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Types of Chemical Reactions

Types of Chemical Reactions

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Types of Chemical Reactions

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  1. Types of Chemical Reactions Writing Chemical Reactions

  2. Types of Reactions • Many chemical reactions have defining characteristics which allow them to be classified as to type.

  3. Types of Chemical Reactions • The five types of chemical reactions in this unit are: • Combination • Decomposition • Single Replacement • Double Replacement • Combustion

  4. Combination Reactions • Two or more substances combine to form one substance. • The general form is A + X AX • Example: • Magnesium + oxygen  magnesium oxide • 2Mg + O2 2MgO

  5. Combination Reactions • Combination reactions may also be called composition or synthesis reactions. • Some types of combination reactions: • Combination of elements • K + Cl2 • One product will be formed

  6. Combination Reactions • K + Cl2 • Write the ions: K+ Cl- • Balance the charges: KCl • Balance the equation: 2K + Cl2  2KCl

  7. Combination Reactions • Some types of combination reactions: • Oxide + water  • Nonmetal oxide + water  acid • SO2 + H2O  H2SO3 • Metal oxide + water  base • BaO + H2O  Ba(OH)2

  8. Combination Reactions • Some types of combination reactions: • Metal oxides + nonmetal oxides • Na2O + CO2 Na2CO3 • CaO + SO2  CaSO3

  9. Predicting products for synthesis reaction • Write down the reactants that are involved in the reactions on the reactant side • of the equations. • Knowing that a synthesis reaction is going to occur, put the reactants together • on the product side of the equation. • When elements make compounds remember that there is a certin ratio • which they must combine in. YOU MUST REMEMBER HOW TO WRITE • CHEMICAL FORMULAS. Lets continue looking at the potassium and chlorine reaction.

  10. Combination Reactions • K + Cl2 • Write down the formula for potassium chloride. • KCl • Notice you are missing a Cl atom on the product side. Where did that Chlorine atom go? No where • Balance the equation: 2K + Cl2  2KCl

  11. One more thing to remember when writing equations • Diatomic elements • Elements that never travel alone. • Last example of potassium and chlorine • Reactants K + Cl2 ProductKCl Notice the Chlorine is never by itself this is a diatomic element How do you remember the diatomic elements?

  12. H.O.Br.F.I.N.Cl H – hydrogen O – oxygen Br – Bromine F – Fluorine I – Iodine N – Nitrogen Cl - Chlorine

  13. Practice Predict the following products. Write a balanced chemical equation for each of the following Copper + oxygen 2Cu + O2 2CuO Silver + Sulphur 2Ag + S Ag2S Iron + Oxygen Fe3O2 3Fe + O2

  14. Decomposition Reactions • One substance reacts to form two or more substances. • The general form is AX  A + X • Example: • Water can be decomposed by electrolysis. • 2H2O  2H2 + O2

  15. Decomposition Reactions • Types of Decomposition Reactions: • Decomposition of carbonates • When heated, some carbonates break down to form an oxide and carbon dioxide. • CaCO3 CaO + CO2 • H2CO3  H2O + CO2

  16. Decomposition Reactions • Types of decomposition reactions: • Some metal hydroxides decompose into oxides and water when heated. • Ca(OH)2 CaO + H2O Note that this is the reverse of a similar combination reaction.

  17. Decomposition Reactions • Types of decomposition reactions: • Metal chlorates decompose into chlorides and oxygen when heated. • 2KClO3 2KCl + 3O2 • Zn(ClO3)2  ZnCl2 + 3O2 • Some of these reactions are used in explosives.

  18. Decomposition Reactions • Some substances can easily decompose: • Ammonium hydroxide is actually ammonia gas dissolved in water. • NH4OH  NH3 + H2O • Some acids decompose into water and an oxide. • H2SO3 H2O + SO2

  19. Summary synthesis V.s Decomposition • Synthesis • Non-metal oxide + H2O Acid • CO2 + H2O HCO3 • Metal oxide + water Base BaO + H2O Ba(OH)2 • Metal oxide + non-metal oxide polyatomic compound • CaO + SO2 CaSO3 • Decomposition • Acid decomposition • HCO3 CO2 + H20 • Hydroxide decomposition • Ba(OH)2 BaO + H2O • Polyatomic compound decomposition • CaSO3 CaO + SO2

  20. Decomposition Reactions • Some decomposition reactions are difficult to predict. • The decomposition of nitrogen triiodide, NI3, is an example of an interesting decomposition reaction.

  21. Single Replacement Reactions • A metal will replace a metal ion in a compound. • The general form is A + BX  AX + B • A nonmetal will replace a nonmetal ion in a compound. • The general form is Y + BX  BY + X

  22. Single Replacement Reactions • Examples: • Ni + AgNO3 • Nickel replaces the metallic ion Ag+. • The silver becomes free silver and the nickel becomes the nickel(II) ion. • Ni + AgNO3 Ag + Ni(NO3)2 • Balance the equation: • Ni + 2AgNO3  2Ag + Ni(NO3)

  23. Single Replacement Reactions • Not all single replacement reactions that can be written actually happen. • The metal must be more active than the metal ion. • Aluminum is more active than iron in Al + Fe2O3 in the following reaction:

  24. Thermite Reaction

  25. Thermite Reaction • Al + Fe2O3 • Aluminum will replace iron(III) as was seen in the video. • Iron (III) becomes Fe and aluminum metal becomes Al3+. • 2Al + Fe2O3 2Fe + Al2O3

  26. Single Replacement Reactions • An active nonmetal can replace a less active nonmetal. • The halogen (F2, Cl2, Br2, I2) reactions are good examples. • F2 is the most active and I2 is the least. • Cl2 +2 NaI  2 NaCl + I2

  27. Double Replacement Reactions • Ions of two compounds exchange places with each other. • The general form is AX + BY  AY + BX • Metathesis is an alternate name for double replacement reactions.

  28. Double Replacement • NaOH + CuSO4 • The Na+ and Cu2+ switch places. • Na+ combines with SO42- to form Na2SO4. • Cu2+ combines with OH- to form Cu(OH)2 • NaOH + CuSO4  Na2SO4 + Cu(OH)2 • 2NaOH + CuSO4  Na2SO4 + Cu(OH)2

  29. Double Replacement • CuSO4 + Na2CO3 • Cu2+ combines with CO32- to form CuCO3. • Na+ combines with SO42- to form Na2SO4. • CuSO4 + Na2CO3  CuCO3 + Na2SO4

  30. Double Replacement • Na2CO3 + HCl  • Na+ combines with Cl- to form NaCl. • H+ combines with CO32- to form H2CO3. • Na2CO3 + 2HCl  2NaCl + H2CO3 • H2CO3 breaks up into H2O and CO2.

  31. Double Replacement • The gas formed was carbon dioxide. • The final balanced reaction is: Na2CO3 + HCl  NaCl + H2O + CO2. • Balance the equation. • Na2CO3 + 2HCl  2NaCl + H2O + CO2

  32. Combustion Reaction • When a substance combines with oxygen, a combustion reaction results. • The combustion reaction may also be an example of an earlier type such as 2Mg + O2 2MgO. • The combustion reaction may be burning of a fuel.

  33. Combustion Reaction • Methane, CH4, is natural gas. • When hydrocarbon compounds are burned in oxygen, the products are water and carbon dioxide. • CH4 + O2 CO2 + H2O • CH4 + 2O2  CO2 + 2H2O

  34. Combustion Reactions • Combustion reactions involve light and heat energy released. • Natural gas, propane, gasoline, etc. are burned to produce heat energy. • Most of these organic reactions produce water and carbon dioxide.

  35. Practice • Classify each of the following as to type: • H2 + Cl2 2HCl • Combination • Ca + 2H2O  Ca(OH)2 + H2 • Single replacement

  36. Practice • 2CO + O2 2CO2 • Combination and combustion • 2KClO3  2KCl + 3O2 • Decomposition

  37. Practice • FeS + 2HCl  FeCl2 + H2S • Double replacement • Zn + HCl  ? • Single replacement • Zn + 2HCl  ZnCl2 + H2