Types of Chemical Reactions

1. Types of Chemical Reactions • There are five main types of chemical reactions that we will be covering in this chapter: • Combination reactions • Decomposition reactions • Single-Replacement reactions • Double-Replacement reactions • Combustion reactions

2. Combination Reactions • In a combination reaction, two or more substances combine to form a single substance. • General Equation: R + S  RS • Example: Mg(s) + O2 (g) MgO(s)

3. Decomposition Reactions • In a decomposition reaction a single compound is broken down into two or more products. • General Equation: RS  R + S • Example: HgO(s) Hg(l) + O2

4. Single-Replacement Reactions • In a single-replacement reaction, one element replaces a second element in a compound. • General Equation: T + RS  TS + R • Example: K(s) + H2O(l) KOH(aq) + H2 (g)

5. Double-Replacement Reactions • Double-replacement reactions involve an exchange of positive ions (or cations) between two reacting compounds. These reactions generally produce a precipitate, a gaseous product, or water. • General Equation: R+S- + T+U- T+S- + R+U- • Example: K2CO3(aq) + BaCl2 (aq) BaCO3 (s) + KCl(aq)

6. Combustion Reactions • In a combustion reaction an element or a compound (often a hydrocarbon) reacts with oxygen, often producing energy as heat and light. The products of these types of reactions are carbon dioxide and water. • Example: The combustion of methane gas in air __CH4 (g) + __O2 (g) __CO2(g) + __H2O(g)

7. Combustion Reactions • The reaction between some elements and oxygen is also an example of a combustion reaction. For example: __Mg(s) + __O2 (g) __MgO(s) • Here Mg acts as the fuel for the reaction and when reacted with O2 (g) will for an oxide – in this case magnesium oxide (MgO(s)). What other type of chemical reaction does this look like?

8. Combination Reactions • Complete and balance the following combination reactions: Be + O2 H2 + O2  H2O

9. Decomposition Reactions • Complete and balance the following decomposition reactions: HI  Mg(ClO3)2  MgCl2 +

10. Single-Replacement Reactions and Activity Series • For example: Mg(s) + Zn(NO3)2 (aq) Mg(s) + AgNO3 (aq)  Mg(s) + LiNO3 (aq) 

11. Single-Replacement Reactions • Complete the equations for the following single replacement reactions that take place in aqueous solution. Balance each equation. If a reaction does not occur (use activity series) write “no reaction.” Fe(s) + Pb(NO3)2 (aq) Cl2 (g) + NaI(aq)  Ca(s) + H2O(l) 

12. Double-Replacement Reactions and the Solubility Rules • Mixing the solutions of two ionic compounds can sometimes result in the formation of an insoluble salt called a precipitate. • In order for a double-replacement reaction to occur there must be fewer ions in solution when the two compounds are mixed. Otherwise all we have is a mixture. • Recall our conductivity demo – when ionic compounds dissociate they can conduct electricity.

13. Double-Replacement Reactions • We also have fewer ions in solution if we produce a gas or water.

14. Double-Replacement Reactions • Write the products for the following reactions. Then balance each equation. NaOH(aq) + Fe(NO3)3(aq) Ba(NO3)2(aq) + H3PO4(aq)

15. Double-Replacement Reactions and the Solubility Rules • Should a precipitate form when aqueous solutions of Na2SO4 (aq) and Ba(NO3)2 (aq) are mixed?

16. Double-Replacement Reactions and the Solubility Rules • Should a precipitate form when aqueous solutions of NaCl(aq) and Zn(NO3)2 (aq) are mixed?

17. Combustion Reactions • Predict the product of the following combustion reaction: S(s) + O2(g) 

18. Combustion Reactions • Write a balanced equation for the complete combustion of each compound. • Heptane (C7H16) • Glucose (C6H12O6)

19. Rates of Reaction • There are three main factors that effect the rates of chemical reactions: • The presence of a catalyst • Temperature • Concentration

20. Catalysts • The presence of a catalyst in a chemical reaction will speed up the reaction while not being used up in the reaction itself. Because a catalyst is neither a reactant nor a product, its formula is written above the yields sign () in a chemical equation.

21. Catalysts • For a reaction to occur between two substances (or reactants) particles of those substances must collide. Not only that, they must hit each other with enough force to cause a change to take place. The amount of energy the particles must have when they collide is called the Activation Energy. • Catalysts are used to lower activation energy.

22. Temperature • Most reactions go faster at high temperatures. • Ex: Baking a cake speeds up the reactions that change the liquid batter into a spongy product. • Lowering the temperatureslows down most reactions. • Batteries tend to last longer if they are kept cool, slowing down the reaction that takes place within them.

23. Concentration • Concentration – The amount of substance present in a certain volume. • Raising the concentration of a reactant will speed up a reaction because there are more particles per volume. • More particles = more collisions = increase in reaction rate.