Chemical Reactions

1. Chemical Reactions SCH3U - Unit 2

2. Chemical Change Chemical Change = any change in which a new substance is formed Evidence of Chemical Change: • Change in colour • Change in odour • Formation of gas/solid • Release/absorption of heat

3. Collision-Reaction Theory • A theory stating that chemical reactions involve collisions and rearrangements of atoms or groups of atoms and that the outcome of collisions depends on the energy and orientation of the collisions • No reaction occurs if: • Molecules don’t have enough energy • Molecules don’t collide in the right orientation

4. Orientation • The molecules must be in a certain 3-D arrangement to allow a reaction e.g. CH2=CH2 + HCl -> CH3CH2Cl

5. Energy Required • A reaction doesn’t occur unless the particles collide with a certain minimum energy called the activation energy of the reaction • Activation energy is the minimum energy required before a reaction can occur. You can show this on an energy profile for the reaction. For a simple over-all exothermic reaction, the energy profile looks like this:

6. Activation Energy Profile

7. Chemical Equations • Chemical Equation = a representation of a chemical reaction that indicates the: • Chemical formulas • Relative number of entities • States of matter of the reactants and products Reactants  Products

8. Chemical Equations In general: • Reactant A + Reactant B  Product C • Reactant = is a chemical that is used up in a chemical reaction • Product = is a product that is created during a chemical reaction.

9. Chemical Formulas • A chemical formula uses subscripts to indicates the number of atoms in a compound • Example: H2O • Has 2 atoms of H • And 1 atom of O • Example: C6H12O6 • Has 6 atoms of C • Has 12 atoms of H • And 6 atoms of O

10. Relative # of Entities • Coefficient = a whole number indicating the ratio of molecules of each substance involved in a chemical reaction • The large number on the left side of a molecule’s formula • Example: Mg + 2Cl MgCl2 • Example: 6 K + N2  2 K3N

11. State of Matter • Solid = (s) • Liquid = (l) • Gas = (g) • Solution = (aq) • Example: 6 K(s) + N2(g) 2 K3N(s)

12. Catalysts • A catalyst is a substance which speeds up a reaction, but is chemically unchanged at the end of the reaction .e.g conc. H2SO4 in many different reactions • Adding a catalyst has exactly this effect on activation energy. A catalyst provides an alternative route for the reaction. That alternative route has a lower activation energy • Draw a standard energy profile and then draw a new line to represent the inclusion of a catalyst

13. 5 Types of Chemical Reactions • Combustion • Synthesis • Decomposition • Single Displacement • Double Displacement

14. 5 Types of Chemical Reactions Generalizations: Combustion: AB + oxygen  oxides of A & B + heat Synthesis: A + B  C Decomposition: AB  A + B Single Displacement: A + BC  AC + B Double Displacement: AB + CD  AD + CB

15. Combustion Reactions • Combustion Reaction: the reaction of a substance with oxygen, producing oxides and energy • Also know as burning • For a combustion reaction to occur 3 things must be present: • Fuel • Oxygen • Heat

16. Combustion Reactions • C

17.  C C C C C C C C C C C C C C C C + O O O O O O O O O O O O O O O O O O O O O O O O O O O O O O O O Type of Reaction: Synthesis Example C + O2 General: A + B  AB

18. Synthesis Reaction Characteristics • Two or more substances (elements or compounds) react to form ONE product. Combination of smaller atoms/molecules into larger molecules. • Usually exothermic (energy is produced) • Can occur naturally or by an initial application of energy (heat, flame, UV light, use of catalyst)

19. Predicting Products of Synthesis Reactions • Metal + oxygen → metal oxide (basic oxide) EX. 2Mg(s) + O2(g) → 2MgO(s) • Nonmetal + oxygen → nonmetallic oxide (acidic oxide) EX. C(s) + O2(g) → CO2(g) • Metal oxide + water → metallic hydroxide (base) EX. MgO(s) + H2O(l) → Mg(OH)2(s) • Nonmetallic oxide + water → acid EX. CO2(g) + H2O(l) → ; H2CO3(aq) • Metal + nonmetal → salt EX. 2 Na(s) + Cl2(g) → 2NaCl(s) • A few nonmetals combine with each other. EX. 2P(s) + 3Cl2(g) → 2PCl3(g) • These two reactions should be remembered: N2(g) + 3H2(g) → 2NH3(g) NH3(g) + H2O(l) → NH4OH(aq)

20.  + Na Na Cl Cl Type of Reaction: Decomposition Example: NaCl General: AB  A + B

21. Hg Hg Hg Hg  + O O O O Type of Reaction: Decomposition Example 2HgO General: AB  A + B

22. Decomposition Reaction Characteristics • ONE reactant produces two or more products. Splitting of large molecules into elements or smaller molecules. • Usually endothermic (requires energy) • Can require energy in the form of heat, electricity, catalyst, UV light • *some decomposition rxns occur at room temperature

23. Predicting Products of Decomposition Reactions • Metallic carbonates, when heated, form metallic oxides and CO2(g). EX. CaCO3(s) → CaO(s) + CO2(g) • Most metallic hydroxides, when heated, decompose into metallic oxides and water. EX. Ca(OH)2(s) → CaO(s) + H2O(g) • Metallic chlorates, when heated, decompose into metallic chlorides and oxygen. EX. 2KClO3(s) → 2KCl(s) + 3O2(g) • Some acids, when heated, decompose into nonmetallic oxides and water. EX. H2SO4 → H2O(l) + SO3(g) • Some oxides, when heated, decompose. EX. 2HgO(s) → 2Hg(l) + O2(g) • Some decomposition reactions are produced by electricity. EX. 2H2O(l) → 2H2(g) + O2(g) EX. 2NaCl(l) → 2Na(s) + Cl2(g)

24. + +  Cl Cl Cl Cl Zn Zn Cu Cu Type of Reaction: Single Displacement Example: Zn + CuCl2 General: AB + C  AC + B

25. Ca Ca  + + Mg Mg S S O O Type of Reaction: Double Displacement Example: MgO + CaS General: AB + CD  AD + CB

26. Chemical Reactions combustion: AB + oxygen oxides of A & B synthesis: A + B  C decomposition: AB  A + B single displacement: A + BC  AC + B double displacement: AB + CD  AD + CB