chemical changes

• metal oxides


  • metals + oxygen -> known as metal oxides
  • known as oxidation reactions because the metals gain oxygen
  • reduction = loss of oxygen
  • oxydation = gain of oxygen
• reactivity series


  • when metals react with other substances, metal atoms form positive ions.
  • reactivity is related to its tendency to form positive ions.
  • metals can be arranged in order od reactivity in a reactivity series.
  • metals: potassium, sodium, lithium, calcium, magnesium, zinc, iron, and copper and be put in order of their reactivity from theirr reactions with water and dilute acids.
• reactivity series


  • non-metals: hydrogen, and carbon, are usually included in the reactivity series

  a more reactive metal can displace a less reactive metal drom a compound
• extraction of metals and reduction 


  • since gold is very unreactive it is found in the earth as the pure metal.
  • most metals are dound as compounds that require chemical reactions to extract the metal.
  • metals less reactive than carbon can be extracted by reduction with carbon
• oxidation and reduction in terms of electrons 


  Oxidation

  Is

  Loss

  Reduction

  Is

  Gain
• reactions of acids with metals


  • acids react with some metals to produce a salt and a hydrogen:
  • acid = metal -> salt + hydrogen
  • these are redox reactions
• neutralisation of acids and salt production 


  acids are neutralised by alkalis and base to produce salts and water

  • acid + alkali -> salt + water
  • acid + base -> salt + water

  acids are neutralised by metal carbonated to produce salts, water, and carbon dioxide.

  • acid​ ​+​ ​metal​ ​carbonate​ ​->​ ​salt​ ​+​ ​water​ ​+​ ​carbon​ ​dioxide
• soluble salts 


  can be made from acids by reacting them with insoluble substances such as metals, metal oxides, hydroxides, or carbonates.

  1. Add​ ​the​ ​chosen​ ​solid​ ​insoluble​ ​substance​ ​to​ ​the​ ​acid​ ​then​ ​the​ ​solid​ ​will​ ​dissolve.
  2. You​ ​know​ ​the​ ​acid​ ​has​ ​been​ ​neutralised​ ​when​ ​excess​ ​solid​ ​sinks​ ​to​ ​the​ ​bottom, so​ ​keep​ ​adding​ ​until​ ​this​ ​happens.
  3. Filter​ ​out​ ​excess​ ​solid​ ​leaving​ ​the​ ​salt​ ​solution,​ ​then​ ​evaporate​ ​some​ ​water,​ ​then leave​ ​the​ ​rest​ ​to​ ​evaporate​ ​slowly.

  this is called crystallisation
• the pH scale and neutralisation 


  • acids produce H+ ions in aqueous solutions
  • alkalis produce OH- ions in aqueous solutions
  • the pH scale mesures the acidity or alkalinity of a solution, and can be measured usinf universal indicator.
  • pH​ ​7​ ​is​ ​neutral
  • pH​ ​<​ ​7​ ​is​ ​acidic
  • pH​ ​>​ ​7​ ​is​ ​alkaline   
• titrations 


  1. Wash​ ​burette​ ​using​ ​dilute​ ​hydrochloric​ ​acid​ ​and​ ​then​ ​water
  2. Fill​ ​burette​ ​to​ ​100cm3​ ​​​with​ ​acid​ ​with​ ​the​ ​meniscus’​ ​base​ ​on​ ​the​ ​100cm3​ ​​​line
  3. Use 25cm3​ pipette to add 25cm3​ of alkali into a conical flask, drawing alkali into the pipette​ ​using​ ​a​ ​pipette​ ​filler
  4. Add a few drops of a suitable indicator to the conical flask
  5. Add​ ​acid​ ​from​ ​burette​ ​to​ ​alkali​ ​until​ ​end-point​ ​is​ ​reached​ ​
  6. The titre is the differencebetween​ ​the​ ​first​ ​​and​ ​second​ ​readings​ ​on​ ​the​ ​burette
• titrations calculations


  1. convert volume into dm3
  2. work out the moles of NaOH
  3. work out mole ratio from equation
  4. work out concentration
• strong and weak acids 


  • strong acid = completely ionised in aquoeus solution
  • e.g. hydrochloric, nitric, sufuric acids.
  • weak acid = partially inoised in aqueous solution
  • ethanolic, citric, carbonic acids
  • the stronger the acid, the lower the pH.
• process of electrolysis 


  • when an ionic substance is melted or dissolved, the ions are free to move about within the liquid or solution.
  • passing a current through substances that are molten or solution means that the solution can be broken down into elements. the substance being broken down is electrolyte.
  • positively charged ions move to the negative electrode (cathode)
  • negatively charged ions move to the positive electrode (anode)
  • ions are discharged at the electrodes.
• electrolysis of molten ionic compounds


  • when a simple ionic compound is electrolysed in the molten state, the metal is produced at the cathode and the non-metal is proudced at the anode.
  • this is because the metal is the positive ion and the non-metal is the negative ion.
• using electrolysis to extract metals


  • metals that are more reactive than carbon, are extracted by electrolysis of molten compounds.
  • large amounts of energy are used in the extraction process.
  • metals that react by carbon can be extracted by electrolysis as well.
• electrolysis of aqueous solutions


  • the ions discharged when an queous solution is electrolysed using electrodes depend on the reactivity of elements involved.
  • at the negative electrode hydrogen is produced unless the metal is less reactive than hydrogen.
  • at the positive electrold, if OH- and halide ions are present, then one of the halide ions will be produced
  • in the aqueous solution, walter molecules break down H+ ions and OH- ions that are discharged
• reactions at electrons as half equations


  • this is an example of a half equation; the small number is always the same as the two larger numbers within the equation.

  2CL$^-$ -> CL$^2$ + 2e$^-$

  2CL$^-$ - 2e$^-$ -> CL$^2$