1. BALANCING EQUATIONS AND TYPES OF CHEMICAL REACTIONS

2. What is a Chemical Equation?  An equation is a short-hand way of writing a chemical reaction.  Reactants are on the left, products are on the right. Mg + 2HCl  MgCl 2 + H 2 reactants products

3. How is this Accomplished? CChemical equations are balanced by adding coefficients in front of the substances in the equation. 1 MgF 2 + 1 Li 2 CO 3  1 MgCO 3 + 2 LiF NNever change the subscript to balance equations

4. Tips for Balancing Equations  1. Balance metals first.  2. Balance nonmetals except H and O.  3. If polyatomic ions are conserved in the reaction, then try balancing them as a unit.  4. Balance any remaining H’s and O’s.

5. Some practice problems:  Here are some practice problems.  1. __NaCl + __BeF 2 --> __NaF + __BeCl 2  2NaCl + BeF 2  2NaF + BeCl 2  2. __FeCl 3 + __Be 3 (PO 4 ) 2 --> __BeCl 2 + __FePO 4  2FeCl 3 + Be 3 (PO 4 ) 2 --> 3BeCl 2 + 2FePO 4 

6.  3. __AgNO 3 + __LiOH --> __AgOH + __LiNO 3  AgNO 3 + LiOH --> AgOH + LiNO 3  4. __CH 4 + __O 2 --> __CO 2 + __H 2 O  CH 4 + 2O 2 --> CO 2 + 2H 2 O  5. __Mg + __Mn 2 O 3 --> __MgO + __Mn  3Mg + Mn 2 O 3 --> 3MgO + 2Mn

7.  7. Na + H 2 O  NaOH + H 2  2Na + 2H 2 O  2NaOH + H 2  8. H 2 SO 4 + Ca(OH) 2  CaSO 4 + H 2 O  H 2 SO 4 + Ca(OH) 2  CaSO 4 +2 H 2 O

8. Types of Chemical Reactions  There are 5 overall types of chemical reactions:  1. Synthesis or Combination  2. Decomposition  3. Single Replacement  4. Double Replacement  5. Combustion

10. Synthesis: A + B  AB  Synthesis – “putting together”  H 2 (g) + Cl 2 (g) ----> 2HCl(g)  C(s) + O 2 (g) ----> CO 2 (g)  CaO(s) + H 2 O(l) ----> Ca(OH) 2 (s)

12. Decomposition: AB  A + B  Decomposition – “breaking apart”  C 12 H 22 O 11 (s) ---->12C(s) + 11H2O(g)  Pb(OH) 2 (s) ----> PbO(s) + H 2 O(g)  2Ag 2 O(s) ----> 4Ag(s) + O 2 (g)

14. Single Replacement: A+ BC  B + AC  Single replacement reactions - When one element displaces another element in a compound  Zn(s) + H 2 SO 4 (aq) ----> ZnSO 4 (aq) + H 2 (g)  2Al(s) + 3CuCl 2 (aq) ---> 2AlCl 3 (aq) + 3Cu(s)  Cl 2 (g) + KBr(aq) ----> KCl(aq) + Br 2 (l)

15. Single Replacement Reactions follow the Activity Series

16. The element by itself has to be more reactive than the one it is trying to replace. Examples:  Mg + Zn(NO 3 ) 2   Mg + AgNO 3   Mg + LiNO 3   Zn + H 2 SO 4   Na + H 2 O   Sn + NaNO 3   Cl 2 + NaBr 

18. Double Replacement: AB + CD  AD + CB  Double Replacement reaction – chemical reaction where two reactant ionic compounds exchange ions to form two new product compounds with the same ions.  AgNO 3 (aq) + NaCl(aq) ----> AgCl(s) + NaNO 3 (aq)  ZnBr 2 (aq) + 2AgNO 3 (aq) ----> Zn(NO 3 ) 2 (aq) + 2AgBr(s)  H 2 SO 4 (aq) + 2NaOH(aq) ----> Na 2 SO 4 (aq) + 2H 2 O(l)

19. Special Type of Double Replacement: Neutralization  A neutralization reaction occurs between an acid and a base.  A base is a metallic hydroxide, such as NaOH, KOH, Ca(OH) 2, Al(OH) 3, etc.  An acid and a base always react to form a salt and water.

20. Example:  HCl + NaOH  NaCl + HOH  H 2 SO 4 + Mg(OH) 2  MgSO 4 + H 2 O  H 3 PO 4 + Al(OH) 3 

22. Combustion: Burning (add O 2 )  CH 4 (g) + 2O 2 (g) ----> 2H 2 O(g) + CO 2 (g)  C 2 H 6 (g) + O 2 (g) ----> H 2 0(g) + CO 2 (g)  C 3 H 8 (g) + O 2 (g) ----> H 2 O(g) + CO 2 (g)  H 2 (g) + O 2 (g) ----> H 2 O(g)  Mg(g) + O 2 (g) ----> MgO(s)

23. Another type of reaction: oxidation- reduction or redox  In these, the charges on some of the atoms involve change because of an electron transfer.  Oxidation Is Loss of electrons.  Reduction Is Gain of electrons.  Remember “OIL RIG”

24.  When oxidation occurs, the charge on the atom or ion is increased.  When reduction occurs, the charge on the atom or ion is decreased.  One thing to remember: The charge on an element by itself is 0.

25. Examples:  Fe +2  Fe +3 + e -  Cl 2 + 2e -  2Cl -  Fe +3 + e -  Fe +2  Cu +2 + 2e -  Cu  Zn  Zn +2 + 2e -