1. Chemical Equations Chp 7 and 8

2. Chemical and Physical Changes  Physical Change –No new substance is formed –Can often be reversed or undone –Ex. Melting, freezing, crushing, etc  Chemical Change –A new substance is produced as compounds switch atoms –Usually are permanent –Ex. Cooking, decaying, rusting, etc

3. Evidence of a Chemical Change  Color change  Bubbles  Solid forms  Flames  Heat  Other visual cues

4. The Chemical Equation  A written representation of a chemical change  Reactants – what compounds/atoms exist before the reaction (before the arrow)  Products – what compounds/atoms exist after the reaction has occurred (after the arrow)  Subscripts – lowerset # after an atom, indicates the amount of that atom present in the compound  Coefficient – large # in front of a compound, indicates that amount of that molecule needed for the chemical reaction to occur

5. Conservation of Matter  Atoms cannot be created or destroyed, they can only switch between compounds –That means, whatever types and #s of atoms are present before a reaction occurs (reactants) must equal the types and #s that are present afterward (products) –Or, equations must be balanced

6. Balancing Equations  Place coefficients in front of compounds to make all atoms equal on each side of the equation (conserve the matter)  The coefficient multiplies by the subscript of each atom within that compound  You may NEVER change a subscript when balancing, that would change the compound –Ex. H 2 O is water, H 2 O 2 is hydrogen peroxide

7. An Example ____ H 2 + _____ O 2 ____ H 2 O ReactantsProducts H = 2H = 2 O = 2O = 1

8. An Example, Continued ____ H 2 + _____ O 2 2 H 2 O ReactantsProducts H = 2H = 2 4 O = 2O = 1 2

9. An Example 2 H 2 + _____ O 2 2 H 2 O 2 H 2 + _____ O 2 2 H 2 O ReactantsProducts H = 2 4H = 2 4 O = 2O = 1 2

10. Tips  If an atom appears in more than one compound on the same side, leave it until the end (it will often work itself out)  Start with atoms that have an odd number of one side and an even number on the other  Use a pencil (it’s okay to go back and erase when things don’t work out)  If you get to the end, and everything but one atoms works, try multiplying all the existing coefficients by 2.

11. Types of Reactions  Synthesis – 2 reactants (or more) form 1 product –A + B C  Decomposition – 1 reactant forms 2 products (or more) –C A + B  Single Displacement – 1 lone element on the reactant side combines with part of a compound, leaving another element alone –A + BC AC + B

12. Types of Reactions, Cont’d  Double Displacement – 2 compounds swap atoms –AB + CD AD + CB –Has 2 subtypes:  Acid base – water and a salt are formed  Precipitation – a solid is formed  Combustion – carbon dioxide and water are produced  Oxidation-Reduction – electrons are exchanged (doesn’t fit anywhere else)

13. Writing Equations from Words  Decide which are your products and which are your reactants –Look for words like creates, yields, forms, etc –Words like and, also, in the presence of usually indicate a plus  Swap and drop your formulas –Check charges from the table, swap if not equal –Remember your diatomic elements (group 17 and all non-noble gases)  Balance your equation –Remember, you can’t have an element on one side, but not on the other

14. An example  Calcium metal reacts in the presence of oxygen gas to form solid calcium oxide. Write a balanced equation for this.  Start with Ca and O, end with CaO  Ca + O -> CaO  Remember that O is diatomic though so its O 2  Check charges: Ca +2, O -2 so we don’t need subscripts  Then balance so:  2 Ca + O 2 -> 2 CaO

15. Aqueous Solution  When most ionic compounds are placed in water, the two ions separate from each other and act independently –Ex. NaCl (aq) Na + + Cl -  Those who don’t can be determined by viewing the rules of solubility on your purple sheet –Ex. Ba(CrO4) (s) Ba(CrO4) (s)

16. Writing a Complete Molecular Equation  All reactants are soluble (ie they dissolve in water)  Reactants swap partners (need to remember to swap and drop on new compounds)  One of the products will be insoluble (check your purple sheet) and one will be soluble –Soluble compounds have (aq) after them because they separate into separate ions in water –Insoluble compounds have (s) after them because they remain together as a solid (or precipitate)  Need to balance the equation if its not balanced

17. An Example  K 2 (CrO 4 ) (aq) + Ba(NO 3 ) 2 (aq)  K 2 (CrO 4 ) (aq) + Ba(NO 3 ) 2 (aq) K(NO 3 ) + Ba(CrO 4 )  K 2 (CrO 4 ) (aq) + Ba(NO 3 ) 2 (aq) 2 K(NO 3 ) (aq) + Ba(CrO 4 ) (s) +1- 1 +2- 2 ?

18. Complete Ionic Equation  Split all (aq) compounds up into their separate ions since they don’t actually stay together in solution.  Ex. K 2 (CrO 4 ) (aq) + Ba(NO 3 ) 2 (aq) 2 K(NO 3 ) (aq) + Ba(CrO 4 ) (s) becomes 2K + + CrO 4 -2 + Ba +2 + 2NO 3 - 2K + + 2NO 3 - + Ba(CrO 4 ) (s)

19. Net Ionic Equation  Eliminate all spectator ions (the ones that always stay apart and are never part of a compound)  Ex. 2K + + CrO 4 -2 + Ba +2 + 2NO 3 - 2K + + 2NO 3 - + Ba(CrO 4 ) (s) Becomes CrO 4 -2 + Ba +2 Ba(CrO 4 ) (s)