1. Unit 6 Chemical Equations

2. Objective
III.A.3(a) - Explain how conservation laws form the basis for balancing chemical reactions and know what quantities are conserved in physical, chemical, and nuclear changes

3. Law of Conservation of Mass
In a chemical reaction, atoms are neither created or destroyed
All atoms in the reactants MUST be accounted for in the products

4. Objective
III.A.3(d) - Use the appropriate symbols for state (i.e., solid, liquid, gaseous, aqueous) and reaction direction when writing chemical equations

5. Chemical Equations Reactants – what goes INTO the rxn
Products – what comes OUT OF the rxn
Reactant A + Reactant B  Product C

6. Balancing Chemical Equations
Done by trial and error 
Can only change coefficients (Coefficients are numbers BEFORE the chemical formula)
Balance elements that only appear once
Keep polyatomic ions together if possible
***NEVER EVER EVER EVER EVER CHANGE A SUBSCRIPT!!!***
Balance hydrogen and oxygen last

7. Diatomic Elements
7 elements exist in nature as diatomic molecules (2 atoms)
Hydrogen (H2)
Oxygen (O2)
Nitrogen (N2)
Fluorine (F2)
Chlorine (Cl2)
Bromine (Br2)
Iodine (I2)

8. Example #1
C2H5OH (l) + O2(g)  CO2(g) + H2O(g)

9. Fe2O3(s) + HNO3(aq)  Fe(NO3)3(aq) + H2O(l)
Example #2
Fe2O3(s) + HNO3(aq)  Fe(NO3)3(aq) + H2O(l)

10. H2S(g) + Pb(NO3)2(aq)  PbS(s) + HNO3(aq)
Example #3
H2S(g) Pb(NO3)2(aq)  PbS(s) HNO3(aq)

12. Objective
III.A.3(c) – Describe what is represented, on a molecular and molar level, by chemical equations

13. What Are these Coefficients of which you speak?
Coefficients represent either molecules, moles, or (when we get to it) volumes in liters
1O2(g) + 2H2(g)  2H2O(g)
1 molecule of O2 reacts with 2 molecules of H2 to produce 2 molecules of H2O OR
1 mole of O2 reacts with 2 moles of H2 to produce 2 moles of H2O

14. Types of Reactions
III.A.3(e) – Classify chemical reactions as being synthesis, decomposition, single replacement, or double replacement

16. Objectives Give general equations for types of reactions
Classify reactions
List 3 types of synthesis and 6 decomposition reactions
List 4 types of single-replacement and 3 types of double-replacement reactions
Predict products of reactions given the reactants

17. Synthesis Reactions
General Formula:
A + X  AX

18. With metals form metal oxides
Synthesis with Oxygen
With metals form metal oxides
Ex – 2Mg(s) + O2(g)  2MgO(s)
4K(s) + O2(g)  2K2O(s)
2Fe(s) + O2(g)  2FeO(s)
4Fe(s) + 3O2(g)  2Fe2O3
With non-metals form non-metal oxides
Ex – S8(s) + 8O2(g)  8SO2(g)
C(s) + O2(g)  CO2(g)

19. With metals produce metal sulfides
Synthesis with Sulfur
With metals produce metal sulfides
Ex- 16Rb(s) + S8(s)  8Rb2S(s)
8Ba(s) + S8(s)  8BaS(s)

20. Synthesis of Metals with Halogens
Form metal halides Group 1: 2M + X2  2MX Ex – 2Na(s) + Cl2(g)  2NaCl(s) Group 2: M + X2  MX2 Ex – Mg(s) + F2(g)  MgF2(s)

21. Metal Oxides with Water
Group 1 & 2 form hydroxides
Ex – K2O(s) + H2O(l)  2KOH(aq)
CaO(s) + H2O(l)  Ca(OH)2(l)

22. Metal Chlorides with Oxygen
Form Metal Chlorates Ex – 2NaCl(s) + 3O2(g)  2NaClO3(s)

23. Non-Metal Oxides with Water
Form oxyacids
Ex – SO2(g) + H2O(l)  H2SO3(aq)
P2O5(s) + 3H2O(l)  2H3PO4(aq)

24. Metal oxides with carbon dioxide
Form metal carbonates EX – CaO(s) + CO2(g)  CaCO3(s)

26. Decomposition Reactions
General Formula
AX  A + X
Only 1 reactant and multiple products

27. Decomposition of Binary Compounds
Breaks down into its component elements
2H2O(l)  2H2(g) + O2(g)

28. Decomposition of Metal Carbonates
Form metal oxides and carbon dioxide
CaCO3(s)  CaO(s) + CO2(g)

29. Decomposition of Metal Hydroxides
Form metal oxides and water
Ca(OH)2(s)  CaO(s) + H2O(g)

30. Decomposition of Metal Chlorates
Form metal oxides and water
Ca(OH)2(s)  CaO(s) + H2O(g)

31. Decomposition of Acids
Form non-metal oxides and water
H2SO4(aq)  SO3(g) + H2O(l)

32. Single Replacement Reactions
General Form
AX + B  A + BX
Element + compound  Different element + different compound

33. Replacement of a Metal in a Compound by Another Metal
Aluminum is more reactive than lead
2Al(s) + 3Pb(NO3)2(aq)  3Pb(s) + 2Al(NO3)3(aq)

34. Replacement of Hydrogen in Water by a Metal
More active metals react with liquid water
2Na(s) + H2O(l)  2NaOH(aq) + H2(g)
Less active metals react with steam
3Fe(s) + 4H2O(g)  Fe3O4(s) +4H2(g)

35. Replacement of Hydrogen in an Acid by a Metal
Metals more active than hydrogen
Mg(s) + 2HCl(aq)  H2(g) + MgCl2(aq)

36. Replacement of Halogens
Each halogen (Group 17) can replace the halogen below it on the periodic table
Cl2(g) + 2KBr(aq)  2KCl(aq) + Br2(l)
F2(g) + 2KCl(aq)  2KF(aq) + Cl2(s)
Br2(l) + KCl  no reaction

37. Double-Replacement Reactions
General Form
AX + BY  AY + BX
Compound A + Compound B  Compound C + Compound D

38. Formation of a Precipitate
Cations switch places
Solid insoluble product forms
2KI(aq) + Pb(NO3)2(aq)  PbI2(s) + 2KNO3(aq)

39. FeS(s) + 2HCl(aq)  H2S(g) + FeCl2(aq)
Formation of a Gas
Insoluble gas forms
Bubbles out of solution
FeS(s) + 2HCl(aq)  H2S(g) + FeCl2(aq)

40. HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
Formation of Water
Neutralization reaction
Acid/Base Reaction
Produce a salt and water
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)

41. Combustion Reactions
React with oxygen to release a large amount of energy in the form of light and heat
C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g)

42. Predicting the Products of Reactions
III.A.3(f) - Predict the products of synthesis, combustion, and decomposition reactions and write balanced equations for these reactions

43. Activity Series of the Elements
Greater activity of a metal indicates how easily it loses electrons
Greater activity of a nonmetal indicates how easily it gains electrons
In a single-replacement reaction, if an element with lower activity is to be replaced, the reaction will take place.

44. The End