1. Chapter 8 Describing Chemical Change Types of Chemical Reactions Reactions in Aqueous Solution

2. Chapter 8.1 Describing Chemical Change Word Equations Chemical Equations Balancing Chemical Equations

3. Word Equations Reaction – one or more substances (the reactants) change into one or more new substances(the products) Reactants  Products  = ? Yields, gives or reacts

4. Word Equations As reactants are converted to products, the bonds holding the atoms together are broken and new bonds are formed. REMEMBER: The atoms are neither created nor destroyed, just rearranged. (Law of Conservation of Mass)

5. Word Equations Rust Iron reacts with oxygen to produce iron(III) oxide (rust) Iron + Oxygen  Iron(III) Oxide (reactants) (yields) (products)

6. Word Equations Hydrogen peroxide reacts to form water and oxygen gas. Hydrogen peroxide  water + oxygen MnO 2 H 2 O 2(aq)  H 2 O (l) + O 2(g)

7. Word Equations Burning of Methane Methane + Oxygen  Carbon Dioxide + Water

8. Chemical Equations Iron + Oxygen  Iron(III) Oxide Fe + O 2  Fe 2 O 3 Now add physical states Fe (s) + O 2 (g)  Fe 2 O 3 (s)

9. Chemical Equations Reactions with a catalyst Catalyst – a substance that speeds up the rate of a chemical reaction, but is not used up in the reaction. A catalyst is written above the arrow

10. Chemical Equations Manganese(IV) oxide catalyzes the decomposition of hydrogen peroxide.

11. Balancing Chemical Reactions Each side of the equation has the same number of atoms of each element. C (s) + O 2(g)  CO 2(g) 1 carbon, 2 oxygen  1 carbon, 2 oxygen

12. Balancing Chemical Reactions H 2(g) + O 2(g)  H 2 O (l) 2 hydrogen, 2 oxygen  2 hydrogen, 1 oxygen Balance 2 H 2(g) + O 2(g)  2 H 2 O (l)

13. Balancing Chemical Reactions Rules 1) Determine the correct formulas 2) Write the formulas for the reactants on the left and the products on the right. Place a  in between. If there are two or more reactants or products, use a + in between. 3) Count the number of atoms of each element.

14. Balancing Chemical Reactions Rules 4) Balance the elements one at a time until you have equal numbers of elements on each side 5) Make sure all numbers are in their smallest whole number ratio.

15. Chapter 8.2 Types of Chemical Reactions Classifying Reactions Combination Reactions Decomposition Reactions Single-Replacement Reactions Double-Replacement Reactions Combustion Reactions Predicting Products of Reactions

16. Classifying Reactions Identify the five general types or reactions: Combination Reactions Decomposition Reactions Single-Replacement Reactions Double-Replacement Reactions Combustion Reactions

17. Combination Reactions Two or more substance combine to form a single substance. General Reaction: R + S = RS Example: 2Mg (s) + O 2(g)  2 MgO (s)

19. Decomposition Reactions A single compound is broken down into two or more substances. General Reaction: RS = R + S Example: 2HgO (s)  2 Hg (l) + O 2(g)

21. Single-Replacement Reactions (Single Displacement Reactions) One element replaces a second element in a compound. General Reaction: T + RS = TS + R Example: 2K (s) + 2H 2 O (l)  2KOH (aq) + H 2(g)

22. http://www.kentchemistry.com/links/Kinetic s/PredictingSR.htm http://www.kentchemistry.com/links/Kinetic s/PredictingSR.htm

23. Double-Replacement Reactions Exchange of positive ions between two reacting compounds. General Reaction: RS + TU = RU + TS R + S - + T + U - = R + U - + T + S - Example: K 2 CO 3(aq) + BaCl 2(aq)  2KCl (aq) + BaCO 3(s)

24. http://www.kentchemistry.com/links/Kinetic s/DRFlash.htm http://www.kentchemistry.com/links/Kinetic s/DRFlash.htm

25. Combustion Reactions An element or compound reacts with oxygen often producing energy as heat or light. General Reaction: C x H y + (x + y/4) O 2  xCO 2 + (y/2)H 2 O Example: CH 4 (g) + 2O 2 (g)  CO 2(g) + 2H 2 O (g)

27. Name each type of reaction 1) 2Mg (s) + O 2(g)  2 MgO (s) 2) 2HgO (s)  2 Hg (l) + O 2(g) 3) CH 4 (g) + 2O 2 (g)  CO 2(g) + 2H 2 O (g) 4) 2K (s) + 2H 2 O (l)  2KOH (aq) + H 2(g) 5) K 2 CO 3(aq) + BaCl 2(aq)  2KCl (aq) + BaCO 3(s)

28. Name each type of reaction 1) 2Mg (s) + O 2(g)  2 MgO (s) Combination 2) 2HgO (s)  2 Hg (l) + O 2(g) Decomposition 3) CH 4 (g) + 2O 2 (g)  CO 2(g) + 2H 2 O (g) Combustion 4) 2K (s) + 2H 2 O (l)  2KOH (aq) + H 2(g) Single Replacement 5) K 2 CO 3(aq) + BaCl 2(aq)  2KCl (aq) + BaCO 3(s) Double Replacement

29. Chapter 8.3 Reactions in Aqueous (aq) Solutions Net Ionic Equations Predicting the Formation of a Precipitate

30. Net Ionic Equations AgNO 3(aq) + NaCl (aq)  AgCl (s) + NaNO 3 (aq) Double Replacement Reaction Most ionic compounds dissociate (separate) into ions (cations and anions) when they dissolve in water. When all ions dissociate, we write the equation with the charges on it. = Complete Ionic Equation

31. Complete Ionic Equation Ag + (aq) + NO 3 - (aq) + Na + (aq) Cl - (aq)  AgCl (s) + Na + (aq) + NO 3 - (aq) The equation can be simplified by crossing out any ions that do not participate in the reaction. You do this by canceling out ions that appear on both sides. Ag + (aq) + NO 3 - (aq) + Na + (aq) Cl - (aq)  AgCl (s) + Na + (aq) + NO 3 - (aq) You are left with the Net Ionic Equation: Ag + (aq) + Cl - (aq)  AgCl (s)

32. Pb (s) + AgNO 3(aq)  Ag (s) + Pb(NO 3 ) 2(aq) Pb (s) + Ag + (aq) + NO 3 - (aq)  Ag (s) + Pb 2+ (aq) + 2NO 3 - (aq) Pb (s) + Ag + (aq)  Ag (s) + Pb 2+ (aq) Need to balance charges Pb (s) + 2Ag + (aq)  2Ag (s) + Pb 2+ (aq)

33. Predicting the Formation of a Precipitate Table F in reference table