium chloride + water + carbon dioxide

• choosing the reaction- chemists need to choose reaction e.g. neutralisation (produce a salt), thermal decomposition (heat breaks a compound into simpler substances), precipitation (an insoluble solid forms when two solutions are mixed) 
• risk assessment- identifies hazrds, asseses who could be harmed and decides on actions to reduce the risks
• calculate quanities of reactants- calculate how much is needed to produce a certain amount of product- v. important bc. you don't want waste
• choosing apparatus and conditions- need suitable apparatus and right conditions, temperatures, catalysts, sizes etc. 
• isolating product- may need to be seperated from reaction mixture- could involve evaporation, filtration and drying 
• purification- crystalisation can be used
• measuring yield and purity- yield tells you overall success, compares what you get to what you thought and purity needs to be measured to put on bottle. 

Three important steps: 

• write out balanced equation
• work out Mr and multiply by number infront of formula
• apply the rule+ divide to get to one, times to get all

E.g. what mass of magnesium is needed to produce 100g of magnesium oxide?

• 2Mg + O2 -> 2MgO
• 2 x Mg + O) = 2x (24 + 16) = 80
• 48g of Mg reacts to give 80g og MgO
• need to divide by 80 to get 1g of MgO
• then times by 100 to get 100g of MgO

Yield- mass of product, yield calculated in this way are theoretical yields

• Isolation and purification use similar techniques
• Filtration- used to seperate an insoluble solid from a liquid- filtration can be used if the product is an insoluble solid that needs to be seperated from a liquid reaction mixture. It can be usedin purification too, e.g. solid impurities. 
• Evaporation and crystallisation is used to seperate a soluble liquid from a solution. Heating up causes solute to evaporate, leaving solid crystals of product. Also useful for purifying product, crystals have a regular structure that impurities can't fit into. The process is often repeated to improve the purity, products are dissolved and crystallised again= recrystallisation. 
• Drying- used to dry product by remaining excess liquid. Product can be dried in a drying oven, products are also dried using desiccators- containers that contain chemicals like silica gel that remove water from their surroundings. 
• Percentage yield compares actual and theoretical yield. 
• ACTUAL- mass of pure, dry product
• THEORETICAL- maximum possible mass of pure product
• PERCENTAGE YIELD- actual yield / theoretical yield x100

• used to check purity of an acid or alkaline product, work by using neutralisation reactions 

Steps:

• add a known volume of alkali to titration flask, along with a few drops of indicator
• fill the burette with acid (burettes measure different volumes and let you sdd solutions drop by drop
• using a burette, add acid to alkali bit by bit, giving conical flask a regular swirl for mixing. 
• slow when near end point (colour change) is about to be reached
• Indicator changes colour when all alkali has been neutralised. 
• Record volume of acid used to neutralise alkali
• solids are weighed out into a titration flask
• Titrations can be carried out with only liquids so solids have to be made into a solution
• If it is a solid lump, it can help to crush it into powder
• put titration flask onto a balance
• Carefully weigh out some powder into the flask
• add solvent to dissolve the powder
• swirl until all the solid has dissolved. 

• some products need to be very pure
• purity of a product will improve as it is being isolated, but to get a really pure product earlier stages are often repeated- filtration, evaportation, crystalisation.
• Purifying and measuring the purity of a product are particularly important steps in the chemical industry e.g. pharmaceuticals and petrochemicals. 
• Titrations can be used to measure the purity of a substance e.g. begin with 0.2g of impure asprin dissolved in 25cm3 of ethanol in a conical flask, find from titration that it takes 9.5cm3 of 4g/dm3 NaOH to neutralise asprin solution.
• Step 1: work out concentration of asprin solution= 4.5 x conc of NaOH x vol of NaOH/ vol. of asprin solution 
• 1dm3=1000cm3. 
• Step 2: work out mass of asprin- mass= conc x volume
• Step 3: calc purity - % purity= calculated mass/ mass of impure substance at the start x100

• Exothermic- reaction which gives out energy to surroundings, usually shown by a rise in temperature
• Endothermic- takes in energy from surroundings, shown by a fall in temperature
• energy level diagrams show if it is exo or endo- show energy levels of products and reactants in a reaction, can use them to work out if eo or endo. 
• Difference in height represents energy taken in or give out. 
• Energy management is important to control reactions
• if a chemical synthesis reaction is exothermic, heat produced has to be removed, otherwise temperature of reaction mixture will increase 
• if temperature increases, rate of reaction will increase and reaction mixture will get hotter- if too hot some reactants/ products could become hases, could increase pressure- could cause an explosion
• If endothermic, heat needs to be provided it mix could becaome too cold, slower rate or reaction or could freeze, could damage equipment and stop the whole process. 
• Rates of reaction- how fast reactants -> products
• slow examples = chemical weathering, working of iron
• Moderate= metal and acid -> producing a stream of bubbles 

• Rates of reaction- how fast reactants -> products
• slow examples = chemical weathering, working of iron
• Moderate= metal and acid -> producing a stream of bubbles  
• Fast= burning, explosions
• Controlling rate of reactions are very important for:
• Safety and risk of explosion
• Economic reasons- changing conditions can be costly, fuel bills, but faster= more product in less time. Often companies choose optimum conditions that give low production costs, but this may mean compromising on the rate of reaction or the yield.                                           Typical rate of reaction graphs:
• quicker reactions have steeper lines and become flat in the leasttime
• Increased rate could be due to= 
• temperature increase
• concentration increase
• catalyst 
• solid reactant crushed up into smaller bits 

• rate of reaction depends on how often and how hard the reaction particles collide with eachother, particles have to collide in order to react and have to hit hard enough to collide. 
• R.O.R  depends on= temp, conc, catalyst, size of lumps/ S.A 
• More collisions increases rate of reaction 
• Temperature= particles are faster = collide more frequently and more energy they have so more collisions will have enough energy to react. 
• Concentration= more particles in same amount of volume, makes collisions between reactant particles more likely
• Smaller solids/ more SA = breaking it up will increase S.A, means more particles around it will have more area to work on, collisions more frequently -e.g. soluble painkillers 
• Catalyst- increases speed of reaction without being used up, works by giving reacting particles on surface to stick to where they can bump into eachother= more successful collisions. 

• R.O.R= amount of reactant used or amount of product formed/ time 
• 1) precipitation and colour change- this is when the product of the reaction is coloured or is a precipitate which clouds the solution, if a precipitate is formed you can observe a mark though the solution and measure how long it takes for it to disappear- the faster the mark disappears, the quicker the reaction. Only works for reactions where the initial solution is see-through. If reactants are coloured and the products are colourless (or vice versa) you can time how long it takes for the solution to lose or gain its colour. The results are very subjective- different people might not agree over the exact point when the mark 'disappears" or changes colour. 
• 2) change in mass- for measuring the speed of reactions that produce gases, can be carried out using a mass balance. As the gas is released, the mass is disappearing is easily measured on the balance. The quicker the reading drops, the faster the reactipn. R.O.R graphs are particularly easy to plot. This is the most accurate of the three methods, but the gas is released in the room.
• 3) Volume of gas given off- involves a gas syringe to measure the volume of gas given off. The more gas given off during a given time interval, the faster the reaction. A graph of gas volume against time could be plotted to give a r.o.r graph. Gas syringes usually give volumes accurate to the nearest cm3. Too vigorous= easily blow the plunger out of the syringe.