1. Chemical Reactions: An Introduction

2. Indicators of a Chemical Reaction A color change A solid forms (precipitate) A gas forms The temperature changes (hot OR cold)

3. Chemical Equations Reactants – what goes INTO the rxn Products – what comes OUT OF the rxn Reactant A + Reactant B Product A

4. Law of Conservation of Mass In a chemical reaction, atoms are neither created or destroyed All atoms in the reactants MUST be accounted for in the products.

5. Physical States After each reactant and product is a symbol representing the physical state of the element or compound (s) – solid (l) – liquid (g) – gas (aq) – aqueous (dissolved in water)

6. Balancing Chemical Equations Done by trial and error  1.Balance elements that only appear once 2.Keep polyatomic ions together if possible 3.***NEVER EVER EVER EVER CHANGE A SUBSCRIPT!!!*** 4.Balance hydrogen and oxygen last

7. Diatomic Molecules 7 elements exist in nature as diatomic molecules (2 atoms) –Hydrogen (H 2 ) –Oxygen (O 2 ) –Nitrogen (N 2 ) –Fluorine (F 2 ) –Chlorine (Cl 2 ) –Bromine (Br 2 ) –Iodine (I 2 )

8. C 2 H 5 OH (l) + O 2 (g)  CO 2 (g) + H 2 O(g)

9. Fe 2 O 3 (s) + HNO 3 (aq)  Fe(NO 3 ) 3 (aq) + H 2 O(l)

10. H 2 S(g) + Pb(NO 3 ) 2 (aq)  PbS(s) + HNO 3 (aq)

11. Types of Reactions

12. Objectives 1.Give general equations for types of reactions 2.Classify reactions 3.List 3 types of synthesis and 6 decomposition reactions 4.List 4 types of single-replacement and 3 types of double-replacement reactions 5.Predict products of reactions given the reactants

13. Synthesis Reactions General Formula: A + X  AX

14. Synthesis with Oxygen With metals form metal oxides Ex – Mg(s) + O 2 (g)  MgO(s) K(s) + O 2 (g)  K 2 O(s) 2Fe(s) + O 2 (g)  2FeO(s) 4Fe(s) + 3O 2 (g)  2Fe 2 O 3 S 8 (s) + 8O 2 (g)  8SO 2 (g) C(s) + O 2 (g)  CO 2 (g)

15. Synthesis with Sulfur With metals produce metal sulfides Ex-16Rb(s) + S 8 (s)  8Rb 2 S(s) 8Ba(s) + S 8 (s)  8BaS(s)

16. Metals with Halogens Group 1: M + X 2  2MX Ex – Na(s) + Cl 2 (g)  2NaCl(s) Group 2: M + X 2  MX 2 Ex – Mg(s) + F 2 (g)  MgF 2 (s)

17. Metal Oxides with Water Group 1 & 2 form hydroxides Ex – K 2 O(s) + H 2 O(l)  2KOH(aq) CaO(s) + H 2 O(l)  Ca(OH) 2 (l)

18. Non-metal Oxide with Water Form oxyacids Ex – SO 2 (g) + H 2 O(l)  H 2 SO 3 (aq) P 2 O 5 (s) + 3H 2 O(l)  2H 3 PO 4 (aq)

19. Decomposition Reactions AX  A + X

20. Decomposition of Binary Compounds Breaks down into its elements Process called electrolysis

21. Decomposition of Metal Carbonates Form metal oxides and carbon dioxide

22. Decomposition of Metal Hydroxides Form metal oxides and water

23. Decomposition of Acids Break down into non-metal oxides and water

24. Single Replacement General Formula A + BX  AX + B

25. Metal Replaces Another Metal Aluminum is more reactive than lead

26. Replacement of Hydrogen in Water by a Metal More Active Metals Less Active Metals

27. Replacement of Hydrogen in an Acid by a Metal Metals more active than hydrogen

28. Replacement of Halogens Each halogen can replace the halogen below it on the periodic table

29. Double Replacement AX + BY  AY + BX Formation of a precipitate Formation of a gas Formation of water

30. Formation of a Precipitate An insoluble product forms

31. Formation of a Gas Insoluble gas forms Example

32. Formation of Water Water forms during reaction

33. Combustion Reaction Substance reacts with oxygen to release heat and light Products are often carbon dioxide and water

34. Neutralization Reaction HA + BOH  AB + HOH Usually these are acid-base reactions Products include salt and water

35. Activity Series

36. Objectives 1.Explain the significance of an activity series 2.Use an activity series to predict if a reaction will take place

37. Metals vs. Nonmetals Greater activity of a metal indicates how easily it loses electrons Greater activity of a nonmetal indicates how easily it gains electrons In a single-replacement reaction, if an element with lower activity is to be replaced, the reaction will take place.