1. Chapter 8
Chemical Equations

2. Parts of a Chemical Equation
O2(g) + Fe(s)  Fe2O3(s)
Products – chemicals produced in a chemical reaction. (products)
Reactants – chemicals reacted together in a chemical reaction.
Products = Fe2O3
Reactants = O2 and Fe
 - yields or produces

3. Parts of a Chemical Equation
O2(g) + Fe(s)  Fe2O3(s)
(g) = gas
(s) = solid
(l) = liquid
examples: water, Hg, Br2, organics
(aq) = aqueous
A homogeneous solution. Most of what we use in class

4. Parts of a Chemical Equation
O2(g) + Fe(s)  Fe2O3(s)
Why O2 not O and why Fe not Fe2?
Oxygen is diatomic.
Meaning that in order to be at its lowest energy, oxygen will bond with other oxygen.
Diatomic Elements – H, N, O, F, Cl, Br, I and probably (At)

5. Let’s Sing There are seven Diatomic Elements That you must know.
(repeat)
Fluorine, Chlorine
Bromine, Iodine
Nitrogen, Oxygen
Hydrogen

6. Other Symbols Things on the reaction arrow
Catalyst – speeds up a reaction.
Inhibitor – slows down a reaction.
Neither are parts of the products or reactants.
Energy – in the form of heat, light or electricity. Probably heat!
NH4Cl Δ

7. Example 1
When solid mercury (II) sulfide is heated with oxygen, liquid mercury metal and gaseous sulfur dioxide are produced.

8. Example 2
Oxygen gas can be made by heating potassium chlorate in the presence of the catalyst manganese dioxide. Potassium chloride is left as a solid residue.

9. Balancing Equations
Balanced Equation – an equation that has the same number and type of atoms on the product and reactant sides.
We balance equations by putting in coefficients.
A coefficient is a multiplier.
You MAY NOT change subscripts, unless they are yours and you made a mistake initially when you wrote the equation.

10. Why do we balance equations?
Law of conservation of matter – matter is neither created nor destroyed.

11. Example 3 Δ
KClO3  O2 + KCl K Cl O

12. Example 4
Al + O2  Al2O3 Al O

13. Example 5
AgNO3 + Cu  Cu(NO3)2 + Ag Ag N O Cu

14. Example 5 (again)
AgNO3 + Cu  Cu(NO3)2 + Ag Ag
NO3 Cu

15. Example 6
FeCl2 + Bi2S3  FeS + BiCl3 Fe Cl Bi S

16. General Hints In general you will want of get rid of odd numbers.
Also, work with combined elements before working on lone elements.

17. Types of Reactions
Chapter 8

18. Five Common Types of Reactions
Synthesis
(Direct combination or combination)
Decomposition
Single Replacement
(Single Displacement)
Double Replacement
(Double Displacement)
Combustion

19. Synthesis
Two or more elements or compounds combine to form a new compound.
General Form:
A + B  AB

20. Specific Examples
Cl2 + Al 

21. Specific Examples
Sr + N2 

22. Decomposition
A compound breaks down into two or more elements and/or compounds.
General Form:
AB  A + B

23. Specific Examples
MgO 

24. Specific Examples
CoBr3 

25. Specific Examples
Carbonates break down into their metal oxide and carbon dioxide.
Na2CO3 

26. Specific Examples
Hydroxides break down into their metal oxide and water.
NaOH 

27. Single Replacement
An element replaces another element inside a compound.
General Form:
A + BC  B + AC or
A + BC  C + BA

28. Activity Series of Metals
An activity series of metals needs to be consulted to determine if a reaction will take place or not.
A more reactive metal will replace a less reactive metal in a compound.
Less commonly, there is also an activity series of non-metals as well.

29. Activity Series of Metals
Lithium Most Reactive
Potassium
Barium
Calcium
Sodium
Magnesium
Aluminum
Zinc
Iron
Nickel
Tin
Lead
Hydrogen
Copper
Mercury
Silver
Gold Least Reactive

30. Activity Series of Halogens
Fluorine Most Reactive
Chlorine
Bromine
Iodine Least Reactive

31. Specific Examples
Au + NaCl 

32. Specific Examples
Cu + AgNO3 

33. Specific Examples
Li + H2O 

34. Specific Examples
Al + Mg(NO3)2 

35. Double Replacement
One element out of a compound replaces another element out of a compound.
General Form
AB + CD  CB + AD

36. Evidence of a Double Replacement Reaction
A precipitate forms.
Insoluble solid
A gas is produced.
Water is produced.

37. Specific Examples
KCl + Ca(NO3)2 

38. Specific Examples
HCl + NaOH 

39. Specific Examples
Na2CO3 + Cd(NO3)2 

40. Specific Examples
HBr + K2CO3

41. Combustion
A hydrocarbon reacts with oxygen to produce carbon dioxide and water. Heat and light are also produced.
General Form:
CxHyOz + O2  CO2 + H2O

42. Specific Examples
Combusting propane (C5H12)
C5H O2  CO2 + H2O

43. Specific Examples Combusting sugar (C12H22O11)
C12H22O O2  CO2 + H2O