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Synthesis

Classification of reactions - Synthesis - Decomposition - Neutralisation - Precipitation - Double decomposition - Reduction Oxidation. Synthesis. Synthesis is the making of a compound by the combination of simpler substances.

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Synthesis

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  1. Classification of reactions- Synthesis- Decomposition- Neutralisation- Precipitation- Double decomposition- Reduction Oxidation

  2. Synthesis • Synthesis is the making of a compound by the combination of simpler substances. • The term is also generally used to describe the making of a new substance. • Example: • Iron + sulphur  Iron sulphide • Fe + S  FeS

  3. Decomposition • The breaking down of one substance into simpler substances. • Example: • Calcium carbonate  calcium oxide + carbon dioxide • CaCO3  CaO + CO2 • This is often carried out by heating.

  4. Neutralisation • The reaction between an acid and a base to produce a neutral salt and water. • Sulphuric acid + sodium hydroxide  Sodium sulphate + water • H2SO4 + 2NaOH  Na2SO4 + 2H2O • Only the hydrogen ions from the acid combine with the hydroxide ions from the base. All of the other ions start and end the reaction in the same ionic form. They are called spectator ions (they “watch” the other ions reacting)

  5. Precipitation • This is the reaction between two solutions resulting in the formation of an insoluble substance, which consequently appears as a solid within the liquid. • It is apparent that a solid has been formed as the liquid becomes turbid (not wholly transparent) • Example: • Copper sulphate + sodium hydroxide  copper hydroxide + sodium sulphate • The copper hydroxide is insoluble and appears as a precipitate.

  6. Double decomposition • This is really a precipitation and is given the name double decomposition to indicate that the ions have “switched round”. In the previous example the hydroxide ions seem to go from the sodium to the copper and the sulphate ions seem to go from the copper to the sodium. • In reality ions in solution are not related to one another and so no such movement takes place. The collision of the copper and the hydroxide ions within the solution forms the insoluble substance, the remaining ions are simply spectators and named as such. To collect the sodium sulphate it would be necessary to remove the water

  7. Oxidation-Reduction: A Reaction • Oxidation: When a substances loses electrons. • Reduction: When a substance gains electrons. • Consider: • Ca(s) + 2H+(aq)  Ca2+(aq) + H2(g). • The neutral Ca(s) has lost two e- to 2 H+ to become Ca2+. • We say Ca has been oxidized to Ca2+ • At the same time 2 electrons are gained by 2 H+ to form H2 . • We say H+ is reduced to H2 .

  8. Redox Reaction with Air • Consider the reaction of Ca with O2: • 2Ca(s) + O2(g)  2CaO(s) • Ca is easily oxidized in air. • On the left there is shiny Ca metal. • On the right we see a white powder – Calcium oxide. • Again, Ca(s) loses electrons and is oxidized to Ca+2 • And the neutral O2 has gained electrons from the Ca to become O2- in CaO. • We say O2 has been reduced to O2-.

  9. Electron Transfer and Terminology • Lose electrons: Oxidation • Gain electrons: Reduction. GER Leo

  10. It Takes Two: Oxidation-Reduction • In all reduction-oxidation (redox) reactions, one species is reduced at the same time as another is oxidized. • Oxidizing Agent: • The species which causes oxidation is called the oxidizing agent. The substance which is oxidized loses electrons to the other. • The oxidizing agent is always reduced • Reducing Agent: • The species which causes reduction is called the reducing agent. • The substance which is reduces gains electrons from the other. • The Reducing agent is always oxidized

  11. Oxidation of Metals with Acids • It is common for metal to produce hydrogen gas when they react with acids. For example, the reaction between Mg and HCl: • Mg(s) + 2HCI(aq)  MgCl2(aq) + H2(g) . • In this reaction, Mg is oxidized and H in HCl is reduced. • Note the change in oxidation state for these species: • Mg0 Mg+2 in MgCl2 • & • H+ in HCl  H0 in H2

  12. Redox reaction with Acid • It is possible for metals to be oxidized with salt: • Fe(s) + Ni(NO3)2 (aq)  Fe(NO3)2 (aq) + Ni (s) . • Molecular Equation • The net ionic equation shows the redox chemistry well: • Fe(s) + Ni+2(aq) Fe2+(aq) + Ni (s) • Net ionic Equation • In this reaction iron has been oxidized to Fe2+ while the Ni+2 has been reduced to Ni0. • What determines whether the reaction occurs ?

  13. Competition For e- Transfer • Consider: Na, Mg, Al, • Metallic character decreases left to right. • Metal tend to give up electrons. • Now consider the reaction: • Na + AlCl3 ??? (NaCl + Al) • To determine if the reaction occurs, the question is to determine which metal has a greater affinity for electrons (or which is willing to lose e- ). • Na is more willing to lose e- than Al • Al is more willing to accept e- (less metallic) • Conclude: The reaction occurs. • 3Na + AlCl3 3NaCl + Al

  14. Summary • Redox - Oxidation/Reduction reaction • Oxidation-Lose electron (LEO) • Reduction-Gain electron (GER) • Activity Series-Table showing elements’ relative ease of oxidation. • Active MetalMetal which prefers to lose e- and therefore prefers the oxidized form. • Noble MetalMetal which do not lose e- and therefore prefers the zero state.

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