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Types of Chemical Reactions

Types of Chemical Reactions. Chemical Reactions. If you add two substances together sometimes you get a chemical reaction. This is when a new substance is made. It is very difficult to reverse the reaction and get the original substances back.

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Types of Chemical Reactions

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  1. Types of Chemical Reactions

  2. Chemical Reactions If you add two substances together sometimes you get a chemical reaction. This is when a new substance is made. It is very difficult to reverse the reaction and get the original substances back. Signs that a chemical reaction has taken place include a change in colour, heat being given off or bubbles of gas being made.

  3. Adding acid to alkali Fireworks Driving a car    Boiling the kettle frying an egg Ice cream melting  x x Dissolving sugar in tea Discuss in pairs which of these are chemical reactions Burning a candle Melting candle wax The changes that are not chemical changes are called physical changes  x x Lighting a match Making ice cubes Ironing clothes  x x Launching a space rocket Printing Baking a cake x  

  4. Objectives • Define and give general equations for synthesis, decomposition, single-replacement, and double-replacement reactions. • Classify a reaction as synthesis, decomposition, single-replacement, single-replacement, double-replacement, or combustion.

  5. Objectives • List three types of synthesis reactions and six types of decomposition reactions. • List four types of single-replacement reactions and three types of double-replacement reactions. • Predict the products of single reactions given the reactants.

  6. Background • Thousands of known chemical reactions occur in various systems. Memorizing the equations for so many chemical reactions would be difficult. It is more useful and realistic to classify reactions according to various similarities and regularities.

  7. Combustion Decomposition Double Replacement The 5 Types of Reactions Synthesis Single Replacement

  8. Synthesis Reactions • In a synthesis reaction, also known as composition reaction, two or more substances combine to form a new compound. • General equation: A + XAX. • Types of synthesis reactions: • Reactions of elements with oxygen and sulfur. • Reactions of metals with halogens. • Synthesis reaction with oxides.

  9. Synthesis Reaction Synthesis reaction – 2 substances combine to form a single product A + B  AB +  2Mg(s) + O2(g) 2MgO(s) Magnesium and oxygen combine to form magnesium oxide. 2H2(g)+ O2(g) 2H2O(l) Hydrogen and oxygen combine to form dihydrogen monoxide

  10. Look at thereactantsrepresentedbelow, whichreactioninvolves elements as reactants? Whichreactioninvolves compounds as reactants? Synthesis reaction Decomposition reaction A synthesis reaction involves the combination smaller molecules A decomposition reaction involves the breaking apart of larger molecules

  11. Recognizingthe types of reactants is key to identifyingthereaction type Lets uslook at two of thereactiontypes:synthesis and decomposition

  12. Synthesis Reactions • Involve the combination of smaller atoms and /or molecules into larger molecules. • They are also called combination reactions • General formula • A + BAB • If you see two elements as reactants, you know the reaction has to be a sysnthesis reaction

  13. Synthesis reactions can also involve combinations of small molecules. • For example, when ammonia and hydrogen chloride vapours combine, they form a white smoke as solid particles of ammonium chloride are formed. Hydrogen chloride + ammonia ammonium chloride HCl + NH3 NH4Cl

  14. Similarly, combination of water and carbon dioxide molecules Synthesis reaction CO2 + H2O H2CO3(aq) Typical example , rainwater

  15. Examples of Reaction of Elements with Oxygen and Sulfur: • Forming Oxides and sulfides: • 2Mg(s) + O2(g) 2MgO(s) • 16Rb(s) + S8(s) 8Rb2S(s) • 8Ba(s) + S8(s) 8BaS(s) • S8(s) + 8O2(g) 8SO2(g) • C(s) + O2(g) CO2(g) • 2C(s) + O2(g) CO(g) • 2H2(g) + O2(g) 2H2O(l)

  16. Forming Oxides and sulfides: • 2Fe(s) + O2(g) 2FeO(s) • 4Fe(s) + 3O2(g) 2Fe2O3(s)

  17. Reactions of Metals with Halogens (most metals react with the halogens (group 17) to form either ionic or covalent compounds. • Group 1 metals with Group 17 elements • 2Na(s) + Cl2(g) 2NaCl(s) • 2K(s) + I2(g) 2KI(s) • Group 2 metals with Group 17 elements • Mg(s) + F2(g) MgF2(s) • Sr(s) + Br2(l) SrBr2(s)

  18. Fluorine is so reactive that it combines with almost all metals: • 2Na(s) + F2(g) 2NaF(s) • 2Co(s) + F2(g) 2CoF3(s) • U(s) +3F2(g) UF6(g) • Practical application with fluorine: • Sodium fluorine added to municipal water supplies. • Cobalt(III) fluoride is a strong fluorinating agent. • Uranium(VI) fluoride is the first step in the production of uranium for use in nuclear power plants.

  19. Synthesis Reactions with Oxides • Oxides of active metals react with water to produce metal hydroxides - example • CaO(s) + H2O(l) Ca(OH)2(s) • Many oxides of nonmetals (upper right portion of the periodic table) react with water to produce oxyacids - example • SO2(g) + H2O(l) H2SO3(aq) • this reacts with oxygen to produce sulfuric acid 2H2SO3(aq) + O2(g) 2H2SO4(aq) • Certain metal oxides and nonmetal oxides react with each other in synthesis reaction to form salts. CaO(s) + SO2(g) CaSO3(s)

  20. Decomposition Reactions • In a decomposition reaction, a single compound undergoes a reaction the produces two or more products. • General equation AX A + X

  21. Decomposition Reactions • It involve the splitting of a large molecule into elements or smaller molecules. • General formula AB A + B • Example, electrolysis of water uses electricity to split water molecules into their elements water hydrogen + oxygen H2O H2 + O2 2H2O 2H2 + O2

  22. Decomposition Reaction Decomposition reaction – A single compound breaks down into 2 or more products. AB  A + B  + 2H2O2(aq) O2(g)+ 2H2O(l) Hydrogen peroxide decomposes into oxygen gas and dihydrogen monoxide. 2NaCl(s)  2Na(s) + Cl2(g) Sodium chloride decomposes into sodium and chlorine gas.

  23. Types of decomposition reactions • Decomposition of Binary Compounds • Decomposition of Metal Carbonates • Decomposition of Metal Hydroxides • Decomposition of Metal Chlorates • Decomposition of Acids • Decomposition of Binary Compounds electricity • 2H2O(l) 2H2(g) + O2(g) (called electrolysis) 2HgO(s) 2Hg(l) + O2(g)

  24. Decomposition of Metal Carbonates • CaCO3(s) CaO(s) + CO2(g) • Decomposition of Metal Hydroxides • Ca(OH)2(s) CaO(s) + H2O(g) • Decomposition of Metal Chlorates • 2KClO3(s) 2KCl(s) + 3O2(g) • Decomposition of Acids • H2CO3(aq) CO2(g) + H2O(l) • H2SO4(aq) SO3(g) + H2O(l)

  25. Single-Replacement Reactions • In a single-replacement reaction, also know as a displacement reaction, one element replaces a similar element in a compound. • A + BX AX + B or • Y + BX BY + X

  26. Single Displacement Reactions • Are chemical changes that involve an element and a compound as reactants. • One element displaces or replaces another element from the compound. • Example, when magnesium ribbon is placed in a solution of silver nitrate Mg + AgNo3 Ag + Mg(NO3)2 Mg + 2AgNO32Ag + Mg(NO3)2

  27. Single Replacement Reaction Single Replacement reaction – A single element takes the place of another element in a compound. A + BC  C + AB +  + AgNO3(aq) + Cu(s)  Ag(s) + CuNO3(aq) Copper replaces silver in silver nitrate to copper (I) nitrate and silver. Zn(s) + 2HCl(aq)  H2(g) + ZnCl2(aq) Zinc replaces hydrogen in hydrogen chloride to yield hydrogen gas and zinc chloride.

  28. Types of single-replacement reactions • Replacement of a Metal in a Compound by Another Metal • Replacement of Hydrogen in Water by a Metal • Replacement of Hydrogen in Acid by a Metal • Replacement of Halogens • Replacement of a Metal in a Compound by Another Metal. • A more active metal will replace a less active metal. • 2Al(s) + 3Pb(NO3)2(aq) 3Pb(s) + 2Al(NO3)3(aq)

  29. Replacement of Hydrogen in Water by a Metal • The most-active metals such as those in Group 1, react vigorously with water to produce metal hydroxides and hydrogen. • example2Na(s) + 2H2O(l) 2NaOH(aq) + 4H2(g) • example 3Fe(s) + 4H2O(g) Fe3O4(s) + 4H2(g) (Less-active metals react with steam to form a metal oxide and hydrogen gas) • Replacement of Hydrogen in an Acid by a Metal • The more-active metals react with certain acidic solutions replacing the hydrogen in the acid. • example Mg(s) + 2HCl(aq) H2(g) + MgCl2(ag)

  30. Replacement of Halogens • One halogen replaces another halogen in a compound. Fluorine is the most-active halogen and can replace any other halogen in their compounds. • Cl2(g) + 2KBr(aq) 2KCl(aq) + Br2(l) • F2(g) + 2NaCl(aq) 2NaF(aq) + Cl2(s)

  31. Double-Replacement Reactions • In double-replacement reactions, the ions of two compounds exchange places in an aqueous solution to form two new compounds. • General equation AX + BY AY + BX • Formation of a Precipitate • The formation of a precipitate occurs when the cations of one reactant combine with the anions of another reactant to form an insoluble or slightly soluble compound. • 2KI(aq) + Pb(NO3)2(aq) PbI2(s) + 2KNO3(aq)

  32. Formation of a Gas • In some double-replacement reactions, one of the products in an insoluble gas that bubbles out of the mixture. • FeS(s) + 2HCl(aq) H2S(g) + FeCl2(aq) • Formation of Water • In some double-replacement reactions, a very stable molecular compound, such as water, is one of the products. • HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

  33. Double Replacement Reaction Double Replacement reaction – Elements in 2 compounds switch places to form 2 new compounds. AB + CD  AD + CB + +  HCl(aq) + NaOH(aq)  HOH(l) + NaCl(aq) Hydroxide and chlorine switch places to hydrogen hydroxide (water) and sodium chloride Na2S(aq) + Zn(NO3)2(aq)  2NaNO3(aq) + ZnS (s) Sulfur and nitrate switch places to form sodium nitrate and zinc sulfide.

  34. Combustion Reactions • In a combustion reaction, a substance combines with oxygen, releasing a large amount of energy in the form of light and heat. • 2H2(g) + O2(g) 2H2O(g) • C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(g)

  35. Combustion Reaction Combustion reaction – A hydrocarbon (H and C) burn in oxygen to produce water and carbon dioxide. Heat is given off as energy. CxHx + O2 H2O + CO2 + heat CH4(g) + 2O2(g)  2H2O (l) + CO2(g) Methane burns in oxygen to produce water and carbon dioxide. C12H22O11(s) + 12O2(g)  11H2O(l) + 12CO2(g) Sucrose burns in oxygen to produce water and carbon dioxide.

  36. Section Review • List five types of chemical reactions. • Complete and balance each of the following reactions identified by type: • synthesis: ______ Li2O • decomposition: Mg(ClO3)2 ______ • single-replacement: Na + H2O ______ • double-replacement: HNO3 +Ca(OH)2 • combustion: C5H12 + O2 ______

  37. Section Review • Classify each of the following reactions: • N2(g) + 3H2(g) 2NH3(g) • 2Li(s) + 2H2O(l) 2LiOH(aq) + H2(g) • 2NaNO3(s) 2NaNO2(s) + O2(g) • 2C6H14(l) + 19O2(g) 12CO2(g) +14H2O(l) • NH4Cl(s) NH3(g) + HCl(g) • BaO(s) + H2O(l) Ba(OH)2(aq) • AgNO3(aq) + NaCl(aq) AgCl(s) + NaNO3(aq)

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