GCSE AQA Chemistry Unit 3

History of the Periodic Table, The Modern Periodic Table, Group 1-The Alkali Metals, Group 7-The Halogens, Transition Elements,Acids and Alkalis, Acids, Alkalis and Titrations, Titration Caluclations,Water, Solubility, Hard Water, Water Quality, Energy, Energy and Fuels, Bond Energies, Energy and Food, Tests for Cautions, Tests for Anions, Tests for Organic Compounds, Instrumental Methods, Identifying Unkown Substances.

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The History of the Periodic Table

  • Until recently you could only categorise elements using:
    • Their chemical properties,
    • Thier relative atomic mass.
  • Newlands Law of Octaves - 1864
    • Order of atomic mass,
    • Noticed every 8th elemtns had a simular property,
    • He set up his table so:
      • Groups of elements had different properties,
      • Mixed metals and non-metals,
      • Didn't leave any gaps.
  • Demitri Mendeleev - 1869
    • Ordered in atomic mass,
    • He left gaps so elements with simular properties were in columns,
    • When elements were discovered they fitted in.
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The Modern Periodic Table

  • Use Periodic Table to work out:
    • The detalied arrangement of electrons in an atom of any element.
    • Predict the chemical properties.
  • Electrons have shells which correspond to energy levels:
    • Energy Level 1 = 2 electrons
    • Energy ELvel 2 = 8 electrons
    • Energy leavel 3= 18 electrons
    • Formula = 2xn(squared)
  • Positive charge of nucleaus attracts electrons
  • Inner electrons get in the way of nuclear charge, reducing attraction = shielding
  • Increased distance + increased shielding = electrons lost easily + harder to gain electrons.
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Group 1 - The Alkali Metals

Li, Na, K, Rb, Cs, Fr

  • As you go down the column:
    • Bigger atoms,
    • More reactive,
    • Higher density,
    • Lower boiling point, 
    • Lower melting point,
    • One outer electron,
    • From 1+ Ions,
    • Form only Ionic Compounds
    • Reacts violently with water, (K ignites)
    • When reacted with water forms a hydroxide solution and hydrogen.
      • e.g 2K + 2H2O ---> 2KOH + H2
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Group 7 - The Halogens

F, Cl, Br, I, At

  • As you go down the column:
    • Less reactive,
    • Higher melting points,
    • Higher boiling points,
    • All non-metals,
    • All produce coloured vapours,
    • For Molecules (pairs of atoms) e.g. Cl2, Br2
    • From Ionic and Covelant Compounds,
    • React with metals to form salts 
      • e.g. 2Fe + 3Br2 ---> 2FeBr3
    • More reactive halogens will displace less reactive ones.
      • e.g. Cl2 + 2KI ---> I2 + 2KCl
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The Transition Elements

  • Found bewteen Groups 2 and 3,
  • Good condictors of heat and electricity,
  • Dense, Shiny and Strong,
  • Less reactive than Group 1,
  • Don't react very much with water and oxygen,
  • Denser, Stronger and Harder than Group 1 elements,
  • Often have more than one ion e.g. Cu+, Cu+2
  • Compounds are very colourful,
  • Make good catalysts,
  • e.g. Iron in the Harber Process
  • Properties due to the way their electron shells fill,
    • As you get further away from the nucleus the energy levels start to overlap bewteen the 3rd anf 4th energy levels
    • The transition elements put their electrons into the overlapping 3rd energy level until it's full.
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Acids and Alkalis

  • Acids (molecules dissociate by) release H+ ions in water, 
  • Acids are proton donors,
  • Alkalis release OH- ions in water,
  • Alkalis are proton acceptors,
  • Protons are hydrated in water:
    • In Acidic Solutions:
      • H+ ions are surrounded by water (H+ aq)
    • In Basic Solutions:
      • Water Molecules can dissociate into H+ and OH-,
      • Ammonium (NH3) can take H+ leaving OH- behind,
      •  KOH releases hydroxide ions straight into the water,
  • Strong Acids ionise compleatly in water - release all H+
  • Weak Acids ionise slightly in water - release some H+
  • pH = concentration of H+ ions in a solution,
  • Ammonia = weak alkali,
  • NaOH and KOH = strong alkali,
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Titration

  • Titrations allow you to find out excatly how much acid is needed to neutralise a quanity of alkali (or vice versa)
  • Indicators:
    • Phenolphthalein = Weak Acids (colourless) and Strong Alkalis (pink)
    • Methyl Orange  = Strong Acids and Weak Alkali
  • Method:
    • Put alkali in flask with indicator
    • Add acid a bit at a time using a burette 
    • Giving regular swirls
    • When the indicator changes colour stop
    • Record the amount of acid to neutralise
    • Repeat
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Titration Calculations

  • Calculate concentration from results of titration experiemnt.
  • Concentration in Moles per dm3
    • How many moles are there in the known substance?
      • No. of moles = concentration x volume
    • Write out a balanced equation for the reaction.
    • How many moles of the unknown are there?
      • Moles of known / moles to neutralise one mole of unknown
    • Work out the concentration of unknown.
      • Concentration = No. of moles / Volume 
  • Concentration in Grams per dm3
    • Work out the relative formula of unknown concentration.
    • Convert concentration of moles (see above) into concentration in grams.
      • Mass in Grams = Concentra. mol/dm3 x Relative formula mass
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The Water Cycle

  • Sun evaporates water from the sea,
  • Water vapour carried up as warm air rises,
  • Vapour cools (due to genral cooling in the lower part of the atmosphere at high altitudes),
  • Condenses, 
  • Forms clouds,
  • Droplets get to heavy,
  • Rain, rain, rain,
  • Water runs back to the sea,
  • At some stage iw will come into contact with rocks (on or under the ground) meaning water in deffierent places will dissolve different minerals
  • The cycle starts over again.
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Water

  • Water - Universal Solvent,
  • Water disolves most ionic compounds:
    • Salts of:
      • Sodium,
      • Potssium,
      • Ammonium,
    • Nitrates,
    • Chlorides - expcept silver and lead
    • Sulphates - barium and lead
  • Water Molecules are Polar:
    • Positive hydrogen side attracts negative ions,
    • Negative side attracts positive ions,
  • Subsatnce with small molecules are soluble in water e.g. CO2, Cl2, SO2
  • Sulfure dioxide dissolves to form an acid which falls as acid rain
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Solubility

  • The solubility of a substance in a given solvent is the number of grams of the solute (usually a solid) that dissolve in 100g of the solvent (liquid) at a particualar temperature.
  • The solubility of (solid) solutes usually increases with temperature.
  • A satureted solution is one that cannot hold any more solid at that temperature - and you have to be able to see solid on the bottom  to be certain that it's saturated.
  • Solibility Curves:
    • Plots the mass of solute dissolves in a satureated solution at var ious temperatures.
    • cooling a saturated solution will cause some solid to crystalise out.
    • Finding out how much will crystalise:
      • The grams dissolved at the particular temperature taken awayfrom eachother.
  • The amount of gas that dissolves dependes on the pressureabove the gas. Higher pressure = more gass dissolving.
  • Fizzy drinks = CO2 dissolved in water (carbonated.) Open lid - fizzes out.
  • Fish ect. need dissolved oxygen - O2 levels decrease due to pollution.
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Hard Water

  • Instead of forming lather with soap you get scum due to the 'hardness minerals' in the hard water reacting with soap.
  • Scale (calicum carbonate) can block the inside of pipes, boiler and kettles. - Scale is a thermal insulator - makes boiling a kettle less efficant.
  • Hardness in caused by Calicum (Ca+2) and Magnesium (Mg+2) ions.
  • Get hard water in certain areas due to the types of rocks there.
  • Ca+2 ions are good for healthy teeth and bones,
  • Sacle forms a protective layer in pipes - stops posionoius metals and protects pipes from rust.
  • Removing Hardness
    • Adding sodium carbonate - carbonate ions join onto Ca or Mg ions and male an insoluble precipate.
    • Ion exchange - water is fed through a ion exhange column where sodium or hydrogen ions replace Ca or Mg.
    • Calicum Carbonate is acid - so descaling products conatin acid.
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Water Quality

  • Water treatment:
    • From resevoir
    • Passes throigh mesh screen to remove big bits e.g. twigs
    • Ozone or cholrine to kill mocrooranisms
    • Chemicals added to make solidsand microoranisms stick together ad fall to the bottom.
      • Iron = remove dissolved phosphates,
      • Bacteris = Nitrates.
    • Filtered through gravel bed to remove all solids
    • Passes through activated carbon to remove bad tates and odors.
    • pH is corrected
    • Water cholrinated to kill an microorganisms left.
  • At home you can use filter containing carbon (remove chlorine taste)and silver (remvoe bugs)
  • Totally pure water = disstilation (to expensive for mass production)
  • The World Health Organisation = 1995 1 billion people - no clean water.
  • Expensive for clean water in some countrys
  • Some puriftying process can damage the enviroment.
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Energy

  • An exothermic reaction is one which gives out energy to the surroundings, using on the form of heat and shown by a rise in temperature.
    • The energy released in the formation of bonds is greater than the energy used to break old bonds.
  • An endothermic reaction is onw which takes in energy from the surroundings, usually in the form of heat and shown by a fall in temperature.
    • The energy required to break old bonds is is greater than the energy released when new bonds are formed.

 Energy transfer can be measured by:

  • Taking the temperature of the reagents, mixing them in a polystryene cup, measuring the temperature at the end of the reaction. 
  • Problem = heat lost to the surroundings.
  • Solution = Lid and cotton wool to give insulation.
  • Works for solid in water and neutalisiation reactions.
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Energy and Fuels

  • Burn fuels to release energy = exothermic
  • Calculated using calometry
    • To  meaure the amount of energy produced accuratly = A Bomb Calorimeter carefully controlls the conditions and meaures the energy changes in a reaction through the change in temperature.
  • Working out Energy per Gram of a fuel:
    • Calculate Temperature Change:
      • Temp Chage = end temp - start temp
    • Calculate Energy Change:
      • 4.2J of energy raises temp of 1g of solution by 1 degree C
      • Energy Change = 4.2 x Temp Change x Volume of water.
      • Measured in kJ (Joules x 1000)
    • Moles:
      • Multilply or divide to get to 1 mole e.g. 0.1 x 10 = 1 Mole
        • Energy change x 10 = Energy change in kJ/mol.
  • Buring fuels effects the enviroment and the economy,
    • Crude oil is expensive and burning fossil fuels causes global warming.
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  • ENDOTHERMIC REACTION
    • Takes in Energy
    • Bond Breaking
    • ^H = Positive (more energy absorbed)
    • Reactants Low - Products High
  • EXOTHERMIC REACTION
    • Releases Energy
    • Bond Making
    • ^H Negative (more energy released)
    • Reactant High -Products Low
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Bond Energies

  • Bond Energy Diagrams show if it an endo- or exo- thermic reaction.
  • Endo = Low reactant to High Products
  • Exo - High reactants to Low products
  • Difference in heigh represents the energy taken in or rleased during the reaction.
  • Activation Energy is the inital rise that represents the minimum energy needed by reacting particles for reaction to occur.
    • Can be lowered by a catalyst. (Lower curve on a diagram)
    • Overall ^H energy change remains the same.
  • Bond Energy Calculations
    • Every chemical bond has a bond energy associated with it.
    • Bond energies calculate the overall energy change.
    • ^H measured in kJ/mol .
    • Negative = Endothermic (more energy released than absorbed)
    • Positive = Exothermic (more energy absorbed than released)
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Energy and Food

  • 1 calorie = amount of energy needed to raise the temp of 1g of water by 1 degree C = 4.2 Joules
  • The dietry information on food lables is in Kilocalories (Calories)
    • 1 Calorie = amount of energy needed to raise temp of 1 Kilogram of water by 1 degrees C = 4200 joules.
  • Different foods produce different amount of energy:
    • Fats and Oil = Large amounts of energy.
    • Carbohydrates = Some energy (much less than fats and oils)
    • Proteins = Same as carbohydrates (not used in bodys for energy)
  • Chemical reactions in your cells need energy all the time.
  • Glucose + Oxygen --> Carbon dioxide + Water + ENERGY
  • More energy than needed = stored as fat
  • Less energy than needed = uses up fat (Calorie-control diets - Low fat diets)
  • Sugar is high in energy and stimulates an apetite.
  • Slow-release carbohudrates fill you up for longer.
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Tests for Cations - Positive Ions

  • Flame Tests:
    • Lithium = Crimson-Red (I Like Raspberry CRumble)
    • Sodium = Yellow-Orange  (The Sun is Yellow-Orange)
    • Potassium = Lilac   (Pots conatin Lilac flowers)
    • Calicum = Brick-Red   (The Cows come home to a Red-Brick House)
    • Barium = Green   (Bears eat Green apples)
  • Precipiate Tests:
    • Calicum = White (Calicum is good for bones and bones are White)
    • Copper = Blue (Copper's have Blue flashing lights on their cars)jie
    • Iron(II) = Sludgey Green (Ironing 2 Save Grandma)
    • Iron(III) = Reddish Brown(I TRIck Rogue Badger's)
    • Alluminum = White but redissolves in excess NaOH and colourless.
    • Magnesium = White
  • Ammonium Compound + NaOH ---> Ammonia (Stinks of Cat wee)
    • Turns damp Red litmus paper Blue.
    • Add Sodium Hydroxide if you smell wee amonia ions are present.
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Tests for Anions - Negative Ions

  • Test for Carbonates:
    • Test for CO2:
      • Acid + Carbonate ---> Salt + Water+ Carbon dioxide
      • Using Lime Water (goes cloudy if present)
    • Some carbonates change colour when they decompose:
      • Method:Spatula of carbonate in test tube heat and allow to cool.
      • Copper Carbonate turns from Green to Black and stays black when it cools.
      • Zinc Carbonate turns from White to Yellow. But when it cools it turns back to White

  • Tests for Sulfates (SO4 -2) and Halides (Cl-, Br-, I-)
    • Sulfate Ions, SO4 -2
      • Add dilute HCl followed by Barium Chloride Solution (BaCl2)
      • A White Precipitate of Barium Sulfate means the origional compound was a sulfate.
      • Ba(+2) +  SO4 (-2) ---> BaSo4 (<-- Solid)
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Tests for Anions Continued...

  •  
    • Chlorine, Bromine and Iodine Ions (Cl-, Br-, I-)
      • Add dilute nitric acids followed by silver nitrate solution.
      • A Chloride = White precipitate of Silver Chloride,
      • A Bromide = Cream precipitate of Silver Bromide,
      • A Iodide = Yellow Precipitate of Silver Iodide.
  • Tests for Nitrates (NO3):
    • Produces Ammonia,
    • Mix mystery compound with alumium powder and sodium hydroxide and heat.
    • Is the mystery compound was a nitrate then it will reduce to ammonia.
    • Test with your nose or damp red litmus paper which should turn blue.
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Tests for Organic Compounds

  • Heat/Burn Test:
    • Org Comp's burn in air to form a Yellow-Orange (Mainly) and/or Blue Flame.
    • Lots of air > Hydrocarbon = Carbon Dioxide + Water.
    • Not enough air . Hydrocarbon = Carbon Monoxide + Carbon (soot)
    • Solid Org Comp's will char - scorched with black marks.
  • Bromine Test:
    • Org Comp is unsaturated = Decolourise Bromine Water.
    • Org Comp is Saturated = Bromine Water stays brown.
    • Test on Margarine (will declourise as has C=C bonds)
  • Find the Empirical Formula of a Compound by Burning it:
    • Burn a known mass in Oxygen.
    • Meaure masses of all products.
    • Find the mass of each element in a compound.
    • Divide these masses by the atomic masses to find the no. of moles.
    • Put it into a ratio.
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Insturmental Methods

  • Machines Analyse Unknown Subsatances
    • Forensic's, Enviromental, Medical, Industry, Athletes Drug Tests,
    • Advantages:
      • More accurate than lab methods,
      • Can detect even the tinitest amount of a substance,
      • Much faster than lab methods,
    • Disadvantages:
      • Expensive to buy, run and maintain,
      • Takes special training to use,
      • Gives results that can often only be interpreted only by comparison with alreadyknown specimens.
  • Atomic Absorbtion Spectroscopy (AAS)
    • Meaures the concentration of a metal in a liquid sample.
    • The patterns of light absorbed by the metals in a sample are analysed - each metal sample produces  a different pattern
    • Method: Liquid sample fed into a flame where it vapourises,
    • Wavelenght to be studied is selected by light,
    • Light falls on a detector - electric current depends on light intensity.
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Instrumenal Methods Continued...

  • Infrared (IR) Spectrometry:
    • Identifies which wavelengths of infared radiation are being absorbed.
    • Pattern of absorbance is unique for every compound.
  • Ultraviolet (UV) Spectroscopy:
    • Indentifies which wavelenghts of UV light are being absorbed
  • Nuclear Magnetic Resonance (NMR) Spectroscopy:
    • Used for Organic Compounds, Shows what atoms hydrogen atoms are connected to from which you find the structure of the molecule.
  • Gas-Liquid Chromotography:
    • Simular principle to paper chromotography, Identifies gas and liquids.
  • Mass Spectrometry:
    • Used for both elements and compounds. Identifies mass of each molecule/paticle. For elements tell you which excatly.
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Chemistry 3...

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Comments

amy

on slide 17 - negative means that it is exothermic because more energy is released rather than being absorbed

positive means that the reaction is endothermic because more energy is absorbed rather than being released

AlmightyBurge

This would be by far the best resource if above was changed and SPELLING/MISTAKES were checked and changed

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