Chemical Changes


Acids and Bases

pH scale 0 - 14

  • lower= more acidic
  • higher=more alkaline
  • neutral=pH 7
  • indicator is a dye that changes colour if above or below a certain pH
  • acids and bases neutralise each other
  • acid= h+ ions in water
  • base=pH greater than 7
  • alkali is base that dissolves in water form OH- ions in water
  • neutralisation: acid+base=salt+water
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  • Used to find out concentration + whats needed to neutarlise it
  • pipette and pipette filler
  • add alkali to conical flask + indicator
  • funnel to fill burette with acid (known concentration)
  • wear safety glasses
  • fill below eye level
  • record intitial volume
  • use burette to add acid slowly
  • until end-point (colour change)
  • indicator changes colour when all alkali is neutralised
  • record final volume
  • to increase accuracy look for anomalous results
  • first titration= rough to get approximate idea
  • repeat more accurately 
  • calculate mean
  • single indicators:
  • (blue in alkalis and red in acids)
  • phenolphthalein (pink in alkalis colourless in acids)
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Strong and Weak Acids

  • Acids produce protons in water
  • Strong acids fully ionise in water = fast rate of reaction so more reactive
  • weak do not fully ionise = reversible reaction = equilibrium to the left
  • pH (the measure of the conc of hydrogen ions)
  • decrease of 1 on pH scale h+ ions x10
  • acid strength tells you proportion of acid molecules ionised in water
  • concentration= how much acid in water
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Reactions of Acids

  • Metal oxides and hydroxides= bases (react with acids to form salt and water)
  • if they dissolve in water=alkalis
  • Acid + metal carbonate = salt + water + carbon dioxide
  • insoluble salts can be mad using an insoluble base:
  • warm dilute acid
  • add insoluble base to acid (in excess)
  • filter out excess to get salt solution
  • heat solution and then leave to form crystals (crystallisation)
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Reactivity Series

  • metals- how easily they loose an electron (form postive ions)
  • most reactive:
  • potassium - K
  • sodium - Na
  • lithium - Li
  • calcium - Ca
  • magnesium - Mg
  • carbon - C (non-metal)
  • zinc - Zn 
  • iron - Fe
  • Hydrogen - H (non-metal)
  • Copper - Cu 
  • acid + metal = salt + hydrogen 
  • speed of reaction indicated by bubbles given off
  • reactivity= how fast reaction is
  • very reactive = explosive
  • metal + water = metal hydroxide + hydrogen (less reactive won't react with water)
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Seperating Metals from Metal Oxide

  • Reacting with oxgen forms an oxide (example of oxidation)
  • oxides are often ores that metals need to be extracted from
  • metal - oxide = reduction reaction
  • some can be reduced with carbon
  • anything higher than carbon can not be reduced by it (have to use electrolysis)
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Redox Reactions

  • if electrons are transferred = redox
  • Oxidation
  • Is
  • Loss
  • Reduction
  • Is
  • Gain (terms of electrons)
  • REDuction + OXidation = redox reaction (happen at same time)
  • displacement = example of a redox reaction
  • (reactive metal "kicks out" less reactive metal)
  • metal ion gains electrons
  • meatl atom loses electrons 
  • use word symbol equations to show this
  • ionic equations = useful parts of reactions
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  • means splitting up with electricity 
  • electric current passed throuh molten or dissolved ionic compound
  • + ions = cathode (negative electrode) and gain electrons = reuced
  • - ions = anode (positive electrode) and losse electrons = oxidised
  • when loose/gain form uncharged element and then discharged from electrolyte
  • molten ionic compounds can be electrolysed ions free to move + carry current
  • metals being ectraced from their ores is very expensive
  • need a lot of energy to melt ore and create current
  • Aluminium exrtracted from ore bauxite (contains aluminium oxide)
  • cryolite added to lower melting point
  • molten mixture= free ions =conduct electricity
  • Al3+ negative electrode + pick up +3 electrons = aluminium atoms
  • O2 goes to positive elctrode -2 electrons = combine to form O2
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Electrolysis of Aqueous Solutions

  • will still have hydrogen and hydroxide ions
  • cathode= h+ ions and metal ions hydrogen gas produced if more reactive than hydrogen
  • if less reactive a solid layer ofp pure metal produced
  • at anode if OH- and halide ions (Cl- Br- I-) molecules of cholorine, bromine or iodine formed
  • no halide ions= oxygen formed
  • cooper sulfate (II) contains Cu2+ SO (2-, 4) H+ and OH- 
  • copper metal produced and coats electrode
  • no halide ions so anode oxygen and water produced
  • chlorine bleaches litmus paper turning it white
  • hyrdogen makes a squeaky pop with lighted splint
  • oxygen relights a glowing splint
  • for half equations make sure elctrons are balanced
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