1. Click to add text : CHEMICAL REACTIONS

2. Chemical reactions: Reactions that produce new substances PRODUCT: substance formed during a chemical reaction (right side of arrow) REACTANT: starting substance(s) in a chemical reaction (left side of arrow) Law of Conservation of Mass must be satisfied!

3. In this unit you should know… 1. How to balance chemical equations 2. Identify the different types of reactions 3. Be able to predict the products for both single and double replacement reactions 4. Determine if a reaction will take place using either the activity series of metals or solubility rules 5. Understand the role of a catalyst in a chemical reaction

4. Evidence of Chemical Reactions Temperature change: endothermic (colder), exothermic (hotter) Color Change Odor Gas Produced (bubbles) Precipitate: formed from 2 liquids

5. Balancing Equations Steps 1)Balance atoms that appear only once on each side. 2)Balance polyatomic ions that appear on both sides as a single unit. 3)Balance hydrogens. 4)Balance oxygens. 5)Never change the subscripts of a compound to balance an equation.

6. Types of Reactions Synthesis Reaction: 1. Two or more substances combine to form a single compound. 2. Usually energy is released (exothermic) 3. Basic reaction: A + B --> AB

7. Synthesis Reaction Examples: Element + Oxygen ----> Oxide Compound Magnesium + Oxygen ---> Magnesium Oxide Mg + O 2 ------> 2 MgO Metal Oxide + Water ---> Hydroxide Compound (base) CaO + H 2 O ---> Ca(OH) 2

8. Decomposition Reactions : 1. Single compound is broken down into two or more simpler products. 2. Usually requires energy. 3. Basic reaction: AB ---> A + B

9. Decomposition Reaction Examples: Metal Carbonate ----> Metal oxide + carbon dioxide Ca CO 3 ----> CaO + CO 2 Metal Hydroxide ----> metal oxide + water Ca(OH) 2 ---> CaO + H 2 O Metal Chlorate ---> metal chloride + oxygen 2KClO 3 ---> 2 KCl + O 2 Oxyacid ---> nonmetal oxide + water H 2 SO 4 ---> SO 3 + H 2 O

10. SINGLE REPLACEMENT REACTION: 1. One element replaces a similar element in a compound. 2. A reactive metal will replace any metal that is less reactive (see pg 288 Activity Series of Metals) 3. Nonmetal will replace other nonmetals.

11. Activity Series: Single Replacement Reactions Only One metal will only replace another if it is HIGHER on the activity series This is because it is a more reactive metal

12. Single Replacement cont. 4. Basic Reaction: A + BC ---> AC + B Y + BX ---> BY + X

13. Single Replacement Examples: Replacement of a metal in a compound by a more reactive metal Use activity series to determine if one metal is strong enough to replace the other one. If not, then no reaction will occur 2Al + 3Fe(NO 3 ) 2 ---> 3Fe + 2Al(NO 3 ) 3 Replacement of Halogens. Cl 2 + 2 KBr ---> 2KCl + Br 2 Metal replacing hydrogen in an acid. Zn + 2HCl ---> ZnCl 2 + H 2

14. DOUBLE-REPLACEMENT REACTIONS: 1.Exchange of positive ions between two compounds. 2.One compound formed is usually a precipitate, gas, or a molecular compound (often water) 3. Basic Equation: AB + CD ---> CB + AD 4. Use the solubility rules to determine whether or not a reaction will take place

15. Double-Replacement Examples: Metal oxide + acid ---> water + salt (metal/nonmetal) MgO + 2 Hcl ---> H 2 O + MgCl 2 Metal carbonate + acid ---> salt + carbon dioxide +water CaCO 3 + 2 HCl ---> Ca Cl 2 + CO 2 + H 2 O Acids + metal Hydroxide ---> salt + water HCl + NaOH ----> NaCl + H 2 O

16. Solubility Rules Overview List of rules used to determine whether or not a reaction will take place Remember! In order for a reaction to take place you must produce a gas or a precipitate from 2 liquids. Solubility rules tell us whether or not a precipitate (solid) is produced You will often see these letters indicating what state of matter a substance is Solid (s) Liquid (l) Gas (g) Aqueous (aq) = soluble in water

17. Double Rep. Reactions: Solubility Rules (see handout- do not have to copy down) 1. Soluble: All salts containing the ammonium or Group IA ions (Li +, Na +, K +, Rb +, Cs + ) 2. Soluble: All salts containing nitrate (NO3-), acetate (C 2 H 3 O 2 - ), and perchlorate (ClO 4 - ) 3. Soluble: All salts containing Group VIIA ions (Cl -, Br -, I - ), except those in Rule 5. 4. Soluble: All salts containing sulfate (SO 4 -2 ). Exceptions are barium sulfate, calcium sulfate, lead II sulfate, and strontium sulfate.

18. Solubility Rules cont. 5. Insoluble: All salts containing silver ion (Ag + ), lead II ions (Pb +2 ), and mercury I ions (Hg 2 +2 ) 6. Insoluble: All salts containing carbonate, chromates, hydroxides, oxides, phosphates, and sulfides Exceptions: Group IIA chromates, except barium chromate are solulbe Group IIA hydroxides, except magnesium hydroxide, are soluble

19. COMBUSTION REACTIONS: 1. Oxygen reacting with another substance. 2. Usually involves hydrocarbons (contain hydrogen & carbon) 3. Heat is always released. 4. Basic Equation: C X H Y + O 2 ---> H 2 O + CO 2 [x & y represent a ratio of carbon & hydrogen]

20. Combustion Examples: 4. Complete combustion: C 3 H 8 + 5O 2 ---> 3CO 2 + 4H 2 O 5. Incomplete combustion: creates carbon monoxide (CO), carbon, & water. [products cannot be predicted]

21. Catalysts A substance that increases the rate of a chemical reaction by lowering activation energies but is not itself consumed in the reaction. Example: Enzymes: allow many chemical rxns to occur at a rate that sustains life at normal living temperatures