1. Ch. 11 Chemical Reactions 11.1 Describing Chemical Reactions

2. I. Equation Basics A. Fe (s) + O 2 (g)  Fe 2 O 3 (s) Reactants Products B. Skeleton equation: does not tell amounts of each part C. Symbols for states of matter: solid (s)gas (g)liquid (l)aqueous (aq) Substance dissolved in water D. Catalyst: substance that speeds up rate of reaction

3. II. Balancing Equations B. Rules: –1. Only use coefficients (#’s in front of substances) –2. Must have same # of each element on each side –3. Equation must be reduced A. Atoms are never lost or gained in a reaction, they are just rearranged.

4. III. Examples A. N 2 (g) + O 2 (g)  N 2 O 5 (g) B. __MnO 2 + HCl  MnCl 2 + H 2 O + Cl 2 C. FeCl 3 + Ca(OH) 2  Fe(OH) 3 + CaCl 2

5. 11.2 Types of Chemical Reactions

6. I. Combination Rxns A. Two or more substances combine to one B. Na (s) + Cl 2 (g) --> NaCl (s)Na (s) + Cl 2 (g) --> NaCl (s) C. If metal combines with non-metal, ionic compound produced D. If non-metals combined, covalent compound produced + () 

7. II. Decomposition Rxns A. Single compound broken into two or more products B. HgO (s)  Hg (l) + O 2 (g) Heat Means Heat Added

8. C. Most decomp. rxns require energy (heat) as a catalyst ***Demo: H 2 O 2 (l)  H 2 O (l) + O 2 (g)H 2 O 2 (l)  H 2 O (l) + O 2 (g) ***Demo: NI 3  N 2 + I 2 *** ***Demo: NI 3  N 2 + I 2 ***

9. III. Single-Replacement A. Atoms of one element replace atoms of a second element in a compound B. Mg + Zn(NO 3 ) 2  Mg(NO 3 ) 2 + Zn C. Reactivity of metals determines whether one atom will replace another D. Mg more reactive than Zn, removes Nitrate ***Demo Thermite Rxn*** +  +

10. IV. Double Replacement A. Two atoms switch places (often forming solid) B. K 2 CO 3 (aq) + BaCl 2 (aq)  2 KCl (aq) + BaCO 3 (s) *** Demo: Silver Nitrate and Magnesium Chloride***

11. V. Combustion Rxns A. Element or compound reacts with O 2, usually forming energy B. Mg (s) + O 2 (g)  MgO (s) C. If compound has C and H, products usually H 2 O and CO 2 D. C x H y + O 2  CO 2 + H 2 O ***Demo: Nitrocellulose***

12. 11.3 Reactions in Aqueous Solution

13. I. Net Ionic Equations A. Equation showing only particles that take part in reaction B. Process: 1. Start with full equation AgNO 3 (aq) + Na 2 S (aq)  Ag 2 S (s) + NaNO 3 (aq) C. 2. Separate ions in aqueous form Ag + (aq) + NO 3 - (aq) + Na + (aq) + S 2- (aq)  Ag 2 S (s) + Na + (aq) + NO 3 - (aq)

14. Ag + (aq) + NO 3 - (aq) + Na + (aq) + S 2- (aq)  Ag 2 S (s) + Na + (aq) + NO 3 - (aq) D. 3. Cross off ions appearing on both sides of reaction as aqueous Net Ionic equation: 2 Ag + (aq) + S 2- (aq)  Ag 2 S (s) E. 4. Balance everything left to get…

15. II. Try This Write the balanced net ionic equation for the following: Step 1 Zn (s) + HCl (aq)  ZnCl 2 (aq) + H 2 (g) Step 2 Zn (s) + H + (aq) + Cl - (aq)  Zn 2+ (aq) + Cl - (aq) + H 2 (g) Step 3 Zn (s) + H + (aq)  Zn 2+ (aq) + H 2 (g) 2 Step 4

16. III. Precipitates A. Solids formed when two aqueous mixtures form insoluble (“non-dissolvable”) compound B. Precipitate formation based on ion solubility rules Solubility Rules: Always Soluble: Alkali metals (1 st column), NH 4 +, NO 3 -, ClO 3 -, ClO 4 -, C 2 H 3 O 2 - Mostly Soluble: Cl -, Br -, I - (except Ag +, Pb 2+, Hg 2 2+ ) F - (except Ca 2+, Ba 2+, Sr 2+, Pb 2+, Mg 2+ ) SO 4 2- (except Ca 2+, Ba 2+, Sr 2+, Pb 2+ ) Mostly Insoluble: O 2-, OH - (except w/ alkali metals, NH 4 +, Ca 2+, Sr 2+, Ba 2+ somewhat soluble) CO 3 2-, PO 4 3-, S 2-, SO 3 2-, C 2 O 4 2-, CrO 4 2- (except w/ alkali metals, NH 4 + )