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Synthesis Reactions

Synthesis Reactions. Type 1: A metal combines with a nonmetal to form a binary salt. Example: A piece of lithium metal is dropped into a container of nitrogen gas. 6Li (s) + N 2 (g) → 2Li 3 N ( aq ). Synthesis Reactions. Type 2: Metallic oxides and water form bases (metallic hydroxides).

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Synthesis Reactions

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  1. Synthesis Reactions Type 1: A metal combines with a nonmetal to form a binary salt. Example: A piece of lithium metal is dropped into a container of nitrogen gas. 6Li (s) + N2 (g) → 2Li3N (aq)

  2. Synthesis Reactions Type 2: Metallic oxides and water form bases (metallic hydroxides). Example: Solid sodium oxide is added to water Na2O (s) + HOH (l) → 2NaOH (aq) Example: Solid magnesium oxide is added to water. MgO (s) + 2HOH (l) → Mg(OH)2 (aq)

  3. Synthesis Reactions Type 3: Nonmetallic oxides and water form acids. The nonmetal retains its oxidation number. Example: Carbon dioxide is bubbled in water. CO2 (g) + H2O (l) → H2CO3 (aq) Example: Dinitrogenpentoxide is bubbled in water. N2O5 (g) + H2O (l) → 2HNO3 (aq)

  4. Synthesis Reactions Type 4: Metallic oxides and nonmetallic oxides form salts. Example: Solid sodium oxide is added to carbon dioxide. Na2O (s) + CO2 (g) → Na2CO3 (aq) Example: Solid calcium oxide is added to sulfur trioxide. CaO (s) + SO3 (g) → CaSO4 (aq)

  5. Decomposition Reactions Type 1: Metallic carbonates decompose into metallic oxides and carbon dioxide Example: A sample of magnesium carbonate is heated. MgCO3 (s) → MgO (s) + CO2 (g)

  6. Decomposition Reactions Type 2: Metallic chlorates decompose into metallic chlorides and oxygen. Example: A sample of magnesium chlorate is heated. Mg(ClO3)2 (aq) → MgCl2 (aq) + 3O2 (g)

  7. Decomposition Reactions Type 3: Ammonium carbonate decomposes into ammonia, water, and carbon dioxide. Example: A sample of ammonium carbonate is heated. (NH4)2CO3 (aq)→ 2NH3 (g) + H2O (l) + CO2 (g)

  8. Decomposition Reactions Type 4: Sulfurous acid decomposes into sulfur dioxide and water. Example: A sample of ammonium carbonate is heated. H2SO3 (aq) → H2O (l) + SO2 (g)

  9. Decomposition Reactions Type 5: Carbonic acid decomposes into carbon dioxide and water. Example: A sample of carbonic acid is heated. H2CO3 (aq) → H2O (l) + CO2 (g)

  10. Decomposition Reactions Type 6: A binary compound may break down to produce two elements. Example: Molten sodium chloride is electrolyzed. 2NaCl(l) → 2Na (s) + Cl2 (g)

  11. Decomposition Reactions Type 7: Hydrogen peroxide decomposes into water and oxygen. Example: 2H2O2(aq) → 2H2O (l) + O2 (g)

  12. Decomposition Reactions Type 8: Ammonium hydroxide decomposes into ammonia and water. Example: NH4OH(aq) → NH3 (g) + HOH (l)

  13. Single Replacement Type 1: Active metals replace less active metals from their compounds in aqueous solution. Example: Magnesium turnings are added to a solution of iron (III) chloride. 3Mg (s) + 2FeCl3(aq)→ 2Fe (s) + 3MgCl2 (aq)

  14. Single Replacement Type 2: Active metals replace hydrogen in water. Example: Sodium is added to water. 2Na (s) + 2HOH(l)→ 2NaOH (aq) + H2 (g)

  15. Single Replacement Type 3: Active metals replace hydrogen in acids. Example: Lithium is added to hydrochloric acid. 2Li (s) + 2HCl(aq) → 2LiCl (aq) + H2 (g)

  16. Single Replacement Type 4: Active nonmetals replace less active nonmetals from their compounds in aqueous solution. Example: Chlorine gas is bubbled into a solution of potassium iodide. Cl2 (g) + 2KI(aq) → I2 (g) + 2KCl (aq)

  17. Single Replacement Type 5: If a less reactive element is combined with a more reactive element in compound form, their will be no resulting reaction. Example: Chlorine gas is bubbled into a solution of potassium iodide. Cl2 (g) + KF (aq) → No Reaction Example: Zinc is added to a solution of sodium chloride. Zn (s) + NaCl (aq) → No Reaction

  18. Double Replacement (Gas)

  19. Double Replacement (Gas) • Example 1: The reaction of Na2SO3 and HCl produces H2SO3: Na2SO3 (aq) + 2HCl (aq) → H2SO3 (aq) + 2NaCl (aq) • Bubbling is observed in this reaction because the H2SO3 (sulfurous acid) is unstable and immediately decomposes to give HOH and SO2 gas: H2SO3 (aq) → HOH (l) + SO2 (g) • The molecular equation for the overall or complete reaction, therefore, is: Na2SO3 (aq) + 2HCl (aq) → HOH (l) + SO2 (g) + 2NaCl (aq)

  20. Double Replacement (Gas) • Example 2: A typical reaction of a carbonate and an acid is: K2CO3 (aq) + 2HNO3 (aq) → HOH (l) + CO2 (g) + 2KNO3 (aq) • Bubbling is also observed in this reaction. Theoretically H2CO3, carbonic acid, is formed, but the acid is unstable and immediately decomposes to form carbon dioxide gas and water according to the following equation: H2CO3 (aq) → HOH (l) + CO2 (g)

  21. Double Replacement (Gas) • Example 3: Ammonium salts and soluble bases react as follows (particularly when the solution is warmed): NH4Cl (aq) + NaOH (aq) → NH3 (g) + HOH (l) + NaCl (aq) • The odor of ammonia gas is noted and moist blue litmus paper held near the mouth of the container will turn blue. Theoretically NH4OH, ammonium hydroxide, is produced (also known as ammonia water). The compound is unstable and decomposes into ammonia gas and water: NH4OH (aq) → NH3 (g) + HOH (l)

  22. Double Replacement (Gas) • Example 4: The odor of rotten eggs and bubbling are noted when an acid is added to a sulfide. A typical reaction producing hydrogen sulfide gas is: FeS (s) + 2HCl (aq) → FeCl2 (aq) + H2S (g)

  23. Acid-Base Reactions • Example 1: Hydrogen sulfide gas is bubbled through excess potassium hydroxide solution. H2S (g) + 2KOH (aq) → K2S (aq) + 2HOH (l)

  24. Acid-Base Reactions • If the base is in excess, all hydrogen ions will react with strong base to produce water. • Example 2: Dilute sulfuric acid is reacted with excess sodium hydroxide. H2SO4 (aq) + 2NaOH (aq) → Na2SO4 (aq) + 2HOH (l)

  25. Acid-Base Reactions • If however, this same reaction were described in terms of mixing equal numbers of moles, then the coefficients for both reactants would be one (the same number of H and OH must be given away). • Example 3: Equal number of moles of sulfuric acid and sodium hydroxide react. H2SO4 (aq) + NaOH (aq) → NaHSO4 (aq) + HOH (l)

  26. Acid-Base Reactions • As the following example demonstrates, it is important to take into account the quantity (concentration and amount) of each reactant. • Example 4: Equal volumes of 0.1 M phosphoric acid and 0.2 M sodium hydroxide are reacted together. H3PO4(aq) + 2NaOH (aq) → Na2HPO4(aq) + 2HOH (l)

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